New Journal of Chemistry January 2009a

www.rsc.org/njc Volume 33 | Number 1 | January 2009 | Pages 1–212

ISSN 1144-0546

Volume 33 | N

umber 1 | 2009 N

JC Pages 1–212

1144-0546(2009)33:1;1-2

New Journal of Chemistry An international journal of the chemical sciences

PAPERT. Yong-Jin Han et al.The solubility and recrystallization of 1,3,5-triamino-2,4,6-trinitrobenzene in a 3-ethyl-1-methylimidazolium acetate–DMSO co-solvent system

www.rsc.org/metallomicsRegistered Charity Number 207890

A new journal from RSC Publishinglaunching in 2009

MetallomicsIntegrated biometal science

This timely new journal will cover the research � elds related to metals in biological, environmental and clinical systems and is expected to be the core publication for the emerging metallomics community. The journal will be supported by an international Editorial Board, chaired by Professor Joseph A. Caruso of the University of Cincinnati/Agilent Technologies Metallomics Center of the Americas.

Metallomics will publish six issues in the � rst year, increasing to 12 issues in 2010. The journal will contain a full mix of research articles including Communications, Reviews, Full papers, and Editorials. From launch, the latest issue will be freely available online to all readers. Free institutional access to previous issues throughout 2009 and 2010 will be available following a simple registration process.

Contact the editor, Niamh O’Connor, at [email protected] for further information or visit the website.

Num

ber 1|

2008M

etallomics

Pages1–100 1754-5692(2008)1:1;1-6

www.rsc.org/metallomics Volume 1 | Number 1 | January 2009 | Pages 1–100

ISSN 1756-5901


1756-5901(2009) 1:1;l-m

060877

Submit your work now!Supporting the

ISSN 1144-0546

PAPERRudi van Eldik et al.Metal ion-catalyzed oxidative degradation of Orange II by H2O2. High catalytic activity of simple manganese salts



www.rsc.org/ibiologyRegistered Charity Number 207890


Integrative BiologyQuantitative biosciences from nano to macro

Integrative Biology provides a unique venue for elucidating biological processes, mechanisms and phenomena through quantitative enabling technologies at the convergence of biology with physics, chemistry, engineering, imaging and informatics.

With 12 issues published annually, Integrative Biology will contain a mix of research articles including Full papers, Reviews (Tutorial & Critical), and Perspectives. It will be supported by an international Editorial Board, chaired by Distinguished Scientist Dr Mina J Bissell of Lawrence Berkeley National Laboratory.

The current issue of Integrative Biology will be freely available to all readers via the website. Free institutional online access to all 2009 and 2010 content of the journal will be available following registration at www.rsc.org/ibiology_registration

Volume

18|N

umber1

|2008Journal of M

aterials Chem

istryPages

1–140

www.rsc.org/ibiology Volume 1 | Number 1 | January 2009 | Pages 1–140 1–1401–140

ISSN 1757-9694

0959-9428(2008)18:1;1-J

Integrative Biology Quantitative biosciences from nano to macro

1757-9694(2009) 1:1;1

0608

58Contact the Editor, Harp Minhas, at [email protected] or visit the website for more details.

CHEMICAL SCIENCE

C1

Drawing together research highlights and news from all RSCpublications, Chemical Science provides a ‘snapshot’ of thelatest developments across the chemical sciences, showcasingnewsworthy articles and significant scientific advances.

EDITORIAL

17

Changes ahead for NJC in 2009

Denise Parent and Sarah Ruthven highlight the changesto NJC for the year ahead, together with the latest newsfrom the RSC.

N J CNew Journal of Chemistry. An international journal for the chemical sciences

www.rsc.org/njc

RSC Publishing is a not-for-profit publisher and a division of the Royal Society of Chemistry. Any surplus made is used to support charitableactivities aimed at advancing the chemical sciences. Full details are available from www.rsc.org

CoverSee T. Yong-Jin Han et al.,pp. 50–56.A very strong inter- and intra-molecular hydrogen bondingsolid, 1,3,5-triamino-2,4,6-trinitrobenzene, can be dissolvedand recrystallized in a 3-ethyl-1-methylimidazolium acetate–DMSOco-solvent system. Imagereproduced with the permission ofLawrence Livermore NationalLaboratory and T. Yong-Jin Han,Philip F. Pagoria, Alexander E. Gash,Amitesh Maiti, Christine A. Orme,Alexander R. Mitchell and LaurenceE. Fried from New J. Chem., 2009,33, 50.

Inside CoverSee Rudi van Eldik et al.,pp. 34–49.Organic dyes from industrialwaste water effluents can causelarge scale pollution of naturalrivers. Simple metal ions catalyzethe oxidative degradation of suchdyes and rapidly clean thepolluted water. Imagereproduced with permission ofErika Ember, Sabine Rothbart,Ralph Puchta and Rudi van Eldikfrom New J. Chem., 2009, 33, 34.

IN THIS ISSUE

ISSN 1144–0546 CODEN NJCHES 33(1) 1–212 (2009)

This journal is �c the Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2009 New J. Chem., 2009, 33, 3–16 | 3

Associate editorsManuscripts should be directed to one of the Editors detailed below.

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Editor (RSC)Sarah Ruthven

Editor (CNRS)Denise Parent

Assistant editorsMarie Cote (CNRS)Sarah Dixon (RSC)

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PublisherEmma Wilson (RSC)

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New Journal of Chemistry (Print: ISSN 1144-0546;electronic: ISSN 1369-9261) is published 12 times a year by the Centre National de la Recherche Scientifique (CNRS), 3 rue Michel-Ange, 75794 Paris cedex 16, France, and the Royal Society of Chemistry (RSC), Thomas Graham House, Science Park, Milton Road, Cambridge, UK CB4 0WF.

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www.rsc.org/njcThe New Journal of Chemistry is a broad-based primary journal encompassing all branches of the chemical sciences. Published monthly, it contains full research articles, letters, opinions and perspectives.

INFORMATION FOR AUTHORS

Co-editor-in-chiefPascal Le Floch, Palaiseau, France

Co-editor-in-chiefJerry Atwood, Columbia, MO, USA

Consulting editorOdile Eisenstein, Montpellier, France

Board membersYasuhiro Aoyama, Kyoto, JapanKumar Biradha, Khargapur, IndiaLaurent Bonneviot, Lyon, FranceFabrizia Grepioni, Bologna, ItalyHelen Hailes, London, UKPeter Junk, Monash, AustraliaBarbara Nawrot, Lodz, Poland

Alan Rowan, Nijmegen, The NetherlandsMichael Scott, Gainesville, FL, USAMichael Veith, Saarbrücken, GermanyVivian Yam, Hong Kong, PR China

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EDITORIAL BOARD

Markus Antonietti, MPI, Potsdam, GermanyMatthias Bremer, Darmstadt, GermanyRobert Crabtree, New Haven, CT, USAFrançois Fajula, Montpellier, FranceJohn A. Gladysz, College Station, TX, USAGeorge Gokel, St Louis, MO, USA

Andrew B. Holmes, Melbourne, AustraliaMiguel Julve, Valencia, SpainHenryk Koslowski, Wroclaw, PolandJean-Pierre Majoral, Toulouse, FranceLuca Prodi, Bologna, ItalyJan Reedijk, Leiden, The Netherlands

David Reinhoudt, Enschede, The NetherlandsKari Rissanen, Jyväskylä, FinlandClément Sanchez, Paris, FranceJeremy K. M. Sanders, Cambridge, UKJean-Pierre Sauvage, Strasbourg, FranceJonathan W. Steed, Durham, UK

ADVISORY BOARD

LETTERS

19

Evidence of crystalline/glassy intermediates in bismuthphosphates

Marie Colmont,* Laurent Delevoye and Olivier Mentre

The wide NMR chemical shift range of 17O provides aprofitable source of information about partially orderedmaterials. In addition, original phosphorous/oxygenthrough-bond correlation experiments have allowed theunambiguous assignment of the 17O resonances.

23

A supramolecular sensing system for AgI at nanomolarlevels by the formation of a luminescentAgI– TbIII–thiacalix[4]arene ternary complex

Nobuhiko Iki,* Munehiro Ohta, Teppei Tanaka,Takayuki Horiuchi and Hitoshi Hoshino

The first example of the detection of AgI ions usingsupramolecular chemistry is demonstrated, in which twothiacalix[4]arene ligands are linked by analyte AgI ions andthen coordinate to TbIII ions to form a luminescent ternarycomplex, AgI2 �TbIII2 �TCAS2, enabling the detection of AgI atconcentrations as low as 3.2 � 10� 9 M.

PAPERS

26

Ionic liquids with dual biological function: sweetand anti-microbial, hydrophobic quaternaryammonium-based salts

Whitney L. Hough-Troutman, Marcin Smiglak,Scott Griffin, W. Matthew Reichert, Ilona Mirska,Jadwiga Jodynis-Liebert, Teresa Adamska, Jan Nawrot,Monika Stasiewicz, Robin D. Rogers* and Juliusz Pernak*

Newly synthesized dual function ionic liquids combine bothanti-microbial and sweetener properties into one compound.

34

Metal ion-catalyzed oxidative degradation of Orange IIby H2O2. High catalytic activity of simple manganesesalts

Erika Ember, Sabine Rothbart, Ralph Puchta andRudi van Eldik*

In an effort to develop new routes for the clean oxidationof non-biodegradable organic dyes, a detailed study of someenvironmentally friendly Mn(II) salts that form very efficientin situ catalysts for the activation of H2O2 in the oxidation ofsubstrates such as Orange II under mild reaction conditions,was performed.


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98

PAPERS

50

The solubility and recrystallization of 1,3,5-triamino-2,4,6-trinitrobenzene in a 3-ethyl-1-methylimidazoliumacetate–DMSO co-solvent system

T. Yong-Jin Han,* Philip F. Pagoria, Alexander E. Gash,Amitesh Maiti, Christine A. Orme, Alexander R. Mitchelland Laurence E. Fried

A highly hydrogen-bonded solid, 1,3,5-triamino-2,4,6-trinitrobenzene (TATB), was dissolved and recrystallizedin various IL systems. Dissolution of TATB in EMImOAcoccurred by forming a very stable s-complex.

57

Supramolecular synthesis of some molecular adductsof 4,40-bipyridine N,N0-dioxide

Kapildev K. Arora, Mayura S. Talwelkar andV. R. Pedireddi*

4,4’-Bipyridine N,N0-dioxide has yielded different types ofsupramolecular assemblies from simple stacked sheet structuresto pseudorotaxane and stair-case type structures dependingupon its interaction with the co-crystallizing agents used in thesupramolecular synthesis.

64

Synthesis and characterisation of bulky guanidines andphosphaguanidines: precursors for low oxidation statemetallacycles

Guoxia Jin, Cameron Jones,* Peter C. Junk,Kai-Alexander Lippert, Richard P. Rose andAndreas Stasch

Reactions of alkali metal amides or phosphides with the bulkycarbodiimide, ArNQCQNAr (Ar = C6H3Pr

i2-2,6), followed

by aqueous work-ups, have yielded several guanidines, abifunctional guanidine and two phosphaguanidines (e.g. seepicture).

76

The hydrogen bond acidity and other descriptorsfor oximes

Michael H. Abraham,* Javier Gil-Lostes,J. Enrique Cometto-Muniz, William S. Cain,Colin F. Poole, Sanka N. Atapattu, Raymond J. Abrahamand Paul Leonard

The hydrogen bond acidity of cyclohexanone oxime andacetone oxime are 0.33 and 0.37, respectively; this placesoximes as about the same hydrogen bond acidity as alcohols.


www.rsc.org/chemcommRegistered Charity Number 207890

Introducing Professor Mike Doyle

Make an impact

Associate Editor for Organic ChemistryMichael P. (Mike) Doyle is Professor and Chair of the Department of Chemistry and Biochemistry at the University of Maryland, College Park. He has been the recipient of numerous awards, including the George C. Pimentel Award for Chemical Education in 2002 and the Arthur C. Cope Scholar Award in 2006. He has written or coauthored ten books, including Basic Organic Stereochemistry, 20 book chapters, and he is the co-author of more than 270 journal publications. The inventor of chiral dirhodium carboxamidate catalysts known as “Doyle catalysts,” his research is focused on applications with metal carbene transformations, Lewis acid catalyzed reactions, and selective catalytic oxidations.

Submit your work to ChemCommProfessor Doyle will be delighted to receive submissions from North America in the field of organic chemistry. Submissions to ChemComm are welcomed via ReSourCe, our homepage for authors and referees.

“ChemComm is an outstanding forum for the communication of significant research in the chemical sciences, and I am honoured to be a member of the editorial family. I continue to be amazed with the breadth of exciting chemistry that is being submitted to ChemComm and the high level of professionalism that is found at ChemComm.”

ISSN 1359-7345

1359-7345(2007)39;1-Y

www.rsc.org/chemcomm Number 39 | 21 October 2007 | Pages 3969 – 4060

Chemical Communications

COMMUNICATIONTakahiko Kojima et al.A discrete conglomerate of a distorted Mo(V)-porphyrin with a directly coordinated Keggin-type polyoxometalate

FEATURE ARTICLESYoshinori Yamanoi and Hiroshi NishiharaAssembly of nanosize metallic particles and molecular wires on electrode surfacesDonnaG.Blackmond and Martin KlussmannAssessing phase behavior models for the evolution of homochirality

0208

78

PAPERS

82

Electrochemical methodology for determinationof imidazolium ionic liquids (solids at room temperature)properties: influence of the temperature

M. P. Stracke, M. V. Migliorini, E. Lissner,H. S. Schrekker, D. Back, E. S. Lang, J. Dupont* andR. S. Goncalves*

Electrochemical impedance spectroscopy for determinationof imidazolium ionic liquid physicochemical properties: theinfluence of the temperature on the Nyquist diagrams.

88

A new family of biocompatible and stable magneticnanoparticles: silica cross-linked pluronic F127 micellesloaded with iron oxides

Zhaoyang Liu,* Jun Ding and Junmin Xue*

A new family of magnetic nanoparticles, silica cross-linkedpluronic F127 micelles loaded with iron oxides having theproperties of high biocompatibility, physical and chemicalstability, high magnetism, and low-cost production, have beensynthesized.

93

Novel thiophene-conjugated indoline dyes for zinc oxidesolar cells

Takuya Dentani, Yasuhiro Kubota, Kazumasa Funabiki,Jiye Jin, Tsukasa Yoshida, Hideki Minoura,Hidetoshi Miura and Masaki Matsui*

The introduction of thiophene ring(s) into D131-type indolinedyes improved cell performance due to their appropriateenergy levels and bathochromic shift in the UV-vis absorptionband on zinc oxide.

102

Gold imidazolium-based ionic liquids, efficient catalystsfor cycloisomerization of c-acetylenic carboxylic acids

Florentina Nea]u, Vasile I. Parvulescu,Veronique Michelet, Jean-Pierre Genet,Alexandre Goguet and Christopher Hardacre

Ionic liquid stabilized gold(III) chloride is shown to be a veryactive catalyst in the cyclization of sterically hindered andunhindered acetylenic carboxylic acid substrates even in theabsence of a base.


www.rsc.org/booksRegistered Charity Number 207890

Chemistry at Oxford:A History from 1600 to 2005A fascinating and unique history of the Oxford Chemistry School!

● Discover how individuals have shaped the school and made great achievements in teaching and research

● Read about the seminal works of Robert Boyle, Hinshelwood, Robinson and Hodgkin

● Discover how separate branches of chemistry (organic, physical, inorganic and biological) have evolved in Oxford

● Get to grips with the unusual character of Oxford University and learn more about its unique history.

This fantastic new book will appeal to all those interested in the history and present day style of science, especially chemistry, education and research at Oxford as contrasted with that to be found elsewhere.

Hardback | 300 pages | ISBN 9780854041398 | 2008 | £54.95

10070

PAPERS

107

Magnetically moveable bimetallic (nickel/silver)nanoparticle/carbon nanotube composites for methanoloxidation

Guan-Ping Jin,* Ronan Baron, Neil V. Rees, Lei Xiao andRichard G. Compton*

The functionalization of carbon nanotubes with both AgNPsand a minute fraction of NiNPs add to the electrocatalyticproperties of the AgNPs, the possibility to easily move them insolution using a magnet. The bi-functionalized carbonnanotubes are then easily recoverable after use.

112

Microwave-assisted facile synthesis of discotic liquidcrystalline symmetrical donor–acceptor–donor triads

Satyam Kumar Gupta, V. A. Raghunathan andSandeep Kumar*

The first examples of columnar phase formingtriphenylene-anthraquinone-based donor–acceptor–donortriads were prepared and characterized by polarizing opticalmicroscopy, differential scanning calorimetry andX-ray diffractometry.

119

Synthesis, crystal structures and luminescence propertiesof lanthanide oxalatophosphonates with athree-dimensional framework structure

Yanyu Zhu, Zhengang Sun,* Yan Zhao, Jing Zhang,Xin Lu, Na Zhang, Lei Liu and Fei Tong

Six new three-dimensional (3D) lanthanideoxalatophosphonates, [Ln(HL)(C2O4)0.5(H2O)2] �H2O(Ln = La (1), Ce (2), Pr (3), Nd (4), Sm (5), Eu (6);H3L = H2O3PCH(OH)CO2H), have been synthesized andstructurally characterized. Compound 6 shows strong redluminescence in the solid state at room temperature.

125

The annular tautomerism of the curcuminoidNH-pyrazoles

Pilar Cornago,* Pilar Cabildo, Rosa M. Claramunt,Latifa Bouissane, Elena Pinilla, M. Rosario Torres andJose Elguero

The structures of six NH-pyrazoles, derived from curcuminand related b-diketones, have been established byX-ray crystallography, and solid state 13C and 15N CPMASNMR spectroscopy.


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‘NJC book of choice’Why not take advantage of free book chapters from the RSC? Through our ‘NJC book of choice’ scheme NJC will regularly highlight a book from the RSC eBook Collection relevant to your research interests. Read the latest chapter today by visiting the NJC website. The RSC eBook Collection o� ers:

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Volum

e31

|Num

ber10|2007

NJC

Pages1693–1832

PERSPECTIVEZhaohua Dai and James W. CanaryTailoring tripodal ligands forzinc sensing

ISSN 1144-0546

New Journal of Chemistry

www.rsc.org/njc

An international journal of the chemical sciences

Volume 31 Number 10 | October 2007 | Pages 1693–1832

1144-0546(2007)31:10;1-8

www.rsc.org/analystRegistered Charity Number 207890

Not just an analyticalchemistry journal

detection

bio-analyticalanalytical

bio-analyticalbio-analytical

The Analyst

High pro� le and cited in MEDLINE, Impact factor: 3.198†

Immediacy index 0.925 - highest ranked general analytical

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0003-2654(2007)132:5;1-J

www.rsc.org/analyst Volume 132 | Number 5 | May 2007 | Pages 377–484

Interdisciplinary detection science

ISSN 0003-2654

CRITICAL REVIEWStepan Shipovskov and Curt T. ReimannElectrospray ionization mass spectrometry in enzymology

PAPERJoseph Wang et al.Discrete micro� uidics with electrochemical detection

†2006 Thompson Scienti� c (ISI) Journal Citation Reports*Based on the ISI JCR 2006 Science Edition for the Analytical Chemistry subject category

PAPERS

136

Neutral 5-nitrotetrazoles: easy initiation with low pollution

Thomas M. Klapotke,* Carles Miro Sabate andJorg Stierstorfer

New synthesis, crystal structures and characterization ofneutral 5-nitro-2H-tetrazole, 1-methyl-5-nitrotetrazole and2-methyl-5-nitrotetrazole are presented. These nitrogen-richcompounds were tested to be highly energetic with increasedsensitivities towards impact, friction and electrical discharge.

148

Probing multivalency for the inhibition of an enzyme:glycogen phosphorylase as a case study

Samy Cecioni, Oana-Andreea Argintaru, Tibor Docsa,Pal Gergely, Jean-Pierre Praly and Sebastien Vidal*

The concept of multivalency was applied to the inhibitionof an enzyme (glycogen phosporylase). Trivalent inhibitorswere synthesized and displayed improved activities incomparison to their monovalent counterparts.

157

The formation of silver nanofibres by liquid/liquidinterfacial reactions: mechanistic aspects

Kun Luo and Robert A. W. Dryfe*

Silver nano-fibres deposited by spontaneous reduction at thewater/organic interface.

164

The role of nucleophilic catalysis in chemistry andstereochemistry of ribonucleoside H-phosphonatecondensation

Michal Sobkowski,* Jacek Stawinski and Adam Kraszewski

Reactions of ribonucleoside 30-H-phosphonates with alcoholsproceed with high stereoselectivity towards the samediastereomer irrespective of the presence or absence ofnucleophilic catalysts.


PAPERS

171

Two polyaminophenolic fluorescent chemosensors forH+ and Zn(II). Spectroscopic behaviour of free ligandsand of their dinuclear Zn(II) complexes

Gianluca Ambrosi, Cristina Battelli, Mauro Formica,Vieri Fusi,* Luca Giorgi, Eleonora Macedi,Mauro Micheloni,* Roberto Pontellini and Luca Prodi

UV-Vis and fluorescence properties of two polyamino-phenolicligands; design of new efficient fluorescent chemosensors forH+ and Zn(II) ions.

181

Dynamic covalent self-assembled macrocycles preparedfrom 2-formyl-aryl-boronic acids and 1,2-amino alcohols

Ewan Galbraith, Andrew M. Kelly, John S. Fossey,Gabriele Kociok-Kohn, Matthew G. Davidson,Steven D. Bull* and Tony D. James*

Reaction of 2-formyl-aryl-boronic acids with1,2-amino alcohols results in dynamic covalent self assemblyto quantitatively afford tetracyclic macrocyclic Schiff baseboracycles containing bridging boron–oxygen–boronfunctionality.

186

N-Inversion in 2-azabicyclopentane derivatives: modelsimulations for a laser controlled molecular switch

Bastian Klaumunzer* and Dominik Kroner

Quantum model simulation of a N-inversion based lasercontrolled molecular switch by IR-ladder-climbing or UV.

196

How does non-covalent Se?SeQO interaction stabilizeselenoxides at naphthalene 1,8-positions: structural andtheoretical investigations

Satoko Hayashi, Waro Nakanishi,* Atsushi Furuta,Jozef Drabowicz, Takahiro Sasamori and Norihiro Tokitoh

Non-covalent G?SeQO 3c–4e interactions are demonstratedto determine the fine structures of 8-G-1-[MeSe(O)]C10H6 andoperate to protect from racemization of the selenoxides: G ofSeMe acts more effectively than G of halogens.

14 | New J. Chem., 2009, 33, 3–16 This journal is �c the Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2009

PAPERS

207

Mechanistic aspects of nitrate ion reduction on silverelectrode: estimation of O–NO2

� bond dissociationenergy using cyclic voltammetry

Mohsin Ahmad Bhat, Pravin Popinand Ingole,Vijay Raman Chaudhari and Santosh Krishna Haram*

Cyclic voltammetric investigations for nitrate ion reductionat silver electrode in alkaline medium, show that reactionfollows a concerted dissociative electron transfer mechanism,with bond dissociation energy of O–NO2

� bond ofca. 48.4 kcal mol� 1.

Call for abstractsThis is your chance to take part in IUPAC 2009. Contributions

are invited for oral presentation by 16 January 2009 and poster

abstracts are welcome until 5 June 2009.

ThemesAnalysis & Detection

Chemistry for Health

Communication & Education

Energy & Environment

Industry & Innovation

Materials

Synthesis & Mechanisms

Plenary speakersPeter G Bruce, University of St Andrews

Chris Dobson, University of Cambridge

Ben L Feringa, University of Groningen

Sir Harold Kroto, Florida State University

Klaus Müllen, Max-Planck Institute for Polymer Research

Sir J Fraser Stoddart, Northwestern University

Vivian W W Yam, The University of Hong Kong

Richard N Zare, Stanford University

For a detailed list of symposia, keynote speakers and to submitan abstract visit our website.

42nd IUPAC CONGRESS Chemistry Solutions2–7 August 2009 | SECC | Glasgow | Scotland | UK

Sponsored by

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AUTHOR INDEX

Abraham, Michael H., 76Abraham, Raymond J., 76Adamska, Teresa, 26Ambrosi, Gianluca, 171Argintaru, Oana-Andreea, 148Arora, Kapildev K., 57Atapattu, Sanka N., 76Back, D., 82Baron, Ronan, 107Battelli, Cristina, 171Bhat, Mohsin Ahmad, 207Bouissane, Latifa, 125Bull, Steven D., 181Cabildo, Pilar, 125Cain, William S., 76Cecioni, Samy, 148Chaudhari, Vijay Raman, 207Claramunt, Rosa M., 125Colmont, Marie, 19Cometto-Muniz, J. Enrique, 76Compton, Richard G., 107Cornago, Pilar, 125Davidson, Matthew G., 181Delevoye, Laurent, 19Dentani, Takuya, 93Ding, Jun, 88Docsa, Tibor, 148Drabowicz, Jozef, 196Dryfe, Robert A. W., 157Dupont, J., 82Elguero, Jose, 125Ember, Erika, 34Formica, Mauro, 171Fossey, John S., 181Fried, Laurence E., 50

Funabiki, Kazumasa, 93Furuta, Atsushi, 196Fusi, Vieri, 171Galbraith, Ewan, 181Gash, Alexander E., 50Genet, Jean-Pierre, 102Gergely, Pal, 148Gil-Lostes, Javier, 76Giorgi, Luca, 171Goguet, Alexandre, 102Goncalves, R. S., 82Griffin, Scott, 26Gupta, Satyam Kumar, 112Han, T. Yong-Jin, 50Haram, Santosh Krishna, 207Hardacre, Christopher, 102Hayashi, Satoko, 196Horiuchi, Takayuki, 23Hoshino, Hitoshi, 23Hough-Troutman, Whitney

L., 26Iki, Nobuhiko, 23Ingole, Pravin Popinand, 207James, Tony D., 181Jin, Guan-Ping, 107Jin, Guoxia, 64Jin, Jiye, 93Jodynis-Liebert, Jadwiga, 26Jones, Cameron, 64Junk, Peter C., 64Kelly, Andrew M., 181Klapotke, Thomas M., 136Klaumunzer, Bastian, 186Kociok-Kohn, Gabriele, 181Kraszewski, Adam, 164

Kroner, Dominik, 186Kubota, Yasuhiro, 93Kumar, Sandeep, 112Lang, E. S., 82Leonard, Paul, 76Lippert, Kai-Alexander, 64Lissner, E., 82Liu, Lei, 119Liu, Zhaoyang, 88Lu, Xin, 119Luo, Kun, 157Macedi, Eleonora, 171Maiti, Amitesh, 50Matsui, Masaki, 93Mentre, Olivier, 19Michelet, Veronique, 102Micheloni, Mauro, 171Migliorini, M. V., 82Minoura, Hideki, 93Miro Sabate, Carles, 136Mirska, Ilona, 26Mitchell, Alexander R., 50Miura, Hidetoshi, 93Nakanishi, Waro, 196Nawrot, Jan, 26Nea]u, Florentina, 102Ohta, Munehiro, 23Orme, Christine A., 50Pagoria, Philip F., 50Parvulescu, Vasile I., 102Pedireddi, V. R., 57Pernak, Juliusz, 26Pinilla, Elena, 125Pontellini, Roberto, 171Poole, Colin F., 76

Praly, Jean-Pierre, 148Prodi, Luca, 171Puchta, Ralph, 34Raghunathan, V. A., 112Rees, Neil V., 107Reichert, W. Matthew, 26Rogers, Robin D., 26Rose, Richard P., 64Rothbart, Sabine, 34Sasamori, Takahiro, 196Schrekker, H. S., 82Smiglak, Marcin, 26Sobkowski, Michal, 164Stasch, Andreas, 64Stasiewicz, Monika, 26Stawinski, Jacek, 164Stierstorfer, Jorg, 136Stracke, M. P., 82Sun, Zhengang, 119Talwelkar, Mayura S., 57Tanaka, Teppei, 23Tokitoh, Norihiro, 196Tong, Fei, 119Torres, M. Rosario, 125van Eldik, Rudi, 34Vidal, Sebastien, 148Xiao, Lei, 107Xue, Junmin, 88Yoshida, Tsukasa, 93Zhang, Jing, 119Zhang, Na, 119Zhao, Yan, 119Zhu, Yanyu, 119

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16 | New J. Chem., 2009, 33, 3–16 This journal is �c the Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2009

Light-responsive azobenzene groups have brought polymers to life

Polymers strut their stuff under the spotlight

Light can be used to make the polymer film move like a robotic arm

©The Royal Society of Chemistry 2009

Chemists in Japan have created light-driven polymer films that walk like inchworms and move like robotic arms.

The films, made by Tomiki Ikeda at the Tokyo Institute of Technology in Yokahama and collaborators, contain a polymer which contracts when visible light shines on it and expands again under UV light.

The polymers respond to light because they have azobenzene groups – which contain N=N double bonds – incorporated into them. Under visible light the N=N bonds have a cis conformation which means the polymer is bent. But when the light source is changed to UV the bonds become trans and the polymer flattens.

To make the polymer walk, the group incorporated it in a laminated film with one pointed end (at the back of the ‘worm’) and one flat end (at the front of the ‘worm’). As the polymer bends the pointed back end is dragged forward then, when the light source is changed to UV, the

January 2009 / Volume 6 / Issue 1 / ISSN 1478-6524 / CSHCBM / www.rsc.org/chemicalscience

Chemical Science

Glowing report for explosive detectionScientists are developing a new method to thwart terrorists

Crinkly tunnels aid gas storageTrifluorolactate crystals may offer an alternative route to hydrogen fuel cells

Porphyrins get energeticThis month’s Instant insight outlines recent advances in the construction of interlocked molecules

Lighting a billion livesNobel peace prize winner Rajendra K Pachauri talks about understanding climate change and giving light to humanity

In this issue

A snapshot of the latest developments from across the chemical sciences

ReferenceT Ikeda et al, J. Mater. Chem., 2009, 19, 60 (DOI: 10.01039/b815289f)

Chem. Sci., 2009, 6, C1–C8 C1

polymer flattens, pushing the front flat end forward. This continuous flattening–bending motion allows the film to move forward like an inchworm.

The robotic arm also requires clever lamination, but this time the polymer layer and laminated sections are alternated which

allows the film to act as a hinge joint and move flexibly. By controlling the intensity of the light and the position on the film where the light is concentrated, the researchers can make the film move as they chose.

‘The polymers function with a minimum of moving parts which minimises friction and surface contact problems,’ says Ikeda. ‘One can envisage applications such as direct light-to-mechanical energy conversion, storage systems and in microfluidic devices.’

Graeme George, an expert in polymer science at Queensland University of Technology, Brisbane, Australia, commends ‘the efficiency of the reversible photo-processes.’ He adds that the time is ripe for further detailed studies of such systems to see if these photo driven polymers offer any challenge to their electroactive counterparts.Ruth Doherty

CS.01.09.C1.indd 20 15/12/2008 15:00:16

Dung

Highlight

Dung

Highlight

French scientists have developed a photochemical method for patterning biomolecules inside glass capillary tubes. The technique could lead to lab-on-capillary devices as cheaper alternatives to lab-on-a-chip medical diagnostics, they claim.

Eric Defrancq, at Joseph Fourier University, Grenoble, and colleagues say their lab-on-capillary vision poses considerable challenges. ‘Retaining the functionality and patterning of biomolecules in the closed environment of a capillary tube is difficult because there is no easy access to the inside surface,’ explains Defrancq.

Defrancq overcame these challenges by grafting patterns of aminooxy groups masked with photocleavable protecting groups to the inside surface of the capillaries. By shining light on the tubes, he removed the protecting groups.

Capillary tubes offer cut price alternative to on-chip diagnostics

Cheaper than chips

Chemical Science

Austrian scientists unravel the secrets behind the dramatic colours of autumn.

Bernhard Kräutler and colleagues at the University of Innsbruck, have shown for the first time that a yellow breakdown product of chlorophyll contributes to the colours of autumn.

The change in autumn leaf colour is a phenomenon that affects the normally green leaves of many deciduous trees and shrubs. Every year, for a few weeks in autumn, a range of colours including intense yellows and reds shape the landscape. So far, these colours have been attributed to carotenoids and flavonoids, explains Simone Moser, a member of the research team. The colours are already present in the leaf, but are not visible due to the predominant green of chlorophyll. As autumn progresses chlorophyll disappears unmasking these hidden colours. But this is not the whole story, according to these researchers.

Why leaves turn red and orange during the autumn is not yet fully understood

Hints behind autumnal tints

Reference S Moser et al, Photochem. Photobiol. Sci., 2008, 7, 1577 (DOI: 10.1039/b813558d)

Biomolecules can be patterned to the inside of glass capillary tubes

Reference N Dendane et al, Lab Chip, 2008, 18, 2161 (DOI: 10.1039/b811786a)

Research highlights

C2 Chem. Sci. , 2009, 6, C1–C8 ©The Royal Society of Chemistry 2009

The exposed aminooxy groups then reacted with aldehyde groups in the peptide and carbohydrate biomolecules, fixing them to the side of the tubes. ‘This method of attaching molecules in patterns allows us to position not just one biomolecule but many,’ explains

The breakdown of chlorophyll is a process that was considered an enigma until about 20 years ago, when the first non-green chlorophyll breakdown product was discovered, says Moser. As these breakdown products were

colourless, they were thought not to contribute to the colours we see in autumn.

These compounds were considered to be the final products of chlorophyll breakdown, but now Kräutler has shown that they may be oxidised to give a yellow-coloured compound. Using leaves from the Katsura Tree, a deciduous tree known for its beautiful autumn leaves, they successfully detected this yellow chlorophyll breakdown product, thus proving its existence.

The similarity in structure between bilirubin, a natural compound reported to help protect cells from damage, and this oxidised breakdown product may suggest they too have important physiological properties. Moser says the team are interested in finding out just what roles, if any, these compounds play in the plant.Sarah Corcoran

Chlorophyll breakdown products may contribute to the colours of autumn

Defrancq.Once in place, the biomolecules

can bind cell proteins or antibodies present in biological fluids that are flushed through the tube. These protein and antibody biomarkers can be used to identify disease risk or progression or measure the effect of treatments. Different biomolecules can be attached in the one tube, which permits multi-analyses to be performed in one experiment.

‘The challenge now is to use these techniques to attach more complex carbohydrates and proteins without losing the recognition properties of the immobilised biomolecule for its target,’ says Defrancq. ‘It is only too easy to lose recognition during the immobilisation process through either the chemistry or the methods used.’ Janet Crombie

CS.01.09.C2.indd 20 15/12/2008 15:04:15

An European Union ban has reduced levels of the marine pollutant tributyltin

The tale of the snail

Reference E Sella and D Shabat, Chem. Commun., 2008, 5701 (DOI: 10.1039/b814855d)

Reference M Rato et al, J. Environ. Monit., 2009, DOI:10.1039/b810188d

Chemical Science

years, albeit at a slow rate,’ says Simon Apte, the leader of the centre of environmental contaminants research at the CSIRO in Sydney, Australia.

Although TBT levels are decreasing, they are still high, Rato comments. ‘Further monitoring surveys should be carried out in order to determine whether these EU measures are sufficient to reduce environmental TBT to a safe level,’ he adds.Rebecca Brodie

©The Royal Society of Chemistry 2009 Chem. Sci., 2009, 6, C1–C8 C3

Gender-switching in mud snails has decreased following a European Union ban on tributyltin (TBT) in ship hull paint.

TBT is an antifouling agent that was used in paint to prevent organisms from growing on the hulls of ships. Unfortunately, it was found that once it enters the water it has a toxic effect on other marine organisms. By 2003, the use of these paints was banned by the EU.

In the mud snails Nassarius reticulatus TBT was found to cause the imposex condition, where females develop male sexual characteristics such as a penis. Milene Rato from the University

of Aveiro, Portugal, and colleagues measured the penis lengths in the female snails to determine TBT pollution levels. ‘The main motivation to conduct this work was to find if the legislation implemented by the EU was effective,’ says Rato.

They found a decrease in the levels of imposex, and therefore a decrease in the levels of TBT, with hotspots being found within harbours that contain marinas and commercial fishing ports. From these results the group concluded that the regulation has had a favourable impact on pollution levels.

‘The data show that some recovery has occurred over the last five

Tributyltin causes imposex in female snails

a reporter group which fluoresces at 510 nanometres when released from the polymer structure. The trigger for the dendrimer breakdown is hydrogen peroxide, one of the natural decomposition products of the explosive.

Most colour-producing tests for TATP require the explosive to be pretreated with acid, say the scientists, so that it decomposes to produce large amounts of hydrogen peroxide. But this new method is sensitive enough to detect the tiny amounts of hydrogen peroxide generated by the small degree of natural decomposition of the

The dendrimer breaks down, and fluoresces, when exposed to hydrogen peroxide

Scientists are developing a new method to thwart terrorists

Glowing report for explosive detection

explosive and, because one molecule of hydrogen peroxide causes each dendrimer to release three fluorescent reporter molecules, a readable detection signal can be obtained for TATP present on the microgram scale.

Using more highly branched polymers each containing more fluorescently-tagged building blocks ‘will significantly increase the detection sensitivity’, says Shabat.

‘The main challenge’, he says, ‘will be to selectively identify TATP in the presence of other “powders” that contain oxidative species’.Freya Mearns

Israeli scientists have developed a sensitive method for detecting TATP – an explosive popular with terrorists.

Triacetone triperoxide, or TATP, is an explosive that has been used by suicide bombers in Israel since the 1980s. It was also employed by the thwarted British ‘shoe bomber’ Richard Reid in December 2001 and is alleged to have been used in the London bombings of July 2005.

The explosive’s ingredients are common chemicals and the material does not contain nitrogen so can pass through many scanners for nitrogenous explosives. Detection methods have been developed for TATP in the past, but now Eran Sella and Doron Shabat from Tel-Aviv University have designed a method that detects the explosive without any sample pretreatment, and also simply amplifies the resulting fluorescent signal. The scientists say that samples could be collected in the real-world using a swab or by vacuum.

Their method uses a type of dendrimer (a repeatedly branched tree-like polymer) that spontaneously breaks down into its separate building blocks following a single trigger event. The dendrimer was designed to consist of three building blocks that each contains

CS.01.09.C3.indd 30 15/12/2008 15:07:51

Japanese scientists have found a new way to store gases based on restraining gas molecules within narrow tunnels.

Gas storage in microporous materials generally relies on physisorption – involving weak Van der Waals interactions – to fill the micropores with gas, explains Toshimasa Katagiri, from Okayama University, who led the research team. Their new storage method involves physically restraining the gas within narrow tunnels (less than one nanometre diameter) running through nanoporous trifluorolactate crystals.

Katagiri and the team suggest that their new tunnel system may be useful for storing gaseous molecules with weak physisorption, such as hydrogen, and could have fuel cell applications.

The internal tunnel surface is serrated, thanks to the trifluoromethyl groups protruding into the tunnel cavity. These protrusions physically restrain the

Trifluorolactate crystals may offer an alternative route to hydrogen fuel cells

Crinkly tunnels aid gas storage

Chemical Science

European chemists have found that using fluorinated solvents in olefin metathesis reactions substantially improves the product yields obtained.

The metathesis reactions of alkenes (olefins) form a vital part of the armoury of transformations available to synthetic organic chemists. They provide a way of breaking and remaking carbon-carbon double bonds – allowing the substituent groups to be swapped – and are usually catalysed by transition metal complexes.

Commercially available catalysts, such as Grubbs’, remain popular among chemists but are often ineffective in more difficult reactions like the multi-step total synthesis of natural products and biologically active molecules.

Karol Grela, from the Polish Academy of Science, Warsaw,

A change of solvent found to dramatically improve important organic reactions

Fluorination gets a good reaction

Reference C Samojłowicz et al, Chem. Commun., 2008, 6282 (DOI: 10.1039/b816567j)

Reference T Katagiri et al, CrystEngComm, 2009, DOI: 10.1039/b814508c

C4 Chem. Sci. , 2009, 6, C1–C8 ©The Royal Society of Chemistry 2009

gas molecules within the tunnel. ‘The unique adsorption–desorption properties of these materials are very inspiring as they show the great potential of engineered hybrid systems where hydrocarbon and fluorocarbon domains alternate,’ comments Giuseppe Resnati, an expert in nanostructured materials at the Polytechnic of Milan, Italy.

The tunnel properties can be optimised to improve storage of a specific gas. Tunnel diameter, for

Poland, and colleagues have found that the yields of reactions using these catalysts can be dramatically improved by using fluorinated aromatic hydrocarbon solvents.

In particular, he reports that it is ‘possible to increase the metathesis reaction yield by up to 18 times by changing the solvent from 1,2-dichloroethane to perfluorotoluene.’

Jie Wu, professor of chemistry at Fudan University, Shanghai, China, adds: ‘This is an excellent improvement in metathesis reactions, which will find applications in the synthesis of advanced natural and biologically active compounds.’

Grela says that uncovering the nature of this effect and improving the recycling efficiency of the valuable catalysts – to satisfy the guidelines of green chemistry – are the next steps in his work.David ParkerFluorinated solvents

improve yields by up to 18 times

example, can be altered by changing the length of the organic chain in the trifluorolactates.

‘We are now trying to grow a perfect single crystal with tunnels. They could act as true molecular sieves for separating gaseous molecules by their size, at room temperature. Such a system would be a key technology for the realisation of a hydrogen fuel cell vehicle with a methanol reforming system,’ says Katagiri. Russell Johnson

Trifluorolactate crystals can be adapted to store different gases

CS.01.09.C4.indd 20 15/12/2008 15:15:33

Instant insightPorphyrins get energeticJonathan Faiz, Valérie Heitz and Jean-Pierre Sauvage, University of Strasbourg, France, outline recent advances in the construction of interlocked molecules inspired by photosynthesis


Artificially recreating photosynthesis – in the quest to find environmentally friendly and renewable energy sources – is a hot topic across many scientific disciplines. It is now known that porphyrin-like units are key features of the reaction centre where photosynthesis occurs, and synthetically reproducing these molecules has become a very active research area.

Porphyrins are planar and highly conjugated cyclic molecules that can complex a variety of metals. They are found in many natural systems, including blood (as hemoglobin in their iron-complexed forms) and in the photosynthetic reaction centre (as magnesium-complexed chlorins, chlorophylls – which are structurally very similar to porphyrins).

The electrochemical and photoactive properties of porphyrins make them ideal for performing energy and electron transfer in a similar fashion to photosynthesis. Importantly their ability to form noncovalent interactions, with a metal centre, can be exploited to form mechanically interlocked systems – such as catenanes and rotaxanes – that incorporate porphyrins in their structure.

One of the most remarkable features of catenanes (two or several interlocked rings) and rotaxanes (two-component assemblies consisting of a central thread encapsulated by a ring and stoppered by two bulky units on each end of the thread to stop the ring slipping off ) is their high flexibility, meaning they can undergo a very large number of different motions. These movements are important in photosynthesis as they facilitate electron transfer. The motions occur both naturally due to the molecules inherent energy (when all the components are not or only very weakly interacting), and when the molecule’s most stable geometry is altered by an external stimulus.

Although seemingly simple when

sketched on paper, the construction of catenanes or rotaxanes is not trivial. This is because attractive forces are needed to hold the components together, to template the reaction, before either the rings are closed (in the case of catenanes) or the stoppers are attached (rotaxanes).

One construction method is the use of transition-metal templates, where the various components of the macrocycle contain 1,10-phenanthroline units that can coordinate to copper(i) ions – holding the components in place. The rotaxanes and catenanes then form around the metal ion, that is removed once the macrocycle is constructed. This method has been used to make a comprehensive range of rotaxanes, with porphyrin stoppers, and catenanes, containing porphyrins rings.

Other templating methods for rotaxanes include the use of hydrogen-bonding or π-stacking interactions to either form a macrocycle around a thread already bearing stoppers or to hold the macrocycle around the thread whilst the stoppers are grafted. The size and metal-binding properties of porphyrins make them ideal for stoppers for rotaxanes. This

route has given access to a wide variety of architectures in which electron transfer can occur between porphyrins and, for example, fullerenes and electron-deficient aromatic macrocycles.

Rotaxanes have also been made with a manganese porphyrin ring that has an olefin-containing backbone threaded through it. The ring can zip along the backbone and catalyse the oxidation of the olefins in the backbone to epoxides – demonstrating the sheer breadth of application of these porphyrin-containing systems.

Mechanically interlocked porphyrin-containing architectures are important synthetic analogues of natural systems as they contain subunits held at predetermined distances and geometries – but not through conventional covalent bonds. In this way, just like natural systems, any intercomponent process that occurs between subunits takes place through the shortest pathways, such as through hydrogen bonds or solvent.

Read Jean-Pierre Sauvage’s tutorial review ‘Design and Synthesis of Porphyrin-Containing Catenanes and Rotaxanes’ in issue 2, 2009 of Chemical Society Reviews

Chem. Sci., 2009, 6, C1–C8 C5

A transition metal ion can hold together the components needed to make a catenane

ReferenceJ Faiz et al., Chem. Soc. Rev., 2009, DOI: 10.1039/b710908n

Chemical Science

Instant Insight_SAUVAGE.indd 40 15/12/2008 15:16:54

What do you think have been the key achievements of the Intergovernmental Panel on Climate Change (IPCC) during your time as chairman?I think the IPCC as a whole has been an extremely significant success story. What I think we have done rather well with the Fourth Assessment Report [an IPCC report on climate change] is to have closed a number of gaps in knowledge. We’ve produced some very clear statements, largely because the scientific basis is now much more robust. This includes a firm statement saying that warming of the climate system is unequivocal and that there’s a high probability that during the last five decades or so, the warming that has taken place is a result of human actions. An extremely significant step that we have been able to take is with respect to disseminating the results of this particular report.

What key areas will the IPCC focus on in the future? We’re actually in the midst of a detailed dialogue within the IPCC in defining what our role and focus should be in the future. The IPCC has decided to continue with the five or six year cycles of comprehensive assessment reports and the fifth assessment report will come out by 2014. This will require some new efforts in terms of developing scenarios of what’s going to happen in the future and running climate models to come up with some of the answers that will form the basis of the next report. In addition, we will carry out work on special reports, which would be in response to a need for focussed and very specific information on subjects of relevance. We’re already working on the production of a special report on renewable energy. So I think essentially we are going to build on what we have achieved so far and try to address demands as they come from our audience from all over the world and ensure that the IPCC plays the role that the world expects it to.

You are also Director General for the Energy and Resources Institute. What research are they currently involved in?This is an institute that I have been with for over a quarter of a century. When I started, all I had was a part time secretary and one room. We are now a fairly large institution with over 750 people and a presence in different parts of the world including the UK, the US, Japan and more recently Africa, the Middle East and Malaysia. We have emerged not merely as an institution that focuses on India or

other developing countries in the region but globally. Apart from the work on energy, climate change and environmental issues, we are also involved in substantial scientific activities, for instance biotechnology research. We do a substantial amount of work at the grass roots level, too – the one thing that I’m now focusing on, and which I think will be my mission for the next 10 years, is what I call ‘lighting a billion lives’. What’s extremely tragic is the fact that 1.6 billion people in the world still don’t have access to electricity or modern forms of energy – that’s a quarter of humanity. If you wait for all these places to be connected to the grid and to get electricity, it will take a long, long time. We have developed a set of solar lanterns and torches, which are really attractive to people in villages in several parts of the world. If we can mobilise the resources for making these available, it can create market based solutions in these villages.

What are your thoughts on the new RSC journal Energy & Environmental Science?I think that any such medium by which knowledge can be created and provided to people is an excellent initiative. I think the focus that you [the journal] have, which embraces all aspects of chemistry, chemical engineering and so on, will be of great value as it will be a major contribution to the creation of knowledge and the production of an area of literature where we still need an enormous amount of expansion and improvement.

You have recently been awarded an honorary doctorate from the University of East Anglia. What advice would you give to young scientists graduating today?Well, I would only say that this is a period of great excitement. We really have to start thinking outside the box. When we have this privilege of getting higher education, we should ensure that we do so because I really believe the world needs to change on a massive scale. That has to be carried out by the people who have the benefit of higher education and the exuberance of youth. So I would tell students that are in the university system right now to just look at the horizons beyond and also look at the world in its entirety. As I said in the acceptance speech for the Nobel Peace Prize, we have a Hindu saying, which is Vasudhaiva kutumbakam, which means ‘the universe is a family’. We have to keep that in focus whatever we learn and whatever we do.

Lighting a billion lives Rajendra K Pachauri speaks to Leanne Marle about shedding light on climate change and giving light to humanity

Interview

Rajendra K Pachauri is the current chairman of the Intergovernmental Panel on Climate Change and the director general of the Energy and Resources Institute in New Delhi, India.

Rajendra K Pachauri


Chemical Science

Chem. Sci., 2009, 6, C1–C8 C7

CT.interview pachauri.indd 15 15/12/2008 15:18:17

An InChI Resolver, a unique free service for scientists to share chemical structures and data, is to be developed via a collaboration between ChemZoo Inc., host of ChemSpider, and RSC Publishing.

Using the InChI – an IUPAC standard identifier for compounds – scientists can share, contribute and search molecular data from many web sources.

©The Royal Society of Chemistry 2009C8 Chem. Sci., 2009, 6, C1–C8

Essential elements

Double debut

InChI collaboration with ChemSpider

Chemical Science

And finally...Materials science researchers joined RSC Publishing last month at a celebration reception at the Fall MRS 2008 meeting. Authors and readers were thanked for their continued support, while RSC journal Soft Matter announced its increase in frequency for 2009 and five years of successful publication.

Delegates were invited to pre-order the latest edition of the bestselling textbook, Nanochemistry by Geoff Ozin, and take part in a prize draw to win a solar powered charger in celebration of the 2008 launch of Energy & Environmental Science.

Looking ahead, preparations are underway for the Third International ChemComm Symposium, which is to be held in China next month. The subject will be organic chemistry and keynote speakers include Professors Peter Kundig, Keiji Maruoka and Susan Gibson.

To find out more visit: www.rsc.org/chemcommsymposia

mice, cytotoxicity of chemical warfare degradation products, and identification and characterisation of metallodrug binding proteins. Visit www.rsc.org/metallomics

Authors from around the globe have submitted work of the highest quality, knowing that they can rely on RSC staff for overseeing a rigorous

peer-review process, efficient manuscript handling and rapid publication.

The current issues of both new journals are freely available to all readers via the website. Free institutional online access to all 2009/2010 content will be available following a simple registration process.

This month sees the debut of two highly interdisciplinary new journals from RSC Publishing: Integrative Biology: Quantitative biosciences from nano to macro and Metallomics: Integrated biometal science.

Integrative Biology is a unique journal focused on quantitative multiscale biology using enabling technologies and tools to exploit the convergence of biology with physics, chemistry, engineering, imaging and informatics. The first issue contains articles on human mammary progenitor cell fate decisions, the analysis of aptamer binding sequence–activity relationships using microarrays, and genome-wide transcriptome analysis of 150 cell

samples and much more. Visit www.rsc.org/ibiology

Metallomics covers the research fields related to metals in biological, environmental and clinical systems and is expected to be the core publication for the emerging metallomics community. First issue articles include a look at the effect of vanadium(IV) in diabetic

The InChI Resolver will give researchers the tools to create standard InChI data for their own compounds, create and use search engine-friendly InChIKeys to search for compounds, and deposit their data for others to use in the future.

‘The wider adoption and unambiguous use of the InChI standard will be an important development for the future of chemistry publishing, and further development of the semantic web,’ comments Robert Parker, managing director of RSC Publishing.

The InChI Resolver will be based on ChemSpider’s existing database of over 21 million chemical compounds and will provide the first stable environment to promote the use and sharing of compound data. ‘With the introduction of the InChI Resolver, we hope to expand the utility and value of both InChI and the ChemSpider service,’ adds Antony Williams of Chemspider.

This collaboration sees RSC Publishing remain at the forefront of chemical information technology.

Chemical Science (ISSN: 1478-6524) is published monthly by the Royal Society of Chemistry, Thomas Graham House, Science Park, Milton Road, Cambridge UK CB4 0WF. It is distributed free with Chemical Communications, Dalton Transactions, Organic & Biomolecular Chemistry, Journal of Materials Chemistry, Physical Chemistry Chemical Physics, Chemical Society Reviews, New Journal of Chemistry, and Journal of Environmental Monitoring. Chemical Science can also be purchased separately. 2009 annual subscription rate: £199; US $396. All orders accompanied by payment should be sent to Sales and Customer Services, RSC (address above). Tel +44 (0) 1223 432360, Fax +44 (0) 1223 426017. Email: [email protected]

Editor: Nina Notman

Deputy editor: Michael Spencelayh

Associate editors: Celia Gitterman, Joanne Thomson

Interviews editor: Elinor Richards

Web editors: James Hodge, Christina Hodkinson, Edward Morgan

Essential elements: Daniel Bradnam and Christine Hartshorne

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CS.01.09.C8.indd 8 16/12/2008 07:29:21

Changes ahead for NJC in 2009DOI: 10.1039/b820900f

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18 | New J. Chem., 2009, 33, 17–18 This journal is �c The Royal Society of Chemistry and the Centre National de la Recherche Scientifique 2009

Evidence of crystalline/glassy intermediates in bismuth phosphates

Marie Colmont,* Laurent Delevoye and Olivier Mentre

Received (in Montpellier, France) 3rd September 2008, Accepted 27th October 2008

First published as an Advance Article on the web 14th November 2008

DOI: 10.1039/b815388b

31P and 17O NMR investigations have been achieved on bismuth

oxide phosphates by a comparison between ordered and semi-

ordered reference compounds; the wide chemical shift range for17O is revealed to be a profitable source of information about

partially ordered materials.

Bi2O3–MO–P2O5, (M = Co, Cu, Cd, Zn, Mn. . .) ternary

systems have been well investigated, leading to the charac-

terization of new bismuth oxide phosphates having particular

structural relationships.1–8 As intensively detailed,9 the

rigid frameworks of these materials can be considered as an

assembly by the edge sharing of O(Bi,M)4 polyhedra, leading

to infinite polycationic ribbons of variable width: one, two,

three, . . . tetrahedra wide, surrounded by isolated PO4 groups.

From XRD/ND crystal structure studies, we classified these

compounds as ‘‘disordered’’ or ‘‘ordered’’, depending on the

competition (or not) between several O4 configurations around

the central P sites. Indeed, in ordered compounds such as

BiM2PO6,3,4,10,11 the Bi3+ cations strictly sit in the middle of

ribbons, whereas the M2+ cations are located at their edges.

Similarly, in disordered compounds, e.g. Bi1.2M1.2PO5.5,7 Bi3+

cations still occupy the middle of the ribbons, whereas the

edges of the ribbons are filled by mixed site Bi3+/M2+. This

statistical distribution leads to a variable orientation of the

PO4 groups, depending on the local nature of its first (Bi, M)

cationic shell. The disorder is all the more important because it

also affects partially filled cationic channels (so-called tunnels

hereafter) surrounded by PO4 groups in between pairs of

ribbons. Of course, the notion of disorder is inexact because

of the existence at the microscopic scale of incommensurate

modulated phenomena (mainly along b*) in most of the

disordered compounds.12–14 This extra information does not

survive over long range scales, e.g., it is not observed in the

XRD of single crystals. Therefore, sometimes only ordered

fragments of the disordered PO4/tunnel interstitial areas can

be assumed from the average crystal structure, on the basis of

plausible interatomic distances.13 However, structural interac-

tions between the edges of the ribbons, PO4 groups and the

tunnel is far from being fully established, probably due to

various phenomena, including anti-phase boundary defects

within tunnels and the probable semi-ordered zones in these

materials. In view of a complementary approach to these

fascinating series and by an easy extension to different com-

pounds, 31P and 17O NMR spectroscopy have been used as

local probes. Indeed, recent technical advances in solid state

NMR has led to the emergence of this technique by adapting it

for use with low natural abundance nuclei having relatively

small gyromagnetic ratios, such as 17O. In addition, it is worth

mentioning that impedance spectroscopy measurements on all

of these materials (ordered and disordered) show low ionic

mobilities due to the strong P–O bonds involved for most of

the oxygen ions. Therefore, only the static aspect is considered

hereafter.

Thus, the present work focuses on a comparison between a

typical ordered and disordered compound, with the aim of

establishing the pertinence, complementarities, and limits of

both 31P and 17O nuclei as probes with regard to the structural

aspects of ordered vs. semi-ordered materials. With that aim in

mind, two compounds have been selected from among the

series:

(i) BiCd2PO6 was chosen as the archetype of ordered

compounds. Its structure is isostructural to BiZn2PO6,15,16

and it is interesting because it crystallizes in the Bbmm space

group, while many members of the BiM2PO6 class (including

the M = Zn term) adopt the less symmetrical Pnma space

group. The coordination around its unique phosphorus posi-

tion is constituted by two independent O2 (2�) and O3 (2�)atoms, while O1 is located in the two tetrahedra-wide ribbons

at the center of a OBi2Cd2 tetrahedron (Fig. 1(a)).

(ii) The simplest disordered compounds have the

BiB1.2MB1.2PO5.5 general formula (M = Mn, Co, Zn). Their

structure (space group Icma) is formed of triple ribbons with

Fig. 1 The structures of (a) BiCd2PO6 and (b) Bi1.2Zn1.2PO5.5. BiCd2PO6

consist of [Cd4Bi2O2]-ordered double ribbons surrounded by six

isolated ordered phosphates. Bi1.2Zn1.2PO5.5 is disordered because of

(1) the presence of mixed Bi3+/Zn2+ sites at the edges of triple

[(Bi0.15Zn0.85)4Bi4O6] ribbons, (2) disordered tunnels, partially occupied

by Zn2+, and (3) multiple PO4 configurations around the same

phosphorus.

UCCS, Unite de Catalyse et Chimie de Lille, UMR-CNRS 8181,Ecole Nationale Superieure de Chimie de Lille, Universite desSciences et Technologies de Lille, BP 90108, 59655 Villeneuve d’Ascq,France. E-mail: [email protected]


LETTER www.rsc.org/njc | New Journal of Chemistry

mixed Bi/M edges (for M = Zn: 15% Bi3+/85% Zn2+). The

partially M-filled tunnels and surrounding disordered PO4

groups are shown in Fig. 1(b). Only one phosphorus position

exists, even if it has finally been split into two close satellites,

P1 (B50%) and P2 (B50%), in the published model. To avoid

any paramagnetic perturbation, the M = Zn compound was

selected. It is noteworthy that the influence of Cd2+ for Zn2+

replacement in BiM2PO6 on the 31P NMR chemical shift has

already been fully quantified on the basis of the empirical z/a2

parameter,15 and no additional contribution is expected

between these two neighboring cations. The possibility of

quantifying the local cationic environment of the PO4 groups

in a Bi(M,M0)PO6 statistical solid solution compounds has

also been enhanced.31PMASNMR: Fig. 2(a) and (b) show the 31P NMR spectra

of BiCd2PO6 against Bi1.2Zn1.2PO5.5, which clearly reveals the

broadening of the signal for the latter due to the multitude of

individual resonances in the disordered compound. It is com-

parable to the IR spectra of ordered vs. disordered compounds

presented elsewhere.13 In that sense, the broad envelope does

not show discrete contributions but rather a continuum. Here,

in addition to the local distortion of each individual PO4

group, the influence of the nature of the neighboring Zn/Bi

cationic shell has to be considered.15 Furthermore, the 31P

double quantum MAS-NMR spectrum shows no particular

privileged out-of-diagonal correlations (Fig. 2(c)) reminiscent

of a glass-like state from the 31P NMR resolution.

Since oxygen occupies both the polycationic regular

sublattice and the disordered interstitial regions, 17O NMR

analysis would be expected to give relevant information about

disorder. Here, samples were enriched via the 17O enrichment

method developed by Flambard et al.17 Due to the presence of

water vapor, the sample was checked by 1H NMR to ensure

that all protons disappeared at the end of the enrichment.

Another difficulty in obtaining 17O NMR spectra is the

presence of the quadrupolar interactions of the nuclei (spin

I = 5/2) that largely broaden signals. This requires suitable

techniques, such as double rotation (DOR),18 multiple-

quantum magic angle spinning (MQ-MAS)19 or satellite

transition magic angle spinning (ST-MAS),20 in order to

remove the anisotropic broadenings that remain under magic

angle spinning conditions.17O MAS NMR: Fig. 3 shows the high resolution MQ-MAS

spectra of (a) BiCd2PO6 and (b) Bi1.2Zn1.2PO5.5. The horizontal

projections (top) correspond to MAS spectra still broadened by

the second order quadrupolar interaction. The vertical projec-

tions reveal 17O isotropic spectra of the two compounds, where

the quadrupolar broadening is removed, i.e., each maximum

peak corresponds to a given oxygen environment. The

resonance at 90 ppm, marked with an asterisk, corresponds

to a spinning sideband of site A on the isotropic dimension.

The two spectra show two groups of resonances, around

Fig. 2 31P MAS-NMR (9.4 T) spectra of (a) BiCd2PO6 and (b)

Bi1.2Zn1.2PO5.5. The spectra were acquired at an MAS speed of

10 kHz, with a short pulse excitation of 1.5 ms (201) and a recycling

delay of 20 s. (c) 31P double quantum MAS-NMR spectrum of

Bi1.2Zn1.2PO5.5. The spinning frequency was 10 kHz. The excitation

and reconversion period was composed of back-to-back 901 pulses21 of

4 ms, which gave a total excitation/reconversion time of 400 ms. Therepetition time was 30 s, preceded by a presaturation period. A total of

16 scans were used and 64 t1 increments were collected. The31P chemical shift was referenced externally to an 85%H3PO4 solution

at 0 ppm.

Fig. 317O MQ-MAS NMR (18.8 T) spectra of (a) BiCd2PO6 and (b)

Bi1.2Zn1.2PO5.5. The spectra were acquired at an MAS speed of

20 kHz, with a recycling delay of 1 s, using the SPAM sequence.24

The excitation and reconversion pulses were set to 3.75 ms and 1.20 ms,respectively, corresponding to an RF field strength of 80 kHz,

followed by a selective 901 pulse of 11 ms (RF field of 8 kHz). For

spectrum (a), each transient was accumulated with 72 scans, and 128 t1

data points were collected using the STATES method. For (b), a total

of 4500 scans were needed and 30 t1 increments were collected. The17O chemical shift was referenced externally to tap water.


180–250 ppm and around 50–130 ppm (see the MAS projec-

tions). The region around 200 ppm (resonance A) is assigned to

O(Bi,M)4 tetrahedral sites in the polycationic ribbons. The

value of the chemical shifts are close to those determined for

OBi4 tetrahedra in related compounds, 195 ppm in Bi2O322 and

265 ppm in a-Bi4V2O11.23 Our assignment was indirectly con-

firmed by a 17O T2 relaxation measurement (using a saturation

recovery pulse sequence) performed on the BiCd2PO6 com-

pound. The A site exhibited a short T2 relaxation time of about

200 ms, maybe due to the presence of Bi quadrupolar nuclei in

its first-neighbour cationic shell (111Cd and 113Cd are non-

quadrupolar). A similar measurement was not possible for the

disordered compound due to the low efficiency of the isotopic

enrichment (probably because the 17O-enriched water had

already been used in previous experiments).

Ordered compound: The second region around 100–160 ppm

is typically in the chemical shift range of oxygens involved in

PO4 groups.25 In BiCd2PO6, it is composed of two resonances,

B (120 ppm, T2 = 5 ms) and C (100 ppm, T2 = 500 ms),corresponding to O2 and O3. This assignment arises from their

proximity or otherwise to quadrupolar Bi nuclei in their

second cationic shell (Table 1). It was checked by a 31P–17O

heteronuclear multiple quantum correlation (HMQC)26

experiment that a correlation existed between the unique31P site and the 17O–B sites (Fig. 4). However, no correlation

signal was detected for the 17O–C sites due to the very short T2

relaxation time (500 ms). It is also noteworthy that both A and

C showed broad isotropic resonances compared to B. So far,

this is not understood in this ‘‘ordered’’ compound. Note the

presence of a broad signal of low intensity in the 17O dimen-

sion (Fig. 4), which is due to an impurity obtained after the

process of enrichment and was not detected by XRD.

Semi-ordered compound: Next, we analyzed a semi-ordered

compound, Bi1.2Zn1.2PO5.5. The17O MQ-MAS NMR spec-

trum is shown in Fig. 2(b). Two isolated regions are high-

lighted in the 2D spectrum. The broadness of the peaks seen in

the isotropic projection is a signature of the high disorder

present in this system, as discussed in the first part of this

work. The assignment of both regions was deduced by analogy

with the 17O spectrum obtained for BiCd2PO6. One region

corresponds to oxygen atoms linked to ribbons (resonance A)

between 160 and 240 ppm. Referring to the structure presented

in Fig. 1(b), two oxygen sites should be distinguishable in the

isotropic dimension, whether they are at the center (O1Bi4) or

at the edge (O2Bi2Zn2) of the ribbon. A close look at the two-

dimensional contours clearly suggests the presence of more

than two sites, probably due to the high sensitivity of the 17O

NMR chemical shift to the cationic environment, even at a

semi-local scale (second shell cationic neighbours). The

presence of additional oxygen sites is easily explained by

the existence of a mixed Bi/Zn cationic site at the edge of

the ribbon. The assignment of each individual resonance is not

yet possible due to the current absence of a large 17O NMR

chemical shift database for these systems. The development of

such a database would require a series of model compounds

to be isotopically enriched for further 17O MAS NMR

analysis. The second option available is to profit from the

recent development of first principles calculations of NMR

parameters using periodic boundary conditions.27 The latter

approach, which is beyond the scope of this Letter, is definitely

more realistic at present.

The chemical shift region centred on 50–150 ppm exhibits

two main resonances in the isotropic projection. These can be

assigned to the oxygen atoms in the PO4 groups. First, it

should be noted that the 17O resonances are spread over a

large chemical shift range, especially in the isotropic dimension

when the second order quadrupolar broadening is removed.

This large distribution of chemical shift values with respect to

the so-called ordered compound, BiCd2PO6, reveals the

important disorder associated with the PO4 groups in these

systems. Nevertheless, some discontinuities associated with

Table 1 The environment of the oxygen atoms (distances in A) inBiCd2PO6 and Bi1.2Zn1.2PO5.5. The first shell is given for the oxygen ofthe ribbons and the first two shells are presented for the oxygen of thePO4 groups.

15,16

BiCd2PO6 Bi1.2Zn1.2PO5.5a

1st shell 2nd shell 1st shell

O1–Bi1 2 � 2.27(2) O2–Bi1 2 � 2.243(1)Cd2 2 � 2.18(2) Bi/Zn2 2 � 2.159(9)

O2–P 1 � 1.43(3) O1–Bi1 4 � 2.299(1)Cd2 2 � 2.22(2) 2 � 3.41(3)Bi1 1 � 3.70(3)

O3–P 1 � 1.51(4)Cd2 1 � 2.11(4) 2 � 3.29(2)Bi1 2 � 3.41(2)Bi1 1 � 3.65(4)

a The coordination of the disordered PO4 groups is not accurately

known.

Fig. 4 The 31P–17O HMQC (18.8 T) spectrum of BiCd2PO6 was

obtained at a MAS speed of 25 kHz by following the pulse sequence

detailed by Massiot et al.26 An echo was applied to the observed 17O

nuclei with respective 901 and 1801 pulses of 10 and 20 ms. Two 901

pulses of 4.5 ms were then applied on either sides of the 17O 1801 pulse.

The evolution delay was set to 3 ms for an evolution under J-coupling.

A total of 512 scans were accumulated with a recycling delay of 1 s.


discrete chemical shift values appear in the isotropic dimen-

sion. This strongly suggests that some PO4 positions are

privileged, as observed from the single crystal X-ray diffrac-

tion data. For example, two major competing PO4 positions

have been located in Bi1.2M1.2PO5.5, while residual electronic

density on Fourier difference maps is reminiscent of a number

of extra orientations. Again, the current 17O NMR chemical

shift database is not sufficient to fully interpret the spectrum in

the P–17O–M region.

This study highlights the informative data provided by 17O

NMR due to its broad range of chemical shifts compared to

the 31P nucleus. In disordered bismuth phosphates, 17O NMR

clearly provides evidence of preferential PO4 orientations, in

good agreement with the semi-ordering deduced from diffrac-

tion data, which shows modulated microdomains with loss of

the order at long range scales. Of course, the methodology and

conclusions developed here are not restricted to our particular

chemical system, but can be generalized to other partially-

ordered systems, solid solutions, composite structures and

so on. Finally, it is worth stating that, to the best of our

knowledge, a successful 31P–17O through-bond correlation has

been presented here for the first time.

Experimental section

The different oxides reported in this Letter were prepared from

stoichiometric mixtures of Bi2O3, MO (M = Cd, Zn) and

(NH4)2HPO4. To avoid the problem of volatile species, they

were removed by solid state reparative methods, implying

several heating–regrinding steps at temperatures from 200 to

700 1C. The purity of the samples was checked by powder

X-ray diffraction using a Siemens D-5000 diffractometer with

back-monochromatized Cu-Ka radiation. NMR experimental

information is given in the Figure captions. 17O enrichment

was achieved by heating samples at 650 1C for 8 h under17O-enriched water vapor.15

Acknowledgements

The FEDER Region Nord Pas-de-Calais, Ministere de

l’Education Nationale, de l’Enseignement Superieur et de la

Recherche, CNRS, USTL and ENSC-Lille are acknowledged

for funding the NMR spectrometers. M. C. thanks the Region

Nord Pas-de-Calais for financial support.

References

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2 M. Ketatni, O. Mentre, F. Abraham, F. Kzaiber and B. Mernari,J. Solid State Chem., 1998, 139, 274.

3 J. Huang, Q. Gu and A. W. Sleight, J. Solid State Chem., 1993,105, 599.

4 F. Abraham, M. Ketatni, G. Mairesse and B. Mernari, Eur. J.Solid State Inorg. Chem., 1994, 31, 313.

5 M. Colmont, M. Huve and O. Mentre, Inorg. Chem., 2006, 45(17),6612.

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A supramolecular sensing system for AgIat nanomolar levels by the

formation of a luminescent AgI–Tb

III–thiacalix[4]arene ternary complexw

Nobuhiko Iki,* Munehiro Ohta, Teppei Tanaka, Takayuki Horiuchi

and Hitoshi Hoshino

Received (in Durham, UK) 22nd September 2008, Accepted 10th November 2008

First published as an Advance Article on the web 1st December 2008

DOI: 10.1039/b816596c

The first example of the detection of AgI ions using supra-

molecular chemistry is demonstrated, in which two thiacalix[4]-

arene ligands are linked by analyte AgI ions and then coordinate

to TbIII

ions to form a luminescent ternary complex, AgI2�TbIII2�

TCAS2, enabling the detection of AgIat concentrations as low

as 3.2�10�9 M.

One of the most significant contributions of supramolecular

chemistry has been the development of a precise strategy to

design fluorescent chemosensors with high selectivities and

sensitivities for heavy metal ions.1 This strategy involves the

covalent joining of a specific binding unit of a metal ion and a

signal-transducing unit (Fig. 1).2 The former is a ligating

group that is carefully selected after considering factors that

will affect its selectivity, such as the affinity of donor atoms to

analyte cations and the stereochemistry of the resulting com-

plex. The latter is a fluorophore, whose photophysical proper-

ties are susceptible to changes such as excimer formation/

dissociation, photoinduced electron transfer, charge transfer

and energy transfer caused by metal binding. The validity of

this strategy, termed the covalent strategy, has been demon-

strated by many of the fluorescent sensors that have been

synthesized.2 For instance, a ratiometric sensor, where pyrene

is attached as a signaling unit to a ligand having N,O donors,

has been designed to enable the detection of AgI ions at

micromolar levels in a 50 : 50 v/v EtOH–water mixture.3

The strategy seems to have been derived on the premise that

two different processes occurs in analyte sensing—recognition

and signaling. Although this strategy is useful, it does not

provide much scope for alternative methods of designing

sensors or sensing systems. In this Letter, we present a system

for sensing AgI ions by the formation of a luminescent

complex using supramolecular chemistry,4 where the analytes

and components are synergistically assembled to function as a

sensor.

Since the development of a facile one-step method to

synthesize thiacalix[4]arene, we have been interested in

its inherent complexing properties and applications.5 For

example, thiacalix[4]arene-p-tetrasulfonate (TCAS, Fig. 2)

reacts with a TbIII ion to form a 1 : 1 complex, TbIII�TCAS (1),

in an aqueous solution at pH 4 8.5 by the ligation of a

bridging sulfur and two adjacent phenol oxygen donors.

Complex 1 exhibits strong luminescence due to the presence

of the TbIII ion, whose excitation energy is transferred from

TCAS in a triplet excited state.6 The luminescence of 1 allows

the detection of the TbIII ion at nanomolar levels.7 Further-

more, TCAS, TbIII and AgI ions form a luminescent ternary

complex, AgI2�TbIII2�TCAS2 (2), at a pH of around 6.8 In this

pH region, only a small fraction of the TbIII ions are com-

plexed by TCAS.6,8 This suggests that AgI can be detected by

measuring the luminescence of complex 2, which is formed in

the presence of TbIII ions and TCAS at a pH of 6 (see the

graphical abstract).

Accordingly, when [TbIII]T = 1.0 � 10�6 and [TCAS]T =

2.0 � 10�6 M at a pH of 6.1 (T = total), the dependence of the

luminescence intensity at 544 nm, assigned as the 5D4 -7F5

transition of TbIII, on the AgI concentration was investigated

(Fig. 3). For a wide range of AgI concentrations, the intensity

increased almost linearly as [AgI]T increased from nanomolar

to sub-micromolar levels. This demonstrates that AgI can be

detected by the formation of ternary complex 2. For higher

[AgI]T levels (42.0 � 10�7 M), the dependence showed a slight

upward convex curve. This can be attributed to the fact that

[AgI]T attains a concentration level equivalent to that of TbIII

and the availability of TbIII ions to form complex 2 is low.

When [AgI]T r 2.0 � 10�8 M, a linear calibration curve was

obtained by least-square fitting, as shown in eqn (1).

Luminescence intensity= 76.5� 108� ([AgI]T/M)+ 437 (1)

Surprisingly, the detection limit (DL) at S/N = 3 was

determined to be 3.2 � 10�9 M (0.35 ppb). This shows that

Fig. 1 Covalent strategy for designing metal sensors.

Graduate School of Environmental Studies, Tohoku University,6-6-07 Aramaki-Aoba, Aoba-ku, Sendai 980-8579, Japan.E-mail: [email protected]; Fax: +81 22-795-7293;Tel: +81 22-795-7222w Electronic supplementary information (ESI) available: Experimentaldetails for sample preparation and ESI-MS of complex 3.


LETTER www.rsc.org/njc | New Journal of Chemistry

the system is more sensitive than covalently designed fluor-

escent sensors, which afford the detection of AgI at the 10�6 M

level.3,9 Notably, the DL of AgI with 2 is lower than that of

flame atomic absorption spectrometry (DL 3 ppb) and as

low as that of inductively-coupled plasma atomic emission

spectroscopy (DL 0.2 ppb).10

The selectivity of this system with regard to AgI ions was

investigated by adding five times the amount of transition

metal cations (M = MnII, FeIII, CoII, NiII, CuII, ZnII, CdII

and PbII) and halide anions (X� = Cl�, Br� and I�) to a

1.0 � 10�7 M AgI ion solution. The luminescence intensity

(at 544 nm), I, was measured and compared to the intensity

measured in the absence of M or X�, I0. As shown in Fig. 4,

the five-fold increase in MnII and ZnII concentration did not

affect the signal intensity of complex 2; however, PbII, CoII

and NiII caused a slight change in its intensity. Notably, CuII

and FeIII ions caused negative interference (�67% and �57%,

respectively). In the TCAS–metal binary systems, CuII and

FeIII ions formed complexes with M : TCAS ratios of 2 : 1 and

1 : 1, respectively, at a pH of 6. If these complexes had been

formed in the present system, 1.75 � 10�7 and 1.5 � 10�7 M of

TCAS would have been available to AgI (1.0 � 10�7 M) to

form complex 2. Therefore, it is likely that CuII and FeIII ions

formed ternary complexes with TCAS and TbIII ions, thereby

reducing the availability of TbIII ions; this results in the for-

mation of an insufficient amount of 2. In addition, such an

M–TbIII–TCAS ternary complex would be non-luminescent

because paramagnetic CuII and FeIII ions readily quench the

excited states of the TCAS ligand. In contrast, CdII caused a

positive deviation (+116%) in the signal. Thus, it follows that

CdII should have formed a luminescent CdII–TbIII–TCAS

ternary complex that is luminescent, since CdII is a non-

quenching ion due to its d10 electronic configuration. In fact,

the CdII–TbIII–TCAS ternary system ([CdII]T = [TbIII]T =

1.0 � 10�6 M, [TCAS] = 2.0 � 10�6 M; pH = 6.5)

yielded a luminescent complex, whose composition was

CdII2�TbIII2�TCAS2 (3), as suggested by electrospray ionization-

mass spectroscopy (ESI-MS) measurements, yielding a peak at

m/z = 1101.5983 that is assignable to [2Cd2+ + 2Tb3+ +

Na+ + 3H+ + 2TCAS8� + H2O]2� (Fig. 5; also see ESIw).In the present system, complex 3, which was formed con-

comitantly, caused an increase in the luminescence. Among

the halide ions, iodide caused a negative (�43%) deviation from

the original intensity, I0, which can be attributed to its strong

ability to form the halo complexes [AgXn](n � 1)� (n = 1–4), as

indicated by their stability constants.11

In metal–ion sensors designed using a covalent strategy, the

roles of each functional group are different (Fig. 1). On the

other hand, in the present AgI sensing system, it is ambiguous

which moiety of 2 is responsible for the functions of binding

and signaling. As shown in the schematic drawing of 2 (Fig. 2),

TCAS has four O and four S donors that form the tetrametal

core, AgI2TbIII

2. Furthermore, there is an antenna present to

absorb photons, the energy from which is eventually trans-

ferred to the TbIII center. Upon excitation, the TbIII center

emits light via an f–f transition. From a structural point of

view, TbIII ions accept two sets of O,S,O donations from the

TCAS ligands. However, it is important to consider that in the

TbIII–TCAS binary system, TbIII does not form a complex

with TCAS at a pH of 6. Thus, analyte AgI is indispensable in

Fig. 2 The structure of TCAS, and complexes 1 and 2.

Fig. 3 Calibration graphs for AgI ions. The inset shows the calibra-

tion curve for the lowest AgI concentrations. Samples: [AgI]T =

0–100 � 10�8 M, [TbIII]T = 1.0 � 10�6 M, [TCAS]T = 2.0 � 10�6 M

and [MES buffer]T = 2 � 10�3 M (pH = 6.11). lex = 323 and

lem = 544 nm.

Fig. 4 The effect of a five-fold increase in concentration of ‘‘foreign

ions’’ added to AgI on the luminescence signal. I and I0 indicate the

luminescence intensity for samples with and without foreign ions,

respectively. Samples: [foreign ion]T = 0 or 5.0 � 10�7 M, [AgI]T =

1.0 � 10�7 M, [TbIII]T = 1.0 � 10�6 M, [TCAS]T = 2.0 � 10�6 M

and [MES buffer]T = 4 � 10�3 M (pH = 5.9). lex = 323 and

lem = 544 nm.


linking two TCAS ligands via S–AgI–S bridges to promote the

coordination of TCAS to TbIII, to form 2. In fact, TCAS

formed a 4 : 2 complex, AgI4�TCAS2, in the binary system at

pH 6.8 In conclusion, multidentate and photon-absorbing

TCAS, luminescent TbIII and analyte AgI, with a linear

coordination geometry, were synergistically assembled to form

a supramolecular structure that is capable of sensing AgI ions

at nanomolar concentrations (see graphical abstract). Since

the sensing function of this system originates from the supra-

molecular nature of complex 2, and not from TCAS and

TbIII individually, complex 2 truly demonstrates the ‘‘supra-

molecular strategy.’’ Here, it is very important to rationally

design molecules so that they form supramolecular assemblies

that display functionalities absent from their individual

components.

Experimental

Procedure for the detection of AgI ions

To a sample solution containing silver(I) nitrate and a

particular foreign ion, if any, appropriate amounts of

aqueous solutions of terbium(III) nitrate, TCAS, pH buffer

(2-morphorinoethanesulfonic acid (MES)) and doubly-

distilled water were added. Before the measurement of its

luminescence spectrum, each sample solution was allowed to

stand for 1 h at room temperature to ensure equilibration. The

luminescence spectra were measured using a Hitachi F-4500

fluorescent spectrometer.

Mass spectrometry

ESI-MS experiments were performed using a Fourier trans-

form ion cyclotron resonance mass spectrometer APEX III

(Bruker). Mass spectra were simulated using the program

iMass for Mac OS X version 1.1.12

Acknowledgements

This study was partly supported by a Grant-in-Aid for

Scientific Research (B) (16350039) from the Japan Society

for the Promotion of Science (JSPS).

References

1 J. M. Lehn, Supramolecular Chemistry, VCH, Weinheim, 1995.2 For reviews, see: A. P. de Silva, H. Q. Nimal Gunaratne,T. Gunnlaugsson, A. J. M. Huxley, C. P. McCoy, J. T.Rademacher and T. E. Rice, Chem. Rev., 1997, 97, 1515;B. Valeur and I. Leray, Coord. Chem. Rev., 2000, 205, 3;L. Prodi, F. Bolletta, M. Montalti and N. Zaccheroni, Coord.Chem. Rev., 2000, 205, 59.

3 R. H. Yang, W. H. Chan, A. W. M. Lee, P. F. Xia, H. K. Zhangand K. Li, J. Am. Chem. Soc., 2003, 125, 2884.

4 Examples of supramolecular sensing systems can be found in thefollowing reviews: E. V. Anslyn, J. Org. Chem., 2007, 72, 687;T. Hayashita, A. Yamauchi, A. J. Tong, J. C. Lee, B. D. Smith andN. Teramae, J. Inclusion Phenom. Macrocyclic Chem., 2004, 50, 87.

5 N. Morohashi, F. Narumi, N. Iki, T. Hattori and S. Miyano,Chem. Rev., 2006, 106, 5291.

6 N. Iki, T. Horiuchi, H. Oka, K. Koyama, N. Morohashi,C. Kabuto and S. Miyano, J. Chem. Soc., Perkin Trans. 2, 2001,2219.

7 T. Horiuchi, N. Iki, H. Oka and S. Miyano, Bull. Chem. Soc. Jpn.,2002, 75, 2615.

8 N. Iki, M. Ohta, T. Horiuchi and H. Hoshino, Chem.–Asian J.,2008, 3, 849.

9 For other examples, see: H. Tong, L. Wang, X. Jing and F. Wang,Macromolecules, 2002, 35, 7169; J. Raker and T. E. Glass, J. Org.Chem., 2001, 66, 6505.

10 J. D. Ingle, Jr. and S. R. Crouch, Spectrochemical Analysis,Prentice Hall, Englewood Cliffs, NJ, 1988.

11 R. M. Smith and A. E. Martell, Critical Stability Constants,Plenum Press, New York, 1976, vol. 4.

12 U. Rothlisberger, iMass for Mac OS X v. 1.1, 2002 (http://home.datacomm.ch/marvin/iMass/).

Fig. 5 Part of the ESI mass spectrum of complex 3, showing

the isotopomer pattern for [2Cd2+ + 2Tb3+ + Na+ + 3H+ +

2TCAS8� + H2O]2�. (a) Observed pattern for a sample ([TCAS]T =

[CdII]T = [TbIII]T = 2.5 � 10�5 M, [HCl]T = 5 � 10�5 M; pH 5.82

(adjusted with NH3)) and (b) simulated pattern.


Ionic liquids with dual biological function: sweet and anti-microbial,

hydrophobic quaternary ammonium-based saltsw

Whitney L. Hough-Troutman,a Marcin Smiglak,a Scott Griffin,a

W. Matthew Reichert,bIlona Mirska,

cJadwiga Jodynis-Liebert,

dTeresa Adamska,

d

Jan Nawrot,eMonika Stasiewicz,

fRobin D. Rogers*

agand Juliusz Pernak*

f

Received (in Montpellier, France) 30th July 2008, Accepted 29th August 2008

First published as an Advance Article on the web 22nd October 2008

DOI: 10.1039/b813213p

The dual nature of ionic liquids has been exploited to synthesize materials that contain two

independent biological functions by combining anti-bacterial quaternary ammonium compounds

with artificial sweetener anions. The synthesis and physical properties of eight new ionic liquids,

didecyldimethylammonium saccharinate ([DDA][Sac]), didecyldimethylammonium acesulfamate

([DDA][Ace]), benzalkonium saccharinate ([BA][Sac]), benzalkonium acesulfamate ([BA][Ace]),

hexadecylpyridinium saccharinate ([HEX][Sac]), hexadecylpyridinium acesulfamate ([HEX][Ace]),

3-hydroxy-1-octyloxymethylpyridinium saccharinate ([1-(OctOMe)-3-OH-Py][Sac]), and

3-hydroxy-1-octyloxymethylpyridinium acesulfamate ([1-(OctOMe)-3-OH-Py][Ace]), are reported,

as well as the single crystal structures for [HEX][Ace] and [1-(OctOMe)-3-OH-Py][Sac].

Determination of anti-microbial activities is described for six of the ILs. While some exhibited

decreased anti-microbial activity others showed a dramatic increase. For two of the ionic liquids,

[DDA][Sac] and [DDA][ACE], oral toxicity, skin irritation, and deterrent activity was also

established. Unfortunately, both ILs received a Category 4 (harmful) rating for oral toxicity and

skin irritation. However, deterrent activity experiments point to use as an insect deterrent, as both

ILs scored either ‘‘very good’’ or ‘‘good’’ against several types of insects.

Introduction

Ionic liquids (ILs) are currently defined as salts that are

composed solely of cations and anions which melt below

100 1C. These salts have been studied for a variety of applica-

tions such as in electrochemistry,1–3 separation science,4–7

chemical synthesis,8–13 and catalysis,14–16 however, until

recently, very few, if any, ILs had been used as liquid materials

themselves.17,18 In addition, those material applications which

have appeared, typically concentrated on a single desirable

property brought by either the cation or the anion. But ILs, by

definition, have at least two discrete types of ions, both of

which can provide a unique property or function. Thus, our

goal has been to explore how to exploit the dual nature of ILs

by preparing materials that possess two functions, particularly

two biological functions. Here, we present the combination of

anti-bacterial quaternary ammonium compounds (QACs)

with artificial sweeteners.

The anti-bacterial properties of QACs were first discovered

during the late 19th century, amongst carbonium dye com-

pounds, such as auramin, methyl violet, and malachite green.19

Initially, QACs were found to be most effective against gram-

positive organisms, until Jacobs and Heidelberger20–23 further

exploited their anti-bacterial properties against other types of

organisms. It was not until 1935 that the full potential of QACs

was recognized by the chemical community, when the synthesis

of benzalkonium chloride, a long-chain QAC, by Domagk24 and

further characterization of its anti-bacterial activities, proved that

QACs were effective against a wider variety of bacterial strains.

Later, in the 20th century, researchers became more

interested in the synthesis of water-soluble QACs for potential

applications as surfactants,25,26 anti-electrostatic agents,27

anti-corrosive agents,28 disinfectants,29 and phase-transfer

catalysts.30 These newly developed water-soluble QACs

showed anti-bacterial action against not only gram-positive

and gram-negative bacteria, but also pathogen species of fungi

and protozoa.31 These discoveries led to applications for

a The University of Alabama, Department of Chemistry and Center forGreen Manufacturing, Tuscaloosa, AL 35487, USA

bUS Naval Academy, Department of Chemistry, Annapolis,MD 21402, USA

cPoznan University of Medical Sciences, Department ofPharmaceutical Bacteriology, Swiecickiego, 4 60-781 Poznan, Poland

d Poznan University of Medical Sciences, Department of Toxicology,Dojazd 30, 60-631 Poznan, Poland

e Institute of Plant Protection, ul. Wegorka 20, 60-318 Poznan,Poland

f Poznan University of Technology, Faculty of Chemical Technology,pl. Sk!odowskiej-Curie 2, 60-965 Poznan, Poland.E-mail: [email protected]

g The Queen’s University of Belfast, QUILL, School of Chemistry andChemical Engineering, Belfast, Northern Ireland BT9 5AG.E-mail: [email protected]

w Electronic supplementary information (ESI) available: Charac-terization data. Fig. S1: ORTEP (50% probability thermal ellipsoids)of the asymmetric unit of [HEX][Ace]. Fig. S2: Close contacts aroundthe cations in [HEX][Ace]. Fig. S3: p-Stacking modes of the polymericcation in [HEX][Ace]. Fig. S4: p-Stacking mode of the dimeric cationin [HEX][Ace]. Fig. S5: ORTEP (50% probability thermal ellipsoids)of the asymmetric unit of [1-(OctOMe)-3-OH-Py][Sac]. CCDC refer-ence numbers 687477 and 687478. For ESI and crystallographic datain CIF or other electronic format see DOI: 10.1039/b813213p


PAPER www.rsc.org/njc | New Journal of Chemistry

QACs in wood preservation32–34 and as preservatives in

common household products,35 especially for general environ-

mental sanitation in hospitals and food production facilities.

Furthermore, QACs have been used as penetration enhancers

for transnasal and transbuccal drug delivery, such as nasal

vaccinations.36 The ability of QACs to penetrate and open cell

membranes has been widely used in drug delivery such as

liposomes, which consists of long alkyl chain QACs, and

non-viral gene delivery.37

We have had specific interest in employing the IL concept to

pair the biological activity of a class of compounds such as

QACs, with a second biological activity inherent in the

counterion.38 One such class of ions, which has also seen

independent use in preparing ‘edible’ ILs, includes non-

nutritive sweeteners such as saccharinate and acesulfame.39,40

Salts of these anions are currently used in food products and

are approved as food additives by most national and global

health agencies. Yet, only a handful of quaternary ammonium

saccharinates and acesulfamates have been reported in the

literature.41 Here we demonstrate the concept of preparing ILs

by pairing the biological activity inherent in the cation with a

separate biological function possessed by the anion with

the synthesis, physical properties, anti-microbial activities,

toxicity, and deterrent activity of new QAC-based ILs.

Results and discussion

Synthesis and characterization

Synthesis. Didecyldimethylammonium saccharinate ([DDA]-

[Sac]), didecyldimethylammonium acesulfamate ([DDA][Ace]),

benzalkonium saccharinate ([BA][Sac]), benzalkonium acesulfa-

mate ([BA][Ace]), hexadecylpyridinium saccharinate ([HEX][Sac]),

hexadecylpyridinium acesulfamate ([HEX][Ace]), 3-hydroxy-

1-octyloxymethylpyridinium saccharinate ([1-(OctOMe)-3-OH-

Py][Sac]) and 3-hydroxy-1-octyloxymethylpyridinium acesulfa-

mate ([1-(OctOMe)-3-OH-Py][Ace]) (Fig. 1) were prepared in

high yield as hydrophobic salts from commercially available

QACs benzalkonium chloride ([BA][Cl]), didecyldimethyl-

ammonium chloride ([DDA][Cl]), and hexadecylpyridinium

chloride ([HEX][Cl]), and from one pyridinium salt, 3-hydro-

xy-1-octyloxymethylpyridinium chloride [1-(OctOMe)-3-OH-

Py][Cl], which was prepared by a nucleophilic substitution

reaction of 3-hydroxypyridine by octyl chloromethyl ether

under anhydrous conditions. Each of the cations was paired

with saccharinate or acesulfamate by a stoichiometric metathesis

reaction in aqueous solution, using sodium saccharin ([Na]-

[Sac]) or potassium acesulfame ([K][Ace]). The hydrophobic

Fig. 1 Structures of the synthesized ILs.

Fig. 2 Packing diagram along the a crystallographic axis for [HEX][Ace]

(top) and overlay of the two cations in the asymmetric unit including the

anions with close contacts to each (bottom).


nature of these salts allowed them to be easily extracted from

the aqueous phase into chloroform.

All of the newly prepared ILs were found to be low melting

solids at room temperature with the exception of [DDA]-

containing salts, the only cation without an aromatic ring,

which were found to be liquid at room temperature. The salts

studied are only sparingly soluble in cold and hot water, but

freely soluble and stable in many organic solvents (e.g., chloro-

form, methanol, ethanol, ethyl acetate, N,N-dimethyl-

formamide (DMF), and dimethyl sulfoxide (DMSO)).

Crystal structures. Single-crystal structures for two of the

compounds, [HEX][Ace] and [1-(OctOMe)-3-OH-Py][Sac],

also confirmed the syntheses. Although not the focus of this

paper, interesting packing behavior was observed which may

provide clues to the low melting nature of these compounds in

particular and QAC ILs in general.

The packing diagram for [HEX][Ace] (Fig. 2) reveals that

the cation tails interdigitate to create charge-rich and hydro-

phobic regions. Closer examination indicates that the two

unique cations are not equivalent with slight differences in

the orientation of the hexadecyl tail groups. This modest

difference leads to completely different packing environments.

One cation p-stacks in a polymeric fashion (Fig. 3) and has

only three close contacts with the anions. The second cation

forms a p-stacked dimer with anions capping each open face.

These cations have five close contacts with the anions.

Fig. 4 illustrates the packing in the structure of [1-(OctOMe)-

3-OH-Py][Sac]. Here the strong hydrogen bonding between the

cation and anion dominates and a single cation/anion pair is

found in the asymmetric unit. These hydrogen bonded ion pairs

stack in alternate directions.

Thermal behavior. The thermal properties of the ILs

(Table 1) were determined by differential scanning calorimetry

(DSC) and thermogravimetric analysis (TGA). All of the

synthesized salts exhibited melting points below 100 1C,

allowing their classification as ILs. Interesting phase transition

behavior was observed for [DDA][Sac], [DDA][Ace] and

[HEX][Ace] which was not found for the other ILs. These

three ILs had a detectable glass transition-type transforma-

tion at �33, �53 and �11 1C, respectively. Following glass

transition, samples [DDA][Sac] and [Hex][Ace] exhibited

consecutive crystallization and melting transitions. On the

contrary, [DDA][Ace] was the only IL obtained that did not

exhibit any other thermal transition besides a glass transition.

As seen in Table 1, all the ILs were found to be thermally

stable to temperatures ranging between 160 and 210 1C. One-

step decomposition was found for [BA][Sac], [DDA][Sac],

[1-(OctOMe)-3-OH-Py][Sac] and [1-(OctOMe)-3-OH-Py][Ace].

The anions [Sac]� and [Ace]� normally display a two-step

decomposition, suggesting that the cations, [BA]+, [DDA]+

and [1-(OctOMe)-3-OH-Py]+, play a role in the decomposition

of these ILs resulting in the single decomposition step observed.

Fig. 3 One cation in [HEX][Ace] p-stacks in a polymeric fashion

(interplanar spacing 3.5 and 3.6 A) (top), while the second cation

forms p-stacked dimers (interplanar spacing 3.4 A) with acesulfamate

anions capping both sides (bottom).

Fig. 4 Packing diagram along the a crystallographic axis for

[1-(OctOMe)-3-OH-Py][Sac] (top) and close up of the hydrogen

bonding and alternate stacking of the ion pairs (bottom).


Two-step decomposition was observed for samples [HEX]-

[Sac], [DDA][Ace] and [HEX][Ace]. Increase in the thermal

stability (first decomposition step) in these salts, over the

thermal stabilities of the starting materials may indicate an

anion stabilizing effect on the parent cations; [HEX]+ and

[DDA]+. Similarly, the stabilizing effect of the anion can be

observed for the sample of [BA][Ace], which is the only sample

that exhibits a three-step decomposition pathway.

Biological properties

Anti-microbial, anti-bacterial and anti-fungal activities. The

minimum inhibitory concentration (MIC) (Table 2) and mini-

mum bactericidal or fungicidal concentration (MBC) (Table 3)

were determined for [BA][Sac], [DDA][Sac], [BA][Ace] and

[DDA][Ace]. (The starting materials, [BA][Cl] and [DDA][Cl],

which inherently exhibit anti-microbial, anti-bacterial and

anti-fungal activities, included in Tables 2 and 3 for com-

parison.) The activities of the ILs approach those of com-

mercially available [BA][Cl] and [DDA][Cl], although the ILs

were not found to be limited to a specific class of bacteria or

fungi. These same observations have been seen in previous

literature,42 where it was found that the anti-microbial

activities for imidazolium chlorides, tetrafluoroborates, and

hexafluorophosphates were independent of the counterion.

It is thought that 1-alkoxymethylpyridinium chlorides are

strongly active against microbes, yet in previous research,43 it

was concluded that the antimicrobial activities depended on the

substituent at the 3-position of the pyridine ring. Unfortu-

nately, [1-(OctOMe)-3-OH-Py][Cl] and the ILs, [1-(OctOMe)-

3-OH-Py][Sac] and [1-(OctOMe)-3-OH-Py][Ace] exhibited no

antimicrobial activity.

Acute oral toxicities. The acute oral toxicities of [DDA][Ace]

and [DDA][Sac] were determined in three male and three

female Wistar rats, where the rats received a dosage of

300 mg/kg b.w. (mg of substance per kg of body weight) and

2000 mg/kg b.w. of each IL. The ILs were suspended in water

Table 1 Thermal propertiesa

Tg Tc Ts�s Tm Tonset5% Tonset

Ionic liquids[BA][Sac] — 16b — 74 164 204[DDA][Sac] �33 15c — 16 187 214[HEX][Sac] — 30c — 66 207 253/412g

[1-(OctOMe)-3-OH-Py][Sac] — — — 95–98e 206 301[BA][Ace] — 30c �36 90 184 187/249/394g

[DDA][Ace] �53 — — — 189 232/426g

[HEX][Ace] �11 5b — 57 212 267/494g

18c

[1-(OctOMe)-3-OH-Py][Ace] — — — 79–81e 203 267

Starting materialsNa[Sac] — 98c — 120 431 459/541g

K[Ace] — — — 68 190 192/260g

[BA][Cl] — 16bd — — 143 169[DDA][Br]f — — — — 166 196[HEX][Cl] — 45c — 73 184 213[1-(OctOMe)-3-OH-Py][Cl] — — — 68–70e 178 247

a Phase transition points (1C) were measured from transition onset temperatures determined by DSC from the second heating cycle at 5 1C min�1,

after initially heating and then cooling of the samples to �100 1C unless otherwise indicated: Tg = glass transition temperature; Tc =

crystallization temperature; Ts–s = solid–solid transition temperature on heating; Tm = melting point on heating. Decomposition temperatures

were determined by TGA, heating at 5 1Cmin�1 under air atmosphere and are reported as (Tonset 5%) onset to 5 wt%mass loss and (Tonset) onset to

total mass loss. b Transition measured on heating cycle. c Transition measured on cooling cycle. d Transition only during first heating e Visual

melting point range via hot-plate apparatus. f Multiple transitions due to presence of water in starting material. g Multiple decomposition steps.

Table 2 MIC valuesa

Ionic liquid Starting materials

Strain [BA][Sac] [DDA][Sac] [BA][Ace] [DDA][Ace] [BA][Cl] [DDA][Cl]

S. aureus 4 4 4 8 2 2S. aureus (MRSA) 4 4 4 4 2 2E. faecium 8 8 8 8 4 4E. coli 16 16 31 16 8 8M. luteus 8 4 8 8 4 2S. epidermidis 4 4 4 4 2 2K. pneumoniae 4 4 8 4 4 4C. albicans 16 16 16 16 8 8R. rubra 16 16 16 16 8 4S. mutans 0.1 31 1 16 2 2Mean value 8.0 10.7 10.0 10.0 4.4 3.8

a In ppm.


prior to intragastric administration. After receiving the dosage

of 300 mg/kg b.w. for [DDA][Ace] or [DDA][Sac], one male

rat died during the first 24 h, while the other 5 rats remained

alive. But when the dosage was increased to 2000 mg/kg b.w.,

all of the rats died between 24 and 96 h after administra-

tion. Death was preceded by decrease in spontaneous motor

activity, excessive excretion from nostrils, and difficulty of

breathing. The above results indicate the acute toxicity range

for both ILs is between 300–2000 mg/kg b.w. in male and

female rats. Thus, these ILs would be classified as category 4

(harmful) toxins according to standard OECD grading.44

Skin irritation. Skin irritation of [DDA][Ace] and [DDA]-

[Sac] was determined on New Zealand albino rabbits. All of

the exposed animals exhibited defined erythema after 1 h. The

erythema had increased to severe and severe eschar formation

was also observed after 24 h. Although no edema occurred, the

skin irritation of these ILs is defined as category 4 (the highest)

by standard OECD grading.45

Deterrent activity. The deterrent activity of [DDA][Ace] and

[DDA][Sac] toward Tribolium confusum (larvae and beetles),

Sitophilus granarius (beetles) and Trogoderma granarium

(larvae) was determined by using a known method, in which

the amount of food consumed is monitored over a specific time

interval. Three deterrent coefficients had to be calculated from

the average amount of food consumed: (a) the absolute coeffi-

cient of deterrency, A = (CC � TT)/(CC + TT) � 100, (b) the

relative coefficient of deterrency, R = (C � T)/(C + T) � 100,

and (c) the total coefficient of deterrency, which is the sum of

the absolute and the relative coefficients, T=A+R.46 In these

Table 3 MBC valuesa

Ionic liquids Starting materials

Strain [BA][Sac] [DDA][Sac] [BA][Ace] [DDA][Ace] [BA][Cl] [DDA][Cl]

S. aureus 31.2 62.5 31.2 16 62.5 31.2S. aureus (MRSA) 31.2 31.2 31.2 31.2 31.2 31.2E. faecium 16 16 31.2 31.2 31.2 31.2E. coli 62.5 16 125 62.5 62.5 31.2M. luteus 62.5 31.2 62.5 62.5 31.2 31.2S. epidermidis 31.2 16 62.5 31.2 16 31.2K. pneumoniae 62.5 16 31.2 31.2 31.2 16C. albicans 31.2 16 31.2 31.2 16 16R. rubra 62.5 31.2 62.5 62.5 31.2 31.2S. mutans 0.5 62.5 16 125 16 16Mean value 39.1 29.9 48.5 48.5 32.9 26.6

a In ppm.

Table 4 Criteria for the estimation of the deterrent activity based onthe total coefficient

Total coefficient Deterrent activity

200–151 Very good150–101 Good100–51 Medium50–0 Weak

Table 5 Feeding deterrent activity

Ionic liquid Relative coefficient Absolute coefficient Total coefficient Deterrent activity

Sitophilus granarius (beetles)[DDA][Ace] 97.5 57.9 155.5 Very good[DDA][Sac] 57.8 56.6 114.5 GoodAzadirachtina 100.0 74.3 174.3 Very goodLSD0.05

b 57.8 28.8 60.1

Trogoderma granarium (larvae)[DDA][Ace] 94.0 85.0 179.0 Very good[DDA][Sac] 94.2 86.1 180.3 Very goodAzadirachtina 100.0 94.2 194.2 Very goodLSD0.05

b 0.3 7.6 7.8

Tribolium confusum (beetles)[DDA][Ace] 96.2 19.1 115.3 Good[DDA][Sac] 95.0 90.7 186.6 Very goodAzadirachtina 100.0 85.0 185.0 Very goodLSD0.05

b 0.6 9.2 9.0

Tribolium confusum (larvae)[DDA][Ace] 95.0 64.1 159.1 Very good[DDA][Sac] 95.3 88.8 184.1 Very goodAzadirachtina 100.0 88.4 188.4 Very goodLSD0.05

b 2.1 29.4 29.1

a Natural deterrent. b The least significant differences at the 5% level of significance.


equations, CC is the average weight of the food consumed in

the control, TT is the average weight of the food consumed in

the no-choice test, and T and C are the average weights of the

food consumed in the choice test.

The total coefficient value T is compared to standard values

for deterrent activity in Table 4, where a value of 0 equals

neutral activity and a value of +150 to +200 corresponds to

very high deterrent activity. The results of deterrent activity

for [DDA][Ace] and [DDA][Sac] are compared to a natural

deterrent, azadirachtin, in Table 5. The ILs received

either ‘very good’ or ‘good’ deterrent activity for all tested

insects. In particular, [DDA][Sac] exhibited the same deterrent

activity toward Tribolium confusum (larvae and beetles) as

azadirachtin and thus, could be classified as a potential

synthetic insect deterrent.

Conclusions

We have prepared ILs of two biologically active ions by

combining anti-microbial QACs cations with sweetener anions.

Some of these ILs demonstrate properties such as limited water

solubility, high thermal stability, and good deterrent activity

against insects; which suggest potential application as an

insecticide. Although oral toxicity and skin irritation values

were higher than hoped, there is still potential use, not only for

these ILs, but also for new, related ILs which can be prepared

by tuning the composition in such a manner to reduce toxicity.

In general, research in the IL field has begun to shift from

random combinations of ions to a design scheme in which

both the cation and anion are chosen based on the desired

physical, chemical, and biological properties. All procedures

performed on these animals were in accordance with esta-

blished guidelines and were reviewed and approved by the

University of Alabama’s Institutional Animal Care and Use

committee. As our fundamental understanding of IL behavior

increases, more control over the resultant properties of the

salts will be possible, and the number of potential applications,

such as those presented here, will continue to grow.

Experimental

Chemicals and microorganisms

Benzalkonium chloride [BA][Cl] (molecular formula

C6H5CH2N(CH3)2RCl where R = C12H25 (60%) and

C14H29 (40%)), didecyldimethylammonium bromide

[DDA][Br] (tech., 75 wt% gel in water), hexadecylpyridinium

chloride [HEX][Cl] (monohydrate, minimum 99%) and

sodium saccharinate Na[Sac] (hydrate, minimum 98%) were

purchased from Sigma Aldrich. Potassium acesulfamate

K[Ace] (Z 99%) was purchased from Fluka.

The following microorganisms were used: bacteria

Staphylococcus aureus ATCC 6538, Staphylococcus aureus

(MARSA) ATCC 43300, Enterococcus faecium ATCC

49474, Escherichia coli ATCC 2592,2 Micrococcus luteus

ATCC 9341, Staphylococcus epidermidis ATCC 12228,

Klebsiella pneumoniae ATCC 4352, and fungi Candida albicans

ATCC 10231, Rhodotorula rubra PhB and Streptococcus

mutans PCM (Polish Collection of Microorganisms) 2502.

The Rhodotorula rubra was obtained from the Department

of Pharmaceutical Bacteriology, Poznan University of

Medical Sciences, Poland.

General synthesis47

Solid (0.001 mol) Na[Sac] or K[Ace] was dissolved in distilled

water and added to hot aqueous solutions containing

0.001 mol of [BA][Cl], [DDA][Br] or [HEX][Cl]. The mixtures

were stirred at 60 1C for 1 h and then cooled to room

temperature. The hydrophobic product was extracted with

chloroform and purified using distilled water washes, until

chloride or bromide ions were no longer detected in the

product phase using AgNO3. The chloroform was evaporated

and the IL was dried under vacuum.

The starting material 3-hydroxy-1-octyloxymethyl-

pyridinium chloride was prepared according to previous

literature.43 Solid (0.03 mol) K[Ace] or Na[Sac] was dissolved

in distilled water and then added to an aqueous solution

containing 0.03 mol [1-(OctOMe)-3-OH-Py][Cl]. The reaction

was completed by gentle heating and stirring in a water bath

for 2 h. The heat was removed and stirring was continued at

room temperature for 24 h. The mixture was filtered, and the

precipitate was washed with cold distilled water (3� 20 mL) to

give an oil or solid IL. The IL was dried under vacuum, and

recrystallized from ethyl acetate and then dried again under

vacuum. Karl–Fischer analysis indicated the water content of

all dried ILs to be less than 500 ppm.

Thermal analysis

Melting points and other thermal transitions of the ILs were

determined by DSC, with a TA Instruments model 2920

Modulated DSC (Newcastle, DE), cooled with a liquid nitro-

gen cryostat. The calorimeter was calibrated for temperature

and cell constants using indium standard (mp 156.61 1C,

DH = 28.71 J g�1). Data were collected at constant atmos-

pheric pressure where the ILs were placed in aluminum pans

with sample sizes from 5 to 15 mg. An empty sample pan was

used as reference. All experiments were performed at a heating

rate and a cooling rate of 5 1C min�1. The DSC was adjusted

so zero heat flow was between 0 and �0.5 mW, and the

baseline drift was less than 0.1 mW over the temperature

range 0–180 1C.

Thermal decomposition temperatures were measured in the

dynamic heating regime using a TGA, 2950 TA Instrument,

under air atmosphere. The amount of IL used was between

2 and 10 mg in each case, and the samples were heated from

40 to 800 1C at a constant heating rate of 5 1C min�1.

Decomposition temperatures (T5%dec) were determined from

onset to 5 wt% mass loss; this provides a more realistic

representation of thermal stability at elevated temperatures.

X-Ray diffraction

Crystalline samples of [HEX][Ace] and [1-(OctOMe)-3-

OH-Py][Sac] were mounted on a glass fiber on a goniometer

head of a Siemens SMART CCD diffractometer equipped

with a Mo-Ka source (l = 0.71073 A) and a graphite

monochromator. Data collection was conducted at �100 1C

which was achieved by streaming cold nitrogen over the


crystal. Final unit cell parameters were determined by least-

squares refinement of the hemispherical data set obtained from

20 s exposures. Data were corrected for Lorentz and polariza-

tion effects and absorption using SADABS.48 The initial

structure solution was carried out using the direct methods

option in SHELXTL version 5.49 The positions of all non-

hydrogen atoms were refined anisotropically. The hydrogen

atoms were added and allowed to refine unconstrained in

order to obtain proper close contact interactions.

Crystal data for [HEX][Ace]. C25H42N2O4S; Mr = 466.67;

triclinic, space group P�1; T = 173 K; a = 7.921(3), b =

13.374(5), c= 25.689(10) A, a= 76.755(7), b= 82.225(7), g=89.260(7)1; Z=4; V=2624.2(17) A3;Dc = 1.181 g cm�3; 7459

independent (Rint = 0.0249) and 5755 observed ([I 4 2s(I)])reflections; GooF = 1.070; R1, wR2 [I 4 2s(I)] = 0.0462,

0.1208; R1, wR2 (all data) = 0.0641, 0.1411

Crystal data for [1-(OctOMe)-3-OH-Py][Sac]. C21H28N2O5S;

Mr = 420.51; triclinic, space group P�1; T = 173 K; a =

8.1626(15), b = 8.7141(16), c = 16.756(3) A, a = 81.872(3),

b = 80.780(3), g = 62.850(3)1; Z = 2; V = 1043.7(3) A3; Dc =

1.338 g cm�3; 2969 independent (Rint = 0.0170) and 2515

observed ([I 4 2s(I)]) reflections; GooF = 1.033; R1, wR2

[I4 2s(I)] = 0.0362, 0.0871; R1, wR2 (all data) = 0.0472, 0.0937

Antimicrobial characteristics

Anti-microbial activity was determined by the tube dilution

method. Bacteria strains were cultured in Mueller–Hinton

broth for 24 h and fungi were cultured on Sabouraud agar

for 48 h. Suspensions of the above microorganisms, at a

concentration of 106 cfu mL�1, were prepared from each

culture. Two milliliters of serial twofold dilutions of IL were

inoculated with the above-mentioned suspension to obtain a

final concentration of (1–5) � 105 cfu mL�1.

Growth of the microorganism (or its lack) was determined

visually after incubation for 24 h at 35 1C (bacteria) or 48 h at

22 1C (fungi). The lowest concentration at which there was no

visible growth (turbidity) was determined to be the minimal

inhibitory concentration (MIC). Then, from each tube con-

tent, 10 mL (calibrated loop) was smeared on an agar medium

with inactivates (0.3% lecithin, 3% polysorbate 80, and 0.1%

L-cysteine) and incubated for 48 h at 35 1C (bacteria) or for

5 days at 22 1C (fungi). The lowest concentration of the IL that

killed 99.9% or more of the microorganism was defined as the

minimum biocidal concentration (MBC).

Acute oral toxicity test

The toxicity was tested according to the method of acute toxic

class.44 Three male (250 � 25 g) and three female (170 � 17 g)

Wistar rats were used for each IL tested. The ILs were first

suspended in distilled water and then administered intra-

gastrically at doses of 300 mg/kg b.w. and 2000 mg/kg b.w.

After the dose was administered, the rats were observed for

14 days.

Skin irritation tests

Each IL was tested on 3 male New Zealand albino rabbits,

where the fur was previously removed from the back of the

rabbit. Half a milliliter of the ILs (100%, pure) was distributed

on two 6 cm3 sites of the same animal. The application site was

then covered with a porous gauze dressing and secured in place

with tape. After a 4 h exposure, the dressing was removed and

the application site was gently washed with water. Observa-

tions were then conducted at 1, 24, 48, and 72 h, where the

test sites were evaluated for erythema and edema using a

prescribed scale.45

Feeding deterrent activity tests

Three species of insects were selected for testing: Tribolium

confusum Duv. (larvae and beetles), Sitophilus granarius L.

(beetles), and Trogoderma granarium Ev. (larvae). Insects were

grown on a wheat grain or whole-wheat meal diet in labora-

tory colonies which was maintained at 26 � 1 1C and 60 � 5%

relative humidity. The laboratory assay was conducted

according to the method developed and standardized for storage

insects feeding activity for both choice and no-choice test.46

Wheat wafer discs (1 cm in diameter � 1 mm thick) were

saturated by dipping in either ethanol (96%) only (control) or

in a 1% ethanol solution of [DDA][Ace] or [DDA][Sac]. After

evaporation of the solvent by air-drying (30 min), the wafers

were weighed and offered as the only food source for the

insects over a five day period. The feeding of the insects was

recorded under three conditions: (a) control test (two control

discs (CC)), (b) choice test (a choice between one treated

disc (T) and one control disc (C)), and (c) no-choice test

(two treated discs (TT)). Each of the three experiments was

repeated five times with 3 beetles of Sitophilus granarius,

20 beetles and 10 larvae of Tribolium confusum, and 10 larvae

of Trogoderma granarium. The number of individual insects

depended on the intensity of their food consumption. The

beetles utilized in the experiments were unsexed, 7–10 days

old, and the larvae were 5–30 days old. After five days of

feeding, the discs were reweighed. The data from the experi-

ments have been statistically corrected by an analysis of

variance.

Acknowledgements

This work was supported by Poznan University of Techno-

logy, BW 32-222/2008.

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ISSN 1144-0546










Volume

18|N

umber1

|2008Journal of M

aterials Chem

istryPages

1–140


ISSN 1757-9694

0959-9428(2008)18:1;1-J


1757-9694(2009) 1:1;1

0608


Metal ion-catalyzed oxidative degradation of Orange II by H2O2. High

catalytic activity of simple manganese salts

Erika Ember, Sabine Rothbart, Ralph Puchta and Rudi van Eldik*

Received (in Montpellier) 6th August 2008, Accepted 25th September 2008


DOI: 10.1039/b813725k

In an effort to develop new routes for the clean oxidation of non-biodegradable organic dyes,

a detailed study of some environmentally friendly Mn(II) salts that form very efficient in situ

catalysts for the activation of H2O2 in the oxidation of substrates such as Orange II under mild

reaction conditions, was performed. The studied systems have advantages from the viewpoint of

green chemistry in that simple metal salts can be used as very efficient catalyst precursors and

H2O2 is used as a green oxygen donor reagent. Oxidations were carried out in a glass reactor over

a wide pH range in aqueous solution at room temperature. Under optimized conditions it was

possible to degrade Orange II in a carbonate buffer solution in less then 100 s using 0.01 M H2O2

in the presence of only 2 � 10�5 M Mn(II) salt. To gain insight into the manganese catalyzed

oxidation mechanism, the formation of the active catalyst was followed spectrophotometrically

and appears to be the initiating step in the oxidative degradation of the dye. High valent

manganese oxo species are instable in the absence of a stabilizing coordinating ligand and lead to

a rapid formation of catalytically inactive MnO2. In this context, the role of the organic dye and

HCO3� as potential stabilizing ligands was studied in detail. In situ UV-Vis spectrophotometric

measurements were performed to study the effect of pH and carbonate concentration of the buffer

solution on the formation of the catalytically active species. Electrochemical measurements and

DFT (B3LYP/LANL2DZp) calculations were used to study the in situ formation of the catalytic

species. The catalytic cycle could be repeated several times and demonstrated an excellent stability

of the catalytic species during the oxidation process. A mechanism that accounts for the

experimental observations is proposed for the overall catalytic cycle.

Introduction

Nowadays, one of the major environmental problems con-

cerns the strong increase in xenobiotic and organic substances

that are persistent in the natural ecosystem. Most of these

compounds have an aromatic structure, which makes them

highly stable and thus difficult to degrade.1 A significant

source of environmental pollution is industrial dye waste due

to their visibility and recalcitrance, since dyes are highly

coloured and designed to resist chemical, biochemical and

photochemical degradation.2 About half of the global produc-

tion of synthetic dyes (700 000 t per year) are classified as

aromatic azo compounds that have a –NQN– unit as chromo-

phore in their molecular structure. Over 15% of textile dyes

are lost in waste water streams during the dyeing operation.3

Azo dyes are known to be largely non-biodegradable under

aerobic conditions and to be reduced to more hazardous

intermediates under anaerobic conditions.4 The decolorization

of wastewater has acquired increasingly importance in recent

years, however, there is no simple solution to this problem

because the conventional physicochemical methods are costly

and lead to the accumulation of sludges.5

One approach to solve these problems would be to develop

low-cost, highly efficient, and environmental friendly oxida-

tion catalysts on the basis of transition metal complexes.6,7

Recently, photodegradation methods based on TiO2 as a

photocatalyst,8 beside Fenton systems,9 emerged as one of

the most promising technologies and received increasing atten-

tion due to their practical and potential value in environmental

protection. However, in some cases they are only successful

under specific pH and temperature conditions.

Several studies were performed during the last few years in

order to find good catalysts for the oxidative degradation of

different organic dyes. From an environmental point of view,

first row transition metals are the most challenging. Highly

effective Fe,10,14 Co,11 Cr12 and Mn13 based oxidation catalyst

were developed. In combination with different oxidizing

agents, the decomposition of stable organic substances was

possible. A novel highly active and environmental benign

catalytic system based on Fe-TAML (TAML = tetraamido

macrocyclic ligand) was recently reported by Chahbane et al.14

In many cases tremendous synthetic efforts are required to

obtain an effective catalytic system and in addition the pre-

sence of high concentrations of oxidizing agents is needed.

Among the possible oxidizing agents, H2O2 is one of the most

commonly used owing to its eco-friendly nature. The use of

H2O2 as a green oxidizing agent in these reactions is justified

by a low organic content of the wastewater to be treated and a

Inorganic Chemistry, Department of Chemistry and Pharmacy,University of Erlangen-Nurnberg, Egerlandstr. 1, 91058 Erlangen,Germany



low reaction temperature, thus requiring the presence of an

adequate catalyst due to the high kinetic activation barrier of

such reactions. Commonly used methods for activation of

H2O2 include the formation of reactive peroxyacids from

carboxylic acids and peroxycarboximidic acid from aceto-

nitrile (Payne oxidation),15 the generation of peroxyisourea

from carbodiimide in the presence of either a weak acid or a

mild base,16 or the use of percarbonate, persulfate or perbo-

rate in strongly basic solution.17 In order to achieve fast

oxidative transformations, the use of large amounts of co-

catalyst additives is often required.18

Among these, the use of percarbonate, a versatile oxidizing

agent, is preferred for environmental reasons.19,20 Oxidation

using environmentally benign oxidants has aroused much

interest,7,21 because chemical industry continues to require

cleaner oxidation, which is an advance over environmentally

unfavoured oxidations and a step up from more costly organic

peroxides.22

In this report, we propose a fast and clean catalytic oxida-

tive degradation of Orange II as model substrate by H2O2 in

aqueous carbonate solution under mild reaction conditions,

pH 8–10 and 25 1C, eqn (1).

ð1Þ

Starting from commercially available Mn(NO3)2 in aqueous

carbonate solution for catalytic applications, various aspects

of the in situ generation of very reactive high valent manganese

intermediates in the presence of H2O2 were studied. Baes and

Mesmer have shown that manganese salts in aqueous solution

are able to form very reactive aquated intermediates.23 More-

over, in an alkaline medium, the introduction of a hydroxy

ligand trans to a water ligand is expected to produce more

labile OH–Mn–H2O species, and their formation (eqn (2)) is

considered to be of major importance for their catalytic

activity.

In the present study, the formation of catalytically inactive

Mn(OH)2 species was observed at higher pH, leading to

deactivation of the producedMn intermediates. The activation

of H2O2 in the presence of manganese salts as a function of pH

and carbonate concentration was therefore monitored using

UV-Vis spectrophotometry. In situ formed, high valent man-

ganese intermediates are known to be highly unstable in the

absence of a spectator ligand. As the study progressed, it was

of importance to investigate the role of the azo dye as a

potential coordinating ligand to stabilize the produced inter-

mediate under different reaction conditions. Electrochemical

measurements and DFT calculations were used to develop

a better understanding of the coordination chemistry of

Orange II. The successful implementation of such catalytic

systems becomes a worthwhile objective when issues such as

environmental compatibility, high atom economy, availability,

and expenses are considered.24

Experimental

Chemicals

Orange II, certified [Acid Orange 7, C.I. 15510, sodium 4-(2-

hydroxy-1-naphthylazo)benzenesulfonate], 99% was supplied

by Sigma–Aldrich and recrystallised from a Et2O/H2O mix-

ture at 4 1C. 2,4,6-Tri-tert-butylphenol (TTBP) 96% was

purchased from Sigma–Aldrich and recrystallised several

times from EtOH/H2O (9 : 1) mixtures prior to use. Hydrogen

peroxide 35 wt% as well as different manganese salt hydrates

used in the experiments, were of analytical grade and provided

by Acros Organics (Germany). Carbonate buffer solutions

were prepared using Millipore Milli-Q purified water.

General procedure

The manganese salts were freshly dissolved in water before

use. To a freshly prepared sodium carbonate solution, an

adequate amount of NaOH was added to adjust the pH of

the solution. Under isothermal conditions, the desired amount

of a concentrated manganese solution was added together with

Orange II, previously dissolved in an aqueous carbonate

solution, and H2O2. In typical measurements, 0.01 M H2O2

was prepared from a 35 wt% solution of H2O2. In addition, to

gain more information on the activation mode of the catalyst,

two further experimental procedures based on different activa-

tion and stabilization modes of the activated catalyst, were

followed. In one, the catalytic active species was generated

in situ in the carbonate buffer solution by addition of the

desired amount of H2O2, followed by the addition of the

corresponding quantity of Orange II to the reaction mixture.

In the other, Orange II was added to the manganese solution

and the formation of an Orange II� � �MnII complex was

observed. The decomposition of the dye was initiated through

the subsequent addition of H2O2. It is important to note that

the catalytic oxidation of the dye by H2O2 could only be

performed in an aqueous carbonate buffer solution. No other

buffer at the same pH, viz. TRIS, TAPS, HEPES or phos-

phate, showed the observed catalytic reaction.

Kinetic study of the manganese catalysed oxidative degradation

of Orange II by H2O2

All kinetic data were obtained by recording time-resolved

UV-Vis spectra using a Hellma 661.502-QX quartz Suprasil

immersion probe attached via optical cables to a 150 W Xe

lamp and a multi-wavelength J & M detector, which records

complete absorption spectra at constant time intervals. In a

thermostated open glass reactor vessel equipped with a mag-

netic stirrer, a 2 � 10�5 M freshly prepared catalyst solution

and 0.01 M H2O2 were added to 40 ml of 5 � 10�5 M dye

at a pH ranging from 8 to 10 at 25 1C. All kinetic measure-

ments were carried out under pseudo-first order conditions


(i.e. 50 r [H2O2]/[Mn2+] r 1000). The pH of the aqueous

carbonate solution was carefully measured using a Mettler

Delta 350 pH meter previously calibrated with standard buffer

solutions at two different pH values (4 and 10). The kinetics

of the oxidation reaction was monitored at 480 nm. First

order rate constants, where possible, were calculated using

Specfit/32 and Origin (version 7.5) software. To estimate the

effect of the catalyst and H2O2 concentrations on the catalytic

reaction at different carbonate concentrations, stopped-flow

kinetic measurements were carried out using an SX.18MV

stopped-flow instrument from Applied Photophysics.

Spectrophotometric titration

UV-Vis spectra were recorded on a Shimadzu UV-2101

spectrophotometer at 25 1C. A 0.88 cm path length tandem

cuvette with two separate compartments (0.44 cm path length

each), was filled with 1 ml 5� 10�5 M Orange II stock solution

in one, and different concentrations of an aqueous Mn(NO3)2solution in the other compartment. The cuvette was placed in

the thermostated cell holder of the spectrophotometer for

10 min. UV-Vis spectra were recorded before and after mixing

the solutions. The resulting spectrum presents the sum of the

two individual spectra before, and that of the reaction mixture,

after mixing. The observed spectral change is a result of

complex-formation between Mn(II) and Orange II.

Cyclovoltammetric measurements

Cyclovoltammetric (CV) measurements were performed in a

one-compartment three-electrode cell using a gold working

electrode (Metrohm) with a geometrical surface of 0.7 cm2

connected to a silver wire pseudo-reference electrode and a

platinum wire serving as counter electrode (Metrohm).

Measurements were recorded with an Autolab PGSTAT

30 unit at room temperature. The working electrode surface

was cleaned using 0.05 mm alumina, sonicated and washed

with water every time before use. The working volume of 10 ml

was deaerated by passing a stream of high purity N2 through

the solution for 15 min prior to the measurements and then

maintaining an inert atmosphere of N2 over the solution

during the measurements. All CVs were recorded for the

reaction mixture with a sweep rate of 50 mV s�1 at 25 1C.

Potentials were measured in a 0.5 M NaCl/NaOH electrolyte

solution and are reported vs. an Ag/AgCl electrode.

IR measurements

IR spectra were recorded as KBr pellets using a Mattson

Infinity FTIR instrument (60 AR) at 4 cm�1 resolution in

the 400–4000 cm�1 range.

Elemental analysis

The measurements were carried out on an elemental analyzer

Euro EA 3000 instrument from Hekaltech Gmbh. The analy-

tical method is based on the complete instantaneous oxidation

of the sample by ‘‘flash combustion’’ at 1000 1C, which

converts all organic and inorganic substances into combustion

products. The resulting combustion gases are swept into the

chromatographic column by the carrier gas (He) where they

are separated and detected by a thermal conductivity detector.

DFT calculations

Unrestricted B3LYP/LANL2DZp hybrid density functional

calculations,25a–c i.e., with pseudo-potentials on the heavy

elements and the valence basis set25d–f augmented with

polarization functions,25g were carried out using the Gaussian

0326 suite of programs. The relative energies were corrected

for zero point vibrational energies (ZPE). The resulting struc-

tures were characterized as minima by computation of vibra-

tional frequencies, and the wave functions were tested for

stability.

Synthesis of insoluble MnCO3

In a 150 ml round flask 3.36 g (0.4 M) NaHCO3 was dissolved

in 100 ml doubly distilled water and the pH of the solution was

set at 8.5 upon addition of small amounts of concentrated

NaOH solution. To the freshly prepared carbonate solution

1 g (0.04 M) Mn(NO3)2 was added. The mixture was stirred at

room temperature for 15 min during which MnCO3� � �H2O

formed as a white precipitate. The product was filtered and

washed several times with large amounts of water. Yield:

0.44 g MnCO3, 96.2%. IR (KBr pellets): n (cm�1) 3421 (m),

1416 (vs), 862 (s), 725 (m). Elemental analysis (%) for

MnCH2O4: calc.: C 9.03, H 1.52; found: C 9.38, H 1.52.

Synthesis of Orange II� � �MnII complex

In a 50 ml Schlenk tube 0.014 g (2 � 10�3 M) Orange II was

dissolved in 20 ml doubly distilled water and an aqueous

solution of 0.01 g (2 � 10�3 M) Mn(NO3)2 was added

dropwise under continuous stirring. The solution mixture

was kept for several hours at room temperature. The formed

precipitate was filtered and dried at room temperature. Yield:

0.018 g Orange II� � �MnII, 87.1%. IR (KBr pellets): n (cm�1)

3527 (vs), 1619 (s), 1511 (s), 1383 (vs), 1262 (m), 1171 (s),

1120 (s), 1034 (s), 1007 (s), 829 (s), 759 (s), 696 (m), 644 (m),

595 (m). Elemental analysis (%) for MnC16H18O10N3SNa:

calc.: C 36.79, H 3.47, N 8.04, S 6.14, O 30.63; found:

C 29.44, H 3.39, N 8.05, S 4.77, O 30.31.

Synthesis of Orange II� � �MnII� � �Orange II complex

An aqueous solution of 0.005 g (1 � 10�3 M) Mn(NO3)2 was

added under continuous stirring to a 0.014 g (2 � 10�3 M)

Orange II water solution at room temperature. The pale

yellow precipitate was collected by filtration and dried in

air. Yield: 0.017 g Orange II� � �MnII� � �Orange II, 93.7%. IR

(KBr pellets): n (cm�1) 3390 (s), 1619 (s), 1570 (m), 1554 (m),

1520 (vs), 1393 (m), 1260 (m), 1169 (vs), 1119 (vs), 1033 (vs),

1007 (s), 828 (s), 758 (s), 695 (m), 644 (m), 593 (m). Elemental

analysis (%) for MnC32H32O15N5S2Na2: calc.: C 43.1, H 3.62,

N 7.23, S 7.19, O 26.91; found: C 42.99, H 3.75, N 7.23, S 7.00,

O 25.67.


General observations

A series of experiments were performed in order to investigate

the in situ generation of the highly reactive manganese catalyst

in the oxidative degradation of Orange II by H2O2 under mild

reaction conditions starting with a simple Mn(II) salt.


Oxidation reactions are in general affected by the protonation

state of the substrate, catalyst and oxidant, and the solvent

used. It is further important to note that the studied organic

dye (Orange II) can exist in either one of two tautomeric

forms, or in an equilibrium mixture, depending on the process

parameters. This kind of rapid dynamic equilibrium is relevant

as one dye species may be more reactive than the other. Azo

dyes containing a hydroxyl group in the ortho position to the

azo group within naphthyl or higher fused ring systems, can

exist as azo and hydrazone tautomers,27 with the relative

amounts varying with reaction parameters such as solvent

and temperature.28 Furthermore, in aqueous solution these

species are in a pH dependent equilibrium with a common

anion, in which the negative charge is delocalised throughout

the molecule (see Scheme 1).29 These are chemically distinct

forms which have characteristically different visible spectra,

the azo form absorbs typically at 400–440 nm and the hydra-

zone form at 475–510 nm (see Fig. 1).30

The absorption spectrum of Orange II in an aqueous

carbonate solution shows under the selected reaction condi-

tions (Fig. 1) one main band at 480 nm, which correspond to

the n - p* transition of the azo form. The other two bands at

300 and 270 nm are attributed to the p - p* transition of the

benzene and naphthalene rings, respectively.31 Orange II, due

to the presence of aromatic groups, is very stable, and in the

presence of a powerful bleaching agent such as H2O2,

degradation of dye solutions occurs slowly under specific

reaction conditions. Surprisingly, the oxidation rate was tre-

mendously accelerated by addition of a simple manganese

salt. The reactivity of the in situ formed intermediate

was comparable with the catalytic activity of some earlier

postulated, well known manganese bleach catalysts13,32 and

manganese porphyrins.33 In our work, the formation and

stabilization of the active catalyst was studied in a carbonate

buffer solution.

Complex-formation between Orange II and Mn2+

ortho-Hydroxy aromatic azo dyes, which are bidentate com-

plexing agents are of considerable practical and theoretical

interest because of their ability to form stable chelate com-

plexes with some metal ions.34 It is known that Orange II can

act as a chelating agent since the hydroxy and sulfonate groups

allow a stabilized complex to be formed.35 Addition of Orange

II to a freshly prepared aqueous carbonate solution of a MnII

salt results in significant changes in the UV-Vis spectrum of

Orange II as shown in Fig. 2.

UV-Vis spectra recorded before and after mixing (ca. 5 s

delay) of 5 � 10�5 M Orange II with 5 � 10�5 M Mn(NO3)2showed a significant increase in absorbance at 480, 310

and 228 nm, respectively. The differences before and after

mixing are not profound at low Mn2+ concentrations. On

increasing the Mn2+ concentration, a continuous increase in

DAl=480 nm = A(dye+Mn(II)) � Adye was observed, indicating

the formation of an Orange II� � �Mn2+ species according to

eqn (3). It should be noted that at higher Mn2+ concentration,

a precipitate started to form. The value of Keq was determined

through a constant variation of the Mn2+ concentration. For

a correct determination of the complex-formation constant,

independent measurements were performed at constant man-

ganese concentration where the Orange II concentration was

continuously varied (see Fig. 3C). Independent measurements

were repeated between five and eight times. Selected data are

Scheme 1 Orange II: R = SO3Na; pKA = 11.4, lmax = 480 nm.

Fig. 1 UV-Vis spectrum of 10�4 M Orange II in carbonate buffer

solution at pH 8.5.

Fig. 2 (Blue curve) UV-Vis spectrum of a 5 � 10�5 M Orange II

carbonate (0.1 M HCO3�) solution at pH 8.5 before mixing with a

5 � 10�5 M Mn(NO3)2 solution at pH 8.5. (red curve) UV-Vis

spectrum recorded directly after mixing (ca. 5 s delay).


shown in Fig. 3A, where the solid line represents a fit of eqn (4)

to the data.

Orange IIþMn2þ ÐKeq

ðOrange II � � �MnIIÞ ð3Þ

Ax � A0 = (AN � A0)Keq[Orange II]/(1 + Keq[Orange II])

DA = DANKeq[Orange II]/(1 + Keq[Orange II]) (4)

The values of A0 and AN represent the absorbances of

Orange II and of the complex Orange II� � �MnII, respectively,

and Ax is the absorbance at any MnII concentration. The value

of Keq was calculated from eqn (4) to be (2.9 � 0.9) � 104 M�1,

indicating a relatively weak coordination of the dye to the metal

center. Experimentally, through addition of a 4 � 10�5 M

Mn(NO3)2 solution to a 5 � 10�5 M Orange II aqueous

carbonate solution (0.2 M HCO3�), a decrease in the pH of

the solution from 8.5 to 8.3 was observed, which suggests

phenolic proton release due to Mn(II) coordination to Orange

II with the formation of a six-membered ring structure instead

of coordination to the terminal sulfonato group. At higher

concentrations (above ca. 10�3 M) Orange II forms dimers and

higher aggregates in aqueous solutions,30,36 and has a marked

effect on the observed spectra, particularly UV-Vis and NMR.37

A Benesi-Hildebrand treatment of the optical data to determine

Keq could not be applied since the concentration of Orange II

and MnII were close to each other.38 Using Job’s method,39 the

stoichiometry of the formed complex could be determined.

According to the data shown in Fig. 3B and D, at lower MnII

concentration the formation of a complex with a stoichiometry

of 1 : 1 can be assumed. On increasing the Orange II concen-

tration further, complexes with a higher stoichiometry are

possibly formed (see Fig. 4A and B).

Similar structures have been reported earlier by Nadtochenko

and Kiwi when a Fe3+ salt was added to an Orange II solution

in acidic medium.40 Bauer also reported a TiIV complex,

where TiIV is coordinated by two oxygen atoms from the

sulfonato group and the oxygen of the carbonyl group of the

Fig. 3 (A) Change in absorbance at 480 nm on addition of different concentrations of Mn2+ to 5 � 10�5 M Orange II in aqueous carbonate

solution (0.2 M HCO3�) at pH 8.5 and 22 1C. (B) Job plot analysis for complex-formation between Orange II and Mn2+ in aqueous carbonate

solution (0.2 M HCO3�) at pH 8.5. (C) Spectral changes at 480 nm on addition of different concentrations of Orange II to a freshly prepared

5 � 10�5 M Mn(NO3)2 carbonate solution (0.2 M HCO3�) at pH 8.5 and 22 1C. (D) Job plot analysis for the complex-formation in aqueous

carbonate solution (0.2 M HCO3�).

Fig. 4 (A) Proposed structure for a 1 : 1 Orange II� � �Mn2+ complex

formed in a carbonate buffer solution at a low concentration of Mn2+.

(B) Proposed structure for a 1 : 2 Orange II� � �Mn� � �Orange II

complex formed in a carbonate buffer solution at a high concentration

of Orange II.


hydrazone tautomer.41 In the enzyme manganese peroxidase,

the double role of Orange II as a stabilizer, forming a complex

with MnIII, and as a substrate that permits the regeneration of

MnII, was recently postulated by Lopez et al.35 Although, the

coordination of organic dyes, viz. Alizarin, Alizarin S42 and

Orange II,3 to several transition metal centres has been known

for years, comparatively little has appeared on their use as

potential stabilizing ligands in oxidative degradation of

organic dyes.

The formed Orange II� � �MnII complex was isolated and the

validity of its composition was confirmed by elemental ana-

lysis. In control experiments the reactivity of the isolated 1 : 1

Orange II� � �MnII and 2 : 1 complexes were studied. The

isolated complexes exhibit the same catalytic activity and

stability under the experimental conditions employed for the

in situ generation of the complex. Due to the weak coordina-

tion mode of the ligand, no differences between the catalytic

activity of the 1 : 1 and 2 : 1 complex were found.

Beside UV-Vis measurements and DFT calculations,

electrochemical measurements were used to study the in situ

formation of highly reactive MnII catalytic species in

the presence of Orange II under the selected experimental

conditions.

CV studies on the complex-formation between Orange II

and Mn2+

CV measurements of a 4 � 10�5 M Mn2+ solution in the

presence of different Orange II concentrations were performed

in order to determine the interaction between the fully aquated

Mn2+ ions and Orange II present in the reaction mixture.

Fig. 5 shows the results of MnII� � �Orange II complex forma-

tion in NaCl electrolyte solution, performed using a standard

three electrode electrochemical setup as described above.

To avoid the oxidation of MnII to MnIV, which precipitates

as MnO2, the potential scan was discontinued at +1.0 V, after

which the reverse scan from +1.0 to �0.8 V was started. The

CVs of Mnaq2+ in the absence of any coordinating substrate

exhibit one quasi-reversible oxidation peak at E = +0.59 V

vs. Ag/AgCl and one quasi-reversible reduction peak at E =

+0.35 V, corresponding to the one electron Mn3+/Mn2+

redox couple. In addition, CV measurements on a freshly

prepared 4 � 10�5 M Orange II electrolyte solution at

pH 8.5 and 22 1C were performed. Orange II, as it can be seen

in Fig. 5, undergoes two electrochemically quasi-reversible,

one-electron reductions with CV half-wave potentials at

Ered1 = �0.19 V and Ered2 = +0.11 V (vs. Ag/AgCl) with a

difference between the cathodic and anodic wave of 0.02 and

0.204 V, respectively. Furthermore, the reduction potential of

Mn3+ decreased from +0.35 V to +0.28 V when Orange II

was added to the solution, indicating the stabilization of

Mn3+ ions. In the presence of a chelating substrate, the

generated Mn3+ complex becomes more stable and the redox

potentials attain lower values.43 When the concentration of

Orange II was increased up to 2 � 10�5 M, the presence of

further reduction peaks along with changes in the oxidation

peak intensity were observed, indicating the formation of

other manganese–Orange II species as specified above.

DFT calculations

To assess the coordination mode of Orange II to the Mn(II)

center, DFT (B3LYP/LANL2DZp) calculations were per-

formed for a series of plausible complexes. Orange II in

aqueous solution under the selected experimental conditions

dissociates into an anionic sulfonate group and a cationic

sodium ion. In the presence of an unsolvated SO3� group

involving charge transfer from the electron-rich sulfonate

group onto the rest of the molecule, may in general not give

satisfactory DFT results.28 Solvent Yellow 14, a model com-

pound for Orange II containing no sulfonate group was

selected for the DFT study of the interaction between the

Mn(II) and the chosen azo dye. A picture of the calculated

conformers of the model compound 1 is shown in Fig. 6.

The optimized geometry of 1a was calculated to be

ca. 5.8 kcal mol�1 lower in energy than that of 1b. Further-

more, the calculated structure of 1a was compared with X-ray

structural data of Solvent Yellow 14.44 A good agreement

between calculated and crystallographically determined struc-

ture was found.

According to the UV-Vis and electrochemical data pre-

sented above, Orange II can coordinate to a fully aquated

Mn2+ center. Different plausible interaction modes of Solvent

Yellow 14� � �MnII (2) and Solvent Yellow 14� � �MnII� � �SolventYellow 14 (3) were studied in detail. Optimized structures of 2

adopting different coordination modes are presented in Fig. 7.

The studied organic dye can coordinate to aquated Mn2+

ion by forming two new bonds, one between Mn2+ and the

deprotonated phenolic OH-group of 1a and the second

between Mn2+ and one of the azo nitrogen atoms, leading

Fig. 5 CVs of a 4 � 10�5 M Mn2+ electrolyte (0.1 M NaCl) solution in the presence of different Orange II concentrations at pH 8.5 (adjusted by

addition of NaOH) and 22 1C.


to the formation of either a planar six-membered (2a) or five-

membered (2b) chelate complex. Furthermore, a second inter-

action mode for 2 involving a hydrogen bond between one

coordinated water molecule and the azo nitrogen atom (2c)

was taken into consideration. The calculated energies indi-

cate that 2a is energetically favoured over 2b by about

3 kcal mol�1. The N–N bond length of 1.30 A for 2a is nearly

identical to that found in the free model molecule 1a (1.28 A),

indicating a weak interaction between the nitrogen atoms and

the positively charged manganese center.

In addition to these structures, DFT calculations were

performed for a further possible interaction of a second dye

molecule with the Mn(II) center leading to the formation of

chelated Mn(II) inner-sphere complexes. Similar transition

metal complexes of ortho-hydroxy azo dyes were prepared

and characterised by Drew and Landquist.45 The introduction

of a second dye molecule is expected to have certain advan-

tages. In addition to the usual stabilization by the chelate-

effect, the introduction of a second molecule of 1a could result

in a protecting effect on the coordination framework. The

optimized structure of 3 is presented in Fig. 8.

The calculated structure of 3 shows a C2-symmetry and the

axial positions are nearly equivalent. The calculated Mn–N

bond lengths in the equatorial plane for the energetically

favoured 2a (2.15 A) and 3 (2.30 A) are comparable with the

X-ray structural data for Mn(II) complexes with nitrogen

containing ligands such as 1,2-bis(imidazol-1-yl)ethane (bim)

(2.213–2.294 A),46a 2-[N,N-bis(2-pyridylmethyl)amoniumethyl]-6-

[N-(3,5-di-tert-butyl-2-oxidobenzyl)-N-(2-pyridylamino)amino-

methyl]-4-methylphenol (H2Ldtb) (2.118–2.237 A)46b and

1,4,7-triazacyclononane (tacn) (2.118–2.146 A).46c

As expected, upon coordination of two dye molecules in 3,

the N–N bond distance becomes longer (1.29 A) than observed

in the crystal structure of 1a due to the partial neutralization

of the delocalized negative charge of the nitrogen atom.

The elongation of the Mn–O bond trans to the azo group

(Mn–O = 2.38 A vs. 2.06–2.27 A for 2, and Mn–O = 2.27/

2.26 A vs. 1.81/2.11 A for 3) exerts a significant trans influence

Fig. 6 Optimized (B3LYP/LANL2DZp) structures of 1a and 1b with a planar geometry and dihedral angles of (a) 180.01 and (b) 178.71 about the

azo group, C–N–N–C.

Fig. 7 Optimized structures of complex 2a, b and c (B3LYP/LANL2DZp).

Fig. 8 Optimized structure of 3 (B3LYP/LANL2DZp).


opposite to the Mn–N bond. The increased lability of the axial

ligand allows the subsequent interaction of the substituted

transition metal atom with an oxidant, leading to the

rapid formation of active oxidizing species. Moreover, DFT

calculations performed by Blomberg et al.47 suggest that in the

presence of weak-field ligands for Mn(II) and Mn(III), five-

coordination is also accessible whereas Mn(IV) has a much

stronger preference for six-coordination.47

Complex-formation between bicarbonate and Mn(II)

The reactions between bicarbonate ions (HCO3�) and differ-

ent manganese species have been studied for several years,

since aquated Mn2+ cations themselves are actually not able

to catalyze H2O2 disproportionation. Depending on the

HCO3� concentration in the reaction mixture, MnII� � �HCO3

�

complexes of different stoichiometry can be formed. Recently,

it was suggested that only the neutral MnII(HCO3�)2 complex

can facilitate H2O2 disproportionation.48 In this study the

complex-formation reaction between Mn2+ and HCO3� was

monitored using UV-Vis spectrophotometric beside CV mea-

surements as a function of carbonate concentration at pH 8.5.

UV-Vis spectra recorded before and after addition of HCO3�

to an aqueous Mn2+ solution showed the formation of a new

broad band at 300 nm as illustrated in Fig. 9A. The time

course of the absorption band formation is shown in Fig. 9B.

It can be seen from Fig. 9B that the rate of formation of the

manganese carbonate intermediate is enhanced at higher

carbonate concentration. The observed first order rate con-

stants following the induction period in Fig. 9B, are directly

proportional to the [HCO3�] in the range 0.01–0.1 M

(see Fig. 10) with a second order rate constant of (3.6 � 0.2) �10�2 M�1 s�1 at 25 1C. Moreover, the observed induction

period is probably related to the displacement of water from

the first coordination sphere of the fully aquated Mn2+ ion by

HCO3� and subsequent rearrangement of the coordinated

ligand, viz. formation of bidentate carbonate complexes. It

should be noted that under these experimental conditions

(high carbonate concentration and pH 8.5) insoluble MnCO3

is formed as a very fine white precipitate at longer reaction

times. Its composition was confirmed by elemental analysis

and IR spectroscopy.

The reactivity of the produced intermediate was tested in

the oxidative degradation of Orange II by H2O2 at pH 8.5 and

25 1C. During the first 200 s, no change in the reactivity of the

in situ formed manganese intermediate occurs. A significant

time dependent loss in catalytic efficiency of the formed

MnII� � �HCO3� intermediate was observed. An irreversible

deactivation of the catalyst occurs within less than 20 min.

On the other hand, no precipitate formation as well as no

deactivation of the catalytically active manganese intermediate

could be observed in the presence of a coordinating organic

substrate, i.e. Orange II, over a long period of time (1–4 days)

in a high carbonate (0.5 M) containing buffer solution under

these conditions. Moreover, the stabilization of the in situ

formed active catalyst in the presence of an organic substrate

is of considerable practical interest, because its successful

implementation could offer a more efficient alternative for

clean oxidation reactions.

CV measurements of freshly prepared aqueous Mn(NO3)2solutions were performed in the presence of different carbo-

nate concentrations in a 0.1 M NaCl electrolyte solution at

pH 8.5 (adjusted by careful addition of NaOH) and 22 1C. In

the presence of a coordinating substrate, the displacement of a

coordinated water molecule from the manganese coordination

sphere takes place. By coordination of a negatively charged

ligand such as HCO3� to a positively charged metal, the peak

potentials are shifted to more negative potentials compared to

the fully aquated Mn2+ (see Fig. 11A and 12).41 On increasing

the carbonate concentration in solution a decrease in the peak

current intensity occurs concomitantly with peak broadening

because of complexation by carbonate. Typical multiple scan

CVs of a 4 � 10�5 M Mn2+ solution in the presence of 0.2 M

Fig. 9 (A) UV-Vis spectra of an aqueous 4 � 10�4 M Mn2+ solution before (black curve) and after (red curve) addition of 0.4 M HCO3� at pH

8.5 and 25 1C. (B) Time course of the band formation at 300 nm of an aqueous 2� 10�4 MMn2+ solution containing different amounts of HCO3�.

Fig. 10 Plot of observed first order rate constant (kobs) for the

formation of Mn2+� � �HCO3� vs. the bicarbonate concentration in

the presence of 4 � 10�4 M Mn2+ at pH 8.5 and 25 1C.


NaHCO3 and 0.1 M NaCl at pH 8.5 and 25 1C is presented in

Fig. 11B. In the presence of a chelating substrate, the gener-

ated Mn3+ complex becomes more stable and the redox

potentials attain lower values. Moreover, at higher carbonate

concentrations in the reaction mixture the presence of a second

oxidation peak at E = +0.41 V, attributed to the formation

of further complexes such as proposed in eqn (5), was

observed.

½MnIIðH2OÞ6�2þ þHCO �

3 ÐK1

½MnIIðH2OÞ5ðHCO3Þ�þ

½MnIIðH2OÞ5ðHCO3Þ�þþHCO �3 Ð

K2

½MnIIðH2OÞ4ðHCO3Þ2�ð5Þ

By plotting the peak potential E as a function of the

hydrogen carbonate concentration (see Fig. 12), the presence

of different complex species at different carbonate concentra-

tions is revealed.

Carbonate concentration dependence of the catalytic reaction

course

The effect of the carbonate concentration on the oxidative

degradation of the dye was studied at a constant pH of 8.5.

The total carbonate concentration was varied between

0.05 and 0.5 M.

In the present case, the catalytic reaction leads to a square

dependence of kobs on the HCO3� concentration (Fig. 13) with

a third rate constant (8.3 � 0.3) � 10�2 M�2 s�1, suggesting

that 2 equivalents of HCO3� are involved in the oxidation

mechanism. It is suggested, among other possibilities, that one

equivalent of HCO3� is required for the formation of the more

reactive [MnII(H2O)5(HCO3�)]+ intermediate, and the second

equivalent of HCO3� is required for the formation of the more

reactive peroxocarbonate species, known to be a versatile

oxidizing agent. It should also be noticed that no oxidation

of Orange II by H2O2 was observed in the absence of a

carbonate buffer. In the view of these findings we decided to

study the influence of carbonate on the manganese catalyzed

oxidation of Orange II by H2O2 and HCO4�, respectively.

The reaction of carbonate with H2O2 at pH 8.5 and 25 1C

Peroxycarbonate ions, known to be several orders of magni-

tude more reactive toward nucleophilic substrates than H2O2

itself,49 are formed in a relatively fast pre-equilibrium

(K = 0.32 � 0.02 M�1)40 between hydrogen carbonate ions

and H2O2 shown in eqn (6).

HCO3� + H2O2 " HCO4

� + H2O (6)

Moreover, the reaction of H2O2 and HCO3� to form the more

electrophilic HOOCO2� (HCO4

�) occurs rapidly (t1/2 E 300 s)

at near neutral pH and 25 1C.50 This step is also regarded to be

a key aspect of several oxidation reactions.51,52 The higher

Fig. 11 (A) Cyclovoltammograms for 4 � 10�5 MMn2+ solution in an aqueous solution of 0.1 M NaCl and different concentrations of HCO3�.

(B) Typical multiple scan CVs of a 4 � 10�5 M Mn2+ solution in the presence of 0.2 M NaHCO3 and 0.1 M NaCl at pH 8.5 and 22 1C.

Fig. 12 Plot of peak potential E as function of [HCO3�]; E vs.

Ag/AgCl electrode, [Mn2+] = 4 � 10�5 M, [HCO3�] = 0.1–05 M

in 0.1 M NaCl electrolyte solution at pH 8.5 and 22 1C.

Fig. 13 Second-order carbonate concentration dependence of kobs.

Experimental conditions: 2� 10�5 MMn(NO3)2, 5� 10�5 M Orange II,

0.01 M H2O2, pH 8.5, 25 1C.


reactivity of peroxycarbonate compared to that of H2O2 is

attributed to carbonate being a better leaving group than

hydroxide.40

By performing the oxidation reactions in the presence of

peroxycarbonate instead of H2O2 in a 0.5 M carbonate con-

taining buffer solution at pH 8.5, no difference in the reactivity

was observed (Fig. 14A).

The Mn(II) catalyzed oxidative degradation of Orange II by

using HCO4� as an oxidizing agent could be significantly

enhanced through increasing the total carbonate concentra-

tion in the reaction mixture. This can be explained in terms of

the equilibrium formulated in eqn (6). Based on our experi-

mental observations and aspects reported in the literature41 for

the Mn(II) catalyzed oxidation reaction by H2O2 in a carbo-

nate containing solution, the reaction sequence presented in

Scheme 2 can be suggested to occur.

Complex-formation between Mn2+

and H2O2 in an aqueous

carbonate solution

Addition of H2O2 to hexaaqua Mn2+ in a carbonate solution

leads to significant spectral changes in the UV-Vis spectra

during the reaction (see Fig. 15A). The initial rapid increase

of the intensity of the broad band at 300 nm, as it is illustra-

ted in Fig. 15A, is attributed to the fast formation of

[MnII(H2O)5(HCO3)]+. An isosbestic point at 330 nm suggests

the formation of a new manganese intermediate by addition of

an oxidizing agent, i.e. H2O2. According to our spectroscopic

observations the formed complex with an absorption band at

400 nm could be attributed to a MnIV–Z2-peroxycarbonate

intermediate.53 Based on spectroscopic observations and data

reported in the literature,54 the formed intermediate can be

regarded as most likely to be a high valent manganese com-

plex. Similar Rh,55 Pt56 and Fe57 peroxycarbonate complexes

have been isolated before and were characterized spectro-

scopically. The time course of the absorption band at 400 nm

at different pH is illustrated in Fig. 15B.

In the absence of any stabilizing ligand, the formed complex

rapidly decomposes with the formation of catalytically

inactive MnIVO2 that precipitates from solution (see Fig. 15B).

The decomposition of the active intermediate is accelerated at

higher pH (see Fig. 15B). To ascertain that the formulated

reaction steps in Scheme 2 are valid under our reaction

conditions, a systematic spectroscopic investigation at differ-

ent pH values was performed. Representative data for the

reaction course at 400 nm at pH 8.5 and 9.5 are presented in

Fig. 15B. Contrary to our expectations, an increase of one unit

in pH resulted in an increase of the induction period and a

decrease in the manganese peroxycarbonate complex forma-

tion rate under the mentioned reaction conditions. This could

be partly due to subsequent formation of Mn(OH)2 precipi-

tates at higher pH and to deprotonation of HCO3� that

becomes significant at pH above 8 to 9.51 This results in a

decrease in the HCO3� concentration in the equilibrium

presented in eqn (6), reducing the concentration of peroxy-

carbonate present in solution.

Reactivity profile as function of pH

The reactivity of the catalytic system is generally influenced by

the protonation state of the substrate, the catalyst and oxidiz-

ing agent. In our work the kinetics was studied in 0.4 M

HCO3� containing buffer solution in the pH range between

8.0 and 9.5 at 25 1C. The pH of the carbonate buffer solution

was adjusted carefully using small amounts of concentrated

NaOH solution to avoid dilution. A typical manganese cata-

lyzed oxidative degradation of Orange II by H2O2 in a

carbonate buffer solution is presented in Fig. 16A. The

catalytic degradation is usually complete within 1–10 min

Fig. 14 (A) Spectral changes observed at 480 nm for the 2 � 10�5 MMn(NO3)2 catalyzed oxidative degradation of 2.5 � 10�5 M Orange II in the

presence of (black curve) 0.01 M H2O2 and (red curve) 0.01 M HCO4�, respectively, at pH 8.5 and 0.5 M total carbonate concentration.

(B). Comparison of the absorbance changes at 480 nm vs. time for the 2 � 10�5 MMn2+ catalyzed oxidative degradation of 5 � 10�5 M Orange II

by 0.01 M H2O2 at pH 8.5 and different carbonate concentrations.

Scheme 2 In situ formation of catalytically active Mn intermediates

in the presence of hydrogen peroxide in a carbonate containing

aqueous solution at pH 8.5 and 25 1C.


depending on the pH of the solution, the catalyst concentra-

tion, and the H2O2 and carbonate concentrations. The decom-

position of the dye was followed by monitoring the spectral

changes at 480 nm. The depletion of the band at 480 nm is in

general correlated with cleavage (heterolytic or homolytic) of

the azo group leading to colorless oxidation products due to

the induced discontinuity in the conjugation of the p-system in

the dye molecule. The inset in Fig. 16A shows the first

spectrum of Orange II before the addition of the catalyst

and H2O2, and the final spectrum recorded after 250 s.

A decrease in the intensity of the two other bands at

270 and 300 nm was observed, showing that further bleaching

also occurs under these reaction conditions. The isolation and

characterization of reaction products is extremely difficult and

requires large synthetic efforts, particularly as different reac-

tion intermediates tend to react further under experimental

conditions. A comparison of the reaction course at different

pH values is shown in Fig. 16B.

The Mn(II) catalyzed decolorization and oxidative decom-

position of Orange II was found to be sensitive to the pH of

the solution. According to our experimental data, an increase

in pH resulted in a slight decrease in the reaction rate under

the above-mentioned reaction conditions and the highest

reactivity is observed at a pH between 8.2 and 8.6

(see Fig. 17). Increasing the pH to 49 leads to a decrease in

the oxidation rate for the bicarbonate-activated peroxide,

which is presumably the result of the deprotonation of

HOOCO2� to form CO4

2�, a less electrophilic oxidant.58

At even higher pH, the decomposition of the peroxide is

accelerated and may reduce the oxidation reaction rate.

Contrary to our expectations, the observed rate constants

for the decolorization reaction of Orange II are similar to

the destruction rate constants of naphthalene and benzene

rings, long-lived intermediates, under the studied conditions

(see Fig. 17). Thus, for a complete oxidation of these stable

molecules higher concentrations of oxidant and catalyst are

required.

A similar screening using MnCl2, Mn(Ac)2 and Mn(SO4)2showed identical catalytic activity in the oxidative degradation

of Orange II by H2O2. In all cases, the manganese catalyzed

oxidative degradation of Orange II is favored by moderate

alkaline pH values and vanishes completely at very high or

very low (strong acidic) values. According to the experimental

observations mentioned above, the manganese catalyzed

oxidative degradation of Orange II by H2O2 in a carbonate

containing solution is considerably inhibited at higher

pH values due to the lower formation of the high valent

manganese Z2-peroxycarbonate complex (see Fig. 15B).

Fig. 15 (A) UV-Vis spectra recorded for the reaction of 2 � 10�4 M Mn(NO3)2 with 0.01 M H2O2 in a 0.5 M HCO3� containing solution at

pH 8.4 and 25 1C. (B) Comparison of typical absorbance at 400 nm vs. time plots at pH 8.5 (black curve) and 9.5 (red curve).

Fig. 16 (A) UV-Vis spectra of a 2 � 10�5 MMn(NO3)2 catalyzed oxidative degradation of 5� 10�5 M Orange II by 0.01 MH2O2 in a 0.4 M total

carbonate containing solution at pH 8.5 and 25 1C. The inset in Fig. 16A shows the first spectrum of Orange II before the addition of the catalyst

and H2O2, and the final spectrum recorded after 250 s. (B) Comparison of absorbance at 480 nm vs. time plots for the 2 � 10�5 M Mn(NO3)2catalyzed oxidative degradation of 5 � 10�5 M Orange II by 0.01 M H2O2 in a 0.4 M total carbonate containing solution at different pH values

and 25 1C.


Effect of the manganese concentration on the oxidative reaction

course

To evaluate the effect of the catalyst concentration on the

manganese catalyzed oxidative degradation of Orange II by

H2O2 under catalytically relevant experimental conditions,

kinetic studies were performed for solutions in which the

carbonate containing water solution with various amounts

of Mn(NO3)2 was added in the presence of 0.01 M H2O2 to a

5 � 10�5 M Orange II solution at 25 1C. The obvious

accelerating ability of the HCO3� ions prompted us to study

the catalytic reaction course in more detail at four different

carbonate concentrations. The in situ produced catalyst con-

centration dependence was studied at 480 nm using in situ

UV-Vis spectroscopic measurements and the kinetic traces

could be adequately fitted to a single exponential function.

Plots of the observed rate constant as a function of [Mn2+] at

different carbonate concentrations are presented in Fig. 18.

As it is evidenced in Fig. 18, the [Mn2+] dependences of the

observed rate constants for the manganese catalyzed oxidative

degradation of Orange II by H2O2 in a low carbonate con-

centration containing solution (0.1–0.3 M HCO3�) are

strongly curved (higher K values, see Table 1) and reach a

limiting value at higher catalyst concentration. In contrast,

similar data at higher carbonate concentrations (0.4–0.5 M

HCO3�) result in a less curved dependence of kobs on the

catalyst concentration, i.e. lower K values (see Table 1). The

observed rate profile can be explained by the general reaction

mechanism proposed in Scheme 2 and simplified in Scheme 3.

The observed rate law for the proposed reaction steps in

Scheme 3 is given by eqn (7). The calculated k and K values

from the non-linear concentration dependences in Fig. 18 are

summarized in Table 1.

kobs ¼kK ½MnðIIÞ�

1þ K ½MnðIIÞ� ð7Þ

Effect of the H2O2 concentration on the manganese-catalyzed

oxidative degradation of Orange II

The effect of H2O2 on the oxidation reaction course was

studied by varying its initial concentration over a wide range,

between 5 and 30 mM (Fig. 19). At lower H2O2 concentrations

(1 and 5 mM) a fast oxidation reaction occurs in the first few

seconds followed by a rapid consumption of H2O2 resulting

finally in a partial and inefficient decolorization of the dye.

This prompted us to study the H2O2 concentration effect on

the catalytic oxidation of the dye at higher concentrations of

H2O2. The kobs values were calculated from a single exponen-

tial fit to the absorbance at 480 nm vs. time plots and showed a

linear dependence on the initial H2O2 concentration over the

studied concentration range.

Stability of the in situ formed catalyst

In control experiments the stability of the in situ generated

catalyst was studied by repeated addition of dye and H2O2 to a

solution of 2 � 10�5 M Mn(NO3)2 at pH 8.5 (0.4 M HCO3�)

and 25 1C (see Fig. 20A and B).

As it can be seen in Fig. 20A, the catalytic cycle could be

repeated several times without any significant loss of activity

during the oxidation reaction, indicating an excellent stability

of the in situ formed catalyst. After the fifth cycle the reactionFig. 18 Mn(NO3)2 concentration dependence of kobs. Reaction con-

ditions: 5 � 10�5 M Orange II, 0.01 M H2O2, pH 8.5 and 25 1C.

Table 1 The constants k and K for theMn(NO3)2 catalyzed oxidationof Orange II by H2O2 at pH 8.5 and 25 1C (see Scheme 3)

[HCO3�]/M k/s�1 10�3K/M�1

0.1 0.0033 34.60.3 0.032 17.60.4 0.051 17.80.5 0.138 15.2

Scheme 3 Proposed reactions steps for the formation of the cata-

lytically active manganese intermediate in the presence of H2O2 in a

carbonate containing solution.Fig. 17 Plot of observed rate constant (kobs) calculated for the

decoloring reaction followed at 480 and 300 nm, respectively. Experi-

mental conditions: 2 � 10�5 M Mn(NO3)2, 5 � 10�5 M Orange II,

0.01 M H2O2, 0.4 M total carbonate and 25 1C.


solution containing the active catalyst was allowed to stay at

ambient temperature for 48 h. Subsequently, the catalytic

activity of the in situ formed manganese complex was evalu-

ated again by performing the oxidation reaction in the

presence of freshly added Orange II and H2O2. The experi-

mental results illustrated in Fig. 20B provide clear evidence for

the high efficiency of the in situ formed catalyst under the

above mentioned experimental reaction conditions.

Mechanistic aspects of the manganese-catalyzed oxidative

degradation of Orange II by H2O2 in carbonate solution

Throughout this study, the oxidation reactions were carried

out in a thermostated open glass reactor vessel at ambient

temperature in aqueous hydrogen carbonate containing solu-

tions. The readily available manganese salts, the mild reaction

conditions and the operation simplicity and practicability

allow for an easy and green oxidative degradation of the

studied organic dye. In control experiments the catalytic

activity of the in situ generated manganese complex was

investigated under an inert atmosphere. By performing the

catalytic reaction in a closed glass reactor under inert reaction

conditions no change in the decomposition reaction rate was

noticed. A comparison of the reaction course carried out

under different experimental conditions is illustrated in

Fig. 21.

By performing the reaction under inert reaction conditions

no significant differences in the decomposition reaction rate was

observed, indicating that HO� or HOO� radical formation

is not prevalent for this oxidation reaction. This is further

supported by the observation that addition of radical traps

such as TTBP had no effect on the reaction course (see Fig. 21).

Taking into account all obtained spectroscopic and kinetic

data, the following reaction schemes can be proposed for the

Mn2+ catalyzed oxidative degradation of Orange II by H2O2

in carbonate solution under catalytically relevant experimental

conditions.

A key feature of the proposed reaction mechanism outlined

in Scheme 4 is that the overall oxidation of Orange II occurs in

a two electron oxidation step leading to the formation of a

relatively stable high-valent MnQO intermediate and transfer

of the oxo group to the substrate. Most of the earlier reported

papers22,59 on the oxidation reaction catalyzed by several

isolated and structurally well defined manganese complexes

have emphasized the formation of a high-valent MnQO

intermediate by the reaction of manganese with the appro-

priate oxidant. According to our observations, HCO3� ions

are involved in two catalytically relevant reactions. HCO3�

ions react with aquated MnII present in solution to form a

catalytically active Mn–HCO3� complex. HCO3

� is also in-

volved in a fast equilibrium with H2O2 to form HOOCO2�,

a versatile heterolytic oxidant. In the following step, through

nucleophilic attack of the oxidizing agent on the MnII center, a

MnII–Z2-peroxycarbonate complex is formed. The remaining

coordination sites in the first shell will be occupied by water

and hydroxyl at a pH between 8 and 10. The principal mode of

the formation of relatively stable high-valent MnQO inter-

mediates is believed to involve the heterolytic cleavage of the

peroxide bond, as shown in Scheme 4. An important role in

the stabilization of the formed MnQO species is played by the

electron donating bicarbonate ions. This may also account for

the unique requirement of HCO3� in the oxidative decom-

position of Orange II catalyzed by simple manganese salts.

The further coordination of the substrate followed by an

oxygen transfer step along with the second electron, leads to

the formation of several oxidation products and finally to the

regeneration of the catalyst. It must be noted that in the

Fig. 19 H2O2 concentration dependence of kobs. Reaction conditions:

5 � 10�5 M Orange II, 2 � 10�5 M Mn(NO3)2, pH 8.5 and 25 1C.

Fig. 20 (A) Spectral changes observed at 480 nm for the repeated addition of 5 � 10�5 M Orange II to a 2 � 10�5 M Mn(NO3)2 solution in the

presence of 0.01 M H2O2 at pH 8.5 and 0.4 M total carbonate concentration. (B) Spectral changes observed at 480 nm for a new addition of

5 � 10�5 M Orange II and 0.01 M H2O2 to a 48 h old reaction mixture containing the catalyst solution under the same experimental conditions as

mentioned in A.


absence of a catalyst, the oxidative degradation of Orange II

by addition of an electrophilic bleaching agent, HOOCO2�,

occurs very slowly under certain reaction conditions. The

oxidation mechanism involves nucleophilic attack of the dye

at the electrophilic oxygen of HOOCO2�. In aqueous solution,

proton transfer can lead to the displacement of HCO3� and

the slow formation of oxidized substrate.

If substrate binding to MnII occurs before the addition of

HOOCO2� to the catalyst solution, following reactions can be

assumed to take place during the reaction cycle under the

chosen experimental conditions.

In line with the concerns mentioned above, the first step in

Scheme 5 involves the prior coordination of Orange II to MnII

and formation of MnII–Orange II complexes of different

stoichiometry, followed by nucleophilic attack of the oxidant

on the MnII center leading to the formation of Orange

II–MnII–peroxycarbonate species. The subsequent scission of

the peroxo bond leads to the formation of high-valent oxo

intermediates, as formulated in Scheme 4. In this case, the

formed MnIVQO intermediate is stabilized by Orange II, an

electron rich organic molecule with chelating capacity. The

importance of Orange II as an equatorial ligand is also to

favor the heterolytic scission of the peroxo bond leading to the

MnIVQO intermediate and bicarbonate.

Conclusions

A fast and environmentally benign method for the oxidative

degradation of Orange II could be achieved using H2O2 in

conjunction with catalytic amounts of relatively non-toxic

manganese salts as catalyst precursors in a carbonate contain-

ing aqueous solution under mild reaction conditions. Screen-

ing and spectroscopic methods allowed us to study the

catalytic reaction course and to identify some key features of

the reaction that reflect upon its mechanism. Our study

revealed that the oxidative degradation of the model substrate

Orange II is catalytic only in carbonate containing aqueous

solution. No other buffer containing aqueous solution could

induce the oxidative degradation of Orange II by H2O2 and

this led to the implication of peroxycarbonate as a key

molecular entity. The reported experimental data suggests that

the in situ formed high-valent manganese intermediate posses-

sing one hydrogen carbonate ligand is able to activate

H2O2, but decomposes rapidly with the formation of neutral

MnCO3, which precipitates from solution as an insoluble

white solid. One of the main factors affecting the process

efficiency was the stabilization of the catalytically active Mn

complex. Furthermore, by addition of Orange II, the forma-

tion of MnII� � �Orange II complexes with different stoichio-

metry was observed. The simultaneous s,p-coordination of

the organic dye is well-precedented, and recent DFT studies

support this type of complex formation.28 The catalytic

Scheme 4 Proposed reaction mechanism for the Mn(II) catalyzed

oxidative degradation of Orange II by H2O2 in a carbonate containing

aqueous solution at pH between 8–9 and 25 1C.

Scheme 5 Proposed reaction mechanism involving first substrate

coordination to MnII in a pre-equilibrium step during the catalyzed

oxidative degradation of Orange II by H2O2 in a carbonate containing

aqueous solution at pH between 8–9 and 25 1C.

Fig. 21 Comparison of typical absorbance at 480 nm vs. time plots of

a 2� 10�5 MMn(NO3)2 catalyzed oxidative degradation of 5� 10�5 M

Orange II by 0.01 M H2O2 in a 0.4 M HCO3� containing solution at

pH 8.5 and ambient temperature performed in the presence of atmos-

pheric oxygen (black curve), inert atmosphere (red curve) and TTBP

(blue curve), respectively.


activity of the formed intermediates was tested under catalytic

reaction conditions.

The kinetic investigations performed at different pH could

provide relevant information about the nature of the oxidizing

agent involved in the reaction. It was found that the pH is a

critical issue for the rate of the oxidation process due to its

influence on the deprotonation of the bicarbonate ions, the

formation of peroxycarbonate in solution, and the deprotona-

tion of aquated Mn2+. The ongoing studies are presently

complemented by investigations on different organic sub-

strates with various functional groups in order to determine

the influence of substrate modification on the catalytic

reaction cycle. DFT studies beside further kinetic and spectro-

scopic investigations should contribute to a better understanding

of the catalytic system.

Acknowledgements

The authors kindly acknowledge fruitful discussions with

Dr Anette Nordskog and Dr Wolfgang von Rybinski, Henkel

KGaA, Dusseldorf, Germany.

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ISSN 1144-0546

Volume 33 | N

umber 1 | 2009 N

JC Pages 1–212

1144-0546(2009)33:1;1-2









Num

ber 1|

2008M

etallomics

Pages1–100 1754-5692(2008)1:1;1-6


ISSN 1756-5901


1756-5901(2009) 1:1;l-m

060877


The solubility and recrystallization of 1,3,5-triamino-2,4,6-trinitrobenzene

in a 3-ethyl-1-methylimidazolium acetate–DMSO co-solvent system

T. Yong-Jin Han,* Philip F. Pagoria, Alexander E. Gash, Amitesh Maiti,

Christine A. Orme, Alexander R. Mitchell and Laurence E. Fried

Received (in Gainesville, FL, USA) 17th June 2008, Accepted 6th August 2008

First published as an Advance Article on the web 18th September 2008

DOI: 10.1039/b810109d

Ionic liquids have previously been shown to dissolve strong inter- and intramolecular hydrogen-

bonded solids, including natural fibers. Much of this solubility is attributed to the anions in ionic

liquids, which can disrupt hydrogen bonding. We have studied the solubility and recrystallization

of 1,3,5-triamino-2,4,6-trinitrobenzene (TATB), a very strong inter- and intramolecular hydrogen-

bonded solid, in various ionic liquid solvent systems. We discovered that acetate-based ionic

liquids were the best solvents for dissolving TATB, while other anions, such as Cl�, HSO4� and

NO3� showed moderate improvements in the solubility compared to conventional organic

solvents. Ionic liquid–DMSO co-solvent systems were also investigated for dissolving and

recrystallizing TATB.

1. Introduction

Ionic liquids (ILs) have recently been shown to be ideal

solvents for dissolving hydrogen-bonded solids, including

cellulose1 and natural fibers.2,3 The use of imidazolium-based

cations with halides as counter-anions has significantly im-

proved the solubilities of these natural products. Successful

dissolution of these highly hydrogen-bonded solids is largely

attributed to the ability of ILs’ anions to act as hydrogen bond

acceptors and disrupt the hydrogen bonds in these materials.1

With a graphite-like crystalline structure,4 1,3,5-triamino-

2,4,6-trinitrobenzene (TATB) is one of the most strongly

hydrogen-bonded solids known. Owing to its inter- and intra-

molecular hydrogen bonds, both in-plane and out-of-plane

(see Fig. 1), the solubility of TATB in conventional organic

solvents is minuscule. With a capacity to dissolve 70 ppm

(0.007% w/v) at room temperature,5 DMSO is the best con-

ventional organic solvent known to dissolve TATB. Super-

acids, such as concentrated sulfuric acid, have been shown to

dissolve up to ca. 240 000 ppm (24% w/v) at room tempera-

ture,5 but due to their highly corrosive nature are often

avoided. There is a need to find a desirable solvent for TATB

as it has become necessary to control the particle size, as well

as the morphology, of TATB crystals.

TATB is of particular interest in the energetic materials

(EM) community due to its extreme insensitivity to impact,

shock and heat, while providing a good detonation velocity.6

This combination of insensitivity with good performance

characteristics makes TATB an ideal insensitive high explosive

(IHE) in numerous applications. TATB has also attracted

researchers from the field of optics, due to its unexpectedly

strong secondary harmonic generation (SHG) efficiency.7 The

source of the nonlinear optical (NLO) property of TATB has

been a topic of intense discussion,8,9 since the original crystal

structure determined by Cady and Larson showned4 TATB to

have a centrosymmetric structure with the space group P�1,

Z = 2, which is incompatible with NLO activity. Some have

attributed the NLO activity of TATB to a small amount of a

second, presently unidentified, TATB polymorph mixed in

with the bulk centrosymmetric TATB crystals.10,11 Therefore,

identifying a suitable solvent system that can increase the

Fig. 1 The crystal structure of TATB, a centrosymmetric structure

with space group P�1, Z = 2. (a) A–B plane view, (b) C plane view.

Chemistry, Materials, Earth and Life Sciences Directorate, LawrenceLivermore National Laboratory, Livermore, CA 94551, USA.E-mail: [email protected]



solubility of TATB may also allow a new opportunity to

crystallize and isolate the potential second polymorph with

exceptional SHG efficiency.

Herein, we report the solubility of highly hydrogen-bonded

TATB in both commercially available and custom synthesized

imidazolium-based ILs. In particular, 3-ethyl-1-methylimida-

zolium acetate (EMImOAc) was extensively investigated in

its pure form, as well as in combination with DMSO as a

co-solvent.

2. Experimental

2.1 Materials

3-Ethyl-1-methylimidazolium chloride (EMImCl), 3-butyl-1-

methylimidazolium chloride (BMImCl), 3-ethyl-1-methyl-

imidazolium acetate (EMImOAc), 3-ethyl-1-methylimidazolium

nitrate (EMImNO3), 3-butyl-1-methylimidazolium hydrogen

sulfate (BMImHSO4) and DMSO were purchased from

Sigma-Aldrich, and used without purification unless otherwise

noted. Vacuum distillation was performed on EMImOAc to

remove any impurities prior to experiments.

3-Allyl-1-methylimidazolium chloride (AllylMImCl) was

synthesized according to a published report.12 3-(Methoxy-

methyl)-1-methylimidazolium chloride (MeOMImCl) was

synthesized using a modification of the procedure reported

for the synthesis of AllylMImCl.

Synthesis of 3-(methoxymethyl)-1-methylimidazolium chloride

(MeOMImCl). Into a 500 mL round-bottomed flask equipped

with a stirrer bar, argon inlet and addition funnel, was

dissolved 1-methylimidazole (25 g, 0.31 mol) in trichloro-

ethylene (100 mL). With stirring, chloromethyl methyl ether

(35 g, 0.43 mol) was added dropwise over a 0.5 h period. The

mixture was warmed and a turbid, two-layer mixture formed.

The mixture was refluxed for 2 h, cooled and poured into a

separating funnel. The organic layer was separated, filtered and

the solvent removed under vacuum at 45 1C to yield a

tan-beige viscous liquid (52 g).

2.2 Solubility measurements

Small scale solubility tests (o10 mg) of TATB in ILs

were monitored with a Nikon optical microscope equipped

with a temperature controlled heating stage, under cross-

polarized light.

Large scale solubility measurements were performed using a

three-necked round-bottomed flask in a silicone oil bath

at a constant temperature of 100 1C. Due to its high density

(1.93 g cm�3) and bright yellow color, visual inspection of

TATB particles in solutions was easily achieved with the aid of

a hand-held flashlight.

2.3 Crystallization

A non-agitated cooling crystallization method was employed

to grow TATB crystals from a DMSO–EMImOAc co-solvent

system. Typically, in a 250 mL round-bottomed flask equipped

with a drying tube and a thermocouple, TATB (4 g) was

added, along with 100 g of DMSO–EMImOAc (80 : 20 w/w).

The solution was slowly heated to 90 1C with constant stirring.

Once all of the TATB had dissolved, the solution was slowly

cooled back down to room temperature without stirring.

Occasionally, an Omega (series 2010) programmable controller

was used to control the cooling rate.

For a typical anti-solvent crystallization, 20 mL of an 80 : 20

DMSO–EMImOAc solution was placed in a four-necked

100 mL round-bottomed flask equipped with an overhead

stirrer, drying tube, thermocouple and septum inlet. To this

was added TATB (0.5 g), and the mixture was stirred and

heated slightly (50 1C) until all of the TATB had dissolved and

a red-orange solution had formed. The temperature of the

sample was maintained at the desired temperature using a

J-KEM temperature controller. The mixture was stirred slowly

as a solution of acetic acid (4 g) in dry DMSO (40 mL) was

added via a syringe and long needle connected to a syringe

pump, set to deliver at 2 mL h�1. The resulting TATB was

collected by suction filtration, and washed with water (25 mL)

and MeOH (10 mL) to yield 0.46 g of a yellow microcrystalline

solid. Raman spectroscopy was used to determine the purity of

the recrystallized TATB.

3. Results

3.1 Solubility of TATB in ILs

There are certain advantages that ILs have over conventional

solvents that make them an attractive alternative for the

dissolution of TATB. ILs, because of their inherent low vapor

pressure and high-temperature stability, have reduced environ-

mental and safety concerns compared to conventional

organic solvents. Also, in theory, the IL is recoverable after

precipitation or distillation of impurities from it.

The solubility of TATB was first investigated in BMImCl,

since BMImCl has previously been shown to dissolve

cellulose,1 Bombyx mori silk fibers2 and wool keratin fibers3

in relatively high concentrations (10, 13.2 and 4 wt%, respec-

tively at 100 1C). A solution of 0.5 wt% of TATB in BMImCl

was stirred rapidly at 100 1C for 20 h. However, at the end of

the 20th hour, there were still TATB particles present in the

flask. The color of the solution was only slightly yellow

(the color of the original TATB powder), signifying that only

a small amount of TATB had dissolved. Similar results were

observed for other ILs with Cl� anions, including EMImCl

and custom synthesized AllylMImCl (see Table 1). As noted

previously, short chain-substituted imidazolium-based ILs

with Cl� anions have been effective in dissolving natural

polymers. The hydrogen bond-accepting Cl� anion is thought

to be the crucial component in disrupting hydrogen bonding in

the biopolymer.1 However, for TATB the hydrogen bond

disruption caused by the Cl� ions was not strong enough to

significantly dissolve TATB particles.

In an attempt to improve the solubility of TATB, a new

imidazolium-based cation, 3-methoxymethyl-1-methylimida-

zolium chloride (MeOMImCl), was synthesized. Unlike the

BMIm and EMIm cations, which may have limited, if any,

hydrogen bond-accepting capability, the ether side chain of

MeOMImCl is a hydrogen bond acceptor13 and may assist in

disrupting the strong hydrogen bonding of TATB. However,

when 0.5 wt% of TATB was added to MeOMImCl and heated

with stirring at 100 1C for 20 h, no significant quantity of


TATB dissolved, similar to BMImCl and EMImCl, which was

visually evident. This result suggests that even with an active

cation, the overall hydrogen bond-accepting capacity of

MeOMImCl is not significant enough to disturb the hydrogen

bonding network of TATB crystals.

The effects of anions other than Cl� in dissolving TATB

were also examined. For this study, BMImHSO4, EMImNO3,

and EMImOAc were selected and tested. These ILs were

chosen for their anions’ hydrogen bond-accepting capability.

As anticipated, BMImHSO4 and EMImNO3 had very similar

results compared to BMImCl, each showing less than 0.5 wt%

solubility of TATB at 100 1C. However, EMImOAc showed a

surprisingly good solubility of TATB. At 100 1C in a large

scale test, EMImOAc was able to dissolve up to 10 � 1.0 wt%

of TATB, confirmed by visual inspection. There was, however,

a noticeable difference in the color of the acetate solution when

compared to the solutions from the other anions (Fig. 2A

inset). When TATB was added to EMImOAc and heated, the

entire mixture turned a dark blood-red color, whereas in the

other ILs, no color change was observed (at times, some

turned slightly yellow, the original color of the TATB crys-

tals). The source of this color change was investigated by

UV-vis spectrophotometry. As seen in Fig. 2A, TATB added

to EMImOAc (diluted 100 times with DMSO) clearly shows a

pronounced peak at lmax = 409 nm, whereas TATB dissolved

in DMSO only shows a TATB absorbance at lmax = 357 nm.

The origin of the peak at lmax = 409 nm can be assigned to a

s-complex. We also carried out a Raman spectroscopy measure-

ment of the s-complex (Fig. 2B). The peaks observed at

807 and 1149 cm�1 correspond well to a previously reported

signature of a s-complex.14 The formation of a s-complex

with TATB was described during the synthesis of TATB by the

vicarious nucleophilic substitution of hydrogen.15 Selig also

assigned the absorption band at 409 nm to a s-complex

between strong bases and TATB.5

3.2 Crystallization of TATB

The primary need for a solvent system that will readily dissolve

TATB is to improve overall processability. More specifically,

high solubility is necessary to produce high quality crystals

from a supersaturated solution of TATB. Therefore, crystal-

lization experiments were performed using EMImOAc via a

non-agitated cooling method. An optical microscope fitted

with a heating stage was used to study the recrystallization of

TATB in EMImOAc. A few drops of TATB particles in

EMImOAc (4 wt% solution) were placed on a cover slip on

a heating stage. The mixture was heated to 100 1C and kept at

this temperature until all of the particles had dissolved (Fig. 3a

and b). The homogeneous solution was cooled slowly by

natural convection. From the saturated solution, crystals

started to emerge when the temperature reached 70 1C. Over

time, single, well-faceted crystals of TATB appeared and grew

larger (Fig. 3c and d). These crystals were far better than the

starting TATB crystals, which were almost all aggregates with

few, if any, well defined facets.

Table 1 The solubility of TATB in various IL systems at 100 1C

IL R X� TATB solubility (wt%)

BMImCl CH3CH2CH2CH2 Cl� o0.5EMImCl CH3CH2 Cl� o0.5AllylMImCl CH2QCH Cl� o0.5MeOMImCl CH3OCH2 Cl� o0.5BMImHSO4 CH3CH2CH2CH2 HSO4

� o0.5EMImNO3 CH3CH2 NO3

� o0.5EMImOAc CH3CH2 CH3COO� 10

Fig. 2 (A) UV-vis spectra of TATB dissolved in DMSO and the

DMSO–EMImOAc system. The inset shows a photograph of TATB

dissolved in DMSO (left) and in EMImOAc (right). (B) Raman

spectra of (a) TATB, (b) EMImOAc in DMSO, (c) TATB dissolved

in EMImOAc–DMSO solution; the peaks at 807 and 1149 cm�1 are

signatures of a s-complex.


Similar recrystallization experiments were performed on a

larger scale (100 mL volume). However, due to the viscosity of

EMImOAc at room temperature, it was difficult to filter and

isolate the recrystallized TATB. In order to lower the viscosity

of EMImOAc, DMSO was added as a co-solvent. When

DMSO was added, the solubility of TATB decreased linearly

with respect to the DMSO concentration (Fig. 4).

However, even with the DMSO concentration as high as

80 wt%, the solubility of TATB was still significantly high,

approximately 4 wt%. Recrystallization experiments were

performed using an 80 : 20 DMSO–EMImOAc solution

(80 wt% DMSO, 20 wt% EMImOAc). Upon heating this

solution with TATB to 90 1C, a color change was once again

observed, signifying the formation of a s-complex. Once the

added amount of TATB had fully dissolved, the solution was

allowed to cool to room temperature by natural convection,

allowing crystals to form. The recrystallized TATB was re-

covered by simple vacuum filtration. It was visually apparent

that the filtrate contained a significant amount of the s-complex.

Therefore, the addition of excess water or another proton

donor (i.e. acetic acid) was necessary in order to fully recover

the remaining TATB dissolved (ca. 1 wt%) in the filtrate. SEM

micrographs of the recrystallized TATB crystals showed good

crystal morphology (Fig. 5a) compared to the starting crystals

(Fig. 5b). The crystal sizes of the recrystallized TATB ranged

from 10–50 mm. On the other hand, the water crash-precipi-

tated crystals showed an irregular crystal morphology, with

crystal sizes that ranged from sub-500 nm–5 mm (Fig. 5c). The

Raman spectrum of the recrystallized TATB crystals con-

firmed that the structure of the recrystallized TATB matched

well with that of the starting material (Fig. 6).

Besides the non-agitated cooling method of TATB crystal-

lization, we also investigated an anti-solvent crystallization

method. As seen from the non-agitated crystallization method

above, a s-complex, which forms when TATB is dissolved in

EMImOAc, requires a proton source to fully convert it back to

TATB at room temperature. Thus, we employed acetic acid as

an anti-solvent to provide the necessary proton for the reac-

tion. Acetic acid, a weak acid, was chosen to limit the rate of

reaction to avoid the crash precipitation of TATB. Preliminary

experiments showed that the concentration and rate of addi-

tion of acetic acid is critical in controlling the overall size of

the recrystallized TATB. The overall morphology of the

TATB crystals formed by this method showed a significant

improvement compared to the starting materials. By carefully

controlling the rate of addition, significantly large TATB

crystals (o500 mm) could be formed via this method.

Fig. 3 Optical images of TATB dissolving and recrystallizing in

EMImOAc: (a) at room temperature, (b) at 100 1C, (c) at 70 1C and

(d) cross-polarized light view of image (c).

Fig. 4 The solubility curve of TATB in the DMSO–EMImAOc

system.

Fig. 5 SEM micrographs of TATB (a) recrystallized from EMImOAc,

(b) starting material and (c) H2O-precipitated material.

Fig. 6 Raman spectra of (a) starting TATB and (b) recrystallized

TATB.


4. Discussion

In order to determine the ineffectiveness of ILs in solubilizing

TATB compared to cellulose, the cohesive energy density

(CED) of the two systems was compared. CED is often

expressed in its square-root form, known as the solubility

parameter, d. For cellulose, a variety of measurements, in-

cluding mechanical and surface free energy measurements,

suggest a value of d B 25 MPa12.16 For TATB, the heat

of sublimation (i.e. cohesive energy per molecule) is

B40.2 kcal mol�1.17 This, coupled with a molar volume of

221.2 A3 in the crystal phase,4 yields a value of dB 35.5 MPa12.

In other words, the CED (= d2) of TATB is approximately

2 times that of cellulose, explaining why the former is more

difficult to dissolve than the latter. The above-obtained solu-

bility parameter can be used as an approximate tool to screen

solvents for TATB. As an example, we note that the conven-

tional water-soluble IL, BMIm+BF4�, has a solubility para-

meter near to d = 26.5 MPa12.18 Given that TATB has a much

higher value of d, we expect the solubility of TATB to be poor

in such prototypical solvents in the absence of chemical

modification. In order to understand the reason for the higher

CED in TATB, we employed first-principles density functional

theory (DFT) using the program DMol3 19–22 to compute the

interlayer van der Waals binding and the intralayer (inter-

molecular) hydrogen bond contribution to the total cohesive

energy. We found that these two energies were comparable,

with the van der Waals contribution being almost 90% that of

the hydrogen bond contribution. The hydrogen bonds them-

selves have an average strength of B3.5 kcal mol�1, slightly

stronger than the hydrogen bonds in cellulose.23

The remarkable solubility of TATB in EMImOAc com-

pared to other ILs was initially very puzzling. Since the cation

of the ILs doesn’t seem to affect the solubility of TATB, we

concluded that the acetate moiety plays a key role in solubiliz-

ing TATB. TATB dissolved in EMImOAc produces a deep red

color, observed at lmax = 409 nm. The origin of this peak can

be assigned to a s-complex. There are many reviews on the

mechanism of the formation of s-complexes in intermolecular

nucleophilic displacement reactions involving electrophilic,

nitro-activated aromatic substrates.24 The mechanism gener-

ally involves the addition of a nucleophile to a position on the

electrophilic aromatic ring that results in the stabilization of

the negative charge by an ortho- and/or para-substituted nitro

group. The structure of the s-complex is shown in Fig. 7 and

consists of a cyclohexadienyl anion, in which the carbon center

that undergoes substitution is converted to an sp3-hybridized

center. The resulting negative charge may be delocalized into

the nitro groups, thus stabilizing the s-complex. Most studies

on the formation of s-complexes investigated 1-substituted-

2,4-dinitro- or 1-substituted-2,4,6-trinitrobenzenes, and very

few s-complexes of fully-substituted aromatic rings are

known. In the TATB molecule, because of its high-symmetry,

there are three equivalent positions at which the acetate anion

may react to produce the s-complex. The s-complex would be

stabilized by the presence of the three nitro groups at the 2-, 4-

and 6-positions relative to the tetrahedral carbon containing

the acetate group. In addition, the stability of the s-complex

relative to TATB would be enhanced by relieving steric

crowding on the TATB ring of the adjacent amino and nitro

groups upon forming the tetrahedral carbon center at the site

of base addition.

TATB added to EMImCl does not form a s-complex. In

order to understand the difference between the action of the

acetate and chloride anions on TATB, we carried out DFT-

based investigations using a state-of-the-art quantum chemical

conductor-like screening model (COSMO)25 and its extension

to real solvents (COSMO-RS).26 This model computes the

chemical potential of a solute in its own environment and in

solvent environments. From the difference between these

chemical potentials, one can estimate the solubility of the

solute in the solvent. Details of the procedure are described

elsewhere.27 Table 2 (column 2) displays the computed solu-

bility of TATB in EMImCl and EMImOAc, respectively. The

computed solubility in EMImCl is in line with the observed

low solubility (i.e. o0.5 wt%) in this solvent. In contrast, the

computed solubility in EMImOAc is 250 times smaller than

the experimentally observed value of 10 wt% at 100 1C. This

result, in conjunction with color changes observed in the

acetate IL solution, indicates that while TATB dissolves in

its pure form in EMImCl, it undergoes some chemical reaction

during its dissolution in EMImOAc. We have also computed

the possibility of a deprotonation mechanism for the observed

solubility in EMImOAc. To do this, we have investigated one

of the simplest reactions, i.e. an NH2 group of TATB loses a

proton to a neighboring anion of the IL (thereby forming an

acetic acid molecule in the acetate IL or a HCl molecule in the

chloride IL), while the unpaired cation of the IL binds to the

ortho position of the deprotonated TATB anion. Column 3 of

Table 2 lists the computed heat of reaction for such a chemical

process using the program Dmol3.20–22 The reaction is highly

endothermic and clearly prohibitive in the chloride IL, but

exothermic and likely to occur in the acetate IL. The above

results can be qualitatively explained based on the stronger

basicity of an OAc� group compared to Cl�. The computed

UV-vis spectrum (using the semi-empirical program

Fig. 7 Schematic of s-complex formation from TATB.

Table 2 COSMO-RS results for TATB dissolution in EMImOAcand EMImCl

SolventTATB solubilityat 100 1C (wt%)a DEdeprotonation/kcal mol�1b

EMImOAc 0.04 �1.6EMImCl 0.10 +22.0

a 1 wt%=1 g 100 mL�1= 10 g L�1. b Energy of the reaction: TATB+

EMIm+Anion� - EMIm+[TATBdeprotonated]� + H-Anion; DE is

positive (negative) for an endothermic (exothermic) reaction.


ZINDO)28 of the deprotonated TATB displayed a sharp peak

at B410 cm�1, which is also in excellent agreement with our

experimental observations. Therefore, we cannot at this point

discount the possibility of alternative chemical pathways of

s-complex formation and the deprotonation mechanism.

We have also attempted to dissolve cellulose in EMImOAc.

Although a recent report29 showed a high solubility of cellu-

lose (Avicel PH-101) in EMImOAc (15 wt%) at 110 1C, the

solubility of our cellulose (Eastman Kodak) in EMImOAc at

100 1C was less than 0.5 wt%. This discrepancy can be

attributed to different cellulose sources and therefore the level

of recalcitrance of the cellulose tested. Our cellulose experi-

ment result supports the idea that the mechanism of solubility

for TATB in EMImOAc is indeed by chemical modification

(rather than by hydrogen bond disruption). It is important to

note that ILs can potentially modify the solute, as with TATB,

to increase its solubility. Therefore, one must be very cautious

in determining the solubility of various solutes in ILs to make

sure that chemical modification of the solute is not the cause of

the observed increase in solubility.

Crystallization of TATB via the non-agitated cooling

method in EMImOAc resulted in an improved morphology

of the TATB crystals obtained compared to the starting

materials. Cooling by natural convection yielded crystals in

the size range of 10–50 mm (Fig. 5a). However, when the

cooling rate was controlled, larger crystals were obtained. As

seen in Fig. 8, when the crystallization solution was cooling at

a rate of 1 1C min�1, 200–500 mm sized crystals were obtained.

One drawback of the cooling crystallization method is that at

room temperature, the amount of recovered TATB is less than

the original amount of TATB added, due to dissolved TATB

in the 80 : 20 solution; the addition of a protic solvent is

necessary to recover the remaining TATB. Thus, an anti-

solvent crystallization scheme was employed as an alternative.

As mentioned above, during our initial attempts to recrystal-

lize of TATB from DMSO–EMImOAc solution, the recovery

of the TATB was rather poor because it is soluble in the

mixture at a 1 wt% level at room temperature. The most

effective anti-solvents of all those tested were hydroxylic

compounds such as organic alcohols and acids. These solvents

provided a proton source that released the acetate from the

s-complex and precipitated the TATB. In addition, some

inorganic acids, such as sodium hydrogen phosphate and boric

acid, were also found to be effective. A limiting factor in the

ability of these inorganic salts to act as an efficient anti-solvent

was their solubility in DMSO. Some had limited solubility,

rendering them less effective in precipitating TATB in good

yields. The addition of gaseous CO2 to the solvent mixture also

precipitated TATB in excellent yields. It is not clear whether

the CO2 acted to neutralize the mixture, making it less basic,

or actually acted as a counter-solvent to precipitate the TATB.

The use of gaseous CO2 is attractive because of its cost and

availability, but it was hard to control the particle size of the

recrystallized TATB when adding CO2 in gaseous form.

The anti-solvent of choice was acetic acid because it was

reasoned that it could be removed from the IL to yield

recovered EMImOAc without contamination from other salts.

Preliminary results show that crystal size and shape are

strongly dependent on the rate of addition and the concentra-

tion of the anti-solvent. Detailed information regarding the

crystallization of TATB via the anti-solvent method will be

published elsewhere.

5. Conclusion

ILs previously shown to dissolve highly hydrogen-bonded

solids were ineffective with TATB. This may be due to the

cohesive energy of TATB, which is almost 2 times that of

cellulose. The remarkable solubility of TATB in EMImOAc

was attributed to the formation of a s-complex or to the

deprotonation of TATB. Owing to the basicity and nucleo-

philicity of the OAc� anion, the s-complex could easily form

in EMImOAc, but not in the presence other anions such as

Cl�. Therefore, the solubility of TATB in EMImOAc is

proportional to the concentration of EMImOAc. The dis-

solved s-complex can revert back to TATB, either via cooling

or by the addition of an anti-solvent. The recrystallized TATB

shows a much improved crystal morphology compared to the

starting material. We are currently performing experiments to

further control the crystallization of TATB by the anti-solvent

crystallization method.

Acknowledgements

This work performed under the auspices of the US Depart-

ment of Energy by Lawrence Livermore National Laboratory

under Contract DE-AC52-07NA27344. The project 06-SI-005

was funded by the Laboratory Directed Research and Develop-

ment Program.

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Supramolecular synthesis of some molecular adducts of 4,40-bipyridineN,N0-dioxidew

Kapildev K. Arora, Mayura S. Talwelkar and V. R. Pedireddi*

Received (in Durham, UK) 8th May 2008, Accepted 8th September 2008

First published as an Advance Article on the web 29th October 2008

DOI: 10.1039/b807853j

Molecular adducts (1a–1e) of 4,40-bipyridine N,N0-dioxide, 1, respectively with cyanuric acid,

trithiocyanuric acid, 1,3,5-trihydroxybenzene (phloroglucinol), 1,3-dihydroxybenzene (resorcinol)

and 1,2,4,5-benzenetetracarboxylic acid have been reported. The major interactions observed in

the structures 1a–1e are N–H� � �O, N–H� � �S, O–H� � �O and C–H� � �O, in the form of homomeric

and heteromeric patterns of the constituents, either as a single or cyclic hydrogen-bonded motifs.

While in the adduct 1a, both homomeric and heteromeric units of both the constituents were

observed, no heteromeric interactions were observed in 1b and 1c. In addition, in 1b, homomeric

aggregation of molecules of 1 occurred in association with water molecules. However, while

heteromeric interactions prevail between the constituents in 1d and 1e, only one of the

co-crystallizing species gave homomeric interactions (4,40-bipyridine N,N0-dioxide in 1d; 1,2,4,5-

benzenetetracarboxylic acid in 1e). Further, in either type of the patterns, the cyclic motifs are

formed as a pair-wise hydrogen bonds comprising of strong and weak hydrogen bonds

(N–H� � �O/C–H� � �O or O–H� � �O/C–H� � �O). In three-dimensions, the ensembles of molecules

yield planar sheets, ladders and pseudorotaxane type assemblies.

Introduction

Design and synthesis of molecular complexes/adducts employ-

ing noncovalent interactions such as hydrogen bonds, which is

broadly defined as supramolecular synthesis, aims at creation

of exotic functional solids and in this connection, exploration

of novel ligands is a continuous process.1 Thus, a variety of

ligands of different molecular dimensions and functional

properties were utilized for the preparation of numerous

supramolecular assemblies of exotic architectures as reported

in the recent literature.2 Among those, 4,40-bipyridine (bpy) is

well studied, especially as a spacer molecule, both in organic

and organic–inorganic hybrid complexes. It was mainly due to

the ability of bpy to form either O–H� � �N only or O–H� � �N/

C–H� � �O pair-wise hydrogen bonds and also dative bonds

with metal ions in conjunction with carboxyl/carboxylate and

many other functional moieties.3

In further exploration and thrust to identify other spacer

molecules, compounds that mimic bpy topologically, for ex-

ample, 1,2-bis(4-pyridyl)ethene and ethane, 1,3-bis(4-pyridyl)-

propane etc., evolved as novel ligands for the preparation of

the tailor-made supramolecular assemblies of desired archi-

tectures and properties.4 Also, in recent times, 4,40-bipyridine

N,N0-dioxide (N-oxide derivative of bpy), 1, has been well

considered in the synthesis of coordination assemblies, but

corresponding organic supramolecular assemblies are limited.5

Since the N-oxide, 1 is a potential hydrogen bond acceptor to

establish interaction with complementary functionalities such as

–OH, –COOH, –NH, –CONH2 etc., it is rather surprising that

1 was not utilized, so effectively, in the supramolecular synthesis

of organic assemblies, as only a few reports are known in the

literature.6 Apart from it, the native structure of 1 itself is not

known in the literature. Thus, we are interested to elucidate the

structure of 1 and also study its application in the molecular

recognition and supramolecular synthesis with different organic

functional moieties such as –OH, –COOH, which are well

known to yield discrete molecular recognition patterns.1c,3b,d

In this direction, our attempts to obtain single crystals of

suitable quality for structure elucidation of 1 are not successful

yet, but co-crystallization experiments of 1 with cyanuric acid,

trithiocyanuric acid, 1,3,5-trihydroxybenzene (phloroglucinol),

1,3-dihydroxybenzene (resorcinol) and 1,2,4,5-benzenetetra-

carboxylic acid, possessing different functional moieties, as

shown in Chart 1, gave molecular complexes in the form of

single crystals. The structural features of these unusual

molecular adducts, unravel by single-crystal X-ray diffraction

methods, are described in this article.


Solid state structure of molecular complex, 1a, of 4,40-bipyridine

N,N0-dioxide, 1 and cyanuric acid (CA)

Co-crystallization of 1 and cyanuric acid, CA, from a methanol

solution gave good quality single crystals, 1a, in a 1:2 ratio of the

reactants and it was characterized by X-ray diffraction methods.

The pertinent crystallographic information is given in Table 1.

Solid State & Supramolecular Structural Chemistry Unit, Division ofOrganic Chemistry, National Chemical Laboratory, Dr. Homi BhabhaRoad, Pune, 411008, India. E-mail: [email protected];Fax: +91 20 25892629; Tel: +91 20 25902097w Electronic supplementary information (ESI) available: ORTEPdiagrams of 1a–1e. Search overview details for Cambridge StructuralDatabase (CSD). CCDC reference numbers 672332–672336. For ESIand crystallographic data in CIF or other electronic format see DOI:10.1039/b807853j



Analysis of molecular packing reveals that in the complex

1a, each molecule of 1 establish interaction with two dimers of

CA, as shown in Fig. 1, by forming two different pair-wise

hydrogen bonding patterns of N–H� � �O (H� � �O, 1.70 and

1.71 A; N� � �O, 2.69 and 2.68 A) and C–H� � �O (H� � �O, 2.38

and 2.44 A; C� � �O, 3.34 and 3.33). Such a recognition pattern

gave a three-dimensional structure, as stacked layers, which is

shown in Fig. 1(b).

However, the arrangement of molecules in a typical sheet is

quite intriguing. Although 1 and CA established heteromeric

pattern, each one in turn form one-dimensional crinkled tapes,

through homomeric pattern by holding the adjacent molecules,

as shown in Fig. 2. While CA molecules form homomeric

patterns through cyclic N–H� � �O hydrogen bonds, with

H� � �O distances being in the range, 1.70–2.02 A (N� � �O,

2.60–2.85 A), molecules of 1 gave such patterns through a

cyclic pattern of C–H� � �O hydrogen bonds and the corres-

ponding H� � �O distances are 2.36 and 2.40 A, (C� � �O, 3.31

and 3.33 A). Further, the molecular tapes of 1 and CA are

arranged alternatively in two-dimensional sheets. In fact, the

homomeric patterns observed for 1 and CA are the most

commonly observed arrangement in many of their molecular

complexes.7 It is interesting to note that pure crystal structure

of CA also is due to the aggregation of such molecular tapes,8

as observed in 1a, held together by single N–H� � �O hydrogen

bonds, as shown in Fig. 2(b). However, such an inference

could not be established about the arrangement of molecules

of 1 as its pure crystal structure is not known. However, since

Chart 1

Table 1 Crystallographic details of crystal structures of molecular adducts, 1a–1e

1a 1b 1c 1d 1e

Formula C10H8N2O2: C10H8N2O2: 1.5(C10H8N2O2): 2(C10H8N2O2): C10H8N2O2:2(C3H3N3O3) 2(C3H3N3S3):2(H2O) C6H6O3 2(C6H6O2):4(H2O) C10H6O8

Mr 446.35 578.74 408.38 668.65 442.33Crystal morphology Blocks Blocks Blocks Rectangular blocks BlocksCrystal color Colorless Colorless Pale-yellow Colorless ColorlessCrystal system Triclinic Monoclinic Triclinic Triclinic MonoclinicSpace group P�1 C2/c P�1 P�1 P21/ca/A 8.218(3) 22.129(8) 10.111(2) 7.129(1) 12.926(5)b/A 9.299(4) 13.217(5) 10.277(2) 10.253(2) 7.948(3)c/A 12.168(5) 8.531(3) 10.405(2) 23.220(4) 19.059(7)a/1 91.93(1) 90 70.61(1) 82.15(1) 90b/1 91.44(1) 105.82(1) 84.88(1) 85.26(1) 106.54(1)g/1 108.10(1) 90 61.60(1) 70.40(1) 90V/A3 882.7(6) 2400.6(15) 902.7(3) 1582.7(5) 1877.0(1)Z 2 4 2 2 4Dc/g cm�3 1.679 1.601 1.502 1.403 1.565T/K 298(2) 298(2) 298(2) 298(2) 273(2)l(Mo-Ka) 0.71073 0.71073 0.71073 0.71073 0.71073m/mm�1 0.138 0.612 0.112 0.109 0.1282y range/1 46.60 46.68 56.54 46.54 56.56Limiting indices �9 r h r 9 �24 r h r 24 �13 r h r 13 �7 r h r 7 �15 r h r 17

�10 r k r 8 �14 r k r 14 �13 r k r 13 �11 r k r 10 �10 r k r 6�13 r l r 11 �9 r l r 7 �13 r l r 13 �25 r l r 25 �25 r l r 24

F(000) 460 1192 426 704 912No. reflns measured 3845 5055 10278 6950 10778No. unique reflns [R(int)] 2526 [0.0281] 1739 [0.0229] 4062 [0.0418] 4524 [0.0333] 4347 [0.0238]No. reflns used 1983 1539 3342 1999 3401No. parameters 345 193 344 465 345Reflection 7.32 9.01 11.80 9.73 12.6GOF on F2 1.043 1.139 1.038 0.821 1.018R1 [I 4 2s(I)] 0.0612 0.0353 0.0558 0.0438 0.0480wR2 0.1520 0.0896 0.1571 0.0971 0.1196Drmax, min/e

� A�3 0.38, �0.44 0.44, �0.39 0.26, �0.34 0.24, �0.23 0.249, �0.288


the majority of N-oxide structures possess the homomeric

patterns of 1, as shown in Fig. 2(a), following the analogy

observed for CA, the pure structure of 1 could be visualized as

a combination of such tapes and this may provide means to

establish the structure of 1 by other methods, such as powder

X-ray diffraction techniques, as it fails to yield single crystals

so far, without additional molecules (either solvent of crystal-

lization or co-crystallizing agent).

Thus, 1a could be visualized as a representative example for

the combination of unity and diversity with the observation of

homomeric and heteromeric patterns of both the co-crystal-

lizing species simultaneously. Also, the dual role of N-oxide 1,

as a spacer and structure directing, could be established, unlike

4,40-bipyridine, which often play a role of spacer, except in the

recently reported assemblies, wherein it acts as a guest.9 In

order to corroborate such features through a large number of

molecular complexes of 1, co-crystallization of it with trithio-

cyanuric acid, TCA, which is an analogue of CA, has been

carried out, expecting formation of an iso-structural complex

with that of 1a, by which relative competition for homomeric

and heteromeric patterns could also be programmed.

Solid state structure of adduct, 1b, of 4,40-bipyridine

N,N0-dioxide, 1 and trithiocyanuric acid (TCA)

N-oxide, 1 gave co-crystals with TCA as a hydrate and it has

been labeled as 1b. Further, the asymmetric unit consists of

1:2 ratio of the reactants, and the important crystallographic

information is given in Table 1. The molecular arrangement in

two- and three-dimensions in the crystal structure of 1b is

shown in Fig. 3.

In 1b, three-dimensional structure is alike in 1a, but through

stacked crinkled sheets (Fig. 3), rather than planar sheets.

Further, in contrast to the structure of 1a, a heteromeric pattern

between the molecules of 1 and TCA is not observed. Instead,

the interaction between 1 and TCA is established through water

molecules. Thus, TCA forms N–H� � �O hydrogen bonds

(H� � �O, 1.66 A, N� � �O, 2.61 A) with water molecules, while 1

forms O–H� � �O hydrogen bonds (H� � �O, 1.86 and 1.91 A with

corresponding O� � �O, 2.73 and 2.70 A), as shown in Fig. 3(b).

Such an ensemble ultimately self-assembles, leading to the

formation of two-dimensional sheets with tapes of TCA mole-

cules separated by the aggregates of 1 and water. Within each

molecular tapes of TCA, the adjacent molecules are held

together by N–H� � �S hydrogen bonds with H� � �S distances of

2.52 and 2.54 A (N� � �S distances of 3.39 and 3.41 A).

To evaluate, further, the nature of the variable hydrogen-

bonding patterns of 1 in the presence of other molecular

entities with potential hydrogen bond donor functionalities,

co-crystallization of 1 with 1,3,5-trihydroxybenzene (THB)

which may be regarded as analogue of CA in its enol form,

as shown below, has been carried out.

Supramolecular assembly in molecular complex, 1c, of

4,40-bipyridine N,N0-dioxide, 1, and 1,3,5-trihydroxybenzene

(phloroglucinol), THB

Co-crystallization of 1 and THB from a methanol solution

gave a molecular complex, 1c in a 3:2 ratio of the reactants 1

Fig. 1 (a) Molecular recognition between 1 and CA in the crystal structure of 1a. (b) Three-dimensional arrangement of molecules in the crystal

structure of 1a, in the form of stacked layers.

Fig. 2 (a) Molecular tapes of 1 and CA through homomeric patterns which are held together by heteromeric in the structure of 1a. (b)

Arrangement of molecules in the crystal structure of CA.


and THB. Analysis of three-dimensional packing reveals

several quite exciting features, especially a pseudorotaxane

type network in the form of a host–guest type assembly, as

shown in Fig. 4.

Although each THB interacts with three molecules of 1,

forming a heteromeric pattern by O–H� � �O hydrogen bonds

with H� � �O distances of 1.71, 1.73 and 1.75 A (O� � �O, 2.63,

2.60, 2.67 A), as shown in Fig. 5(a), the homomeric patterns

formed by both 1 and THB play a crucial role in the formation

of ultimate exotic structure in 1c. The homomeric pattern of

THB is shown in Fig. 5(b) and the corresponding patterns of 1

are shown in Fig. 6.

The molecules of THB were found to be yielding a molecular

tape, through homomeric pattern, constituted by C–H� � �Ohydrogen bonds (see Fig. 5(b)), which is, in fact, unknown

either in its pure structure or in its molecular complexes.10

Further, two molecules of 1 in the asymmetric unit of 1c also

form molecular tapes independently. Interestingly, while one of

these remains like infinite tapes, the tapes belonging to the

second molecule are held together by cyclic C–H� � �O hydrogen

bonding patterns constituting layers with void space (Fig. 6). In

those cavities the tapes of THB molecules fit like a thread,

yielding a pseudorotaxane type structure (Fig. 4(b)). Earlier,

in our investigations on 1,10-phenanthroline complexes, we

demonstrated the feasibility of such structures entirely engraved

by noncovalent interactions.11

Thus, molecular complex, 1c further demonstrates the

elegancy of noncovalent synthesis to mimic the ensembles

known to exist for decades, often, being synthesized by con-

ventional means. Looking at the tapes formed by THB, it

appears that such tapes could be even possibly synthesized by

dihydroxybenzene as well, which may possibly also can yield a

pseudorotaxane type structure as observed in 1c. Hence, co-

crystallization of 1 with 1, 3-dihydroxybenzene (DHB) has

been carried out.

Molecular complex, 1d, of 4,40-bipyridine N,N0-dioxide, 1 and

1,3-dihydroxybenzene, DHB

N-Oxide, 1 and DHB form co-crystals, 1d, in a 1:1 ratio along

with two molecules of water and crystallize in triclinic space

group, P�1. The three-dimensional arrangement of these mole-

cules is indeed quite interesting with a stair-case type structure.

A typical arrangement is shown in Fig. 7.

A detailed analysis of the arrangement reveals that both the

symmetry independent molecules of 1, form homomeric pat-

terns independently, as observed in 1a and 1c, yielding mole-

cular tapes through C–H� � �O hydrogen bonds (H� � �O, 2.40

and 2.51; 2.47 and 2.51 A with corresponding C� � �O, 3.31 and

3.38 A; 3.40 and 3.40 A). Infinite tapes corresponding to a

particular symmetry are only shown in Fig. 8(a), for the

purpose of clarity, while the tapes of the other symmetry

independent molecules is shown in the inset of Fig. 8(a). The

tapes correspond to both the symmetry independent mole-

cules, are held together by two water molecules through

O–H� � �O (H� � �O, 1.67 A; O� � �O, 2.77 A) and C–H� � �O

Fig. 3 (a) Stacking of layers comprising of molecules of 1 and TCA in the crystal structure of 1b. (b) Two-dimensional arrangement of molecules

showing the molecular tapes of TCA separated by the molecules of 1 and water, which are held together by O–H� � �O hydrogen bonds.

Fig. 4 (a) Pseudorotaxane type network in the crystal structure of 1c, with void space being filled by a molecular tape of 1. Schematic

representation is shown as inset. (b) A typical pseudorotaxane network with molecules of 1 as rings and molecules of THB as rods.


(H� � �O, 2.54 A; C� � �O, 3.31 A) hydrogen bonds, constituting

cavities. The water molecules, in turn, are held together by

O–H� � �O hydrogen bond with a H� � �O distance of 1.94 A

(O� � �O, 2.83 A). In the cavities, two DHB molecules, which

are held together by C–H� � �O (H� � �O, 2.90 and 2.91 A; C� � �O,

3.51 and 3.52 A) hydrogen bonds are situated. These DHB

molecules are further glued to the tapes of 1 by O–H� � �O and

C–H� � �O hydrogen bonds. Such adjacent ensembles are

further held together, as shown in Fig. 8(a), by water mole-

cules connecting the two molecular tapes corresponding to

the same symmetry molecules by O–H� � �O and C–H� � �Ohydrogen bonds. A schematic representation of the arrange-

ment is shown in Fig. 8(b).

Thus, in complex 1d, only the molecules of 1 aggregated to

yield homomeric patterns, while DHB remains as monomers

forming interactions with 1 yielding heteromeric patterns.

Taking into account the facile formation of ladders and

stair-case type structures by 4,4 0-bipyridine (bpy) with

–COOH functionality, and in particular, the recent reports

of preparation of such architectures by co-crystallizing bpy

with 1,2,4,5-benzenetetracarboxylic acid (BTCA),12 further

studies have been directed to create supramolecular assembly

of 1 and BTCA.

Supramolecular assembly in molecular complex, 1e, of

4,40-bipyridine N,N0-dioxide, 1, and

1,2,4,5-benzenetetracarboxylic acid, BTCA

Co-crystallization of 1 and BTCA gave a 1:1 molecular com-

plex, 1e. In this structure (Fig. 9), in two-dimensional arrange-

ment, each molecule of 1 interacts with BTCA forming

heteromeric pattern through the formation of O–H� � �O/

C–H� � �O hydrogen bonding patterns, H� � �O, 1.60/2.30;

1.49/2.59 A (O� � �O, 2.53/3.23; 2.48/3.53 A). But, molecules

of 1 did not undergo homomeric aggregation, in the structure

of 1e. In contrast, molecules of BTCA show homomeric

recognition pattern through well known R22(8) hydrogen

bonding pattern, with H� � �O distances of 1.67 and 1.70 A

(O� � �O, 2.63 and 2.65 A), via the remaining –COOH groups,

that did not interact with the molecules of 1.

Thus, the arrangement ultimately could be visualized as

sheets with layers of molecules of BTCA stuffed by the

molecules of 1 with appreciable void space. However, in

three-dimensional arrangement, the adjacent layers are

arranged in such a manner that molecules from the adjacent

layers effectively fill the void space; thus, 1e could not yield a

channel structure. It is noteworthy to mention that among all

the structures studied in this series (1a–1e), molecules of 1 did

not undergo homomeric recognition only in the structure of 1e,

perhaps, due to the strong interaction between –COOH and

N - O moieties, thus exhibiting the ability of 1 also to

perform the role of spacer, like its analogue bpy, and suggests

the importance of the complementarity between the functional

groups undergoing the molecular recognition process.

Fig. 5 (a) Molecular recognition between 1 and THB, yielding heteromeric patterns. (b) Homomeric pattern of THB.

Fig. 6 Homomeric patterns of N-oxide, 1 in the crystal structure of 1c.

Fig. 7 Three-dimensional packing of molecules in the crystal structure of 1d.


Conclusions

In this study, we have shown the ability of 4,40-bipyridine

N,N0-dioxide to yield different types of supramolecular assem-

blies from simple stacked sheet structures to pseudorotaxane

and stair-case type structures depending upon its interaction

with the co-crystallizing agents. Unlike its analogue, bpy, the

N-oxide shows preference for the homomeric patterns, although

its spacer role is visualized in the structure 1e. Further, ob-

servation of the homomeric patterns formed by 1 in the crystal

structures, 1a–1d, and also in some of the examples found in the

literature, it may be possible to extrapolate it to predict the

three-dimensional structure of 1 as a stacked sheets with each

sheet as an aggregation of molecular tapes formed by the

mutual recognition of the adjacent molecules through C–H� � �Ohydrogen bonds. Thus, we strongly believe that this can be a

good starting model to determine the three-dimensional struc-

ture of 1 by other techniques such as powder X-ray diffraction

methods or by computational procedures and we have already

initiated process in this direction.

Experimental

Preparation of molecular adducts of the molecular complexes,

1a–1e

All the chemicals used in this study were obtained from

commercial suppliers and used as such without any further

purification. The solvents employed for the crystallization

purpose were of spectroscopy grade of highest available

purity. Co-crystals have been prepared by dissolving 4,40-

bipyridine N,N0-dioxide, 1, and cyanuric acid, trithiocyanuric

acid, 1,3,5-trihydroxybenzene, 1,3-dihydroxybenzene and

1,2,4,5-benzenetetracarboxylic acid in 1:1 or 1:2 ratio either

in CH3OH or H2O as solvent and slowly evaporating the

obtained solution. Single crystals were obtained over a period

of 48 h in all the cases. In typical preparation, 0.0941 g

(0.5 mmol) of 1 and 0.127 g (0.5 mmol) of 1,2,4,5-benzenetetra-

carboxylic acid were dissolved in 15 mL of CH3OH by gently

warming on a water bath. The resultant solution was kept for

evaporation at ambient conditions by protecting the conical

flask from external mechanical disturbances and within 48 h,

colorless and good quality crystals of 1e, were obtained that

are suitable for studies by single-crystal X-ray diffraction

methods.

Crystal structure determination of 1a–1e

Good quality single crystals of 1a–1e have been chosen by

viewing under microscope and glued to a glass fiber using an

adhesive to mount on a goniometer of Bruker single crystal

X-ray diffractometer equipped with APEX CCD detector. The

data collection was smooth in all the cases without any

complications and all the crystals were found to be stable

throughout data collection period. The intensity data were

processed using Bruker suite programmes, SAINT,13 followed

Fig. 8 Arrangement of molecules within a two-dimensional layer in the crystal structure of 1d.

Fig. 9 (a) Two-dimensional arrangement of molecules in the crystal structure of 1e. (b) Stacking of sheets in three-dimensions.


by absorption correction by SADABS.14 The structures were

solved by using XS and refined by least-square methods using

XL.15 All the non-hydrogen atoms were refined by anisotropic

methods and the hydrogen atoms were either refined or placed

in the calculated positions. All the structural refinements

converged to good R-factors as listed in Table 1.

Acknowledgements

We thank Department of Science and Technology (DST) for

the financial support and also greatly acknowledge Professor

Judith A. K. Howard (Durham, UK) for her generous support

to Mayura by awarding scholarship. Also one of us (K. A.)

thanks Council of Scientific and Industrial Research (CSIR),

for the Research Fellowship.

References

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5 A search performed on Cambridge Structural Database (CSD)using version 1.10, retrieved 140 entries possessing 4,40-bipyridineN,N0-dioxide, in which 118 are found to be organometallic while 22are only the organic molecular complexes (see ESIw).

6 (a) R. Thaimattam, D. S. Reddy, F. Xue, T. C. W.Mak, A. Nangiaand G. R. Desiraju, J. Chem. Soc., Perkin Trans. 2, 1998, 1783;(b) L. S. Reddy, N. J. Babu and A. Nangia, Chem. Commun., 2006,1369; (c) M. T. Messina, P. Metrangolo, W. Panzeri, T. Pilati andG. Resnati, Tetrahedron, 2001, 57, 8543.

7 (a) A. Ranganathan, V. R. Pedireddi, G. Sanjayan, K. N. Ganeshand C. N. R. Rao, J. Mol. Struct., 2000, 522, 87; (b) M. T. Messina,P. Metrangolo, W. Panzeri, T. Pilati and G. Resnati, Tetrahedron,2001, 57, 8543; (c) J. A. Bis, P. Vishweshwar, D. Weyna andM. J. Zaworotko, Mol. Pharm., 2007, 4, 401; (d) T. C. Lewis,D. A. Tocher and S. L. Price, Cryst. Growth Des., 2005, 5, 983;(e) V. R. Pedireddi and D. Belhekar, Tetrahedron, 2002, 58, 2937.

8 (a) G. C. Verschoor and E. Keulen, Acta Crystallogr., Sect. B, 1971,27, 134; (b) H. Dietrich, C. Scheringer, H. Meyer, K.-W. Schulte andA. Schweig, Acta Crystallogr., Sect. B, 1979, 35, 1191.

9 V. R. Pedireddi and J. PrakashaReddy, Tetrahedron Lett., 2003,44, 6679.

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11 K. K. Arora and V. R. Pedireddi, Cryst. Growth Des., 2005, 5,1309.

12 (a) N. Shan, A. D. Bond and W. Jones, Cryst. Eng., 2002, 5, 9;(b) J. J. Kane, R.-F. Liao, J. W. Lauher and F. W. Fowler, J. Am.Chem. Soc., 1995, 117, 12003.

13 SAINT, Version 6.02; Bruker AXS, Inc., Analytical X-raySystems, 5465 East Cheryl Parkway, Madison, WI 53711-5373,2000.

14 SADABS [Area-Detector Absorption Correction]; Siemens Indus-trial Automation, Inc.: Madison, WI, 1996.

15 G. M. Sheldrick, SHELXTL, University of Gottingen, Gottingen,Germany, 1997.


Synthesis and characterisation of bulky guanidines and

phosphaguanidines: precursors for low oxidation state metallacyclesw

Guoxia Jin,b Cameron Jones,*a Peter C. Junk,a Kai-Alexander Lippert,b

Richard P. Roseab

and Andreas Stascha

Received (in Durham, UK) 29th May 2008, Accepted 20th August 2008


DOI: 10.1039/b809120j

Reactions of alkali metal amides or phosphides with the bulky carbodiimide, ArNQCQNAr

(Ar = C6H3Pri2-2,6), followed by aqueous work-ups, have yielded several guanidines,

ArNC(NR2)N(H)Ar (R = cyclohexyl (GisoH) or Pri (PrisoH); NR2 = cis-NC5H8Me2-2,6

(PipisoH)), a bifunctional guanidine, {ArNCN(H)Ar}2{m-N(C2H4)2N} (Pip(GisoH)2), and two

phosphaguanidines, ArNC(PR2)N(H)Ar (R = cyclohexyl (CyP-GisoH) or Ph (PhP-GisoH)).

A very bulky guanidine, ArNC{N(Ar)SiMe3}N(H)Ar (ArSi-Giso), and an aryl coupled

bifunctional guanidine, {ArN(H)C(NPri2)NC6H2Pri2-2,6-}2 (PrisoH)2, have been prepared by

other routes. All compounds have been crystallographically characterised and shown to exist in

a number of isomeric forms in the solid state. These appear to be largely retained in solution.

The deprotonation of GisoH with BunLi in either hexane or THF led to crystallographically

characterised dimeric and monomeric complexes respectively, viz. [Li{Li(k2-N,N0-Giso)2}] and

[Li(THF)(Z1-N,Z3-Ar-Giso)]. Deprotonation of PrisoH and Pip(GisoH)2 with K[N(SiMe3)2] gave

the unsolvated polymer, [{K(Z1-N,Z6-Ar-Priso)}N], and the solvated complex,

[{K(THF)2}{Pip(Giso)2}{K(THF)3}], respectively.

Introduction

The coordination chemistry of anionic amidinate

([RNC(R)NR]�, R = H, alkyl, aryl etc.) and guanidinate

([RNC(NR2)NR]�) ligands has been extensively studied, giving

rise to numerous complexes incorporating metals from across the

periodic table.1 In these, the ligands have displayed an impressive

array of coordination modes which depend upon the nature and

bulk of the substituents (R), and the metal involved. This

structural diversity is one of the main factors that have led

to such complexes finding many applications in catalysis,2–4

materials science5 and synthesis,1 to name but a few.

Recent developments in this area have concentrated on the use

of very bulky amidinates to stabilise low nuclearity s- and

p-block metal complexes which show significant potential as,

for example, lactide polymerisation catalysts.2 Of most note here

is the Piso� ligand, [ArNC(But)NAr]�, which incorporates

sterically demanding 2,6-diisopropylphenyl (Ar) substituents at

its N-centres and a tert-butyl group on the backbone carbon.

The spatial profile and ligating abilities of this ligand have

been likened to those of b-diketiminates, the most commonly

utilised examples of which also possess N–Ar substituents,

e.g. [(ArNCR)2CH]� (R = Me or But).6 Although complexes

of b-diketiminates are widely used in catalytic processes, they are

perhaps more notable for their capacity to kinetically stabilise

complexes containing low oxidation state metal centres. A salient

illustration of this is the synthesis and structural characterisation

of the homologous series of monomeric, N,N0-chelated group 13

metal(I) complexes, [:M{(ArNCMe)2CH}] (M = Al, Ga, In or

Tl),7 which have shown remarkable further chemistry.

In contrast to b-diketiminates, bulky amidinates (e.g. Piso�)

had rarely been employed in the preparation of low oxidation

state metal complexes. In 2005, we began to address this

paucity with the preparation of the group 13 metal(I) com-

plexes [:M(Piso)] (M = In or Tl).8 However, unlike their

b-diketiminate counterparts, [:M{(ArNCMe)2CH}], the Piso�

ligand in these complexes is localised and chelates the metal

centre in an Z1-N,Z3-arene-fashion. In addition, the analogous

GaI and AlI complexes could not be stabilised. These results

suggested that related, but bulkier ligands would need to be

accessed to enforce N,N0-chelation and allow stabilisation of

lighter group 13 metal(I) centres. To this end, the very large

guanidinate ligand, [ArNC(NCy2)NAr]� (Giso�; Cy = cyclo-

hexyl), was developed and used in the syntheses of the

remarkably stable monomeric four-membered heterocycles,

[:M(k2-N,N0-Giso)] (M = Ga or In; N.B. the Al(I) heterocycle

has not yet been accessed),9 the coordination chemistry of

which was later explored.10 In addition to the increased steric

bulk of Giso� over Piso�, the greater stabilising ability of the

guanidinate can be attributed to the fact that it is a more

N-electron rich donor than the amidinate, a result of it

possessing a zwitterionic resonance form containing two nega-

tively charged N-donor centres, viz. [Cy2N+QC(N�Ar)2].

a School of Chemistry, PO Box 23, Monash University, 3800 VIC, Australiab School of Chemistry, Main Building, Cardiff University, Cardiff,UK CF10 3AT

w Electronic supplementary information (ESI) available: ORTEP dia-grams for 2 and 3. Crystallographic data (excluding structure factors)for the structures of 1–12. CCDC reference numbers 704662 (1),704663 (2), 704664 (3), 704665 (4�2CHCl3), 704666 (5), 704667 (6),704668 (7), 699384 (8�hexane), 704669 (9), 704670 (10), 704671 (11),704672 (12�2THF). For ESI and crystallographic data in CIF or otherelectronic format see DOI: 10.1039/b809120j



Over the last three years we have extended our application

of Giso�, and a range of other Ar-substituted guanidinate and

phosphaguanidinate ([ArNC(PR2)NAr]�)11 ligands, to the

stabilisation of heterocyclic complexes containing low oxida-

tion state metal centres from all blocks of the periodic table

(e.g. Mg(I),12 Ge(I),13 As(I),14 various d-block metal(I)15 and

f-block metal(II)16 species) with considerable success. More-

over, we have used these ligands in the synthesis of a variety of

gallyl–metal complexes, including examples exhibiting un-

precedented Ga–Zn17 and Ga–Sn18 bonds. In all these studies,

the ligands have been prepared by the deprotonation of

neutral guanidines or phosphaguanidines with alkali metal

reagents. Although some preliminary details of the synthesis of

the neutral ligand precursors have been previously been des-

cribed by us,9–16 it seemed that a full report of the preparation

and characterisation of these compounds would aid other

researchers seeking to harness their unique properties for their

own purposes. The value of this is highlighted by the fact that

prior to our involvement in this field, only one guanidine

bearing 2,6-diisopropylphenyl substituents at its N-centres,

viz. ArNC{N(H)Ar}2, had appeared in the literature.19 Here,

we report on the synthesis, structures and properties of eight

N–Ar substituted guanidines and phosphaguanidines, and

some of their alkali metal derivatives.


(i) Synthesis of bulky guanidines and phosphaguanidines

A number of synthetic routes are known for the preparation

of guanidines.1 One of the most versatile of these involves

the addition of metallated amides to carbodiimides

(RNQCQNR), followed by aqueous work-up. Here, this

route has been employed to synthesise the guanidines GisoH

(1), PrisoH (2), PipisoH (3), as well as the bifunctional

guanidine, Pip(GisoH)2 (4), in high to quantitative yields

(Scheme 1). In all preparations, THF was used as the solvent

and the initial addition reactions were carried out at either

ambient temperature and/or under reflux conditions.

It appears that this route does have steric and electronic

limitations, as the attempted addition of some amides to

the carbodiimide (ArNQCQNAr) were not successful. For

example, lithiated cis-2,6-dimethylpiperidine adds to the

carbodiimide to give compound 3, whereas lithiated 2,2,6,6-

tetramethylpiperidine does not react with ArNQCQNAr in

THF at reflux. Moreover, M[N(SiMe3)2] (M = Li, Na or K)

do not react with ArNQCQNAr under similar conditions,

though these reagents are known to add to smaller carbodii-

mides at room temperature.20

Although considerably less sterically demanding than some

of the amide precursors mentioned above, lithium carbazolyl

did not react with ArNQCQNAr in THF at reflux, and only

carbazole and the carbodiimide were recovered after work-up.

This lack of reactivity probably derives from the lower nucleo-

philicity of the aromatic carbozyl anion, relative to the bulkier

amides used in the preparation of 1–3.

Interest in the coordination chemistry of phosphaguanidi-

nates, [RNC(PR 02)NR]�, has recently begun to escalate.1d,21

One of the main reasons behind this is that the phosphino

group of these ligands is pyramidal, unlike the planar amino

substituent of guanidinates. Therefore, the zwitterionic reso-

nance form of these ligands, [R02P+QC(N�R)2], does not play

a significant role in their chemistry. As a result, phospha-

guanidinates are coordinatively versatile, and in many of their

complexes the phosphino group acts a P-lone pair donor.1d,21

Despite this emerging importance, there had been no reports

of N–Ar substituted phosphaguanidinates or phospha-

guanidines in the literature. We have reversed this situation

with the synthesis of CyP-GisoH, 5, and PhP-GisoH, 6, via the

addition of the relevant lithium phosphide to ArNQCQNAr

(Scheme 1). Aqueous work-ups of these compounds were

performed under an inert atmosphere to prevent oxidation

of the phosphorus atom. However, we have found that the

products can be handled in moist air as solids or in solution

without significant oxidation occurring, as judged by 31P

NMR spectroscopy. It is noteworthy that the addition of

phosphines to smaller carbodiimides to form phospha-

guanidines, in the presence of catalytic amounts of s-block

amide or alkyl bases, has recently been reported.22,23

Although the addition of metal amides to ArNQCQNAr is

a versatile route to bulky guanidine compounds, its limitations

centre on the bulk of the reacting amide complex (as men-

tioned above). Because of this, a different approach was used

Scheme 1 Reagents and conditions: (i) ArNQCQNAr, THF, 20 1C

or reflux; (ii) H2O.


to synthesise the exceedingly bulky guanidine, ArSi–GisoH 7

(Scheme 2). This involved lithiation of the known guanidine,

ArNC{N(H)Ar}2, the product of which was subsequently

quenched with Me3SiCl in THF at reflux to give 7 in

good yield.

One further bifunctional guanidine has been prepared in this

study via a route not involving carbodiimide addition. Though

this synthesis was originally not intended, it is moderately

yielding, reproducible and thus is included here. In an attempt

to form a Mn(II) complex of Priso�, K[Priso] was reacted with

commercially available MnI2 in THF. This, instead led to the

isolation of the aryl-coupled guanidine, (PrisoH)2 8, in a 30%

yield (Scheme 3) without aqueous work-up. When the reaction

was repeated with a pure sample of [MnI2(THF)3], compound

8 was not obtained. Presumably, the commercially sourced

MnI2 initially employed, was contaminated with significant

amounts of higher oxidation state manganese species. It is

believed that the reaction of the impure MnI2 with K[Priso] led

to the oxidative coupling of two Priso� anions through aryl

para-positions on each. This seems reasonable in light of the

fact that we have recently shown that Priso� can coordinate

the Rh(COD) fragment (COD = 1,5-cyclooctadiene) solely

through one aryl substituent in a Z5-cyclohexadienyl fashion.15

A Mn(4II)-Priso complex in which the ligand exhibits this

cyclohexadienyl binding mode can easily be envisaged as an

intermediate in the oxidative coupling that gave 8. The possi-

bility that 8 was alternatively formed via the oxidative cou-

pling of two Priso� anions by a diiodine contaminant in the

impure sample of MnI2 was examined and discounted.

(ii) Structural and spectroscopic properties of prepared

compounds

The crystal structures of all compounds 1–8 have been deter-

mined (see Fig. 1–6 for the molecular structures of 1, 4–8;

those of 2 and 3 can be found in ESIw). The compounds

display solid state structures comparable to those of previously

characterised guanidines and phosphaguanidines.1,21 Each of

the guanidines, 1–3, possesses a close to planar backbone

amino (–NR2) fragment which in no case is co-planar with

the CN3 core of the molecule. Therefore, any interaction of the

amino N-lone pair within the p-system of the largely localised

guanidine CN3 backbone must be limited. It is noteworthy

that the –NR2 fragments of 4 are significantly more distorted

from planar than those of the monofunctional guanidines.

Similarly, the two phosphaguanidines display distorted pyra-

midal phosphorus centres, the lone pairs of which are direc-

tional and therefore cannot be involved with the p-system of

their localised CN2P cores. The bond lengths and angles

within these core fragments (see Table 1) are consistent with

these descriptions.

Several different isomeric forms of the compounds have

been identified in this study. To allow comparisons with

related amidines, the backbone unit (R2N or R2P) has been

defined as the lower priority in determining the stereo-configu-

ration of the compounds (see refs. 1d and 1e for a description

of the four isomeric and tautomeric forms of amidines, viz.

Z-anti, Z-syn, E-anti and E-syn). The guanidines, 1–3 (see

Fig. 1 for the structure of 1), and the phosphaguanidine, 5

(Fig. 3), exist in the Z-anti-form which is common for guani-

dines but not for uncoordinated phosphaguanidines which

normally occur in the solid state in their E-syn-form.1d,21

Indeed, this is the isomer adopted by the phosphaguanidine,

6, in the solid state (Fig. 4). In contrast, the extremely bulky

guanidine, 7 (Fig. 5), crystallises in the rarely observed Z-syn-

form, probably because of steric buttressing of its aryl groups

by the larger N(Ar)SiMe3 substituent. It is of note that the

Z-syn-isomer of amidines with very bulky backbone C-substitu-

ents have been previously reported, e.g. (tript)C{N(H)R}(NR)

(tript = 9-triptycenyl, R = Cy or Pri).24 Both the bifunctional

amidines, 4 and 8 (Fig. 2 and 6, respectively), exist in the

solid state as Z-anti-,Z-anti-isomers, as has been previously

documented for bifunctional amidines.1

Often, amidines and guanidines will be present in solution in

more than one of their four possible isomeric forms. This can

Scheme 2 Reagents and conditions: (i) BunLi, THF; (ii) Me3SiCl,

THF reflux.

Scheme 3 Reagents and conditions: (i) K[N(SiMe3)2], THF; (ii) MnI2,

THF.

Fig. 1 Molecular structure of 1 (25% thermal ellipsoids are shown;

hydrogen atoms, except H(1), omitted for sake of clarity).


lead to complicated NMR spectra for such compounds.

However, the guanidines and phosphaguanidines, 1–6, display

relatively simple 1H and 13C{1H} NMR spectra, which are

suggestive of only one, or predominantly one, isomer occur-

ring in solution. These spectra imply that each compound has

two chemically inequivalent Ar substituents, and that both

alkyl or aryl groups on the backbone –ER2 (E = N or P)

groups are equivalent. If the compounds retain their solid state

isomeric forms in solution, which seems likely, the latter

observation requires their –ER2 groups to partially rotate on

the NMR timescale, thus leading to the compounds possessing

averaged mirror planes incorporating their ECN2 fragments.

Although the isomeric forms adopted by the guanidines,

1–4, in solution cannot be certain without two-dimensional

NMR experiments, some insight into the solution conforma-

tions of the phosphaguanidines, 5 and 6, can be gained from

their 1H NMR spectra. That for 5 shows only one isomer, the

NH resonance of which exists as a doublet (3JPH = 14.1 Hz;31P{1H} NMR: d �2.9 ppm). The spectrum of 6 reveals the

compound to exist as two isomers in solution in an approxi-

mately 90 : 10 ratio. The NH resonance of the major isomer

(31P{1H} NMR: d �18.5 ppm) is a singlet, while that for

the minor isomer (31P{1H} NMR: d �13.3 ppm) is a doublet

(3JPH = 18.2 Hz). In an excellent paper on phosphaguanidinate

solution behaviour, Coles et al. have shown that isomer

interconversion can readily occur by one or more of a number

of possible pathways.21e Importantly, they also showed that

the closely related phosphaguanidine, Cy2PC{N(H)Pri}(NPri),

Fig. 2 Molecular structure of 4 (25% thermal ellipsoids are shown; hydrogen atoms, except H(1), omitted for sake of clarity). Symmetry

operation:0 �x + 1, �y + 1, �z.






is present in solution in both its E-syn- (major) and Z-anti-

(minor) forms (14 : 1 ratio at 298 K). The NH resonance of the

E-syn-form shows no coupling to the P-centre, while that of

the minor Z-anti-isomer does (3JPH = 14.5 Hz). Accordingly,

we conclude that compound 5 exists solely as its Z-anti-form

in solution (as in the solid state), whereas the major solution

state isomeric form of 6 is E-syn (as in the solid state), and the

minor form is Z-anti.

Many of the signals in the 1H NMR spectrum of ArSi-

GisoH, 7, are very broad and suggest one or more dynamic

processes are occurring in solution. Despite efforts, the spec-

trum could not be resolved, and thus we could not shed light

on the nature of the dynamic behaviour. One possibility,

however, is that it involves a restricted rotation of the Ar

and/or SiMe3 groups about the N–C or N–Si bonds of 7. In

this respect, it should be noted that similar solution dynamic

behaviour has been observed for the closely related com-

pound, ArNC{N(H)Ar}2, an exhaustive variable-temperature

NMR study of which showed this behaviour to be derived

from restricted rotation of its three Ar groups.19 Another

possibility for 7 is that there is a fluxional interconversion

between two or more isomers of the compound, which is

occurring at close to the NMR timescale. This seems less

likely, however, when the imposing sterics of the compound

are taken into account.

Little information could be gained from the solution NMR

spectra of the bifunctional guanidine, 8. These are very

complicated and point towards more than one isomer existing


hydrogen atoms, except H(3), omitted for sake of clarity). Selected

bond lengths (A) and angles (1): Si(1)–N(1) 1.7762(16), N(1)–C(1)

1.410(2), C(1)–N(2) 1.285(2), C(1)–N(3) 1.383(2); N(2)–C(1)–N(3)

130.87(17), N(2)–C(1)–N(1) 116.38(16), N(3)–C(1)–N(1) 112.73(16),

C(1)–N(1)–C(5) 119.99(15), C(1)–N(1)–Si(1) 119.94(12),

C(5)–N(1)–Si(1) 119.81(12).

Fig. 6 Molecular structure of 8 (25% thermal ellipsoids are shown; hydrogen atoms, except H(1) and H(6), omitted for sake of clarity). Selected

bond lengths (A) and angles (1): N(1)–C(1) 1.391(3), C(1)–N(3) 1.290(4), C(1)–N(2) 1.378(4), N(4)–C(44) 1.290(3), N(5)–C(44) 1.383(4),

N(6)–C(44) 1.387(4); N(3)–C(1)–N(2) 121.1(2), N(3)–C(1)–N(1) 122.2(3), N(2)–C(1)–N(1) 116.7(3), C(1)–N(2)–C(17) 120.2(2), C(1)–N(2)–C(14)

119.8(2), C(17)–N(2)–C(14) 115.5(2), C(44)–N(5)–C(45) 119.5(2), C(44)–N(5)–C(48) 119.9(2), C(45)–N(5)–C(48) 116.0(2), N(4)–C(44)–N(5)

120.8(3), N(4)–C(44)–N(6) 122.0(3), N(5)–C(44)–N(6) 117.3(2).

Table 1 Selected bond lengths (A) and angles (1) for 1–6 (E = N or P)

1 2 3 4 5 6

ArNQC 1.290(2) 1.2911(16) 1.285(2) 1.287(2) 1.2909(19) 1.311(2)ArN–C 1.384(3) 1.3910(16) 1.394(2) 1.373(2) 1.375(2) 1.346(2)C–ER2 1.388(2) 1.3807(16) 1.385(2) 1.398(2) 1.8708(17) 1.8798(18)

ArN–CQN 121.26(17) 122.02(11) 124.10(17) 124.67(15) 123.26(14) 121.51(16)R2E–CQN 121.66(18) 120.57(11) 119.91(16) 119.84(15) 121.76(11) 119.98(13)R2E–C–N 117.08(17) 117.42(10) 115.99(16) 115.48(14) 114.96(11) 118.51(13)P

angles about E 353.3 357.0 353.5 341.9 302.3 304.4


in solution. For example, several overlapping N–H resonances

were seen in its 1H NMR spectrum, where only one would be

expected if it retained its solid state Z-anti-, Z-anti-isomeric

form in solution. As a result, the spectra proved difficult to

assign.

(iii) Metallation of bulky guanidines and phosphaguanidines

The guanidines and phosphaguanidines prepared here (with

the exception of ArSi-Giso 7), can be easily deprotonated

by standard metallation procedures. The reactions of these

ligands with one equivalent of BunLi or K[N(SiMe3)2] proceed

rapidly and near quantitatively in common solvents such as

hexane, toluene, THF or diethyl ether at ambient temperature

or below. The solvent and metal involved in the reaction can

have a striking bearing on the nuclearity of the formed

complex, and the conformation adopted by the guanidinate

or phosphaguanidinate ligand. This is important as it can

influence the product obtained from, for example, further salt

metathesis reactions of these alkali metal complexes with other

metal halides. In this study, we have structurally and spectro-

scopically characterised four lithium or potassium salts of the

ligands prepared above.

The lithiation of GisoH, 1, with BunLi in hexane led to the

solvent free dimeric complex, 9, whilst in THF the monomeric

solvated complex, 10, was formed (Scheme 4). In contrast,

metallation of PrisoH, 2, with K[N(SiMe3)2] in toluene

afforded the polymeric, solvent free complex, 11, whereas

metallation of Pip(GisoH)2 with the same reagent in THF

gave the solvated complex, 12 (Scheme 4). The NMR spectro-

scopic data for 9–11 are more symmetrical than their solid

state structures (vide infra) would suggest and imply that

fluxional processes are occurring in solution that are rapid

on the NMR timescale. This is not uncommon for alkali-metal

amidinates and guanidinates,1 and therefore no efforts were

made to investigate these dynamic behaviours by variable

temperature NMR studies. Once crystallised from the reaction

mixture, compound 12 has negligible solubility in normal

deuterated solvents (including D8-THF) and therefore no

meaningful NMR spectroscopic data could be obtained for

this compound.

The molecular structure of 9 is depicted in Fig. 7 and shows

it to be dimeric with two different lithium coordination

environments. Li(1) is coordinated by two chelating Giso�

ligands that have largely localised N(1)–C(1)–N(2) fragments.

The Li(1)–N bond lengths of 2.072(2) A (to N(2) and N(2)0)

and 2.240(5) A (to N(1) and N(1)0), although different, lie

within the normal range for amidinate and guanidinate N–Li

interactions.25 The two more distant N-atoms (N(1) and

N(1)0) also coordinate the bent two-coordinate Li(2) centre

with short interactions (1.954(4) A). The coordination sphere

of the both Li atoms is completed by agostic interactions to

ligand hydrogen atoms; Li(1) has two such interactions (both ca.

2.23 A), whereas Li(2) has four (from ca. 2.03 A to ca. 2.27 A).

When these close contacts are taken into account, both

Li-centres can be thought of as having heavily distorted

octahedral geometries. A survey of the Cambridge Crys-

tallographic Database revealed two similar dimeric lithium ami-

dinates, [Li{k2-N,N0-(SiMe3)NC(R)N(SiMe3)}2{Li(OEt2)}]

(R = C6H5CF3-4 or C6H5F-2),26 though the non-chelated

Li centre of both is further coordinated by an ether molecule.

Scheme 4 Reagents and conditions: (i) BunLi, hexane (Cy = cyclo-

hexyl); (ii) BunLi, THF; (iii) K[N(SiMe3)2], toluene; (iv) K[N(SiMe3)2],

THF.


hydrogen atoms omitted for sake of clarity). Selected bond lengths (A)

and angles (1): N(1)–C(1) 1.394(3), N(2)–C(1) 1.323(3), N(3)–C(1)

1.409(3), N(1)–Li(2) 1.954(4), N(1)–Li(1) 2.240(5), N(2)–Li(1)

2.072(2); N(2)–C(1)–N(1) 114.3(3), N(2)–Li(1)–N(1) 63.78(12),

N(2)0–Li(1)–N(1) 119.6(2), N(1)0–Li(2)–N(1) 121.7(4). Symmetry

operation:0 �x, y, �z + 1/2.


The molecular structure of monomeric 10 is shown in Fig. 8.

The localised guanidinate ligand is acting as an amide that

coordinates the Li atom in an Z1-fashion through N(2).

In addition, there is an approximately Z3-interaction of the

Li-centre with the Ar-substituent of N(1). The coordi-

nation sphere on the Li(1) is completed by one THF molecule.

A similar coordination mode (but minus the coordi-

nated THF) has been reported for the thallium(I) complex,

[Tl(Z1-N,Z3-Ar-Giso)].8

Like the structure of 10, the guanidinate moieties of the

potassium complexes, 11 and 12 (Fig. 9 and 10, respectively),

adopt the Z-anti-configuration but with more localised

coordinated NCN fragments. In addition, the arene-K inter-

actions in both are close to Z6-, as opposed to the Z3-Ar-Li

coordination seen in 10. In 11, this leads to a one-dimensional

polymeric structure in which one Ar-group of each ligand

bridges two K-centres. Compound 12 is monomeric, and in

addition to arene and N-attachments, one K-centre is coordi-

nated by two THF molecules, while the other is ligated by

three. All the distances to the K-centres in both complexes are

in the normal range.25

Conclusion

In conclusion, the synthesis and characterisation of a variety

of guanidine, bifunctional guanidine and phosphaguanidine

compounds, all bearing 2,6-diisopropylphenyl N-substituents,

have been described. In the solid state, the Z-anti-isomeric

form is observed for all guanidines, except in one extremely

bulky example, ArSi-GisoH 7. The sterics of this necessitate

it occurring as the rarely observed Z-syn-isomer. Of the

phosphaguanidinates, the bulkier example, CyP-GisoH 5,

crystallises in the Z-anti-form, while PhP-GisoH, 6, adopts

the E-syn-conformation. In solution, most of the described

compounds appear to retain their stereochemistry, though in

some cases isomer mixtures were observed. Several of the

prepared compounds have been deprotonated with alkali

metal reagents and the resulting salts crystallographically

characterised. In the case of the deprotonation of GisoH 1

with BunLi, the nuclearity and guanidinate coordination mode

displayed by the formed complexes are dependent upon the

reaction solvent employed. We are currently systematically

exploring the use of bulky guanidinates and phosphaguanidi-

nates, prepared from the neutral compounds 1–8, for the

stabilisation of low oxidation metallacycles incorporating

metals from all blocks of the periodic table.

Experimental

General considerations

All manipulations were performed under an inert atmosphere

(dinitrogen or argon) using Schlenk or glove box techniques.

Aqueous organic work-ups were carried out in air, except those

for the phosphaguanidines, 5 and 6. Melting points were

determined in sealed capillaries under a dinitrogen atmosphere,

except those for the guanidines, 1–4, which were determined in

open capillaries. Reaction solvents were dried over potassium

or Na/K alloy prior to use, except dichloromethane and chloro-

form which were used as received. Mass spectra were recorded

at the EPSRC National Mass Spectrometric Service, Swansea

University. Microanalyses were obtained from either Medac

Ltd or Campbell Microanalytical, Ottago. IR spectra were



and angles (1): K(1)–N(1) 2.755(3), K(1)–Ar centroid 3.077(1), K(1)0–Ar

centroid 2.945(1), C(1)–N(2) 1.329(5), C(1)–N(3) 1.402(5), N(1)–C(1)

1.340(5); N(2)–C(1)–N(1) 121.7(3), N(2)–C(1)–N(3) 115.1(3),

N(1)–C(1)–N(3) 123.2(3), C(1)–N(3)–C(26) 122.0(3), C(1)–N(3)–C(29)

121.5(3), C(26)–N(3)–C(29) 114.7(3), C(1)–N(1)–K(1) 128.2(2). Sym-

metry operation:0 x � 1/2, �y + 1/2, �z.



and angles (1): N(1)–C(1) 1.3149(16), C(1)–N(2) 1.3587(16), C(1)–N(3)

1.4092(16), Li(1)–N(2) 1.943(3), Li(1)–C(2) 2.290(3), Li(1)–C(3)

2.458(3), Li(1)–C(7) 2.591(3), O(1)–Li(1) 1.889(3); N(1)–C(1)–N(2)

121.56(11), N(1)–C(1)–N(3) 117.46(11), N(2)–C(1)–N(3) 120.98(11),

C(1)–N(3)–C(32) 117.05(10), C(1)–N(3)–C(26) 120.82(10),

C(32)–N(3)–C(26) 115.51(10), C(1)–N(2)–Li(1) 117.59(11).


recorded using a Nicolet 510 FT-IR spectrometer as Nujol

mulls between NaCl plates. 1H and 13C{1H} NMR spectra were

recorded on either Bruker DXP400, Bruker DPX300, Jeol

Eclipse 300 or Bruker WM250 spectrometers and were refer-

enced to the resonances of the solvent used. 31P{1H} NMR

spectra were recorded on a Jeol Eclipse 300 spectrometer and

were referenced to external 85% H3PO4. Cy2NH, Pri2NH,

cis-2,6-dimethylpiperidine and piperazine were obtained com-

mercially, dried over molecular sieves, and distilled under

dinitrogen prior to use. K[N(SiMe3)2] was prepared by treating

(SiMe3)2NH with KH in toluene at 20 1C. ArNQCQNAr27

and ArNC{N(H)Ar}219 were synthesised according to literature

procedures. All other reagents were obtained from commercial

sources and used as received.

Preparation of GisoH 1

BunLi (5.33 cm3 of a 1.6 M solution in hexanes, 8.52 mmol)

was added to a solution of Cy2NH (1.58 g, 1.73 cm3,

8.69 mmol) in THF (40 cm3) at 20 1C over 5 min and the

resultant solution stirred for 1 h. ArNQCQNAr (3.00 g,

8.27 mmol) was then added, the suspension stirred for

15 min, followed by heating at reflux for 1.5 h (or alternatively

stirred at room temperature for 4 h). All volatiles were

removed under reduced pressure and diethyl ether (40 cm3)

and H2O (10 cm3) added to the residue. The mixture was

stirred for 30 min to give two clear solution phases. The

organic phase was separated and the aqueous layer was

extracted with CH2Cl2 (3 � 30 cm3). The combined organic

phases were dried (MgSO4), filtered, and volatiles evaporated

from the filtrate under vacuum. The oily residue solidified

upon standing to give 1 as colourless crystals (yield 4.40 g,

98%). The product can be recrystallised from hot hexane

(yield 80%); mp 140–141 1C. 1H NMR (300 MHz, 298 K,

CDCl3): d 0.91 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 1.08–0.90

(m, 8 H, CH2), 1.21 (d, J=6.8 Hz, 6 H, CH(CH3)2), 1.36 (d, J=

6.8 Hz, 6 H, CH(CH3)2), 1.38 (d, J = 6.8 Hz, 6 H,

CH(CH3)2), 1.47–1.70 (m, 8 H, CH2), 2.05 (m, 4 H,

CH2CHN), 2.97 (tt, J = 11.7, 3.3 Hz, 2 H, CHN), 3.22 (sept,

J = 6.8 Hz, 2 H, CH(CH3)2), 3.32 (sept, J = 6.8 Hz, 2 H,

CH(CH3)2), 4.95 (s, 1 H, NH), 6.89–7.17 (m, 6 H, ArH); 1H

NMR (250 MHz, 298 K, C6D6): d 0.96 (d, J = 6.8 Hz, 6 H,

CH(CH3)2), 1.12–1.32 (m, 6 H, CH2), 1.44 (d, J= 6.8 Hz, 6 H,

CH(CH3)2), 1.50–1.68 (m, 2 H, CH2), 1.54 (d, J = 6.8 Hz, 12

H, CH(CH3)2), 1.75–1.92 (m, 8 H, CH2), 2.21–2.44 (m, 4 H,

CH2), 3.23 (tt, J = 11.7, 3.3 Hz, 2 H, CHN), 3.60 (sept, J =

6.8 Hz, 4 H, CH(CH3)2), 5.32 (s, 1 H, NH), 7.06–7.44 (m, 6 H,

ArH); 13C{1H} NMR (75.5 MHz, 298 K, CDCl3): d 21.6

(CH(CH3)2), 22.5 (CH(CH3)2), 24.9 (CH2), 26.0 (CH(CH3)2),

26.1 (CH(CH3)2), 27.1 (CH(CH3)2), 28.6 (CH(CH3)2), 29.0

(CH2), 32.6 (CH2), 58.0 (HCN), 121.6, 122.8, 123.5, 126.9,

135.9, 140.0, 145.5, 145.6, (ArC), 148.0 (CN3),13C{1H} NMR

(75.5 MHz, 298 K, C6D6): d 21.7 (CH(CH3)2), 22.3

(CH(CH3)2), 25.2 (CH2), 26.2 (CH(CH3)2), 27.3 (CH(CH3)2),

28.7 (CH(CH3)2), 29.3 (CH(CH3)2), 32.9 (CH2), 39.8 (CH2),

58.2 (HCN), 122.6, 123.3, 123.7, 127.1, 136.1, 139.9, 145.5,

145.6 (ArC), 148.5 (CN3); IR (Nujol): n/cm�1 = 3384 (m),

1614 (s), 1583 (s), 1259 (m), 1163 (m), 1110 (m), 1072 (m),

986 (m), 954 (w), 894 (m), 799 (m), 761 (m), 700 (w);

MS/APCI: m/z (%) = 544.7 (MH+, 100).

Preparation of PrisoH 2

A procedure analogous to that used to prepare 1 was em-

ployed, but using Pri2NH (colourless crystals: crude yield

99%; ca. 90% after recrystallisation); mp 144–145 1C; 1H

NMR (250 MHz, 298 K, CDCl3): d 0.89 (d, J = 6.8 Hz, 6 H,

CH(CH3)2), 1.12 (overlapping d, J = 6.8 Hz, 18 H,

CH(CH3)2), 1.20 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 1.37 (d,

J = 6.8 Hz, 6 H, CH(CH3)2), 3.17 (sept, J = 6.8 Hz, 2 H,

CH(CH3)2), 3.25 (sept, J = 6.8 Hz, 2 H, CH(CH3)2), 3.49

(sept, J = 6.8 Hz, 2 H, CH(CH3)2), 4.80 (s, 1 H, NH),

6.80–7.18 (m, 6 H, ArH); 1H NMR (400 MHz, 298 K,

C6D6): d 0.98 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 1.36 (over-

lapping d, J = 6.8 Hz, 18 H, CH(CH3)2), 1.51 (d, J = 6.8 Hz,

6 H, CH(CH3)2), 1.54 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 3.68

(sept, J = 6.8 Hz, 6 H, CH(CH3)2), 5.25 (s, 1 H, NH),

7.09–7.39 (m, 6 H, ArH), 13C{1H} NMR (75.5 MHz, 298 K,

CDCl3): d 21.6 (CH(CH3)2), 21.9 (NCH(CH3)2), 22.7

(CH(CH3)2), 24.7 (CH(CH3)2), 25.7 (CH(CH3)2), 28.2

Fig. 10 Molecular structure of 12 (25% thermal ellipsoids are shown; hydrogen atoms and isopropyl groups omitted for sake of clarity). Selected

bond lengths (A) and angles (1): K(1)–O(1) 2.681(3), K(1)–O(2) 2.698(3), K(1)–O(5) 2.780(3), K(1)–N(1) 2.823(2), K(2)–O(3) 2.646(3), K(2)–O(4)

2.710(3), K(2)–N(5) 2.735(3), K(1)–Ar centroid 3.007(1), K(2)–Ar centroid 2.915(1), N(1)–C(1) 1.328(4), N(2)–C(1) 1.321(4), N(3)–C(1) 1.437(4),

N(4)–C(30) 1.426(4), N(5)–C(30) 1.336(4), N(6)–C(30) 1.322(4); N(2)–C(1)–N(1) 124.6(3), N(2)–C(1)–N(3) 114.8(3), N(1)–C(1)–N(3) 120.6(2),

C(1)–N(1)–K(1) 123.88(18), C(30)–N(5)–K(2) 129.42(18), N(6)–C(30)–N(5) 123.6(3), N(6)–C(30)–N(4) 115.2(2), N(5)–C(30)–N(4) 121.2(2).


(CH(CH3)2), 28.8 (CH(CH3)2), 47.7 (HCN), 121.7, 122.7,

123.6, 126.9, 135.4, 139.8, 145.3, 145.8 (ArC), 148.3 (CN3);13C{1H} NMR (100.6 MHz, 298 K, C6D6): d 22.2 (CH(CH3)2),

22.5 (NCH(CH3)2), 23.1 (CH(CH3)2), 25.5 (CH(CH3)2), 26.0

(CH(CH3)2), 28.8 (CH(CH3)2), 29.7 (CH(CH3)2), 48.3 (HCN),

123.2, 123.8, 124.3, 127.7, 136.1, 140.3, 145.9, 146.2 (ArC),

149.1 (CN3); IR (Nujol): n/cm�1 = 3364 (m), 1608 (s), 1580

(s), 1303 (m), 1245 (m), 1184 (m), 1154 (m), 1109 (m), 1046

(m), 1002 (m), 932 (m), 828 (m), 798 (m), 767 (m), 714 (m);

MS/APCI: m/z (%) = 464.4 (MH+, 100).

Preparation of PipisoH 3

A procedure analogous to that used to prepare 1 was em-

ployed, but using cis-2,6-dimethylpiperidine (colourless crys-

tals: crude yield 98%; ca. 88% after recrystallisation); mp

128–130 1C. 1H NMR (400 MHz, 298 K, CDCl3): d 0.89 (d,

J E 6.1 Hz, 6 H, NCH(CH3)), 1.11 (d, J = 6.8 Hz, 6 H,

CH(CH3)2), 1.16 (d, J = 6.8 Hz, 12 H, CH(CH3)2), 1.18–1.76

(m, 6 H, CH2), 1.26 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 3.12

(sept, J = 6.8 Hz, 2 H, CH(CH3)2), 3.16 (sept, J = 6.8 Hz, 2

H, CH(CH3)2), 3.70 (mc, 2 H, NCH(CH3)), 4.85 (s, 1 H, NH),

6.88–7.18 (m, 6 H, ArH); 1H NMR (400MHz, 298 K, C6D6): d1.01 (d, J E 6.0 Hz, 6 H, NCH(CH3)), 1.28–1.73 (m, 6 H,

CH2), 1.42 (d, J = 6.8 Hz, 6 H, CH(CH3)2), 1.49 (d, J = 6.8

Hz, 6 H, CH(CH3)2), 1.54 (d, J = 6.8 Hz, 6 H, CH(CH3)2),

1.56 (d, J= 6.8 Hz, 6 H, CH(CH3)2), 3.55 (sept, J= 6.8 Hz, 4

H, CH(CH3)2), 4.11 (mc, 2 H, NCH(CH3)), 5.34 (s, 1 H, NH),

7.11–7.39 (m, 6 H, ArH); 13C{1H} NMR (75.5 MHz, 298 K,

CDCl3): d 14.4 (CH2), 20.8 (NCH(CH3)), 21.5 (CH(CH3)2),

22.9 (CH(CH3)2), 24.1 (CH(CH3)2), 25.6 (CH(CH3)2), 28.2

(CH(CH3)2), 28.9 (CH(CH3)2), 30.0 (CH2), 48.3 (HCN), 121.8,

122.6, 123.5, 126.7, 135.1, 139.4, 145.3, 145.6 (ArC), 149.8

(CN3);13C{1H} NMR (100.6 MHz, 298 K, C6D6): d 15.0

(CH2), 21.4 (NCH(CH3)), 22.2 (CH(CH3)2), 23.4 (CH(CH3)2),

25.0 (CH(CH3)2), 25.9 (CH(CH3)2), 28.9 (CH(CH3)2), 29.8

(CH(CH3)2), 30.7 (CH2), 49.0 (HCN), 123.3, 123.7, 124.3,

127.6, 135.9, 140.0, 145.8, 146.1 (ArC), 150.7 (CN3); IR

(Nujol): n/cm�1 = 3378 (m), 1616 (s), 1579 (s), 1303 (m),

1258 (m), 1183 (m), 1145 (m), 1169 (m), 1079 (m), 1023 (m),

934 (m), 803 (m), 765 (m), 755 (m); MS/APCI: m/z (%) =

476.4 (MH+, 100).

Preparation of Pip(GisoH)2 4

BunLi (5.00 cm3 of a 1.6 M solution in hexanes, 8.00 mmol) was

added to a solution of piperazine (0.339 g, 3.94 mmol) in THF

(40 cm3) at 20 1C over 5 min and the resultant solution stirred

for 1 h. ArNQCQNAr (2.93 g, 8.08 mmol) was then added

and the mixture stirred for 30 min, before being heated at

reflux for 2 h. After cooling to ambient temperature, water

(ca. 3 cm3) was added and volatiles removed under reduced

pressure. More water (ca. 30 cm3) and CH2Cl2 (60 cm3) were

then added to the residue and the mixture vigorously stirred

until two clear solution phases were formed. The organic

phase was separated and the aqueous layer was extracted with

CH2Cl2 (3� 30 cm3). The combined organic phases were dried

(MgSO4), filtered and volatiles removed from the filtrate under

reduced pressure. The residue was recrystallised from CHCl3at –30 1C to give 4 as colourless crystals (yield: 1.88 g, 75%);

mp 196–198 1C; 1H NMR (400 MHz, 298 K, CDCl3): d 0.89

(br d, J=6.8 Hz, 12 H, CH(CH3)2), 1.00 (d, J=6.8 Hz, 12 H,

CH(CH3)2), 1.15 (br d, J = 6.8 Hz, 12 H, CH(CH3)2), 1.23 (d,

J= 6.8 Hz, 12 H, CH(CH3)2), 2.89 (br s, 8 H NCH2), 3.04 (mc

of overlapping sept., J=6.8 Hz, 8 H, CH(CH3)2), 4.93 (s, 2 H,

NH), 6.90–7.16 (m, 12 H, ArH); 13C{1H} NMR (100.6 MHz,

298 K, CDCl3): d 22.8 (CH(CH3)2), 23.3 (CH(CH3)2), 24.5

(CH(CH3)2), 25.7 (CH(CH3)2), 28.8 (CH(CH3)2), 28.9

(CH(CH3)2), 47.3 (NCH2), 123.1, 123.4, 124.3, 127.3, 134.1,

140.1, 144.4, 145.2 (ArC), 150.8 (N3C); IR (Nujol): n/cm�1 =3391 (m), 1623 (s), 1585 (m), 1261 (m), 1196 (m), 1145 (m),

1109 (m), 1041 (m), 988 (m), 935 (m), 840 (m), 799 (m), 759

(m); MS/APCI: m/z (%) = 811.4 (MH+, 100).

Preparation of CyP-GisoH 5

BunLi (4.00 cm3 of a 1.6 M solution in hexanes, 6.40 mmol) was

added to a solution of Cy2PH (1.27 g, 6.40 mmol) in THF

(20 cm3) at 0 1C over 5 min. The resultant solution was stirred

for 1 h at room temperature. A solution of ArNQCQNAr

(2.54 g, 6.28 mmol) in THF (15 cm3) was then added to the

mixture which was subsequently heated at reflux for 1.5 h.

After cooling, degassed water (1 cm3) was added, the mixture

vigorously stirred for 1 h, and all volatiles removed under

reduced pressure. The residue was extracted with warm hexane

(2 � 50 cm3). The extract was dried over MgSO4, then filtered

and concentrated to ca. 15 cm3. Slow cooling of the filtrate to

�30 1C overnight yielded colourless crystals of 5 (yield: 2.85 g,

81%); mp 150–152 1C. 1H NMR (400 MHz, 298 K, CDCl3): d0.81 (d, J= 6.7 Hz, 6 H, CH(CH3)2), 1.09 (d, J= 6.7 Hz, 6 H,

CH(CH3)2), 1.10–1.25 (m, 8 H, CH2), 1.23 (d, J= 6.7 Hz, 6 H,

CH(CH3)2), 1.25 (d, J = 6.7 Hz, 6 H, CH(CH3)2), 1.58–2.04

(m, 14 H, CHP and CH2), 3.00 (sept., J = 6.7 Hz, 2 H,

CH(CH3)2), 3.18 (sept., J = 6.7 Hz, 2 H, CH(CH3)2), 5.44 (d,

JPH = 14.1 Hz, 1 H, NH), 6.92–7.18 (m, 6 H, ArH); 13C{1H}

NMR (100.6 MHz, 298 K, CDCl3): d 22.3 (CH2), 22.4 (CH2),

25.2 (CH(CH3)2), 26.1 (CH(CH3)2), 27.0 (CH(CH3)2), 27.8

(CH(CH3)2), 27.9 (CH(CH3)2), 28.0 (CH(CH3)2), 28.1 (CH2),

28.8 (CH2), 29.1 (CH2), 29.1 (CH2), 32.2 (d, J = 20 Hz, CH2),

33.7 (d, J = 13.2 Hz, CH2), 123.3, 123.4, 123.5, 128.4, 133.9,

139.0, 145.2, 147.2 (ArC), 160.1 (d, J = 13.1 Hz, backbone

PCN2);31P{1H} NMR (121 MHz, 298 K, C6D6): d –2.9 (s); IR

(Nujol): n/cm�1 = 3354 (NH), 1620 (m), 1592 (m), 1568 (s),

1324 (m), 1259 (s), 1173 (m), 1109 (m), 1043 (m), 934 (m),

884 (m), 852 (m), 799 (s), 756 (s); MS/EI: m/z (%) = 560.4

(M+, 4), 517.4 (M+ � C3H7, 100). Accurate mass (EI), m/z:

calc. for M+: 560.4254, found: 560.4251.

Preparation of PhP-GisoH 6


was added to a solution of Ph2PH (0.85 g, 4.56 mmol) in THF

(10 cm3) at �70 1C over 5 min then warmed to room

temperature and stirred for 2 h. To the resultant red solution

was added ArNQCQNAr (1.61 g, 4.43 mmol) in THF

(10 cm3) at �70 1C. The mixture was subsequently heated

under reflux for 1.5 h. It was then cooled to room temperature

and ca. 0.3 cm3 degassed H2O was added with stirring.

Volatiles were removed in vacuo and the residue extracted

into diethyl ether (80 cm3) and filtered. The filtrate was


concentrated and stored at –30 1C to give colourless blocks of

6 (yield: 1.66 g, 68%); mp 160–162 1C. 1H NMR (400 MHz,

298 K, CDCl3): d 0.83 (4 � overlapping d, J = 6.8 Hz, 24 H,

CH3), 2.78 (sept, J= 6.8 Hz, 2 H, CH), 3.15 (sept, J= 6.8 Hz,

2 H, CH), 5.60 (s, 1 H, NH), 6.85–7.51 (m, 16 H, Ar–H);13C{1H} NMR (100.6 MHz, 298 K, CDCl3): d 21.9, 22.1, 24.4,25.4 (CH(CH3)2), 28.5, 28.8 (CH(CH3)2), 122.2, 122.8, 123.0,

123.1, 127.6, 128.0, 128.4, 129.1, 129.7, 137.4, 138.7, 145.9,

146.1, 146.4 (ArC), 155.6 (J= 16.1 Hz, PCN2);31P{1H} NMR

(121 MHz, 298 K, CDCl3,): d �18.5; MS/APCI, m/z (%): 549

(M+, 100); IR (Nujol): n/cm�1 = 1607 (s), 1579 (s), 1434 (m),

1258 (s), 1185 (m), 1099 (m), 742 (m), 693 (m); C37H45N2P

requires: C 80.99%, H 8.27%, N 5.10%, found: C 80.84%, H

8.38%, N 5.25%.

Preparation of ArSi-GisoH 7


was added to a solution of ArNC{N(H)Ar}2 (1.35 g, 2.50 mmol)

in THF (15 cm3) at room temperature over 5 min. The solution

was then stirred for 1 h. Me3SiCl (0.36 g, 2.85 mmol)

was added at room temperature and the mixture subsequently

heated at reflux for 2.5 h. All volatiles were removed under

reduced pressure and the residue was extracted into warm

hexane (60 cm3). The solution was concentrated under reduced

pressure to ca. 12 cm3 and cooled to 4 1C to obtain colourless

crystals of 7 (yield 0.96 g, 63%); mp 257–258 1C. 1H NMR

(400 MHz, 298 K, CDCl3): d �0.3 (v br s, 9 H, Si(CH3)3),

0.92–0.64 (m, 12 H, CH(CH3)2), 1.14 (d, J = 6.8 Hz, 12 H,

CH(CH3)2), 1.17 (d, J = 6.8 Hz, 12 H, CH(CH3)2), 2.81 (sept,

J = 6.8 Hz, 2 H, CH(CH3)2), 3.54–3.14 (m, 4 H, CH(CH3)2),

5.59 (s, 1 H, NH), 7.28–6.54 (m, 9 H, ArH); 13C{1H} NMR

(75.5 MHz, 298 K, CDCl3): only resonances of one aryl

substituent are resolved. Others, as well as those for the SiMe3group, are too broad to be detected. d 23.2, 26.0, 28.2,

(CH(CH3)2 and CH(CH3)2), 122.4, 126.6, 134.2, 145.5

(ArC), 146.8 (CN3); IR (Nujol): n/cm�1= 3378 (m), 1618

(s), 1578 (s), 1378 (m), 1246 (m), 1223 (m), 1107 (m), 1007 (m),

970 (m), 843 (m), 822 (m), 752 (m); MS/EI: m/z (%) = 611

(M+, 14), 596 (M+ � CH3, 5), 568 (M+ � C3H7, 15), 539

(M+ � SiMe3, 15), 496 (M+ � SiMe3 � C3H6, 18).

Preparation of (PrisoH)2 8

A solution of K[N(SiMe3)2] (0.65 g, 3.24 mmol) in THF

(10 cm3) was added to PrisoH 2 (1.50 g, 3.24 mmol) in THF

(10 cm3) at 20 1C and the mixture stirred for 1 h. A solution of

impure MnI2 (Aldrich Chemical Company, 1.00 g, 3.24 mmol)

in THF (20 cm3) was then added at �78 1C and the reaction

mixture slowly warmed to room temperature overnight. All

volatiles were removed in vacuo and the residue was extracted

with hexane (40 cm3). Filtration, concentration and slow

cooling overnight to �30 1C yielded colourless crystals of 8

(yield 0.44 g, 30%); mp 223–225 1C. 1HNMR (400MHz, 298 K,

C6D6): d 0.81–1.01 (m of overlapping br d, 18 H, CH(CH3)2),

1.05–1.47 (m of overlapping broad d, 54 H, CH(CH3)2),

3.11–3.49 (m of overlapping br sept, 12 H, CH(CH3)2),

4.74–4.91 (m, 2 H, NH), 6.88–7.29 (m, 10 H, ArH); 13C{1H}

NMR (100.6 MHz, 298 K, C6D6): d 21.83, 21.87, 22.70, 22.75,

24.73, 25.69, 28.12, 28.28, 28.81, 28.94, 31.63 (CH(CH3)2 and

CH(CH3)2), 47.60 (CHN), 121.32, 121.41, 121.89, 122.68,

123.55, 123.61, 139.69, 139.78, 140.12, 145.59, 145.78 (ArC),

148.24, 148.33 (CN3), NB: more than one isomer present. Only

major resonances reported with tentative assignments; IR

(Nujol): n/cm�1 = 3364 (m), 1611 (s), 1589 (s), 1376 (s),

1342 (m), 1261 (m), 1184 (m), 1111 (m), 999 (m), 870 (m),

802 (m), 761 (m), 715 (m); MS/EI: m/z (%) = 924.7 (M+, 14),

881.7 (M+ –C3H7, 52). Accurate mass (EI), m/z: calc. for M+:

924.7691, found: 924.7688; CHN: C62H96N6 requires: C

80.46%, H 10.45%, N 9.08%; found: C 79.77%, H 10.99%,

N 9.31%.

Preparation of [Li{Li(Giso)2}] 9


was added over 5 min to a solution of GisoH 1 (0.58 g, 1.07

mmol) in hexane (20 cm3) at 20 1C. The resultant solution was

stirred for 30 min then concentrated under reduced pressure to

ca. 8 cm3. It was then stored at 4 1C overnight to afford

colourless crystals of 9 (yield: 0.42 g, 71%); mp 190–192 1C

(melts with slow decomposition); 1H NMR (300 MHz, 298 K,

C6D6): d 0.85–1.17 (m, 16 H, CH2), 1.18 (d, J = 6.8 Hz, 24 H,

CH(CH3)2), 1.49 (d, J = 6.8 Hz, 24 H, CH(CH3)2), 1.78–1.48

(m, 16 H, CH2), 2.03 (m, 8 H, CH2CHN), 3.35 (br t, J E11 Hz, 4 H, CHN), 3.57 (br sept, J= 6.8 Hz, 8 H, CH(CH3)2),

7.26–6.94 (m, 12 H, ArH); 13C{1H} NMR (75.5 MHz, 298 K,

C6D6): d 23.7, 25.0, 26.9, 27.9, 28.7, (CH2), CH(CH3),

CH(CH3)), 35.1 (CH2), 58.6 (HCN), 120.8 (ArC), 124.0

(ArC), 141.3 (ArC), 150.2 (br, ArC), 160.6 (v br, CN3);7Li NMR (155.5 MHz, 298 K, C6D6): d 2.6 (s); IR (Nujol):

n/cm�1 = 1612 (s), 1583 (s), 1236 (s), 1156 (m), 1110 (m), 1027

(m), 933 (m), 895 (m), 792 (m), 748 (m); MS/EI: m/z (%) =

543.7 (GisoH+, 5), 500 (GisoH+ � C3H6, 62).

Preparation of [Li(THF)(Giso)] 10


was added over 5 min to a solution of GisoH 1 (1.71 g, 3.14

mmol) in THF (20 cm3) at 0 1C. The solution was then stirred

for 1 h and volatiles removed under reduced pressure. Hexane

(15 cm3) was added to the residue and the resultant solution

concentrated to ca. 6 cm3. This was filtered and cooled to

�30 1C to yield large colourless crystals of 10. Concentration

of the supernatant solution at room temperature yielded

another crop of 10 (yield 1.52 g; 78%); mp 208–210 1C. 1H

NMR (400 MHz, 298 K, C6D6): d 0.97 (mc, 4 H, THF-CH2),

1.27 (br mc, 12 H, CH(CH3)2), 1.30–1.45 (m, 6 H, CH2), 1.63

(d, 3JHH = 6.8 Hz, 12 H, CH(CH3)2), 1.71 (mc, 2 H, CH2),

1.93 (mc, 4 H, CH2), 2.06 (mc, 4 H, CH2), 2.52 (mc, 4 H, CH2),

2.63 (mc, 4 H, THF-OCH2), 3.38 (mc, 2 H, NCH), 3.85 (sept, J

= 6.8 Hz, 4 H, CH(CH3)2), 6.98 (t, J = 7.5 Hz, 2 H, p-ArH),

7.26 (d, J= 7.5 Hz, 4 H,m-ArH); 13C{1H} NMR (100.6 MHz,

298 K, C6D6): d 22.8 (CH(CH3)2), 23.4 (CH2), 25.3 (CH2), 26.3

(CH2, THF), 27.1 (CH(CH3)2), 32.3 (CH2), 56.5 (NCH), 66.7

(OCH2), 118.8, 122.2, 140.1, 149.7 (ArC), 156.6 (backbone,

CN3);7Li NMR (116.8 MHz, 298 K, C6D6): d 1.64 (s);

IR (Nujol): n/cm�1 = 3378 (m), 1618 (s), 1578 (m), 1378

(m), 1246 (s), 1228 (m), 1198 (m), 1107 (m), 1008 (m), 970 (m),

843 (m), 822 (m), 752 (m); MS/EI: m/z (%) = 543.7

(GisoH+, 100).


Preparation of [{K(Priso)}N] 11

Toluene (30 cm3) was added to a mixture of PrisoH 2 (1.02 g,

2.20 mmol) and K[N(SiMe3)2] (0.45 g, 2.26 mmol) and the

resultant suspension stirred vigorously for 4 h at room tem-

perature. All volatiles were removed under reduced pressure

and the residue washed with hexane (15 cm3). Recrystallisation

from a toluene solution at�30 1C yielded colourless crystals of

11 (yield 0.95 g, 86%); mp 4300 1C. 1H NMR (400 MHz, 298

K, C6D6): d 1.08 (d, J = 6.8 Hz, 12 H, CH(CH3)2), 1.43 (d,

J = 6.8 Hz, 12 H, CH(CH3)2), 1.51 (d, J = 6.8 Hz, 12 H,

CH(CH3)2), 3.41 (sept, J= 6.8 Hz, 2 H, CH(CH3)2), 3.68 (two

overlapping sept, J = 6.8 Hz, 4 H, CH(CH3)2), 6.76–7.18 (m,

6 H, ArH); 13C{1H} NMR (100.6 MHz, 298 K, C6D6): d 22.7

(NCH(CH3)2), 24.0 (CH(CH3)2), 24.2 (CH(CH3)2), 27.4

(CH(CH3)2), 47.0 (HCN), 117.3, 122.5, 141.1, 145.5 (ArC),

153.2 (CN3); IR (Nujol): n/cm�1 = 1613 (s), 1584 (m), 1261

(m), 1152 (m), 1098 (m), 933 (m), 779 (m); MS/EI: m/z (%) =

501.3 (M+, 3), 420.3 (M+ � K � C3H6, 100). Accurate mass

(EI), m/z: calc. for M+: 501.3480, found: 501.3484.

Preparation of [{K(THF)2}{Pip(Giso)2}{K(THF)3}] 12

A solution of K[N(SiMe3)2] (0.336 g, 1.68 mmol) in THF

(15 cm3) was added to a solution of Pip(GisoH)2 4 (0.65 g,

0.801 mmol) in THF (25 cm3) at 20 1C. The resultant mixture

was stirred for 1 h, concentrated to ca. 15 cm3 and then cooled

to �30 1C to afford colourless crystals of 12 (yield 0.46 g,

52%); mp 4300 1C; IR (Nujol): n/cm�1 = 1620 (s), 1584 (m),

1238 (m), 1195 (m), 1050 (m), 987 (m), 840 (m), 799 (m), 758

(m), 736 (m); MS/APCI: m/z (%) = 811.4 (Pip(GisoH)2H+,

100). N.B. The very low solubility of 12 in common deuterated

solvents precluded the acquisition of meaningful NMR data.

X-Ray crystallography

Crystals of 1–12 suitable for X-ray structural determination

were mounted in silicone oil. Crystallographic measurements

were made using a Nonius Kappa CCD diffractometer. The

structures were solved by direct methods and refined on F2 by

full-matrix least squares (SHELX97)28 using all unique data.

Hydrogen atoms have been included in calculated positions

Table 2 Crystal data for compounds 1–12

Compound 1 2 3 4�2CHCl3 5 6

Empirical formula C37H57N3 C31H49N3 C32H49N3 C56H80Cl6N6 C37H57N2P C37H45N2PMr 543.86 463.73 475.74 1049.96 560.82 548.72T/K 123(2) 150(2) 150(2) 150(2) 150(2) 150(2)Crystal system Monoclinic Orthorhombic Monoclinic Monoclinic Monoclinic TriclinicSpace group P21/n Pbca P21/c P21/c P21/n P�1a/A 12.265(3) 18.397(4) 19.236(4) 13.042(3) 10.960(2) 10.847(2)b/A 17.424(4) 15.542(3) 16.327(3) 12.030(2) 26.459(5) 10.942(2)c/A 15.775(3) 20.168(4) 19.536(4) 18.989(4) 12.934(3) 14.156(3)a/1 90 90 90 90 90 96.42(3)b/1 90.43(3) 90 106.01(3) 91.97(3) 112.61(3) 101.60(3)g/1 90 90 90 90 90 102.50(3)V/A3 3371.2(12) 5767(2) 5898(2) 2977.7(10) 3462.6(12) 1585.4(6)Z 4 8 8 2 4 2Dc/Mg m�3 1.072 1.068 1.072 1.171 1.076 1.149m(Mo-Ka)/mm�1 0.062 0.062 0.062 0.328 0.105 0.114F(000) 1200 2048 2096 1120 1232 592No. reflections collected 38 325 19 664 21 080 10 581 14 712 10 001No. independent reflns 7339 5343 11498 5520 7517 5420Rint 0.1167 0.0382 0.0500 0.0272 0.0299 0.0306Final R1 (I 4 2s(I)) 0.0629 0.0435 0.0603 0.0465 0.0479 0.0464Final wR2 (all data) 0.1584 0.1069 0.1527 0.1147 0.1227 0.1159

Compound 7 8�hexane 9 10 11 12�2THF

Empirical formula C40H61N3Si C68H110N6 C74H112Li2N6 C41H64LiN3O C31H48KN3 C82H132K2N6O7

Mr 612.01 1011.62 1099.58 621.89 501.82 1392.14T/K 123(2) 150(2) 123(2) 150(2) 123(2) 150(2)Crystal system Monoclinic Monoclinic Monoclinic Monoclinic Orthorhombic MonoclinicSpace group P21/c P21/c C2/c P21/c P212121 P21/na/A 35.510(7) 15.467(3) 20.707(4) 18.807(4) 11.462(2) 24.500(5)b/A 9.9877(2) 22.956(5) 12.031(2) 11.783(2) 11.997(2) 16.473(3)c/A 21.966(4) 19.774(4) 26.802(5) 18.773(4) 21.319(4) 20.494(4)a/1 90 90 90 90 90 90b/1 104.59(3) 111.87(3) 104.33(3) 113.20(3) 90 93.48(3)g/1 90 90 90 90 90 90V/A3 7456(3) 6516(2) 6469(2) 3823.8(13) 2931.6(10) 8256(3)Z 8 4 4 4 4 4Dc/Mg m�3 1.090 1.031 1.129 1.080 1.137 1.120m(Mo-Ka)/mm�1 0.093 0.059 0.064 0.063 0.204 0.168F(000) 2688 2240 2416 1368 1096 3040No. reflections collected 58 137 16 169 18 157 24 208 22 486 27 961No. independent reflns 16 069 11 404 5651 8282 5060 14 497Rint (0.0810) (0.0423) (0.1428) (0.0304) (0.1036) (0.0408)Final R1 (I 4 2s(I)) 0.0546 0.0755 0.0669 0.0501 0.0737 0.0713Final wR2 (all data) 0.1380 0.1828 0.1301 0.1236 0.1371 0.1858


(riding model) for all structures, with the exception of the

methyl hydrogens of C(21) and C(24) in the structure of 11

which were not included in the refinement. Two crystallo-

graphically independent molecules were refined in the asym-

metric units of the crystal structures of 3 and 7. No significant

geometric differences were found between the two molecules in

each structure and therefore only the metrical parameters for

one molecule from each structure are reported here. The Flack

parameter for the crystal structure of compound 11 is 0.01(6).

Crystal data, details of data collections and refinement are

given in Table 2.

Acknowledgements

We gratefully acknowledge financial support from the Aus-

tralian Research Council (fellowships for C. J. and A. S.), the

Leverhulme Trust (fellowship for A. S.), the Erasmus scheme

of the European Union (travel grant for K. L.), the Royal

Society (fellowship for G. J.), and the US Air Force Asian

Office of Aerospace Research and Development. Thanks also

go to the EPSRC Mass Spectrometry Service, Swansea.

References

1 For general references on the structure and reactivity of amidinateand guanidinate complexes, see: (a) J. Barker and M. Kilner,Coord. Chem. Rev., 1994, 133, 219; (b) F. T. Edelmann, Coord.Chem. Rev., 1994, 137, 403; (c) P. J. Bailey and S. Price, Coord.Chem. Rev., 2001, 214, 91; (d) M. P. Coles, Dalton Trans., 2006,985; (e) P. C. Junk and M. L. Cole, Chem. Commun., 2007, 1579,and references therein.

2 N. Nimitsiriwar, V. C. Gibson, E. L. Marshall, A. J. P. White,S. H. Dale and M. R. J. Elsegood, Dalton Trans., 2007, 4464.

3 S. R. Foley, Y. Zhou, G. P. A. Yap and D. S. Richeson, Inorg.Chem., 2000, 39, 924.

4 S. Dagorne, I. A. Guzei, M. P. Coles and R. F. Jordan, J. Am.Chem. Soc., 2000, 122, 274.

5 J. Barker, N. C. Blacker, P. R. Phillips, N. W. Alcock,W. Errington and M. G. H. Wallbridge, J. Chem. Soc., DaltonTrans., 1996, 431.

6 L. Bourget-Merle, M. F. Lappert and J. R. Severn, Chem. Rev.,2002, 102, 3031.

7 (a) C. Cui, H. W. Roesky, H.-G. Schmidt, M. Noltemeyer, H. Haoand F. Cimpoesu, Angew. Chem., Int. Ed., 2000, 39, 4274;(b) N. J. Hardman, B. E. Eichler and P. P. Power, Chem. Commun.,2000, 1991; (c) M. S. Hill and P. B. Hitchcock, Chem. Commun.,2004, 1818; (d) M. S. Hill, P. B. Hitchcock andR. Pongtavornpinyo, Dalton Trans., 2005, 273.

8 C. Jones, P. C. Junk, J. A. Platts, D. Rathmann and A. Stasch,Dalton Trans., 2005, 2497.

9 C. Jones, P. C. Junk, J. A. Platts and A. Stasch, J. Am. Chem. Soc.,2006, 128, 2206.

10 S. P. Green, C. Jones and A. Stasch, Inorg. Chem., 2007, 46, 11.11 G. Jin, C. Jones, P. C. Junk, A. Stasch and W. D. Woodul, New J.

Chem., 2008, 32, 835.12 S. P. Green, C. Jones and A. Stasch, Science, 2007, 318, 1754.13 S. P. Green, C. Jones, P. C. Junk, K.-A. Lippert and A. Stasch,

Chem. Commun., 2006, 3978.14 S. P. Green, C. Jones, G. Jin andA. Stasch, Inorg. Chem., 2007, 46, 8.15 C. Jones, D. P. Mills and A. Stasch, Dalton Trans., 2008,

4799.16 D. Heitmann, C. Jones, P. C. Junk, K.-A. Lippert and A. Stasch,

Dalton Trans., 2007, 187.17 C. Jones, R. P. Rose and A. Stasch, Dalton Trans., 2007, 2997.18 S. P. Green, C. Jones, K.-A. Lippert, D. P. Mills and A. Stasch,

Inorg. Chem., 2006, 45, 7242.19 R. E. Boere, R. T. Boere, T. Masuda and G. Wolmershauser, Can.

J. Chem., 2000, 78, 1613.20 Z. Lu, G. P. A. Yap and D. S. Richeson, Inorg. Chem., 1999, 38,

5788.21 See for example (a) M. P. Coles and P. B. Hitchcock, Chem.

Commun., 2002, 2794; (b) J. Grundy, M. P. Coles and P. B.Hitchcock, Dalton Trans., 2003, 2573; (c) N. E. Mansfield,M. P. Coles and P. B. Hitchcock, Dalton Trans., 2005, 2833;(d) N. E. Mansfield, M. P. Coles and P. B. Hitchcock, DaltonTrans., 2006, 2052; (e) N. E. Mansfield, J. Grundy, M. P. Coles,A. G. Avent and P. B. Hitchcock, J. Am. Chem. Soc., 2006, 128,13879.

22 W.-X. Zhang, M. Nishiura and Z. Hou, Chem. Commun., 2006,3812.

23 M. R. Crimmin, A. G. M. Barrett, M. S. Hill, P. B. Hitchcock andP. A. Procopiou, Organometallics, 2008, 27, 497.

24 R. J. Baker and C. Jones, J. Organomet. Chem., 2006, 691, 65.25 As determined from a survey of the Cambridge Crystallographic

Database, May, 2008.26 C. Knapp, E. Lork, P. G. Watson and R. Mews, Inorg. Chem.,

2002, 41, 2014.27 K. Ogawa and M. Akazawa, Jpn. Pat. Appl., JP 91-208987910517,

1993.28 G. M. Sheldrick, SHELX-97, University of Gottingen, 1997.


The hydrogen bond acidity and other descriptors for oximes

Michael H. Abraham,*aJavier Gil-Lostes,

aJ. Enrique Cometto-Muniz,

b

William S. Cain,bColin F. Poole,

cSanka N. Atapattu,

cRaymond J. Abraham

d

and Paul Leonardd

Received (in Durham, UK) 9th July 2008, Accepted 20th August 2008


DOI: 10.1039/b811688a

The solvation descriptors for cyclohexanone oxime and acetone oxime have been obtained from

measurements on water–solvent partitions, and gas–liquid chromatographic retention data. These

yield values of 0.33 and 0.37 for the Abraham hydrogen bond acidity, A, in reasonable agreement

with a value of 0.37 for cyclohexanone oxime obtained by our recent NMR method. The other

descriptors E, S, B, L and V have also been obtained for cyclohexanone oxime and acetone

oxime, and have been estimated for a number of other oximes as well. The value for A, the

overall or effective hydrogen bond acidity of the oximes is reasonably close to the 1 : 1 hydrogen

bond acidity, a2H = 0.39 to 0.46, that can be deduced from previous literature measurements on

oximes, and to the 1 : 1 hydrogen bond acidity, a2H = 0.43 for another NOH compound,

N,N-dibenzylhydroxylamine, that again can be deduced from literature measurements.

Introduction

The oximes were important derivatives of aldehydes and

ketones, often used for identification in the 19th and early

20th century. Their use as derivatives has declined, but a

number of oximes are important. Nifuroxime is a drug, and

diacetylmonooxime is a cholinesterase reactivator. In order to

predict physicochemical and biochemical properties of the

oximes, a knowledge of their Abraham descriptors1,2

(or solvation parameters) is needed. One of the key descriptors

is the overall, or effective, hydrogen bond acidity, A, in which

we were particularly interested, especially as we have recently

developed a new method for the experimental determination

of this parameter.3 In this work, we showed that the difference

(Dd) in the 1H NMR chemical shift of a protic hydrogen in

DMSO vs. CDCl3 solvent is directly related to the hydrogen

bond acidity. This correlation was valid over 54 compounds

and 72 protic hydrogens varying from cyclohexane to the OH

proton of phenol. An important advantage of the NMR

method is that it allows the determination of A values for

individual protic hydrogens in multifunctional solutes.

As we have pointed out,1 the overall or effective hydrogen

bond acidity, A, is the important type of acidity when con-

sidering processes in which a solute is in dilute solution and

surrounded by solvent molecules, or is present in the gas phase

as an isolated molecule. A related acidity is the 1 : 1 hydrogen

bond acidity, a2H, in which a solute complexes with a hydrogen

bond base in an inert solvent such as tetrachloromethane.1,4

The defining equations for a2H are eqn (1),4 where K is the 1 : 1

complexation constant for an acid against a reference base B,

eqn (2) in which logK is put on a general scale of hydrogen

bond acidity KAH, and finally eqn (3) in which KA

H is

transformed into the a2H scale. In eqn (2), LB and DB are

the fitting coefficients.

A–H + :B - A–H� � �B; K (1)

logK (for an acid against a reference base B) = LB log

KAH + DB (2)

a2H = (1.1 + KA

H)/4.636 (3)

The term (1.1 + KAH) serves to define the origin of the scale

where a2H = 0 for zero acidity, and the factor 4.636 is used

only to provide a suitable range of the scale. A number of

equations on the lines of eqn (2) were constructed for various

reference bases.

The only acid–base measurements that seem to have been

made on oximes are those of Ossart et al.,5 who measured 1 : 1

complexation constants for a number of oximes against the

base tetrahydrofuran in tetrachloromethane. The 1 : 1 com-

plexation constants, K, in units of mol�1 dm3, are in Table 1,

together with the corresponding values of a2H that we have

deduced from the LB and DB values for the base tetrahydro-

furan4 in Table 2, through eqn (2) and (3). Feuer et al.6 have

measured 1 : 1-complexation constants for the NOH com-

pound N,N-dibenzylhydroxylamine against a number of

hydrogen bond bases in tetrachloromethane, as shown in

Table 2, where we give the deduced values of a2H.

aDepartment of Chemistry, University College London, 20 GordonStreet, London, UK WC1H OAJ. E-mail: [email protected]: [email protected]

bChemosensory Perception Laboratory, Department of Surgery(Otolaryngology), University of California, San Diego, La Jolla,CA 92093-0957, USA. E-mail: [email protected]

cDepartment of Chemistry, Wayne State University, Detroit,MI 48202, USA. E-mail: [email protected]: [email protected]

dChemistry Department, The University of Liverpool, P.O. Box 147,Liverpool, UK L69 3BX. E-mail: [email protected]: [email protected]



Results

The complexation constants of Ossart et al.5 can be trans-

formed into KAH and then into a2

H values through eqn (2) and

(3). The deduced values of a2H range from 0.39 to 0.46 as

shown in Table 1. Similarly, the complexation constants of

Feuer et al.6 yield the a2H values given in Table 2. No equation

on the lines of eqn (2) has been constructed for benzene as a

reference base, and so we are left with three independent

values of a2H for N,N-dibenzylhydroxylamine. There is not

very good agreement, but we can say that the 1 : 1 hydrogen

bond acidity of N,N-dibenzylhydroxylamine is around 0.43

units. Once a2H is known, the general equation, eqn (4),7 can

be used to estimate the 1 : 1 complexation constant of the

oximes or of the hydroxylamine with any base for which the

1 : 1 hydrogen bond basicity b2H has been determined.8–11

logK = (7.354a2Hb2

H) � 1.094 (4)

Of more practical utility is the overall hydrogen bond acidity,

A, which is one of the descriptors in our linear free energy

relationships, LFERs, eqn (5) and (6).1,2

SP = c + eE + sS + aA + bB + vV (5)

SP = c + eE + sS + aA + bB + lL (6)

In eqn (5) and (6), the independent variables are solute

descriptors as follows. E is the solute excess molar refractivity

in units of (cm3 mol�1)/10, S is the solute dipolarity/polariz-

ability, A and B are the overall or summation hydrogen bond

acidity and basicity, V is the McGowan characteristic volume12 in units of (cm3 mol�1)/100 and L is the logarithm of the gas

to hexadecane partition coefficient at 25 1C. Eqn (5) is used for

transfer of solutes from one condensed phase to another, and

eqn (6) is used for processes that involve the transfer of solutes

from the gas phase to a solvent phase. The dependent variable,

SP, is a set of solute properties in a given system. For example,

SP in eqn (5) could be the water-to-octanol partition coeffi-

cient, as logPoct, and SP in eqn (6) could be a gas-to-solvent

partition coefficient or some measure of gas chromatographic

retention. The coefficients in eqn (5) and 6 are evaluated

through multiple linear regression analysis (MLRA).

The use of eqn (5) and (6) in the determination of descrip-

tors has been described in detail,2 and numerous examples are

available.13–16 In brief, equations on the lines of eqn (5) and (6)

are set up for a number of physicochemical processes, using

solutes whose descriptors are known. The SP values for the

investigated compound are then obtained by experiment for

the same processes under exactly the same conditions as used

in the calibration experiments. There are six descriptors that

are required for any compound. However, V can be calculated

from atomic and bond contributions,1,12 and E can then be

obtained by one of a variety of methods. If the refractive index

of the liquid compound at 20 1C is available, E can be

obtained directly. Otherwise E can be calculated by addition

of fragments, either by hand or by a commercial program,17 or

can be obtained from a calculated refractive index.18

Cyclohexanone oxime and acetone oxime are solids, but a

number of lower oximes are liquids whose refractive index has

been measured,19 and for which we have calculated E, see

Table 3. Also included are values of E calculated from the

ACD refractive index,18 and values of E calculated through

the PharmaAlgorithm (PHA) program.17 The ACD values are

all too low, but the PHA values show good agreement with the

experimental values. We take the PHA value of 0.58 for

cyclohexanone oxime and a value of 0.39 for acetone oxime

(slightly larger than that for butanone oxime).

This then leaves four descriptors, S, A, B and L to be

obtained by experiment. In principle, if four values of SP are

obtained in four calibrated systems, we have four unknowns

(S, A, B and L) that can be deduced from four equations. In

practice, it is much better to have a larger number of equations

and then to find the best solution of the equations by trial-and-

error, the best solution being the values of the descriptors that

provide the best fit of calculated and experimental SP values.

We used the procedure in Microsoft ‘Solver’ to obtain the best

fit descriptors. We can extend the number of equations

through eqn (7), where Ps is a water-to-solvent partition

coefficient, Ks is the corresponding gas-to-solvent partition

coefficient, and Kw is the corresponding gas-to-water partition

coefficient. In the case of a solvent such as octanol, that takes

Table 2 Values of LB and DB in eqn (2), the 1 : 1 complexationconstant, K, in tetrachloromethane and derived values of a2

H forN,N-dibenzylhydroxylamine

Base LB DB K (ref. 6) a2H

Triethylamine 1.0486 0.0517 14 0.462Diethyl ether 0.7129 �0.3206 2.3 0.444Dimethyl sulfoxide 1.2399 0.2656 11 0.372Benzene N/A N/A 0.5Tetrahydrofuran 0.8248 �0.1970

Table 3 Some experimental and calculated values of E for oximes

Oxime Z(20) V E(exptl.)aACD(calc.)

PHA(calc.)

Formaldehyde oxime 0.3650 0.37Acetaldehyde oxime 1.4264 0.5059 0.390 0.300 0.40Propanal oxime 1.4303 0.6468 0.366 0.293 0.40Butanal oxime 1.4367 0.7877 0.357 0.288 0.40Isobutanal oxime 0.7877 (0.37) 0.41Acetone oxime 0.6468 (0.39) 0.296 0.38Butanone oxime 1.4431 0.7877 0.383 0.292 0.38Pentan-2-one oxime 1.4455 0.9286 0.369 0.290 0.37Pentan-3-one oxime 1.4465 0.9286 0.375 0.290 0.37Hexan-2-one oxime 1.4470 1.0695 0.354 0.288 0.37Heptan-4-one oxime 1.4475 1.2104 0.335 0.288 0.37Cyclopentanone oxime 0.8200 (0.58) 0.59Cyclohexanone oxime 0.9609 (0.58) 0.728 0.58

a Values in parenthesis are estimated.

Table 1 Values of the 1 : 1 complexation constant, K, for someoximes against tetrahydrofuran in tetrachloromethane, and the corres-ponding values of a2

H

Oxime K (ref. 5) a2H

Acetaldehyde oxime 3.75 0.44Acetone oxime 3.51 0.43Butanone oxime 4.08 0.45Cyclohexanone oxime 2.45 0.39Acetophenone oxime 4.24 0.45Benzophenone oxime 4.49 0.46Benzaldehyde oxime (b) 4.65 0.46


up a considerable amount of water when in equilibrium with

water, both logPs and logKs refer to the water-saturated

octanol. Then eqn (7) can be applied provided that logKw as

obtained for pure water is the same for water saturated with

octanol. There is a considerable amount of experimental

evidence that logKw is indeed the same, within any realistic

experimental error, for water and octanol saturated water,20

and so eqn (7) can be applied to wet octanol as well as to

solvents that take up only very small quantities of water.

logPs = logKs � logKw (7)

If we allow the value of logKw to float, we have increased the

number of ‘descriptors’ to be determined from four to five.

However, the logPs values for the four solvents listed in

Table 4 Coefficients in the equations used to calculate descriptors for cyclohexanone oxime, and the corresponding observed and calculatedvalues (Ps is the water-to-solvent partition coefficient, Ks is the corresponding gas-to-solvent partition coefficient, Kw is the corresponding gas-to-water partition coefficient, k is the gas to stationary phase partition coefficient, tr

0 is the retention time relative to the standard)

SP

System SP c e s a b v/l Obs Calc

Water–octanol logPs 0.088 0.562 �1.054 0.034 �3.460 3.814a 0.988 1.031Water–chloroform logPs 0.327 0.157 �0.391 �3.191 �3.437 4.191a 0.821 0.944Water–hexane logPs 0.361 0.579 �1.723 �3.599 �4.764 4.344a �0.599 �0.773Water–toluene logPs 0.143 0.527 �0.720 �3.010 �4.824 4.545a 0.260 0.232Gas–waterb logKw �0.994 0.577 2.549 3.813 4.841 �0.869a 5.115 5.011Gas–octanol logKs �0.198 0.002 0.709 3.519 1.429 0.858 6.103 6.181Gas–chloroform logKs 0.116 �0.467 1.203 0.138 1.432 0.994 5.936 6.141Gas–hexane logKs 0.292 �0.169 0.000 0.000 0.000 0.979 4.516 4.423Gas–toluene logKs 0.121 �0.222 0.938 0.467 0.099 1.012 5.375 5.423Gas–waterc logKw �1.271 0.822 2.743 3.904 4.814 �0.213 5.115 4.979CW-20M log tr

0 �3.270 0.144 1.420 1.950 0.000 0.467 0.824 0.752OV-275 log tr

0 �2.822 0.355 1.650 1.797 0.325 0.341 1.106 1.133Hp-Innowax log tr

0 �2.675 0.033 1.290 1.703 �0.051 0.386 0.765 0.704DEGS log tr

0 �3.296 0.327 1.568 1.882 0.297 0.424 0.964 0.939HP-5 80 log k �1.927 �0.051 0.360 0.303 0.000 0.636 1.258 1.215100 log k �1.970 �0.022 0.329 0.243 0.000 0.573 0.916 0.869120 log k �2.008 0.000 0.305 0.200 0.000 0.518 0.613 0.570160 log k �2.552 0.050 0.229 0.145 0.000 0.389 �0.557 �0.589SPB-Octyl 80 log k �2.645 0.165 0.062 0.000 0.000 0.703 0.600 0.543100 log k �2.719 0.181 0.057 0.000 0.000 0.644 0.267 0.219120 log k �2.738 0.189 0.076 0.000 0.000 0.578 �0.016 �0.063160 log k �1.980 0.174 0.059 0.000 0.000 0.431 0.084 0.036180 log k �1.996 0.182 0.060 0.000 0.000 0.391 �0.104 �0.147200 log k �1.965 0.186 0.048 0.000 0.000 0.350 �0.250 �0.302240 log k �1.979 0.192 0.052 0.000 0.000 0.287 �0.530 �0.581Rtx-440 80 log k �2.452 �0.038 0.505 0.389 0.000 0.667 1.001 0.990100 log k �2.537 0.000 0.461 0.316 0.000 0.613 0.647 0.630120 log k �2.584 0.021 0.427 0.271 0.000 0.559 0.337 0.317160 log k �2.419 0.046 0.336 0.211 0.000 0.427 �0.168 �0.176180 log k �2.398 0.048 0.312 0.192 0.000 0.382 �0.368 �0.376200 log k �2.403 0.067 0.288 0.181 0.000 0.346 �0.549 �0.550220 log k �2.479 0.077 0.270 0.174 0.000 0.323 �0.730 �0.739240 log k �2.393 0.098 0.226 0.156 0.000 0.284 �0.842 �0.854DB-1701 160 log k �2.119 �0.007 0.553 0.575 0.000 0.409 0.238 0.331180 log k �2.078 �0.001 0.511 0.488 0.000 0.362 0.024 0.106200 log k �2.083 0.020 0.471 0.419 0.000 0.328 �0.164 �0.092220 log k �2.070 0.039 0.428 0.356 0.000 0.295 �0.333 �0.270Rxi-50 160 log k �2.104 0.124 0.592 0.283 0.000 0.390 0.264 0.279180 log k �2.110 0.145 0.536 0.258 0.000 0.352 0.059 0.062200 log k �2.118 0.160 0.486 0.250 0.000 0.319 �0.114 �0.127220 log k �2.111 0.169 0.446 0.216 0.000 0.288 �0.297 �0.296240 log k �2.093 0.181 0.402 0.192 0.000 0.259 �0.446 �0.44480 log k �2.192 0.090 0.807 0.398 0.000 0.623 1.448 1.409120 log k �2.236 0.117 0.713 0.302 0.000 0.505 0.778 0.755140 log k �2.242 0.143 0.648 0.269 0.000 0.455 0.504 0.479HP-Innowax 160 log k �2.568 0.215 1.157 1.544 0.000 0.356 0.634 0.645180 log k �2.383 0.202 0.998 1.363 0.000 0.299 0.367 0.374200 log k �2.350 0.204 0.926 1.198 0.000 0.265 0.133 0.142220 log k �2.334 0.209 0.854 1.071 0.000 0.237 �0.077 �0.067DB-225 160 log k �2.784 0.055 0.980 0.853 0.000 0.340 �0.210 �0.120180 log k �2.833 0.074 0.909 0.776 0.000 0.311 �0.354 �0.372200 log k �2.826 0.091 0.842 0.691 0.000 0.278 �0.600 �0.586220 log k �2.775 0.096 0.754 0.612 0.000 0.251 �0.731 �0.754a These coefficients are for v, the remainder are for l. b Eqn (5). c Eqn (6).


Table 4 then yield four extra logKs values, and in addition we

have two equations, one from eqn (5) and one from eqn (6) for

logKw, making an extra six equations. In Table 4 are given the

systems that we have used for cyclohexanone oxime, the coeffi-

cients in eqn (5) and (6), and the observed and calculated SP

values. The extra equations lead to a total of 53 equations for

which the SP values can be fitted with a standard deviation, SD,

of only 0.063 log units with the descriptors shown in Table 5.

For acetone oxime, we have the GLC data obtained at UCL.

We also have an equation derived from the retention indices, I,

obtained by Zenkevich21 on Porapak Q for a large number of

volatile compounds. Application of eqn (6) yielded eqn (8).

I = 154.68 � 69.354E + 38.611B + 175.622L (8)

N = 214, R2 = 0.9873, SD = 28.7, F = 2702.6

In eqn (8), N is the number of compounds, R is the correlation

coefficient, SD is the standard deviation and F is the F-statis-

tic. There is also a set of GLC data on a Perkin–Elmer column

that includes acetone oxime.22 The relevant equation is

eqn (9), making a total of 16 equations for acetone oxime.

Details of the calculations for acetone oxime are in Table 6;

the standard deviation between observed and calculated values

is only 0.040 log units.

I = 83.84 � 19.68E + 63.46S + 118.44A + 11.85B

+ 196.853L (9)

N = 48, R2 = 0.9880, SD = 13.9, F = 713.13

The 1H NMR spectra of oximes in CDCl3 and DMSO

solvents have been recorded previously. There is exchange

between the NH and OH protons in hydroxylamines in

DMSO solution which was noted by Feuer et al.6 in their

measurements of the self-association of these compounds in

this solvent. However the OH chemical shift in oximes in

DMSO solution is independent of concentration and this was

used by Kurtz and D’Silva23 in their estimation of the pKa of

twenty oximes in DMSO solvent. The 1H NMR data of ca.

forty oximes in CDCl3 solution, including acetone and cyclo-

hexanone oxime are given in the Aldrich Spectral catalogue.24

The OH proton chemical shift is always very deshielded,

for example acetone oxime 9.97 ppm, cyclohexanone oxime

9.78 ppm. Very similar shifts are obtained in DMSO solution:

10.12,23 10.14 (this work) for acetone oxime, and 10.02,23

10.05 (this work) for cyclohexanone oxime. The values

for chloroform are for relatively concentrated solutions

(8/10%, weight to volume,24 i.e. for cyclohexanone oxime

0.9 mol dm�3). The chemical shift of the OH proton in oximes

in CDCl3 solvent is known to be concentration dependent6 due

to intermolecular hydrogen bonding; thus a dilution experi-

ment was performed in CDCl3 solution on cyclohexanone

oxime to obtain the N dilution chemical shift required for

this study. The oxime concentration was decreased until

the OH chemical shift showed very little change with concen-

tration (Table 7). The concentrations were measured by

using the integral of the a-CH2 protons of the oxime with

respect to the residual CHCl3 peak. The results are given in

Table 7. The plot of d(OH) vs. concentration is linear until a

dilution of ca. 0.06 mol dm�3 is reached when the plot is

essentially independent of concentration. Thus the value of

4.45 ppm may be regarded as the N dilution chemical shift in

this experiment. However the OH peak of the oxime at the

lowest concentration measured was a broad peak of intensity

2H, with respect to the a-CH2 protons of the oxime

(see above). This value was interpreted as due to the oxime

OH (intensity 1) plus an equal amount of water protons

present despite careful drying of the CDCl3 solvent over

molecular sieves. There is rapid exchange on the NMR time

scale between the oxime OH proton and the water protons

to give the broad peak observed. The chemical shift of this

peak is therefore the weighted average of the chemical shifts of

the oxime OH and the water protons. Thus eqn (10) applies

where dobs, d and d2 are the observed chemical shift and

the chemical shifts of the oxime OH and the water protons

at these concentrations and n1 and n2 the mole fractions of the

two species.

dobs = n1d1 + n2d2 (10)

The N dilution chemical shift of water in CDCl3 solvent is

1.56 ppm25 and inserting this in eqn (10) with dobs = 4.45 ppm

and n1 = n2 = 1/2 gives the N dilution value for the OH

shift in cyclohexanone oxime as 7.34 ppm. This value,

when inserted into the A vs. Dd, eqn (3), gives an A value

of 0.37.

Table 5 Solvation descriptors for cyclohexanone and acetone oxime

Oxime E S A B V L logKw

Cyclohexanone oxime 0.58 0.90 0.33 0.61 0.9609 4.320 5.11Acetone oxime 0.39 0.66 0.37 0.56 0.6488 2.557 4.46

Table 6 Observed and calculated values for acetone oxime (seeTable 4 for definitions)

SP

System SP Obs. Calc.

Water–octanol logPs 0.120 0.154Water–chloroform logPs �0.351 �0.264Water–hexane logPs �1.725 �1.740Water–toluene logPs �0.960 �1.002Gas–watera logKw 4.464 4.472Gas–octanol logKs 4.584 4.580Gas–chloroform logKs 4.113 4.137Gas–hexane logKs 2.739 2.744Gas–toluene logKs 3.504 3.484Gas–waterb logKw 4.464 4.452CW-20M log tr

0 �0.287 �0.354OV-275 log tr

0 0.058 0.129HP-Innowax log tr

0 �0.227 �0.217DEGS log tr

0 �0.152 �0.181Porapak Q21 I/100 5.980 6.009See text22 I/100 6.700 6.748

a Eqn (5). b Eqn (6).

Table 7 d(OH) vs. concentration of cyclohexanone oxime in CDCl3

Conc. (mol dm�3 � 10�2) 2.00 6.97 9.26 11.76 20.0d(OH) 4.45 4.68 5.77 6.27 8.82


Discussion

The descriptors for cyclohexanone oxime have been derived

from fits to 53 equations and can be regarded as soundly based.

Those for acetone oxime are based on 16 equations, and so

should also be quite reliable. The value of the hydrogen bond

acidity descriptor, A, is 0.33 or 0.37 for cyclohexanone oxime

and 0.37 for acetone oxime, as compared to the 1 : 1 hydrogen

bond acidity 0.39 and 0.43, respectively, see Table 1, and 0.43

for the NOH compound, N,N-dibenzylhydroxylamine, see

Table 2. For alcohols, A and a2H do not differ too much :

0.37 and 0.32 for propan-1-ol, 0.33 and 0.33 for isopropanol,

and 0.31 and 0.32 for tert-butanol. Hence, for N,N-dibenzyl-

hydroxylamine we expect A to be near 0.43 units. The hydro-

gen-bond acidity of the two types of NOH compound, the

oximes and the hydroxylamines, are thus quite close.

The value of 0.37 for the hydrogen bond acidity of cyclo-

hexanone oxime by the NMRmethod is a little higher than the

value of 0.33 from the GLC and partition measurements.

However, the NMR method is rendered more difficult than

usual because of the large concentration dependence of the

chemical shift in CDCl3, and the necessity of obtaining the N

dilution chemical shift of the oxime from the observed shift

due to the oxime and water. For other acyclic oximes, we

suggest that an A-value of 0.35 could be taken.

In the calculation of the descriptors for the oximes, we used

the method of fitting by trial-and-error. If, for a given oxime,

we have four unknown descriptors S, A, B and L, then four

equations of the type of eqns (5) and (6) would suffice to yield

values for the four descriptors. It is obviously better to have

more equations, but then the solution can only be obtained by

trial-and-error. We used the ‘Solver’ add-on programme in

Microsoft Excel to obtain the best-fit descriptors. Inspection

of Table 4 shows that the various equations that can be used in

the calculation of descriptors have very different coefficients.

The larger the coefficient the more accurately can the corre-

sponding descriptor be obtained. Several of the GLC phases

have reasonably large values of the s- and a-coefficients,

because they are dipolar and are hydrogen bond bases and

so they are useful in the determination of the S and A

descriptors: note that the solvent hydrogen bond basicity is

complementary to the solute hydrogen bond acidity. However,

the values of the a-coefficients for the GLC phases are never

more than 2.0, whereas a number of other processes, including

partitions from water to non-polar solvents, have a-coeffi-

cients numerically almost twice as large. It is therefore an

advantage to include water-to-solvent partitions in the set of

equations when calculating S and A. Of course, since there are

no commercially available GLC stationary phases with any

significant hydrogen bond acidity (the b-coefficients are zero),

it is then absolutely essential to include other processes such as

water to solvent partitions in order to obtain the B descriptor.

For a few other oximes, water-to-octanol partition coeffi-

cients26 and retention data by Zenkevich21 are available, and

we give in Table 8 approximate values for descriptors, with A

fixed at 0.35 for the acyclic oximes, and at 0.33 for

cyclopentanone oxime.

Reversed phase HPLC systems have been used instead of

water-to-solvent systems in the calculation of descriptors,27

but this is only possible if rather unusual HPLC systems are

used. Du et al.28 and Valko et al.29 have shown that most of the

common isocratic elution and gradient elution systems have

similar coefficients, with rather small a-coefficients. Hence if

HPLC systems are used, it is preferable to include some

water-to-solvent partition systems as well as GLC systems.

Probably the best set of experimental data to use in order to

obtain all the descriptors is a combination of retention data on

GLC stationary phases and partition coefficients in a number

of water-to-solvent partition systems, as we have used here.

Experimental

Partition coefficients

Cyclohexanone oxime and acetone oxime were used as

received. Solvents were pre-equilibrated with water, and the

water saturated with the solvent and the solvent saturated with

water were used in the experiments. Dilute solutions of the

oximes in water were gently shaken with the organic solvent

and left to equilibrate at 25 1C for 30 min. Portions of the

organic layer and the aqueous layer were carefully withdrawn

using hypodermic syringes and directly injected into a Perkin-

Elmer F-33 gas chromatograph with a stationary phase of

Carbowax 20M at 101 1C. The volumes withdrawn (Vo and

Vw) were arranged so that the area under the GC peaks was

almost the same for the aqueous and organic layers. The ratio

of the areas (Ao/Aw) could then be taken as the ratio of the

quantities of oxime in the withdrawn volumes (Qo/Qw). Then

the partition coefficient, P, is given by P = (Qo/Vo)/

(Qw/Vw) = (Ao/Vo)/(Aw/Vw). The partition coefficients in each

water-to-solvent system are given in Table 9; this includes a

value for the water-to-octanol partition coefficient from the

MedChem data base.26 From the replicate measurements we

Table 8 Approximate solvation descriptors for some oximes

Oximes E S A B V L logKw

Cyclopentanone oxime 0.580 0.94 0.33 0.61 0.8200 3.700 5.23Acetaldehyde oxime 0.390 0.50 0.35 0.54 0.5059 1.931 3.98Propanal oxime 0.366 0.52 0.35 0.54 0.6468 2.498 3.92Butanal oxime 0.357 0.58 0.35 0.54 0.7877 3.149 3.96Isobutanal oxime 0.370 0.59 0.35 0.57 0.7877 2.992 4.13Butanone oxime 0.383 0.71 0.35 0.56 0.7877 3.173 4.40

Table 9 Partition coefficients for cyclohexanone oxime and acetoneoxime between water and various solvents

Solvent logP logP takenCyclohexanone oxime

Octanol 0.988 0.988Toluene 0.260 0.260Chloroform 0.805, 0.818, 0.839 0.821Hexane �0.570, �0.596, �0.630 �0.599

Acetone oxime

Octanol 0.1226

Toluene �0.980, �0.982 �0.981Chloroform �0.297 �0.297Hexane �1.784, �1.669, �1.751 �1.725

�1.738, �1.682


estimate that the standard deviation is about 0.03–0.04 log

units. In the GC experiments, a flame ionisation detector was

used; we encountered no particular problems in the analysis of

the aqueous solutions.

GLC retention data

At UCL, four GLC stationary phases were each calibrated

using 45–65 solutes of known descriptors: CW-20M at 101 1C,

DEGS at 87 1C, HP-Innowax at 100 1C and OV-275 at 89 1C.

The obtained coefficients are in Table 4, together with coeffi-

cients for all the other equations used. Cyclohexanone oxime

or acetone oxime were then injected onto a given phase

together with standard compounds as references, and reten-

tion data obtained under the same conditions as the calibra-

tion. The coefficients in Table 4 refer to log tr0, where tr

0 is the

retention time relative to the standard. The internal standards

were heptanol for CW-20M, DEGS, and HP-Innowax and

hexanol for OV-275. A number of secondary standards were

also used. At Wayne State, retention factors at 20 1C intervals

over the temperature range 60–140 or 180–240 1C were

obtained with an Agilent Technologies HP-6890 gas chromato-

graph (Palo Alto, CA, USA) fitted with a split/splitless injector

and flame ionization detector. Nitrogen was used as carrier gas

at a constant linear velocity of 40 cm s�1 and methane

was used to determine the column hold-up time. Measure-

ments were made for seven different stationary phases on

30 m � 0.25 mm I.D. open-tubular columns with a film thick-

ness of 0.25 mm for 60–140 1C and 1.00 mm for 180–240 1C.

The system constants at each temperature were determined by

calibration using 60–100 varied compounds exactly as before30

and are summarized with the retention factors for cyclo-

hexanone oxime in Table 4; k in log k is the gas to stationary

phase partition coefficient.

NMR experiments

These were conducted exactly as described before.3 All the

compounds and solvents were obtained commercially. The

CDCl3 and DMSO solvents were commercial samples (Sigma-

Aldrich). The CDCl3 was bought in 1 ml ampoules and used

directly in the experiments. Solutions of B10 mg mL�1

concentration were run with TMS as internal standard in

DMSO solvent. The 1H spectra were obtained on a Bruker

Avance 400 MHz NMR spectrometer operating at 400.13 MHz.

Typical running conditions were 128 transients, spectral width

3300 Hz and 32 K data points, giving an acquisition time of

5 s. The FIDs were zero-filled to 64 K. The spectra were first

order, and the assignments were straightforward.

Acknowledgements

This work was supported in part by Philip Morris USA, Inc.,

and Philip Morris International.

References

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2005, 7, 295–301.17 The Absolv data base, PharmaAlgorithms Inc., 591 Indian Road,

Toronto, ON, M6P 2C4, Canada.18 Advanced Chemistry Development, Inc., 90 Adelaide Street West,

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21, 823–832.21 I. G. Zenkevich, Zh. Anal. Khim. (USSR), 1998, 53, 932–945;

I. G. Zenkevich, J. Anal. Chem. USSR, 1998, 53, 816–828 (Englishtranslation).

22 Compilation of Gas Chromatographic Data ASTM Data Series NoDS 25A, eds. O. E. Schupp III and J. S. Lewis, ASTM, 1916 RaceStreet, Philadelphia, PA 19013, 1967.

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Electrochemical methodology for determination of imidazolium ionic

liquids (solids at room temperature) properties: influence of the

temperaturew

M. P. Stracke,aM. V. Migliorini,

aE. Lissner,

aH. S. Schrekker,

aD. Back,

b

E. S. Lang,bJ. Dupont*

aand R. S. Goncalves*

a

Received (in Gainesville, FL, USA) 17th July 2008, Accepted 29th August 2008


DOI: 10.1039/b812258j

A set of six imidazolium ionic liquids (1a–b, 2a–c, 3), that were solids at room temperature, were

characterized by electrical impedance spectroscopy to obtain information about their polarization

resistance (Rp), conductivity (s) and charge transfer activation energy (Ea). These experiments

were performed at different temperatures in a glass micro-cell, equipped with three platinum

electrodes. The observed conductivities were due to charge transfer processes of molecular oxygen

at the electrode surface and mass transfer processes within the IL matrix. Higher temperatures

resulted for all ionic liquids in increased conductivities. X-Ray diffraction of the ionic liquids 2a–c

suggested that a higher degree of supramolecular two-dimensional organization, higher density, is

related to an easier oxygen-electrode approximation, lower Ea. Two distinct temperatures ranges

were observed. The larger conductivity increases in the higher temperature range were explained

by melting (ILs 1–2) and fluxional behavior/reorientation phenomena of the ionic liquids and are

due to enhanced oxygen diffusion (IL 3). In general, the understanding of imidazolium ionic

liquid electrochemical properties could facilitate the development of new applications.

1. Introduction

The discovery of air- and water-stable imidazolium room-

temperature ionic liquids (RTILs) by the suitable choice of the

anion initiated intensive research efforts towards their appli-

cation.1 Further attractive physical and chemical properties of

the imidazolium RTILs include,2–6 a negligible vapor pressure;

low inflammability; thermal stability; liquidity over a wide

temperature range; easy recycling; and being a good solvent

for a wide variety of organic and inorganic chemical com-

pounds. Besides, imidazolium RTILs are ‘‘designable’’ as

structural modifications in both the cation (especially the

1 and 3 positions of the imidazolium ring) and anion permit

the tuning of properties such as, e.g., miscibility with water

and organic solvents,7 melting point and viscosity.3 This

adaptability is also responsible for the easy preparation of

task-specific imidazolium ionic liquids, ionic liquids that con-

tain a specific functionality covalently incorporated in either

the cation or anion.8–11 As a result, applications of imidazolium

RTILs are numerous and found in the fields of, for instance,

extraction and separation processes,4,12,13 synthetic chemistry,4,6

catalysis (organometallic,5,6,14,15 transition-metal nanoparticle,14–19

bio20), and materials science.4,21

Another important imidazolium RTIL research area is in

the field of electrochemistry, which is due to their chemical and

electrochemical stability, wide electrochemical windows, and

high electrical conductivities and ionic mobilities.3–6,22–24

Electrochemical applications of imidazolium RTILs as

electrolytes are found in, e.g., fuel cells,25 electrodeposition,26

capacitors,27–29 solar cells,30,31 batteries32 and water electro-

lysis for hydrogen generation.33 However, the use of imida-

zolium RTILs could suffer from sealing problems due to

leakage issues. Possible alternatives are, e.g., imidazolium RTIL

polymer homologues such as gel34 or solid35 polyelectrolytes,

and imidazolium RTILs confined in silica-derived networks

(ionogels)36 and polymers.27 Without doubt, the direct appli-

cation of imidazolium ionic liquids (ILs), that are solids at

room temperature, instead of imidazolium RTILs, would be

another attractive option. As a consequence, we were inter-

ested in the electrochemical properties of imidazolium ILs

(solids at room temperature). In general, understanding the

physicochemical properties of ILs is of great importance

to provide information about their application scope.37,38

Herein, we report the results obtained with the imidazolium

ILs 1–3, presented in Fig. 1, which can be divided in two

classes: (1) hydrophilic ILs 1a–b and 3, and (2) hydrophobic

ILs 2a–c. Electrical impedance spectroscopy (EIS), a non-

destructive technique, was used to determine their temperature

dependent polarization resistance (Rp), conductivity (s) and

charge transfer activation energy (Ea).

a Laboratory of Electrochemistry, Laboratory of Molecular Catalysisand Laboratory of Technological Processes and Catalysis, Instituteof Chemistry, Universidade Federal do Rio Grande do Sul, Av. BentoGoncalves 9500, P.O. Box 15003, CEP: 91501-970 Porto Alegre-RS,Brazil. E-mail: [email protected]; E-mail: [email protected];Fax: +55-51-3308-7304; Fax: +55-51-3308-7304;Tel: +55-51-3308-6321; Tel: +55-51-3308-7236

bDepartamento de Quımica, Laboratorio de Materiais Inorganicos,Universidade Federal de Santa Maria, CEP: 97105-900 SantaMaria-RS, Brazil

w Electronic supplementary information (ESI) available: Experimentalsection. CCDC 607218 (2a: room temperature), 607812 (2b: roomtemperature) and 671958 (2c: 100 K). For ESI and crystallographicdata in CIF or other electronic format see DOI: 10.1039/b812258j



2. Experimental

2.1 Imidazolium ionic liquids

The de-aerated imidazolium ILs 1a–b,39,40 2a–c41–44 and 345

were prepared according to known procedures, and the NMR

spectral data were in agreement with the literature data.

Recrystallizations were performed to obtain high purity ILs

as white solids at room temperature.

2.2 Electrical impedance spectroscopy

The device used to perform the electrical impedance measure-

ments of the room-temperature ionic solids consisted of a

home-made glass micro-cell (Fig. 2) with a free area of

0.65 cm2, equipped with three platinum wire electrodes. This

micro-cell was inserted in a three-way round-bottom flask

allowing the control of the gas atmosphere and humidity.

The working electrode was located at the center of the micro-

cell, the counter electrode was placed at the full length of the

inner wall, and the reference electrode was located in between

the working and counter electrodes. A computer-controlled

potentiostat Autolab PGSTAT 30 was connected to the ionic

solid in the glass-cell by the corresponding electrodes, and the

temperature were kept under control. The electrical impedance

spectra were measured over the frequency sweep range from

50 kHz to 5 Hz and the amplitude of the applied sine wave

voltage was 10 mV. The experimental data were corrected

by the software, taking into consideration the influence of

connecting cables and other parasite capacitances, to obtain

the polarization resistance (Rp) of the samples. The RP values

were obtained from the intercepts of the electrode impedance

arc on the real impedance axis and were used to calculate the

corresponding conductivities (s).

2.3 Differential scanning calorimetry

The melting points of the ILs 1a–b and 2a–c were determined

using a TA Instruments DSC 2010 differential scanning

calorimeter, equipped with a manual cooling unit. The DSC

instrument was calibrated using an indium primary standard.

An average sample weight of 7–12 mg was sealed in an

aluminium pan in a nitrogen-filled glove box. The DSC

measurements were carried out under a nitrogen atmosphere.

The melting points (Tm, determined at the maximum of the

endothermic peaks) were determined on heating in the second

heating run.

2.4 X-Ray diffraction studies

Crystallographic data were collected at room temperature

and/or �100 1C on a Bruker Kappa Apex II CCD diffracto-

meter using Mo-Ka radiation (l = 0,71073 A). The experi-

mental set-up did not allow full rotations. Hence, the data sets

are of lower coverage. Crystal structures were refined with full-

matrix least squares on F2 using all data (SHELXTL crystal

structure solution software). Non-hydrogen atoms were

refined anisotropically. Hydrogen atoms were fixed on geo-

metrically ideal positions during the refinement. The free

refinement of hydrogen atom parameters gave low data/

parameter ratios and led to high correlations. Relevant crystallo-

graphic data, and collection and refinement details, are

compiled in Table S5 of ESI.w The structures presented in

Fig. 6 were obtained from the original X-ray data using the

DIAMOND software (version 2.1c, Crystal Impact GbR,

http://www.crystalimpact.com/diamond/).

3. Results and discussion

3.1 Impedance spectrum analysis

A home-made glass micro-cell (Fig. 2), equipped with three

platinum wire electrodes, was used for the electrical impedance

spectroscopy measurements of the imidazolium ILs 1–3

(Fig. 1). Furthermore, the partial oxygen pressure of the gas

atmosphere was kept constant. The Nyquist plots of IL

[PhC3MIm][NTf2] 2b at different temperatures are presented

in Fig. 3(a). As for most of the ILs 1–3, electrical impedance

spectroscopy measurements with 2b afforded partial semi-

circles. An equivalent circuit is proposed taking into account

that there exists a semicircle corresponding to one time constant

that represents an electrochemical circuit with two resistances

and one parallel combination of phase constant (CPE, Fig. S1,

ESIw). In contrast, complete semicircles were observed with IL

[C2O2MIm][Cl] 3 at higher temperatures (Fig. 3(b)). Perfor-

mance of these electrical impedance spectroscopy measure-

ments under vacuum resulted in confuse and irreproducible

Fig. 1 Imidazolium ionic liquids (solid at room temperature) applied

in this work.

Fig. 2 Illustration of the home-made glass micro-cell: (a) top view;

(b) section.


Nyquist plots. This strongly suggested that a charge transfer

process of molecular oxygen at the platinum electrode surface

was responsible for the observed phenomena, which is repre-

sented by the equilibrium reaction of Scheme 1.

This was further supported by the Nyquist plots obtained

when the experiments were performed under a pure argon

atmosphere and a pure oxygen atmosphere (Fig. S2, ESIw).The charge transfer process was not observed under an argon

atmosphere. However, this process did take place in the

presence of a pure molecular oxygen atmosphere.

The polarization resistance (RP) values were determined by

fitting the obtained impedance semicircles. The Rp values

represent the polarization resistances related to the charge

transfer process of oxygen on the platinum electrode surface,

since the platinum electrode is inactive under the applied

conditions. For all ILs 1–3, RP decreased with increasing

temperatures. Eqn (1) was used to convert RP into the

conductivity (s) of oxygen within the ILs, where l (0.1 cm)

and A (3.14 � 10�2 cm2) represent the length and active

surface area of the working platinum electrode, respectively.

These conductivities were due to charge transfer processes and

transport phenomena of molecular oxygen and were not

related to the ionic conductivities of the bulk ILs. This strategy

was applied to determine the activation energies of the oxygen

Fig. 3 (a) Nyquist diagram of [PhC3MIm][NTf2] 2b at 19 1C (’),

26 1C (K), 42 1C (m) and 60 1C (window); (b) Nyquist diagram of

[C2O2MIm][Cl] 3 at 5 1C (’), 17 1C (K), 31 1C (m), 34 1C (E) and

42 1C (.).

Scheme 1 Charge transfer processes of molecular oxygen at the

platinum electrode surface.

Fig. 4 Conductivities of (a) IL [PhC3MIm][NTf2] 2b and (b) IL

[C2O2MIm][Cl] 3 at different temperatures.

Fig. 5 Arrhenius conductivity plot of IL [PhC3MIm][NTf2] 2b.


redox processes on the platinum electrode surface as

described below.

s = l/RPA (1)

Fig. 4(a) and (b) show the conductivities of the ILs

[PhC3MIm][NTf2] 2b and [C2O2MIm][Cl] 3 at different

temperatures. The same conductivity–temperature correlation

was observed for all ILs 1–3. Higher temperatures resulted in

higher conductivities. As such, the transport of the species

involved in the charge transfer reaction is temperature depen-

dent. Furthermore, the conductivity was characterized by two

distinct temperature dependences: (1) small conductivity in-

creases in the lower temperature range; and (2) large conducti-

vity increases in the higher temperature range. In case of the ILs

2b and 3, these dependences showed their intersection at 42 and

35 1C, respectively, which indicates that the transport processes

are differently affected below and above this temperature.

It was found that the experimental conductivity data of the

lower temperature range fitted the conventional Arrhenius

eqn (2), where Ea is the activation energy for the charge

transfer process. Fig. 5 shows the Arrhenius conductivity plot

of IL 2b and the activation energies calculated from the

Arrhenius formula are presented in Table 1. The hydrophilic

ILs showed the higher charge transfer activation energies,

which decreased in the order: [C2O2MIm][Cl] 3 4[C10MIm][Mes] 1b 4 [C4MIm][Mes] 1a 4 [Ph2C2MIm][NTf2]

2a4 [PhC2MIm][NTf2] 2c4 [PhC3MIm][NTf2] 2b. Now, it is

important to remember that these activation energies were

measured in the presence of atmospheric oxygen. The values of

51.6 kJ mol�1 (0.53 eV) to 110 kJ mol�1 (1.14 eV) are very

close to those observed for charge transfer processes of oxygen

at polycrystalline oxide surfaces,46 LSCF-SDC composite47

and multi-metallic electrodes.48 Apparently, the determined IL

charge transfer processes are due to electrochemical reactions

of molecular oxygen at the electrode surface.49

ln s = ln s0 � (Ea/RT) (2)

3.2 X-Ray diffraction studies

The charge transfer processes involving molecular oxygen

should be the same for all ILs 1–3. As a consequence, it is

reasonable to infer that the IL crystalline structure influences

the transport of molecular oxygen inside the crystal. The

crystal data concerning the ILs 2a–c are listed in Table S5 (ESIw)

and their structures at room temperatures are presented in

Fig. 6. The electrostatic interactions of IL 2a generate a

tri-dimensional structural organization (Fig. 6(a)). In contrast,

the existing interactions in the ILs 2b and 2c are of

two-dimensional nature, generating structures in the form of

layers as can be verified in Fig. 6(b) and (c). This structural

organization at room temperature is reflected by the

density of these ILs at room temperature, which decreases

in the order: [PhC3MIm][NTf2] 2b (d = 1.553 g cm�3) 4[PhC2MIm][NTf2] 2c (d = 1.539 g cm�3) 4 [Ph2C2MIm]-

[NTf2] 2a (d= 1.47 g cm�3). This suggests that a higher degree

of two-dimensional organization as in IL 2b results in a more

dense packing. However, the observed activation energies

decrease in exactly the opposite order: [Ph2C2MIm][NTf2]

2a 4 [PhC2MIm][NTf2] 2c 4 [PhC3MIm][NTf2] 2b. As a

consequence, it is possible to infer that the diffusion of

molecular oxygen is faster in two-dimensional organized ILs.

Importantly, it is not possible to verify the formation of

structures in the form of channels or tunnels.

3.3 Differential scanning calorimetry

The melting points of the ionic liquids 1–3 were determined by

differential scanning calorimetry (DSC) to check if there exists

a correlation with their temperature dependent conductivities

(Table 1). Most of these ionic liquids have melting points that

are close to their intersection temperatures as determined from

the conductivity plots (e.g. Fig. 4). This indicates that the

change from the slowly to the faster changing conductivity is

most likely due to the changeover from the solid to the liquid

state. In strong contrast, IL [C2O2MIm][Cl] 3 showed an

intersection temperature (42 1C, Fig. 4(b)) far below

its melting point (197 1C). As a consequence, IL 3 was not

melted at the beginning of the second temperature range. This

behavior allows us to infer that the faster increasing conduc-

tivity in the second temperature range of IL 3 should be

associated with the oxygen diffusion inside the crystal arrays.

A possible explanation could be an increase in fluxional

behavior/reorientation phenomena in the solid state, which

enhances the molecular oxygen diffusion.50,51 This was further

supported by the low degree of organization of IL 2c observed

at 25 1C by X-ray diffraction, and high quality data were only

obtained at 100 K due to a more defined organization.

4. Conclusions

In conclusion, electrical impedance spectroscopy is a suitable

analytical tool for the determination of important imida-

zolium IL properties, including polarization resistance (Rp),

conductivity (s) and activation energy (Ea) for a charge

transfer reaction involving molecular oxygen. Increased tem-

peratures result in higher conductivities, showing two distinct

temperature ranges. The detected conductivities were due to

charge transfer processes of molecular oxygen at the platinum

electrode surface and mass transfer processes of oxygen inside

the IL matrix. Comparison of the oxygen charge transfer

process activation energies with the X-ray diffraction data of

2a–c suggests that the oxygen mobility in the ionic liquids

(solids at room temperature) is affected by their nature

of structural supramolecular organization: tri-dimensional

Table 1 Activation energy, conductivity intersection and meltingpoint of the ILs 1–3

Entry IL Eaa/kJ mol�1 Ea

a/eV Tisb/1C Tm

c/1C

1 [C4MIm][Mes] 1a 78.1 0.81 63 772 [C10MIm][Mes] 1b 79.0 0.82 31 573 [Ph2C2MIm][NTf2] 2a 73.7 0.76 40 624 [PhC3MIm][NTf2] 2b 51.6 0.53 42 505 [PhC2MIm][NTf2] 2c 56.6 0.58 28 416 [C2O2MIm][Cl] 3 110 1.14 42 197 (204)d

a Activation energy calculated from the Arrhenius formula. b Tem-

perature at the intersection of the low- and high-temperature range of

the temperature dependent conductivity. c Melting point determined

by differential scanning calorimetry on heating. d Ref. 42.


vs. two-dimensional. Furthermore, the changeover from the

solid to the liquid state and fluxional behavior/reorientation

phenomena in the solid state are most likely the responsible

factors for the faster increasing conductivity in the second

temperature range. As such, electrical impedance spectroscopy

could accelerate the discovery of new electrochemical ionic

liquid (solid at room temperature) applications and the sub-

stitution of ionic liquids where beneficial.

Acknowledgements

The authors thank the CNPq for financial support.

Fig. 6 X-Ray diffraction crystal structures of (a) [Ph2C2MIm][NTf2] 2a (room temperature); (b) [PhC3MIm][NTf2] 2b (room temperature) and

(c) [PhC2MIm][NTf2] 2c (100 K).


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A new family of biocompatible and stable magnetic nanoparticles:

silica cross-linked pluronic F127 micelles loaded with iron oxides

Zhaoyang Liu,* Jun Ding and Junmin Xue*

Received (in Montpellier, France) 18th June 2008, Accepted 4th September 2008


DOI: 10.1039/b810302j

A new family of magnetic nanoparticles, silica cross-linked pluronic F127 micelles loaded with

iron oxides having the properties of high biocompatibility, physical and chemical stability, high

magnetism, and low-cost production, have been synthesized.

1. Introduction

Iron oxide (IO) nanoparticles are emerging as promising

candidates for various biomedical applications, such as mag-

netic resonance imaging (MRI),1 targeted drug delivery,2

hyperthermia treatment,3 the labelling and sorting of cells,4

and the separation of biochemical products,5 due to their

superparamagnetic properties. Most of these applications

require the nanoparticles to be biocompatible, water soluble,

and physically and chemically stable in a physiological environ-

ment.6 To date, the most promising synthetic strategy for

iron oxide nanoparticles is based on the high temperature

decomposition of iron salts in the presence of organic solvents,

which can produce monodisperse and highly crystalline IO

nanoparticles.7 However, their uses in biomedicine are quite

limited because the particles synthesized through this route

can only be dispersed in hydrophobic solvents. Before bio-

medical applications are possible, they have to be transferred

into an aqueous medium. Moreover, the reactivity of iron oxide

particles have been shown to greatly increase as their dimen-

sions are reduced to the nano scale. Therefore, it is necessary

to engineer the surface of IO nanoparticles to improve

their biocompatibility, solubility and stability in physiological

environments for various biomedical applications.8

Several natural and synthetic polymers have been employed

to coat the surface of IO nanoparticles to transfer their surface

wettability. These polymers include dextran,9 lipids,10 dendri-

mers,11 polyethylene glycol (PEG)12 or polyethylene oxide

(PEO),13 and polyvinylpyrrolidone (PVP).14 All the polymers

used are known to be biocompatible and able to promote the

dispersion of IO nanoparticles in an aqueous medium. How-

ever, these polymer coatings are not robust and can be

detached from particle surfaces easily under in vivo condi-

tions.10,12 To improve the stability of the polymer coatings

in vivo, a cross-linking technique has been developed.15 For

example, the stability of dextran coatings on IO nanoparticles

can be improved when the dextran polymer chains are chemi-

cally cross-linked. Although cross-linking is considered a

promising method for strengthening the polymer coatings of

IO nanoparticles,16–19 this method requires multiple synthetic

steps (multi-pot).18 Therefore, a simple and one-pot method is

in demand.

In this work, we have developed a simple and one-pot

method to fabricate a new family of biocompatible and stable

magnetic nanoparticles that are coated with a hybrid layer of

silica cross-linked pluronic F127 (SCL-P@IO). Pluronic F127

(PF127) is an ABA-type triblock copolymer consisting of

hydrophobic poly(propylene oxide) (PPO) and hydrophilic

PEO. The PEO blocks present a high biocompatiblity by

effectively preventing aggregation, the adsorption of proteins,

the adhesion to tissues, and recognition by the reticulo-

endothelial system in vivo.20 Silica has been extensively studied

as an inorganic coating for a long period of time due to its rich

surface chemistry, which means it is able to conjugate biofunc-

tional moieties easily. Compared to polymer coatings, silica

coatings are more chemically stable and resistant to diffusion

for encapsulated components. PF127 copolymer and silica

have also been selected as suitable surface coating candidates

because they are highly safe in vivo and have been approved by

the Food and Drug Administration.20,21

A double-layer coating of PEO and silica on the surface of

IO nanoparticles is proposed, as shown in Fig. 1a. Like

conventional mesoporous silica synthesis using PF127 surfac-

tants,22 the silica was controlled so as to be deposited on the

interior PEO blocks of the PF127 micelle, leaving the exterior

PEO blocks stretched out in aqueous media. Therefore, a

double-layer structure of PEO and silica on the surfaces of

the IO nanoparticles was formed (Fig. 1a). The advantages of

the resulting magnetic nanoparticles are follows: (1) high

biocompatibility and stability due to the presence of PEO

blocks on the surface of the nanoparticles, (2) high chemical

and physical stability due to the robust and dense silica cross-

linked micelles, and (3) simple and low-cost synthesis, since

silica cross-linking is much easier in comparison with tradi-

tional polymer cross-linking, and since both PF127 and the

silica precursor are commercially available. Therefore, the

present magnetic nanoparticles possess great potential in a

variety of biomedical applications.

2. Experimental

2.1 Materials

Dioctyl ether, oleic acid, iron(III) chloride, tetraethoxysilane

(TEOS), diethoxydimethylsilane (Me2Si(OEt)2, DEDMS) and

Department of Materials Science and Engineering, NationalUniversity of Singapore, Singapore, 117576, Republic of Singapore.E-mail: [email protected]. E-mail: [email protected];Fax: +65 67763604; Tel: +65 65164655



pluronic F127 (PF127) were purchased from Aldrich

Chemical Co. Fibroblasts 3T3 were purchased from ATCC.

RPMI-1640 medium, Dulbecco’s modified Eagle’s medium

(DMEM), fetal bovine serum (FBS), L-glutamine, penicillin

and streptomycin were purchased from Sigma. All other

solvents and chemicals were purchased from Aldrich and used

as received.

2.2 Synthesis of IO nanoparticles (Fe3O4)

FeCl3�6H2O (3.3 mmol) and sodium oleate (10 mmol) were

dissolved in a mixture of ethanol (25 mL), de-ionized water

(20 mL) and hexane (45 mL). After refluxing for 4 h at 62 1C,

the iron oleate complex was washed with de-ionized water

three times. The iron oleate complex (3.3 mmol) and oleic acid

(1.67 mmol) were dissolved in dioctyl ether (20 mL) at 70 1C.

After heating for 1.5 h at 290 1C, ethanol (30 mL) was added

to the mixture, and the nanoparticles were collected by

centrifugation at 6000 rpm. The nanoparticles were washed

with hexane and ethanol three times. Finally, the nanoparticles

were dispersed in hexane (40 mL) and oleic acid (100 ml).

2.3 Synthesis of SCL-P@IO nanoparticles

A hexane solution (0.2 mL) of IO nanoparticles was added to

an aqueous solution (8 mL) of PF127 (1.3 g). After 3 h of

stirring at room temperature, the mixture was dried under

nitrogen gas. The obtained powder could be readily redis-

persed in de-ionized water (8 mL) with shaking. Then, a 0.3 M

HCl solution (0.4 g) and TEOS (0.2 g) were added to the

mixture with stirring. After stirring at room temperature for

48 h, DEDMS (0.1 g) was added. Stirring was continued for

another 3 h.

2.4 Cell culture

The cell viability of the nanoparticles was carried out by

testing the viability of 3T3 fibroblasts after incubation with

DMEM, supplemented with 10% FBS, 1 mM L-glutamine and

100 IU mL�1 penicillin via an 3-(4,5-dimethylthiazol-2-yl)-2,5-

diphenyl tetrazolium bromide (MTT) assay. Mouse macro-

phage cells (RAW 264.7) were used to assess the cellular

uptake of the nanoparticles. The RAW 264.7 cells were

cultured in RPMI-1640 medium, supplemented with 10%

FBS, 2 mM L-glutamine, 100 IU mL�1 penicillin and

100 mg mL�1 streptomycin. The cell concentrations were

determined by hemacytometry, and the Fe concentrations

were determined by Thermal Jarrell Ash Duo Iris inductively-

coupled plasma optical emission spectrometer (ICP-OES).

2.5 Characterization

TEM measurements were taken using a JEOL JEM 3010

instrument. The hydrodynamic diameters of the nanoparticles

were measured by a Malvern Zeta Sizer Nano S-90 dynamic

light scattering (DLS) instrument. TGA analyses were per-

formed by a TA Instruments Q500 thermogravimetric analyzer.

IR studies were run on an ATI Mattson Infinity Series

FT-IR spectrophotometer. Magnetic properties were mea-

sured on a superconducting quantum interference device

(SQUID, Quantum Design, USA) and a Lakeshore 7300 series

vibrating sample magnetometer (VSM). The fluorescence

emission spectra of pyrene were measured by a fluorescence

photometer (FP-777 Jasco) at 254 nm excitation. 29Si NMR

spectra were recorded with an Advance 500 Bruker spectro-

meter at 99.36 MHz. The pulse length was 6 ms (theta = p/6)with 6 s repetitions.


3.1 TEM, XRD and29Si NMR

In a typical synthesis, 10.5 nm monodisperse and hydrophobic

Fe3O4 nanoparticles were synthesized separately (Fig. 1b).23

The as-synthesized Fe3O4 nanoparticles were coated with a

layer of oleic acid, which made them only soluble in hydro-

phobic solvents (Fig. 1b inset photo). With the addition of

PF127 surfactants, these polymeric surfactants self-assembled

into a micellar structure, encapsulating the IO nanoparticles in

their cores. The hydrophobic PPO block in the middle of the

PF127 associated with the alkyl tail of the oleic acid through a

hydrophobic interaction, while the two hydrophilic PEO

blocks stretched out in aqueous media (Fig. 1a). Then, TEOS

Fig. 1 a: Synthetic scheme for silica cross-linked pluronic F127

micelles loaded with IO nanoparticles (SCL-P@IO). Left: an as-

synthesized hydrophobic IO nanoparticle. Right: an IO nanoparticle

with its surface covered by a double-layer of PEO and silica. b: A TEM

of as-synthesized hydrophobic IO nanoparticles. The inset photo

shows that the IO nanoparticles can only be dispersed in hexane. c:

A TEM of SCL-P@IO nanoparticles. The right bottom inset is a

higher magnification TEM of typical SCL-P@IO nanoparticles. The

inset photo shows that the SCL-P@IO nanoparticles can be readily

dispersed in water. d: Size distribution histogram of the nanoparticles

in c. e: XRD patterns of IO and SCL-P@IO nanoparticles. f: 29Si

NMR spectra corresponding to samples (a) with and (b) without the

addition of DEDMS, respectively.


was added as a silica precursor to polymerize and cross-link

the PF127 micellar shells.

Fig. 1c shows a TEM image of the SCL-P@IO nanoparti-

cles. The formation of well-defined core/shell morphologies is

readily confirmed, since the Fe3O4 within the micelle cores is

more electron-dense than the silica/PF127 hybrid shell.

A higher magnification TEM of typical SCL-P@IO nano-

particles is shown in the right bottom of Fig. 1c. These

nanoparticles have a number-average diameter of around

21 nm, as shown in the histogram in Fig. 1d, with an ultrathin

(about 5 nm) silica shell deposited outside the Fe3O4 nano-

particles. A size distribution analysis, as shown in the

histogram in Fig. 2a, reveals that the spherical NPs are

monodisperse and have an average size of 7.9 � 1.5 nm. The

inset photo shows that the resulting nanoparticles can be

readily dispersed in water. The XRD patterns of the IO and

SCL-P@IO nanoparticles (Fig. 1e) can be assigned to the

(220), (311), (400), (422), (511) and (440) reflections of the

spinel structure of magnetite (JCPDS no. 19-0629). The broad

band at 20–301 is due to the presence of amorphous

silica, while the labelled peaks are associated with Fe3O4

nanocrystals.

The condensation of silicate during the preparation was

studied by NMR techniques. Fig. 1f (a) and (b) correspond to

the sample with and without the addition of DEDMS, respec-

tively. In the 29Si NMR spectra, three peaks at ca. �94, �104and �114 ppm are assigned to Q2, Q3 and Q4 species with

progressively increasing cross-linking (condensation).24 After

automatic calculation of the integrated area, the Q4 : Q3 ratio

decreased from 0.91 (b) to 0.83 (a), indicating a slightly lower

condensation of silicate with the addition of DEDMS.

3.2 DLS and FT-IR

The obtained SCL-P@IO nanoparticles were characterized by

DLS. The DLS measurements of the IO and SCL-P@IO

nanoparticles are shown in Fig. 2a. The hydrodynamic

diameter of the SCL-P@IO nanoparticles is 43.3 nm, which

is larger than that measured by TEM in Fig. 1c. This is

because the light scattering measurement includes the PEO

chains stretching out into the aqueous solution, which cannot

be observed by TEM due to the low contrast of the PEO

polymers.25 This result also confirms the double-layer struc-

ture of PEO and silica on the surface of SCL-P@IO nano-

particles instead of just a single layer of silica; that is, the silica

is deposited on the interior PEO chains, while the exterior

PEO chains are stretched out into the aqueous solution, as

suggested in Fig. 1a. The surface coating of the SCL-P@IO

nanoparticles was further studied by IR analysis. As shown

in Fig. 2b, silica formation is confirmed, since a band at

1080 cm�1, assigned to Si–O–Si, is observed for these

SCL-P@IO nanoparticles.26 After calcination at 700 1C, the

characteristic band of C–H at 1726 cm�1 disappears com-

pletely, while the band assigned to thermally stable silica is still

observed.

3.3 Magnetism characterization

The magnetic properties of the IO and SCL-P@IO nanopar-

ticles were examined at room temperature by using a VSM. As

shown in Fig. 3a, the saturated magnetization of the IO and

SCL-P@IO nanoparticles were 63.1 and 28.3 emu g�1, res-

pectively. The reduction in saturated magnetization for the

SCL-P@IO nanoparticles accounts for the diamagnetic

properties of the silica and the PF127 shell surrounding

the IO cores. The inset photo in Fig. 3a shows the magnetic

manipulation ability of the SCL-P@IO nanoparticles.

When an external magnet is placed beside the glass vial, the

aqueous dispersion of SCL-P@IO nanoparticles could be

directed towards the magnet. This efficient magnetism will

allow these nanoparticles to be useful in many biomedical

applications, such as targeted delivery and separation. The

magnetization change of the SCL-P@IO nanoparticles with

storage time was monitored by VSMmeasurements. As shown

in Fig. 3b, the saturated magnetization of the SCL-P@IO

nanoparticles was almost constant (at around 28 emu g�1)

during 90 d storage, suggesting that the PF127/silica hybrid

coating is dense enough to be non-permeable, preventing the

encapsulated IO cores from degrading and leading to lower

magnetism.

3.4 Cell viability and uptake

The biocompatibility of the SCL-P@IO nanoparticles was

examined by the cell viability of 3T3 fibroblast lines through

Fig. 2 a: Hydrodynamic sizes of IO and SCL-P@IO nanoparticles in

aqueous solution determined by DLS measurements. b: IR spectra of

PF127 triblock copolymer (top), SCL-P@IO nanoparticles (middle)

and SCL-P@IO nanoparticles after calcination at 700 1C (bottom).

The IR spectra of the SCL-P@IO nanoparticles are characteristic of

both the silica network (Si–O–Si stretch at 1080 cm�1) and the

copolymer (C–H stretch at 1730 cm�1). This C–H band disappears

after calcination.

Fig. 3 a: Magnetization curves (M–H) of IO and SCL-P@IO nano-

particles. The inset photo shows that the SCL-P@IO nanoparticles

can be driven by an external magnet. b: The magnetization change of

the SCL-P@IO nanoparticles with storage time, as observed by VSM

measurements.


an MTT assay. As shown in Fig. 4a, the SCL-P@IO nano-

particles were biologically inert up to an iron concentration of

1000 mg mL�1. This result indicates that the SCL-P@IO

nanoparticles are highly non-toxic and biocompatible. When

injected into the bloodstream, nanoparticles are often con-

sidered as an intruder by the innate immunity system, and can

be readily recognized and become engulfed by the macrophage

cells. The nanoparticles will then be removed from the blood

circulation system, and lose their efficiency in diagnostics and

therapeutics.

Fig. 4b shows the uptake of the oleic acid-stabilized (Oleic

acid-IO) and PEO-stabilized (SCL-P@IO) nanoparticles

by macrophage cells with an initial Fe concentration of

0.23 mg mL�1. The Oleic acid-IOs were quickly internalized

into the cells within 24 h, with an uptake of 163 pg Fe cell�1.

The amount taken in by the cells decreased with time because

of the rapid growth and division of the macrophage cells.

After grafting with PEO, the SCL-P@IO nanoparticles’

uptake by macrophage cells was much lower, at only

3 pg Fe cell�1, compared with Oleic acid-IO nanoparticles.

The very low uptake of SCL-P@IO nanoparticles could be

due to surface PEO grafting of the SCL-P@IO nanoparticles

lowering the adsorption of the proteins and decreasing the

possibility of macrophage recognition.

3.5 Stability test

The stability of the SCL-P@IO nanoparticles under physio-

logical conditions was also studied by DLS measurements.

The size change of the SCL-P@IO nanoparticles with incuba-

tion time in phosphate-buffered saline (PBS) plus 10% FBS

were monitored. As shown in Fig. 5a, the sizes of SCL-P@IO

nanoparticles were almost constant over 90 d incubation,

indicating that these nanoparticles did not aggregate and were

fairly dispersed in the physiological medium. This is mainly

attributed to the anti-aggregation and anti-biofouling proper-

ties of the PEO blocks of PF127 on the surfaces of SCL-P@IO

nanoparticles. The stability of micelles under dilution is also

a big issue because they are highly diluted in vivo after

administration. Pyrene is an effective fluorescent probe to test

micelle stability. The ratio between the first (375 nm) and third

(386 nm) emission intensities (Im1/Im3) of the pyrene spectrum

depends on the environmental polarity of the solvent.27 Here,

the fluorescence spectra of pyrene molecules encapsulated in

silica-cross-linked PF127 micelles were measured as a function

of PF127 surfactant concentration (different dilution levels) to

study the micelle stability. As shown in Fig. 5b, the low and

almost constant Im1/Im3 ratio (about 1.2 : 1) meant that the

silica deposition effectively cross-linked the PF127 micellar

chains and stabilized the micelles from dissociation under

dilution.

4. Conclusions

In conclusion, we have fabricated a new family of magnetic

nanoparticles based on silica cross-linked PF127 block co-

polymer micelles loaded onto IO nanoparticles. The advan-

tages of high biocompatibility, physical and chemical stability,

high magnetism, and the low-cost of production of these new

magnetic nanoparticles make them promising for a wide range

of biomedical applications, such as bioimaging, bioseparation

and drug delivery.

Acknowledgements

This work was supported by the Singapore MOE’s ARF Tier 1

funding WBS R-284-000-050-133.

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dilution.


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Novel thiophene-conjugated indoline dyes for zinc oxide solar cells

Takuya Dentani,aYasuhiro Kubota,

aKazumasa Funabiki,

aJiye Jin,

b

Tsukasa Yoshida,cHideki Minoura,

cHidetoshi Miura

dand Masaki Matsui*

a

Received (in Montpellier, France) 28th May 2008, Accepted 21st August 2008


DOI: 10.1039/b808959k

The application of a series of thiophene-conjugated indoline dyes for zinc oxide solar cells,

prepared by the one-step cathode deposition template method, was examined. The introduction

of thiophene ring(s) into D131-type indoline dye improved the cell performance due to their

appropriate energy levels and bathochromic shift in the UV-vis absorption band on zinc oxide.

It is important for the oxidation potential (Eox) of dyes to have a more positive value than

ca. 0.25 V vs. Fc/Fc+ in acetonitrile in order to show a high (470%) incident photon-to-

current efficiency.

Introduction

Organic dyes, such as coumarins,1 styryls,2 polyenes,3 dimethyl-

fluorenyl-containing derivatives4 and indoline derivatives,5

have been reported to act as good sensitizers for titanium

oxide. Bathochromic organic dyes, such as squaryliums,6

phthalocyanines7 and heptamethinecyanines,8 have also been

reported to sensitize semiconductors. In particular, D149

has been reported to show the highest solar-light-to-electricity

conversion efficiency (Z) of 9.0% among organic dyes.5a

One promising approach to improve the performance of

sensitizers is the expansion of the p-conjugation system to

absorb more photons. The introduction of ethylene and

thiophene units into chromophores is a good methodology

to expand p-conjugation.9 On the other hand, a convenient

preparation process for zinc oxide thin films has been

reported.10 The key point of this method is the formation

of porous zinc oxide films at low temperature (o70 1C).

Indoline dyes D131, D102 and D149, in which cyanoacrylic,

monorhodanic and double rhodanic acids are used as

anchor moieties, respectively, are known to show good

performances.5f We report herein the application of novel

thiophene-conjugated indoline dyes having a series of anchor

moieties to zinc oxide dye-sensitized solar cells.


Synthesis of indoline dyes

Thiophene-conjugated indoline dyes 20–28 were synthesized,

as shown in Scheme 1. Compound 1 was allowed to react with

NBS (2) to give 3, followed by a reaction with thiophene

boronic acids esters 4–7 to provide 8–11, which were formy-

lated to give 12–15. These compounds were allowed to react

with cyanoacetic, mono- and double-rhodanic acids 16–19 to

provide 20–28. D131, D102 and D149 were prepared in a

similar way.

UV-vis absorption and fluorescence spectra

The UV-vis absorption and fluorescence spectra of 20–28,

D131, D102 and D149 are shown in Fig. 1, Fig. 2 and

Fig. 3. The results are also listed in Table 1. All the indoline

dyes showed first and second absorption bands at around 500

and 400 nm, respectively. The first absorption maximum

(lmaxfirst) of monothiophene derivatives 20, 23, 24, 25 and 26

were more bathochromic than thiophene-free derivatives

D131, D102 and D149. Interestingly, no further bathochromic

shift was observed for di- and trithiophene derivatives 21, 22,

27 and 28 compared to monothiophene derivatives 20, 23, 24,

25 and 26, respectively. The molar absorption coefficients at

the first absorption band (efirst) of 20–28 were less than those of

thiophene-free derivatives D131, D102 and D149. The half-

widths of 20–28 (99–146 nm) were larger than those of D131,

D102 and D149 (65–79 nm). No marked difference in efirst

among mono-, di- and trithiophene derivatives was observed,

being in the range 37 700 to 47 900. The second absorption

maximum (lmaxsecond) of thiophene derivatives 20–28 showed

a bathochromic shift, and at the same time, their molar

absorption coefficients (esecond) were larger with increasing

numbers of thiophene units. No remarkable differences in

the UV-vis absorption spectra between 20 and 23, and between

25 and 26, were observed. The fluorescence maximum (Fmax)

showed a bathochromic shift by the introduction of a

thiophene unit.

Electrochemical properties

The oxidation potentials (Eox) ofD131,D102,D149 and 20–25

were measured by using an Ag/Ag+ electrode in acetonitrile to

compare the energy levels of Eox and Eox � E0–0 of the dyes,

the I�/I3� potential, and the conduction band of zinc oxide.

aDepartment of Materials Science and Technology, Faculty ofEngineering, Gifu University, Yanagido, Gifu 501-1193, Japan.E-mail: [email protected]; Fax: +81 58 293 2794;Tel: +81 58 293 2601

bDepartment of Chemistry, Faculty of Science, Shinshu University,3-1-1 Asahi, Matsumoto, Nagano 390-8621, Japan

c Environmental and Renewable Energy System Division, GraduateSchool of Engineering, Gifu University, Yanagido, Gifu 501-1193,Japan

dChemicrea Co. Ltd., 2-1-6 Sengen, Tsukuba, Ibaragi 305-0047,Japan



Fc/Fc+ was used as a standard. The Eox of ferrocene was

observed at +0.13 V vs. Ag/Ag+ in acetonitrile. Fig. 4 shows

that the Eox of 20 was observed at +0.38 V vs. Ag/Ag+ in

acetonitrile, corresponding to +0.25 V vs. Fc/Fc+ in aceto-

nitrile. The I�/I3� potential level was observed at +0.09 V vs.

Ag/Ag+ in acetonitrile, corresponding to �0.04 V vs. Fc/Fc+

in acetonitrile.

The potential level of Eox � E0–0, where E0–0 represents the

intersection of the normalized absorption and fluorescence

spectra in solution, is considered to correspond to the LUMO

energy level.9 The E0–0 of 20 was observed at 589 nm,

corresponding to 2.11 eV. Therefore, the Eox � E0–0 value of

20 was calculated to be�1.86 V vs. Fc/Fc+ in acetonitrile. The

energy levels of free indoline dyes measured in solution

differed from those of adsorbed ones. Unfortunately, ferro-

cene and indoline dyes on a zinc oxide-coated ITO electrode

did not give distinct redox responses due to a slow charge

transfer process. Hence, the Eox of ferrocene and indoline dyes

could not be determined. The Eox and Eox � E0–0 of all the

indoline dyes are listed in Table 1.

Scheme 1 The synthesis of indoline dyes 20–28.

Fig. 1 UV-vis absorption and fluorescence spectra of indoline dyes

D131, 20, 21, 22 and 23 at a concentration of 1.0 � 10�5 mol dm�3 in

chloroform. Solid and dotted lines represent UV-vis absorption and

fluorescence spectra, respectively.


D102 and 24 at a concentration of 1.0� 10�5 mol dm�3 in chloroform.

Solid and dotted lines represent UV-vis absorption and fluorescence

spectra, respectively.


UV-vis absorption and IR spectra of 20

The UV-vis absorption spectra of 20 are shown in Fig. 5. The

lmaxfirst of 20 in chloroform and on zinc oxide were observed at

519 and 449 nm, respectively. Thus, large hypsochromic shift

of lmaxfirst was observed on zinc oxide. The lmax

first of 20 in the

presence of an equimolar amount of triethylamine (TEA) in

chloroform was observed at 470 nm, there being slightly more

bathochromic than that on zinc oxide.

FTIR spectra of 20 are shown in Fig. 6. The IR spectrum of

20 in a potassium bromide disk showed an absorption band at

around 1680 cm�1, which was assigned to a carbonyl stretch-

ing absorption. When indoline dye 20, adsorbed onto a zinc

oxide film, was scraped off and its IR spectrum was measured

in a potassium bromide disk, the absorption band at around

1680 cm�1 disappeared and new absorption was observed at

around 1600 cm�1. This spectrum is similar to that of the

triethylammonium salt of 20, in which the absorption band at

around 1600 cm�1 is assigned to the asymmetric stretch

absorption of the carboxylate anion. It was also observed

that indoline dye 20 showed negative solvatochromism

in solution (lmaxfirst = 516 (toluene), 519 (chloroform),


D149, 25, 26, 27 and 28 at a concentration of 1.0 � 10�5 mol dm�3 in

chloroform. Solid and dotted lines represent UV-vis absorption and

fluorescence spectra, respectively.

Table 1 Optical and electrochemical properties of indoline dyes

Compound lmax/nm (e)a Fmax/nma RFIb Eox vs. Fc/Fc

+ in MeCN/V Eox � E0–0 vs. Fc/Fc+ in MeCN/V

D131 463 (55 400) 591 83 +0.41 �2.00325 (15 600)

20 519 (43 300) 659 77 +0.25 �1.86373 (27 300)

21 517 (37 700) 712 27 +0.23 �1.81393 (38 300)

22 519 (41 700) 701 4 +0.22 �1.83409 (47 100)

23 523 (47 300) 653 203 +0.25 �1.85373 (29 100)

D102 514 (54 700) 621 68 +0.37 �1.83368 (25 200)

24 548 (41 400) 702 80 +0.25 �1.75388 (37 100)

D149 550 (68 000) 636 100 +0.30 �1.79395 (32 000)

25 571 (43 500) 717 78 +0.24 �1.67410 (34 800)

26 568 (45 600) 713 49 —c —c

408 (40 500)27 564 (42 000) 743 10 —c —c

405 (41 900)28 550 (47 900) 727 3 —c —c

412 (48 900)

a Measured on 1.0 � 10�5 mol dm�3 of substrate in chloroform at 25 1C. b Relative fluorescence intensity. c Not measured due to low solubility.

Fig. 4 The electrochemical measurement of 20 in acetonitrile (2 ml)

containing tetrabutylammonium perchlorate (0.1 mol dm�3). Ag/Ag+

in acetonitrile was used as a reference electrode. Platinum wire

was used as the working and counterelectrode. The scan rate was

100 mV s�1.

Fig. 5 The UV-vis absorption spectra of 20 in chloroform and on

zinc oxide.


529 (dichloromethane), 465 (DMSO), 453 (acetonitrile) and

464 nm (methanol)). These results indicate that the hypso-

chromic shift of 20 on zinc oxide is mainly attributed to the

formation of a bidentate complex between the carboxylate and

zinc. The polar zinc oxide surface could also be attributed to

the hypsochromic shift.

Photoelectrochemical properties

The cell performance ofD131-type indoline dyes was examined.

The normalized UV-vis absorption spectra on zinc oxide

and action spectra are shown in Fig. 7. The results are also

listed in Table 2. The cell performance was improved in the

presence of cholic acid (CA), a co-adsorbate that can inhibit

the aggregation of dyes on zinc oxide due to carboxylic acid

and hydrophobic moieties. The UV-vis absorption spectra of

20 and 23 in the absence and presence of CA are depicted in

Fig. 7(a). In the case of 20, a broad absorption at around

530 nm decreased in the presence of CA, indicating the

prevention of aggregation on zinc oxide. Meanwhile, only

slight differences in the absorption bands between the absence

and presence of CA were observed for 23. The Z values of 20

and 23 in the absence of CA were 3.13 and 3.30%, respectively

(Table 2, runs 3 and 7). Those in the presence of CA were 3.78

and 3.36%, respectively (Table 2, runs 2 and 6). Thus, the Zvalues of 20 and 23 were improved by 21 and 2% in the

presence of CA, respectively. These results suggest that the

hexyl group in 23 is very effective in inhibiting aggregation on

zinc oxide. In the cases of 21 and 22, aggregation formation

decreased in the presence of CA, resulting in an improved cell

performance (Table 2, runs 4 and 5). The absorption bands of

20, 21, 22 and 23 on zinc oxide were more bathochromic than

that of D131, as shown in Fig. 7(b). The action spectra show

the sensitization of zinc oxide by 20, 21, 22 and 23 at around

550 nm, whereas no sensitization was observed for D131 at

around 550 nm, as depicted in Fig. 7(c). The incident photon-

to-current efficiency (IPCE) in the presence of CA was in the

following dye order: 20 (83.1%) 4 23 (78.2%), D131 (77.8%)

4 21 (69.1%) 4 22 (55.5%) (Table 2, runs 1, 2, 4–6). The

short-circuit photocurrent densities (Jsc) of 21 (8.15 mA cm�2),

20 (8.09 mA cm�2), 23 (7.42 mA cm�2) and 22 (6.69 mA cm�2)

were higher than that of D131 (5.55 mA cm�2). The fill factor

(ff) was lowered by introducing a thiophene unit. Con-

sequently, the Z value was in the following order of dyes:

20 (3.78%) 4 23 (3.36%), 21 (3.19%) 4 D131 (2.60%) 4 22

(2.08%). Thus, an improvement in cell performance was

successfully observed for a series of D131-type thiophene-

conjugated indoline dyes. The improved cell performance of

20, 21 and 23, compared with D131, mainly came from the

bathochromic shift in the absorption band and a high IPCE

(470%) to increase Jsc.

Next, the cell performance of D102 and 24 was examined

(Table 2, runs 8–10). The UV-vis absorption and action

spectra are shown in Fig. 8. The absorption band of 24 was

more bathochromic than that of D102. The absorption spec-

trum of 24 in the absence of CA clearly showed a broad

absorption at around 600 nm, suggesting the formation of

aggregates. Fig. 8(b) shows the sensitization of zinc oxide by

24 at around 630 nm. However, the IPCE value of 24 was

lower than that of D102 so as not to increase Jsc. The open-

circuit voltage (Voc) and ff of 24 were lower than those of

D102 (Table 2, runs 8 and 9). Thus, no improvement in

cell performance was observed for D102-type thiophene-

conjugated indoline dyes.

Fig. 6 FTIR spectra of 20: (a) 20, (b) 20 in the presence of zinc oxide

and (c) the triethylammonium salt of 20.

Fig. 7 (a) Normalized UV-vis absorption spectra of 20 and 23 on zinc

oxide in the absence and presence of CA, (b) normalized UV-vis

absorption spectra of D131, 20, 21, 22 and 23 on zinc oxide in the

presence of CA and (c) action spectra of D131, 20, 21, 22 and 23 in the

presence of CA.


Finally, the cell performance ofD149-type indoline dyes was

examined. The UV-vis absorption and action spectra of D149,

25, 26, 27 and 28 are shown in Fig. 9. In this case, the

difference in the UV-vis absorption bands of 25 and 26 in

the absence and presence of CA was small compared with

those in the cases of 20 and 23, as shown in Fig. 9(a). The Zvalues of 25 and 26 in the presence of CA were higher than

those in the absence of CA (Table 2, runs 12–15). Fig. 9(b)

shows that 25, 26, 27 and 28 are more bathochromic than

D149. Fig. 9(c) indicates that though the sensitization of zinc

oxide was observed for 25, 26, 27 and 28 at around 670 nm,

Fig. 8 (a) Normalized UV-vis absorption spectra of D102 and 24 on

zinc oxide in the absence and presence of CA, and (b) action spectra of

D102 and 24 in the presence of CA.

Table 2 Physical properties of indoline dyes

Run Compound CAa lmax/nm Abs.b IPCE (%) Jsc/mA cm�2 Voc/V ff Zc (%)

1 D131 2 405 3.36 77.8 5.55 0.66 0.71 2.602 20 2 449 2.48 83.1 8.09 0.69 0.68 3.783 20 0 459 1.96 71.4 7.09 0.65 0.68 3.134 21 2 461 2.16 69.1 8.15 0.63 0.62 3.195 22 2 457 2.44 55.5 6.69 0.59 0.53 2.086 23 2 450 2.07 78.2 7.42 0.66 0.68 3.367 23 0 452 2.07 76.2 7.58 0.65 0.67 3.308 D102 2 476 3.20 77.1 9.00 0.65 0.66 3.889 24 2 514 1.94 48.6 7.33 0.62 0.63 2.8310 24 0 539 1.54 42.4 6.44 0.59 0.58 2.2011 D149 2 521 3.08 81.2 11.08 0.68 0.57 4.2312 25 2 547 1.76 43.4 6.85 0.54 0.64 2.3513 25 0 555 1.65 35.9 5.38 0.50 0.62 1.6814 26 2 546 1.25 37.5 5.85 0.62 0.68 2.4515 26 0 561 1.21 37.7 5.55 0.57 0.65 2.0716 27 2 546 1.26 29.7 4.40 0.59 0.66 1.7117 28 2 542 1.45 26.1 3.67 0.56 0.67 1.36

a Equivalents of cholic acid with respect to dye. b Absorbance at absorption maximum on zinc oxide. c Action spectra and I–V characteristics

under AM 1.5 irradiation (100 mW cm�2).

Fig. 9 (a) Normalized UV-vis absorption spectra of 25 and 26 on zinc

oxide in the absence and presence of CA, (b) normalized UV-vis

absorption spectra of D149, 25, 26, 27 and 28 on zinc oxide in the

presence of CA, and (c) action spectra of D149, 25, 26, 27 and 28 in the

presence of CA.


their IPCE values were lower than that of D149. Thus, no

improvement in cell performance was observed for the series

of D149-type thiophene-conjugated indoline dyes (Table 2,

runs 11, 12, 14, 16 and 17).

Relationship between IPCE and Eox, Eox � E0–0

To examine why only D131-type thiophene-conjugated indo-

line dyes showed improved cell performances, the relationship

between IPCE and energy levels was examined. Fig. 10(a)

shows the relationship between IPCE and Eox. Indoline dyes

D131, 20, 21, 23, D102 and D149, of which the Eox levels were

more positive than the ca. +0.25 V vs. Fc/Fc+ in acetonitrile,

showed high (470%) IPCE values. The potential level of

I�/I3� was observed at �0.04 V vs. Fc/Fc+ in acetonitrile.

Fig. 10(b) shows the relationship between IPCE and Eox �E0–0. It was also found that indoline dyes D131, 20, 21, 23,

D102 and D149 showed high IPCE values. It is reported that

the potential levels of the conduction band of titanium oxide

and the I�/I3� redox are �0.5 and +0.4 V vs. NHE, respec-

tively, there being an energy gap of 0.9 V.9,11 The conduction

band level of zinc oxide is similar to that of titanium oxide.

Therefore, the level of zinc oxide is considered to be �0.94 V

vs. Fc/Fc+ in acetonitrile, which is much more positive than

the Eox � E0–0 levels of all the indoline dyes, the energy gap

between Eox � E0–0 and the conduction band levels being

larger than 0.7 V. It is suggested that an energy gap larger than

0.2 V between Eox and I�/I3�, and Eox � E0–0 and the

conduction band levels, respectively, are required.9 Thus,

though the Eox � E0–0 level of all the indoline dyes are

sufficiently negative, their Eox levels are critical for the sensi-

tization cycle to proceed. No marked difference in the

Eox levels among the thiophene-conjugated derivatives 20

(+0.25 V), 21 (+0.23 V), 22 (+0.22 V), 23 (+0.25 V), 24

(+0.25 V) and 25 (+0.24 V) was observed in solution.

However, their Eox level on zinc oxide could differ from that

in solution. As the Eox level of D131 (+0.41 V vs. Fc/Fc+ in

MeCN) was more positive than those of D102 (+0.37 V) and

D149 (+0.30 V) in solution, those ofD131-type derivatives 20,

21, 22 and 23 might be more positive than those of D102- and

D149-type derivatives 24, 25, 26, 27 and 28 on zinc oxide. The

Eox levels of D102, 20, 21 and 22 were observed at +0.37,

+0.25, +0.23 and +0.22 V vs. Fc/Fc+ in acetonitrile,

respectively. This suggests that the Eox level can negatively

shift with increasing numbers of thiophene units on zinc oxide.

Therefore, the Eox levels of D131-type mono- and dithiophene

derivatives 20, 21 and 23 could be more positive than those of

D102- and D149-type derivatives, and the redox potential of

I�/I3� on zinc oxide could show an improved cell perfor-

mance. As a result, indoline dyes 20, 21 and 23 could show

better performances than D131 due to larger Jsc values. The

Eox levels of D102- and D149-type thiophene-conjugated

indoline dyes 24, 25, 26, 27 and 28 might be too negative on

zinc oxide, despite their bathochromic shift in the UV-vis

absorption spectrum on zinc oxide. In order to improve the

performance of indoline dyes, it is important to design deri-

vatives of them having more positive Eox level.

Conclusion

A series of D131-, D102- and D149-type thiophene-conjugated

indoline dyes were examined as sensitizers for zinc oxide solar

cells, prepared by the one-step cathode deposition template

method. Among the series of thiophene-conjugated indoline

dyes, D131-type indoline dyes improved cell performance.

This could have been due to their positive Eox levels. In order

to improve the performance of D102- and D149-type indoline

dyes, it is important to design derivatives of them having more

positive Eox levels.

Experimental

General

Melting points were measured with a Yanagimoto MP-52

micro-melting-point apparatus. NMR spectra obtained using

a JEOL JNM-ECX 400P spectrometer. EI and FAB MS

spectra were recorded on a JEOL MStation 700 spectrometer.

UV-vis absorption and fluorescence spectra were acquired on

Hitachi U-3500 and F-4500 spectrophotometers, respectively.

Cyclic voltammetry was carried out using an EG&G Princeton

Applied Research Potentiostat/Galvanostat (Model 263A)

driven by the M270 software package. One-step cathode electro-

deposition was undertaken using a Hokuto-Denko HSV-100

potentiostat system. The photoelectrochemical measurements

of solar cells were performed on a Bunko-Keiki CEP-2000

system. The I–V curve measurements of solar cells were

performed on an EKO Instruments I–V curve tracer MP-160

and Grating spectroradiometer LS-100.

Electrochemical measurements

The electrochemical measurements of indoline dyes D131, 20,

21, 22, 23, D102, 24, D149, 25, ferrocene and potassiumFig. 10 The relationship between IPCE and energy levels: (a) IPCE

vs. Eox and (b) IPCE vs. Eox � E0–0.


iodide, were performed in acetonitrile. The oxidation potential

(Eox) was measured by using three small-sized electrodes.

Ag/Ag+ was used as a reference electrode. Platinum wire

was used as the working and counterelectrode. Acetonitrile

solutions (2 ml) of dyes containing tetrabutylammonium

perchlorate (0.1 mol dm�3) were prepared. Dry argon

gas was introduced into the solution for 10 min. The electro-

chemical measurements were then performed at a scan rate

of 100 mV s�1.

Preparation of the zinc oxide solar cell

An aqueous potassium chloride solution (300 ml, 0.1 mol dm�3)

was electrolyzed at �1.0 V vs. SCE with bubbling oxygen

gas at 70 1C for 30 min. Platinum was used as a counter-

electrode. To the pre-electrolyzed film was added an

aqueous solution of zinc chloride. The concentration of zinc

chloride was adjusted to 5 mmol dm�3. Then, the film was

again electrodeposited in the solution at �1.0 V vs. SCE at

70 1C for 20 min with bubbling oxygen gas. To the electro-

deposited film was added an aqueous solution of eosin Y

(0.050 mmol dm�3). The film was electrodeposited at �1.0 V

vs. SCE at 70 1C for 30 min with bubbling oxygen gas. The film

was kept in a dilute aqueous potassium hydroxide solution

(pH 10.5) for 24 h to remove adsorbed eosin Y. The film was

then dried at 100 1C for 1 h. The thin film was immersed in a

chloroform solution of dye (1 � 10�4 mol dm�3) and kept at

ambient temperature for 1 h to adsorb dyes 20–28 onto the

zinc oxide. In the cases of D131, D102 and D149, the film was

immersed in an acetonitrile–tert-butyl alcohol 1 : 1 mixed

solution (0.5 mmol dm�3). Then, the film was washed with

chloroform. In the cases of D131,D102 andD149, the film was

washed with an acetonitrile–tert-butyl alcohol 1 : 1 mixed

solution. The films were dried under an air atmosphere at

ambient temperature. The film was used as the working

electrode. A platinum spattered film was used as the counter-

electrode. The cell size was 5.0 � 5.0 mm. Thermosetting resin

was put around the cell. An acetonitrile–ethylene carbonate

(v/v = 1 : 4) mixed solution containing tetrabutylammonium

iodide (0.5 mol dm�3) and iodine (0.05 mol dm�3) was used as

the electrolyte.

Photoelectrochemical measurements

Action spectra were measured under monochromatic light

with a constant photon number (5 � 1015 photon cm�2 s�1).

I–V characteristics were measured under illumination with

AM 1.5 simulated sun light (100 mW cm�2) through a shading

mask (5.0 � 4.0 mm).

Synthesis of dyes

Materials. 1,2,3,3a,4,8b-Hexahydro-4-[4-(2,2-diphenylethenyl)-

phenyl]cyclopent[b]indole (1) was supplied from Chemicrea

Co. Ltd. N-Bromosuccinimide (NBS, 2) and 2-(4,4,5,5-tetra-

methyl-1,3,2-dioxaborolan-2-yl)thiophene (4) were purchased

from Wako Pure Chemical Industries Ltd. 5-(4,4,5,5-Tetra-

methyl-1,3,2-dioxaborolan-2-yl)-2,20-bithiophene (5), 5-(4,4,5,5-

tetramethyl-1,3,2-dioxaborolan-2-yl)-2,2 0,5 0,200-terthiophene

(6) and cyano acetic acid (16) were purchased from Aldrich

Co. Ltd. Rhodanine-3-acetic acid (17) was purchased from

Tokyo Kasei Co. Ltd. Compound 19 was synthesized in the

similar procedure to that described for 18.12 3-Hexyl-2-

(4,4,5,5-tetramethyl-1,3,2-dioxaborolan-2-yl)thiophene (7)13

and 3,5-(di-tert-butyl)benzylamine14 were synthesized as des-

cribed in the literature. D131, D102 and D149 were prepared

in a similar way, as described in the literature.5c,e

Synthesis of 3. To a dry acetone solution (23 ml) of 1

(980 mg, 2.37 mmol) was added NBS (423 mg, 2.38 mmol)

at 0 1C under an argon atmosphere. The mixture was stirred at

room temperature for 3 h. The reaction mixture was poured

into water (20 ml) and extracted with chloroform (3 � 50 ml).

The extract was washed with brine (2 � 50 ml) and dried over

anhydrous sodium sulfate. The solvent was removed in vacuo.

The crude product was purified by silica gel column chromato-

graphy (chloroform–hexane = 1 : 3) to afford 3 as a pale

yellow solid. Yield 96%, mp 81–83 1C. dH (400 MHz, CDCl3,

Me4Si): 1.42–1.49 (1 H, m), 1.61–1.65 (1 H, m), 1.79–1.87

(3 H, m), 1.96–2.02 (1 H, m), 3.76–3.79 (1 H, m), 4.65–4.69

(1 H, m), 6.83 (1 H, d, J = 8.4 Hz), 6.92 (1 H, s), 6.99–7.01

(4 H, m), 7.09 (1 H, d, J= 8.4 Hz), 7.16 (1 H, s) and 7.24–7.40

(10 H, m). m/z (EI) = 493 (M+ + 2, 100), 491 (M+, 98), 464

(69), 462 (67), 413 (55), 384 (52) and 178 (42).

Synthesis of 8–11. To a THF solution (10 ml) of 3 (492 mg,

1.0 mmol) were added boronic acid esters 4–7 (1.20 mmol),

tetrakis(triphenylphosphine)palladium(0) (60 mg, 0.05 mmol)

and a 2 M aqueous potassium carbonate solution (0.8 ml). The

mixture was refluxed (8: 12 h, 9: 20 h, 10: 20 h and 11: 20 h)

under an argon atmosphere. After cooling, chloroform

(100 ml) was added to the reaction mixture and it was then

filtered through Celite. The filtrate was next poured into

water (50 ml). The chloroform layer was washed with brine

(3 � 50 ml) and dried over anhydrous sodium sulfate. The

solvent was removed in vacuo and the crude product purified

by silica gel column chromatography (8: chloroform–hexane =

4 : 3 � 1, chloroform � 1; 9: chloroform–hexane = 1 : 1 � 1,

chloroform–hexane = 2 : 5 � 1; 10: chloroform–hexane =

1 : 1 � 1, chloroform–hexane = 3 : 5 � 2; 11: chloroform–

hexane = 8 : 11 � 1, chloroform–hexane = 1 : 3 � 2) to give

8–11 as a yellow solid. The physical and spectral data are

shown below.

8. Yield 72%, mp 204–206 1C. dH (400 MHz, CDCl3,

Me4Si): 1.46–1.51 (1 H, m), 1.62–1.67 (1 H, m), 1.79–1.94

(3 H, m), 2.00–2.07 (1 H, m), 3.81–3.86 (1 H, m), 4.70–4.73

(1 H, m), 6.93 (1 H, s), 6.98–7.06 (6 H, m), 7.15 (1 H, s),

7.16–7.17 (1 H, m) and 7.26–7.39 (12 H, m). m/z (EI) = 495

(M+, 100), 466 (24) and 248 (8).


Me4Si): 1.39–1.47 (1 H, m), 1.55–1.59 (1 H, m), 1.69–1.84

(3 H, m), 1.91–2.00 (1 H, m), 3.69–3.73 (1 H, m), 4.58–4.61

(1 H, m), 6.91–7.03 (8 H, m), 7.08–7.10 (3 H, m) and 7.21–7.34

(12 H, m). m/z (FAB) = 578 (MH+).


Me4Si): 1.48–1.59 (1 H, m), 1.62–1.66 (1 H, m), 1.82–1.93

(3 H, m), 2.03–2.06 (1 H, m), 3.80–3.88 (1 H, m), 4.67–4.78

(1 H, m), 6.94 (1 H, s), 6.98–7.11 (10 H, m), 7.17 (1 H, d,


J = 3.4 Hz), 7.21 (1 H, d, J = 4.8 Hz) and 7.27–7.41

(12 H, m). m/z (FAB) = 660 (MH+).

11. Yield 56%, mp 57–61 1C. dH (400 MHz, CDCl3, Me4Si):

0.89 (3 H, t, J = 6.8 Hz), 1.30–1.34 (6 H, m), 1.42–1.52

(1 H, m), 1.59–1.64 (3 H, m), 1.78–1.90 (3 H, m), 1.97–2.07

(1 H, m), 2.58 (2 H, t, J = 7.6 Hz), 3.78–3.82 (1 H, m),

4.66–4.69 (1 H, m), 6.73 (1 H, s), 6.92–7.04 (7 H, m) and

7.22–7.40 (12 H, m). m/z (FAB) = 580 (MH+).

Synthesis of 12–15. To DMF (4 ml) was added phosphorous

oxychloride (352 mg, 2.30 mmol) at 0 1C. To the solution was

then added a DMF solution (14 ml) of 8–11 (0.84 mmol) at

0–5 1C. The mixture was heated at 75 1C (12: 2 h, 13: 4 h, 14:

20 h and 15: 2 h). After the reaction was complete, the reaction

mixture was poured into ice–water (100 ml) and neutralized

with aqueous sodium hydroxide. The product was extracted

with chloroform (3 � 50 ml). The extract was washed with

brine (2 � 50 ml) and water (2 � 50 ml), and dried over

anhydrous sodium sulfate. The solvent was removed in vacuo

and the product purified by silica gel column chromatography

(2 � chloroform) to afford 12–15 (12: orange solid, 13: red

solid, 14: red solid and 15: orange solid). The physical and

spectral data are shown below.


Me4Si): 1.47–1.53 (1 H, m), 1.65–1.67 (1 H, m), 1.82–1.89

(3 H, m), 2.07–2.11 (1 H, m), 3.81–3.85 (1 H, m), 4.76–4.79

(1 H, m), 6.94 (1 H, s), 6.96 (1 H, d, J = 8.5 Hz), 7.01 (2 H, d,

J = 8.8 Hz), 7.05 (2 H, d, J = 8.8 Hz), 7.23–7.39 (13 H, m),

7.68 (1 H, d, J = 4.1 Hz) and 9.80 (1 H, s). m/z (EI) = 523

(M+, 100), 495 (56), 373 (52) and 344 (31).


Me4Si): 1.44–1.53 (1 H, m), 1.59–1.66 (1 H, m), 1.77–1.91

(3 H, m), 2.00–2.09 (1 H, m), 3.79–3.81 (1 H, m), 4.70–4.73

(1 H, m), 6.93 (1 H, s), 6.95 (1 H, d, J = 8.5 Hz), 6.99 (2 H, d,

J = 8.9 Hz), 7.03 (2 H, d, J = 8.9 Hz), 7.08 (1 H, d, J =

3.9 Hz), 7.18 (1 H, d, J = 3.9 Hz), 7.24–7.40 (13 H, m), 7.62

(1 H, d, J = 4.1 Hz) and 9.82 (1 H, s). m/z (FAB) =

606 (MH+).


Me4Si): 1.46–1.54 (1 H, m), 1.62–1.68 (1 H, m), 1.79–1.86

(3 H, m), 1.98–2.02 (1 H, m), 3.77–3.99 (1 H, m), 4.46–4.69

(1 H, m), 6.38–7.40 (23 H, m), 7.57–7.58 (1 H, m) and 9.78

(1 H, s). m/z (FAB) = 688 (MH+).

15. Yield 96%, mp 66–68 1C. dH (400 MHz, CDCl3, Me4Si):

0.89 (3 H, t, J = 7.0 Hz), 1.30–1.39 (6 H, m), 1.45–1.48

(1 H, m), 1.61–1.73 (3 H, m), 1.79–1.90 (3 H, m), 2.02–2.05

(1 H, m), 2.91 (2 H, t, J = 7.0 Hz), 3.79–3.83 (1 H, m),

4.73–4.77 (1 H, m), 6.93 (1 H, s), 6.94 (1 H, d, J = 8.2 Hz),

7.00 (2 H, d, J = 8.9 Hz), 7.04 (2 H, d, J = 8.9 Hz), 7.06

(1 H, s), 7.23–7.40 (12 H, m) and 9.95 (1 H, s). m/z (FAB) =

608 (MH+).

Synthesis of 20, 21, 22 and 23. In the cases of 20, 21 and 23,

to an acetonitrile solution (6 ml) of 12, 13 and 14 (0.30 mmol)

were added cyano acetic acid (100 mg, 1.18 mmol)

and piperidine (46 mg, 0.54 mmol). In the case of 22, to

an acetonitrile–chloroform (1 : 1) mixed solution (80 ml) of

14 (107 mg, 0.16 mmol) were added cyano acetic acid (97 mg,

1.14 mmol) and piperidine (776 mg, 9.11 mmol). The mixture

was then refluxed (20: 2 h, 21: 2 h, 22: 23 h and 23: 17 h). After

cooling, the solvent was removed in vacuo and the residue

dissolved in chloroform. To the solution was added 1 M

aqueous hydrochloric acid (0.4 ml) and water (50 ml), and

the mixture stirred at room temperature for 30 min. The

chloroform layer was separated, washed with water (3 � 50 ml)

and dried over anhydrous sodium sulfate. The solvent was

removed in vacuo and the product purified by silica gel column

chromatography (20: chloroform–methanol = 8 : 1 � 1, 10 :

1 � 2; 21: chloroform–methanol = 10 : 1 � 3; 22:

chloroform–methanol = 10 : 1 � 1, 8 : 1 � 3; 23: chloroform–

methanol = 10 : 1 � 3) to afford 20, 21, 22 and 23 as a purple

solid. The physical and spectral data are shown below.

20. Yield 64%, mp 268–271 1C. dH (400 MHz, DMSO-d6,

Me4Si): 1.23–1.32 (1 H, m), 1.59–1.68 (2 H, m), 1.79–1.84 (2 H,

m), 1.99–2.08 (1 H, m), 3.82–3.86 (1 H, m), 4.86–4.90 (1 H, m),

6.97 (1 H, d, J = 8.7 Hz), 7.02 (2 H, d, J = 8.5 Hz), 7.07

(1 H, s), 7.11 (2 H, d, J = 8.5 Hz), 7.19–7.22 (2 H, m),

7.28–7.36 (5 H, m), 7.41–7.48 (4 H, m), 7.57–7.59 (2 H, m),

7.95 (1 H, d, J = 3.4 Hz) and 8.43 (1 H, s). m/z (FAB) =

591.2106 (MH+, C39H31N2O2S requires 591.2106).


Me4Si): 1.28–1.35 (1 H, m), 1.58–1.69 (2 H, m), 1.81–1.86

(2 H, m), 2.00–2.07 (1 H, m), 3.81–3.85 (1 H, m), 4.81–4.84

(1 H, m), 6.98 (1 H, d, J = 9.2 Hz), 7.00 (2 H, d, J = 9.1 Hz),

7.06 (1 H, s), 7.09 (2 H, d, J = 9.1 Hz), 7.20–7.21 (2 H, m),

7.28–7.53 (13 H, m), 7.82 (1 H, br s) and 8.28 (1 H, br s). m/z

(FAB) = 673.1974 (MH+, C43H33N2O2S2 requires 673.1983).


Me4Si): 1.29–1.36 (1 H, m), 1.61–1.70 (2 H, m), 1.84–1.87

(2 H, m), 2.00–2.05 (1 H, m), 3.81–3.85 (1 H, m), 4.80–4.83

(1 H, m), 6.98 (1 H, d, J = 6.3 Hz), 7.00 (2 H, d, J = 7.8 Hz),

7.06 (1 H, s), 7.08 (2 H, d, J = 7.8 Hz), 7.20–7.22 (2 H, m),

7.28–7.47 (13 H, m), 7.60–7.62 (2 H, m), 7.98 (1 H, d, J=3.4 Hz)

and 8.49 (1 H, s). m/z (FAB) = 755.1831 (MH+,

C47H35N2O2S3 requires 755.1861).


Me4Si): 0.84 (3 H, t, J = 6.0 Hz), 1.20–1.30 (7 H, m),

1.57–1.65 (4 H, m), 1.76–1.84 (2 H, m), 1.97–2.06 (1 H, m),

2.74 (2 H, t, J = 7.2 Hz), 3.78–3.82 (1 H, m), 4.82–4.86

(1 H, m), 6.93 (1 H, d, J = 8.2 Hz), 6.99 (2 H, d, J = 8.6 Hz),

7.04 (1 H, s), 7.07 (2 H, d, J = 8.6 Hz), 7.17–7.19 (2 H, m),

7.26–7.34 (5 H, m), 7.38–7.45 (4 H, m), 7.50 (1 H, s), 7.54

(1 H, s) and 8.25 (1 H, s). m/z (FAB) = 675.3117 (MH+,

C45H43N2O2S requires 675.3045).

Synthesis of 24, 25, 26, 27 and 28. To an acetic acid solution

(4 ml) of 12–14 (0.30 mmol) was added rhodanine derivatives

17–19 (0.32 mmol). The mixture was heated at 120 1C and

ammonium acetate (0.88 mmol) added, after which it

was refluxed for 2 h. After cooling, the reaction mixture was

poured into water (20 ml). The resulting precipitate was

filtered and washed with water, and the crude product purified

by silica gel column chromatography (24: chloroform–methanol=

8 : 1 � 3; 25: chloroform–methanol = 8 : 1 � 2; 26:


chloroform–methanol = 8 : 1 � 9; 27: chloroform–methanol =

8 : 1 � 5; 28: chloroform–methanol = 8 : 1 � 3) to afford

compound 24, 25, 26, 27 and 28 as a purple solid. The physical

and spectral data are shown below.


Me4Si): 1.29–1.33 (1 H, m), 1.59–1.68 (2 H, m), 1.81–1.84

(2 H, m), 2.00–2.05 (1 H, m), 3.81–3.85 (1 H, m), 4.69 (2 H, s),

4.84–4.88 (1 H, m), 6.95 (1 H, d, J= 8.5 Hz), 7.01 (2 H, d, J=

8.3 Hz), 7.06 (1 H, s), 7.09 (2 H, d, J = 8.3 Hz), 7.19–7.21

(2 H, m), 7.28–7.34 (5 H, m), 7.39–7.49 (4 H, m), 7.57–7.58

(2 H, m), 7.77 (1 H, d, J = 3.9 Hz) and 8.10 (1 H, s). m/z


25. Yield 31%, mp 4 300 1C. dH (400 MHz, DMSO-d6,

Me4Si): 1.14 (3 H, t, J = 6.9 Hz), 1.25–1.37 (1 H, m),

1.60–1.68 (2 H, m), 1.80–1.84 (2 H, m), 2.01–2.08 (1 H, m),

3.82–3.86 (1 H, m), 3.98–4.00 (2 H, m), 4.69 (2 H, s), 4.84–4.87

(1 H, m), 6.96 (1 H, d, J = 8.5 Hz), 7.00 (2 H, d, J = 9.3 Hz),

7.07 (1 H, s), 7.08 (2 H, d, J = 9.3 Hz), 7.20–7.55 (13 H, m),

7.68 (1 H, d, J = 3.9 Hz) and 8.02 (1 H, s). m/z (FAB) =

824.1757 (MH+, C46H38N3O4S4 requires 824.1745).


Me4Si): 1.24–1.29 (1 H, m), 1.25 (18 H, s), 1.59–1.87 (2 H, m),

1.85–1.87 (2 H, m), 2.01–2.09 (1 H, m), 3.83–3.87 (1 H, m),

4.61 (2 H, s), 4.85–4.88 (1 H, m), 5.19 (2 H, s), 6.98 (1 H, d,

J= 8.5 Hz), 7.00 (2 H, d, J= 8.8 Hz), 7.07 (1 H, m), 7.10 (2 H,

d, J= 8.8 Hz), 7.19–7.48 (14 H, m), 7.53 (1 H, s), 7.57 (1 H, d,

J = 3.9 Hz), 7.72 (1 H, d, J = 4.1 Hz) and 8.07 (1 H, s). m/z



Me4Si): 1.22–1.35 (1 H, m), 1.23 (18 H, s), 1.57–1.66 (2 H, m),

1.78–1.85 (2 H, m), 1.95–2.05 (1 H, m), 3.73–3.79 (1 H, m),

4.33 (2 H, s), 4.74–4.78 (1 H, m), 5.10 (2 H, s), 6.91 (1 H, d,

J = 8.5 Hz), 6.99 (2 H, d, J = 8.3 Hz), 7.03 (1 H, s), 7.04

(2 H, d, J = 8.3 Hz), 7.17–7.54 (18 H, m), 7.62 (1 H, s) and

7.94 (1 H, s). m/z (FAB) = 1080.3016 (MH+, C63H58N3O4S5requires 1080.3031).


Me4Si): 1.22–1.33 (1 H, m), 1.26 (18 H, s), 1.59–1.68 (2 H, m),

1.80–1.84 (2 H, m), 1.97–2.03 (1 H, m), 3.78–3.83 (1 H, m),

4.46 (2 H, s), 4.77–4.81 (1 H, m), 5.17 (2 H, s), 6.96 (1 H, d,

J = 8.5 Hz), 6.98 (2 H, d, J = 9.1 Hz), 7.05 (1 H, s), 7.06

(2 H, d, J = 9.1 Hz), 7.19–7.54 (20 H, m), 7.71 (1 H, s) and

8.02 (1 H, s). m/z (FAB) = 1162.2795 (MH+, C67H60N3O4S6requires 1162.2908).

Acknowledgements

This work was financially supported in part by Grants-in-Aid

for Science Research (no. 19550185) from the Japan Society

for the Promotion of Science (JSPS).

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9 K. Hara, T. Sato, R. Katoh, A. Furube, Y. Ohga, A. Shinpo,S. Suga, K. Sayama, H. Sugihara and H. Arakawa, J. Phys. Chem.B, 2003, 107, 597.

10 T. Yoshida, M. Iwaya, H. Ando, T. Oekermann, K. Nonomura,D. Schlettwein, D. Wohrle and H. Minoura, Chem. Commun.,2004, 400.

11 R. Katoh, A. Furube, T. Yoshihara, K. Hara, G. Fujihashi,S. Takano, S. Murata, H. Arakawa and M. Tachiya, J. Phys.Chem. B, 2004, 108, 4818.

12 J. D. Mee, US Pat., 5679795, 1997.13 G. Bidan, A. De Nicola, V. Enee and S. Guilerez, Chem. Mater.,

1998, 10, 1052.14 M. Matsui, M. Wang, K. Funabiki, Y. Hayakawa and

T. Kitaguchi, Dyes Pigm., 2007, 74, 169.


Gold imidazolium-based ionic liquids, efficient catalysts for

cycloisomerization of c-acetylenic carboxylic acids

Florentina Neat-u,a Vasile I. Parvulescu,a Veronique Michelet,b Jean-Pierre Genet,b

Alexandre Goguetcand Christopher Hardacre

cd

Received (in Durham, UK) 22nd July 2008, Accepted 8th September 2008


DOI: 10.1039/b812580e

Ionic liquid stabilized gold(III) chloride is shown to be a very active catalyst in the cyclization

of sterically hindered and unhindered acetylenic carboxylic acid substrates even in the absence

of a base.

Introduction

Ionic liquids (ILs) have received significant attention recently

because they exhibit several advantages over molecular

solvents with respect to their environmental impact.1 Catalytic

reactions in ILs have been examined for at least 20 years and

have been used for a wide variety of carbon–carbon bond

forming reactions, which were the first to be undertaken in this

media, and a number of good reviews cover the area of ILs.2–6

The extensive interest stems from the fact that the properties

of ionic liquids may be tuned in order to suit a particular

application by varying the cation–anion combination system-

atically and thereby are useful engineering solvents. In addi-

tion, for chemical reactions, the ionic liquid provides an ionic

environment which can significantly alter the reactivity and

selectivity of processes compared with molecular solvents.7

The transition metal-catalyzed cyclization of 4-alkynoic

acids constitutes a major route8 for the construction of

5-membered exocyclic lactones and has been the subject of a

large number of investigations.9–14 Recently, a very attractive

route to perform this reaction under mild conditions in the

presence of gold was reported.15,16 In our laboratory, the

catalytic properties of AuCl and AuCl3 for gem-substituted

substrates, in the absence of base, was described15 as well as

the use of two heterogeneous systems Au2O317 and Au/beta.18

These systems were found to be active for both substituted and

unsubstituted substrates. In addition, the optimized hetero-

geneously catalyzed system was found to be recyclable.22

Whilst gold has been shown to be highly active, it commonly

undergoes significant deactivation due to the ease by which it

may change its oxidation state and, in the case of nanoparti-

cles, the catalyst particle size leading to instability in the form

of the active catalyst.19 Therefore, modalities which can

enhance the stability are extremely important for practical

applications.20 These can be in the form of the solvent used or

the nature of the support.21 For example, stabilization of gold

in the form of nanoparticles in ionic liquids by imidazolium

derivatives has been reported.22–25

In this report the behavior of a 4 wt% Au/beta catalyst and

a series of 1,3-dialkylimidazolium tetrachloroaurate salts in

ionic liquids as catalysts for the cyclization of acetylenic

substrates has been studied. Gold has recently been reported

as a catalyst in ionic liquids in the hydration of phenyl-

acetylene26,27 and the syntheses of substituted 3(2H)-furanones,27

2,5-dihydrofurans28 and substituted indoles.29 However, with

the exception of the Co2(CO)8-catalyzed intramolecular and

intermolecular Pauson–Khand annelation using 1-butyl-3-

methylimidazolium hexafluorophosphate ([C4mim][PF6]) and

tetrafluoroborate ([C4mim][BF4]) ionic liquids as solvents, no

other related reactions to the cyclization of acetylenic sub-

strates have been reported in ILs, to date.30

Experimental

Catalysts preparation

1. Au-beta catalyst. The catalyst was prepared by stirring

1 g of beta zeolite (PQ Corporation) for 3 h with 100 cm3 of 1 M

NH4NO3 at 353 K.22 The slurry was filtered off and carefully

washed with deionized water, dried for 6 h at 333 K and

calcined for 24 h at 773 K. Deposition of gold was performed

using the deposition–precipitation method. The support

(1 g) was added to 100 cm3 of an aqueous solution of HAuCl4(2.1 � 10�3 M) at 343 K which had previously adjusted to a

pH = 8.5 with 0.2 M NaOH. The temperature of the slurry

was maintained at 343 K under vigorous stirring for 3 h.

Thereafter, the sample was filtered off, washed with deionized

water to remove the free chloride and then dried under

vacuum at 333 K for 24 h. The resultant catalyst contained

4 wt% Au as determined by ICP-AES analysis. This catalyst

has been used as a reference material in these experiments.

2. Ionic liquids synthesis. Gold ionic liquids were

synthesized using the method reported previously by Hasan

et al.25 Four 1-alkyl-3-methylimidazolium tetrachloroaurate

([Cnmim][AuCl4], n = 2, 4, 6, 18) ionic liquids were prepared

by adding a 10% molar excess of [Cnmim]Cl to tetrachloroauric

aDepartment of Chemical Technology and Catalysis, University ofBucharest, B-dul Regina Elisabeta 4-12, 030016 Bucharest, Romania.E-mail: [email protected]; Fax: +40 21 4010241

b Laboratoire de Synthese Selective Organique et Produits Naturels,ENSCP, UMR 7573, 11 rue P. et M. Curie, F-725231 Paris Cedex 05,France. E-mail: [email protected]; Fax: +33 1 44071062

c CenTACat School of Chemistry and Chemical Engineering, Queen’sUniversity, Stranmillis Road, Belfast, Northern Ireland,UK BT9 5AG

dQUILL School of Chemistry and Chemical Engineering, Queen’sUniversity, Stranmillis Road, Belfast, Northern Ireland, UK BT9 5AG.E-mail: [email protected]; Fax: +44 28 90974687



acid (HAuCl4�4H2O, Alfa Aesar) resulting in the rapid

formation of a yellow solid. Following heating to above

its melting points with stirring for 0.5 h, the product was

purified by recrystallization from benzene–acetonitrile in the

volume ratio of 4 : 1. [Cnmim]Cl were prepared in house using

standard literature methods.31 Trihexyltetradecylphospho-

nium hydrogen sulfate ([P66614][HSO4]) was supplied by Cytec.

1-Butyl-3-methylimidazolium bis{(trifluoromethyl)sulfonyl}-

imide ([C4mim][NTf2]) and trihexyltetradecylphosphonium

bis{(trifluoromethyl)sulfonyl}imide ([P66614][NTf2]) were

formed by metathesis from [C4mim]Cl and [P66614]Cl (Cytec),

respectively, according to literature methods.32 The base func-

tionalized ionic liquid 5-diisopropylamino-3-oxapentyl)-

dimethylethylammonium bis{(trifluoromethyl)sulfonyl}imide

(BIL) shown in Fig. 1 was synthesized as previously re-

ported.33 In all cases prior to reaction the ionic liquids were

dried under vacuum at 50 1C overnight. All ionic liquids

contained o0.16 wt% water determined by Karl–Fischer

analysis and o5 ppm halide by suppressed ion chromato-

graphy. All other reagents were used as received.

Catalyst characterization

The solid catalysts were characterized by nitrogen adsorption–

desorption isotherms at 77 K (Micromeritics ASAP 2000) after

out-gassing the samples at 393 K for 12 h.

The XPS spectra were recorded using a Kratos Axis

UltraDLD spectrometer with monochromatic Al-Ka radia-

tion. The data were analyzed using Casa-XPS (v2.3.13)

employing a Shirley-background subtraction prior to fitting.

EXAFS data were collected at the Synchrotron Radiation

Source in Daresbury, UK, using station 9.3. The spectra were

recorded at the Au LIII edge using a double crystal Si(111)

monochromator. Scans were collected and averaged. Data

were processed using EXCALIB which was also used to

convert raw data into energy vs. absorption data. EXBROOK

was used to remove the background. The analysis of the

EXAFS was performed using EXCURV98.34 The gold con-

centration was determined by ICP-AES.

Catalytic tests

Typically 0.26 mmol of acetylenic acid was stirred with 2.5 mol%

[Cnmim][AuCl4] dissolved in 0.5 g [Cnmim]Cl in air at

room temperature, until completion of the reaction. After

the completion of the reaction, the reaction products

were extracted three times with diethyl ether and the

solvent completely removed under vacuum to give the

corresponding lactone. 1H and 13C NMR were recorded

on a Bruker AV 300 instrument operating at 300 Hz to

identify the products. The measured NMR spectra for

3-phenyl-5-methylene-g-butyrolactone (2a), 3-n-butyl-3-ethoxy-

carbonyl-5-methylene-g-butyrolactone (2b), 3-methoxycarbonyl-

5-methylene-3-(3 0-phenylprop-2 0-enyl)-g-butyrolactone (2c),3-allyl-3-methoxycarbonyl-5-methylene-g-butyrolactone (2d), 3-

benzyl-3-ethoxycarbonyl-5-methylene-g-butyrolactone (2e) and

3-methoxycarbonyl-5-methylene-g-butyrolactone (2f) were in

good agreement were those reported previously.19–21 In all

cases, no cyclization is observed in the absence of any of the

gold catalysts irrespective of the solvent used.


Catalysts characterization

The beta zeolite used in these experiments had a surface area

of 464 m2 g�1 and a pore volume of 0.96 cm3 g�1. After the

deposition of gold (4 wt%) the surface area decreased to

383 m2 g�1 and the pore volume to 0.80 cm3 g�1. TEM

analysis of this material shown an uniform size distribution

with an average of 3 nm. The characterization of IL-stabilized

gold(III) chloride was examined using TEM, XPS and

EXAFS. Fig. 2 shows the XPS spectra of the Au 4f photo-

electron emission corresponding to the mixture of the catalyst

([C6mim][AuCl4]) with the ionic liquid [C6mim]Cl after the

reaction. The binding energies of Au 4f7/2 and Au 4f5/2 levels

were located at 90.0 eV and 86.3 eV, respectively. EXAFS of

[C6mim][AuCl4] dissolved in [C6mim]Cl showed a single peak

at 0.22 nm associated with 4 chlorine atoms in the first

coordination shell. During reaction, this peak decreases

slightly and a small decrease in the white line of the XANES

is observed. This variation may indicate that chlorine is being

replaced by a lighter element such as coordination by the

substrate during reaction, as would be expected. TEM analysis

is in agreement with the XPS and EXAFS measurements and

showed no nanoparticle formation. These observations are

Fig. 1 Schematic of the cation of the basic ionic liquid (BIL) used as

solvent for the cyclization of the functionalized acetylenic substrates

1a and 1b.

Fig. 2 The XPS spectra in the Au 4f region of the [C6mim][AuCl4]

after reaction.


consistent with the presence of a dissolved Au(III) species with

B4 chlorines in the first coordination sphere.35

Catalytic tests

Table 1 shows the activity of Au/beta in a range of ionic

liquids for the cyclization of the functionalized acetylenic

substrates 1a and 1b.

Although cycloisomerization of substrate 1a did occur in the

BIL and [P66614][HSO4] with conversions of 75 and 70%,

respectively, (Table 1, entries 1, 2), the ILs could not be

separated from the reaction products due to the high solubility

of the ionic liquid–substrate mixture in diethyl ether. In the

case of the reaction performed in [C4mim][NTf2] and

[P66614][NTf2], the ionic liquid could only be partially sepa-

rated from the reactants and products and showed conversion

of the substrate of 80% and 70%, respectively (Table 1, entries

3, 4). In each case, the conversions were obtained from NMR

determination in the ionic liquid. [C6mim]Cl was found to

separate efficiently from diethyl ether and the cycloisomeriza-

tion of 1a (Table 1, entry 5) resulted in the desired compound

2a in 79% isolated yield.

Due to the ease of workup [C6mim]Cl was also examined as

a medium for the cyclization of 1b over Au/beta. From an

analysis of the 1H NMR, a conversion of 91% was obtained

with an isolated yield of the lactone of 58% (Table 1,

entry 6). Similar activities and selectivities were also found in

conventional molecular solvents, such as acetonitrile (Table 1,

entry 7).

The heterogeneously catalyzed reaction results were com-

pared with the use of the ionic liquid as a catalyst in the form

of the tetrachloroaurate based ionic liquids. Since [C6mim]-

[AuCl4] is a solid at room temperature, [C6mim]Cl was used as

solvent. The results are summarised in Table 2.

To date, it has only been possible to exclude a base from the

homogeneous reaction conditions if the two substituents on a

tetrahedral centre were of a significant size, as understood by

the Thorpe–Ingold effect.36 In the case of the ionic liquid

catalyst system, all the g-acetylenic carboxylic acids examined

were cleanly transformed to the corresponding g-alkylideneg-butyrolactones. Moreover, similar activity and selectivity

was found for the ionic liquid catalyst compared with

the homogeneous AuCl catalysts system in acetonitrile.15

Irrespective of the alkenyl side chain length the lactones were

isolated in 85–96% yields (Table 2, entries 1–5) even at room

temperature. In all cases, the catalytic amount of gold ionic

liquid used in these reactions was equivalent to the amount

used under heterogeneous conditions. No side reactions were

observed on the alkenyl side chains during the course of the

reaction.

Using the gold based ionic liquid system, complete trans-

formations of 2-prop-2-ynylmalonic acid monomethyl ester 1f

with an 84% isolated yield (Table 2, entry 6) and of 2-phenyl-

pent-4-ynoic acid 1a with an 96% isolated yield (Table 2, entry 7)

after 1 h at room temperature were obtained. Similar results

were found for the reaction of 1a in the presence of [C2mim]-

[AuCl4], [C4mim][AuCl4] and [C18mim][AuCl4] at room

temperature. These results for unsubstituted substrates are

comparable with those obtained under heterogeneous Au2O3

conditions (3 h),21 or homogeneous AuCl/K2CO3 conditions

(2 h)20 and show significant advantages over other homo-

geneous gold chloride based catalysts.19 In the case of the

homogeneous catalysts, the formation of degradation pro-

ducts or the corresponding methylketone was observed for

the sterically unhindered substrates and the reaction only

occurred in the presence of a base. In the ionic liquid catalyzed

reactions, excellent reactivity was found without the need for

additives. Furthermore, the IL-catalysts were recycled three

times without any loss in conversion or yields.

A comparison of the homogeneous reaction (Table 2, entries

7 and 1) with that of the heterogeneous reaction (Table 1,

entries 5 and 6) indicates that the former is more active

requiring a lower temperature and resulting in a higher con-

version/yield after 1 h. This is reflected in the turnover

frequencies (TOF) for the homogeneous catalysts compared

with the supported catalysts which are B19.8 h�1 (at RT) and

B1.3 h�1 (at 40 1C), respectively, for the formation of 2b, for

example. The lower rate is unlikely to be due to mass transfer

Table 1 Cyclization of functionalized acetylenic substrate overAu/beta in a range of ionic liquids

Entry R1 R2 Solvent Product Yielda (Conv.) (%)

1 Ph H BIL 2a (75)b

2 Ph H [P66614][HSO4] 2a (70)b

3 Ph H [P66614][NTf2] 2a (70)b

4 Ph H [C4mim][NTf2] 2a (80)b

5 Ph H [C6mim]Cl 2a 79 (83)6 CO2Et n-Bu [C6mim]Cl 2b 58 (91)7 CO2Et n-Bu CH3CN 2b 60 (90)

a Isolated yield. b IL inseparable from the system.

Table 2 Au-catalyzed cyclization of functionalized carboxylic acids

Entry R1 R2 Product Temp./1C Time/h Yielda (%)

1 CO2Et n-Bu 2b RT 1 962 CO2Me Cinnamyl 2c 40 2 903 CO2Me Allyl 2d RT 2 914 CO2Et Bn 2e RT 1 855 CO2Et Bn 2e RT 2 956 CO2Me H 2f RT 1 847 Ph H 2a RT 1 96

a Isolated yield.


limitations in the heterogenous case which has been observed

in other solid catalyzed reactions in ionic liquids37 as the

acetonitrile reaction (Table 1, entries 6 and 7) also shows

reduced activity. The higher TOF may be expected due to the

inaccessibility of the bulk gold atoms; however, even taking

into account the lower dispersion of the heterogeneous cata-

lyst, which contains 3 nm gold particles, the homogeneous

catalyst has a higher intrinsic activity compared with the

heterogeneous system.

Scheme 1 shows a proposed mechanism for the gold based

ionic liquid catalyzed reaction. From the XPS and EXAFS,

the active catalyst is thought to be Au3+ which activates the

acetylenic group.20 The ionic liquid provides a medium which

eliminates the use of a base. This role of the ionic liquid may

be to shift the acid equilibrium of the starting material to

favour the carboxylate anion, either by hydrogen bonding

with the chloride, for example, or via stabilizing the anion in

the ionic environment. This increase in concentration of the

active intermediate allows reaction and activation of the

carbon–carbon triple bond by the Au3+ centre and nucleo-

philic addition to the alkyne. After the cyclization a hydrogen

mediated demetalation ends the catalytic cycle. It should be

noted that the scheme does not preclude the possibility that the

cation of the ionic liquid may play an important role in

creating the active catalyst. For example, imidazolium based

ionic liquids have been shown to form carbenes with homo-

geneous catalysts in C–C bond forming reactions.10

Conclusions

Gold chloride based ionic liquids have been shown to be very

active catalysts in the cyclization of acetylenic substrates. The

ionic liquid medium prevents nanoparticle formation via

stabilization of the gold in the form of isolated species by

chlorine coordination. The ionic liquid also allows activation

of the carboxylic acid in the starting material and provides a

medium which eliminates the need for added base in order to

cyclize unhindered acetylenic carboxylic acid.

Acknowledgements

This work was financially supported by the Consiliul National

al Cercetarii Stiintifice din Invatamantul Support (CNCSIS)

and the Centre National de la Recherche Scientifique (CNRS),

a Portfolio partnership from the EPSRC and an EU transna-

tional grant. CCLRC are thanked for providing EXAFS

beamtime.

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Scheme 1 The mechanism of cyclization of acetylenic carboxylic acid

using Au3+ derived catalysts. It should be noted that representation of

the gold as [Au] is for illustrative purposes only and will exist in the

form of a [AuClx]y+ complex.


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Magnetically moveable bimetallic (nickel/silver) nanoparticle/carbon

nanotube composites for methanol oxidationw

Guan-Ping Jin,*ab Ronan Baron,a Neil V. Rees,a Lei Xiaoa

and Richard G. Compton*a

Received (in Durham, UK) 22nd August 2008, Accepted 18th September 2008

First published as an Advance Article on the web 31st October 2008

DOI: 10.1039/b814630f

Multi-walled carbon nanotubes (CNTs) functionalized both by nickel and silver nanoparticles

were obtained using a single step chemical deposition method in an ultrasonic bath. The new

composite material was characterized by means of scanning electron microscopy (SEM), X-ray

diffraction (XRD) and cyclic voltammetry (CV). The electroactivity of the bi-functionalized CNTs

multi-walled carbon nanotubes was assessed in respect to the electrooxidation of methanol. It was

found that the carbon nanotube supported silver nanoparticles have significantly higher catalytic

properties than the bulk metal of the same surface area. Furthermore, it was shown that the

presence of only a very small proportion of magnetic nickel nanoparticles (1.5% of the total

number of metallic nanoparticles) allows the bi-functionalized carbon nanotubes to be moved

magnetically in solution, making them easily recoverable after use whilst keeping an optimal

electrocatalytic surface area.

1. Introduction

The synthesis of magnetic nanomaterials has attracted con-

siderable attention in the last few years. One of the reasons for

that interest is that magnetic material can be moved using a

magnet and which makes the considered material recoverable

both for economical and environmental issues.1–4 Nanomater-

ials that can be magnetically driven are also very promising

for medical applications, such as drug delivery5 or complex

biomanipulations.6

The functionalization of carbon materials with magnetic

particles is particularly attractive due to the properties and

wide use of carbon materials. In particular, magnetic nano-

particles have been synthesized on the surface of carbon

nanotubes using various different methodologies.6–10

As far as the use of magnetic nanomaterials in electro-

chemistry is concerned, there have been only a limited number

of reports on the matter. Most notably Willner et al. studied

the use of hydrophobic magnetic nanoparticles capable of

blocking an electrode surface11–13 and Wang et al. addressed

the use of magnetic nickel nanoparticles both for on-demand

control of electrocatalytic processes and to reduce electrode

surface fouling.14–16

A lot of interest is devoted to methanol electrocatalytic

oxidation as it can be used in fuel cells and both bulk silver and

silver nanoparticles are known to be efficient electrocatalysts

for methanol oxidation in alkaline solutions17,18 with less

poisoning observed than at platinum materials.19 However,

even though silver is much less expensive than platinum its

cost remains an issue.

In order to get the benefits of both the magnetic properties

of magnetic nanoparticles and the catalytic properties of

AgNPs we designed a hybrid material, which contains both.

NiNPs and AgNPs were synthesized on CNTs using a one-

pot chemical deposition in an ultrasonic bath using a

methodology recently developed in our laboratory for the

synthesis of NiNPs on glassy carbon microspheres.20 To

date, this publication is the first report of the bi-functionaliza-

tion of CNT with NiNPs and AgNPs. In addition, the

electrocatalytic oxidation of methanol at the new hybrid

material was studied.

2. Experimental

2.1 Reagents and equipment

Bamboo-like multi-walled carbon nanotubes (CNTs, diameter

30 � 10 nm, 5–20 mm length, o95% purity) were purchased

from NanoLab (Brighton, MA, USA). Nickel(II) chloride

(NiCl2, 99.9%) was obtained from Alfa Aesar (Heysham,

UK). L-Ascorbic acid (99.7%) and silver nitrate were supplied

by BDH (Poole, UK). Nafion was purchased from Aldrich

(Poole, UK). Acetonitrile (ACN) was supplied by Sigma-

Aldrich (Gillingham, UK). All the reagents were used without

further purification. All solutions were prepared using purified

water from Vivendi UHQ grade water system with a resistivity

of not less than 18.2 MO cm.

Electrochemical measurements were recorded using an

Autolab PGSTAT 30 computer-controlled potentiostat with

a standard three-electrode setup. Either a home-made 4 mm

aDepartment of Chemistry, Physical and Theoretical ChemistryLaboratory, Oxford University, South Parks Road, Oxford,UK OX1 3QZ. E-mail: [email protected];Fax: 0044-1-865 275410; Tel: 0044-1-1865 275413

bAnhui Key Laboratory of Controllable Chemistry Reaction &Material Chemical Engineering, School of Chemical Engineering,Hefei University of Technology, Hefei, 230009, P. R. China.E-mail: [email protected]; Fax: 0086-0551-2902450;Tel: 0086-551-2901450

w Electronic supplementary information (ESI) available: Videos S1–3.See DOI: 10.1039/b814630f



diameter disc paraffin-impregnated graphite electrode21 or a

0.08 mm diameter disk bulk silver served as working electro-

des. Paraffin-impregnated graphite electrodes have similar

behaviour than other common graphite electrodes and have

been chosen here because they are easy to fabricate and it is

easy to renew their surface by a polishing step. A platinum

wire was used as a counter electrode, and a silver wire used as

the reference electrode completed the cell assembly. The

paraffin-impregnated graphite electrode surface was renewed

by successive mechanical polishing steps on alumina powders

(Micropolish II, Buehler) of 1 to 0.3 mm in diameter. The

electrode was sonicated for 5 min in deionized water after each

polishing step. All experiments were carried at a temperature

of 20 � 1 1C. All the solutions were degassed with nitrogen

prior to the electrochemical recordings.

Scanning electron microscopy (FEG-SEM, tungsten fila-

ment as electron source, acceleration voltage 20 keV) images

and energy dispersion X-ray spectra analysis were performed

using a JEOL 6300 F instrument. X-Ray diffraction patterns

(XRD) were collected on a PANalytical X’Pert instrument

with 40 kV and 40 mA settings.

Sonication was obtained using a D-78224 Singen/Htw

sonicator (50/60 Hz, 80 W, UK).

2.2 Ultrasonic synthesis of silver and nickel nanoparticles

on CNTs

The nickel and silver nanoparticles were synthesized onto

the surface of CNTs using the following protocol: The CNTs

were sonicated in conc. HClO4 + HNO3 (3:7, v:v) for

7 h in order to oxidize their surface, they were then filtered

and extensively washed with deionized water to pH 7, and

dried in air. Then, 2.9 mg NiCl2, 1.7 mg AgNO3 and 2.0 mg

oxidized CNTs were added to 60 mL of acetonitrile in an

airtight glass flask. The mixture was sonicated for one hour.

4.0 mg of L-ascorbic was then added in the flask and the pH

was adjusted to 5.2 using 1 M NaOH. The reaction was

allowed to proceed for 5 min at 65 1C under sonication.

Finally, the products were separated by centrifugation,

washed with acetonitrile and deionized water to remove any

unreacted species. The multi-walled carbon nanotubes deco-

rated with silver and nickel nanoparticles (AgNPs,NiNPs/

CNTs) were allowed to air-dry for 24 h prior to use. Multi-

walled carbon nanotubes decorated only with silver (AgNPs/

CNTs) or nickel nanoparticles (NiNPs/CNTs) were obtained

using the same method.

2.3 Modification of the paraffin-impregnated graphite

electrodes with CNTs

Films of CNTs on the surface of paraffin-impregnated

graphite electrodes were obtained as follows: 2 mg of CNTs

decorated with nanoparticles was suspended in 2 mL of

Nafion (0.05%) and acetonitrile solution to form a ‘‘casting’’

suspension. The casting suspension was then briefly sonicated

for 2 min in order to disperse the CNTs decorated with

nanoparticles. Some of the suspension was then pipetted onto

the surface of a freshly polished paraffin-impregnated graphite

electrodes and let to dry in air.

2.4 Movement of the AgNPs,NiNPs/CNTs composite

material

The AgNPs,NiNPs/CNTs (AgNPs/CNTs, NiNPs/CNTs)

composite materials were pipetted onto a 1.0 mm thick glass

surface with some water. A magnet (NdFeB alloy rod magnet,

purchased from e-magnets UK Ltd, Sheffield, UK) was posi-

tioned directly below the glass surface and moved by an

electronic motor. The directed movement of the nanocom-

posites was observed and recorded using an optical microscope

with a Bressler Visiomar video camera (160� magnification

using a 320 � 240 pixel frame).


3.1 Synthesis and microscopic characterization of the

nanoparticle-modified CNTs

The synthesis of the silver and nickel nanoparticles on the

surface of the CNTs was obtained following the steps noted in

Scheme 1. First the CNTs are treated with concentrated nitric

and perchloric acids to generate carboxylic groups on their

edges and defects. The negatively charged sites chelate silver

and nickel cations added to the solution. It is then expected

that the addition of ascorbic acid as a mild reducing agent in

the presence of ultrasound results in the production of small

nickel and silver nanostructures on the surface of the CNTs.

Similar experimental routes were followed for the synthesis of

nickel or silver nanoparticles separately on the CNTs.

A scanning electronic microscopy analysis of the samples

reveals that silver and nickel nanoparticles of 100 nm in

diameter in average are obtained on the CNTs (Fig. 1). The

EDX (Fig. 2) and XRD spectra of the samples (Fig. 3) confirm

the presence of both silver and nickel nanoparticles. It can be

noticed that the peaks for Ni (111) and Ag (111) are larger

than the other peaks, which reveals that the most common

nanoparticles on the CNTs are face-centered cubic (fcc)

nickel and silver. The average crystallite size calculated using

Scherrer’s equation from the width at half peak maximum for

the NiNPs is 30 � 25, and 25 � 15 nm, respectively for the

Scheme 1 Preparation of multi-walled carbon nanotubes (CNTs)

functionalized both by nickel and silver nanoparticles using a single

step chemical deposition method in an ultrasonic bath and subsequent

immobilization of the new composite material on an electrode surface.


NiNPs/CNTs and the AgNPs,NiNPs/CNTs and 20 � 10 nm

for the AgNPs of the AgNPs/CNTs and the AgNPs,NiNPs/

CNTs. It can be further stated on the nature of the silver and

nickel nanoparticles that those particles are distinct and do not

crystallise in the form of alloys, as it is well known that Ni and

Ag are immiscible in both the solid and liquid phases.22

3.2 Electrode modification and characterization

The electrode surfaces were modified by the decorated CNTs

in a Nafion film by following the protocol described in the

Experimental section. The electrode surfaces can be charac-

terized electrochemically by oxidizing the metallic nanoparti-

cles. Some metal oxides have a well-defined and specific

reduction potential and the corresponding reduction peak

can be used to estimate the surface area of specific metals.

Fig. 4 shows the cyclic voltammograms that were obtained in

0.1 M NaOH for the various modified electrodes. It can

be observed that the voltammogram corresponding to the

AgNPs,NiNPs/CNTs modified electrode corresponds to the

superposition of the characteristic features of both the AgNPs/

CNTs and the NiNPs/CNTs modified electrodes. Using the

literature values of 790 and 270 mC cm�2 for the charge passed

per unit area of surface area of nickel and silver,23–25 we can

estimate the total surface area of each of the metals for the

AgNPs,NiNPs/CNTs/Nafion material. The average loading

of the nickel and silver nanostructures on the CNTs were

estimated to be, respectively in the order of 3.4 � 10�2 and

2.2 cm2 mg�1. Which then shows that, for nanoparticles of

about the same size, the NiNPs represent only 1.5% of the

total number of nanoparticles.

3.3 Electrocatalysis

The electroactivity of the mono- and bi-functionalized multi-

walled carbon nanotubes was assessed and compared with the

Fig. 1 SEM images of the multi-walled carbon nanotubes functionalized both by nickel and silver nanoparticles (AgNPs,NiNPs/CNTs). The

square on Fig. 1(A) indicates where the EDX spectrum in Fig. 2 was obtained.

Fig. 2 EDX spectra of (A) AgNPs,NiNPs/CNTs, (B) NiNPs/CNTs

and (C) AgNPs/CNTs.

Fig. 3 XRD spectra of (a) Ni/CNTs, (b) AgNPs/CNTs and (c)

AgNPsNi/CNTs.

Fig. 4 Cyclic voltammetry (30th cycle) in 0.1 M NaOH at a 4 mm

diameter paraffin-impregnated graphite electrode modified with

(a) 20 mL CNTs/Nafion, (b) 20 mL NiNPs/CNTs/Nafion, (c) 20 mLAgNPs/CNTs/Nafion and (d) 80 mL AgNPs,NiNPs/CNTs/Nafion

casting solutions. Scan rate: 50 mV s�1.


electroactivity of bulk silver macroelectrodes in respect

to the electrooxidation of methanol. The cyclic voltam-

mograms obtained for the second cycle are shown in Fig. 5.

We choose to present the second cycle as a non-negligable

decrease in intensity (ca. 25%) is observed from the first cycle

to the second one. Measurements show a decrease of less

than 10% is then observed between the second and the

twentieth cycle. The electrochemical characterization of the

modified electrode surfaces, conducted independently as

described above, provides valuable data to compare the

electrocatalytic efficiency of the different materials. Indeed,

the catalytic currents obtained can be normalized with the

total metal surface area to provide the current density per unit

of electroactive surface (Table 1). The data reported in Table 1

show that the AgNPs/CNTs/Nafion- and AgNiNPs/CNTs/

Nafion-modified electrodes have similar properties, with a

current density about more than ten times higher than the

current density obtained at the bulk silver macroelectrode.

Such a higher electrocatalytic property can partially be

explained by the higher substrate diffusion that we expect to

observe at dispersed nanoparticles.26–31 Furthermore it has to

be said from the results reported in Table 1 that the experi-

mental results show that NiNPs and AgNPs have similar

electrocatalytic properties towards the electrooxidation of

methanol.

3.4 Characterization of magnetically driven movement of the

AgNPs,NiNPs/CNTs composites

The possibility of magnetically recovering the electrocatalytic

nanomaterial was explored by assessing the possibility to move

it with a magnet using the setup described in the Experimental

section. As it can be seen in Fig. 6 and Video S1 (ESIw), withthe shift of a magnet positioned directly below the glass

surface, the AgNPs,NiNPs/CNTs composites move from right

bottom to middle, then, to left up in the video screen,

suggesting an obvious movement for AgNPs,NiNPs/CNTs

composites on the surface of glass. The same test was

undertaken for the NiNPs/CNTs and AgNPs/CNTs nano-

composites. A similar response can be seen for NiNPs/CNTs

composites (Video S2, ESIw). However, no move was observed

for AgNPs/CNTs nanocomposites (Video S3, ESIw). It is thenpossible to conclude that the bifunctionalisation of the CNTs

provides the possibility to magnetically drive them to a specific

location in a solution. This added property to the new

electroactive nanomaterial allows, for example, the recovery

of the catalyst once the reaction has taken place.

4. Conclusions

Multi-walled carbon nanotubes functionalized either by nickel or

silver nanoparticles or by both were obtained using a single step

chemical deposition method in an ultrasonic bath. The electro-

activity of the bi-functionalized CNT multi-walled carbon nano-

tubes was assessed in respect to the electrooxidation of methanol.

It was found that they have significantly higher catalytic proper-

ties than the bulk silver of the same surface area. Furthermore, it

was shown that the addition of a minute fraction (1.5%) of

NiNPs in respect to the total number of nanoparticles adds to

their electrocatalytic properties the possibility to easily move

them in solution using a magnet. The bi-functionalized carbon

nanotubes are then easily recoverable after use.

Table 1 Catalytic performance for methanol oxidation of variouselectrode architectures; A is the metal total surface area estimated byelectrochemical oxidation of the surface, Ip is the peak current valuemeasured for a voltammetric scan obtained at 50 mV s�1 in 0.56 Mmethanol and 0.1 M NaOH and Jp is the corresponding currentdensity in respect to the total surface of silver

Electrode A/cm2 Ip/mA Jp/mA cm�2

Ag bulk 0.16 � 0.05 9.28 � 10�2 0.58 � 0.2NiNPs/CNTs/Nafion 0.29 � 0.1 3.00 10.3 � 2AgNPs/CNTs/Nafion 0.30 � 0.1 3.3 11.0 � 2AgNPs,NiNPs/CNTs/Nafion 0.56 � 0.2 4.14 7.4 � 2

Fig. 6 Optical microscopy images of the AgNPs,NiNPs/CNTs material, taken at 10 s time-intervals and following the movement of a magnet

going forward (images (A) to (C)) and then going backward (images (D) to (F)). These pictures were extracted from Video S1 (ESIw).

Fig. 5 Second cycle cyclic voltammetric curves obtained in 0.56 M

methanol and 0.1 M NaOH at a 0.08 mm diameter silver electrode

(curve a) and at a 4 mm diameter paraffin-impregnated graphite

electrode modified with casting solutions made of 80 mL CNTs/Nafion

(curve b), 80 mL NiNPs/CNTs/Nafion (curve c), 80 mL AgNPs/CNTs/

Nafion (curve d) and 120 mL AgNPs,NiNPs/CNTs/Nafion (curve e).

Scan rate: 50 mV s�1.


Acknowledgements

G.-P. J. gratefully acknowledges financial support from

Natural Science Foundation of Anhui Province of China

(No. 070415210), Science and Technology Program Founda-

tion of Hefei City (No. 20071032), and Doctor Foundation of

Hefei University of Technology (2005). R. B. and N. V. R. are

grateful to EPSRC for funding.

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Microwave-assisted facile synthesis of discotic liquid crystalline

symmetrical donor–acceptor–donor triads

Satyam Kumar Gupta, V. A. Raghunathan and Sandeep Kumar*

Received (in Durham, UK) 23rd May 2008, Accepted 11th July 2008

First published as an Advance Article on the web 8th September 2008

DOI: 10.1039/b808750d

We report the synthesis and characterization of two series of novel triphenylene–anthraquinone-

based symmetric discotic liquid crystalline trimers. These triads were prepared using microwave

dielectric heating. Conventional heating under similar reaction conditions failed to produce

desired products. To the best of our knowledge, these are the first donor–acceptor–donor triads

in which all the three components represent discotic mesogenic moieties. Chemical structures of

these discotic oligomers have been characterized by spectral techniques and elemental analysis.

The thermotropic liquid crystalline properties of these donor–acceptor–donor triads were

investigated by polarizing optical microscopy and differential scanning calorimetry. They exhibit

a columnar mesophase over a wide range of temperature. The columnar hexagonal mesophase

structure of these discotic oligomers has been elucidated with the help of X-ray diffraction studies.

Introduction

The notable improvement in the performance of electronic

devices based on organic semiconductors has attracted great

interest in recent years.1 The improved efficiency of organic

devices has origins ranging from appropriate molecular design

to well-defined structured layers essential for effective charge

transport. Recently there have been tremendous efforts to

achieve both p-type (hole conducting) and n-type (electron

conducting) properties in organic semiconducting materials

which are crucial for molecular electronics. One elegant

approach for such materials is to covalently link electron

donor and electron acceptor components at molecular level.

These kinds of materials are expected to behave as intrinsic,

non-composite p/n-type semiconductors. Such chemical tailor-

ing could lead to the development of other molecular archi-

tectures and it is envisaged that the combination of covalent

chemistry and self-assembly will be crucial for the develop-

ment of nano-engineered functional materials for electronic

applications.1 Among the diverse semiconductors, discotic

liquid crystals (DLCs) play an important role in the design

of electronic devices.2 Discotic liquid crystals are unique

nanostructures with remarkable electronic and optoelectronic

properties. Due to the co-facial stacking of aromatic cores,

disc-like molecules self organize into one dimensional colum-

nar wire and these columns in turn arrange themselves in

various two-dimensional lattices. The transport along the

columnar axis is much faster than between the columns. Due

to their relatively high charge carrier mobility, tendency to

form highly order films of various thickness and self healing of

defects owing to their dynamic nature, discotic mesogens have

been considered as attractive candidates for applications in

organic electronic devices such as photovoltaic solar cells, light

emitting diodes and field effect transistors.2

Microwave-assisted high-speed chemical synthesis has

attracted a considerable amount of attention in the past decade.

Almost all types of organic reactions have been performed

using the efficiency of microwave-flash heating. This is not

only due to the fact that reactions proceed significantly faster

and more selectively than under conventional thermal condi-

tions but also because of the operational simplicity, high yield

of products and cleaner reactions with easier work-up. A large

number of review articles provide extensive coverage of the

subject.3 Recently we and others have reported the synthesis

of a variety of liquid crystalline materials using microwave

dielectric heating.4

Very recently a great deal of attention is being paid to liquid

crystal oligomers.5 The physical properties of liquid crystalline

oligomers are significantly different from those of conven-

tional low molar mass liquid crystals. Their purification and

characterization are simple, and due to the restricted motion

of their components liquid crystal oligomers provide and

stabilize a variety of fluid phases with fascinating functions.

Further, an oligomeric approach provides a wide flexibility in

molecular design towards multifunctional liquid crystals.

However, compared to the number of calamitic oligomers,

discotic oligomers are rare. In this context we are interested in

the design and synthesis of novel functional discotic oligo-

meric materials and their mesophase behavior. Our molecular

design is such that it contains the well studied electron rich

triphenylene moiety6 and electron deficient anthraquinone7 as

the hole and electron transporting components, respectively.

These molecular double-cables, owing to their incommensu-

rate core sizes, may stack one on top of the other in the

columns to give columnar versions of double cable polymers,8

which could eventually provide side-by-side percolation

pathways for electrons and holes in solar cells. Here, we

report the synthesis and mesomorphism of novel triphenylene–

anthraquinone–triphenylene discotic liquid crystalline symmetric

Raman Research Institute, C.V. Raman Avenue, Sadashivanagar,Bangalore, 560 080, India. E-mail: [email protected];Fax: +91 80 23610492; Tel: +91 80 23610122



trimers. To the best of our knowledge, these are the first

donor–acceptor–donor triads in which all the three compo-

nents represent discotic mesogenic moiety.

Experimental

General information

Chemicals and solvents (AR quality) were used as received

without any further purification. Microwave irradiation was

performed in an unmodified household microwave oven.

(LG, MS-192W). However, commercial microwave reactors

for organic reactions are now available which provides

adequate mixing and control of reaction parameters such as

temperature and pressure. Column chromatographic separa-

tions were performed on silica gel (230–400 mesh). Thin layer

chromatography (TLC) was performed on aluminum sheets

precoated with silica gel (Merck, Kieselgel 60, F254). Chemi-

cal structure characterization of the compounds was carried

out through a combination of 1H NMR, 13C NMR (Bruker

AMX 400 spectrometer) and elemental analysis (Carlo-Erba

EA1112 analyzer). 1H NMR spectra were recorded using

deuterated chloroform (CDCl3) as solvent. Tetramethylsilane

(TMS) was used as an internal standard. The transition

temperatures and associated enthalpy values were determined

using a differential scanning calorimeter (DSC; Perkin-Elmer,

Model Pyris 1D) which was operated at a scanning rate of

5 1C min�1 both on heating and cooling cycles. The apparatus

was calibrated using indium (156.6 1C) as a standard. The

textural observations of the mesophase were carried out using

polarizing light microscopy (Olympus BX51) provided with a

heating stage (Mettler FP82HT) and a central processor

(Mettler FP90). X-Ray diffraction studies (XRD) were carried

out on unoriented samples using Cu-Ka (l = 1.54 A) radia-

tion from a Rigaku Ultrax 18 rotating anode generator

(5.4 kW) monochromated with a graphite crystal. The samples

were held in sealed Lindemann capillary tubes (0.7 mm

diameter) and the diffraction patterns were collected on a

two-dimensional Marresearch image plate.

Synthesis of trimers

Rufigallol 2, 1,5-dihydroxy-2,3,6,7-tetraalkoxy-9,10-anthra-

quinone 3, hexaalkoxytriphenylene 4, monohydroxypenta-

alkoxytriphenylene 5 and o-bromo-substituted triphenylene

6 were prepared as reported by us previously.9 All the trimers

were prepared following same method which involves alkyla-

tion of 1,5-dihydroxy-2,3,6,7-tetraalkoxy-9,10-anthraquinone

3 with terminal bromo-substituted triphenylene 6 using micro-

wave dielectric heating. A typical procedure for the synthesis

of a representative example 7a10 is given below. The suffix

number in the series 7a and 7b, represents the number of

carbon atoms in the peripheral chains attached with central

anthraquinone moiety (R in the structure 7, Scheme 1).

A mixture of compound 6a (n = 9) (300 mg, 0.30 mmol),

3 (R= C10H21) (43 mg, 0.05 mmol) and Cs2CO3 (200 mg, 0.61

mmol) in NMP (0.5 mL) was irradiated in a microwave oven

for 30 s. The vial was removed from the oven and left to stand

for about 1 min and again irradiated for 30 s. This process was

repeated for 20 times until the reaction was complete (TLC

monitoring). The cooled reaction mixture was then poured

into an excess of distilled water and extracted with chloroform.

The organic extract was dried over anhydrous sodium sulfate,

concentrated and the product was purified by repeated column

chromatography over silica gel (eluent: 4% ethyl acetate in

hexane). Solvent was then removed in rotary evaporator. The

residue was then dissolved in dichloromethane and the result-

ing solution was added to cold methanol to afford 7a10

(34 mg, 25%). 1H NMR (400 MHz, CDCl3): d 7.83 (s, 12 H),

7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.4 Hz,

4 H), 4.06 (t, J=6.2Hz, 8 H), 1.94 (m, 32H), 1.77 (q, J=7.5Hz,

4 H), 0.8–1.6 (m, 190 H). 13C NMR (100 MHz, CDCl3):

d 181.2, 157.5, 153.9, 149.1, 147, 132.7, 123.7, 107.6, 107.1,

77.3, 77, 76.7, 75.9, 74.7, 74.1, 69.8, 69.2, 31.9, 31.7, 30.4, 29.5,

29.4, 26.1, 25.9, 22.7, 21.3, 18.5, 15.9, 14.0. Elemental analysis:

Calc. for C174H276O20, C 77.75, H 10.35. Found: C 77.32,

H 10.53%. All other compounds give satisfactory spectral and

elemental analysis data in accordance with their chemical

structure. Selected data for compound 7a6: 1H NMR: d 7.83

(s, 12 H), 7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14 (t, J =

6.4 Hz, 4 H), 4.06 (t, J = 6.5 Hz, 8 H), 1.94 (m, 32 H), 1.77

(q, J = 7.5 Hz, 4 H), 0.8–1.6 (m, 158 H). Elemental analysis:


H 9.98%. 7a7: 1H NMR: d 7.83 (s, 12 H), 7.59 (s, 2 H), 4.23

(t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.4 Hz, 4 H), 4.06 (t, J =

6.5 Hz, 8 H), 1.94 (m, 32 H), 1.78 (q, J= 7.8 Hz, 4 H), 0.8–1.6

(m, 166 H). Elemental analysis: Calc. for C162H252O20, C

77.22, H 10.08; Found C 76.91, H 10.04%. 7a8: 1H NMR:

d 7.83 (s, 12 H), 7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14

(t, J = 6.5 Hz, 4 H), 4.06 (t, J = 6.2 Hz, 8 H), 1.94 (m, 32 H),

1.78 (q, J = 6.9 Hz, 4 H), 0.8–1.6 (m, 174 H). Elemental

analysis: Calc. for C166H260O20, C 77.40, H 10.17. Found: C

77.13, H 9.82%. 7a100: 1H NMR: d 7.83 (s, 12 H), 7.61

(s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.5 Hz,

4 H), 4.06 (t, J = 6.2 Hz, 8 H), 1.94 (m, 32 H), 1.78 (q, J =

6.9 Hz, 4 H), 0.8–1.6 (m, 190 H). Elemental analysis: Calc. for

C174H276O20, C 77.75, H 10.35. Found: C 77.32, H 10.89%.

7a14: 1H NMR: d 7.83 (s, 12 H), 7.59 (s, 2 H), 4.23 (t, J=6.5 Hz,

24 H), 4.14 (t, J= 6.3 Hz, 4 H), 4.06 (t, J= 5.6 Hz, 8 H), 1.94

(m, 32 H), 1.77 (q, J = 7.8 Hz, 4 H), 0.8–1.6 (m, 222 H).

Elemental analysis: Calc. for C190H308O20, C 78.35, H 10.66.

Found: C 77.96, H 10.71%. 7b6: 1H NMR: d 7.83 (s, 12 H),

7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.15 (t, J = 6.4 Hz,

4 H), 4.06 (t, J = 6.4 Hz, 8 H), 1.94 (m, 32 H), 1.78 (q, J =


C154H236O20, C 76.83, H 9.88. Found: C 76.37, H 9.89%. 7b7:1H NMR: d 7.83 (s, 12 H), 7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz,

24 H), 4.14 (t, J = 6.8 Hz, 4 H), 4.06 (t, J = 6.3 Hz, 8 H),

1.94 (m, 32 H), 1.78 (q, J = 7.4 Hz, 4 H), 0.8–1.6 (m, 158 H).

Elemental analysis: Calc. for C158H244O20, C 77.03, H 9.98.

Found: C 76.62, H 10.36%. 7b10: 1H NMR: d 7.83 (s, 12 H),

7.59 (s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.3 Hz,

4 H), 4.06 (t, J = 6.5 Hz, 8 H), 1.94 (m, 32 H), 1.77 (q, J =


C170H268O20, C 77.58, H 10.26. Found: C 77.31, H 10.23%.

7b100: 1H NMR: d 7.84 (s, 12H, Ar–H), 7.59 (s, 2 H), 4.23

(t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.3 Hz, 4 H), 4.06 (t, J =

6.5 Hz, 8 H), 1.94 (m, 32 H), 1.77 (q, J= 7.8 Hz, 4 H), 0.8–1.6

(m, 182 H). Elemental analysis: Calc. for C170H268O20, C 77.58,


H 10.26. Found: C 77.14, H 10.00%. 7b12: 1H NMR: d 7.84

(s, 12 H), 7.59 (s, 2 H), 4.23 (t, J = 6.3 Hz, 24 H), 4.14 (t, J =

6.3 Hz, 4 H), 4.06 (t, J = 6.2 Hz, 8 H), 1.94 (m, 32 H), 1.77

(q, J = 7.8 Hz, 4 H), 0.8–1.6 (m, 190 H). Elemental analysis:


H 10.30%. 7b14: 1H NMR: d 7.83 (s, 12 H), 7.59

(s, 2 H), 4.23 (t, J = 6.5 Hz, 24 H), 4.14 (t, J = 6.4 Hz,

4 H), 4.06 (t, J = 6.3 Hz, 8 H), 1.94 (m, 32 H), 1.78 (q, J =


C186H300O20, C 78.21, H 10.59. Found: C 78.16, H 10.53%.


Synthesis

The synthesis of the novel symmetrical trimers was achieved as

shown in Scheme 1. The unequal reactivity of the six phenolic

groups of rufigallol 2, two of which are less reactive by virtue

of being intramolecularly hydrogen bonded to the adjacent

quinone carbonyls, was exploited. Etherification of rufigallol 2

under mild conditions produced 1,5-dihydroxy-2,3,6,7-tetra-

alkoxy-9,10-anthraquinone 3 without alkylating the hydrogen

bonded C-1 and C-5 positions. These tetraalkoxy derivatives

were further alkylated by o-bromo-substituted triphenylenes

with the help of microwave dielectric heating as shown in

the Scheme 1, under mild basic conditions to furnish the

symmetrical trimers within 10 min, which is simple, efficient,

rapid and economic. All attempts to etherify the intramolecu-

larly hydrogen bonded C-1 and C-5 positions with bulky

o-bromo-substituted triphenylene failed under classical

thermal heating conditions even by using strong basic condi-

tions and prolonged reaction times (24 h). For instance,

heating the same reaction mixture in DMF at 100 1C for

48 h or heating a mixture of 3 and 6 in DMF and NaOH or

K2CO3 for 48 h did not furnish any product.

Thermal behavior

The thermal behavior of all the compounds was investigated

by polarizing optical microscopy (POM) and differential scan-

ning calorimetry (DSC). In the case of materials which were

mesomorphic, classical textures of discotic columnar meso-

phases appeared upon cooling from the isotropic liquid as

shown in Fig. 1. These textures are similar to the known

textures for Colh phases. All the trimers contain two identical

triphenylenes substituted with five hexyloxy peripheral chains

linked to the central anthraquinone moiety through a

12- (7a series) or a 10- (7b series) methylene spacer. In both

the series the peripheral alkyl chain lengths around the

anthraquinone core varies from hexyloxy to tetradecyloxy.

The transition temperature and associated enthalpy data

obtained from the heating and cooling cycles of DSC are

collected in Table 1. The peak temperatures are given in 1C

and the numbers in parentheses indicate the transition en-

thalpy (DH) in J g�1. The compound 7a0, without any

peripheral alkyl chains (OR = H) around the central core of

the trimer, does not exhibit any liquid crystalline property. It

melts from crystalline solid state to isotropic liquid state at

Scheme 1 Synthetic route of triphenylene–anthraquinone trimers. 7a Series: OR = H (7a0); R = n-C6H13 (7a6); R = n-C7H15 (7a7); R =

n-C8H17 (7a8); R = n-C10H21 (7a10); R = 3,7-dimethyloctyl (7a100); R = n-C14H29 (7a14); 7b Series: R = n-C6H13 (7b6); R = n-C7H15 (7b7);

R = n-C10H21 (7b10); R = 3,7-dimethyloctyl (7a100); R = n-C12H25 (7a12); R = n-C14H29 (7a14).


39.7 1C on heating and on cooling it crystallizes slowly over a

period of time at room temperature. This could be because the

absence of alkyl chains around the core does not provide the

space filling effect of alkyl chains which is crucial for exhibiting

mesophase behavior in discotic liquid crystals. The highest

homologue of the series 7a14 also does not display any liquid

crystalline property, it passes from crystalline solid state to

isotropic liquid state at 47 1C on heating and on cooling the

isotropic liquid crystallizes at 18.4 1C. This could be because

the longer alkyl chains around the central anthraquinone core

may hinder the self-assembly of molecules. All other members

of the 7a series 7a6, 7a7, 7a8, 7a10 and 7a100 display enantio-

tropic mesophase behavior. In their DSC thermograms, they

display a soft solid to mesophase transition followed by

mesophase to isotropic transition on heating. Upon cooling

they show only isotropic to mesophase transition and the

mesophase remains stable down to room temperature or

partially solidified at low temperature. As a typical example

the DSC thermogram of compound 7a6 is shown in Fig. 2. On

increasing the alkyl chain length around the anthraquinone

core the mesophase to isotropic transition temperatures of the

trimers decrease as shown in the Fig. 3. This could be because

the longer alkyl chains introduce more intracolumnar disorder

and hence core–core unstacking becomes easier.

In series 7b only two trimers 7b6 and 7b7 were found to

display enantiotropic liquid crystalline properties. Compound

7b100 shows monotropic phase behavior. Other trimers 7b10,

7b12 and 7b14 of the series do not exhibit any liquid crystalline

property. They show only crystalline to isotropic and isotropic

to crystalline transitions on heating and cooling, respectively.

This is not surprising as the spacer connecting the donor with

Fig. 1 Optical micrograph of 7a6 at 80 1C on cooling from the

isotropic liquid (crossed polarizer, magnification � 200)

Table 1 Phase transition temperatures (peak, 1C) and associated enthalpy changes (J g�1 in parentheses) of novel symmetrical trimers

Compounda First heating scan First cooling scan

7a6 ss 59.1 (1.6) Colh 104.1(6.0) I I 99.3 (6.4) Colh7a7 ss 37 (1.9) g0 67.4 (1.5) Colh 89.6 (2.4) I I 81.4 (2.4) Colh 6.4 (0.9) ss7a8 ss 47.3 (8.6) Colh 83.0 (2.5) I I 72.1 (3.1) Colh7a10 ss 51.2 (8.6) Colh 59.1 (0.9) I I 53 (2.6) Colh 32.6 (0.9) ss7a100 ss 45.6 (10.6) Colh 69 (2.3) I I 57 (2.6) Colh7a14 Cr 47 (34.2) I I 18.4 (23.0) Cr7b6 ss 41.5 (9.6) Colh 69.3 (1.2) I I 59.7 (1.4) Colh7b7 ss 45 (20.4) Colh 65.8 (2.6) I I 58.6 (2.8) Colh7b10 Cr 47(2.4) Cr0 63 (31.1) I I 31.4 (1.0) Cr0 21.4(0.4) Cr7b100 Cr 70.7 (36.5) I I 45.4 (6.9) Colh7b12 Cr 44.4 (18.1) I I 8.6 (11.2) Cr7b14 Cr 60.1 (33.1) I I 29.7 (28.4) Cr

a See Scheme 1 for chemical structures. ss: semisolid; Cr: crystal; Colh: hexagonal columnar phase; I: isotropic phase.

Fig. 2 DSC thermogram of the trimer 7a6 on heating and cooling

cycles (scan rate 10 1C min�1).

Fig. 3 Variation of phase transition temperatures of 7a6–7a10

with number of carbon atoms in the peripheral alkyl chains of

anthraquinone.


the acceptor is short, and so long peripheral substitution

around the central core can disturb their packing. The absence

of a mesophase in compounds 7a14, 7b10, 7b12 and 7b14

clearly indicates that, when the peripheral chain lengths of the

central core are either equal to or longer than the spacer

length, then these symmetrical trimers do not exhibit any

liquid crystalline property. Shorter alkyl chains around the

central core stabilize the mesophase in these discotic trimers. If

we compare the mesophase stability between compound 7a10

and 7a100 having the same mass unit around the central core,

the compound 7a100 shows a wider mesophase range of

23.4 1C compared to 7.9 1C of compound 7a10. Similarly on

comparing 7b10 and 7b100 we find that compound 7b10 does

not exhibit any liquid crystalline property but compound

7b100 displays a monotropic phase behavior. The above

trimers contain the same mass units around the central core

but the alkyl chain lengths around the central core are

different. Both 7a100 and 7b100 contain shorter chain lengths

i.e. 3,7-dimethyloctyl as compared to 7a10 and 7b10 with

longer decyl chains.

X-Ray diffraction studies

In order to reveal the mesophase structure and hence the

supramolecular organization of these compounds, X-ray dif-

fraction experiments were carried out using unoriented sam-

ples. X-Ray diffraction patterns for all the trimers were

recorded in the columnar phase 10 1C below the clearing

temperature while cooling from the isotropic phase. The

X-ray diffraction patterns of the mesophase exhibited by all

the samples belonging to both the series is supportive of a

discotic hexagonal columnar arrangement. As a typical exam-

ple, the X-ray diffraction pattern of compound 7a6 and its

one-dimensional intensity vs. theta (y) graph derived from the

pattern are shown in the Fig. 4. Qualitatively all the com-

pounds show similar X-ray diffraction patterns. As can be seen

from the figure, in the small angle region, two sharp peaks, one

very strong and one weak reflection are seen whose d-spacings

are in the ratio of 1 : 1/O3, consistent with a two-dimensional

hexagonal lattice. In the wide-angle region two diffuse reflec-

tions are seen. The broad one centered at 4.62 A corresponds

to the liquid-like order of the aliphatic chains. The reflection at

higher y value and well separated from the previous one is due

to the stacking of the molecular cores one on the top of the

other. The diffuse nature of this peak implies that the stacking

of the discs within each column is correlated over short

distances only. The average stacking distance (core–core

separation) was found to be 3.66 A and falls in the range

observed for a number of materials exhibiting a discotic

columnar phase. The discotic molecules stack one on top of

the other to form the columns and these columns in turn

arrange themselves on a two-dimensional hexagonal lattice for

both the series of compounds. The intercolumnar distances, a,

calculated using the relation a = d10/cos301, where d10 is the

spacing corresponding to the strong peak in the small angle

region, for all the compounds, are listed in Table 2. In both

the series it is evident that as alkyl chain lengths increase

the diameter of the cylindrical columns formed by the

discotic molecules also increases, as shown in Fig. 5. The

Fig. 4 X-Ray diffraction pattern of the trimer 7a6 at 85 1C and its

intensity vs. y profile.

Table 2 Values of d-spacings, and of inter- (dinter) and intracolumnar(dintra) distances (A) of the trimers derived from their diffractionpatterns

Compound d-Spacing/A dinter/A dintra/A

7a6 17.71 20.45 3.587a7 17.80 20.56 3.657a8 18.15 20.96 3.667a10 18.59 21.47 3.667a100 18.13 20.93 3.737b6 17.33 20.01 3.667b7 17.43 20.13 3.65H6TP 19.5 22.52 3.56H6AQ 18.19 21.0 3.6

Fig. 5 Variation of d-spacing value with respect to side chain length.


intercolumnar distances varies from 20.5–21.5 A, whereas the

intracolumnar distance is constant at around 3.7 A, which is

usually observed for discotic columnar mesophases. In these

unoriented samples, we do not observe any additional small

angle peak for the formation of a superlattice arising from the

ideal top-on-top stacking of the trimer molecules which could

lead to the formation of columnar double cables. Therefore, it

was concluded that the triphenylene and anthraquinone sub-

units arrange themselves statistically to form a columnar

hexagonal phase. The intercolumnar distance for hexahexyl-

oxytriphenylene (H6TP) and hexahexyloxyanthraquinone

(H6AQ) is 22.52 A,10 and 21.0 A,11 respectively, but the

intercolumnar distance of the symmetrical trimer 7a6 is

20.45 A, which is less than the corresponding monomers. This

minor shrinkage of the intercolumnar distance in the trimer is

expected upon covalent linking the two molecules. On com-

paring the X-ray diffraction results of 7a6 with 7b6 and 7a7

with 7b7 (Table 2), it is evident that the intercolumnar distance

is decreasing with decreasing spacer length. This is due to

shortening of hexagonal lattice with decreasing the length of

spacer linking the discotic moieties. As expected, the inter-

columnar distance of 7a10 is larger than that of 7a100, since

7a10 contains longer alkyl chains around the central anthra-

quinone core than 7a100, although they contain the same mass

units around the anthraquinone core. However, the intra-

columnar distance of 7a10 is less than 7a100 because of the

steric effect exerted by the branched alkyl chains around the

central core of 7a100, which will hinder the discotic cores

coming closer in columns.

Absorption spectra

As the trimers contain both electron donor and acceptor

moieties, it is expected that they may show charge transfer

absorption. However, the UV-vis spectrum of the trimer 7a10

(Fig. 6) does not show any additional absorption band as

compared to the separate hexaalkoxytriphenylene 4 and hexa-

alkoxyanthraquinone (RF6C4) and is essentially a sum of

donor and acceptor units. The colour of the trimer 7a10 also

almost matches with the colour of the acceptor. This implies

that there is no or very weak charge transfer interaction

between donor and acceptor units. Similar behaviour has

previously been reported for other non-liquid crystalline as

well as liquid crystalline donor–acceptor dimers.9a,12

Conclusions

In conclusion, we have synthesized two series of novel symme-

trical liquid crystalline trimers based on anthraquinone and

triphenylene moieties using microwave irradiation. The etheri-

fication of H-bonded hydroxyl groups of tetraalkoxyanthra-

quinones with bulky o-bromo-substituted triphenylenes failed

to produce the desired triads under classical reaction condi-

tions. The mesophase behavior of the symmetrical trimers was

studied by polarizing optical microscopy and differential

scanning calorimetry and they exhibit a columnar mesophase

over a wide range of temperature. Hexagonal columnar struc-

ture of the mesophase of these donor–acceptor–donor triads

was established by X-ray diffraction studies. Longer spacer

length, smaller peripheral alkyl chain length and branching in

peripheral alkyl chains of the anthraquinone favor liquid

crystalline property in these symmetrical trimers.

References

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Fig. 6 UV-vis spectra of a chloroform solution of 7a10 and of its

individual monomers (hexaalkoxyanthraquinone (RF6C4) and hexa-

alkoxytriphenylene).


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Synthesis, crystal structures and luminescence properties of lanthanide

oxalatophosphonates with a three-dimensional framework structurew

Yanyu Zhu, Zhengang Sun,* Yan Zhao, Jing Zhang, Xin Lu, Na Zhang,

Lei Liu and Fei Tong

Received (in Montpellier, France) 20th August 2008, Accepted 3rd October 2008


DOI: 10.1039/b814400a

Six new three-dimensional (3D) lanthanide oxalatophosphonates, [Ln(HL)(C2O4)0.5(H2O)2]�H2O

(Ln = La (1), Ce (2), Pr (3), Nd (4), Sm (5), Eu (6); H3L = H2O3PCH(OH)CO2H), have been

synthesized under hydrothermal conditions and structurally characterized by single-crystal X-ray

diffraction as well as by infrared spectroscopy, elemental analysis and thermogravimetric analysis.

Compounds 1–6 are isomorphous and they exhibit a complex three-dimensional (3D) open-

framework structure with a one-dimensional channel system along the c-axis. The interconnection

of the lanthanide(III) ions by phosphonate ligands results in a lanthanide phosphonate layer, and

these layers are further bridged by oxalate anions to form a 3D open-framework. Compound 6

shows strong red luminescence in the solid state at room temperature.

Introduction

Metal phosphonates as a class of inorganic–organic hybrid

materials have attracted a great deal of research interest as a

result of their ability to form interesting structures with

potential applications as catalysts, ion exchangers, sorbents,

meso-/microporous materials, or intercalation chemistry.1–5

Usually, the metal phosphonates adopt layered or pillared

layered structures, with the organic groups filling in between

the inorganic layers.6–10 Other structural types have also been

observed in some phosphonates, among which the open-

framework and porous structures are of particular inter-

est.11–16 The strategy of attaching functional groups such as

amine, hydroxyl and carboxylate groups to the phosphonic

acid has proven to be effective for the isolation of a variety

of metal phosphonates with open-framework and micro-

porosity.17–21 Based on 2-hydroxyphosphonoacetic acid, proline-

N-methylphosphonic acid and DL-(a-aminoethyl)phosphonic

acid, a series of metal phosphonates with two-dimensional

(2D) layer and three-dimensional (3D) open-framework struc-

tures have also been isolated in our laboratory.22

Recently, many research activities have focused on the

synthesis of inorganic–organic hybrid compounds by incor-

porating organic ligands in the structures of metal phospho-

nates.23–26 The direct use of two types of ligands in the

preparation, such as a phosphonic acid and a carboxylic acid,

has been found to be another effective method for the

exploration of hybrid open-frameworks. Among these studies,

the oxalate moiety, C2O42�, was found to be a good candidate

and has been successfully incorporated into phosphonate

frameworks with transition metals and main group ele-

ments.27–30 Although some progress has been made in the

construction of metal oxalatophosphonates as mentioned

above, less progress has been achieved in the synthesis of

lanthanide oxalatophosphonates.31–33 Lanthanide phospho-

nates normally have low solubility in water and other organic

solvents, hence introducing a second ligand such as C2O42�

into the lanthanide phosphonate system can improve the

solubility and crystallinity of the lanthanide phosphonate. In

addition, the coordination of two types of ligands with the

lanthanide ion may reduce or eliminate water molecules from

the coordination sphere of the lanthanide(III) ion, hence

increasing the luminescent intensity and lifetime of the

materials.34 In this paper, we selected 2-hydroxyphosphono-

acetic acid (H3L) as the phosphonate ligands and oxalate

as the second metal linker. Hydrothermal reactions of the

above two ligands with lanthanide(III) chlorides afforded six

new lanthanide oxalatophosphonate hybrids with 3D open-

framework structures, namely, [Ln(HL)(C2O4)0.5(H2O)2]�H2O

(Ln = La (1), Ce (2), Pr (3), Nd (4), Sm (5), Eu (6); H3L =

H2O3PCH(OH)CO2H). The luminescent property of com-

pound 6 has also been studied.


Synthesis

By using 2-hydroxyphosphonoacetic acid as the phosphonate

ligand and oxalate as the second metal linker, six new lantha-

nide(III) oxalatophosphonates have been synthesized under

hydrothermal conditions. The compounds 1–6 were obtained

as a pure phase materials by adjusting the synthetic conditions.

Institute of Chemistry for Functionalized Materials, Faculty ofChemistry and Chemical Engineering, Liaoning Normal University,Dalian, 116029, P. R. China. E-mail: [email protected];Fax: +86 411 82156858w Electronic supplementary information (ESI) available: Fig. S1:Simulated XRD pattern of compound 1 and experimental powderXRD patterns of compounds 1–6. Fig. S2: Experimental powder XRDpattern of compound 1 and of dehydrated samples after calcination at150 and 180 1C. Table S1: Selected bond lengths (A) for compounds1–6. Table S2: Selected bond angles (1) for compounds 1–6. CCDCreference numbers 670477–670481 (1–5) and 686602 (6). For ESIand crystallographic data in CIF or other electronic format seeDOI: 10.1039/b814400a



The molar ratio of the starting materials and the pH of the

reaction mixture play an important role in the formation of

these six compounds. NaOH was employed as the inorganic

base to adjust the pH of the reaction mixture. It was found

that pure phases of compounds 1–6 can be obtained with good

yields when the molar ratio of LnCl3�6H2O, H3L, H2C2O4�2H2O, NaOH, H2O and the pH are 1 : 4 : 4 : 4 : 888 and

1.5–2.5. In addition, the reaction temperature was very

important for the formation of suitable single crystals for

X-ray diffraction. The compounds 1–6 were obtained at

120–140 1C under hydrothermal conditions. The powder

XRD patterns of compounds 1–6 and the simulated XRD

patterns of compound 1 are shown in the supplementary

material (Fig. S1, ESIw). The diffraction peaks on the patterns

correspond well in position, confirming these six compounds

are isomorphous, and showing their phase purity. The differ-

ences in reflection intensities are probably due to preferred

orientation in the powder samples.

Description of the crystal structures

Compounds 1–6 are isomorphous and feature three-dimen-

sional open frameworks, hence only the structure of 3 will be

discussed in detail as a representation. The ORTEP diagram

for compound 3 is shown in Fig. 1. Crystallographic data and

structural refinements for compounds 1–6 are summarized in

Table 1.

As shown in Fig. 1, the Pr(III) ion is nine-coordinated by two

phosphonate oxygen atoms, two carboxylate oxygen atoms,

and one hydroxyl oxygen atoms from three HL2� anions, two

oxygen atoms from one oxalate anion as well as two aqua

ligands. The Pr–O distances range from 2.372(3) to 2.573(3) A,

which are comparable to those reported for other praseo-

dymium(III) phosphonates.23,33 The asymmetric unit contains

half of an oxalate ion which lies about an inversion centre.

The oxalate anion is bidentate, and it chelates with

two different Pr(III) ions by using its four carboxylate

oxygen atoms. Each oxalate anion forms two Pr–O–C–C–O

five-membered chelating rings. The pentadentate HL2� ligand

is bidentate with Pr1 and Pr1D and monodentate with Pr1B.

Each HL2� anion chelates with Pr1D ion by using its one

carboxylate oxygen atom (O6) and one hydroxyl oxygen

atom (O4), and one carboxylate oxygen atom (O5) and one

phosphonate oxygen atom (O2) chelate with Pr1 ion.

One phosphonate oxygen atom (O3) is unidentate, whereas

the remaining one (O1) is protonated and noncoordinated.

Such configuration is favorable because of the formation of

stable six-atom rings (P–O–Pr–O–C–C) and five-atom rings

(O–Pr–O–C–C). It is noted that the Ln–O (hydroxyl oxygen)

distances are longer than the other Ln–O distances in com-

pounds 1–6, attributed to the presence of the hydroxyl proton

(Table S1, ESIw).The HL2� anion acts as a bridging ligand to link Pr(III) ions

into a 1D chain of {Pr(HL)}+ along the b axis (Fig. 2(a)).

The dihedral angle between two chelating rings sharing a

common Pr(III) ion is 74.61(10)1. These chains are cross-linked

by bridging HL2� anions to form a praseodymium phosphonate

layer in the bc plane (Fig. 2(b)), the layers are interconnected

by sharing Pr(III) ions into a pillared-layered architecture with

the oxalate groups acting as pillars (Fig. 3). The result of

connections in this manner is the formation of a 1D channel

system along the c axis (Fig. 4(a)). The channel is formed by

21-atom rings composed of five Pr(III) ions, three HL2� anions

and two oxalate anions (Fig. 4(b)). The dimensions of the

channels are estimated to be 11.2 A (Pr1a–P1e) � 10.2 A

(Pr1–C3c) based on the structural data. The oxygen atoms

from coordinated water molecules are oriented toward the

channel center, and lattice water molecules are located inside

the channels. The structure of 3 can be viewed as the praseo-

dymium phosphonate layer being connected via oxalate anions

to form a complex 3D open-framework structure.

Fig. 1 ORTEP representation of a selected unit of compound 3. The

thermal ellipsoids are drawn at the 30% probability level. All H atoms

and lattice water molecules are omitted for clarity: Pr(1)–O(3)B,

2.372(3) A; Pr(1)–O(2), 2.481(3) A; Pr(1)–O(7), 2.493(3) A;

Pr(1)–O(6)A, 2.522(3) A; Pr(1)–O(5), 2.559(3) A; Pr(1)–O(9), 2.537(4)

A; Pr(1)–O(8)C, 2.571(3) A; Pr(1)–O(10), 2.542(4) A; Pr(1)–O(4)A,

2.573(3) A. Symmetry codes: A: �x + 2, y + 1/2, �z + 3/2; B: �x +

2, �y,�z+2; C:�x+ 1,�y, �z+ 1; D: �x+ 2, y� 1/2,�z+3/2.

Fig. 2 (a) 1D chain {Pr(HL)}+ and (b) A 2D praseodymium(III)

phosphonate layer in compound 3.


IR spectra

The IR spectra of the six compounds have many similar

features corresponding to the common groups, thus only the

spectrum of compound 3 will be discussed (Fig. 5). The

IR spectrum for compound 3 was recorded in the region

4000–400 cm�1. The broad band in the range 3550–3000 cm�1

corresponds to the O–H stretching vibrations of water mole-

cules, hydroxyl groups and phosphonate groups. There are

two strong bands centered at 1650 and 1575 cm�1, which are

assigned to the asymmetrical and symmetrical stretching

vibrations of C–O groups when present as COO� moieties.35

Strong bands between 1200 and 900 cm�1 are due to stretching

vibrations of the tetrahedral CPO3 groups, as expected.36

Additional intense and sharp bands at low energy (617, 572

and 430 cm�1 etc) are found. These bands are probably due to

bending vibrations of the tetrahedral CPO3 groups.

Thermal properties

Except for the final weight loss temperature and total weight

losses, the TGA curves of compounds 1–6 are very similar,

with three main continuous weight losses. Herein, we use

compound 1 as an example to illuminate the weight losses in

detail. As shown in Fig. 6, the first step corresponds to the loss

Table 1 Crystal data and structure refinements for compounds 1–6

1 2 3 4 5 6

Empirical formula C3H9O11PLa C3H9O11PCe C3H9O11PPr C3H9O11PNd C3H9O11PSm C3H9O11PEuM 390.98 392.19 392.98 396.31 402.42 404.03Crystal system Monoclinic Monoclinic Monoclinic Monoclinic Monoclinic MonoclinicSpace group P21/c P21/c P21/c P21/c P21/c P21/ca/A 7.1991(6) 7.2260(7) 7.2300(7) 7.2326(6) 7.2351(6) 7.2293(9)b/A 13.3838(11) 13.2918(13) 13.2239(13) 13.1558(11) 13.0550(10) 13.0429(16)c/A 10.2926(8) 10.2745(10) 10.2535(10) 10.2337(8) 10.1999(8) 10.1980(13)b/1 98.8980(10) 99.4930(10) 99.8690(10) 100.1070(10) 100.4270(10) 100.541(2)V/A3 979.77(14) 973.32(16) 965.82(16) 958.63(14) 947.51(13) 945.4(2)Z 4 4 4 4 4 4Dc/g cm�3 2.651 2.676 2.703 2.746 2.821 2.839m/mm�1 4.576 4.894 5.263 5.636 6.420 6.858GOF on F2 1.017 1.063 1.074 1.039 1.090 1.099R1 [I 4 2s(I)]a 0.0210 0.0239 0.0278 0.0220 0.0206 0.0243wR2 [I 4 2s(I)]a 0.0547 0.0526 0.0726 0.0543 0.0512 0.0589R1 (All data)a 0.0233 0.0296 0.0318 0.0242 0.0226 0.0255wR2 (All data)a 0.0560 0.0555 0.0753 0.0557 0.0523 0.0595

a R1 =P

(||Fo| � |Fc|)/P

|Fo|; wR2 = [P

w(|Fo| � |Fc|)2/P

wFo2]1/2.

Fig. 3 A ball-and-stick and polyhedral view of compound 3 along the

b axis.

Fig. 4 (a) View of the framework for compound 3 along the c-axis showing the voids in the structure. (b) A 21-atom rings in compound 3.

All H atoms and lattice water molecules are omitted for clarity. Symmetry codes: a: �x+ 2, y � 1/2, �z+ 3/2; b: x, �y � 1/2, z � 1/2; c: �x+ 1,

y � 1/2, �z + 1/2; d: �x + 1, �y, �z + 1; e: x � 1, y, z � 1.


of one lattice water molecule and two aqua ligands. The

weight loss started at 50 1C and was completed at 145 1C.

The observed weight loss of 14.0% is close to the calculated

value (13.8%). The second step between 345 and 430 1C can be

attributed to decomposition of oxalate and phosphonate units.

The third step corresponds to the further decomposition of the

phosphonate group. The observed total weight loss at 735 1C

is about 41.5%, and the final products are not identified.

However, we suspect they are mainly LaPO4. The total weight

loss of 41.5% is close to the calculated value (40.2%) if the

final product is assumed to be LaPO4. The observed total

weight losses of compounds 2–6 are 41.4, 39.0, 37.0, 38.6,

40.3%, respectively. Considering the thermal stability of the

compounds, X-ray powder diffraction studies were performed

for the as-synthesized compound 1 and the samples calcined at

150 and 180 1C. The XRD patterns for the calcined samples fit

well with that of the as-synthesized samples, indicating that the

structure of these six compounds can be kept after dehydra-

tion process (Fig. S2, ESIw).

Photoluminescent properties

It is well-known that the lanthanides, especially europium and

terbium, can absorb ultraviolet radiation efficiently through an

allowed electronic transition to convert to the excited state5D4, and these excited states are deactivated to the multiplet7FJ states radiatively via emission of visible radiation as

luminescence. The solid-state luminescence property of com-

pound 6 was investigated at room temperature. Compound 6

emits red light upon excitation at 396 nm, and its luminescence

spectrum is depicted in Fig. 7. These emission bands arise from5D0 - 7FJ (J = 1, 2 and 4) transitions, typical of Eu(III)

ions.34,37 The 5D0 -7F1 transition (593 nm) corresponds to a

magnetic dipole transition, and the intensity of this emission

for 6 is medium-strong. The most intense emission in

the luminescent spectrum is the 5D0 - 7F2 transitions at

617 nm, which are the so-called hypersensitive transitions

and are responsible for the brilliant-red emission of compound

6.38 The emission spectrum of 6 shows a weak emission band

at 695 nm, which can be attributed to the 5D0 - 7F4

transition. The results indicate that compound 6 is a good

candidate as a red-light luminescent material.

Conclusions

By using 2-hydroxyphosphonoacetic acid as the phosphonate

ligand and oxalate as the second metal linker, six new

lanthanide(III) oxalatophosphonates with a general formula

[Ln(HL)(C2O4)0.5(H2O)2]�H2O (Ln = La (1), Ce (2), Pr (3),

Nd (4), Sm (5), Eu (6); H3L = H2O3PCH(OH)CO2H) have

been synthesized and structurally characterized. Compounds

1–6 are isomorphous and the structure of these compounds

features a 3D open-framework with a one-dimensional

channel system along the c-axis. The interconnection of the

lanthanide(III) ions by phosphonate ligands results in a lantha-

nide phosphonate layer, and these layers are further bridged

by oxalate anions to form 3D open-frameworks. Compound 6

is a new example of luminescent rare-earth oxalatophospho-

nates characterized by a significant red luminescence. The

results of our study indicate that by introduction of oxalate

Fig. 5 IR spectra of compounds 1–6.

Fig. 6 TGA curves of compounds 1–6. Fig. 7 Solid-state emission spectrum of compound 6 at room

temperature.


as the second ligand, we can obtain lanthanide oxalatophos-

phonates with well characterized crystal structures as well as

strong luminescence.

Experimental

Materials

2-Hydroxylphosphonoacetric acid (H3L) solution was

obtained from Taihe Chemical Factory (48.0 wt%). The

lanthanide(III) chlorides were prepared by the dissolving

corresponding lanthanide oxides (General Research Institute

for Nonferrous Metals, 99.99%) in hydrochloric acid followed

by recrystallization and drying. All other chemicals were used

as received without further purification.

Physical measurements

Elemental analyses (carbon and hydrogen) were performed

using a PE-2400 elemental analyzer. La, Ce, Pr, Nd, Sm, Eu

and P were determined by using an inductively coupled plasma

(ICP) atomic absorption spectrometer. IR spectra were

recorded on a Bruker AXS TENSOR-27 FT-IR spectrometer

with KBr pellets in the range 4000–400 cm�1. The X-ray

powder diffraction data were collected on a Bruker AXS D8

Advance diffractometer using Cu-Ka radiation (l=1.5418 A)

in the 2y range of 9–601 with a step size of 0.021. The

luminescence analysis was performed on a JASCO FP-6500

spectrofluorimeter (solid). TG analysis was performed on a

Perkin–Elmer Pyris Diamond TG–DTA thermal analysis

system in static air with a heating rate of 10 K min�1 from

50 to 800 1C.

Synthesis

[La(HL)(C2O4)0.5(H2O)2]�H2O (1). A mixture of LaCl3�6H2O (0.18 g, 0.50 mmol), H3L (0.50 ml, 2.00 mmol),

H2C2O4�2H2O (0.25 g, 2.00 mmol), and NaOH (0.08 g,

2.00 mmol) was dissolved in 8 mL distilled water. The resulting

solution was stirred for about 1 h at room temperature, sealed

in a 20 mL Teflon-lined stainless steel autoclave, and heated at

140 1C for 4 days under autogenous pressure. After the

mixture was cooled slowly to room temperature, colorless

block crystals were obtained in ca. 40.0% yield based on La.

C3H9O11PLa (390.98): calc.: C 9.21, H 2.30, P 7.93, La 35.53;

found: C 9.28, H 2.22, P 7.85, La 35.45%. IR (KBr) data: 3540

(br), 1644 (m), 1581 (s), 1432 (w), 1373 (w), 1309 (w), 1209 (s),

1178 (w), 1070 (s), 935 (m), 781 (w), 719 (w), 619 (w), 580 (w),

526 (w) cm�1.

[Ce(HL)(C2O4)0.5(H2O)2]�H2O (2). The procedure was the

same as that for 1 except that LaCl3�6H2O was replaced by

CeCl3�7H2O (0.19 g, 0.50 mmol). Yield: 81.0% (based on Ce).

C3H9O11PCe (392.19): calc.: C 9.18, H 2.29, P 7.90, Ce, 35.73;

found: C 9.11, H 2.21, P 7.95, Ce 35.65%. IR (KBr) data: 3465

(br), 2221 (w), 1648 (m), 1583 (s), 1425 (w), 1373 (w), 1311 (w),

1213 (s), 1074 (s), 937 (w), 781 (w), 721 (w), 615 (w), 588 (w),

524 (w) cm�1.

[Pr(HL)(C2O4)0.5(H2O)2]�H2O (3). The procedure was the


PrCl3�6H2O (0.18 g, 0.50 mmol). Yield: 76.0% (based on Pr).

C3H9O11PPr (392.98): calc.: C 9.16, H 2.29, P 7.89, Pr 35.86;

found: C 9.23, H 2.20, P 7.96, Pr 35.94%. IR (KBr) data: 3473

(br), 2915 (w), 1650 (m), 1575 (s), 1432 (m), 1371 (m), 1363

(m), 1317 (m), 1214 (s), 1064 (s), 927 (m), 831 (w), 777 (w), 696

(w), 617 (w), 572 (w), 516 (w) cm�1.

[Nd(HL)(C2O4)0.5(H2O)2]�H2O (4). The procedure was the


NdCl3�6H2O (0.18 g, 0.50 mmol). Yield: 55.0% (based on Nd).

C3H9O11PNd (396.31): calc.: C 9.08, H 2.27, P 7.82, Nd 36.40;

found: C 9.15, H 2.35, P 7.91, Nd 36.49%. IR (KBr) data:

3484 (br), 2915 (w), 1652 (s), 1577 (s), 1432 (m), 1369 (m), 1319

(m), 1209 (s), 1064 (s), 968 (w), 931 (m), 835 (w), 786 (w), 781

(w), 694 (w), 619 (m), 580 (m), 520 (m) cm�1.

[Sm(HL)(C2O4)0.5(H2O)2]�H2O (5). A mixture of SmCl3�6H2O (0.19 g, 0.50 mmol), H3L (0.50 ml, 2.00 mmol),

H2C2O4�2H2O (0.25 g, 2.00 mmol), and NaOH (0.08 g,

2.00 mmol) in 8 mL distilled water was sealed in an autoclave

equipped with a 20 mL Teflon liner, and then heated at 120 1C

for 4 days. After the mixture was cooled slowly to room

temperature, pale yellow block crystals were obtained in ca.

86.0% yield based on Sm. C3H9O11PSm (402.42): calc.: C 8.95,

H 2.25, P 7.69, Sm 37.36; found: C 9.03, H 2.33, P 7.63, Sm

37.28%. IR (KBr) data: 3477 (br), 3290 (br), 2925 (w), 1660

(s), 1575 (s), 1433 (m), 1366 (m), 1318 (m), 1212 (s), 1072 (s),

930 (m), 783 (w), 704 (w), 617 (m), 524 (m) cm�1.

[Eu(HL)(C2O4)0.5(H2O)2]�H2O (6). The procedure was the

same as that for 5 except that SmCl3�6H2O was replaced by

EuCl3�6H2O (0.19 g, 0.50 mmol). Yield: 45.0% (based on Eu).

C3H9O11PEu (404.03): calc.: C 8.92, H 2.24, P 7.67, Eu 37.61;

found: C 8.98, H 2.31, P 7.58, Eu 37.69%. IR (KBr) data: 3475

(br), 3292 (br), 2920 (w), 1670 (s), 1579 (s), 1435 (m), 1361 (m),

1317 (m), 1217 (s), 1180 (m), 1070 (s), 974 (w), 921 (m), 783

(w), 707 (w), 621 (m), 580 (w), 526 (w), 482 (m) cm�1.

Crystallographic determinations

Data collections for compounds 1–6 were performed on the

Bruker Smart APEX II X-diffractometer equipped with

graphite-monochromated Mo-Ka radiation (l = 0.71073 A)

at 293 � 2 K. An empirical absorption correction was applied

using the SADABS program. All structures were solved by

direct methods and refined by full-matrix least squares fitting

on F2 by SHELXS-97.39 All non-hydrogen atoms were refined

anisotropically. Hydrogen atoms of organic ligands were

generated geometrically with fixed isotropic thermal para-

meters, and included in the structure factor calculations.

Acknowledgements

This research was supported by grants from the Natural

Science Foundation of Liaoning Province of China

(20062140).

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39 G. M. Sheldrick, SHELXS-97, Program for the solution of crystalstructures, University of Gottingen, Germany, 1997.


The annular tautomerism of the curcuminoid NH-pyrazolesw

Pilar Cornago,*aPilar Cabildo,

aRosa M. Claramunt,

aLatifa Bouissane,

a

Elena Pinilla,bM. Rosario Torres

band Jose Elguero

c

Received (in Montpellier, France) 16th July 2008, Accepted 10th October 2008

First published as an Advance Article on the web 2nd December 2008

DOI: 10.1039/b812018h

The structures of four NH-pyrazoles, (E)-3,5-bis[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-1H-

pyrazole (3), (E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-5(3)-methyl-1H-pyrazole (4),

(E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-4,5(3)-dimethyl-1H-pyrazole (5) and (E)-3(5)-

[b-(3,4-dimethoxyphenyl)-ethenyl]-4-methyl-5(3)-phenyl-1H-pyrazole (8), have been determined by

X-ray crystallography. Compounds that have a phenol residue crystallize forming sheets that are

stabilized by a complex pattern of hydrogen bonds between a unique tautomer (4), or by a 2 : 1

mixture of both tautomers (5) (these tautomers being identical in the case of 3). Pyrazole 8, which

lacks OH groups, crystallizes in cyclic dimers that are stabilized by N–H� � �N hydrogen bonds.

The tautomerism in solution and in the solid state was determined by 13C and 15N CPMAS NMR

spectroscopy. For compounds 4, 5 and 8, the solid state results agree with those observed by

crystallography; the most abundant tautomer in solution coincides with the tautomer present in

the solid state (4 and 8) or with the most abundant tautomer in the crystal (5).

Introduction

Turmeric is a spice derived from the rhizomes of Curcuma

longa, which is a member of the ginger family.1 The bright

yellow color of turmeric comes mainly from polyphenolic

pigments known as curcuminoids. Curcumin (1) (Scheme 1)

is the principal curcuminoid found in turmeric, and is gen-

erally considered to be its most active constituent. In addition

to its use as a spice and a pigment, turmeric has been used in

India for medicinal purposes for centuries. More recently,

evidence that 1 may have anti-inflammatory and anti-cancer

activities has renewed scientific interest in its potential to

prevent and treat disease. 1 is also an effective scavenger of

reactive oxygen and nitrogen species in vitro. In addition to its

direct antioxidant activity, 1 has been found to inhibit PLA2,

COX-2 and 5-LOX activity in cultured cells. It has also been

found to inhibit NF-kB-dependent gene transcription, and to

inhibit the induction of COX-2 and iNOS in cell culture and

animal studies.2 1 has been found to induce cell cycle arrest

and apoptosis in a variety of cancer cell lines grown in

cultures. The ability of 1 to induce apoptosis in cultured

cancer cells has generated scientific interest in its potential to

prevent some types of cancer. Oral administration of 1 has

been found to inhibit the development of chemically-induced

cancer in animal models of oral, stomach, liver and colon

cancer.

We have devoted a series of papers to the annular tauto-

merism of NH-pyrazoles 2 (2a vs. 2b),3,4 and decided to study

those derived from 1 and related b-diketones.Pyrazole 3, which is derived from 1, has been prepared many

times since 1991.5–11 It has been described as a pale yellow

solid that melts at 211–2145 or 2157 1C.

The activity of the curcuminoid pyrazoles covers domains

such as anti-inflammatory (5-lipooxygenase and cyclooxygen-

ase inhibitors)5,8 and anti-tumoral (anti-angiogenic)6–8 agents,

and drugs for the treatment of Alzheimer’s disease (AD;

potent g-secretase inhibitors, potent ligands for fibrillar

Ab42 aggregates, tau aggregation inhibitors and depolymeriz-

ing agents for tau aggregates).10,11 Particularly promising for

treating reduced cognitive functions is 4,40-[(1-phenyl-1H-

pyrazole-3,5-diyl)di-(1E)-2,1-ethenediyl]bis(2-methoxyphenol)

(CNB-001), the product obtained by reacting 1 with phenyl-

hydrazine.12 In the last of these applications, curcumin-

derived pyrazoles were synthesized in order to minimize

the metal chelation properties of 1. The reduced rotational

freedom and the absence of stereoisomers were anticipated

to enhance the inhibition of g-secretase. Accordingly, the

replacement of the 1,3-dicarbonyl moiety by isosteric hetero-

cycles, such as pyrazoles, turned these compounds into very

interesting candidates for AD research.

Scheme 1 The structure of curcumin (1) and the tautomerism of

pyrazoles 2.

aDepartamento de Quımica Organica y Bio-Organica, Facultad deCiencias, UNED, Senda del Rey 9, E-28040 Madrid, Spain.E-mail: [email protected]; Fax: +34 913988372;Tel: +34 913987323

bDepartamento de Quımica Inorganica I, Facultad de CienciasQuımicas, Universidad Complutense de Madrid (UCM), 28040Madrid, Spain

c Instituto de Quımica Medica, CSIC, Juan de la Cierva 3, E-28006Madrid, Spainw CCDC reference numbers 690489–690492. For crystallographic datain CIF or other electronic format see DOI: 10.1039/b812018h



The aim of this paper is to determine and discuss the

structure, tautomerism and possible proton transfer in the

solid state (SSPT) of six NH-pyrazoles by using a combination

of X-ray crystallography and 13C/15N NMR spectroscopy.

The nomenclature used in the text and in the experimental is

not in accordance with IUPAC rules. For all of the com-

pounds with phenolic hydroxyl groups, 3–6, the phenol system

has the highest priority; however, using IUPAC nomenclature

here would be at the expense of comparability and clearness.

For instance, compound 4 would be 2-methoxy-4-[(E)-2-

(5-methyl-1H-pyrazol-3-yl)vinyl]phenol under IUPAC rules,

rather than (E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)ethenyl]-

5(3)-methyl-1H-pyrazole. In order to prioritize comparability

over correct nomenclature, we have named all of the com-

pounds as pyrazole derivatives.


Synthesis

All of the compounds discussed in this work (Scheme 2) are

reported in the experimental section. They were prepared by

the reaction of hydrazine with the corresponding b-diketone,the most common method of synthesizing pyrazoles,13 which

in the case of 3 was 1.14

X-Ray structure determination

The structures of pyrazoles 3 (derived from 1), 4, 5 and 8 have

been determined by X-ray crystallography.

Concerning tautomerism, in the case of 3, tautomers 3a and

3b are identical. In the case of 4, the only tautomer present

is 3-(3-methoxy)-4-hydroxy-styryl-5-methyl-1H-pyrazole (4a).

In the case of 5, there is a 2 : 1 mixture of 3-(3-methoxy)-4-

hydroxy-styryl-4,5-dimethyl-1H-pyrazole (5a) and 3,4-di-

methyl-5-(3-methoxy)-4-hydroxy-styryl-1H-pyrazole (5b). In

the case of 8, the only observed tautomer is 3-phenyl-4-

methyl-5-(3-methoxy)-4-hydroxy-styryl-1H-pyrazole (8b). The

main data are collected in Table 1 and Table 2. A charac-

teristic feature of the geometry of NH-pyrazoles is that the

angle centered at N1 (the atom bearing the NH proton) is

always larger than that centered at N2, about 112 and 1041,

respectively.15

Crystals of sufficient quality for X-ray diffraction analysis

were obtained for compounds 3 (1 : 1 H2O/EtOH), 4 (1 : 1 : 1

CH2Cl2/hexane/EtOH), 5 (1 : 1 : 1 CH2Cl2/hexane/EtOH) and

8 (1 : 1 : 1 CH2Cl2/hexane/EtOH) from their respective solvent

mixtures. Table 1 shows selected bond lengths and angles for

each of these compounds, and Table 2 shows the distances and

angles of the intermolecular hydrogen bonds.

One crystallographically-independent molecule was identi-

fied in the structural determination of 3, where the pyrazole

and phenyl rings were co-planar, with bond distances and

angles within normal ranges (Fig. 1). The intermolecular

hydrogen bonds led to layers parallel to (1 0 1), as shown in

Fig. 2.

Scheme 2 The structures of the NH-pyrazoles.

Table 1 The bond lengths (A) and angles (1) for compounds 3, 4, 8and the three crystallographically-independent molecules of 5

3 4 5(1) 5(2) 5(3) 8

N1–N2 1.354(3) 1.365(3) 1.349(4) 1.352(4) 1.359(3) 1.351(3)N2–C3 1.347(4) 1.339(3) 1.341(5) 1.348(5) 1.349(5) 1.339(4)C3–C4 1.399(4) 1.398(4) 1.388(6) 1.410(5) 1.401(6) 1.415(4)C4–C5 1.373(4) 1.369(3) 1.381(6) 1.366(5) 1.374(6) 1.379(4)C5–N1 1.353(4) 1.340(3) 1.332(6) 1.346(5) 1.331(5) 1.360(4)C3–C6 1.445(4) 1.453(3) — 1.463(5) 1.450(6) 1.377(3)C5–C6 — — 1.446(7) — — 1.444(4)C6–C7 1.327(4) 1.325(3) 1.309(1) 1.310(6) 1.304(4) 1.333(4)C7–C8 1.467(4) 1.472(3) 1.460(1) 1.457(5) 1.475(6) 1.460(4)C3–C15 — — 1.484(6) — — —C5–C15 1.450(4) 1.484(2) — 1.490(5) 1.509(6) —C15–C16 1.322(4) — — — — —C16–C17 1.463(4) — — — — —C10–O2 1.376(4) 1.367(3) 1.359(6) 1.367(4) 1.372(5) 1.372(3)O2–C14 1.433(4) 1.419(3) 1.424(6) 1.430(5) 1.442(5) 1.416(4)C11–O1 1.368(4) 1.369(3) 1.372(5) 1.361(5) 1.363(5) 1.371(3)C15–O1 — — — — — 1.418(4)C19–O4 1.366(4) — — — — —O4–C23 1.416(4) — — — — —C20–O3 1.382(4) — — — — —N2–N1–C5 112.2(3) 112.7(2) 112.2(4) 111.8(3) 111.6(3) 112.4(2)N1–N2–C3 105.3(2) 104.4(2) 104.4(3) 105.2(3) 104.7(3) 105.3(2)

Table 2 The bond lengths (A) and angles (1) for the hydrogen bondsin compounds 3, 4, 5 and 8

Compound D–H� � �A dD–H dH� � �A dD� � �A +D–H� � �A

3 O3–H3� � �O4 1.10 1.98 2.647(4) 115.1N1–H1B� � �O3a 1.06 1.93 2.864(4) 144.7O1–H1A� � �N2b 1.17 1.79 2.811(4) 142.7O3–H3� � �O1c 1.10 2.26 2.825(4) 108.7

4 O1–H1A� � �N2d 0.99 1.86 2.832(3) 167.5N1–H1B� � �O2e 1.07 2.17 2.962(3) 128.6

5 O13–H113� � �N21 1.16 1.81 2.782(5) 137.3N12–H12� � �N23 1.10 1.82 2.914(5) 175.6O11–H111� � �O13f 0.92 2.03 2.813(4) 141.3O12–H112� � �N22g 1.14 1.57 2.673(4) 159.4N11–H11� � �O11g 1.08 2.01 2.951(5) 144.3N13–H13� � �O12i 1.02 1.93 2.853(4) 148.6

8 N1–H1� � �N2j 0.90(4) 2.07(4) 2.872(3) 147(4)

Symmetry transformations used to generate equivalent atoms: a �x+

2, y � 12, �z + 1

2.b �x + 1, y + 1

2, z +32.

c x + 1, y, z � 1. d �x + 1,

�y + 2, �z + 1. e x, �y + 32, z � 1

2.f �x + 4, y � 1

2, �z + 32.

g �x +

1, y � 12, �z + 1

2.h �x + 4, y + 1

2, �z + 32.

i �x + 1, y + 12, �z +

12.

j �x + 2, �y, �z + 1.


Fig. 3 shows an ORTEP representation of the asymmetric

unit of compound 4, a non-planar molecule with a dihedral

angle of 19.0(1)1 between the pyrazole and phenyl rings.

Dimers (O1–H1A–N2) linked by hydrogen bonds

(N1–H1B–O2) led to layers parallel to (1 0 0), as shown in

Fig. 4.

The asymmetric unit of compound 5 is presented in Fig. 5.

The crystal consists of three crystallographically-independent,

almost planar molecules, held together by hydrogen bonds

that form a trimer, which, through additional hydrogen

bonding, forms layers parallel to (–1 0 3), as shown in Fig. 6.

Fig. 7 shows the non-planar molecule of compound 8, with

a dihedral angle of 15.7(1)1 between the pyrazole and

the phenyl ring at the 3-position, and 36.5(1)1 between the

pyrazole and the phenyl ring of the styryl group at the

5-position. Molecules of 8 are centrosymmetrically linked by

hydrogen bonds (Table 2), giving rise to dimers, and these

species are within van der Waals distances (Fig. 8).

The cyclic N–H� � �N hydrogen-bonded motifs (cyclamers) of

NH-pyrazoles have been studied on several occasions.4d,16,17

These motifs are characteristic of NH-pyrazoles lacking sub-

stituents that bear hydrogen bonding functional groups, such

as –OH or –CO2H. These groups, as well as solvent molecules

like H2O and ROH, participate in the hydrogen bonding

network that determines the secondary structure of the crys-

tals, destroying the (N–H� � �N)n hydrogen bonds.18–20 In three

of the compounds described in the present paper, those

bearing phenol groups (3, 4 and 5) form several hydrogen

bonds involving the OH group: 3 (O–H� � �N, N–H� � �O,

O–H� � �O), 4 (O–H� � �N, N–H� � �O) and 5 (O–H� � �N,

N–H� � �O, O–H� � �O, N–H� � �N; present as two molecules of

tautomer 5a and one molecule of tautomer 5b). In the case of

8, which lacks phenol groups, the compound crystallizes as a

dimer. This kind of cyclamer is characteristic of NH-pyrazoles

that are substituted with phenyl groups at the 3- and 5-posi-

tions,16 to which compound 8 is clearly related.

Fig. 1 The X-ray molecular structure of compound 3 (ORTEP plot,

35% probability for the ellipsoids).

Fig. 2 The view along the a axis of 3, showing the formation of layers

due the intermolecular hydrogen bonds.



Fig. 4 The view along the b axis of 4, showing the formation of layers



NMR study

We have reported the 1H, 13C and 15N NMR results concerning

compounds 3–8 in Table 3, Table 4 and Table 5, respectively.

These data have been collected with the aim of determining the

tautomeric equilibrium constants by simple integration.

Although it has been pointed out that only 1H NMR signal

intensities are reliable for the determination of populations, in

our experience, 13C and 15N signals can also been used in

connection with signals related by tautomerism, i.e. carbon or

nitrogen atoms linked to the same substituents.3c The assign-

ments of the signals were based on standard 2D experiments,

on the values of coupling constants (auto-consistency) and by

comparison with other NH-pyrazoles where tautomerization

is blocked.21

We have illustrated with one example the kind of spectra

that we obtained (Fig. 9). The spectrum corresponds to

compound 5 in HMPA-d18, concentration 0.10 M and tem-

perature 268 K (Table 4). The region of the methyl groups

shows two narrow signals corresponding to the most abundant

tautomer, and two broad signals corresponding to the less

abundant one, as expected by simple consideration of the

energy profile.

For compounds whose structure had not been determined

by crystallography, we relied on CPMAS NMR results: 6b and

7b were the only tautomers present in the solid state

(see Table 4 and Table 5). We are aware that solid state

NMR and single crystal X-ray diffraction do not show exactly

the same properties, for instance, static vs. dynamic disorder.3b

To avoid further complications, we used fine powders for

Fig. 5 The X-ray molecular structure of compound 5 (ORTEP plot, 40% probability for the ellipsoids).

Fig. 6 The view along the a axis of 5, showing the formation of layers




Fig. 8 The view along the a axis of 8, showing the formation of

dimers.


Table

3The

1H

NMR

chem

icalshifts

(d)and

1H–1H

couplingconstants

ofcompounds3–8(J/H

z)in

DMSO-d

6andHMPA-d

18solutionsa

Compound

R1

R2

R3

Solvent

Conc./M

T/K

NH

R2

H3

H4

H6

OMe

OR

3H7

H8

R1

Tautomerism

3*

HH

DMSO

0.12

300

12.80

6.61(H

)6.76

6.93

7.13

3.82

9.17(H

)7.03

6.91

—Average

4CH

3H

HDMSO

0.07

300

12.40

6.20(H

)6.74

6.91

7.12

3.81

9.15(H

)6.95

6.88

2.19(M

e)Average

HMPA

0.07

300

13.26

6.11(H

)6.85/7.16

3.80

10.26(H

)6.85/7.16

2.23(M

e)B50%

a

HMPA

0.07

300

13.20

6.11(H

)6.85/7.16

3.80

10.24(H

)6.85/7.16

2.15(M

e)B50%

b

HMPA

0.10

276

13.34

6.19(H

)6.85

6.85

7.06

3.80

10.42(H

)7.20

6.87

2.25(M

e)B50%

a

HMPA

0.10

276

13.27

6.11(H

)6.85

6.85

7.06

3.80

10.34(H

)6.92

6.83

2.14(M

e)B50%

b

5CH

3CH

3H

DMSO

0.07

300

12.29

2.03(M

e)6.75

6.91

7.13

3.83

9.08(H

)6.95

6.86

2.10(M

e)Average

HMPA

0.10

300

13.16

2.04(M

e)6.87

6.92

6.97

3.80

10.28(H

)7.21

6.82

2.04(M

e)35%

a

HMPA

0.10

300

13.10

2.04(M

e)6.87

6.92

6.97

3.80

10.20(H

)7.21

6.82

2.07(M

e)65%

b

HMPA

0.10

268

13.25

2.05(M

e)6.87

6.96

7.00

3.81

10.44(H

)7.25

6.88

2.05(M

e)35%

a

HMPA

0.10

268

13.20

2.05(M

e)6.87

6.96

7.00

3.81

10.38(H

)7.25

6.88

2.07(M

e)65%

b

6C6H

5H

HDMSO

0.11

300

12.96

6.88(H

)6.78

6.96

7.15

3.84

9.10(H

)7.10

6.95

7.80(o)

36%

a

7.43(m

)7.31(p)

DMSO

0.11

300

13.18

6.88(H

)6.78

6.96

7.15

3.84

9.21(H

)7.10

6.95

7.80(o)

64%

b

7.43(m

)7.31(p)

7C6H

5H

CH

3DMSO

0.06

300

13.00

6.87(H

)6.96

7.06

7.19

3.83

3.78(M

e)7.14

7.03

7.80(o)

40%

a

7.43(m

)7.32(p)

DMSO

0.06

300

13.21

6.87(H

)6.96

7.06

7.19

3.83

3.78(M

e)7.14

7.03

7.80(o)

60%

b

7.43(m

)7.32(p)

8C6H

5CH

3CH

3DMSO

0.05

300

12.94

2.29(M

e)6.95

7.07

7.25

3.84

3.77(M

e)7.14

7.06

7.65(o)

Richin

b

7.45(m

)7.34(p)

HMPA

0.06

268

13.94

2.34(M

e)7.09

7.09

7.25

3.88

3.84(M

e)7.45

7.13

7.70(o)

b

7.45(m

)7.31(p)

aThecouplingconstants

were,

onaverage:

3JH3–H4=

8.0

Hz,

4JH4–H6=

2.0

Hz(notalwaysobserved)and

3JH7–H8trans=

16.5

Hz.


Table 4 The 13C NMR chemical shifts (d) and 1H–13C coupling constants (J/Hz) in DMSO-d6 and HMPA-d18 solutions, and under CPMASconditionsa

Compound R1 R2 R3 Solvent Conc./M T/KCa Cb Cc R2 C1 C2 C3

TautomerismC4 C5 C6 C7 C8 OCH3 R1

3 * H H DMSO 0.12 300 151.0 99.3 142.0 — (H) 147.9 146.8 115.6 No tautomerism120.1 128.4 109.5 129.8 112.9 (C8) 55.6

118.4 (C80)CPMAS — 300 150.2 95.5 142.8 — (H) 147.4 145.1 114.5 No tautomerism

N.o.b 127.1 106.3 129.9 111.9 53.5

56.5

4 CH3 H H DMSO 0.36 300 149.6 101.3 140.5 — (H) 147.9 146.6 115.7 Average

119.9 128.6 109.5 129.0 117.4 55.6 11.6 (Me)

HMPA 0.10 276 150.9 100.0 142.4 — (H) 148.7 148.6 115.7 B50% a

119.4 128.4 110.5 128.7 113.2 55.9 10.8 (Me-5)

138.4 101.6 148.9 — (H) 148.7 146.9 115.7 B50% b

119.6 128.4 110.5 129.4 119.3 55.9 13.9 (Me-3)

CPMAS — 300 151.5 101.1 142.4 — (H) 148.8 143.3 115.3 a

120.5 129.9 113.2 129.9 113.2 55.9 9.9 (Me-5)

115.3

5 CH3 CH3 H DMSO 0.07 300 141.6 110.4 141.6 8.1 (Me) 147.9 146.6 115.6 Average

119.8 128.8 109.6 127.9 114.9 55.7 10.6

HMPA 0.08 300 147.5 109.9 135.7 8.4 (Me) 149.0 148.7 116.1 35% a

119.6 129.0 111.5 127.7 118.7 56.3 11.9 (br)

HMPA 0.08 300 138.3 109.9 145.8 8.4 (Me) 149.0 148.7 116.1 65% b

119.6 129.0 111.5 128.4 112.7 56.3 11.9 (br)HMPA 0.10 268 147.5 110.0 135.7 8.8 (br, Me) 148.8 148.6 115.8 35% a

119.5 128.8 110.7 127.6 118.5 55.9 9.1 (br)HMPA 0.10 268 138.3 109.9 145.8 8.4 (Me) 148.8 148.6 115.8 65% b

119.5 128.8 110.7 128.3 112.4 55.9 12.1CPMAS — 300 145.9 110.2 138.6 9.8 (Me) 148.8 146.6 121.8 66% a

123.4 130.9 105.5 128.9 117.0 55.3 11.2 (br)137.5 112.0 146.6 9.8 (Me) 148.8 146.6 121.8 34% b

123.4 130.9 105.5 128.9 119.0 55.3 11.2 (br)

6 C6H5 H H DMSO 0.11 300 151.4 100.4 140.3 — (H) 147.9 146.6 115.3 36% a

122.1 128.1 109.5 130.1 118.4 55.5 132.0 (i)125.0 (o)128.7 (m)127.5 (p)

DMSO 0.11 300 142.6 99.5 151.0 — (H) 147.9 147.1 115.6 64% b

120.2 128.1 109.5 130.1 112.7 55.6 133.6 (i)125.1 (o)128.7 (m)127.5 (p)

6b C6H5 H H CPMAS — 300 144.0 103.5 152.6 — (H) 148.3 116.0116.0 129.0 112.3 129.0 113.5 54.0 133.2 (i)

126.4 (o)129.0 (m)129.0 (p)

7 C6H5 H CH3 DMSO 0.11 300 151.3 99.2 142.8 — (H) 149.0 149.0 111.9 40% a

119.4 129.4 108.9 128.9 113.6 55.51 (C1) 133.7 (i)55.45 (C2) 125.0 (o)

128.6 (m)127.4 (p)

DMSO 0.11 300 142.4 99.8 150.9 — (H) 149.0 149.0 111.9 60% b

119.9 129.4 108.9 129.7 113.6 55.51 (C1) 133.7 (i)55.45 (C2) 125.0 (o)

128.6 (m)127.4 (p)


CPMAS NMR, obtained by grinding the same batch of

crystals that we used for X-ray crystallography.

Percentages of tautomers and equilibrium constants

Although some exceptions are known, the assumption of the

identical nature of the most stable tautomer in solution and

the tautomer present in the crystal is one of the most basic

tenets in tautomerism.3b,3c,17b The results in Table 6 confirm

this principle for compounds 4, 5 and 8, and allow us

to conclude that in the solid state, 6 should crystallize as

6b and 7 as 7b, or at least in cyclamers where 6b and 7b are

predominant.

Compound 5 exists in the solid state as a 66% 5a/34% 5b

mixture and in HMPA as a 35% 5a/65% 5b mixture, thus

being an exception to the rule of similarity between solution

and solid state. However, the difference in energy at 300 K

between the two situations is only of 3.2 kJ mol�1.

Conclusions

The structure, tautomerism and absence of SSPT have been

determined for six NH-pyrazoles by a combination of X-ray

crystallography and 13C/15N NMR spectroscopy. Two of the

conditions required to observe SSPT in NH-pyrazoles are the

identity (or, at least, strong similarity) of the substituents at

the 3- and 5-positions, and the formation of cyclic structures,

cyclamers, linked by N–H� � �N hydrogen bonds. Compound 3

has the same substituent at both positions (tautomer 3a is

identical to tautomer 3b), but crystallizes in a complex

network of hydrogen bonds involving the OH groups. Com-

pound 8 crystallizes as a dimer, but with only one tautomer

present (8a). Thus, none of the compounds of Table 6 display

SSPT. Finally, compound 5 is the only known example of an

NH-pyrazole that crystallizes as a 2 : 1 mixture of two

tautomers (there are examples of 2 : 2 and 3 : 1 mixtures,

but in cyclic tetramers17b,22).

Experimental

The melting points of pyrazoles 3–8 were determined by

differential scanning calorimetry (DSC) on a Seiko DSC

220C connected to a Model SSC5200H Disk Station; for the

other compounds, a hot stage microscope was used. Thermo-

grams (sample size 0.003–0.0010 g) were recorded at a scan-

ning rate of 2.0 1C min�1. Thin-layer chromatography

(TLC) was performed using Merck silica gel (60 F254) and

compounds were detected with a 254 nm UV lamp. Silica gel

(60–320 mesh) was employed for routine column chromato-

graphy separations. Elemental analyses for carbon, hydrogen

and nitrogen were carried out by the Microanalytical Service

of the Universidad Complutense of Madrid on a Perkin-Elmer

240 analyzer.

Table 4 (continued )

Compound R1 R2 R3 Solvent Conc./M T/KCa Cb Cc R2 C1 C2 C3

TautomerismC4 C5 C6 C7 C8 OCH3 R1

7b C6H5 H CH3 CPMAS — 300 143.1 96.3 149.8 — (H) 148.8 148.8 110.6120.9 129.1 108.0 129.1 110.6 53.5 (C1*) 132.1 (i)

56.1 (C2*) 125.2 (o)129.1 (m)126.2 (p)

8 C6H5 CH3 CH3 DMSO 0.31 300 141.7 110.7 147.1 9.4 (Me) 149.1 148.8 111.9 Aver.119.9 130.0 109.1 128.4 114.5 55.6 (C1) 133.2 (i) Rich in b

55.5 (C2) 127.1 (o)128.5 (m)127.3 (p)

HMPA 0.06 268 139.7 110.4 149.5 10.1 (Me) 149.8 149.3 112.1 b

120.2 130.9 109.1 128.8 113.2 55.9 (C1) 135.9 (i)55.9 (C2) 127.2 (o)

128.6 (m)126.9 (p)

CPMAS — 300 140.7 112.5 148.8 9.2 (Me) 148.8 148.8 112.5 b

124.7 130.2 110.7 130.2 117.1 54.7 134.6 (i)128.5 (o)130.2 (m)126.7 (p)

a The 1J coupling constants are not reported; their average values are: pyrazole C4–Hb = 175 Hz; phenyl CH = 159 Hz except C4–H and

C6–H = 156 Hz; olefin C–H = 155 Hz; OCH3 = 144 Hz; C–Me substituents: 126.5 Hz. The other couplings (Hz) are: 2J = 2.2 (C1), 2J = 4.5

(C7), 2J = 5.9 (Cb–Me4); 3J = 8.4 (C1), 3J = 7.3 (C2), 3J = 5.8 (C4), 3J = 6.8 (C5), 3J = 6.0 (C6), 3J = 4.5 (C7), 3J = 2.4 (Cb–H).b Not observed.


General procedure for the preparation of pyrazole derivatives

(3–8)

Compounds 3–8 were prepared by reacting the corresponding

b-diketones23 (1 mmol) with hydrazine hydrate 98% (1.5 mmol)

in acetic acid (5 mL). After heating at reflux for 2 h, the

reaction mixture was poured into water, and the precipitate

filtered off, washed with water and dried. The solid was

purified by column chromatography using ethyl acetate as

the eluent.

(E)-3,5-Bis[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-1H-

pyrazole (3)

3 was prepared from purified commercially available 1. The

compound was obtained as a colourless solid after recrystalli-

zation from H2O/EtOH (1 g, 2.74 mmol, 63%). Mp: 217.1 1C,

lit.: 211–214 1C5 or 215 1C.8 Anal. calc. for C21H20N2O4

(364.14): C, 69.22; H, 5.53; N, 7.69; found: C, 68.79; H,

5.53; N, 7.70%.

(E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-5(3)-methyl-

1H-pyrazole (4)

4 was prepared from (E)-6-(4-hydroxy-3-methoxyphenyl)hex-

5-ene-2,4-dione.23 The compound was obtained as a colourless

solid after recrystallization from CH2Cl2/hexane/EtOH

(251 mg, 1.1 mmol, 85%). Mp: 141.6 1C. Anal. calc. for

C13H14N2O2 (230.11): C, 67.26; H, 6.44; N, 12.11; found: C,

67.81; H, 6.13; N, 12.17%.

(E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-4,5(3)-

dimethyl-1H-pyrazole (5)

5 was prepared from (E)-6-(4-hydroxy-3-methoxyphenyl)-3-

methylhex-5-ene-2,4-dione.23 The compound was obtained as

a colourless solid after recrystallization from CH2Cl2/hexane/

EtOH (180 mg, 0.73 mmol, 61%). Mp: 176.1 1C. Anal. calc.

for C14H16N2O2 (244.12): C, 68.46; H, 6.61; N, 11.35; found:

C, 68.83; H, 6.60; N, 11.47%.Table

5The

15N

NMR

chem

icalshifts

(d)in

DMSO-d

6andHMPA-d

18solutions,andunder

CPMASconditions

Compound

R1

R2

R3

Solvent

Conc./M

T/K

N–H

–N=

%a

%b

PTa

3*

HH

CPMAS

—300

�180.8

�100.6

50

50

No

4CH

3H

HHMPA

0.10

276

�180.5

N.o.b

B50

B50

No

�173.6

4a

CH

3H

HCPMAS

—300

�177.7

�100.9

100

0No

5CH

3CH

3H

HMPA

0.08

300

�185.6

(major)

N.o.

30

70

No

�175.9

CPMAS

—300

�187.8

�111.2

66

34

No

�172.0

(major)

�103.6

(major)

6b

C6H

5H

HCPMAS

—300

–181.5

�105.3

0100

No

8b

C6H

5CH

3CH

3HMPA

0.06

268

–182.2

N.o.

0100

No

CPMAS

—300

–181.3

�98.7

0100

No

aProtontransfer

bNotobserved

Fig. 9 The methyl group region of the 13C NMR spectrum of 5.


(E)-3(5)-[b-(4-hydroxy-3-methoxyphenyl)-ethenyl]-5(3)-phenyl-

1H-pyrazole (6)

6 was prepared from (E)-5-(4-hydroxy-3-methoxyphenyl)-1-

phenylpent-4-ene-1,3-dione.23 The compound was obtained as



for C18H16N2O2 (292.12): C, 73.95; H, 5.54; N, 10.11; found:

C, 73.95; H, 5.54; N, 10.11%.

(E)-3(5)-[b-(3,4-dimethoxyphenyl)-ethenyl]-5(3)-phenyl-1H-

pyrazole (7)

7 was prepared from (E)-5-(3,4-dimethoxyphenyl)-1-phenyl-

pent-4-ene-1,3-dione.23 The compound was obtained as a

colourless solid after recrystallization from CH2Cl2/hexane/


for C19H18N2O2 (306.37): C, 74.48; H, 5.92; N, 9.14; found: C,

74.21; H, 5.82; N, 9.16%.

(E)-3(5)-[b-(3,4-dimethoxyphenyl)-ethenyl]-4-methyl-5(3)-

phenyl-1H-pyrazole (8)

8 was prepared from (E)-5-(3,4-dimethoxyphenyl)-2-methyl-1-

phenylpent-4-ene-1,3-dione.23 The compound was obtained as



for C20H20N2O2 (320.39): C, 74.97; H, 6.29; N, 8.74; found: C,

74.28; H, 6.14; N, 8.77%.

X-Ray data collection and structure refinement (compounds 3, 4,

5 and 8)

Data collection for all of the compounds was carried out at

room temperature on a Bruker Smart CCD diffractometer

Table 6 The tautomeric composition of 3–8 (Sty: Ar–CHQCH–)a

Compound Tautomers X-Ray CPMAS DMSO HMPA

3 a, b: 3,5-BisSty 3a = 3b 3a = 3b 3a = 3b No PTb N. M.4 a: 3-Sty-5-Me 4a 4a Average rich in 4a B50% 4a

b: 3-Me-5-Sty B50% 4b

5 a: 3-Sty-5-Me 66% 5a 66% 5a Average rich in 5b 35% 5a

b: 3-Me-5-Sty 34% 5b 34% 5b 65% 5b

6 a: 3-Sty-5-Ph N. M. 6b 36% 6a N. M.b: 3-Ph-5-Sty 64% 6b

7 a: 3-Sty-5-Ph N. M. 7b 40% 7a N. M.b: 3-Ph-5-Sty 60% 7b

8 a: 3-Sty-5-Ph 8b 8b Average rich in 8b 8b

b: 3-Ph-5-Sty

a N. M. means not measured. b Proton transfer

Table 7 The crystal and structure refinement data for compounds 3, 4, 5 and 8

Crystal data 3 4 5 8

Empirical formula C21H20N2O4 C13H14N2O2 C14H16N2O2 C20H20N2O2

Formula weight 364.39 230.26 244.29 320.38Crystal system Monoclinic Orthorhombic Monoclinic OrthorhombicSpace group P2(1)/c Pbca P2(1)/c PbcaUnit cell dimensions a/A 8.2394(10) 13.2563(15) 8.519(2) 13.2363(13)

b/A 14.0198(17) 7.6962(9) 12.964(4) 8.2769(8)c/A 16.306(2) 22.855(3) 34.615(10) 30.673(3)b (1) 101.060(3) — 94.607(7) —

Volume/A3 1848.7(4) 2331.7(5) 3810.6(19) 3360.4(6)Z 4 8 12 8Density (calculated)/Mg m�3 1.309 1.312 1.277 1.267Absorption coefficient/mm�1 0.092 0.090 0.087 0.083Scan technique o and j o and j o and j o and jF(000) 768 976 1560 1360Range for data collection (1) 1.93 to 25.00 1.78 to 27.00 1.18 to 25.00 1.33 to 25.00Index ranges �9, �16, �18 to 9, 16, 19 �13, �9, �29 to 16, 9, 29 �9, �15, �41 to 10, 15, 41 �15, �9, �36 to 10, 9, 32Reflections collected 13 998 19 397 28 784 16 484Independent reflections 3244 2541 6720 2954Observed reflections [I 4 2s(I)] 1418 1248 2855 1655Rint 0.1198 0.0889 0.0905 0.0708Completeness to y (%) 99.6 100.0 100.0 99.9Data/restraints/parameters 3244/0/245 2541/0/156 6720/2/497 2954/0/224Goodness-of-fit on F2 0.912 1.034 0.984 1.074R1a 0.0539 0.0508 0.0769 0.0507wR2b (all data) 0.1808 0.1768 0.2486 0.1848Largest differential peak andhole/eA�3

0.232 and �0.278 0.214 and �0.247 0.950 and �0.377 0.193 and �0.192

a R1 =P

||Fo| � |Fc||/P

|Fo|.b wR2 =

P[w(Fo

2 � Fc2)2]/

P[w(Fo

2).


using graphite-monochromated Mo-Ka radiation (l =

0.71073 A) operating at 50 kV and 30 mA. In all cases, the

data were collected over a hemisphere of the reciprocal space

by the combination of three exposure sets. Each frame ex-

posure time was either 10 or 20 s, covering 0.31 in o. The cellparameters were determined and refined by a least-squares fit

of all reflections collected. The first 100 frames were re-

collected at the end of the data collection to monitor crystal

decay, and no appreciable decay was observed. A summary of

the fundamental crystal and refinement data is given in

Table 7. The structures of all the compounds were solved by

direct methods and conventional Fourier synthesis, and re-

fined by full matrix least-squares on F2 (SHELXL-97).24 All

non-hydrogen atoms were refined anisotropically.

In all cases, the hydrogen atoms were calculated, included

and refined as riding on their respective carbon-bonded atom

with a common anisotropic displacement. The rest of the

hydrogen atoms, i.e. those bonded to nitrogen or oxygen

atoms, were located in a Fourier difference synthesis, and in

all cases were included and refined as riding on their respective

bonded atoms for 3, 4 and 5, while for 8, its coordinates were

refined and the thermal factors kept constant. The longer O–H

bond distances in some of the hydroxyl groups are due to the

formation of hydrogen bonds.25

The largest peaks and holes in the final difference map were

0.232 and �0.278, 0.214 and �0.247, 0.950 and �0.377, and0.193 and �0.192 eA�3 for 3, 4, 5 and 8, respectively. The final

R1 and wR2 values were 0.0539 and 0.1808, 0.0508 and 0.1768,

0.0769 and 0.2486, and 0.0507 and 0.1848 for 3, 4, 5 and 8,

respectively.

NMR spectroscopy

Solution NMR spectra. Solution NMR spectra were re-

corded on a Bruker DRX 400 (9.4 T; 400.13 MHz for 1H,

100.62 MHz for 13C and 40.56 MHz for 15N) spectrometer

fitted with a 5 mm inverse detection H–X probe and equipped

with a z-gradient coil at 300 K. 1H and 13C NMR chemical

shifts (d) are referenced to Me4Si; for15N NMR, nitromethane

(0.00) was used as an external standard. Typical parameters

for the 1H NMR spectra were: a spectral width of 5787 Hz, a

pulse width of 7.5 ms, an attenuation level of 0 dB and a

resolution of 0.34 Hz per point. Typical parameters for the 13C

NMR spectra were: a spectral width of 21 kHz, a pulse width

of 10.6 ms, an attenuation level of �6 dB, a relaxation delay of

2 s and a resolution of 0.63 Hz per point; WALTZ-16 was used

for broadband proton decoupling and the FIDs were multi-

plied by an exponential weighting (lb = 1 Hz) before Fourier

transformation. 2D inverse proton detected heteronuclear

shift correlation spectra, gs-HMQC (1H–13C) and gs-HMBC

(1H–13C), were acquired and processed using standard Bruker

NMR software. Typical parameters for these spectra were:

a spectral width of 5787 Hz for 1H and 20.5 kHz for 13C, a

1024 � 256 data set, number of scans = 2 (gs-HMQC) or 4

(gs-HMBC), a relaxation delay of 1 s, and a delay for the

evolution of 13C–1H coupling constants of 3 ms (gs-HMQC)

or 60 ms (gs-HMBC). The FIDs were processed using zero

filling in the F1 domain, and a sine-bell window function in

both dimensions was applied prior to Fourier transformation.

In the gs-HMQC experiments, GARP modulation of 13C was

used for decoupling. 15N NMR spectra were acquired using

2D inverse proton detected heteronuclear shift correlation

spectroscopy. Typical parameters for the gs-HMQC

(1H–15N) spectra were: a spectral width of 5787 Hz for 1H

and 12.5 kHz for 15N, a 1024 � 256 data set, number of

scans = 4, a relaxation delay of 1 s and a 7 ms delay for the

evolution of the 15N–1H coupling. The FIDs were processed

using zero filling in the F1 domain, and a sine-bell window

function in both dimensions was applied prior to Fourier

transformation. A Bruker BVT 3000 temperature unit was

used to control the temperature of the cooling gas stream and

an exchanger was used to achieve low temperatures.

Solid state. Solid state 13C (100.73MHz) and 15N (40.60MHz)

CPMAS NMR spectra were recorded on a Bruker WB 400

spectrometer at 300 K using a 4 mm DVT probe head.

Samples were carefully packed in 4 mm diameter cylindrical

zirconia rotors with Kel-F caps. Operating conditions in-

volved 90 3.2 ms 1H pulses and a decoupling field strength of

78.1 kHz in a TPPM sequence. The non-quaternary suppres-

sion (NQS) technique to observe only the quaternary carbon

atoms was employed. 13C spectra were initially referenced

to a glycine sample and then the chemical shifts were recalcu-

lated to Me4Si (for the carbonyl atom, dglycine = 176.1).

Similarly, 15N spectra were initially referenced to 15NH4Cl

and then recalculated to nitromethane, using the relationship:

d 15Nnitromethane = d 15NNH4Cl� 338.1. Typical acquisition

parameters for the 13C CPMAS were: a spectral width of

40 kHz, a recycle delay of 15–75 s, an acquisition time of

30 ms, a contact time of 2 ms and a spin rate of 12 kHz.

Typical parameters for the 15N CPMAS were: a spectral width

of 40 kHz, a recycle delay of 15–75 s, an acquisition time of 35

ms, a contact time of 7–8 ms and spin rate of 6 kHz.

Acknowledgements

This work has been financed by the Spanish MEC (CTQ2006-

02586 and CTQ2007-62113).

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Neutral 5-nitrotetrazoles: easy initiation with low pollutionw

Thomas M. Klapotke,* Carles Miro Sabate and Jorg Stierstorfer

Received (in Durham. UK) 21st July 2008, Accepted 8th September 2008


DOI: 10.1039/b812529e

5-Nitro-2H-tetrazole (1), 1-methyl-5-nitrotetrazole (2) and 2-methyl-5-nitrotetrazole (3) were

synthesized starting from the corresponding 5-amino-substituted tetrazoles in good yields and

purities. The compounds were fully characterized by analytical and spectroscopic methods and

their solid state structures were determined by low temperature X-ray diffraction techniques. Due

to the potential of tetrazoles as energetic materials an extensive computational study (CBS-4M)

was performed in order to estimate the energies of formation (DfU1) of the molecules, which are

highly endothermic (1, 2527 kJ kg�1; 2, 2253 kJ kg�1 and 3, 2006 kJ kg�1). The EXPLO5

software was used to calculated the corresponding detonation velocities (Ddet) and detonation

pressures (pdet) (1, Ddet = 9457 m s�1 and pdet = 390 kbar; 2, Ddet = 8085 m s�1 and

pdet = 257 kbar and 3, Ddet = 8109 m s�1 and pdet = 262 kbar) by combining the DfU1 values

of the materials with the (X-ray calculated) densities and molecular formulas, giving performances

comparable to commonly used secondary explosives (e.g., RDX). Lastly, all three neutral

compounds can be easily initiated by impact (o2 J) and with high detonation velocities and

excellent combined oxygen and nitrogen contents offer a more powerful and environmentally

friendly alternative to commonly used primary explosives in initiating devices.

Introduction

In the continuous search for novel green energetic materials1

with high nitrogen but low carbon content,2,3 several groups

around the world are currently investigating HEDMs (High

Energy Dense Materials) based on tetrazoles.4 These energetic

materials have variable application such as in low-smoke

producing pyrotechnic compositions,5 gas generators,6 pro-

pellants,7 high explosives8 and primers in primer charges

(PC).9 Tetrazole derivatives,10,11 tetrazolate12,13 and tetra-

zolium14,15 salts are of special interest. One of the most

promising class of molecules in this regard are 5-substituted

tetrazoles16 (Fig. 1) due to the fact that their properties can be

controlled by selection of the substituent at the carbon atom.

While electron donating groups (EDGs) such as NH217 or

OH18 yield rather stable compounds, electron withdrawing

groups (EWGs) such as NO219 and CN20 destabilize the ring

system and increase the sensitivity of the materials. Also

protonation/alkylation of the tetrazole ring is directed by the

electronegativity of the substituent. While EWGs direct the

protonation/alkylation to the nitrogen atom labelled as N2 of

the tetrazole ring (see crystal structure labels),21 EDGs favor

substitution at N1.22 However, there are other factors that

contribute to the explosivity of tetrazoles. For example, the

high sensitivity of 5-azidotetrazole (C)23 and 5-nitriminotetra-

zole (D)24,25 is better explained due to the energetic nature of

the azide and nitramine groups rather than based on the

electronic influence of these groups on the ring system.

The combination of a tetrazole ring with energetic groups

containing oxygen such as nitro groups (R–NO2),26 nitrate

esters (R–O–NO2)27 or nitramines (R2N–NO2)

28 is of parti-

cular interest. Energetic materials based on tetrazoles show the

desirable compromise in properties with high nitrogen con-

tents on the one hand, and surprising kinetic and thermal

stabilities due to aromaticity on the other.

The interesting energetic properties of tetrazole-based

energetic materials have been mainly investigated in view of

the properties of such compounds for use as propellants and/

or secondary explosives1,29 and it has only been until recent

times that metal salts with 5-substituted tetrazole ligands have

been studied as prospective primary explosives.30,31 Primary

explosives are characterized by easy initiation when submitted

Fig. 1 Structural formulas of neutral 5-substituted tetrazoles.

Prof. Dr. Thomas M. Klapotke, Energetic Materials Research,Department of Chemistry and Biochemistry, University of Munich(LMU), Butenandtstr. 5-13, D-81377 Munich, Germany.E-mail: [email protected]; Fax: +49 89 2180 77492w CCDC reference numbers 689202 (1), 689201 (2) and 689203 (3). Forcrystallographic data in CIF or other electronic format see DOI:10.1039/b812529e



to heat or shock and have the ability to transmit the detona-

tion to less sensitive (secondary) explosives. For these reasons,

they are used in initiating devices. Typical detonation

velocities for this class of compounds are in the range

3500–5500 m s�1, much lower than those of secondary

explosives (5500–9000 m s�1).32 Commonly used primary

explosives are based on lead (e.g., lead azide or styphnate)

with the accompanying environmental impact and thus, recent

efforts have focused on the development of more environmen-

tally friendly metal-based alternatives.30,31

In this work we would like to present the syntheses, full

analytical, spectroscopic and structural characterization of

neutral 5-nitrotetrazoles. In addition, the energetic properties

of the compounds were assessed revealing easy initiation by

impact and detonation velocities, which are almost twice as

high as those of commonly used primary explosives.5


Syntheses

5-Nitro-2H-tetrazole (1) was prepared starting from 5-amino-

1H-tetrazole (5-At) by a modified literature procedure accord-

ing to Scheme 1.33 5-At was diazotized according to a

previously published procedure in our group34 to yield ammo-

nium 5-nitrotetrazolate. ‘‘In situ’’ formation of the potassium

salt by reaction with potassium hydroxide in ethanol and

subsequent treatment with diluted hydrochloric acid and

extraction with ether yields the desired compound. The solvent

needs to be removed using vacuum since the product (1)

absorbs water on time and the compound needs to be stored

under nitrogen. On the other hand, methylation of sodium

5-aminotetrazolate using dimethyl sulfate22 yields a separable

mixture of the two (1-methyl and 2-methyl) isomers.35

Both compounds can be treated similarly and diazotization

with two equivalents of sodium nitrite in the presence of a

non-nucleophilic acid (e.g., sulfuric acid) yields 1-methyl-5-

nitrotetrazole (2) and 2-methyl-5-nitrotetrazole (3) as crystal-

line compounds. Similar reactions are also found in the

literature by using N2O5.36

2 and 3 are extracted from the

reaction mixture using CH2Cl2. The selection of the acid for

the diazotation process is of utmost importance since it affects

the yield of the nitro-compound. For example using hydro-

chloric acid 1- or 2-methyl-5-chlorotetrazoles are obtained as

the main product.

Lastly, 1 is readily soluble in most common solvents such as

ether, THF, MeCN, acetone, water, DMSO and DMF

whereas 2 and 3 show also good solubility in MeOH, EtOH,

acetone, MeCN, ethyl acetate, THF, CH2Cl2 and DMSO

and DMF.

NMR spectroscopy

All three neutral 5-nitrotetrazoles were characterized by ana-

lytical and spectroscopic methods. The elemental analysis of 1

was omitted due to the risk of explosion found in similar

compounds with a high sensitivity30 on the one hand and to

the hygroscopicity of the material on the other. The 1H NMR

spectra of 1–3 measured in DMSO-d6 show two signals

corresponding to the ring proton (1, broad, d = 6.29 ppm)

and to the methyl group protons (2, d = 3.68 ppm and 3,

d = 4.50 ppm). The electron withdrawing character of the

–NO2 group shifts the proton resonances to low field in

comparison to 5-amino-1H-tetrazole and 1-methyl- and

2-methyl-5-amino-1H-tetrazole (i.e., –NH2 group).37 Table 1

contains summarized the 13C and 15N NMR chemical shifts

and the 15N–1H coupling constants for all three compounds.

The proton coupled as well as the proton decoupled 15N NMR

spectra (with full NOE) were also recorded. As already

observed in the 1H NMR spectra, the methyl group resonance

(in this case carbon resonance) of the 1-methyl isomer (2, d =

33.1 ppm) is shifted to higher field in comparison with that of

the 2-methyl isomer (3, d = 41.9 ppm). Similarly, the ring

carbon atom signal is also to be found at higher field for

2 (d = 157.6 ppm) than for 3 (d = 166.4 ppm). This carbon

atom shows the lowest field resonance for 1 (d = 168.4 ppm),

which is in keeping with salts containing the 5-nitrotetrazolate

anion.30,34

Scheme 1 Syntheses of neutral 5-nitrotetrazoles: a = ref. 34;

b = (i) KOH, (ii) HCl (2 M); (c) (i) NaOH, (ii) Me2SO4; d = 2 eq.

NaNO2, H2SO4.

Table 1 15N and 13C NMR resonances of compounds 1–3 with protonation (1, PIS) and methylation (2 and 3, MIS) induced shiftsa and couplingconstantsb

Compound N1 N2 N3 N4 N5 C1 C2

1 �69.6 [�3.4] 19.6 [5.2] 19.6 [5.2] �69.6 [�3.4] �29.8 [�4.5] 168.4 —2 �155.7 [�89.5] �0.7 [�15.1] 6.7 [7.7] �54.8 [11.4] �37.6 [�12.3] 157.6 33.1

2J(N–H) = 2.1 3J(N�H) = 1.83 �97.9 [�31.7] �76.6 [�91.0] 5.3 [9.1] �55.1 [11.1] �33.5 [�8.2] 166.4 41.9

3J(N–H) = 1.7 2J(N–H) = 2.1 3J(N�H) = 1.7NaNTc �66.2 14.4 14.4 �66.2 �25.3 169.2 —

a PIS and MIS values are shown in square [] brackets and given in ppm. b Coupling constants (J) are given in Hz. c NaNT = Sodium

5-nitrotetrazolate (see ref. 30).


The quadrupolar moment of the 14N nucleus results in

signals at approximately +14 (N2/3), �24 (NO2) and �66(N1/4) in the 14N NMR spectrum of compound 1, which are

broad (n1/2 B 300 Hz, B60 Hz and B320 Hz, respectively).

The 15N NMR spectra show comparable resonances to those

observed in the 14N NMR spectra but much sharper (Fig. 2).

The proton induced shifts (PIS, 1) and methyl induced shifts

(MIS, 2 and 3)37 are useful for identifying the protonation/

alkylation site as well as assigning the resonances of the

nitrogen atoms and are also tabulated in Table 1. Comparison

of the resonances observed for the 5-nitrotetrazolate anion in

sodium 5-nitrotetrazolate with those of the compounds in this

study shows unexpected shifts. The nitrogen atoms labelled as

N2 and N3, which are equivalent due to fast exchange in the

NMR in solution of 1, show the largest (positive) PIS value,

indicative of protonation taking place on these two nitrogen

atoms as observed (in the solid state) in the crystal structure of

the compound (see X-ray discussion). The remainder of the

PIS values are small in value and negative. The MIS effect in

compounds 2 and 3 is much more unexpected and the methyl-

ated nitrogen atoms (N1 for 2 and N2 for 3) show the largest

(negative) MIS values (B�90 ppm for both). The next nitrogen

atom close to the methyl group (i.e., at two bonds) feels the

effect of the alkyl group much more weakly (MIS = �31.7 ppmfor N1 in 3) but can still be used to assign the resonances of

this atom. Lastly, the fast exchange of the proton in 1 results in

broadening of the resonance corresponding to the protonated

nitrogen atom, whereas 2 and 3 show coupling constants to the

methyl groups, which are slightly larger for the nitrogen atoms

directly attached to the methyl group (2J(N�H) = 2.1 Hz)

than for those at three bonds of the methyl group protons

(3J(N�H) = 1.7–1.8 Hz).

Molecular structures

Suitable single crystals of 1 and 2 were picked from the

crystallization mixture and mounted in Kel-F oil and trans-

ferred to the N2 stream of an Oxford Xcalibur3 diffractometer

with a Spellman generator (voltage 50 kV, current 40 mA) and

a KappaCCD detector. The data collections were performed

using the CrysAlis CCD software,38 the data reduction with

the CrysAlis RED software.39 The data for compound 3 were

collected on a Nonius Kappa CCD diffractometer under an N2

stream as well. Data collection and reduction was done by the

Bruker ‘‘Collect’’ and the ‘‘HKL Denzo and Scalepack’’

software.40 The structures were solved with SIR-92 (2, 3),41

and SHELXS-97 (1),42 refined with SHELXL-9743 and finally

checked using the PLATON software.44 The non-hydrogen

atoms were refined anisotropically and the hydrogen atoms

were located and freely refined. The absorptions of 1 and 2

were corrected by a SCALE3 ABSPACK multi-scan

method.45 All relevant data and parameters of the X-ray

measurements and refinements are given in Table 2.46

One of the two crystallographically independent formula

units found in the crystal structure of compound 1 is repre-

sented in Fig. 3. Protonation of the NT� anion occurs at N3,

Fig. 2 15N NMR spectra of compounds 1–3.

Table 2 Crystallographic data and refinements

1 2 3

Formula CHN5O2 C2H3N5O2 C2H3N5O2

Mr/g mol�1 115.07 129.09 129.09Crystal system Monoclinic Monoclinic MonoclinicSpace group P21 (No. 4) P21/n (No. 14) P21/c (No. 14)Color/habit Colorless plates Colorless rods Colorless rodsSize/mm 0.03 � 0.18 �

0.270.04 � 0.08 � 0.10 0.03 � 0.18 �

0.27a/A 5.3358(4) 10.0578(4) 6.331(1)b/A 9.4799(7) 9.7055(4) 4.993(1)c/A 8.3190(8) 16.5331(6) 16.388(3)b/1 106.989(9) 101.701(4) 97.13(3)V/A3 402.44(6) 1580.36(11) 514.0(2)Z 4 12 4Dc/g cm�3 1.899 1.628 1.668m/mm�1 0.174 0.143 0.146F(000) 232 792 264lMoKa/A 0.71073 0.71073 0.71073T/K 123 200 200y Range/1 4.0, 32.5 3.9, 25.0 3.2, 26.0Data set �7: 7; �14: 14;

�12:12�11: 10; �11: 11;�19: 15

�7: 7; �6: 6;�20: 19

Reflectionscollected

5975 7250 3454

Independentreflections

1452 2766 1003

Rint 0.059 0.064 0.037Observedreflections

742 1139 862

No. parameters 153 258 94R1 (obs) 0.0274 0.0664 0.0364wR2 (all data) 0.0497 0.2102 0.0946GooF 0.83 0.91 1.08Dr/e A�3 �0.23, 0.21 �0.27, 0.85 �0.22, 0.14CCDC 689202 689201 689203

Fig. 3 Formula unit of 1 with the labelling scheme. Hydrogen atoms

shown as spheres of arbitrary radius and thermal displacements set at

50% probability.


which is in contrast with 5-amino-46 or 5-azidotetrazole.23 In

Table 3 there are summarized the angles and distances of the

two 5-nitrotetrazolate rings, which, within the limits of error,

are not significantly different. The main difference is observed

in the twist of the nitro groups in respect to the tetrazole rings.

Whereas in one of the two molecules the nitro group is almost

coplanar to the ring (O4–N10–C1–N6 = �178.2(3)1), in the

other, this is significantly deviated (O2–N5–C1–N1 =

�164.9(3)1). So, one of the formula units is similar to metal

NT salts, which show small torsion angles between 2 and 51,30

whereas the other one is more similar to NT salts with

nitrogen-rich bases (0–101).34 A plausible explanation for this

could be that in NT salts, there is a negative charge, which is

delocalized all around the tetrazole ring and over the nitro

group, making them virtually coplanar. Proof for this is the

relatively longer C1–N1 distances (B1.445(4) A) in 1. There-

fore one would expect larger torsion angles due to the lack of

delocalization in both formula units. The smaller torsion angle

for one of the two units can be explained by hydrogen-bonding

effects (see discussion below).

Regardless of the expected planarity of 1 the presence of a

proton surrounded by many electronegative atoms forces the

compound to form hydrogen bonds, which prohibits layering.

These hydrogen bonds are formed by the protonated nitrogen

atom (N2 or N7) as the donor atom and either tetrazole ring

nitrogen atoms or, in one instance with one of the nitro groups

oxygen atoms (O3). A report of hydrogen-bridges is given in

Table 4. The interaction between N7 and O3 (N7� � �O3i =

3.015(4) A; symmetry code: (i) �x, 0.5 + y, �z) ‘‘fixes’’ thenitro group in such a way that it is coplanar to the tetrazole

ring and forms the C1,1(6) motifs represented in Fig. 4

(at the primary level) (Table 5). Similarly, the dimer

pairs formed by two crystallographically related rings

Table 3 Selected bond lengths [A] and angles [1] for compounds 1–3

1 (A) 1 (B) 2 (A) 2 (B) 2 (C) 3

O1–N5 1.229(3) 1.227(3) 1.208(6) 1.194(6) 1.196(5) 1.223(2)O2–N5 1.221(3) 1.215(3) 1.191(6) 1.202(5) 1.185(5) 1.222(2)N5–C1 1.440(4) 1.450(4) 1.450(7) 1.481(7) 1.438(6) 1.445(2)N1–C1 1.319(4) 1.310(4) 1.326(6) 1.316(6) 1.342(6) 1.321(2)N1–N2 1.323(4) 1.319(4) 1.336(6) 1.345(5) 1.354(6) 1.319(2)N2–N3 1.318(3) 1.322(3) 1.297(6) 1.317(6) 1.325(6) 1.329(2)N3–N4 1.324(4) 1.318(3) 1.356(7) 1.363(6) 1.317(7) 1.317(2)N4–C1 1.333(4) 1.336(4) 1.303(6) 1.291(6) 1.295(6) 1.331(2)N1(2)–C2 1.461(6) 1.478(7) 1.487(6) 1.461(2)

O2–N5–O1 125.5(3) 126.0(3) 127.6(6) 127.2(5) 125.6(5) 125.1(1)O1–N5–C1 117.9(3) 117.2(3) 115.8(5) 117.2(5) 117.4(5) 117.4(1)O2–N5–C1 116.6(3) 116.7(3) 116.6(5) 115.7(5) 117.0(5) 117.5(1)N1–C1–N5 123.2(3) 122.3(3) 125.9(5) 123.2(5) 124.8(4) 121.9(1)C1–N1–N2 100.1(3) 99.0(3) 106.8(4) 106.5(4) 107.3(4) 99.9(1)N3–N2–N1 114.7(3) 115.4(3) 106.6(5) 106.2(4) 104.9(4) 114.3(1)N2–N3–N4 105.7(2) 105.5(2) 111.4(5) 110.8(4) 111.5(5) 106.1(1)N3–N4–C1 105.0(2) 104.3(3) 103.6(5) 103.5(4) 106.6(5) 104.6(1)N4–C1–N5 122.1(3) 121.9(3) 122.5(5) 123.5(5) 125.4(5) 123.0(1)N1–C1–N4 114.5(3) 115.8(3) 111.6(5) 113.1(5) 109.7(5) 115.1(1)C1–N1–C2 106.8(4) 131.9(5) 131.4(5)N1–N2–C2 123.6(1)N2–N1–C2 121.1(5) 121.5(5) 121.3(5)N3–N2–C2 122.1(1)

Table 4 Geometry for selected hydrogen bonds in the structure of 1

D–H� � �A D–H (A) H� � �A (A) D� � �A (A) D–H� � �A (1)

1

N2–H1� � �N4 0.97(4) 1.88(4) 2.837(4) 171.(3)N7–H2� � �O3a 0.85(4) 2.22(4) 3.015(4) 158.(3)N7–H2� � �N3b 0.85(4) 2.53(4) 3.057(4) 121.(3)

a Symmetry codes for 1: �x, 0.5 + y, �z. b 1 � x, 0.5 + y, 1 � z.

Fig. 4 View of the unit cell of 1 along the a-axis showing the graph-

sets in the structure (dotted lines). Symmetry codes: (ii) 1 � x, 0.5 + y,

1 � z; (iii) 1 + x, y, 1 + z; (iv) 1 � x, �0.5 + y, 1 � z; (v) 2 � x, 0.5 +

y, 1 � z; (vi) 1 � x, 0.5 + y, �z; (vii) x, 1 + y, z; (viii) 1 + x, y, z.


(N7� � �N3ii = 3.057(4) A; symmetry code: (ii) 1 � x, 0.5 + y,

1 � z) yield a C1,1(4) graph-set. Lastly, the third hydrogen

bond found in the structure forms only finite patterns of the

type D1,1(2) at the primary level, which combine with the

other two hydrogen bonds yielding larger dimeric interactions

with the label D3,3(X) (X = 7, 9) at the secondary level. This

results in a highly efficient packing as can be deduced from the

high density of the compound (1.899 g cm�3).

The unit cell of 2, which crystallizes in the monoclinic space

group P21/c contains twelve molecules. For better clearness

only one molecule of the asymmetric unit is shown in Fig. 5.

Since hydrogen bonds are not present in the structures

of 2 and 3, the densities (2: 1.628, 3: 1.668 g cm�3) are

significantly lower than that observed for 1 (1.899 g cm�3).

3, (in Fig. 6), crystallizes monoclinic in the space group P21/c

with four molecules in the unit cell. The molecular geometry

of 2 as well as of 3 is particularly comparable to that of

1 and other 5-substituted tetrazoles. All C–N and N–N bond

lengths lie between single and double bonds, whereby the

shortest distance (1.30–1.33 A) is observed between the

nitrogen atoms N2 and N3. In both cases the NO2

group is co-planar with the tetrazole ring, which confirms

previously published assumptions.47 The distances between

the atoms C1 and N5 are between 1.43 and 1.48 A, which

are in the range of typical C–N single bonds. The same trend

can be found for the N1–C2 and N2–C2 bond lengths

(1.46–1.49 A).

Thermal and energetic properties

In order to assess the thermal and energetic properties of

neutral 5-nitrotetrazoles 1–3 the thermal stability (decomposi-

tion points from DSC measurements), as well as the sensitivi-

ties to friction, impact, electrostatic discharge and thermal

shock of all three compounds were experimentally assessed

(Tables 6 and 7) using standard BAM tests.50–55 In addition,

for all three CHNO compounds the constant volume energies

Table 5 Graph-set matrix for medium to strong hydrogen bonds inthe crystal structure of 1. First level motifs on-diagonal and secondlevel graph sets off-diagonal

H-Bond N2–H1� � �N4 N7–H2� � �O3a N7–H2� � �N3b

N2–H1� � �N4 D1,1(2)N7–H2� � �O3a D3,3(9) C1,1(6)N7–H2� � �N3b D3,3(7) C1,1(4)

a Symmetry codes for 1: �x, 0.5 + y, �z. b 1 � x, 0.5 + y, 1 � z.



30% probability.



50% probability.

Table 6 Physico-chemical properties, initial safety data and predictedperformance of compounds 1–3

1 2 3

Formula CHN5O2 C2H3N5O2 C2H3N5O2

Molecular mass/g mol�1

115.05 129.08 129.08

Impact sensitivitya/J o 1 2 1Friction sensitivityb/N o 5 82 40Electrical dischargec/J — 0.50 0.20N (%)d 60.9 54.3 54.3N + O (%)e 88.6 79.0 79.0O (%)f �7.0 �43.4 �43.4Thermal shockg

Deflagration Combustion CombustionCombustion Very good Good GoodSmokeless + + +DSCh/1C 98 (mp), 130

(decomp.)45 (mp), 155(decomp.)

75 (mp), 150(decomp.)

Densityi/g cm�3 1.899 1.628 1.668DfHm1

j/kJ mol�1 281 278 247DfU1

k/kJ kg�1 +2527 +2253 +2006

Calculated values using EXPLO5

�DEUm1l/J g�1 �5744 �5588 �5368

TEm/K 4804 4226 4071

pn/kbar 390 257 262Do/m s�1 9457 8085 8109Gas vol.p/L kg�1 779 766 763

a BAMmethods, see ref. 50–55. b BAMmethods, see ref. 50–55. c OZM

electric spark tester, see ref. 57–59. d Nitrogen content. e Nitrogen +

oxygen content. f Oxygen balance.60 g Fast heating behavior. h Decom-

position temperature from DSC (b = 5 1C). i Estimated from X-ray

diffraction. j Calculated molar enthalpy of formation. k Energy of

formation. l Energy of explosion, EXPLO5 V5.02. m Explosion tempera-

ture. n Detonation pressure. o Detonation velocity. p Assuming only

gaseous products.


of combustion were calculated using quantum chemical

methods (see Computational Methods section). Initially, we

measured experimentally the combustion data for 3 using

oxygen bomb calorimetry, however the high sensitivity of

the compound did not allow reproducible values to be ob-

tained. The material explodes rather than burning in the

aerobic conditions of the measurements leading to erroneous

values. The heats and energies of formation of 1–3 were back-

calculated from the combustion data and subsequently used in

conjunction with the molecular formula and density (from

X-ray) to predict the performance (detonation pressure and

velocity) for each compound using the EXPLO5 computer

code.61

Fig. 7 shows typical DSC thermographs of compounds

1–3. Slow heating in a DSC apparatus (b = 5 1C min�1) of

samples of B1.5 mg of each energetic material gives rapid

decomposition at temperatures above 130 1C for all three

compounds. All three materials show highly exothermic

decompositions following to endothermic peaks at 98 (1), 45 (2)

and 75 (3) 1C corresponding to the melting of the compounds.

The difference in area between endothermic and exothermic

peaks gives a feeling for the energy released upon decomposi-

tion. For all 1–3 the decomposition releases much more energy

than that required for melting. Particularly, compound 1

shows only a small melting endotherm followed by highly

energetic decomposition. It is interesting to note the effect

of the substituent in the tetrazole ring in the melting and

decomposition points. The presence of the ring proton in 1

results in a higher melting point than for 2 and 3 due to the

possibility of forming classical hydrogen bonds in 1, however,

the compound is much more sensitive (i.e., less stable) and

decomposes at lower temperatures. The substitution pattern of

the methyl group in 2 and 3 also accounts for the lower

melting point of 2 in comparison to 3 due to the formation

of a less effective packing, as suggested by the lower crystal

density of 2 (1.628 g cm�3) in comparison to 3 (1.668 g cm�3).

Furthermore, the presence of the methyl group results in

an increase (20–25 K) of the decomposition temperatures

(B150 1C) as observed for methylated 5-aminotetrazoles.37

Further studies on the decomposition of 3 can be found in

literature.48 In addition to DSC analysis, the response to thermal

shock of 1–3was tested by placing a small sample (B0.5–1.0 mg)

of compound in the flame. This resulted in an vigorous reaction

(deflagration) in the case of 1 and normal burning in the case of

2 and 3, in all cases smokeless. By comparison with typical

primary explosives such as lead azide or styphnate, which both

explode in the flame, the compounds studied here are less

sensitive to thermal shock and show a similar response to

classical secondary explosives such as TNT or RDX.

Data collected for initial safety testing of compounds 1–3

are summarized in Table 6. The impact and friction sensiti-

vities as well as the electrostatic sensitivity were determined.49

The impact sensitivity tests were carried out according to

STANAG 448950 modified according to instruction51 using a

BAM (Bundesanstalt fur Materialforschung)52 drophammer.53

The friction sensitivity tests were carried out according to

STANAG 448754 modified according to instruction55 using

the BAM friction tester. Compound 1 is very sensitive towards

impact (o1 J) and extremely friction sensitive (o5 N). 2 and 3

are also very sensitive towards impact (2, 2 J and 3, 1 J) but

less sensitive towards friction (2, 82 N and 3, 40 N). Grinding

of the compounds in a mortar results in rattling and (in some

instances) a loud explosion. According to the ‘‘UN Recom-

mendations on the transport of dangerous goods’’,56

compounds 1–3 are classified as ‘‘very sensitive’’ regarding

the impact sensitivity values. The compounds in this study

are significantly more sensitive to friction and impact than

nitrogen-rich salts of 5-nitro-2H-tetrazole34 and the impact

sensitivity approaches that of alkali metal salts of 5-nitro-2H-

tetrazole.30 Comparison of the energetic compounds of the

materials in this study with those of commonly used high

explosives are useful to assess the potential of the materials

described here. All three materials have impact sensitivity

values, which are comparable to lead azide (2.5–4.0 J, pure

product). As for the friction sensitivity, 1 has a value between

that of the primary explosives lead azide (0.1–1.0 N, pure

product) and tetrazene (7 N), whereas 2 and 3 have similar

sensitivity to the secondary explosive PETN (60 N).5 In

addition, the sensitivity towards electrostatic discharge of 2

and 3 was tested using an electric spark tester ESD 2010EN

(OZM Research) operating with the ‘‘Winspark 1.15 software

package’’.57 Due to the hygroscopicity of compound 1, this

compound was omitted from this study. The electrical spark

sensitivities of microcrystalline materials (5–100 mm)58 were

determined to be 0.50 � 0.05 (2) and 0.20 (3) � 0.04 J. These

values can be compared to those of commonly used secondary

Table 7 Comparison of energetic properties of compounds 1–3 withRDX

1 2 3 RDX

Density/g cm�3 r 1.899 1.630 1.670 1.820Oxygen balance (%) O �7.0 �43.4 �43.4 �21.6Energy of formation/kJ kg�1 DfU +2527 +2253 +2006 +67Heat of detonation/kJ kg�1 Qv �5744 �5588 �5368 �5902Detonation temperature/K Tex. 4804 4226 4071 3986Detonation pressure/kbar P 390 257 262 299Detonation velocity/m s�1 D 9457 8085 8109 8796Volume of detonation gases/L kg�1 V0 779 766 763 932

Fig. 7 DSC thermographs of 5-nitrotetrazoles 1–3 at a heating rate

b = 5 1C min�1.


explosives, e.g. RDX (0.15 J), PETN (0.19) and TNT (0.57).

However, the ESD sensitivities determined are higher

than those of modern insensitive explosives such as TATB

(1,3,5-triamino-2,4,6-trinitrobenzene).59 Lastly, the high

sensitivities of 1–3 can be attributed not only to the high

endothermicity of the materials (see discussion below) but also

to the only slightly negative oxygen balance, in particular in

the case of compound 1.

In addition to safety considerations, performance of

HEDMs is of utmost importance. Using the molecular for-

mula, density (from X-ray) and energy of formation, the

EXPLO5 computer code61 can be used to calculate the deto-

nation velocity and pressure of CHNO-based explosive

materials. The program is based on the chemical equilibrium,

steady-state model of detonation. It uses the Becker–

Kistiakowsky–Wilson’s equation of state (BKW EOS) for

gaseous detonation products and Cowan–Fickett’s equation

of state for solid carbon.61,62 The calculation of the equili-

brium composition of the detonation products is done by

applying modified White, Johnson and Dantzig’s free energy

minimization technique. The program is designed to enable

the calculation of detonation parameters at the CJ point. The

BKW equation in the following form was used with the

BKWN set of parameters (a, b, k, y) as stated below the

equations and Xi being the mol fraction of ith gaseous

product, ki is the molar covolume of the ith gaseous pro-

duct.40,41 The results of the EXPLO5 calculations for neutral

5-nitrotetrazoles 1–3 are presented in Tables 6 and 7, with the

corresponding values for commonly used RDX for compari-

son purposes.

pV/RT = 1 + xebxx = (kP

Xiki)/[V (T + y)]a; a = 0.5,

b = 0.176, k = 14.71, y = 6620.

Further physico-chemical properties of all three compounds

are tabulated in Table 7. Compounds 1–3 have high nitrogen

contents in the range between B50 and 60%, excellent com-

bined oxygen and nitrogen balances in the range between

B80 and 90% and slightly negative oxygen balances approxi-

mately in the range between that of dinitroglycol (C2H4N2O6,

O = �0.0%) and that of nitromethane (CH3NO2, O =

�39.3%). As expected, the densities, calculated from the

X-ray measurements, are lower in the case of the methylated

derivatives 2 and 3 (B1.65 g cm�3) but still comparable to

TNT (1.654 g cm�3), while 1 has an exceptionally high density

of 1.899 g cm�3 comparable to b-HMX (1.900 g cm�3).5 The

energies of formation of 1–3 were back-calculated from the

energies of combustion on the basis of their combustion

equations (see below), Hess’s Law, the known standard heats

of formation for water and carbon dioxide and a correction

for change in gas volume during combustion. No corrections

for the non-ideal formation of nitric acid (typically B5% of

the nitrogen content reacts to form HNO3) were made.

As pointed out above, all three compounds are highly en-

dothermic with energies of formation above +2000 kJ kg�1.

1 has the most positive value of all at +2527 kJ kg�1 and all

three compounds show similar endothermicities to nitrogen-

rich salts with the 5,50-azotetrazolate anion ([N4C–N =

N–CN4]2�).12,63

1: CHN5O2 (s) + 0.25 O2 (g) - CO2 (g) + 0.5 H2O (l)

+ 2.5 N2 (g) (1)

2, 3: C2H3N5O2 (s) + 1.75 O2 (g) - 2 CO2 (g) + 1.5

H2O (l) + 2.5 N2 (g) (2)

The methylated derivatives 2 and 3 have calculated detonation

velocities of B8100 m s�1, higher than TNT (6900 m s�1),

lower than RDX (8800 m s�1) and similar to 5,50-azotetrazole

salts,12,63 regardless of the high sensitivity of the compounds.

On the other hand, 1 although being very sensitive to impact

and friction and thus classifying as a primary explosive has an

astonishingly high calculated detonation velocity of 9457 m s�1,

which is comparable to some of the highest performing

secondary explosives known to date such as HMX (octogen,

9100 m s�1), CL-20 and octanitrocubane (B10 000 m s�1)5

and also higher than the primary explosive 5-azido-1H-tetra-

zole regardless of the lower endothermicity of the –NO2 group

in comparison to the –N3 substituent.21 Here it is necessary to

mention that the previous study of Koldobskii and coworkers

on compound 1, reports a experimental density value of

1.73 g cm�3,33 which is much lower than our calculated value

of 1.899 g cm�3 and therefore affects strongly the detonation

parameters. The detonation pressures have accordingly high

values (390 kbar for 1 and B260 kbar for 2 and 3), which are

comparable to HMX (octogen) (pdet. = 384 kbar) and RDX

(hexogen) (pdet. = 299 kbar) respectively.5

Decomposition gases

Using the calculated heats of formation, the calculated density

(from X-ray) and the molecular formula the ICT code64 was

used to predict the heats of explosion as well as the decom-

position gases formed upon explosion/decomposition of com-

pounds 1–3. Table 8 contains tabulated results of these

calculations together with the predicted values for two com-

monly used high explosives, namely lead azide (primary

Table 8 Predicted decomposition gases and heats of explosion of compounds 1–3 and comparison with commonly used high explosives (using theICT code)

Compoundab CO2 H2O N2 CO H2 NH3 CH4 HCN C Pb DHexc

1 276.61 75.54 607.47 17.11 0.05 1.34 — — 21.40 — 16212 100.72 185.39 533.13 17.25 0.47 11.16 0.80 0.44 150.35 — 15583 100.45 186.28 533.22 16.21 0.43 11.06 0.66 0.43 150.97 — 1512Pb(N3)2 — — 288.56 — — — — — — 711.44 391RDXd 292.09 232.40 373.83 22.90 0.21 5.26 0.16 0.30 72.28 — 1592

a The amount of gases formed at 298 K is given in g kg�1 (i.e., grams of gas per kilogram of energetic compound). b —, the decomposition product

was not predicted by the code. c Heat of explosion in cal g�1. d Measured at a density of 1.76 g cm�3.


explosive) and RDX (secondary explosive), for comparative

purposes.

As expected from the high nitrogen content of the materials

molecular nitrogen is predicted to be the major product of the

decomposition of compounds 1–3 (B500–600 g kg�1). By

comparison, the decomposition of RDX is expected to pro-

duce much lower amounts of environmentally-friendly nitro-

gen (374 g kg�1), which is however still the main expected

product followed by the formation of carbon dioxide (292 g

kg�1). In keeping with the logic that RDX derives its energy

from both oxidation of the carbon backbone and the forma-

tion of nitrogen. The second main product predicted upon

decomposition of 1 is also carbon dioxide (277 g kg�1), which

is anticipated to form in larger amounts than for compounds 2

and 3 (B100 g kg�1) fitting with the better oxygen balance of

1. 2 and 3 are nevertheless expected to generate larger amounts

of water than 1 (B185 vs. 75 g kg�1). Apart from carbon soot,

which is predicted to form in relatively large amounts for 2 and

3 (B150 g kg�1), the rest of the decomposition gases (CO, H2,

NH3, CH4 and HCN) are foreseen to form in marginally low

quantities (o20 g kg�1). The amount of highly toxic gases

(i.e., CO and HCN) expected from the explosion of 1–3 are

then comparable to those formed upon explosion of RDX and

in contrast with the large amounts of highly toxic lead powder

predicted for the explosion of lead azide (711 g kg�1). Lastly,

the heats of explosion have all values above 1500 cal g�1, is

larger for the more energetic compound 1, are comparable to

the secondary explosive RDX and much larger than the

primary explosive lead azide (391 cal g�1).

Computational methods. Due to the high sensitivity of all

compounds studied here, bomb calorimetric measurements

could only be performed with small amounts of the materials

and doubtful combustion data was obtained. Therefore we

decided to estimate the thermodynamic data by quantum

chemical methods. All calculations were carried out using

the Gaussian G03W (revision B.03) program package.65 The

enthalpies (H) and free energies (G) were calculated using the

complete basis set (CBS) method described by Petersson and

coworkers in order to obtain very accurate values. The CBS

models use the known asymptotic convergence of pair natural

orbital expressions to extrapolate from calculations using a

finite basis set to the estimated complete basis set limit. CBS-4

begins with a HF/3-21G(d) geometry optimization; the zero

point energy is computed at the same level. It then uses a large

basis set SCF calculation as a base energy, and a MP2/

6-31+G calculation with a CBS extrapolation to correct

the energy through second order. A MP4(SDQ)/6-31+(d,p)

calculation is used to approximate higher order contributions.

In this study we applied the modified CBS-4M method

(M referring to the use of minimal population localization)

which is a re-parametrized version of the original CBS-4

method and also includes some additional empirical correc-

tions.66,67 The enthalpies of the gas-phase species M were

computed according to the atomization energy method

(eqn (3)) (Tables 9–12).68

DfH1(g, M, 298) = H(Molecule, 298) �P

H1(Atoms, 298)

+P

DfH1(Atoms, 298) (3)

From the gas-phase enthalpies of formation DfH1(g) the

enthalpies of the solid state were calculated using the enthal-

pies of sublimation by the equation:

DfH1(s) = DfH1(g) � (DsubH) (4)

For a solid compound the enthalpy of sublimation (DsubH) can

be approximated on the basis of TROUTON’s rule 72 if the

melting temperature (Tm in K) is known:

DsubH [J mol�1] = 188 Tm [K] (5)

With the known enthalpies of formation of carbon

dioxide (DfH1298(CO2(g)) = �393.8 kJ mol�1) and water

(DfH1298(H2O(g)) = �241.9 kJ mol�1) the enthalpies of

formation of 1–3 can now be calculated. From these values,

the energy of formation (DfU1298) can easily be obtained from

the combustion eqns (1)–(3) according to eqn (6) with Dnbeing the change of moles of the gaseous components

(Dn: 1 = �3.25; 2, 3 = 2.75) in eqns (7) and (8).

DUm = DHm � DnRT (6)

1: C (s) + 0.5 H2 (g) + 2.5 N2 (g) + O2 (g)

- CHN5O2 (s) (7)

2, 3: 2 C (s) + 1.5 H2 (g) + 2.5 N2 (g) + O2 (g)

- C2H3N5O2 (s) (8)

Table 9 CBS-4M results

Pt. Gp. �H298/a.u. �G298/a.u. NIMAG

5-Nitro-1H-tetrazole Cs 462.190261 462.227578 05-Nitro-2H-tetrazole Cs 462.195276 462.232789 01-Methyl-5-nitrotetrazole Cs 501.427898 501.468919 02-Methyl-5-nitrotetrazole Cs 501.437871 501.479892 0H 0.500991 0.514005C 37.786156 37.803062N 54.522462 54.539858O 74.991202 75.008515

Table 10 Literature values for atomic DH1f298/kcal mol�1

ref. 69 NIST70

H 52.6 52.1C 170.2 171.3N 113.5 113.0O 60.0 59.6

Table 11 Enthalpies of the gas-phase species M

M M DfH1(g)/kcal mol�1

5-Nitro-1H-tetrazole CHN5O2 +87.15-Nitro-2H-tetrazole CHN5O2 +84.01-Methyl-5-nitrotetrazole C2H3N5O2 +80.72-Methyl-5-nitrotetrazole C2H3N5O2 +74.5

Table 12 Enthalpies of sublimation of compounds 1–371

Tm/K DHsub/kcal mol�1

1 374 16.82 318 14.33 348 15.6


As can be seen from Table 13 compounds 1–3 are formed

strongly endothermically (1: 281, 2: 278, 3: 247 kJ mol�1).

These values are slightly higher than that of 5-amino-1H-

tetrazole (DfH1(s) = 208 kJ mol�1) and in the same range

observed for 5-nitriminotetrazole (264 kJ mol�1) and

1-methyl-5-nitriminotetrazole (260 kJ mol�1). The enthalpy

of formation of energetic materials are governed by the

molecular structure of the compound. Therefore, heterocycles

with a higher nitrogen content (e.g. imidazole (DfH1(s) = 58.6

kJ mol�1),72 1,2,4-triazole (DfH1(s) = 109.3 kJ mol�1),73 1H-

1,2,3,4-tetrazole (DfH1(s) = 237.4 kJ mol�1))74 show trends in

increasing heats of formation.

Experimental

CAUTION! The 5-nitrotetrazoles described here are energetic

compounds, which are sensitive towards heat, impact, friction

and electrostatic discharge. Although we experienced no diffi-

culties in the synthesis of these materials, proper protective

measures (safety glasses, face shield, leather coat, earthened

equipment and shoes, Kevlars gloves and ear plugs) should be

used when undertaking work involving 1–3 on small and in

particular on larger scales.

General method

All reagents and solvents were used as received (Sigma-

Aldrich, Fluka, Acros Organics) unless stated otherwise.

Melting points were measured with a Linseis PT10 DSC75

and checked with a Buchi Melting Point B-450 apparatus

(uncorrected). DSC measurements were performed at a heat-

ing rate of 5 1C min�1 in closed aluminum sample pans with a

1 mm hole in the top for gas release under a nitrogen flow of

20 mL min�1 with an empty identical aluminum sample pan as

a reference. NMR spectra were recorded with a Jeol Eclipse

270, Jeol EX 400 or a Jeol Eclipse 400 instrument. All chemical

shifts are quoted in ppm relative to TMS (1H, 13C) and

MeNO2 (14N, 15N). Infrared (IR) spectra were recorded using

a Perkin-Elmer Spektrum One FT-IR instrument.76 Transmit-

tance values are qualitatively described as ‘‘very strong’’ (vs),

‘‘strong’’ (s), ‘‘medium’’ (m) and ‘‘weak’’ (w). Raman spectra

were measured using a Perkin-Elmer Spektrum 2000R NIR

FT-Raman instrument equipped with a Nd:YAG laser

(1064 nm). The intensities are reported as percentages of the

most intense peak and are given in parentheses. Elemental

analyses were performed with a Netsch Simultaneous Thermal

Analyzer STA 429.

Synthesis of 5-nitro-2H-tetrazole (1)

Anhydrous ammonium 5-nitrotetrazolate (0.44 g, 3.32 mmol)

and potassium hydroxide (0.19 g, 3.32 mmol) were dissolved in

3.7 mL water. The solution was stirred at reflux until no more

ammonia gas was evolved. At this point, the reaction mixture

was cooled by means of an ice-bath and cold B25% sulfuric

acid (2 mL) was added dropwise using a plastic syringe. The

solution was then extracted with ether (4 � 6 mL) and the

ether extracts were combined and washed to remove the excess

of acid with water (6 mL). The organic phase was then dried

with magnesium sulfate and filtered and the solvent was

stripped under high vacuum (B10�3 mbar) yielding the pure

product as a slightly yellow semicrystalline solid (0.27 g, 72%),

which was carefully (!!) scratched out using a plastic spatula

and analyzed. DSC (5 1C min�1, 1C): 98 (mp), 4130 (de-

comp.); IR ~n/cm�1 (KBr, rel. int.): 3443 (s), 2013 (w), 1629

(m), 1565 (vs), 1443 (m), 1401 (m), 1320 (s), 1262 (vw), 1192

(w), 1103 (w), 1047 (m), 1022 (m), 840 (s), 666 (w), 534 (vw);

Raman ~n/cm�1 (rel. int.): 3316 (2), 3261 (2), 1572 (14), 1492

(9), 1446 (100), 1433 (90), 1396 (15), 1358 (7), 1317 (13), 1200

(13), 1186 (14), 1142 (30), 1094 (42), 1069 (13), 1043 (11), 1027

(26), 837 (22), 775 (13), 736 (4), 592 (3), 532 (11), 444 (22), 256

(17), 240 (16), 154 (6); 1H NMR (DMSO-d6, 400.18 MHz, 25

1C, TMS) d/ppm: 6.29 (1H, NH); 13C{1H} NMR (DMSO-d6,

100.63 MHz, 25 1C, TMS) d/ppm: 168.4 (1C, C–NO2);14N

NMR (DMSO-d6, 40.55 MHz, 25 1C, MeNO2) d/ppm: +14

(2 N, n1/2 B300 Hz, N2/3), �24 (1 N, n1/2 B60 Hz, NO2), �66(2 N, n1/2 B320 Hz, N1/4); 15N NMR (DMSO-d6, 40.55 MHz,

25 1C, MeNO2) d/ppm: +19.6 (2 N, s, N2/N3), �29.8 (1 N, s,

NO2),�69.6 (2 N, s, N1/4);m/z (FAB�, xenon, 6 keV,m-NBA

matrix): 113.9 (100, CN5O2�); EA (CHN5O2, 115.07): calc. C

10.44, H 0.88, N 60.87; found: not determinable due to high

sensitivity; BAM drophammer: o1 J, friction tester: o5 N,

flame: deflagration.

Synthesis of 1-methyl-5-nitrotetrazole (2)

A suspension of 1-methyl-5-aminotetrazole (2.00 g, 20 mmol)

in 1 M sulfuric acid (10 mL) and 30 mL of water was added at

0 1C to 30 mL water containing sodium nitrite (2.76 g,

40 mmol). After stirring at room temperature for 12 h and

filtration of the precipitated bis(1-methyltetrazolyl)triazene,

the solvent was evaporated. Dry acetone (80 mL) was added

to this and the precipitated Na2SO4 was removed by filtration.

After evaporating the acetone the crude product was recrys-

tallized from a small amount of ethanol (1.52 g, yield 59%);

DSC (5 1C min�1, 1C): 45 1C (mp), 155 1C (decomp.); IR

(KBr, cm�1): ~n = 3038 (w), 2860 (w), 1550 (vs), 1481 (s), 1467

(m), 1408 (s), 1364 (s), 1328 (vs), 1280 (w), 1209 (m), 1073 (m),

1025 (w), 846 (s), 720 (s), 535 (w), 430 (m); Raman (1064 nm,

200 mW, 25 1C, cm�1): ~n = 3054 )2), 2978 (6), 2964 (6), 1530

(20), 1508 (9), 1469 (100), 1463 (44), 1447 (44), 1425 (15), 1412

(11), 1329 (24), 1261 (10), 1208 (13), 1103 (11), 1085 (23), 1028

(21), 923 (7), 779 (2), 740 (2), 709 (3), 679 (4); 1H NMR ([d6]-

DMSO, 25 1C, ppm) d: 3.68 (s, 3H, CH3);13C NMR ([d6]-

DMSO, 25 1C, ppm) d: 157.6 (CN4), 33.1 (CH3);14N NMR

(DMSO-d6, 40.55 MHz, 25 1C, MeNO2) d/ppm: �37 (1 N,

n1/2 B60 Hz, NO2),15N NMR ([d6]-DMSO, 25 1C) d = 4. 5

(N3), �14.1 (N6), �18.29 (N2, t, 3JNH = 1.9 Hz), �71.15(N4), �157.16 (N5), �168.38 (N1, d, 2JNH = 2.2 Hz), –289.13

(N7, 1JNH = 102.7 Hz), �329.66 (N8, 1JNH = 69.4 Hz); m/z

(DEI+): 130 (19) [M +H]+, 129 (65) [M]+, 100 (1), 83 (8), 55

Table 13 Solid state enthalpies (DfH1) and energies of formation(DfU1)

DfH1(s)/kcalmol�1

DfH1(s)/kJmol�1 Dn

DfU1(s)/kcalmol�1

M/gmol�1

DfU1(s)/kJkg�1

1 +67.2 +281.4 �4 +69.5 115.07 +2527.12 +66.5 +278.4 �5 +69.5 129.08 +2252.83 +58.9 +246.6 �5 +61.9 129.08 +2006.4


(15), 54 (17), 53 (100), 46 (38), 43 (28), 40 (7), 39 (5), 28 (45), 18

(8), 15 (4); EA (C2H3N5O2, 129.08): calcd.: C 18.61, H 2.34, N

54.26; found: C 18.39, H 2.28, N 52.80; BAM drophammer:

2 J; friction tester: 82 N, ESD: 0.50� 0.05 J, flame: combustion.

Synthesis of 2-methyl-5-nitrotetrazole (3)

To 20 mL of an aqueous sodium nitrite (2.76 g, 0.04 mol)

solution, a solution of 2-methyl-5-aminotetrazole (2.00 g, 0.02 mol)

in 20 mL 1N sulfuric acid was added at 0 1C. The reaction

mixture was stirred for 8 h and the precipitated bis(2-methyl-

tetrazolyl)triazene precipitated was removed by filtration.

Afterwards the product was extracted three times with 20 mL

of CH2Cl2. The organic phases were combined, dried over

MgSO4 and evaporated. The crude product was recrystallized

from acetone yielding single crystals suitable for XRD ana-

lysis. (1.68 g, yield 65%); DSC (5 1C min�1, 1C): 75 1C (mp),

150 1C (decomp.); IR (KBr, cm�1): ~n = 3022 (m), 1610 (m),

1565 (s), 1510 (m), 1468 (m), 1412 (s), 1285 (s), 1160 (m), 1001

(w), 880 (m), 788 (m), 750 (m), 670 (w), 610 (w), 530 (w);

Raman (1064 nm, 200 mW, 25 1C, cm�1): ~n= 3052 (12), 2967

(60), 1555 (28), 1486 (40), 1468 (26), 1418 (100), 1369 (11),

1335 (8), 1322 (14), 1287 (12), 1209 (30), 1075 (12), 1043 (40),

1026 (44), 841 (16), 776 (17), 715 (46), 547 (10), 436 (30), 378

(18), 307 (16), 218 (16); 1H NMR ([d6]-DMSO, 25 1C, ppm) d:4.50 (s, 3H, CH3);

13C NMR ([d6]-DMSO, 25 1C, ppm) d:166.4 (CN4), 41.9 (CH3);

14N NMR (DMSO-d6, 40.55 MHz,

25 1C, MeNO2) d/ppm: �34 (1 N, n1/2 B50 Hz, NO2);15N

NMR ([d6]-DMSO, 25 1C) d = 5.3 (N3), �33.5 (N5), �55.1(N4), �76.6 (N2, 2JNH = 2.1 Hz), –97.9 (N1, 3JNH = 1.7 Hz);

m/z (DEI+): 130 (2) [M+H]+, 129 (2) [M]+, 115 (1) [M+H

� CH 3]+, 101 (15), 58 (89), 43 (100) [HN3]

+, 42 (6), 28 (3)

[N2]+, 18 (29); EA (C2H3N5O2, 129.08): calcd.: C 18.61,

H 2.34, N 54.26; found: C 18.88, H 2.35, N 52.99; BAM

drophammer: 1 J; friction tester: 40 N, ESD: 0.20 � 0.04 J,

flame: combustion.

Conclusions

From this combined experimental and theoretical study the

following conclusions can be drawn:

Convenient procedures for the synthesis of three highly

energetic neutral 5-nitrotetrazoles (1–3) are presented, which

allow the materials to be obtained at low cost using facile

routes, good yields and excellent purities. The full characteri-

zation of the compounds by analytical and spectroscopic

methods is described in detail. In addition we determined the

molecular structure of the compounds in the solid state by

X-ray diffraction methods. The energetic properties of the

compounds were assessed by means of standard tests and

quantum chemical calculations (CBS-4M). 1–3 are calculated

to be strongly endothermic with heats of formation between

247 and 282 kJ mol�1. The compounds show highly exother-

mic decomposition peaks (DSC) and are easily initiated by

impact. Furthermore, 1–3 have high performances (EXPLO5

code) comparable to commonly used secondary explosives

regardless of their ease of initiation and 1 classifies as a

primary explosive in regard to its impact and friction sensitivity

values and still has a calculated detonation velocity, which is

almost twice as large as that of common primary explosives

(e.g., lead azide) and puts it first in the list of high performing

primers. Unfortunately, 1 absorbs water, which limits its

application. Furthermore 1–3 are thermally too unstable for

use as conventional high explosives. The predicted products

(ICT code) formed upon explosion of 1–3 are expected to be

less harmful than those expected from the decomposition of

commonly used high explosives, which suggest their potential

(2 and 3) as environmentally friendly alternatives with high

performance and low initiation barriers (by impact) for use in

energetic applications (e.g., as ingredients for energetic fillers

in high explosive compositions).

Acknowledgements

Financial support of this work by the Ludwig-Maximilian

University of Munich (LMU), the Fonds der Chemischen

Industrie (FCI), the European Research Office (ERO) of the

U.S. Army Research Laboratory (ARL) under contract nos.

N-62558-05-C-0027, 9939-AN-01 and W911NF-07-1-0569

and the Bundeswehr Research Institute for Materials, Explo-

sives, Fuels and Lubricants (WIWEB) under contract nos.

E/E210/4D004/X5143 and E/E210/7D002/4F088 is gratefully

acknowledged. The authors are indebted to and thank

Dr Betsy Rice and Dr Gary Chen for many helpful discussions

and support of our work. We also acknowledge Mr. Stefan

Huber for help with the sensitivity tests. The authors acknowl-

edge collaborations with Dr M. Krupka (OZM Research,

Czech Republic) in the development of new testing and

evaluation methods for energetic materials and with

Dr M. Sucesca (Brodarski Institute, Croatia) in the develop-

ment of new computational codes to predict the detonation

parameters of high-nitrogen explosives.

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Probing multivalency for the inhibition of an enzyme: glycogen

phosphorylase as a case studyw

Samy Cecioni,a Oana-Andreea Argintaru,a Tibor Docsa,b Pal Gergely,b

Jean-Pierre Pralyaand Sebastien Vidal*

a

Received (in Montpellier, France) 22nd July 2008, Accepted 21st August 2008

First published as an Advance Article on the web 23rd October 2008

DOI: 10.1039/b812540f

Glycogen phosphorylase is involved in the hepatic glucose production and appears an emerging

biological target for the treatment of type 2 diabetes. Two distinct trivalent inhibitors of GP were

synthesized either through Cu(I)-assisted 1,3-dipolar cycloaddition or through formation of a tris-

oxadiazole derivative. A biological study of the inhibiting properties of these trivalent inhibitors

of GP have shown that the valency of the molecules influences slightly the inhibition of the

enzyme whereas the presence of a spacer arm between the core and the pharmacophore moieties

does not. The possible modes of binding of these multivalent inhibitors to the enzyme are

discussed.

Introduction

Glycogen phosphorylase and diabetes

Diabetus mellitus affects about 3% of the world population

and up to 6% for the adult population of developed countries.

Diabetes, one of the major causes of death worldwide, is

characterized by elevated glycaemia causing heart and kidney

failures as well as visual impairment problems.1 Type 2

diabetes is non-insulino-dependent and arises from insulin

signalling inefficiency resulting in insufficient or even late

insulin secretion. A series of biological targets have been

identified for anti-diabetic therapy2 such as peroxisome

proliferator-activated receptors a/g (PPARs a/g),3 glucagon-

like peptide-1 (GLP-1),4 dipeptidyl peptidase IV (DPP-IV)5 or

protein tyrosine phosphatase 1B (PTP 1B).6 Glycogen phos-

phorylase (GP) has recently appeared as an enzyme of interest

for the treatment of type 2 diabetes.7 This enzyme catalyses

glycogen depolymerisation to release glucose-1-phosphate

according to the schematic equation:

(Glucose)n - (Glucose)n�1 + Glucose-1-phosphate

The inhibition of GP is expected to slow down glycogenolysis

and to lower the production of glucose from the liver therefore

allowing for a better control over hyperglycaemia. GP has

been extensively studied and crystallographic data analyses are

displaying a large number of sites for the inhibition of this

enzyme.8 A series of GP inhibitors have been described pre-

viously with various heterocyclic structures9 but our work10 is

focusing on carbohydrate-based inhibitors of GP7b,c,f which

are capable of binding selectively at the catalytic site of GP.

These glycomimetic approaches are based on structural modi-

fications at the molecular level for improving the binding to

the enzyme and therefore affording valuable GP inhibitors.

Multivalency and inhibition of enzymes

Another approach for the inhibition of GP could take advan-

tage of multivalency. This strategy may provide additional

opportunities in the field of drug discovery for the design of

potent enzyme inhibitors particularly by reaching higher

affinities and probably better selectivities. A few examples of

multivalent inhibition of an enzyme are reported in the

literature where multimeric species of a specific drug are

capable of improving the inhibition in comparison to the

monomeric molecule.11 The binding of one ligand subunit

from the multivalent molecule to the enzyme generates an

increase in local concentration of ligands, thus creating an

apparent cooperativity causing an enhancement in inhibition.

Dimeric inhibitors of influenza virus neuraminidase have been

developed by MacDonald et al.12 and displayed up to a

100-fold increase in inhibition along with improved pharmaco-

kinetic properties. In an additional study of the same

group,13 a set of trimeric and tetrameric inhibitors displayed

improved antiviral activities and long-lasting protective activ-

ities against influenza virus. More recently, the group of

J. Gervay-Hague has described the synthesis of a series of

trivalent zanamivir derivatives via click chemistry although no

biological activity has been reported yet.14 Inhibition of

acetylcholinesterase by dimeric molecules resulted in up to

3000-fold increases in potency and selectivity compared to

the monomeric inhibitor.15 Glycosidases inhibition16 with

tethered dimeric azasugars was investigated and the molecules

displayed interesting inhibitions of these enzymes but more

aUniversite de Lyon, Lyon, Universite Lyon 1, Villeurbanne, CNRS,UMR5246, Institut de Chimie et Biochimie Moleculaires etSupramoleculaires, Laboratoire de Chimie Organique2 - Glycochimie, 43 Boulevard du 11 Novembre 1918, F-69622Villeurbanne, France. E-mail: [email protected];Fax: +33 472 448 349; Tel: +33 472 448 349

bCell Biology and Signalling Research Group of the HungarianAcademy of Sciences, Department of Medical Chemistry, ResearchCentre for Molecular Medicine, University of Debrecen, Nagyerdeikrt. 98, Debrecen, H-4032, Hungary. E-mail: [email protected];Fax: +36 52 412 566; Tel: +36 52 412 345

w Electronic supplementary information (ESI) available: Determina-tion of Ki values and detailed atom numbering of molecules. See DOI:10.1039/b812540f



importantly a selectivity for two enzymes out of the seven

tested.16a However, tetravalent 1-azafagomine inhibitors dis-

played no improved inhibition compared to the monovalent

structure but rather a strong decrease in activity.16b Finally,

bivalent inhibitors of tetrameric b-tryptase constructed on a

cyclodextrin scaffold have shown increased inhibitions for this

enzyme.17

Inspired by the above mentioned results for the multivalent

neuraminidase inhibition and by the fact that a multimeric

enzyme offers additional possibilities for improved inhibition

through multivalency, we have designed synthetic routes to

trivalent carbohydrate-based GP inhibitors. A single case of a

dimeric inhibitor of GP has been reported with a bis(5-

chloroindole-2-carboxamide) derivative inhibiting human liver

GPa (HLGPa) with an IC50 value of 6 nM compared to 12.5 mMfor the parent monovalent inhibitor (2000-fold increase).18

This result highlights a productive and cooperative binding

of both ends of the bivalent inhibitor to two binding sites. The

co-crystallization of HLGPa with the divalent inhibitor

demonstrated that a single molecule of inhibitor was capable

of interacting with each monomeric unit of GP by linking the

two binding sites across the interface of GP homodimeric

structure. This result encouraged us to further investigate this

approach for the design of multivalent inhibitors of GP.

A closer look at the modes of binding of multivalent

inhibitors to an enzyme reveals several possibilities as depicted

in Fig. 1. A dimeric enzyme such as GP can interact with two

monomeric molecules of inhibitor in a 1:1 complex (2:2 at the

molecular level) providing a reference IC50mono value for

monovalent inhibitors. When considering the same inhibitor

repeated three times on a molecular scaffold, four main

possible cases can then be envisaged. The inhibition of a

trivalent inhibitor must be divided by three in order to

consider the contribution of each residue in its comparison

to a monovalent inhibitor. If a 1:1 complex is formed in

solution, a simple statistical effect can be invoked if no positive

effect is observed and the IC50 value observed will be similar to

1/3 IC50mono (Fig. 1, Case 1). Nevertheless, the IC50 measured

can be improved with a lower value in comparison to 1/3

IC50mo (Fig. 1, Case 2). If each binding site of the enzyme is

occupied by a ligand of the multivalent inhibitor, the forma-

tion of 3:1 complexes would afford aggregates of proteins. If

the size of the aggregates remains small enough to maintain a

good solubility of the complex, the IC50 observed will be

similar to 1/3 IC50mono (Fig. 1, Case 3). Nevertheless, if the

size of the multivalent inhibitor-enzyme clusters becomes large

enough to cause their precipitation, the quantity of enzyme

present in the solution will diminish and therefore the IC50

value measured will be lower than 1/3 IC50mono (Fig. 1, Case

4). In this case, the IC50 value measured will be a ‘‘virtual’’

value because of the lower quantity of the enzyme available in

solution.


Synthesis of trivalent inhibitors

We have recently reported the preparation of 3-C-glycosyl-

5-aryl-1,2,4-oxadiazoles which displayed good inhibition

towards GP.10b In this context, we synthesized an alkyne-

terminated 3-C-glycosylated 1,2,4-oxadiazole which could then

be involved in a 1,3-dipolar cycloaddition with a tris-azido-

functionalized derivative to obtain a multivalent GP inhibitor

candidate. A more condensed trimeric inhibitor was also

prepared by direct coupling of an amidoxime with a tris-acyl

chloride and subsequent dehydrative cyclization to the corres-

ponding trivalent C-glycosylated oxadiazole. These inhibitors

were designed in order to determine the influence of a spacer

arm between the core and the pharmocophore ligands on the

inhibition of the enzyme.

Synthesis of the trivalent inhibitor with a spacer arm

The perbenzoylated glucosyl cyanide 119 was reacted with

hydroxylamine hydrochloride in pyridine to afford the desired

amidoxime 2 (Scheme 1). In our previous work,10b the crude

product obtained was rather difficult to purify by silica gel

column chromatography. The amidoxime 2 could be obtained

pure without chromatography simply by diluting the crude

product in ethyl acetate and then washing the organic layer

with 1 M aqueous HCl to remove pyridine and excess of

hydroxylamine, followed by saturated aqueous NaHCO3 and

brine. This simple chromatography-free purification process

afforded the expected amidoxime 2 in 99% yield and high

purity. The formation of the O-acyl-amidoxime 3 was

achieved with 4-pentynoic acid in the presence of EDCI/HOBt

as coupling agents.

We previously observed that reactions times lasting from a

few hours to a few days were required for the thermal

cyclodehydration of O-acyl-amidoximes. In order to optimize

both time and yield, we performed this reaction under TBAF

catalysis20 and/or microwaves activation21 (Table 1). We

observed that the use of TBAF catalysis at room temperature

provided the cyclic oxadiazole 4 within 1 day (entry 1) while

thermal activation combined with TBAF catalysis drastically

shortened the reaction time to 10 minutes (entry 2). NoFig. 1 Possible modes of binding for a mono- and trivalent inhibitors

with a dimeric enzyme.


reaction or decomposition of the starting material was ob-

served when applying microwaves activation without TBAF

catalysis (entries 3 and 4). Nevertheless, the association of

TBAF catalysis and microwaves activation performed on a

short timescale (entry 5 and 6) provided a result comparable to

that observed under conventional heating (entry 2). These

results underline the beneficial influence of TBAF catalysis.

The alkyne-terminated oxadiazole 4 was then engaged in a

Huisgen’s Cu(I)-catalyzed 1,3-dipolar cycloaddition22 reaction

under microwaves activation with benzyl azide to afford the

desired 1,4-disubstituted 1,2,3-triazole 6 in excellent yield.

Debenzoylation of compounds 4 and 6 afforded two hydro-

xylated GP inhibitor candidates 5 and 7. Similarly, the reac-

tion of 1,3,5-tris(azidomethyl)benzene23 with the alkyne

derivative 4 under microwaves activation and Cu(I) catalysis

afforded the cycloadduct 8. The saponification of the benzoate

esters provided the fully hydroxylated macromolecule 9.

Synthesis of the trivalent inhibitor without spacer arm

We next prepared a more condensed trifunctional macro-

molecule were the C-glucosyl-oxadiazole moiety was directly

attached to a benzene ring (Scheme 2). Condensation of

amidoxime 2 with 1,3,5-benzenetricarbonyl trichloride af-

forded the corresponding triester 10 in 73% yield. At first,

compound 10 was subjected to cyclodehydration under

thermal conditions (reflux in 1,4-dioxane). The product obtained

was not the expected tris-oxadiazole 13 but the bis-oxadiazole

11 with one unreacted O-acyl amidoxime moiety as evidenced

by mass spectrometry (m/z = 2035.4 [M + H]+). Interest-

ingly, molecular ions could be observed neither for compound

13 nor the mono-oxadiazole intermediate. Saponification of

the ester groups of 11 resulted in the concomitant cleavage of

the O-acyl amidoxime function and afforded the benzoic acid

derivative 12 whose structure was clearly demonstrated by

mass spectrometry (m/z= 581 [M �H]�) and NMR analyses.

The triple thermal cyclodehydration of 10 was then performed

under microwaves activation and TBAF catalysis for

40 minutes. The tris-oxadiazole derivative 13 was isolated in

72% yield as the only product of the reaction highlighting

again the positive influence of TBAF catalysis and microwaves

activation for this cyclodehydration process. Deprotection

under Zemplen conditions afforded the expected hydroxylated

trivalent GP inhibitor candidate 14.

Inhibition of glycogen phosphorylase

The inhibition of GP was determined, as previously repor-

ted,10b for the three monovalent C-glycosylated oxadiazoles

(5, 7 and 1510b) and the two trivalent derivatives 9 and 14

(Fig. 2, Table 2). The enzymatic assays were performed at two

concentrations and most molecules displayed poor inhibition

properties at a concentration of 625 mM and moderate to good

inhibition at higher concentration (2.5 mM). In addition,

Ki values could be estimated only for trivalent derivatives 9

and 14 (see ESIw).The alkyne-terminated C-glycosylated oxadiazole derivative

5 displayed no inhibition at 625 mM and poor activity at

Scheme 1 Reagents and conditions: (a) NH2OH�HCl, C5H5N, 50 1C, 5 h, 99%; (b) HCRC(CH2)2CO2H, EDCI, HOBt, CH2Cl2/DMF (9:1),�8 1Cthen r.t., 16 h, 67%; (c) PhMe, TBAF 10 mol%, mW (150 1C, 200 W, 5 min), 97%; (d) NaOMe, MeOH then Amberlite IR-120 (H+ form);

(e) PhCH2N3, CuI, Et3N, mW (110 1C, 150 W, 15 min), 88%; (f) C6H3(CH2N3)3, CuI, Et3N, mW (110 1C, 150 W, 15 min), 98%.

Table 1 Cyclodehydration of O-acyl-amidoxime 3 to the 1,2,4-oxadiazole 4 in toluene

Entry Catalyst T/1C Microwave condition Time Yield (%)

1 10% TBAF 25 None 24 h 992 10% TBAF 110 None 10 min 973 None 150 100 W 1 h No reaction4 None 175 200 W 2 h Decomposition5 10% TBAF 150 200 W 5 min 976 10% TBAF 150 200 W 30 min 66


2.5 mM. The inhibition properties disappeared completely

when a spacer arm was added such as in the structure of 7.

Interestingly, the trivalent analogue 9 of the non-active deri-

vative 7 was now inhibiting GP with values of 30% at 625 mMand 56% at 2.5 mM. The valency of the molecule is therefore

responsible for an increase in inhibition from 0 to 56% when

comparing 7 and 9 at 2.5 mM. The C-glycosylated oxadiazole

derivative 15 bearing a phenyl group on the 5-position of the

oxadiazole ring displayed 10% inhibition at 625 mM.10b The

inhibition was again increased to 35% at 625 mM for trivalent

analogue 14. We anticipated that the distance between the core

and the carbohydrate moiety would influence for the binding

to the enzyme. Nevertheless, this was not the case based on the

inhibition measured for 9 and 14.

The increase of valency from monovalent to trivalent species

is responsible for an increase in inhibition of GP. The inhibi-

tion per residue for trivalent molecules is always similar the

Scheme 2 Reagents and conditions: (a) C6H3(COCl)3, 1,4-dioxane, r.t., 24 h, 73%; (b) 1,4-dioxane, 100 1C, 4 days; (c) NaOMe, MeOH then

Amberlite IR-120 (H+ form); (d) PhMe, TBAF 30 mol%, mW (150 1C, 200 W, 40 min), 72%.

Fig. 2 Structure of monovalent and trivalent GP inhibitors tested.

Table 2 Inhibition of GP observed for monovalent and trivalentinhibitors at two concentrations

Inhibition (%)

Inhibitor Valency At 625 mM At 2.5 mM Ki/mM

5 1 0 22 � 4 n.d.a

7 1 0 0 n.d.a

15 1 10 n.d.a n.d.a

9 3 30 � 5 56 � 5 480 � 45b

14 3 35 � 5 62 � 5 535 � 50b

a n.d. = not determined. b Estimated.


corresponding monovalent analogue in accordance with the

statistical effect model proposed (Fig. 1, Case 1). The 3:1

complex non soluble mode of binding (Fig. 1, Case 4) can be

ruled out for these trivalent inhibitors of GP to the dimeric

enzyme since no precipitate was observed under the concen-

trations of trivalent inhibitors and enzyme used for the

inhibition studies. The inhibition observed for the trivalent

species was never better than 1/3 of the inhibition (in %)

observed for the corresponding monovalent molecules. In

conclusion, two modes of binding are possible with either

1:1 or 3:1 complexes (Fig. 1, Case 1 or Case 3, respectively)

resulting in an observed inhibition close to 1/3 of the inhibition

for the parent monovalent inhibitor.

In the present study, the expected binding sites of these

glucose-based multivalent inhibitors are the catalytic site of

GP homodimer which are separated from each other by a long

distance and pointing into opposite directions.8d The structure

of the trivalent inhibitors tested did not permit such an intra-

molecular interaction with both catalytic sites on the same

GP dimer, but rather an interaction with two independent

GP dimers. In comparison, the bis(5-chloroindole-2-carboxamide)

derivative is binding simultaneously at each indole binding site

near the interface between the monomeric units of the GP

dimer.17 The linker is composed of 12 atoms between two

indole aromatic units which are therefore available for inter-

acting with the binding site of each monomer of GP.

The designed trivalent molecules could also bind to the enzyme

on a different site. The large aromatic appendage present in the

aglycon, composed of oxadiazole, phenyl and triazole rings,

might interact with the surface of the protein through hydro-

phobic interactions. The stability of the inhibitor-protein com-

plex would therefore be lower than complex involving an internal

binding site such as the catalytic site. The observed inhibitions

would therefore be weak as currently observed in the present

study. Nevertheless, we do not possess any experimental data

confirming or denying such a mode of interaction.

Conclusions

In conclusion, we have designed two kinds of multivalent

inhibitors of GP based on the acylation of an amidoxime

intermediate followed by thermal dehydrative cyclization to

the corresponding oxadiazole. The introduction of an alkyne

residue at the 5-position of the oxadiazole ring allowed the

coupling to a trivalent azido-functionalized benzene ring

leading to an extended trivalent inhibitor candidate. The

enzyme inhibition assays revealed poor to moderate inhibitory

effect of these analogues. But, more important was the fact

that multivalent inhibitors were always superior to their

monovalent counterparts. This study provides one of the few

examples of multivalent inhibition for an enzyme, even though

the inhibitions observed remain modest.

Experimental

General methods

Thin-layer chromatography (TLC) was carried out on aluminum

sheets coated with silica gel 60 F254 (Merck). TLC plates

were inspected by UV light (l = 254 nm) and developed by

treatment with a mixture of 10% H2SO4 in EtOH/H2O

(1:1 v/v) followed by heating. Silica gel column chromato-

graphy was performed with Gedurans silica gel Si 60 (40–63 mm)

purchased from Merck (Darmstadt, Germany). Reactions

under microwave activation were performed on a CEM

Discover system. HRMS (LSIMS) mass spectra were recorded

in the positive mode using a Thermo Finnigan Mat 95 XL

spectrometer. MS (ESI) mass spectra were recorded in the

positive mode using a Thermo Finnigan LCQ spectrometer.1H and 13C NMR spectra were recorded at 23 1C using Bruker

Advance DRX300 or DRX500 spectrometers with the residual

solvent as the internal standard. The following abbreviations

are used to explain the observed multiplicities: s, singlet; d,

doublet; dd, doublet of doublet; ddd, doublet of doublet of

doublet; t, triplet; td, triplet of doublet; q, quadruplet; m,

multiplet; br, broad; p, pseudo. Structure elucidation was

deduced from 1D and 2D NMR spectroscopy which allowed,

in most cases, complete signal assignments based on COSY,

HSQC, and HMBC correlations. NMR solvents were

purchased from Euriso-Top (Saint Aubin, France). Atom

numbering of the molecules is presented in the ESI.w

Syntheses

1,3,5-Tris(azidomethyl)benzene. A solution of 1,3,5-tris(bro-

momethyl)benzene (3.07 g, 8.6 mmol) and sodium azide

(3.36 g, 51.6 mmol) in DMF (100 mL) was stirred at 65 1C

for 24 hours. The solution was cooled to room temperature

then poured into water (400 mL). The aqueous layer was

extracted with Et2O (3 � 250 mL). The combined organic

layers were washed with water (2 � 400 mL) and brine

(300 mL). The organic layer was dried (Na2SO4), filtered

and evaporated with extreme care (water-bath at room tem-

perature, reduced pressure and Plexiglas shield) to afford 1,3,5-

tris(azidomethyl)benzene (2.05 g, 98%) as a colorless oil. Rf =

0.83 (PE/EtOAc, 8:2). 1H NMR (300 MHz, CDCl3) d = 4.39

(s, 6H, CH2N3), 7.25 (s, 3H, H-ar).

C-(2,3,4,6-Tetra-O-benzoyl-b-D-glucopyranosyl)-formami-

doxime (2). A solution of 2,3,4,6-tetra-O-benzoyl-b-D-gluco-pyranosyl cyanide 1 (3.00 g, 4.96 mmol) and hydroxylamine

hydrochloride (0.86 g, 12.4 mmol) in pyridine (10 mL) was

stirred at 50 1C for 5 hours. The mixture was diluted with

EtOAc (250 mL) and washed with 100 mL portions of water,

1 M HCl, saturated NaHCO3, water and brine successively.

The organic layer was dried (MgSO4), filtered and evaporated

to obtain the pure amidoxime 2 (3.24 g, 99%) as a white foam.

Rf = 0.48 (PE/EtOAc, 1:1). 1H NMR (300 MHz, CDCl3)

d 4.21 (ddd, 1H, J = 9.7 Hz, J = 5.1 Hz, J = 2.7 Hz, H-5),

4.31 (d, 1H, J = 9.8 Hz, H-1), 4.47 (dd, 1H, J = 5.1 Hz, J =

12.4 Hz, H-6a), 4.62 (dd, 1H, J= 2.7 Hz, J= 12.4 Hz, H-6b),

4.76 (bs, 2H, NH2), 5.69 (t, 1H, J = 9.8 Hz, H-2), 5.73 (t, 1H,

J = 9.8 Hz, H-4), 5.96 (t, 1H, J = 9.8 Hz, H-3), 7.24–7.43

(m, 10H, H-ar), 7.47–7.57 (m, 2H, H-ar), 7.81–8.04 (m, 8H, H-ar).

O-(Pent-40-ynoyl)-3-C-(2,3,4,6-tetra-O-benzoyl-b-D-gluco-pyranosyl)-formamidoxime (3). A solution of 4-pentynoic acid

(27 mg, 0.27 mmol) in CH2Cl2/DMF (4 mL, 9:1) was cooled to

�8 1C before addition of 1-hydroxybenzotriazole (HOBt)


(36.5 mg, 0.27 mmol), 1-(3-dimethylaminopropyl)-3-ethyl-

carbodiimide hydrochloride (EDCI) (52 mg, 0.27 mmol) and

amidoxime 2 (146 mg, 0.23 mmol). The mixture was kept at

�8 1C for 30 minutes then stirred at room temperature for

16 hours. The solvents were evaporated off and the crude

product was purified by flash silica gel column chromato-

graphy (PE then PE/EtOAc, 1:1) to afford theO-acylamidoxime

3 (110 mg, 67%) as a white foam. Rf = 0.53 (PE/EtOAc, 1:1).

[a]D = �2.9 (c = 1.00/CH2Cl2).1H NMR (300 MHz, CDCl3)

d 1.94 (t, 1H, J= 2.5 Hz, H-50), 2.40–2.63 (m, 4H, H-20 H-30),

4.31 (ddd, 1H, J= 2.7 Hz, J= 5.1 Hz, J= 9.8 Hz, H-5), 4.54

(d, 1H, J = 9.8 Hz, H-1), 4.53–4.59 (m, 1H, H-6a), 4.67

(dd, 1H, J = 2.7 Hz, J = 9.5 Hz, H-6b), 5.36 (s, 2H, NH2),

5.76 (t, 1H, J = 9.8 Hz, H-4), 5.80 (t, 1H, J = 9.8 Hz, H-2),

6.01 (t, 1H, J = 9.8 Hz, H-3), 7.25–7.57 (m, 12H, H-ar),

7.83–8.07 (m, 8H, H-ar). 13C NMR (75 MHz, CDCl3) d 14.3

(C-30), 32.1 (C-20), 63.0 (C-6), 69.2 (C-4), 69.3 (C-50), 70.0

(C-2), 73.7 (C-3), 75.6 (C-1), 76.7 (C-5), 82.5 (C-40), 128.40,

128.45, 128.5 (3s, 8C, CH-ar), 128.77, 128.82, 129.4 (3s, 4C,

CIV-ar), 129.8, 129.9, 130.00, 130.04 (4s, 8C, CH-ar), 133.4,

133.5, 133.7 (3s, 4C, CH-ar), 153.7 (NQCR–NH2), 165.3,

165.6, 165.7, 166.3 (4s, 4C, COPh), 169.0 (C-10). ESI-MS

(positive mode) m/z: 719.2 [M + H]+, 741.3 [M + Na]+,

786.9 [M + HCOOH + Na]+, 1436.9 [2M + H]+, 1458.9

[2M + Na]+, 1504.5 [2M + HCOOH + Na]+. HR-ESI-MS

(positive mode) m/z: calcd. for C40H34N2O11 [M + H]+

719.2241, found 719.2244.

5-(But-100-yn-400-yl)-3-C-(20,30,40,60-tetra-O-benzoyl-b-D-gluco-pyranosyl)-1,2,4-oxadiazole (4). In a CEM Discover 5 mL vial

was introduced a solution of O-acylamidoxime 3 (369 mg,

0.5 mmol) and TBAF (50 mL, 50 mmol, 1 M in THF) in toluene

(5 mL). The reaction vial was heated at 150 1C for 5 min upon

microwave irradiation (200 W). The solvent was evaporated

and the residue purified by flash silica gel column chromato-

graphy (PE then PE/EtOAc, 1:1) to afford the oxadiazole 4

(348 mg, 97%) as a white foam. Rf = 0.65 (PE/EtOAc, 7:3).

[a]D =+8.3 (c= 1.12/CH2Cl2).1H NMR (300 MHz, CDCl3)

d 1.84 (t, 1H, J= 2.6 Hz, H-100), 2.65 (td, 2H, J= 2.6 Hz, J=

7.5 Hz, H-300), 3.10 (t, 2H, J = 7.5 Hz, H-400), 4.33 (ddd, 1H,

J = 3.4 Hz, J = 5.1 Hz, J = 9.8 Hz, H-50), 4.53 (dd, 1H, J =

5.2 Hz, J = 12.4 Hz, H-60a), 4.65 (dd, 1H, J = 3.0 Hz, J =

12.4 Hz, H-60b), 5.09 (d, 1H, J = 9.5 Hz, H-10), 5.82 (t, 1H,

J= 9.5 Hz, H-40), 6.00 (m, 2H, H-20 H-30), 7.29–7.58 (m, 12 H,

H-ar), 7.81–8.02 (m, 8H, H-ar). 13C NMR (75 MHz, CDCl3) d16.0 (C-300), 26.2 (C-400), 63.3 (C-60), 69.4 (C-40), 70.2 (C-100),

70.6 (C-20), 72.4 (C-10), 74.1 (C-3 0), 77.0 (C-50), 80.8 (C-200),

128.3 (s, 2C, CH-ar), 128.4 (s, 4C, CH-ar), 128.5 (s, 2C, CH-

ar), 128.7, 128.7, 128.8, 129.5 (4s, 4C, CIV-ar), 129.7, 129.8,

129.8, 129.9 (4s, 8C, CH-ar), 133.2, 133.3, 133.4, 133.5 (4s, 4C,

CH-ar), 164.6, 165.2, 165.8, 166.3 (s, 4C, COPh), 166.2 (C-3),

179.0 (C-5). ESI-MS (positive mode) m/z: 701.1 [M + H]+,

723.2 [M + Na]+, 1400.9 [2M + H]+, 1422.9 [2M + Na]+.

HR-ESI-MS (positive mode) m/z: calcd. for C40H32N2O10Na

[M + Na]+ 723.1955, found 723.1954.

3-C-(b-D-Glucopyranosyl)-5-(but-100-yn-400-yl)-1,2,4-oxadiazole (5).

A solution of benzoylated oxadiazole 4 (261 mg, 0.37 mmol)

and NaOMe (5 mg, 0.09 mmol) in CH2Cl2/MeOH (5 mL, 2:3)

was stirred at room temperature for 4 hours. The solution

was neutralized with a cation exchange resin (Amberlite IR-120,

H+ form) and the resin washed with MeOH (3 � 5 mL).

The filtrate was evaporated off and the residue was dissolved in

MeOH then precipitated with CH2Cl2. The resulting solid was

washed with CH2Cl2 (2 � 5 mL) and dried under vacuum

to afford the hydroxylated oxadiazole 5 (103 mg, 98%) as a

white foam. Rf = 0.26 (CH2Cl2/MeOH, 9:1). [a]D = +9.6

(c = 0.51/H2O).1H NMR (300 MHz, CD3OD) d 2.36 (t, 1H,

J = 2.6 Hz, H-100), 2.73 (td, 2H, J = 2.6 Hz, J = 7.3 Hz, H-300),

3.18 (t, 2H, J= 7.3 Hz, H-400), 3.42–3.53 (m, 3H, H-30 H-40 H-50),

3.65–3.73 (m, 2H, H-20 H-60a), 3.87 (dd, 1H, Jo1.0 Hz, J= 12.0

Hz, H-60b), 4.44 (d, 1H, J = 9.7 Hz, H-10). 13C NMR (75 MHz,

CD3OD) d 16.6 (C-300), 27.1 (C-400), 62.8 (C-60), 71.2 (C-40), 71.3

(C-100), 73.3 (C-20), 74.8 (C-10), 79.2 (C-30), 82.3 (C-200), 82.6 (C-50),

169.2 (C-3), 180.5 (C-5). ESI-MS (positive mode) m/z: 285.0

[M + H]+, 307.1 [M + Na]+, 590.9 [2M + Na]+. HR-ESI-MS

(positive mode) m/z: calcd. for C12H16N2O6Na [M + Na]+

307.0906, found 307.0907.

5-[200-(10 0 0-Benzyl-10 0 0,20 0 0,30 0 0-triazol-40 0 0-yl)ethyl]-3-C-

(20,30,40,60-tetra-O-benzoyl-b-D-glucopyranosyl)-1,2,4-oxa-diazole (6). In a CEM Discover 5 mL vial was introduced a

solution of benzyl azide (110 mg, 0.828 mmol), alkyne 4

(193 mg, 0.276 mmol), copper iodide (26 mg, 0.138 mmol)

and DIPEA (240 mL, 1.38 mmol) in toluene (5 mL). The

solution was sonicated for 1 min then heated at 110 1C for 15

min upon microwave irradiation (150 W). The solvent was

evaporated off and the residue purified by flash silica gel

column chromatography (PE/EtOAc, 1:1) to afford the cyclo-

adduct 6 (202 mg, 88%) as a colorless oil. Rf = 0.24

(PE/EtOAc, 1:1). [a]D = +1.4 (c = 1.02/CH2Cl2).1H NMR

(500 MHz, CDCl3) d 3.22 (bs, 2H, H-200), 3.28 (bs, 2H, H-100),

4.37 (ddd, 1H, J=3.0 Hz, J= 5.3 Hz, J= 9.7 Hz, H-50), 4.56

(dd, 1H, J = 5.3 Hz, J = 12.4 Hz, H-60a), 4.68 (dd, 1H, J =

3.0 Hz, J = 12.4 Hz, H-60b), 5.10 (d, 1H, J = 9.7 Hz, H-10),

5.48 (s, 2H, NCH2Ph), 5.86 (t, 1H, J = 9.7 Hz, H-40), 5.97

(t, 1H, J = 9.7 Hz, H-20), 6.06 (t, 1H, J = 9.7 Hz, H-30),

7.24–7.56 (m, 18H, H-50 0 0 H-ar), 7.80–8.03 (m, 8H, H-ar). 13C

NMR (125 MHz, CDCl3) d 21.8 (C-200), 26.6 (C-100), 53.6

(NCH2Ph), 62.9 (C-60), 68.1 (C-40), 69.8 (C-20), 71.7 (C-10),

73.4 (C-30), 76.5 (C-50), 122.2 (C-50 0 0), 127.8, 127.9, 128.0,

128.4, 128.5, 128.9, 129.7, 129.8 (8s, 25C, CH-ar), 133.4, 133.6,

133.8, 133.9 (4s, 4C, CIV-ar), 134.7 (CIV-ar), 145.3 (C-40 0 0),

164.7, 165.0, 165.6, 165.3 (4s, 4C, COPh), 165.6 (C-3), 180.1

(C-5). ESI-MS (positive mode) m/z: 834.1 [M + H]+, 856.1

[M+Na]+, 1666.4 [2M+H]+. HR-ESI-MS (positive mode)

m/z: calcd. for C47H40N5O10 [M + H]+ 834.2775, found

834.2781.

5-[200-(10 0 0-Benzyl-10 0 0,20 0 0,30 0 0-triazol-40 0 0-yl)-ethyl]-3-C-(b-D-glucopyranosyl)-1,2,4-oxadiazole (7). A solution of benzoy-

lated cycloadduct 6 (128 mg, 0.153 mmol) and NaOMe

(5 mg, 0.09 mmol) in CH2Cl2/MeOH (5.5 mL, 10:1) was

stirred at room temperature for 4 hours. The solution was

neutralized with a cation exchange resin (Amberlite IR-120,

H+ form) and resin washed with MeOH (3 � 5 mL). The

filtrate was evaporated off and the residue purified by

flash silica gel column chromatography (CH2Cl2 then


CH2Cl2/MeOH, 8:2 then EtOAc/MeOH, 8:2) to afford the

hydroxylated cycloadduct 7 (63 mg, 98%) as a white foam.

Rf = 0.68 (EtOAc/MeOH, 4:1). [a]D = +5.5 (c = 0.95/

MeOH). 1H NMR (300 MHz, CD3OD) d 3.18–3.23 (m, 2H,

H-200), 3.24–3.36 (m, 2H, H-100), 3.44–3.52 (m, 3H, H-30 H-40

H-50), 3.68–3.74 (m, 2H, H-20 H-60a), 3.88 (dd, 1H, Jo1.0 Hz,

J = 11.4 Hz, H-60b), 4.44 (d, 1H, J = 9.7 Hz, H-10), 5.55

(s, 2H, NCH2Ph), 7.28–7.38 (m, 5H, H-ar), 7.79 (s, 1H, H-50 0 0).13C NMR (75 MHz, CD3OD) d 23.3 (C-200), 27.2 (C-100), 54.9

(NCH2Ph), 62.8 (C-60), 71.3, 79.2, 82.6 (3s, 3C, C-30 C-40

C-50), 73.4 (C-2 0), 74.9 (C-10), 123.9 (C-50 0 0), 129.1, 129.5,

130.0 (3s, 5C, CH-ar), 136.8 (CIV-ar), 146.9 (C-40 0 0), 169.2

(C-3), 180.9 (C-5). ESI-MS (positive mode) m/z: 418.1 [M +

H]+, 440.1 [M + Na]+, 856.6 [2M + Na]+. HR-ESI-MS

(positive mode) m/z: calcd. for C19H23N5NaO6 [M + Na]+

440.1546, found 440.1549.

1,3,5-Tris-40-200-[30 0 0-C-(20 0 0 0

,30 0 0 0

,40 0 0 0

,60 0 0 0

-tetra-O-benzoyl-b-D-glucopyranosyl)-1 0 0 0,2 0 0 0,4 0 0 0-oxadiazol-5 0 0 0-yl]-ethyl-1 0,20,3 0-

triazol-1 0-ylmethylbenzene (8). In a CEM Discover 5 mL vial

was introduced a solution of 1,3,5-tris(azidomethyl)benzene

(POTENTIALLY EXPLOSIVE, 4.9 mg, 20 mmol), alkyne 4

(63 mg, 90 mmol), copper iodide (1.9 mg, 10 mmol) and DIPEA

(17 mL, 100 mmol) in toluene (6 mL). The solution was

sonicated for 1 min then heated at 110 1C for 15 min upon

microwave irradiation (150 W). The solvent was evaporated

off and the residue purified by flash silica gel column chromato-

graphy (PE/EtOAc, 1:1 then EtOAc) to afford the tris-

cycloadduct 8 (46 mg, 98%) as a colorless oil. Rf = 0.67

(EtOAc). 1H NMR (300 MHz, CDCl3) d 3.10–3.32 (m, 12H,

H-100 H-200), 4.34–4.40 (m, 3H, H-50 0 0 0

), 4.51 (dd, 3H, J = 5.1

Hz, J= 12.4 Hz, H-60 0 0 0

a), 4.66 (dd, 3H, J= 2.7 Hz, J= 12.4

Hz, H-60 0 0 0

b), 5.12 (d, 3H, J = 9.4 Hz, H-10 0 0 0

), 5.31 (s, 6H,

NCH2C6H3), 5.87 (t, 3H, J= 9.4 Hz, H-40 0 0 0

), 5.97 (t, 3H, J=

9.4 Hz, H-20 0 0 0

), 6.05 (t, 3H, J = 9.4 Hz, H-30 0 0 0

), 6.93 (s, 3H,

H-2 H-4 H-6), 7.23–7.50 (m, 36H, H-ar), 7.74–8.00 (m, 24H,

H-ar). 13C NMR (75 MHz, CDCl3) d 22.5 (s, 3C, C-100), 26.6

(s, 3C, C-200), 53.1 (NCH2C6H3), 63.3 (s, 3C, C-60 0 0 0

), 69.3

(s, 3C, C-40 0 0 0

), 70.7 (s, 3C, C-20 0 0 0

), 72.3 (s, 3C, C-10 0 0 0

), 74.1

(s, 3C, C-30 0 0 0

), 77.0 (s, 3C, C-50 0 0 0

), 122.2 (s, 3C, C-50), 127.0

(s, 3C, C-2, C-4, C-6), 128.39, 128.43, 128.5 (3s, 24C, CH-ar),

128.69, 128.71, 128.8, 129.5 (4s, 12C, CIV-ar), 129.7, 129.78,

129.82, 129.9 (4s, 24C, CH-ar), 133.2, 133.4, 133.6 (3s, 12C,

CH-ar), 137.0 (s, 3C, C-1, C-3, C-5), 145.6 (s, 3C, C-40), 164.9,

165.2, 165.8, 166.2 (4s, 12C, COPh), 166.4 (s, 3C, C-30 0 0),

180.0 (s, 3C, C-50 0 0). ESI-MS (positive mode) m/z: 1173.5

[M + 2H]2+.

1,3,5-Tris-40-200-[30 0 0-C-(b-D-glucopyranosyl)-10 0 0,20 0 0,40 0 0-oxa-diazol-50 0 0-yl]ethyl-10,20,30-triazol-10-ylmethylbenzene (9). A

solution of benzoylated tris-cycloadduct 8 (114 mg, 49 mmol)

and NaOMe (5 mg, 92 mmol) in CH2Cl2/MeOH (5.5 mL, 10:1)

was stirred at room temperature for 6 hours. The solution was

neutralized with a cation exchange resin (Amberlite IR-120,

H+ form) and resin washed with MeOH (3 � 5 mL). The

filtrate was evaporated off and the residue was dissolved in

MeOH then precipitated with PE. The resulting solid was

washed with PE (5 � 5 mL), dissolved into pure water and

freeze-dried to afford the hydroxylated tris-cycloadduct 9

(45 mg, 84%) as a white foam. 1H NMR (300 MHz, D2O) d3.02–3.19 (m, 6H, H-100), 3.20–3.30 (m, 6H, H-200), 3.51–3.86

(m, 18H, H-20 0 0 0

H-30 0 0 0

H-40 0 0 0

H-50 0 0 0

H-60 0 0 0

a H-60 0 0 0

b), 4.54

(d, 3H, J = 8.8 Hz, H-10 0 0 0

), 5.39 (s, 6H, NCH2C6H3), 6.98

(s, 3H, H-2 H-4 H-6), 7.73 (s, 3H, H-50). 13C NMR (125 MHz,

D2O) d 22.0 (s, 3C, C-100), 26.3 (s, 3C, C-200), 53.4 (s, 3C,

NCH2C6H3), 61.1 (s, 3C, C-60 0 0 0

), 72.1 (s, 3C, C-20 0 0 0

), 73.2

(s, 3C, C-10 0 0 0

), 69.6, 77.1, 80.6 (3s, 9C, C-30 0 0 0

C-40 0 0 0

C-50 0 0 0

),

124.2 (s, 3C, C-50), 127.3 (s, 3C, C-2 C-4 C-6), 137.2 (s, 3C, C-1

C-3 C-5), 146.5 (s, 3C, C-40), 167.5 (s, 3C, C-30 0 0), 181.0 (s, 3C,

C-50 0 0). ESI-MS (positive mode) m/z: 1118.2 [M + Na]+.

HR-ESI-MS (positive mode) m/z: calcd. for C45H57N15NaO18

[M + Na]+ 1118.3904, found 1118.3918.

N,N0,N00-1,3,5-Tris(benzoyloxy)-C-(20,30,40,60-tetra-O-benzoyl-

b-D-glucopyranosyl)tricarboximidamide (10). A solution of

1,3,5-benzenetricarbonyl trichloride (102 mg, 0.38 mmol) and

amidoxime 2 (794 mg, 1.24 mmol) in 1,4-dioxane (15 mL) was

stirred at room temperature for 24 hours. The solvent was then

evaporated off and the mixture was diluted with EtOAc (150mL).

The organic layer was washed by 100 mL portions of saturated

NaHCO3, water and brine successively. The organic layer was

dried (MgSO4), filtered and evaporated. The crude product was

purified by flash silica gel column chromatography (EtOAc) to

afford the O-acylamidoxime 10 (571 mg, 73%) as a white foam.

Rf = 0.75 (PE/EtOAc, 3:7). [a]D = �48.6 (c = 1/CH2Cl2).1H

NMR (300 MHz, CDCl3) d 4.23–4.33 (m, 3H, H-50), 4.49–4.57

(m, 6H, H-10 H-60a), 4.58–4.65 (m, 3H, H-60b), 5.46 (bs, 6H,

NH2), 5.70 (t, 3H, J = 9.6 Hz, H-20), 5.76 (t, 3H, J = 9.6 Hz,

H-40), 5.98 (t, 3H, J = 9.6 Hz, H-30), 7.22–7.58 (m, 36H, H-ar),

7.80–8.04 (m, 24H, H-ar), 8.52 (s, 3H, H-2 H-4 H-6). 13C NMR

(75 MHz, CDCl3) d 62.9 (C-60), 69.0 (C-40), 70.0 (C-20), 73.5

(C-30), 75.6 (C-10), 76.8 (C-50), 128.3, 128.4, 128.5, 128.7, 129.3,

129.7, 129.8, 129.90, 129.94, 130.2 (10s, 20C, CH-ar), 133.3,

133.4, 133.6, 134.1 (4s, 4C, C-ar), 154.5 (H2NCQNO), 161.5

(NOCQO), 165.2, 165.58, 165.63, 166.1 (4s, 4C, COPh). LSIMS

(positive mode, thioglycerol) m/z: 2071.6 [M + H]+. HR-ESI-

MS (positive mode) m/z: calcd. for C114H91N6O33 [M + H]+

2071.5627, found 2071.5651.

3,5-Bis[30-C-(b-D-glucopyranosyl)-10,20,40-oxadiazol-50-yl]-benzoic acid (12). A solution of tris-O-acylamidoxime 10

(532 mg, 256 mmol) in 1,4-dioxane (12 mL) was stirred at

100 1C for 4 days. The reaction was then cooled to room

temperature and the solvent evaporated off. The crude pro-

duct 11 was used without further purification. A solution of

crude 11 (239 mg) and NaOMe (10 mg) in CH2Cl2/MeOH

(5 mL, 1:1) was stirred at room temperature for 5 hours. The

solvent was evaporated off and the crude mixture was purified

by flash reverse-phase silica gel chromatography (H2O then

H2O/MeOH 7:3) to afford the benzoic acid derivative 12

(95 mg, 58% over two steps) as a white foam. Rf = 0.20

(EtOAc/MeOH 1:1). [a]D = +7.1 (c = 0.41/MeOH). 1H

NMR (300 MHz, CDCl3) d 3.62–3.76 (m, 3H, H-300, H-400,

H-500), 3.81–3.89 (m, 2H, H-200, H-600a), 3.99 (m, 1H, H-600b),

4.75 (m, 1H, H-100), 8.77 (s, 2H, H-ar), 8.87 (s, 1H, H-ar). 13C

NMR (75 MHz, CDCl3) d 61.2 (C-600), 69.7, 77.1, 80.8 (3s, 3C,

C-300, C-400, C-500), 72.1 (C-200), 73.4 (C-100), 124.7 (s, 2C,

CIV-ar), 126.9 (CIV-ar), 129.8 (CH-ar), 133.0 (s, 2C, CH-ar),


139.5 (CO2H), 168.4 (C-30), 175.8 (C-50). ESI-MS (negative

mode) m/z: 581 [M � H]�. HR-ESI-MS (negative mode) m/z:

calcd for C23H25N4O14 [M � H]� 581.1367, found 581.1369.

1,3,5-Tris[30-C-(200,300,400,600-tetra-O-benzoyl-b-D-glucopyra-nosyl)-10,20,40-oxadiazol-50-yl]benzene (13). In a CEMDiscover

5 mL vial was introduced a solution of O-acylamidoxime

10 (454 mg, 0.22 mmol) and TBAF (70 mL, 70 mmol, 1 M in

THF) in toluene (5 mL). The reaction vial was heated

at 150 1C for 40 min upon microwave irradiation (200 W).

The solvent was evaporated and the residue purified by

flash silica gel column chromatography (PE then PE/EtOAc,

1:1) to afford the trivalent oxadiazole 13 (317 mg, 72%) as a

white foam. Rf = 0.79 (PE/EtOAc, 1:1). [a]D = �35.1 (c =

1.00/CH2Cl2).1H NMR (300 MHz, CDCl3) d 4.41 (ddd, 3H,

J = 3.1 Hz, J = 5.1 Hz, J = 9.5 Hz, H-500), 4.59 (dd, 3H,

J = 5.2 Hz, J = 12.5 Hz, H-600a), 4.71 (dd, 3H, J = 2.9 Hz,

J = 12.5 Hz, H-600b), 5.22 (d, 3H, J = 9.5 Hz, H-100), 5.91

(t, 3H, J = 9.7 Hz, H-400), 6.04–6.15 (m, 6H, H-200 H-300),

7.26–7.52 (m, 36H, H-ar), 7.82–8.02 (m, 24H, H-ar), 8.96

(s, 3H, H-2). 13C NMR (75 MHz, CDCl3) d 63.3 (s, 3C,

C-600), 69.4 (s, 3C, C-400), 70.8 (s, 3C, C-200), 72.6 (s, 3C, C-100),

74.1 (s, 3C, C-300), 77.4 (s, 3C, C-500), 126.2 (s, 3C, C-1 C-3

C-5), 128.5, 128.5, 128.6 (3s, 24C, CH-ar), 129.9, 130.0

(2s, 24C, CH-ar), 131.3 (s, 3C, C-2 C-4 C-6), 133.3, 133.4,

133.6, 133.6 (4s, 12C, CH-ar), 164.9, 165.2, 165.9, 166.3

(4s, 12C, COPh), 167.6 (s, 3C, C-30), 174.1 (s, 3C, C-50).

ESI-MS (positive mode) m/z: 2039.6 [M + Na]+.

1,3,5-Tris[30-C-b-D-glucopyranosyl)-10,20,40-oxadiazol-50-yl]-benzene (14). A solution of benzoylated tris-oxadiazole 13 (296

mg, 0.15 mmol) and NaOMe (15 mg, 0.3 mmol) in CH2Cl2/

MeOH (8 mL, 3:5) was stirred at room temperature for 3

hours. The solution was neutralized with a cation exchange

resin (Amberlite IR-120, H+ form) and the resin washed with

MeOH (3 � 5 mL). The filtrate was evaporated off and the

residue was dissolved in MeOH then precipitated with

CH2Cl2. The resulting solid was washed with CH2Cl2 (2 �5 mL) and dried under vacuum to afford the hydroxylated tris-

oxadiazole 14 (114 mg, 99%) as a white foam. Rf = 0.63

(CH2Cl2/MeOH, 9:1). [a]D = +10.6 (c = 1.00/H2O). 1H

NMR (300 MHz, D2O) d 3.61–3.84 (m, 15H, H-200 H-300

H-400 H-500 H-600a), 3.97 (d, 3H, J = 11.9 Hz, H-600b), 4.67

(d, 3H, J = 9.4 Hz, H-100), 8.74 (s, 3H, H-2 H-4 H-6). 13C

NMR (75 MHz, D2O) d 59.0 (s, 3C, C-600), 71.0 (s, 3C, C-100),

67.4, 69.9, 74.9, 78.6 (4s, 12C, C-200 C-300 C-400 C-500), 123.5

(s, 3C, C-1 C-3 C-5), 129.2 (s, 3C, C-2 C-4 C-6), 166.5 (s, 3C,

C-30), 172.0 (s, 3C, C-50). ESI-MS (negative mode) m/z:

812.9 [M + HCO2�]�. ESI-MS (positive mode) m/z: 791.7

[M + Na]+. HR-ESI-MS (positive mode) m/z: calcd. for

C30H36N6NaO18 [M + Na]+ 791.1984, found 791.1981.

Glycogen phosphorylase inhibition measurements

Glycogen phosphorylase b was prepared from rabbit skeletal

muscle according to the method of Fischer and Krebs,24 using

dithiothreitol instead of L-cysteine, and recrystallized at least

three times before use. Kinetic experiments were performed in

the direction of glycogen synthesis as described previously.25

Kinetic data for the inhibition of rabbit skeletal muscle

glycogen phosphorylase were collected using different concen-

trations of a-D-glucose-1-phosphate (2–20 mM), constant

concentrations of glycogen (1% w/v) and AMP (1 mM), and

various concentrations of inhibitors. Inhibitors were dissolved

in dimethyl sulfoxide (DMSO) and diluted in the assay buffer

(50 mM triethanolamine, 1 mM EDTA and 1 mM dithio-

threitol) so that the DMSO concentration in the assay should

be lower than 5%. The enzymatic activities were presented in

the form of double-reciprocal plots (Lineweaver–Burk) apply-

ing a nonlinear data analysis program. The inhibitor constants

(Ki) were determined by Dixon plots, by replotting the

slopes from the Lineweaver–Burk plots against the inhibitor

concentrations.26,27 The means of standard errors for all

calculated kinetic parameters averaged to less than 10%.

Ki values for compounds 9 and 14 were also estimated and

found to be 490 � 45 mM and 535 � 50 mM, respectively.

The poor solubility of inhibitors 5, 7 and 15 limited the

concentrations used in the kinetic studies. The inhibition of

glycogen phosphorylase was therefore determined at 625 mMand 2.5 mM concentrations of these inhibitors and given in

Table 2.

Acknowledgements

The authors wish to thank Universite Claude Bernard Lyon 1

and CNRS for financial support. A stipend to S. C. and

financial support from Region Rhone-Alpes )Cluster de

Recherche Chimie* are gratefully acknowledged. CNRS is

also thanked for additional funding through )Programme

Interdisciplinaire: Chimie pour le Developpement Durable*.

This work was also supported by grants from the Hungarian

Science Research Fund (OTKA K60620) and the Hungarian

Ministry of Health (ETT 083/2006).

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The formation of silver nanofibres by liquid/liquid interfacial reactions:

mechanistic aspects

Kun Luo and Robert A. W. Dryfe*

Received (in Montpellier, France) 9th June 2008, Accepted 22nd August 2008

First published as an Advance Article on the web 23rd October 2008

DOI: 10.1039/b809654f

The liquid/liquid interfacial reaction (LLIR) between silver nitrate in aqueous solution and

ferrocene in organic solution has been investigated: the resultant silver deposit is found to contain

long, well-defined nanometre scale fibres, together with thin silver nanowire networks. In situ

optical microscopy and ex situ scanning electron microscopy indicate that the 1D growth of the

interfacial deposits is due to recrystallisation of the structure formed initially. Geometric factors

are found to exert a larger effect on the 1D growth of silver by LLIRs compared to the

electrochemical mechanism previously suggested by Scholz et al.

1. Introduction

Nanostructures (i.e. structures with at least one dimension in

the range of 1 to 100 nm) have attracted increasing attention

because of their unusual chemical and physical properties.

There has been particular interest in methods of forming

one dimensional (1D) nanostructures, including nanowires,

because such structures provide a better model system for

investigating the dependence of electronic transport, optical

and mechanical properties on size confinement and dimen-

sionality.1 Strategies for achieving 1D growth have been

summarized by Xia et al.,2 these include: (i) use of the

intrinsically anisotropic crystallographic structure of a solid;3

(ii) introduction of a liquid/solid interface to reduce the

symmetry of a seed;4 or use of supersaturation control to

modify the growth habit of a seed;5 (iii) use of various

templates with 1D morphologies to direct the formation of

nanostructures6–9 (iv) assembly of zero dimensional nano-

structures (i.e. nanoparticles);10 (v) use of appropriate capping

reagent(s) to kinetically control the growth rates of various

facets of a seed.11,12 Another interface, the liquid/liquid (L/L)

interface, can also be used to limit the growth of materials, as

in (ii) above, or to assemble the symmetry of nanoparticles

(NPs), as in (iv) above, where liquid/liquid interfacial reactions

(LLIRs) are involved.

Metal NPs can be grown at the L/L interface either electro-

chemically, by applying a voltage across the L/L interface

when sufficient electrolytes are present in each phase,13 or by

spontaneous chemical reaction where the electron exchange

between redox couples present in the oil and water phases is

normally accompanied with biphasic ion exchange. Using the

former approach, gold NPs,14 platinum NPs15 and pyrrole

oligomers16 have been prepared electrochemically at the inter-

face between immiscible electrolyte solutions. The second

approach, using spontaneous deposition, can be traced back

to Faraday’s formation of colloidal gold at the water/carbon

disulfide boundary. In this case, particles can either be formed,

or pre-formed particles can be assembled,17 at the L/L inter-

face. A surprising variation in particle morphology has been

reported at the L/L interface. Most interfacial deposits appear

to consist of spherical NPs, which assemble to form films, or

aggregate into larger structures if no stabilising ligands are

present.18,19 The intrinsic difficulty in studying the larger-scale

structure is in finding an appropriate microscopic technique to

probe particle morphology in situ. Recent studies have des-

cribed the preparation of well-defined metal and metal oxide

NPs at the toluene/water interface, with X-ray scattering being

used to study the nanometre-scale assembly of Au NPs into an

ordered interfacial film.20,21 Silver deposition by LLIR, the

focus of this manuscript, has been described in a number of

previous reports. Silver assembly (as opposed to formation) in

the presence of surfactants at the water/dichloromethane

interface produces a ‘‘metal liquid-like film’’,22 whose struc-

ture has been described as micron-scale flocs of silver NPs.23,24

Assembly of silver NPs at an aqueous/chloroform interface in

the presence of thiol species has also been described.25 Other

reports have suggested that more unusual structures are ob-

served for Ag assembly and/or deposition at the L/L interface.

Agitation of aqueous silver hydrosols, during their assembly at

the water/toluene interface, has been reported to form

‘‘2D networks of uniform diameter nanowires’’.26 The forma-

tion of silver deposits, by interfacial reduction with an organic

phase electron donor, gives rise to intergrown ‘‘whisker’’

structures, although in this case the geometry is not uniform.27

The latter article noted that a transition between 1D and 2D

growth could be tuned according to the experimental condi-

tions of the spontaneous LLIR. By choosing appropriate

organic solvents and concentrations of the reagents in the

two phases, either silver whiskers (with radii from about 50 nm

to 50 microns) or ultrathin Ag films were observed. Herein, we

present further investigations into the spontaneous LLIR

between Fc in various organic solvents and aqueous AgNO3

solutions, where long and well-defined Ag nanofibres were

found under appropriate reaction conditions. The morpholo-

gical evolution and reaction mechanism are also discussed,

based on the micrographic observations.

School of Chemistry, University of Manchester, Oxford Road,Manchester, UK M13 9PL. E-mail: [email protected];Fax: +44 (0)161 275 4734



2. Experimental

Silver nitrate (AgNO3, BDH Chemicals, GPR), ferrocene

(Fc, 99%, Alfa Aesar), 1,2-dichoroethane (DCE, Rathburn,

HPLC), nitrobenzene (NB, 99%, Sigma) and toluene

(99%, Fisher Scientific) were used directly without further

treatment. Typically, 3.3 mM of silver nitrate solution was

prepared with deionised water from an Elga ‘‘Purelab Ultra’’

(Elga, Marlow, UK) system. Ferrocene was dissolved in DCE,

or other organic solvents, as the organic phase for the inter-

facial reactions. The aqueous solution of AgNO3 and one of

the organic Fc solutions were placed together in a glass tube

with dimensions 75 mm (height) � 25 mm (diameter), follow-

ing the sequence that the higher density phase was added prior

to the light one. The mixture was then kept still at ambient

temperature. The interfacial deposit was collected after 48 h of

reaction, and transferred onto glass slides and dried in air. It

was then washed with acetone and deionised water separately,

and dried at ambient temperature before further analysis.

Copper grids with holey carbon films (S147-4, Agar Scientific)

were employed to collect samples for analysis via transmission

electron microscopy (TEM) and high resolution transmission

microscopy (HRTEM). The samples were rinsed with both

acetone and deionised water, and were dried in air in order to

remove the remained contaminants.

In situ optical microscopy was performed by a Leica DMIL

optical microscope fitted with a Sony CCD-IRIS camera on

an anti-vibration system (Active vibration isolation system

TS-200, HWL Scientific Instruments GmbH). The X-ray

diffraction (XRD) analysis was carried out using an Oxford

Diffraction System (Xcalibur 2, Mo-Ka = 0.7093 A), and XL

30 FEG Philips and ESEM XL30 Philips electron microscopes

were employed at 15 kV for scanning electron microscopy

(SEM). TEM and HRTEM were performed with a Tecnai F30

FEG-TEM system operating at 300 kV.

3. Results

The deposition of Ag resulting from the LLIR between

AgNO3 in water and Fc in organic solvent can be written as:27

Fc(o) + Ag+(w) + X�(w) - Fc+(o) + X�(o) + Ag(s) (1)

where the subscripts ‘‘s’’, ‘‘w’’, and ‘‘o’’ in the reactions

represent interfacial, aqueous and organic phases, respectively,

and the anion X� is added to balance the charge since no Fc+

transfer to the aqueous phase was believed to occur (in the

experiments reported herein, X� is nitrate). The reaction was

monitored in situ by an optical microscope placed on the active

anti-vibration system. As shown in Fig. 1(a), at the beginning

of the reaction (ca. 1 min), only separate particles were

observed at the L/L interface. Many particles were rapidly

generated, and started aggregating, after about five minutes of

contact between the two phases (Fig. 1(b)). After 10 min,

fractal-like aggregates appeared at the interface, as illustrated

by Fig. 1(c), and the fluidity at the interface became obstructed

after 25 min of reaction, owing to the appearance of large Ag

agglomerates (see Fig. 1(d)). After 24 h of reaction (Fig. 1(e)),

the L/L interface became somewhat solidified, and a grey

coloured deposit was seen. The interfacial deposits became a

bit denser compared to Fig. 1(e) when the reaction time was

extended to 48 h (shown in Fig. 1(f)), but no visible fibre-like

deposit was found under the optical microscopy. The process

was also investigated ex situ by SEM, and the morphological

evolution of the Ag deposit is illustrated by Fig. 2, although an

important point to note here is that the extraction and drying

of the sample could induce changes in morphology. The

interfacial deposit seen after twenty five minutes of reaction

appeared as micron-scale ‘‘flakes’’ with a few nuclei on the

surface as displayed in Fig. 2(a), indicative of a possible

destruction of an originally compact 2D interfacial layer

during the sample collection. After 1 h of the reaction, some

1D Ag deposits can be differentiated from others (see

Fig. 2(b)). After 4 h of reaction (Fig. 2(c)), some of the

‘‘microflakes’’ are found with holes and irregular edges in

the background and co-exist with the 1D Ag nanostructures,

which are not seen at shorter or longer reaction times. After 48

h, long Ag nano-scale fibres are observed, as illustrated in

Fig. 2(d), together with some short 1D nanostructures and

‘‘microflakes’’. The inset shows that the nanofibre is rather

smooth and well-defined at a larger magnification. The Ag

nanofibres shown in Fig. 2(d) are measured and give an

average diameter of 171 � 4 nm (N = 13) with a mean aspect

ratio of ca. 174, where the largest aspect ratio from the other

micrographs is observed to be ca. 450. Fig. 2(e) further reveals

that the growth of the nanofibres originates from defects, such

as independent nuclei or the edges of the ‘‘microflakes’’ etc.

TEM and HRTEM micrographs offer the means of observa-

tion under higher magnification. Fig. 3(a) suggests that some

of the ‘‘microflakes’’ observed under SEM are actually com-

posed of networks of thinner 1D nanostructures, termed

‘‘nanowires’’ in the following text, with an average diameter

of 14.8 � 3.7 nm (N = 163), where the distribution of the

diameter values visible in Fig. 3(a) is shown in Fig. 3(b).

Fig. 3(a) also illustrates the evolution from the 2D film to

1D nanostructures, where a few branches of nanowires are

observed to extend from a piece of film higlighted by the circle

in the figure. The triangular highlights in Fig. 3(c) suggest that

Fig. 1 Optical micrographs recorded during the formation of Ag

interfacial deposits by LLIR: (a) at ca. 1 min, (b) at 5 min, (c) at

25 min, (d) at 1 h, (e) at 24 h, (f) at 48 h. The length of the scale bars in

the figure is 50 microns.


the growth of the nanowires originates from the triangular

nuclei, and the 1D extension from the nuclei is ‘‘welded’’ on

meeting other wires as shown by the appearance of lattice

fringes under TEM (circular highlight). The HRTEM image in

Fig. 3(d) indicates that the wires are nanocrystalline. From the

micrographs presented in Fig. 2 and 3, one can summarize that

the microflakes (the 2D growth owing to the presence of the

L/L interface) are formed in the early stages of the LLIR, a

process followed by 1D growth at the active sites of the

interfacial layer after a long-term LLIR, such as the edge of

the layer or the tips of triangular nuclei. The intermediate

stage (holey microflakes with irregular edges displayed in

Fig. 2(c)) at 4 h of the LLIR, suggests the appearance of a

parallel process which might be either the dissolution of the

as-formed microflakes; alternatively the secondary aggrega-

tion or growth of the nuclei may occur, to form thin nanowire

networks shown in Fig. 3(a). Considering that complete

microflakes and nanofibres are seen in the SEM micrographs

at 48 h of the LLIR, and the nanowire networks can only be

seen under TEM, the visible networks in Fig. 3(c) would

appear to arise from the dissolution of the as-formed 2D

layers. The LLIR process is therefore viewed as involving:

(i) formation of Ag nuclei at the L/L interface; (ii) agglomera-

tion of the as-prepared Ag nuclei to form 2D flakes due to the

constraint of the L/L interface; (iii) dissolution of some of the

initial Ag nuclei while secondary nucleation or growth occurs

elsewhere. Some of the larger 2D structures can even be

dissolved if the nuclei are depleted; (iv) the structural defects,

including the tips of the triangular nuclei, independent nuclei

and the edges of the microflakes, offer active sites for new

nucleation, leading to 1D growth. The diameter of the 1D

nanostructures is normally dependant on the size of the

protuberance of the defects, for example larger nuclei for the

nanofibres, compared to the corner of the smaller triangular

nuclei for the nanowires.

Different organic solvents, such as NB and toluene, were

used instead of DCE in the LLIR process. A thin and

Fig. 2 SEM micrographs of the Ag interfacial deposits collected at different times during the LLIR process: (a) at 25 min, (b) at 1 h, (c) at 4 h,

(d) at 48 h, the inset shows the diameter of the nanowire is around 100–200 nm; (e) at 48 h, 1D growth stems from the defects of the interfacial deposit.

Fig. 3 TEM and HRTEM micrographs of the Ag interfacial deposit

of the LLIR reaction between 3.3 mM AgNO3 aqueous solution and

5 mM Fc in DCE at 48 h: (a) nanowire networks in the deposit, the

highlighted part shows the 1D growth from a piece of a 2D thin film,

(b) distribution of the diameter of the nanowires, (c) the connection of

the nanowires, the parts highlighted with ‘‘&’’ indicate that nanowires

protrude from portions that were originally nanoparticulate. The

other parts highlighted with ‘‘J’’ denote the ‘‘welding’’ positions of

the nanowires; (d) HRTEM image of the nanowires, which indicates

the nanowires are composed of crystalline silver.


continuous film was observed at the NB/water interface in a

Teflon container after 24 h of LLIR between 3.3 mM AgNO3

solution and 1 mM Fc in NB (see Fig. 4(a)), in accordance

with the results described by Scholz and Hasse.27 In contrast, a

radial pattern gradually appeared after a period of time in a

glass container (see Fig. 4(b)), in possible association with

random vibrations. The pattern was quite stable over time,

and exhibited a ‘‘self-recovery’’ capability from external dis-

ruption of the interface. If the LLIR was forced to stop ‘‘half

way’’ by the depletion of the silver salt, a ring-like thin film

extending from the surface of the glass tube with some

irregular deposit in the centre was seen, indicative of the

adhesion of Ag nuclei on the hydrophilic surface of glass (as

shown in Fig. 4(c)). Fig. 4(d) shows that no pattern appeared if

the LLIR was carried out on an active anti-vibration table for

48 h. Hence, the pattern shown in Fig. 4(b) can be interpreted

in terms of: (i) the adsorption of the initial Ag nuclei on the

surface of glass tube; (ii) growth of the Ag deposit to form a

ring-like interfacial deposit layer as shown in Fig. 4(c); (iii) this

process continues and consequently the interfacial layer is able

to cover the whole L/L interface; (iv) vibrations cause standing

waves to form at the L/L interface, and the interfacial layer is

then easily folded since its outer edge is ‘‘pinned’’ to the

surface of the glass container. However, when the Ag nuclei

are adjacent to the Teflon surface, no adsorption occurs, and

the Ag layer formed by the LLIR ‘‘floats’’ on the L/L inter-

face, hence no such standing waves are set up. Consequently

no radial pattern appeared after 48 h of reaction.

The microstructure of the interfacial deposits in water/NB

and water/toluene was observed under SEM. Fig. 5(a) shows

that the interfacial deposit collected from the NB/water inter-

face displays an analogous microstructure to that in Fig. 2(d).

Long and well-defined nanofibres and other short 1D nano-

structures are observed with the presence of ‘‘microflakes’’.

In contrast, no 1D growth is seen after the LLIR between

3.3 mM AgNO3 solution and 5 mM Fc in toluene, but only

smoother ‘‘microflakes’’ are shown in Fig. 5(b) under SEM.

The X-ray diffraction patterns of the Ag interfacial deposits

formed by the LLIRs between 3.3 mM AgNO3 solution and

5 mM Fc in DCE, NB and toluene are presented in Fig. 6.

Reflections assigned to Ag (111), (200), (220), (311), (222)

planes are marked in the plot (Fm3m, a= 4.08 A, JCPDF No.

02-1167). The crystallite size is estimated from the broadening

of X-ray diffraction peaks by Scherrer’s equation:28

Bcrystallite = kl/(L cos y) (2)

where l is the wavelength of the X-ray, y is the Bragg angle,

L is the average crystallite size measured in a direction

perpendicular to the surface of the specimen, and k is a

constant taken to be 0.9. The crystallite size calculated from

the (111) reflection of the interfacial deposit of DCE/water

system is 2.2 nm, which is the same as that in the toluene/water

system. The crystallite size for the interfacial deposit from the

NB/water system is 3.1 nm. Other polar and nonpolar organic

solvents, such as 1,2-dichlorobenzene and silicone oil (data not

shown), were also employed to perform the LLIR with the

same concentrations of reactants, and the morphology of the

interfacial deposits also follows the same trend, in that polar

organic solvents favour the formation of 1D Ag deposits at

L/L interfaces.

The effect of varying the concentrations of the reagents was

investigated. As shown in Fig. 7(a) and (b), a lower concentra-

tion of AgNO3 solutions was employed to react with 5 mM Fc

in DCE. The 0.33 mM AgNO3 (c+Ag/cFc = 0.066) exhibited a

tendency to 1D growth, while the 0.07 mM AgNO3 solution

(c+Ag/cFc = 0.014) basically formed Ag aggregates. If the

concentration of Fc was varied, as illustrated in Fig. 7(c) and

(d), 1D growth could be seen but was not fully developed

in the case of 3.3 mM AgNO3 solution reacted with 0.5 mM

(c+Ag/cFc = 6.6) and with 0.1 mMFc in DCE (c+Ag/cFc = 33).

The influence of the organic solvent on the resultant composi-

tion of the aqueous phase was also investigated via the visible

absorbance of the aqueous phase (Fig. 8). The toluene/water

system shows the highest transfer of Fc+ to water, whereas the

NB/water displays the weakest spectral response. The transfer is

also found to be proportional to the concentration of Fc

employed in the DCE solutions (data not shown).

4. Discussion

Scholz and Hasse27 have proposed an electrochemical mecha-

nism for the deposition of silver via LLIR, where the reaction

occurs at the L/L interface (see eqn (1), above). In the

treatment of Scholz et al., the nuclei at the L/L interface were

viewed as disc-shaped microelectrodes, where the current (i) to

the equivalent disc-shaped silver/electrolyte interface was

described by:

idisc ¼ nFAdiscca4Da

prð3Þ

where n is the number of electrons transferred, F is the

Faraday constant, Adisc is the surface area of the disc, ca and

Fig. 4 The Ag film obtained by interfacial reaction between 3.3 mM

AgNO3 aqueous solution and 1 mM Fc in NB (after reaction for 24 h):

(a) in Teflon container, (b) in glass container, (c) stopped half-way

owing to the depletion of AgNO3 in glass container, (d) the LLIR in a

glass container on an active anti-vibration table.


Da the bulk concentration and diffusion coefficient of

species a, respectively, and r is the radius of the disc. When

the same crystal is growing in a 1D mode, a cylinder electrode

was used by Scholz and Hasse to approximate the current flow

across the silver/liquid interface:

icyl ¼ nFAcylca2Da

r ln tð4Þ

where t = Dat/r2, and t is the time. Acyl and r are the surface

area and radius of the cylinder, respectively. Since the oxida-

tive process (oxidation of Fc) should balance the reductive one

(reduction of Ag+) at all times, Scholz and Hasse assumed the

1D Ag structure must protrude into the organic phase for the

above-named fluxes to balance under conditions of excess

silver ion, since equating (3) and (4) leads to:

Acyl

Adisc¼

cAgþðwÞcFcðoÞ

2DAgþ

DFc

ln tp

ð5Þ

Fig. 5 SEM micrographs of the Ag interfacial deposits at different LLIR systems: (a) 3.3 mM AgNO3 aqueous solution with 5 mM Fc in NB,

(b) 3.3 mM AgNO3 aqueous solution with 5 mM Fc in toluene.

Fig. 6 X-Ray diffraction patterns of the Ag interfacial deposits by the

LLIR between 3.3 mM AgNO3 aqueous solution and 5 mM Fc in (a)

DCE, (b) NB and (c) toluene.

Fig. 7 SEM micrographs of the Ag interfacial deposits as a function of reagent concentration: (a) 0.33 mM AgNO3 aqueous solution with 5 mM

Fc in DCE, (b) 0.07 mM AgNO3 aqueous solution with 5 mM Fc in DCE, (c) 3.3 mM AgNO3 aqueous solution with 0.5 mM Fc in DCE,

(d) 3.3 mM AgNO3 aqueous solution with 0.1 mM Fc in DCE.


where DAg+ and DFc are the diffusion coefficients of Ag+ in

the aqueous phase and Fc in the organic phase, respectively.

Eqn (5) indicates that the higher concentration ratio of Ag+ to

Fc increases the ratio of Acyl to Adisc and favours 1D growth,whereas the converse case should lead to 2D films. Note that

the Ag+ reduction can occur at a site distant from Fc

oxidation, if the electrical conductivity of the deposit is

sufficient.

The data presented in Fig. 7 act as experimental tests of

eqn (5): from inspection of the deposits, it is clear that the

initial concentration ratio of the LLIR is not the sole factor

controlling deposit morphology (cf. eqn (5)), and the concen-

tration of AgNO3 itself actually exerts a large effect on the

final deposit. Combined with the observation (in Fig. 3) that

no 1D Ag nanostructures were found in the first 25 min of the

LLIR, which also suggests that a secondary crystallisation step

is involved in the 1D growth, this leads to the conclusion that

the actual mechanism is more complicated than the simple

electrochemical process suggests. However, the experiments

on the effect of organic solvents demonstrate that the use of

non-polar solvents suppressed the formation of 1D structures.

The morphological evolution seen here suggests that 2D thin

films are formed initially at the L/L interface, followed with a

transformation from 2D to 1D growth, in the case of more

polar organic solvents.

The spontaneous LLIR between Fc in the organic phase and

Ag+ in the aqueous solution initially generates Ag nuclei. An

associated transfer of Fc+ from the organic phases (i.e. DCE,

NB or toluene) to water, or of nitrate in the reverse direction,

must occur to preserve electroneutrality. The visible absor-

bance band (see Fig. 8) centred on 620 nm is attributed to Fc+,

on the basis of a previous report,29 but the extent of transfer is

a function of the polarity of the solvent: the least polar solvent

(toluene) is least able to solvate the Fc+ nitrate ion pair, hence

in the toluene case, transfer of Fc+ to the aqueous phase

predominates.

This observation, combined with the change to 2D morpho-

logy on using the less polar solvent, leads to the following

mechanism being postulated for the 1D growth mode. Ag

nuclei are initially generated by the spontaneous LLIR, and

form a 2D interfacial layer because of the constraint of the L/L

interface. After that, a transformation from 2D layers to 1D

nanostructures occurs as a higher flux of reactants to

the deposit can be sustained by radial diffusion.30 Con-

sequently, parts of the 2D structures appear to dissolve,

accompanied with the emergence of 1D nanostructures. The

1D growth is observed to occur at active sites with high

surface energy, such as independent nuclei, the edges of 2D

structures or the corner of the small triangular crystals,

which show surprisingly well-defined long nanofibres without

any branches. Nanowires from the Ag triangular crystals

can even connect to form nanowire networks. The evolution

of the 1D process is, however, suppressed in less polar

solvents since it requires the (unfavourable) formation of a

ferrocenium nitrate ion pair in the organic phase, by transfer

of the nitrate, or the transfer of the ferrocenium to the aqueous

phase. The latter process is more favourable, but the extent of

transfer depends on the distance from the interface where the

ferrocenium ion is formed. In the case of a 2D Ag deposit,

the ferrocenium ion is formed adjacent to the interface

and is readily transferred. By contrast, if 1D growth occurs,

the ferrocenium ion may be formed some distance from the

aqueous phase (a distance determined by the length of

the structure). We therefore suggest that the driving force

behind the morphological change observed in the deposit is the

solvation of the ions formed by the LLIR.

5. Conclusions

The liquid/liquid interfacial reaction (LLIR) between silver

nitrate in aqueous solution and ferrocene in various organic

solvents has been investigated: long and well-defined silver

nanofibres and thin nanowire networks were obtained in more

polar media. In situ optical microscopy and ex situ scanning

electron microscopy indicate that the 1D growth of the inter-

facial deposits is due to directed recrystallization, where geo-

metric factors associated with the flux to the growing deposit,

and energetic factors, associated with the solvation of the ions

generated, play an important role.

Acknowledgements

The authors thank the financial support from the UK

Engineering & Physical Science Research Council (EPSRC,

grant EP/C509773/1).

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The role of nucleophilic catalysis in chemistry and stereochemistry of

ribonucleoside H-phosphonate condensationwMichal Sobkowski,*

aJacek Stawinski

band Adam Kraszewski

a

Received (in Montpellier, France) 24th July 2008, Accepted 10th September 2008


DOI: 10.1039/b812780h

The efficiency and stereoselectivity of condensation of ribonucleoside 30-H-phosphonates with alcohols were

investigated as a function of amines used for the reaction. It was found that irrespective of the presence or absence

of nucleophilic catalysts, the Dynamic Kinetic Asymmetric Transformation (DYKAT) was the major factor

responsible for the stereoselective formation of the DP(SP) isomers of the H-phosphonate diesters, and a

mechanistic rationalization of this observation was proposed. In addition, studies on the reactions carried out in

the presence of various bases led to the conclusion that certain sterically hindered pyridines, e.g. 2,6-lutidine, may

act as nucleophilic catalysts in the condensation of ribonucleoside 30-H-phosphonates with alcohols.

Introduction

P-Chiral oligonucleotide analogues (e.g. phosphorothioates,2

phosphoramidates,3 methylphosphonates,4 or boranophos-

phates5) having defined configuration at the phosphorus atom

find diverse applications in investigations of nucleic acid inter-

actions with other biologically important molecules, for exam-

ple proteins, RNA, and DNA.6 Such P-chiral oligonucleotides

may also be considered as potential drugs for nucleic acid-based

therapies,7 that could permit a more precise tuning of oligo-

nucleotide interactions with the biological targets than is possible

with the currently used pools of P-diastereomers. This should

also relieve problems of potential variation of therapeutic and

toxic effects resulting from different ratios of P-diastereomers

produced in various batches of oligonucleotide drugs.

There are several strategies to stereocontrolled synthesis of

P-chiral oligonucleotides.8 One of them, stereoselective (or more

precisely, diastereoselective) condensation of ribonucleosideH-phos-

phonates,9 attracted our attention due to its simplicity and high

efficiency. It makes use of commercially available H-phosphonate

synthons which are condensed with nucleosides under standard

reaction conditions commonly used for the synthesis of H-phos-

phonate diesters to provide DP diastereomersz as major products.

Recently, we have proposed a Dynamic Kinetic Asymmetric

Transformation (DYKAT) as a possible mechanism for the

stereoselectivity observed in these reactions.1 According to this

model, diastereomers of nucleoside H-phosphonic–pivalic

mixed anhydrides 2 exist in a rapid equilibrium, and one of

them, namely the LP(SP) diastereomer, is significantly more

reactive towards nucleosides (or alcohols) than the other one

(Fig. 1 and Chart 1). To simplify mechanistic considerations,

in our earlier studies the role of nucleophilic and base catalysis

by the amines was consciously neglected. However, since the

participation of nucleophilic catalysis in condensation of

H-phosphonates is a well-established phenomenon,11–15 it

was important to examine and to assess its impact on the

asymmetric induction in the reactions investigated. In this

paper we present studies on the role of nucleophilic catalysis

in the chemistry and stereochemistry of condensation of

ribonucleoside H-phosphonates with alcohols.


In routine condensations of nucleoside H-phosphonates

pyridine or quinoline (either neat or diluted with non-basic

solvent) is used as a basic component of the reaction mixture.

Both of these weakly basic heterocyclic amines (pKa 5.2 and

4.9, respectively) secure fast and quantitative formation

of H-phosphonate diesters due to their ability to act as

nucleophilic catalysts.11,12 In contrast to this, in the presence

of more powerful nucleophilic catalysts, e.g. NMI or DMAP,ynucleoside H-phosphonates are prone to P-acylation that

compromises the diester formation.12 Also strongly basic

tertiary amines (e.g. TEA, pKa 11.0) are usually avoided since

these can promote undesired base-catalysed bis-acylation of

H-phosphonate monoesters,13,15 while in the presence of less

basic tertiary amines (e.g. DMA, pKa 5.1) the condensations

a Institute of Bioorganic Chemistry, Polish Academy of Sciences,Noskowskiego 12/14, 61-704 Poznan, Poland.E-mail: [email protected]; Fax: +48 61 8520 532;Tel: +48 61 852 8503

bDepartment of Organic Chemistry, Arrhenius Laboratory, StockholmUniversity, S-106 91 Stockholm, Sweden

w Stereochemistry of internucleotide bond formation by the H-phos-phonate method. Part 4.1

z For the compounds presented in this paper the DP descriptor refersto a structure in which the P–H bond is directed to the right in theFischer projection, and in the LP one, to the left. The full DP/LP

notation is described in ref. 10.

y Abbreviations: DABCO, 1,4-diazabicyclo[2.2.2]octane; DIPEA, di-isopropylethylamine; DMAP, 4-(N,N-dimethylamino)pyridine; DMA,N,N-dimethylaniline; DTBP, 2,6-di-tert-butylpyridine; EDIPP,4-ethyl-2,6-diisopropyl-3,5-dimethylpyridine; HMTA, hexamethylene-tetramine; Lut, 2,6-lutidine; MPO, 4-methoxypyridine N-oxide; NMI,N-methylimidazole; PvCl, pivaloyl chloride; Py, pyridine; TEA,triethylamine.



are effective but sluggish (at least 10 times slower than those

for pyridine), presumably due to lack of nucleophilic cata-

lysis.12 Thus, it was somewhat surprising that 2,6-lutidine

(pKa 6.7), which is usually considered as poorly nucleophilic

base,16,17 promoted condensations of ribonucleoside H-phos-

phonates with similar efficiency as pyridine or quinoline.1

Moreover, the stereochemistry of the reactions performed in

the presence of 2,6-lutidine was the same as that with pyridine.

The above called into question the commonly accepted

non-nucleophilic character of 2,6-lutidine and prompted us

to consider the involvement of nucleophilic catalysis as an

additional process in the DYKAT mechanism (Fig. 1).

Since the involvement of P–N+ adducts of type 3 (Fig. 1) in

ribonucleoside H-phosphonate diester formation has to be

crucial for the rate of condensation as well as for stereo-

chemical outcome of the reaction (Fig. 2), we undertook

investigations to pinpoint the cases in which amines acted

solely as base catalysts or as base and nucleophilic catalysts

during H-phosphonate condensations. To this end, the reac-

tions of H-phosphonate 1 with ethanol were carried out in the

presence of selected tertiary amines, various pyridine deriva-

tives, and strong nucleophilic catalysts. The obtained data

(Table 1) showed that irrespective of significant differences in

the yields and stereoselectivity observed for different amines,

the same DP(SP) diastereomer of diester 4 was always formed

as the main product. In the light of our earlier studies,1,18 these

results might suggest that in the absence of nucleophilic

catalysts, the previously described DYKAT mechanism oper-

ated at the level of the mixed anhydride 2 (Fig. 2, Path B),

while in the nucleophile-catalyzed reactions, an analogous

DYKAT took place at the level of adducts of type 3 (Fig. 2,

Path A2).zAdditionally, these experiments confirmed the earlier find-

ings11–15 that neither powerful nucleophilic catalysts nor

strongly basic tertiary amines could promote quantitative

condensations of H-phosphonates with alcohols. However,

in contrast to the literature data,11–15 there was no (or very

little) side product formation, and the 31P NMR spectra of the

reaction mixtures revealed only presence of the expected

diester 4 and unreacted monoester 1 (Fig. 3). This lack of

by-product formation was tentatively attributed to low

Fig. 1 Putative routes of the reaction during stereoselective condensation of ribonucleoside H-phosphonate monoester 1 with alcohols and

nucleosides according to the DYKAT mechanism in the absence (curved arrows) and in the presence (central pathways) of a nucleophilic catalyst.

Chart 1

z Involving a rapid 3-DP " 3-LP equilibrium in which the morereactive diastereomer 3-DP was esterified preferentially.


concentration of the amines in the reaction mixtures (0.3 M or

ca. 2.5%).

In order to find sources for the incomplete condensations

that have been carried out in the presence of the amines

examined herein, the reactivity of pivaloyl chloride towards

nucleosides was investigated in separate experiments. It was

found that TEA and pyridine derivatives when used alone did

not promote significant acylation of nucleosides, however, in

the presence of strong nucleophilic amines (e.g. DMAP) or

TEA–pyridine mixtures, pivaloyl chloride was rapidly con-

sumed in the acylation of 50-OH or N3-H functions of

uridine.41 These side reactions could compete with the forma-

tion of the mixed anhydride 2 and, at least partly, could

be responsible for incomplete condensations. However,

H-phosphonate condensations were also not quantitative in

the presence of tertiary aliphatic amines alone (Table 1, entries

25–27), i.e. under the conditions in which the acylation of

nucleoside components was negligible.41 This issue was

addressed in additional experiments, which indicated that

the mixed anhydride 2 might undergo deacylation by pivalic

Fig. 2 Possible stereochemistry of esterification of the more reactive LP(SP) diastereomer of ribonucleoside H-phosphonic—pivalic mixed

anhydride 2 in the presence (Path A) and absence (Path B) of nucleophilic catalysts.

Table 1 Diastereomeric excess (de) of the DP(SP) diastereomer of the H-phosphonate diester 4b (Fig. 1, B = Ura) formed in the presence ofvarious amines

Entry Amine pKa (H2O)a pKa (DMSO) pKa (ACN) pKHBb dec,d (DP) Yield of diester (%)d

Strong nucleophilic catalysts1 MPO 2.119 3.520 12.421 62% 272 HMTAe 5.2 1.922 57% 663 NMI 7.0 14.323 2.724 60% 844 DABCOe 8.7 8.925 18.326 2.622 53% 425 DMAP 9.7 7.927 17.728 2.829 53% 85

Heteroaromatic amines6 Pyrazine 0.7 1.229 39% 557 Pyrimidine 1.2 1.429 39% 708 Tetramethylpyrazine 3.6 59% 1009 Quinoline 4.9 12.028 1.929 63% 10010 1,10-Phenanthrolinee 4.9 66% 10011 2,6-Di-tert-butyl-pyridine 5.030 1.031 47% 7012 Pyridine 5.2 3.227 12.528 1.929 62% 10013 2-Picoline 5.9 4.027 13.932 2.029 69% 10014 4-Picoline 6.0 3.827 14.532 2.129 68% 10015 Neocuproinee 6.2 64% 10016 2,5-Lutidine 6.4 68% 10017 3,4-Lutidine 6.5 4.327 14.732 2.229 63% 10018 2,6-Lutidine 6.7 4.427 14.432 2.129 70% 10019 2,4-Lutidine 6.7 4.527 15.032 70% 10020 2,4,6-Collidine 7.5 15.028 2.329 68% 10021 EDIPP (7.6)f 52% 9222 (�)-Nicotine 8.0 65% 10023 (�)-Nicotine 8.0 64% 100

Tertiary amines24 DMA 5.1 2.533 11.428 0.534 56% 10025 N-Methylmorpholine 7.4 15.635 1.722 70% 9126 TEA 11.0 9.036 18.828 2.022 75% 7427 DIPEA 11.4 1.122 71% 89

a Aqueous pKa data, unless otherwise indicated, are taken from ref. 37. b Hydrogen bonding basicity. pKHB = logK(formation of HB complex); larger

values correspond to greater basicity.38 c One should note that the difference between de values, for instance de 52% and de 75%, corresponds to

over two-fold increase of the stereoselectivity measured as a ratio of diastereomers (i.e. B3 : 1 vs. B7 : 1, respectively). d Determined via

integration of the corresponding 31P NMR signals. e For structure, see Chart 1. f Estimated, assuming an additive and similar methyl and ethyl

groups effect on the pKa39 and a linear correlation between a,a0-steric hindrance and pKa

40 of substituted pyridines.


acid with regeneration of the starting H-phosphonate mono-

ester 1 and formation of pivalic anhydride (a poor activator of

H-phosphonates42). The rate of deacylation was found to

correlate well with the ability of an amine conjugated acid to

form hydrogen bonds (quantified as a pKHB value8) rather

than with the amine basicity expressed by pKa. A plausible

rationale, which could account for the obtained results in-

volved an increased contribution of general acid catalysis

during decomposition of the mixed anhydride 2 by amines

having high pKHB (Fig. 4).43

Thus, it can be tentatively concluded that pivaloyl chloride

promoted coupling of H-phosphonates with alcohols in the

presence of strongly nucleophilic amines, and/or those of high

H-bonding basicity, did not go to completion due to con-

sumption of the condensing agent (PvCl) in the acylation

of nucleosides or due to formation of unreactive pivalic

anhydride via a partial deacylation of the mixed anhydride 2.

In contrast, pyridine and most of its derivatives investigated

secured quantitative condensations (with an exception of

pyridine derivatives bearing branched substituents in both

a positions)** despite significant differences in their pKa

(3.6–8.0) and considerably high pKHB values (1.9–2.3).

Although it might be argued that the high pKHB of pyridines

should be associated with high catalytic activity of their

conjugate acids which should lead to deacylation of the mixed

anhydride 2 (Fig. 4), apparently it was not the case in the

reactions discussed. A plausible explanation of the excellent

yields obtained for the most of the pyridines examined could

be the participation of nucleophilic catalysis, i.e. the involve-

ment of intermediate phosphonopyridinium adducts of type 3

(Fig. 2) which, as monofunctional entities, should undergo a

nucleophilic attack at the phosphorus centre only. While this is

readily understandable in the case of pyridine derivatives with

unhindered endocyclic nitrogen atoms (i.e. having at least one

a position unsubstituted), the question might arise, whether

this could hold also for a,a0-dimethylpyridines?

Although 2,6-lutidine and its derivatives are usually con-

sidered as poor nucleophiles,16,17 they can act as nucleophiles

under mild conditions undergoing, for instance, N-alkylation

with alkyl halides,44 alkyl iodonium triflate45 or radical

cations,46 or N-sulfonation with triflic anhydride.47 Notably,

in phosphorus chemistry the lack of nucleophilic properties of

2,6-lutidine was observed for P(V) compounds, e.g. for phos-

phoroiodidates,16 while whether nucleophilic catalysis by this

base may operate for H-phosphonates, remains to be deter-

mined. Since P(V) and P(III) compounds differ significantly in

electrophilicity,48 and the steric hindrance around the phos-

phorus atom in H-phosphonates is clearly lower than that

in P(V) compounds, significant differences in their reactivity

towards hindered pyridine derivatives cannot be excluded.

To get a better insight into this problem, condensations of

H-phosphonate 1 were performed in the presence of EDIPP

(a peralkylated 2,6-diisopropylpyridine derivative, pKa

B7.6)49 and DTBP (2,6-di-tert-butylpyridine, pKa 5.0) for

which the nucleophilicity might be safely excluded on steric

grounds. The yields of H-phosphonate diester 4 obtained in

these reactions (92 and 70%) were similar to those found for

trialkyl amines, while low stereoselectivity (de ca. 50%) was

similar to that observed for tertiary aniline derivatives

(e.g. DMA, de 56%). In contrast, all the other pyridine

derivatives, including a,a0-dimethylpyridines, differed only

slightly in stereoselectivity and invariably gave quantitative

condensations of H-phosphonate 1. Thus, it seems reasonable

to assume that the main route for H-phosphonate diester

formation in the presence of 2,6-lutidine derivatives could still

involve the nucleophilic catalysis (preventing in this way

deacylation of the mixed anhydride 2, and in consequence,

the yield deterioration), and that only bulky alkyl substituents

in a,a0 positions were able to suppress the nucleophilic proper-

ties of pyridine.

In additional experiments the condensations of uridine

H-phosphonate 1 with ethanol performed in the presence of

mixtures of 2,6-lutidine with more nucleophilic amines

(pyridine, NMI, DMAP, MPO) were investigated (Fig. 5). In

neither case were any specific effects due to the nucleophilic

amine noted, and the yields and stereoselectivity of the con-

densations were proportional to the weighted average of the

values obtained for each amine used separately. This lent

support to the aforementioned assumption that the same

mechanism (i.e. nucleophilic catalysis) was operating for 2,6-

lutidine and for other amines of known nucleophilic character.

Fig. 331P NMR spectra of the reaction of H-phosphonate 1 with

ethanol (3 equiv.) promoted by PvCl (1.5 equiv.) in DCM containing

3 equiv. of 2,6-lutidine or TEA. The minor signal (ca. 1.5%) at�3.2 ppmin the upper spectrum is in the region of P-acylated compounds.

Fig. 4 A possible participation of general acid catalysis in deacyla-

tion of the mixed anhydride 2 by pivalic acid.

8 The pKHB measures the relative strength of the acceptor in hydro-gen-bonded complex formation with a reference acid (H-bondingbasicity). pKa and pKHB may be unrelated.38

** Two heteroaromatic amines, pyrazine and pyrimidine, were appar-ently too weakly basic (pKa 0.7 and 1.2, respectively) to be efficientpromoters of the condensations investigated since a significant detri-tylation was observed during the course of reactions, even in thepresence of 6 equiv. of an amine (c E 5%). These amines were thusexcluded from further investigations.


Kinetic quenching experiments

To probe the involvement of nucleophilic catalysis in the

DYKAT mechanism, kinetic quenching experiments for

various amines were carried out using large excess of methanol.

Under such reaction conditions we expected to observe

significant changes in the ratio of diastereomers (with a

possible reversal of stereoselectivity1) of the produced

H-phosphonate diester 4a as a result of substantial increase

in the rate of esterification of the reactive intermediates (mixed

anhydride 2 and amine adduct 3). Indeed, a remarkable

decrease in stereoselectivity was observed for the reactions

involving pyridine or 2,6-lutidine as bases (Table 2).

Such results can be interpreted as a partial change of the

DYKAT into the Dynamic Thermodynamic Resolution

(DYTR) mechanism of the asymmetric induction due to

acceleration of the esterification at high concentration of

MeOH.1 In contrast, in the presence of the poorly nucleophilic

amines, the stereoselectivity under the kinetic quenching

conditions decreased only slightly.

Thus, it seems that the nucleophilic catalysis (Table 2,

entries 1 & 2) speeded up the esterification of intermediates 3

more efficiently than their epimerization, while for the base

catalysed reactions (entries 3–6 & 9) or in the presence of

highly basic amines (e.g. TEA, pKa 11.0; entries 7 & 8), the rate

of epimerization was always significantly higher than that of

esterification, even in the presence of large excess of an

alcohol. Interestingly, it seems that the behaviour of a given

amine in a kinetic quenching experiment might be exploited as

a marker of its nucleophilic properties towards H-phospho-

nates, according to the following rule of thumb: the higher the

stereoselectivity of ribonucleoside H-phosphonate conden-

sation in neat methanol, the lower the nucleophilicity of the

amine used for the reaction.

Conclusions

In the previous paper in this series we reported that stereo-

selectivity in condensations of ribonucleoside H-phosphonates

1 with alcohols originated from the Dynamic Kinetic Asym-

metric Transformation (DYKAT).1 The data presented in this

paper confirmed this conclusion and suggested that the equili-

brium between the diastereomers of nucleoside H-phosphonic—

pivalic mixed anhydride (2-DP " 2-LP) was significant

for the stereochemical outcome of the reaction only in the

absence of nucleophilic catalysis. In the presence of nucleo-

philic amines, however, the DYKAT mechanism was

governed most likely by the 3-DP " 3-LP equilibrium between

the putative P–N+ intermediates. In most instances this path

was also essential for quantitative yield of the condensation.

Pyridine derivatives (excluding those with a large steric

hindrance around the nitrogen atom) secured practically

quantitative yields of the condensations along with reasonable

high stereoselectivity (de 60–70%). Noteworthy, pyridine

derivatives with methyl groups in the a positions (e.g. 2,6-

lutidine) also provided fast, clean and highly stereoselective

condensations, and thus indicated that the esterification of

H-phosphonate monoesters in the presence of these bases

might proceed with the intermediacy of the P–N+ adducts

of type 3 (i.e. involving nucleophilic catalysis; Fig. 1 and 2). To

the best of our knowledge this would be the first documented

example of manifestation of nucleophilic properties of

2,6-dimethylpyridines in SN2(P) reactions.

For practical purposes, among investigated bases, 2,6-luti-

dine was found to be the amine of choice (quantitative yield of

condensations, high stereoselectivity, and easy availability)

Fig. 5 Diastereomeric excess (solid bars) of the DP(SP) diastereomer

of H-phosphonate diester 4b formed in the presence of mixtures of

amines, and the total yield of diester 4b (a sum of diastereomers, open

bars). Reaction conditions: 0.05 mmol of 1 (B = Ura) + EtOH

(3 equiv.) + amines (the number of molar equivalents specified on the

x axis) + PvCl (1.5 equiv.) in DCM (0.5 mL).

Table 2 Comparison of the yield and the ratio of diastereomers of the methyl uridine H-phosphonate diester 4a (Fig. 1) formed under standardand kinetic quenching conditions

Entry Amine

‘‘Standard’’ 3 equiv. of MeOH 3 equiv.of amine [1] = 100 mM

‘‘Kinetic quenching’’ 2500 equiv. of MeOH,30 equiv. of amine [1] = 10 mM

de (DP)a Yield of diester (%)a de (DP)

a Yield of diester (%)a

1 Pyridine 63 100 �10b 1002 2,6-Lutidine 68 100 12 1003 EDIPP 47 100 66 1004 DMA 56 96 42 1005 TEA 69 73 51 936 Proton sponge 49 95 45 957 Pyridine + TEA 1:1c 62 89 60 888 2,6-Lutidine + TEA 1:1c 68 80 61 919 DMA + TEA 1:1c 60 84 58 94

a Determined via integration of the corresponding 31P NMR signals. b Advantage of the LP diastereomer. c 3 + 3 equiv. or 15 + 15 equiv.


and is advised to be used in stereoselective ribonucleoside

30-H-phosphonate diesters formation.

Experimental section

Methods and materials

31P NMR spectra were recorded at 121 MHz on a Varian

Unity BB VT spectrometer. 31P NMR experiments were

carried out in 5 mm tubes using 0.5 mL of the reaction

mixture and the spectra were referenced to 2% H3PO4 in

D2O (external standard). The quantities of phosphorus-

containing compounds were determined via integration of

the corresponding 31P NMR signals. Diastereomeric excess

was calculated with accuracy of �1.5 percentage points

(an average of 3 measurements).

Commercial (Sigma-Aldrich, Alfa Aesar, Merck, POCh-

Poland) reagents and were used as purchased unless otherwise

noted. Ethanol was distilled over magnesium. Dichloro-

methane and pyridine were refluxed over P2O5, distilled, and

stored over 4 A molecular sieves until they contained below

20 ppm of water (Karl Fischer coulometric titration, Metrohm

684 KF coulometer). Anhydrous triethylamine (TEA) was

distilled and kept over CaH2. Commercial p-methoxypyridine

N-oxide (MPO) hydrate was rendered anhydrous by co-

evaporation with dry acetonitrile (1�) and dry toluene (2�).Other liquid amines were refluxed for 1 hour with 2,4,6-

triisopropylbenzenesulfonyl chloride (TPS-Cl) and distilled

under reduced pressure. 4-Ethyl-2,6-diisopropyl-3,5-dimethyl-

pyridine (EDIPP)49 was a gift from Prof. A. T. Balaban,

Texas A&M University. Uridine H-phosphonate 150 was

obtained according to the published method. Racemization

of S-(�)-nicotine was done according to ref. 51. Immediately

prior to all reactions, H-phosphonate 1 was rendered anhy-

drous by dissolving in toluene (3 mL/0.05 mmol) and evapora-

tion of this solvent under reduced pressure. After drying under

vacuum (15 min, 0.5 Torr), the flask was filled with air, dried

up by passing through Sicapent (Merck).

General procedure for condensation ofH-phosphonates of type 1

with alcohols

Nucleoside H-phosphonate 1 (0.05 mmol) was dissolved in

0.5 mL of DCM and amine (3 equiv.) and EtOH or MeOH

(3 equiv.) were added, followed by PvCl (1.5 equiv.). The

reaction mixture was transferred to an NMR tube and the 31P

NMR spectra were recorded within 1 hour.

Kinetic quenching experiments

Nucleoside H-phosphonate 1 (0.05 mmol) was dissolved in

DCM (0.3–0.5 mL), and TEA (0.2 equiv.) and PvCl

(1.2 equiv.) were added successively. The formation of the

mixed anhydride 2 was confirmed by 31P NMR spectroscopy

(dP 1.47 & 1.56) and the reaction mixture was utilized within

one hour (no degradation products were found within that

period).

The solution (0.3 mL) of the intermediate 2 (0.5 mmol;

generated as described above) was added dropwise by a

syringe to a septum-sealed flask containing vigorously stirred

methanol (5 mL) and an amine (2.4% v/v or w/v). After ca.

30 s toluene (15 mL) was added and the mixture was evapo-

rated almost to dryness under vacuum at temperature o40 1C

(such procedure was obligatory for strongly basic tertiary

amines in order to avoid transesterification52 of the product).

The oily residue was dissolved in DCM (0.5 mL) and analyzed

by 31P NMR spectroscopy.

Acknowledgements

The financial support from the Polish Ministry of Science and

Higher Education is gratefully acknowledged.

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Two polyaminophenolic fluorescent chemosensors for H+

and Zn(II).

Spectroscopic behaviour of free ligands and of their dinuclear Zn(II)

complexesw

Gianluca Ambrosi,aCristina Battelli,

aMauro Formica,

aVieri Fusi,*

aLuca Giorgi,

a

Eleonora Macedi,aMauro Micheloni,*

aRoberto Pontellini

aand Luca Prodi

b

Received (in Durham, UK) 17th June 2008, Accepted 18th September 2008


DOI: 10.1039/b810228g

The UV-Vis and fluorescence optical properties of the two polyamino-phenolic ligands

3,30-bis[N,N-bis(2-aminoethyl)aminomethyl]-2,2 0-dihydroxybiphenyl (L1) and 2,6-bis{[bis-

(2-aminoethyl)amino]methyl}phenol (L2) were investigated in aqueous solution at different pH

values as well as in the presence of Zn(II) metal ion. Both ligands show two diethylenetriamine

units separated by the 1,10-bis(2-phenol) (BPH) or the phenol (PH) for L1 and L2, respectively.

Both ligands are fluorescence-emitting systems in all fields of pH examined, with L1 showing a

higher fluorescence emission than L2. In particular, the emission of fluorescence mainly depends

on the protonation state of the phenolic functions and thus on pH. The highest emitting species

is H3L3+ for both systems, where the BPH is monodeprotonated (in L1) and the PH is in the

phenolate form (in L2). On the contrary, when BPH and PH are in their neutral form both

ligands show the lowest fluorescence, since H-bonds occurring between the phenol and the closest

tertiary amine functions decrease fluorescence. The Zn(II)-dinuclear species are also fluorescent

in the pH range where they exist; the highest emitting species being [Zn2(H�2L1)]2+ and

[Zn2(H�1L2)]3+ which are present in a wide range of pH including the physiological one.

Fluorescence experiments carried out at physiological pH highlighted that, in the case of L1, the

presence of Zn(II) ion in solution produces a simultaneous change in lem with a drop in

fluorescence due to the formation of the [Zn2(H�2L1)]2+ species, while, in the case of L2, it gives

rise to a strong CHEF effect (a twenty-fold enhancement was observed) due to the formation of

the [Zn2(H�1L2)]3+ species. These results, supported by potentiometric, 1H and 13C NMR

experiments, are of value for the design of new efficient fluorescent chemosensors for both H+

and Zn(II) ions.

Introduction

The development of chemosensors is in continuous expansion

due to their usefulness in many fields; they have a wide range

of applications, such as environmental monitoring, process

control, food and beverage analysis, medical diagnosis and

others.1–8 Due to their use in many disciplines, they are very

attractive for chemists, biologists, physicists and material

scientists. For example, in biochemistry, clinical and medical

sciences, and cell biology, freely mobile sensor molecules are

employed extensively in microscopy, offering the possibility of

performing real-space measurements.9,10

Among the different chemosensors, the fluorescence-based

ones present many advantages: fluorescence measurements are

usually very sensitive, low-cost, easily performed and versatile,

offering submicrometer spatial resolution and submillisecond

temporal resolution.11–19 The versatility of fluorescence-based

sensors originates also from the wide number of parameters

that can be tuned in order to optimize the convenient signal. In

most cases, changes in luminescence intensity represent the

most directly detectable response to target recognition; more

recently, however, other properties such as excited-state life-

time and fluorescence anisotropy have also been preferred as

diagnostic parameters, since they are less affected by environ-

mental and experimental conditions.

Phenol and poly-phenols show well known optical proper-

ties which mainly depend on their protonation degree;20,21 in

our lab, several polyamino-phenolic ligands of different topol-

ogies have been synthesized. In this study, we wanted to

extend our knowledge to the spectroscopic properties of two

of them to identify their possible applications as chemosensors

for suitable guests. In this case, we focused our attention on

the two previously synthesized amino-phenolic ligands L1 and

L2 (Chart 1). They were chosen for several reasons: they have

similar topology; they both show two diethylenetriamine

(dien) units separated by a phenolic aromatic spacer, the

1,10-bis(2-phenol) group (BPH) and the phenol for L1 and

a Institute of Chemical Sciences, University of Urbino,P.za Rinascimento 6, I-61029 Urbino, Italy

bDepartment of Chemistry, University of Bologna, Via Selmi 2,Bologna, Italy. E-mail: [email protected]

w Electronic supplementary information (ESI) available: Fig. S1:Location of acidic hydrogen atoms in the protonated species of L2.See DOI: 10.1039/b810228g



L2, respectively; in addition, they easily form dinuclear species

with transition metal ions. The molecular skeleton of both

ligands affords the formation of preorganized dinuclear

Zn(II) species where the two Zn(II) ions can cooperate in

binding guests; in particular, it has been demonstrated that

in some dinuclear species such as the [Zn2(H�2L1)]2+ and

[Zn2(H�1L2)]3+ ones, the two zinc ions show, in both systems,

an equal coordination environment, are displaced at fixed

different distances and are able to add guests to saturate the

coordination requirement of the two zinc ions (see Scheme 1).

Although zinc is an essential metal ion in human life and

plays a fundamental role in many biological functions, for

example in the alkaline phosphatase or carbonic anhydrase

enzymes,22 excess zinc can be very harmful, as it can lead to

many health problems.23 For this reason, easy recognition of

the zinc ion is key mainly and as a result many fluorescent

molecular sensors have been developed in recent years, also to

allow its in vivo mapping.24

In this work, we have studied the NMR, UV-Vis and

fluorescence properties of the free ligands as well as of their

zinc complexes in aqueous solution. The aim has been to

detect if the optical properties of these systems are affected

by pH as well as by the presence of Zn(II) in solution.

Experimental

Synthesis

Ligand 3,30-bis[N,N-bis(2-aminoethyl)aminomethyl]-2,2 0-di-

hydroxybiphenyl (L1) and 2,6-bis{[bis-(2-aminoethyl)amino]-

methyl}phenol (L2) were prepared as previously described.25

EMF measurements

Equilibrium constants for protonation and complexation

reactions with L2 were determined by pH-metric measure-

ments (pH = �log[H+]) in 0.15 M NaCl at 298.1 � 0.1 K,

using the fully automatic equipment that has already been

described; the EMF data were acquired with the PASAT

computer program.26 The combined glass electrode was cali-

brated as a hydrogen concentration probe by titrating known

amounts of HCl with CO2-free NaOH solutions and determin-

ing the equivalent point by Gran’s method,27 which gives the

standard potential E1 and the ionic product of water (pKw =

13.73(1) at 298.1 K in 0.15 M NaCl, Kw = [H+][OH�]). At

least three potentiometric titrations were performed for each

system in the pH range 2–11, using different molar ratios of

Zn(II)/L2 ranging from 1:1 to 2:1. All titrations were treated

either as single sets or as separate entities, for each system; no

significant variations were found in the values of the deter-

mined constants. The HYPERQUAD computer program was

used to process the potentiometric data.28

Spectroscopic experiments

1H and 13C NMR spectra were recorded on a Bruker Avance

200 instrument, operating at 200.13 and 50.33 MHz, respec-

tively, and equipped with a variable temperature controller.

The temperature of the NMR probe was calibrated using

1,2-ethanediol as calibration sample. For the spectra recorded

in D2O, the peak positions are reported with respect to HOD

(4.75 ppm) for 1H NMR spectra, while dioxane was used as

reference standard in 13C NMR spectra (d = 67.4 ppm).

Fluorescence spectra were recorded at 298 K with a Varian

Cary Eclipse spectrofluorimeter. UV absorption spectra were

recorded at 298 K with a Varian Cary-100 spectrophotometer

equipped with a temperature control unit.

The fluorescence quantum yields (Ff) of the highest fluores-

cent species were calculated as reported in ref. 29 using

2-aminopyridine as standard reference.


Solution studies

Ligands L1 and L2 as well as the Zn(II)/L systems were studied

by fluorescence spectroscopy in aqueous solution at different

pH values to investigate the fluorescence properties of both

ligands and how these are affected by protonation and the

presence of Zn(II) ion. 1H and 13C NMR experiments on the

free L1 as well as those reported for the Zn(II)/L1 system25a

aided in understanding the role played by both protonation

and Zn(II). The fluorescence quantum yields (Ff) of the highest

fluorescent species are reported in Table 1.

Similar 1H NMR studies carried out on L2 and Zn(II)/L2

system are reported in refs. 30 and 31, respectively. Moreover,

further studies on the UV-Vis absorption properties of both

L and Zn(II)/L systems were performed in aqueous solution

in addition to those already reported.25,30,31

Chart 1 Ligands together with labels for the NMR resonances.

Scheme 1 Coordination scheme for Zn(II) in the [Zn2(H�2L1)]2+ and

[Zn2(H�1L2)]3+ complexes.

Table 1 Fluorescence quantum yield (Ff) of the main fluorescentspecies in 0.15 mol dm�3 NaCl at 298.1 K

Ff

H3L13+ 0.34

H3L23+ 0.01

[Zn2H�2L1]2+ 0.24

[Zn2H�1L2]3+ 0.08


L1 and L2 in aqueous solution at different pH values

Basicity. The basicity of L1 in 0.15 mol dm�3 NaCl aqueous

solution at 298.1 K was potentiometrically studied and the

results obtained reported in ref. 25a; the protonation constants

of ligand L2 were determined under these ionic conditions and

the stepwise basicity constants of L2 are reported in Table 2.

The basicity of L2 is similar to that previously reported using

NMe4Cl as ionic medium thus the discussion can be outlined

in the same way.30

Fluorescence of L1 at different pH values. Emission spectra

performed at different pH values gave information on the

interaction between the dien and the 2,20-biphenol (BPH) units

and on the behavior of the ligand in its excited states.

As reported in the literature, BPH shows emission of

fluorescence depending on the degree to which it is deproto-

nated;20 in particular, it shows the most intense fluorescence in

its monodeprotonated form and the least intense in its neutral

one (more than six times lower), while the dianionic form,

although fluorescent, is obtainable only at very high pH values

(pH 4 15).20,32

The fluorescence spectra of L1 (lexc = 287 nm) recorded in

aqueous solution in the pH range 2–12 are reported in Fig. 1;

the trend of the fluorescence emission intensity (E) vs. pH

(lexc = 287 nm) is reported in Fig. 2(a) together with the

maximum absorption (� � �) and the emission (---) wavelength

trend. Fig. 2(b) reports the trend of the absorption titration at

l = 308 nm (K) together with the distribution curves for the

species of L1 (—) as a function of pH.

Excitation of L1 acid solution at pH 2 (lexc = 287 nm) gives

rise to a fluorescence emission band of very low intensity

(lem = 403 nm) attributed to the BPH fluorophore. The

intensity of the fluorescence emission of the compound is

highly dependent on the protonation state of the ligand

(see Fig. 1 and 2(a)); however the shape and the lem of the

spectra are substantially pH-independent. In this pH range,

the free BPH group shows a similar fluorescent behavior

produced by the monoanionic excited state of BPH.20,32

Taking into account that the fluorescence of L1 is due to the

BPH fluorophore, this suggests that also in L1 the changes in

fluorescence emission reflect only the ground states acid–base

equilibrium.33 For this reason, no indication of the excited

state proton transfer reaction was found and, as reported for

free BPH, the fluorescence is due to the monoanionic excited

state of BPH in L1.

In the fluorescence spectra, the emission remains substan-

tially very low and constant (Fig. 2(a)) at acidic pH values

(2 r pH r 5) while it starts increasing at pH 5 in concomi-

tance with the appearance of the H3L13+ species in solution,

reaching a maximum intensity at pH 7.4–8.4 with the complete

formation of the H3L13+ species. A small decrease can be

Table 2 Basicity and equilibrium constants for the complexationreactions of L2 with Zn(II) ion determined in 0.15 mol dm�3 NaCl at298.1 K

Reaction logK

L + H+ = HL+ 10.04(1)a

HL+ + H+ = H2L2+ 9.87(1)

H2L2+ + H+ = H3L

3+ 9.12(1)H3L

3+ + H+ = H4L4+ 7.59(1)

H4L4+ + H+ = H5L

5+ 2.50(3)Zn2+ + L + 2H+ = ZnH2L

4+ 28.42(1)Zn2+ + L + H+ = ZnHL3+ 23.82(2)Zn2+ + L = ZnL2+ 14.67(2)Zn2+ + L = Zn(H�1L)

+ + H+ 5.05(2)2Zn2+ + L = Zn2(H�1L)

3+ + H+ 17.17(1)2Zn2+ + L + H2O = Zn2(H�1L)OH2+ + 2H+ 8.34(3)2Zn2+ + L + 2H2O = Zn2(H�1L)(OH)2

+ + 3H+ �1.63(3)Zn2(H�1L)

3+ + OH� = Zn2(H�1L)OH2+ 4.90Zn2(H�1L)OH2+ + OH� = Zn2(H�1L)(OH)2

+ 3.76

a Values in parentheses are the standard deviations on the last

significant figure.

Fig. 1 Fluorescence spectra of L1 at different pH values.

Fig. 2 Fluorescent emission titration (lexc = 287 nm, lem = 403 nm)

(E), absorption wavelength trend (� � �), and emission wavelength trend

(---) (a); absorption titration at l = 308 nm (K) and distribution

curves of the species (—) (b) as function of pH in aqueous solution:

[L1] = 5.0 � 10�5 M, I = 0.15 M NaCl, T = 298.1 K.


observed in the alkaline range up to pH 10.5 at which the less

protonated species appear in solution; on the contrary, when

the pH is increased to 11 emission rises again, reaching

maximum intensity at pH 12 with the presence in solution of

the monoanionic H�1L1� species.

Bearing in mind the previous studies on BPH,20,32 the trend

of the emission intensity in the range of pH 2–8 can be easily

explained by the deprotonation of the neutral BPH unit to

form its monoanionic species that occurs with the formation of

the H3L13+ species. In other words, in the protonated species

H5L15+ and H4L1

4+, BPH is present in its neutral form while

in the H3L13+ species it has lost one of the acidic protons

forming the highest emitting species (F = 0.34, Table 1).

These results are in agreement with those already obtained by

UV-Vis absorption studies which revealed that the deproto-

nation of one of the hydroxyl functions of BPH occurred in

the pH range involving the passage from H4L14+ to H3L1

3+

species.25a This was highlighted, as reported in Fig. 2(b), by

the change in absorption at 308 nm which increases when the

monoanionic form of BPH is present in solution and is further

underlined by the variation in the trend of the maximum of the

absorption wavelength (Fig. 2(a)) as a function of pH; both

figures highlight that the changes take place in the field of pH

where the H3L13+ species forms. Although the absorption and

emission wavelength maxima as well as the absorption at

308 nm remain constant, increasing the pH to form lesser

protonated species than H3L13+, there is a small decrease

(about 25% at pH 10) in fluorescence intensity occurring with

the formation of the H2L12+, HL1

+ and neutral L1 species

(Fig. 2(a)). This trend could be explained by the formation of a

H-bond network involving BPH and the closer nitrogen

atoms. In fact, as reported for free BPH,20,32 the formation

of an intramolecular H-bond interaction occurring between

the two oxygen atoms of BPH in stabilizing the hydrogen

atom in the monoanionic species gives the greatest fluores-

cence intensity, while, on the contrary, the formation of

intermolecular H-bonds with H-accepting molecules, such as

water, gives rise to a very fast nonradiating process through-

out the H-bond, thus leading to a decrease in the fluorescence

(this occurs for example in the neutral form of excited BPH).

In addition, it has been demonstrated that in the presence of

strong proton-accepting molecules such as triethylamine

(TEA), the formation of H-bonding between TEA and a

hydrogen atom of BPH once again leads to a decrease in

fluorescence.20b In our case, it is presumable that similar

H-bonding between the BPH oxygen and the closer nitrogen

atoms of the dien units is formed, thus decreasing fluorescence;

however, the formation of this type of H-bond cannot be the

favourite situation since only a slight drop in fluorescence was

observed. Moreover, two different H-bonds could be sug-

gested in the case of L1: via OH� � �N as well as via O�� HN+;

in other words, in the H2L12+, HL1

+ and L1 species, a partial

stabilization of the acidic hydrogen atom of the monoanionic

BPH unit could also take place with the closer N atom

(c in Scheme 2), but also a partial ammonium character of

the closer N atom could give rise to the same quenching

H-bond effect with the BPH unit (b in Scheme 2). In any case,

the form (a) shown in Scheme 2 is the favoured form and it is

the only one present in the H3L13+ as well as in the H�1L1

�

species where the highest fluorescence is reached. In addition,

the absence of fluorescence changes even at highly alkaline pH

values once again demonstrates that the full deprotonation of

BPH in L1 is not reachable under our experimental conditions.

L11H and

13C NMR studies at different pH values. In order

to obtain further structural information about the distribution

of acidic protons in the protonated species of L1, 1H and 13C

NMR spectra were recorded over the pH range of the

Scheme 2 Possible H-bond interactions for the neutral L1 species.

Fig. 3 Experimental NMR chemical shifts in aqueous solution of L1

as a function of pH: 1H NMR (a); 13C NMR (b).


potentiometric, UV-Vis and fluorescence measurements.1H–1H and 1H–13C NMR 2D correlation experiments were

performed to assign all the signals. The trends for the chemical

shift of the 1H and 13C NMR resonances are reported in

Fig. 3(a) and (b). The 1H NMR spectrum recorded at pH 12,

where the H�1L1� species is prevalent in solution, exhibits two

triplets at 2.67 and 2.85 ppm corresponding to the resonance

of the hydrogen atoms H2 and H1, respectively, one singlet at

3.76 ppm due to the hydrogen atoms H3, one triplet at

6.91 ppm for the resonances of H6 and two doublets at 7.31 and

7.42 ppm for H5 and H7, respectively. This spectral feature

indicates a C2v symmetry mediated on the NMR time-scale

which is preserved throughout the pH range investigated. In

agreement with this symmetry, the 13C NMR spectrum

recorded at the same pH value shows only nine signals at d37.7 (C1), 53.0 (C3), 55.0 (C2), 117.8 (C6), 126.6 (C4), 129.7

(C8), 130.5 (C5), 130.6 (C7) and 158.0 (C9). At lower pH,

where the species L1, HL1+, H2L1

2+ and H3L13+ are form-

ing (pH= 11–7), the main shift is exhibited by the protons H1

which show a marked downfield shift, suggesting that the four

protonation steps take place mainly on the primary amine

functions. This hypothesis is confirmed by the trend of the 13C

NMR resonances which mainly shows an upfield shift in the

signal of the carbon atom C2, in agreement with the b-effect ofthe protonation of the polyamines.34 However, in this pH

range, slight shifts in other 1H NMR resonances could be seen:

for example, the resonance of H7 first moves downfield up to

pH 9 then decreases with the formation of the H3L13+ species,

while H3 moves upfield; this suggests little changes in charge

density on both the tertiary amine groups and BPH unit

occurring in this pH range that can be correlated with the

formation of H-bonding involving the BPH oxygen and the

closer nitrogen atoms in the L1, HL1+ and H2L12+ species, in

agreement with the fluorescence experiments reported above.

In the pH range 4–6 the H4L4+ species is prevalent and, as

demonstrated both by UV and fluorescence experiments, the

fifth protonation step occurs at the BPH group. This was also

confirmed in the NMR experiments by the downfield shift of

the H6 signal in the para position to the phenolic oxygens and

by the upfield shift of the H7 protons in the 1H NMR spectra,

as well as by the accompanying upfield shift of C4, C8 and C9

and downfield shift of C6 in the 13C NMR spectra. The strong

upfield shift exhibited by the signal of H7 could be related not

only to a protonation process of the BPH unit but also to a

change in the angle between the two aromatic rings that

probably is affected by the protonation degree of L1 leading

the formation of a new H-bond network involving the neutral

BPH and the unprotonated tertiary amine functions, as

depicted in Fig. 4 for the H4L14+ species; this almost entirely

quenches the fluorescence (see above). The protonation step

giving the H5L5+ species, occurring below pH 4, basically

causes a downfield shift of protons H2 and H3 together with

an upfield shift in the signals of the carbon atoms C1 and C4,

suggesting that it takes place on the tertiary amine groups.

Once again the H7 and H5 resonances, both of which shift

downfield, are perturbed by this protonation step, highlighting

the formation of a H-bond network with the closer amine

functions on the BPH unit different from the previous one; this

Fig. 4 Location of acidic hydrogen atoms in the protonated species

of L1.

Fig. 5 Fluorescence emission titration (lex = 280 nm) (E), absorp-

tion wavelength trend (� � �), and emission wavelength trend (---) (a);

absorption titration at l = 290 nm (K) and distribution curves of the

species (—) (b); of L2 as function of pH in aqueous solution: [L2] =

5.0 � 10�5 M, I = 0.15 M NaCl, T = 298.1 K.


affects the chemical shift, modifying the electron density of the

neutral BPH unit as well as the angle between the two

aromatic rings.

A protonation scheme arising from NMR experiments is

summarized in Fig. 4.

Fluorescence of L2 at different pH values. The same fluores-

cence experiments were carried out on the ligand L2

and compared to the previous UV-Vis and NMR studies

performed in aqueous solution at different pH values.30 The

trend in fluorescence emission intensity (E) versus pH (lexc =280 nm) is reported in Fig. 5(a) together with the maximum

absorption (� � �) and emission (---) wavelength trends. Fig. 5(b)

reports the trend for the absorption titration at l = 290 nm

(K) together with the distribution curves of the species of L2

(—) as a function of pH obtained by potentiometry.

The acidic solutions of L2 up to pH 6 are barely fluorescent,

as also reported for the free neutral phenol (PH), while

fluorescence increases with the formation of the H3L23+

species, reaching maximum emission at pH 8 together with

the maximum presence in solution of the H3L23+ species; at

higher pH values, the emission drops, reaching a plateau at pH

higher than 11 with the formation of the neutral L2 species.

The fluorescence of ligand L2, which is lower compared to that

of L1 (see Table 1), is highly dependent on the protonation

state of the ligand as seen before for L1. The acidic proton

distribution in the several protonated species of L2 obtained

by UV-Vis, potentiometry and NMR studies was previously

reported and the scheme is reported in Fig. S1 of the ESI;w the

most fluorescent H3L23+ species is the one in which the phenol

is deprotonated (i.e. phenolate) and the four acidic protons are

located on the primary amine functions; this is the same

situation found for the H3L13+ species where there are no

H-bond interactions with the closest amine functions, thus

affording the highest emission quantum yield also in the

H3L23+ species. It should be noted that in the free PH, the

anion presents a much lower fluorescence intensity than the

neutral species.21 In this case the opposite behaviour was

observed; this could be explained (see also below) by a

decrease in the solvation via H-bond network of the phenolate

oxygen atom by the water molecules in the H3L23+ species in

comparison with the free PH anion.21a In other words, the

presence of the two protonated dien units linked to the PH

group modifies the accessibility of the solvent molecules to the

phenolate oxygen atom decreasing its quenching effect and

thus increasing the emissive relaxation decay of the PH anion.

As reported, an acidic proton redistribution was observed in

the less protonated species involving at least a tertiary amine

function that becomes protonated. This ammonium group,

found mainly in the neutral L2 species, is stabilized viaH-bond

with the close phenolate oxygen atom (see Fig. S1, ESIw). Forthis reason, as for ligand L1, the formation of H-bonding with

the amine function leads to a decrease in its fluorescence.

This H-bond interaction, which is also monitorable through

the UV-Vis spectra (see Fig. 5(b)), is also highlighted by the

change in the maximum of the absorption and emission

wavelengths (Fig. 5(a)) as a function of pH. lmax and lemshifted in different directions, increasing the Stokes shift when

the phenol becomes phenolate (pH Z 5, lmax and lem shift

towards lower and higher energies, respectively), while an

opposite trend was observed at higher pH values (lmax and

lem shift towards higher and lower energies, respectively) with

the formation of the neutral zwitterionic L2 species in which a

strong H-bond between the closest tertiary ammonium and

phenolate groups was suggested. Taking into account the

trend and shift in both lmax and lem, it can be suggested that

the fluorescence is yielded by light emission decay from the

phenolate excited state of all L2 species to different ground

states, characterized by the formation of strong intramolecular

H-bonds.

In conclusion, although L1 is a much more efficient fluores-

cent system than L2 (Table 1), both ligands show fluorescence

emission depending on the protonation state of the aromatic

functions. In particular, the highest emitting species are due to

the monodeprotonated form of BPH of L1 as well as to the

phenolate species of PH of L2, both of which are achieved in

the H3L3+ species; on the contrary, the neutral BPH and PH

species are very low fluorescence emitters. The presence of the

closer tertiary amine function affects the emission quantum

yield in some species by forming intramolecular H-bonding

with the close phenol oxygen atom of both systems. The

H-bonding induces a nonradiative relaxation process of

the excited species, yielding a decrease in the fluorescence in

both ligands. This H-bonding is weaker in L1 via OH� � �N as

well as O�� HN+, and for this reason only a relatively low

efficiency of fluorescence quenching could be observed, while it

takes place strongly via O�� HN+ in L2 giving an almost

total quenching of the fluorescence of L2. Taking into account

these results, both ligands behave as chemosensors of H+ in

that they are able to change their optical absorption and

fluorescence properties as a function of pH.

Coordination of Zn(II)

The coordination behaviour of both systems towards Zn(II)

was potentiometrically studied and the results obtained are

reported in ref. 25 and 31; as for basicity, the Zn(II)/L2 system

had been studied in NMe4Cl ionic medium,31 thus we per-

formed new potentiometric measurements to obtain the stabi-

lity constants for the Zn(II)/L2 system under the same

experimental conditions as the Zn(II)/L1 system (0.15 mol dm�3

NaCl aqueous solution at 298.1 K). The potentiometrically

determined stability constants for the equilibrium reactions of

L2 with Zn(II) are reported in Table 2. The species formed as

well as the values of the stability constants evaluated are

similar to those previously reported and thus the discussion

can be outlined in the same way. The main difference found

was the formation of the [Zn2(H�1L2)(OH)2]+ species in this

ionic medium which was not previously detected. The addition

of the second OH� anion to [Zn2(H�1L2)OH]2+ is quite high

(logK = 3.76) suggesting that it is probably bound in a bridge

disposition between the Zn(II) ions. The distribution diagrams

for the Zn(II)-complexed species for both 2Zn(II)/L systems are

reported in Fig. 6 for L1 and in Fig. 8 for L2 as a function of

pH. However, the results previously discussed can be summar-

ized in this way: (i) the dinuclear species are prevalent in

solution and the only species existing at pH higher than 7 is a

L/Zn(II) with a 1:2 molar ratio; (ii) the most prevalent species


are [Zn2(H�2L1)]2+ and [Zn2(H�1L2)]

3+ for L1 and L2,

respectively; (iii) these dinuclear species have similar molecular

skeletons indicating a preorganized dinuclear Zn(II) species in

which the two Zn(II), similarly coordinated, can cooperate in

binding suitable guests (see Scheme 1).

Fluorescence and UV-Vis of the 2Zn(II)/L1 system at differ-

ent pH values. Emission and absorption spectra were per-

formed at different pH values using Zn(II)/L1 at a 2 to 1

molar ratio. The trend in fluorescence emission intensity (E)

versus pH (lexc = 283 nm) is reported in Fig. 6(a) together

with the maximum absorption (� � �) and emission (---) wave-

length trends. Fig. 6(b) reports the trend for the absorption

titration at l = 295 nm (K) together with the distribution

curves for the species of the 2Zn(II)/L1 system (—) as a

function of pH. Moreover, fluorescence titration was carried

out by adding increasing amounts of Zn(II) to a HEPES buffer

(pH = 7.4) solution of L1 and the spectra are reported in

Fig. 7. The fluorescence quantum yield of the highly emitting

species is reported in Table 1.

The UV-Vis absorption spectra of solutions containing

2Zn(II)/L1 recorded at different pH values were discussed

previously;25a they showed spectral profiles indicating the

deprotonation of BPH and simultaneous coordination of the

Zn(II) ions; in these new experiments, some further aspects can

be discussed. The absorption lmax shifts toward lower energy

when monitored from acidic (free ligand) to basic pH values

(Zn(II)-complexes); up to the presence in solution of the Zn(II)-

mononuclear species it moves from 280 to 305 nm, while the

appearance in solution of the dinuclear species, at approxi-

mately pH 6, gives rise to a change in the lmax which shifts

from 305 nm in the presence of the mononuclear [Zn(HL1)]3+

species to 298 nm with the complete formation of the more

stable [Zn2(H�2L1)]2+ species at pH = 7.4. This lmax is

preserved also at higher pH values where only dinuclear

Zn(II)-complexed species are present in solution. The shift in

lmax observed from the mono- to the di-nuclear species can be

ascribed to the full deprotonation of BPH which loses both

acidic hydrogen atoms in the Zn(II)-dinuclear species, afford-

ing the bi-negative form of BPH; this result is in agreement

with the studies previously reported for the Zn(II)-dinuclear

species of L1. The changes in absorption from the mono-

negative BPH to the bi-negative species are also visible in

Fig. 6(b), where a change in absorptivity can also be observed

when the dinuclear species appear in solution. These changes

are in agreement with a change in the protonation degree of

BPH and thus to its full deprotonation and simultaneous

coordination of each Zn(II) ion by one phenolate oxygen

atom of the BPH unit as already reported. The fluorescence

experiments gave rise to analogous results, with fluorescence

increasing at values starting from acidic pH and reaching

maximum intensity in the field of pH 7.4–8.4 with the maxi-

mum presence in solution of the [Zn2(H�2L1)]2+ species, then

decreasing at higher pH values and reaching a plateau at pH

4 11 with the presence in solution of the di-hydroxylated

[Zn2(H�2L1)(OH)2] species (Fig. 6). It is interesting to note

that, unlike the free L1, the lem changes (lexc = 283 nm) by

changing the pH, and as in the absorption experiments

the change occurs at the pH values where there is the forma-

tion of the Zn(II)-dinuclear species. Specifically, lem shifts from

403 nm (free ligand) to 379 nm with the formation of the

[Zn2(H�2L1)]2+ species, while remaining constant in the other

dinuclear species. Once again, this trend can be related to the

full deprotonation of BPH, as retrieved in the crystal structure

of the [Zn2(H�2L1)(H2O)2]2+ previously reported, which

produces changes in the ground as well as in the excited state

of BPH. Moreover, the formation of the hydroxylated

[Zn2(H�2L1)OH]+ and [Zn2(H�2L1)(OH)2] species produces

a drop in fluorescence emission without changing the lem(Fig. 6(a)); this is due to an increase in electron density




species (–) (b); as a function of pH in aqueous solution: [L1] = 5.0 �10�5 M, [Zn(II)] = 10�4 M, I = 0.15 M NaCl, T = 298.1 K.

Fig. 7 Fluorescence spectra of the Zn(II)/L1 system in aqueous buffer

(HEPES, 5 � 10�2 M) solution at pH = 7.4, obtained by adding

several amounts of Zn(II) up to 2 equivalents with respect to [L1] =

5.0 � 10�5 M.


of the BPH unit by coordinating the OH� species which,

as reported in similar cases, increases a thermal relaxation

negatively affecting emission decay mechanisms.20b The

change in lem occurring with the formation of the dinuclear

[Zn2(H�2L1)]2+ is well highlighted by titrating a buffer

(pH = 7.4) solution of L1, adding increasing amounts of

Zn(II) up to 2 equivalents (Fig. 7); as shown in the figure, the

lem shifts toward higher energy by adding Zn(II) but,

at the same time, the fluorescence of the new species

formed decreases by about 30% and thus none chelation-

enhanced fluorescence (CHEF) effects were observed for

this system.

Fluorescence and UV-Vis of the 2Zn(II)/L2 system at different

pH values. Analogous fluorescence and absorption experi-

ments were performed at different pH values using Zn(II)/L2

at a 2 to 1 molar ratio; the results are reported in Fig. 8.

UV-Vis absorption spectra of solutions containing 2Zn(II)/

L2 at different pH values show, as previously reported,

spectral profiles due to the deprotonated form of PH. How-

ever, also in this case, some further aspects can be discussed.

Observing the lmax of the spectra from acidic to alkaline pH

values (Fig. 8(a)), a shift in the lmax from 273 (free ligand) to

286 nm (complexed ligand) occurs at pH 4 5 with the

appearance in solution of the [Zn2(H�1L2)]3+ species, in

agreement with the deprotonation of PH as previously demon-

strated. The value of lmax 286 nm is enough preserved also at

higher pH values where only a little decrease is shown with the

appearance in solution of the dihydroxylated species (lmax =

282 at pH = 12). On the contrary, in the 6–8 pH range, where

the [Zn2(H�1L2)]3+ species is prevalent in solution, a slight

increase in absorption with respect to the free ligand can

be observed, while a marked increase is visible at higher

pH values with the formation of the hydroxylated

[Zn2(H�1L2)OH]2+ species (see Fig. 8). This finding could

be explained by a different disposition of L2 in forming the

Zn–O–Zn cluster system (O is the phenolate oxygen atom) in

the [Zn2(H�1L2)]3+ and [Zn2(H�1L2)OH]2+ species. In fact,

while it was demonstrated that the hydroxylated

[Zn2(H�1L2)OH]2+ species shows the OH� displaced in a

bridged disposition between the two Zn(II) ions,30 on the

contrary, a coordination environment without secondary brid-

ging ligands could be hypothesized in the [Zn2(H�1L2)]3+

species. In the latter, the fifth coordination site of each Zn(II)

ion could be saturated by a water molecule or by a chloride

anion of the ionic medium. This may be the reason for the

increase in absorption of the [Zn2(H�1L2)OH]2+ with respect

to the [Zn2(H�1L2)]3+ species.

Analysis of the fluorescence experiments gives additional

information; examining the maximum of lem (lexc = 275 nm)

from acidic to alkaline field of pH, a shift of the lem is

observable (see Fig. 8) at pH 4 5; lem moves from 354 nm,

typical of the free ligand, reaching a constant value (308 nm) at

pH 6, with the full formation of the [Zn2(H�1L2)]3+ species.

This change in lem is coupled with an increase in fluorescence,

which shows its highest emission in the range of the

[Zn2(H�1L2)]3+ species. These changes are in agreement with

the simultaneous deprotonation of the phenolic oxygen atom

due to the Zn(II) complex formation and its bridging coordi-

nation between the two Zn(II) ions, as phenolate. At pH4 9, a

further change in the lem can be highlighted, since it shifts

from 308 to 325 nm in concomitance with the appearance of

the [Zn2(H�1L2)OH]2+ species in solution; this occurs without

observing any significant change in fluorescence intensity. This

result may be related, as above, to a different disposition of

the secondary ligands in the two complexed [Zn2(H�1L2)]3+

and [Zn2(H�1L2)OH]2+ species that could be responsible

of the different lem in the dinuclear [Zn2(H�1L2)]3+ and

[Zn2(H�1L2)OH]2+ species. As previously discussed for

the [Zn2(H�2L1)(OH)2] species, the increase in the total

electron density of the complex in the dihydroxylated

[Zn2(H�1L2)(OH)2]+ species affects fluorescence at higher

pH values.

The Zn(II)-L2 dinuclear complexes showed very interesting

fluorescent properties; in fact, although free L2 exhibits emit-

ting species in the same range of pH of the Zn(II)-dinuclear

one, the fluorescence intensity of the latter is higher, giving a

strong CHEF effect. This effect, occurring to L2 in the

presence of Zn(II), is highlighted in Fig. 9, which reports the

fluorescence spectra of L2 obtained by adding several amounts

of Zn(II) in aqueous buffer pH = 7.4 solution. At this pH

value, the species formed in the presence of Zn(II) is the

[Zn2(H�1L2)]3+ species. As can be observed in Fig. 9, the free

ligand shows low fluorescence emission with a lem centered

at 347 nm; by adding Zn(II), the emission increases and lemshifts toward higher energy. The spectra preserve the same

profile when adding up to 2 equivalents of Zn(II), reaching a

constant emission and lem of 308 nm, in concomitance

with the complete formation of the [Zn2(H�1L2)]3+ species.




species (—) (b); as a function of pH in aqueous solution: [L2] = 5.0 �10�5 M, [Zn(II)] = 1.0 � 10�4 M, I = 0.15 M NaCl, T = 298.1 K.


(see Fig. 8(b)). At this pH the emission quantum yield is more

than twenty-fold higher for the [Zn2(H�1L2)]3+ species than

for the free ligand; furthermore, a similar CHEF effect was

also found in other fields of pH such as 8 and 10, at which the

dinuclear species are formed, thus highlighting the sensing role

of L2 towards Zn(II) in aqueous solution in a biologically

important range of pH. For this reason, L2 can be considered

a potential chemosensor for Zn(II).

Conclusions

The studies highlighted that the intensity of the fluorescence of

both ligands depends on the protonation state of the phenolic

functions, and in the case of L2, the lem is also affected by

protonation while this does not occur for L1. For this beha-

vior, both ligands are suitable chemosensors of H+ in that

they are able to change their optical absorption and fluores-

cence properties as function of pH.

For both systems the most fluorescent species is the same:

H3L3+ in which the BPH unit of L1 is in its mono-deproto-

nated form, while PH is present as phenolate form in L2; on

the contrary, when BPH and PH are in their neutral form both

ligands show the lowest fluorescence. While these results are in

agreement with those found for free BPH (the more fluores-

cent species is the monoanionic species of BPH), this finding is

opposite to that for free PH where the neutral species is the

most fluorescent. This can be explained by a lower solvation of

the phenolate oxygen atom in the H3L23+ species which limits

the quenching effect occurring via H-bond with the water

molecules. The presence of the tertiary amine function close

to the phenol oxygen affects the emission quantum yield of

those species in which the formation of intramolecular

H-bonds is possible, highlighting that the formation of

H-bonds has a quenching effect to the fluorescence in these

systems.

The Zn(II)-dinuclear species are fluorescent in the field of

pH where they exist; the highest emitting species are the

[Zn2(H�2L1)]2+ and the [Zn2(H�1L2)]

3+ species, respectively;

they are prevalent in a wide range of pH including the

physiological one. In the [Zn2H�2L1]2+, the presence of the

dianionic form of BPH produces a blue shift of lem in the

fluorescence experiments, in comparison with the free ligand.

The interaction with guests such as OH� perturbs the emission

but not the absorption of the dinuclear species.

In the [Zn2(H�1L2)]3+ species a slight blue shift in lem can

also be observed, as well as, a decrease in fluorescence brought

about by the addition of an anionic guest such as OH�.

The main result retrieved is that both L1 and L2 sense the

Zn(II) in aqueous solution at physiological pH 7.4 by fluores-

cence; at this pH the [Zn2(H�2L1)]2+ and [Zn2(H�1L2)]

3+

species are prevalent in solution. The [Zn2(H�2L1)]2+ species

shows a simultaneous change in the lem with a drop in

fluorescence, but real and efficient sensing was obtained by

using ligand L2 which, in the presence of two equivalents of

Zn(II), gives rise to a strong CHEF effect (a twenty-fold

increase) with the formation of the [Zn2(H�1L2)]3+ species;

in this case, a similar CHEF effect was also found in other

fields of pH such as 8 and 10, highlighting the sensing role of

L2 towards Zn(II) in aqueous solution in a biologically im-

portant range of pH.

Concluding, both systems behave as chemosensors for both

H+ and Zn(II) and their investigation has given much useful

information for the design of more efficient systems. More-

over, taking into account that both [Zn2(H�2L1)]2+ and

[Zn2(H�1L2)]3+ dinuclear species show the highest fluores-

cence intensity and that they are the most suitable hosting

species for guests, they are a very interesting platform for the

sensing of guest species.

Acknowledgements

The authors thank the Italian Ministero dell’Istruzione

dell’Universita e della Ricerca (MIUR), PRIN2007 for finan-

cial support.

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Dynamic covalent self-assembled macrocycles prepared from 2-formyl-

aryl-boronic acids and 1,2-amino alcoholsw

Ewan Galbraith, Andrew M. Kelly, John S. Fossey, Gabriele Kociok-Kohn,

Matthew G. Davidson, Steven D. Bull* and Tony D. James*

Received (in Durham, UK) 2nd September 2008, Accepted 18th September 2008


DOI: 10.1039/b815138e

Reaction of 2-formyl-aryl-boronic acids with 1,2-amino alcohols results in dynamic covalent self

assembly to quantitatively afford tetracyclic macrocyclic Schiff base boracycles containing

bridging boron–oxygen–boron functionality.

Introduction

The development of boronic acid based saccharide sensors

that rely on the dynamic covalent interaction of boronic acids

with diols has been widely investigated.1–9 Boronic ester

formation with diols has also been used for the construction

of discrete macrocycles and cages.10 The reversible nature

of boronic acid complexation with diols makes this type of

interaction highly suitable for the reversible self-assembly of

multicomponent systems. With these types of reversible

systems any errors that occur during the assembly process may

be corrected because equilibration of the reactive species

results in formation of a thermodynamically favoured

product. A number of boracycles have been prepared

that employ a combination of facile imine formation and

boronic acid esterification to afford multicomponent macro-

cycles.11–24 For example, Severin has prepared a series

of self assembled macrocycles/cages by combining 3- or 4-for-

myl-phenyl-boronic acids with bis or tris primary amines

and pentaerythritol (tetraol).25,26 Nitschke has also prepared

a macrocycle derived from pentaerythritol, 2-formyl-phenyl-

boronic acid and para-diaminobenzene, as well as a cage

compound arising from self assembly of cyclotricatechylene,

meta-xylylenediamine and 2-formyl-phenyl-boronic acid.27

Farfan has prepared boracycles from boric acid, 4-diethyla-

mino salicylaldehyde and (R)-phenylglycinol 6d. This

complex was formed in two steps involving reaction of

4-diethylamino salicylaldehyde and (R)-6d to produce an

imine, followed by reflux with boric acid in toluene under

Dean–Stark conditions for 18 h to produce the observed

complex.28

We now report herein that simple room temperature mixing

of 2-formyl-aryl-boronic acids with 1,2-amino alcohols results

in dynamic covalent self assembly to afford stable tetracyclic

macrocyclic Schiff base complexes that contain a rigid brid-

ging boron–oxygen–boron functionality.


We have recently reported the development of versatile three-

component derivatization protocols for determining the

enantiomeric excess of chiral primary amines, diols or

diamines.29–35 For the case of amines, this approach involves

derivatization of a chiral amine 1 with 2-formyl-phenyl-

boronic acid 2 and enantiopure BINOL (S)-3 in CDCl3 to

quantitatively afford a mixture of diastereoisomeric imino-

boronate esters (S,S)-4 and (S,R)-5. The diastereoisomeric

ratio of (S,S)-4:(S,R)-5 is then determined by 1H NMR

spectroscopic analysis, and since no kinetic resolution occurs

this value is an accurate reflection of the enantiomeric excess

of the parent amine (Scheme 1).

We reasoned that this type of three-component derivatiza-

tion protocol might also be useful for analyzing the enantio-

purity of chiral 1,2-amino alcohols. Therefore, (S)-leucinol 6b

was treated with 2-formyl-phenyl-boronic acid 2 and (S)-

BINOL 3 in CDCl3 and its 1H NMR spectrum acquired after

ten minutes. The resultant 1H NMR spectra revealed the

presence of a complicated mixture of interconverting products

Scheme 1 Three-component protocol for determining the enantio-

meric purity of chiral amines by 1H NMR spectroscopic analysis.

Department of Chemistry, University of Bath, Bath, UK BA2 7AY.E-mail: [email protected]. E-mail: [email protected];Tel: +44 1225 383810w CCDC reference numbers 694358–694361 [(S,2R,4S)-7, 8a, 8f and8h]. For crystallographic data in CIF or other electronic format seeDOI: 10.1039/b815138e



that was clearly unsuited for carrying out ee determination.

However, on standing overnight, the crude reaction product

fractionally crystallised to afford the expected oxazolidine-

boronate ester (S,2R,4S)-7, whose structure was subsequently

confirmed by X-ray crystallographic analysis (Scheme 2).

In order to investigate this complexation reaction further, it

was decided to determine what products would be formed

when 2-formyl-phenyl-boronic acid 2 was individually reacted

with either (S)-BINOL 3 or (R)-valinol 6a. Two-component

mixing of 2-formyl-phenyl-boronic acid 2 with (S)-BINOL 3

in CDCl3 resulted in no reaction occurring. However, reaction

of 2 with (R)-valinol 6a at room temperature in chloroform

resulted in exclusive formation of a new boracycle (R,R)-8a in

quantitative yield (Scheme 3). The structure of symmetrical

boracycle (R,R)-8a was confirmed by X-ray crystallographic

analysis (Fig. 1), which revealed it to be the condensation

product of two equivalents of 2-formyl-phenyl-boronic acid

2 with two equivalents of (R)-valinol 6a, with concomitant

elimination of five molecules of water. This complexation

reaction results in formation of the densely packed central

core of boracycle (R,R)-8a which comprises two fused seven

membered rings formed from two tetrahedral sp3-boron

atoms, two imino alcohol fragments, and a central oxygen

atom that bridges both boron atoms. This architecture results

in its central fused bicyclic ring structure being further

appended by two five-membered rings formed from two

imino-boronate ester linkages that confer sp3 character on

the boron atoms. The scope and limitation of this four-

component condensation reaction was then investigated via

treatment of a series of five chiral amino alcohols 6b–f with

2-formyl-phenyl-boronic acid 2, which resulted in clean for-

mation of their respective boracycles 8b–f in 84–96% isolated

yield (Scheme 3).

The reversible nature of macrocycle formation of these

boracycles 8a–f was confirmed by adding one equivalent of

amino alcohol (S)-6a to macrocycle (S,S)-8b in chloroform.

Mass spectrometry indicated that this solution now contained

a mixture of three macrocycles, (S,S)-8a (M + H 431 m/z),

(S,S)-8b (M + H 445 m/z) and a mixed macrocycle derived

from (S)-6a and (S)-6b (M + H 417 m/z) in a statistical

1:1:2 ratio.

Norman and coworkers have previously reported the synth-

esis of achiral boracycle 8g derived from condensation of

2-aminophenol with 2-formyl-phenyl-boronic acid 2 in

ethanol at reflux.36 Attempts to repeat this condensation

reaction using our mild complexation conditions at room

temperature resulted in no reaction occurring. However,

heating 2-aminophenol 6g (or 4-methyl-2-aminophenol 6h)

with 2-formyl-phenyl-boronic acid 2 at reflux in 95:5 ethanol:

benzene under Dean–Stark conditions did result in quantita-

tive formation of the boracycles 8g (or 8h). Comparison of the

X-ray crystal structures of boracycle (S)-8a with that of

boracycle 8h (Fig. 2) revealed that whilst they belong to the

Scheme 2 Formation and X-ray crystal structure of boronic ester

(S,2R,4S)-7.

Scheme 3 Condensation of 2-formyl-phenylboronic acid 2 with chiral

amino alcohols 6a–f and achiral amino alcohols 6g, h affords four-

component boracycles 8a–h.

Fig. 1 Crystal structure of macrocycle 8a. (a) Viewed along the

boron–boron axis. (b) Viewed perpendicular to the boron–boron axis.


same class of bridging boracycle, their three dimensional

architectures are very different. In the case of boracycle 8a,

the central bridging oxygen atom lies on the opposite side to

the other two oxygen atoms about the plane bisected by the

two boron atoms. This results in the alkyl side-chains of their

amino alcohol fragments adopting a conformation that creates

the walls of a potential binding cavity centred around its

bridging oxygen atom, with its aryl rings acting as buttressing

elements to contribute structural rigidity. Conversely, for the

case of macrocycle 8h, the presence of the more rigid amino-

phenol fragments results in the three oxygen atoms now being

presented on the same face of the plane bisected by the boron

atoms. This, in turn, results in the aryl rings of the boronic

acid fragment forming the walls of a cavity centred around the

bridging oxygen atom, with its aminophenol derived frag-

ments now adopting the role of buttressing substituents to

confer structural rigidity.

We have also varied the nature of the boronic acid template

used for supramolecular assembly, demonstrating that com-

plexation of 2-formyl-furanyl-boronic acid 9 with chiral

aminoalcohols 6a–e in chloroform quantitatively affords their

corresponding four-component boracycles 10a–e in 85–92%

isolated yield (Scheme 4). 11B NMR spectroscopic analysis of

these macrocycles reveals that the boron atoms of the furan

derived boracycles 10a–e (d 4.6–5.4 ppm) have more tetra-

hedral character than their corresponding phenyl derived

boracycles 8a–f (d 10.5–11.5 ppm). This increased tetrahedral

character may be a consequence of the need to incorporate a

more geometrically constrained five-membered furan ring into

these complexes. It may also explain why reaction of achiral

amino alcohols 6g–h with 2-formyl-furanyl-boronic acid 9 did

not result in clean formation of their corresponding four

component boracycles, which may be precluded by the oppos-

ing steric demands of incorporating tetrahedral sp3 boron

atoms and vicinal sp2 aryl carbon atoms into the central

boracyclic core of the macrocyclic ring system.

Conclusions

In conclusion, a range of covalent self-assembled macrocycles

8 and 10 containing bridging O–B–O–B–O have been prepared

and fully characterised. Their ease of preparation suggests that

this class of boracycle is well suited for the reversible self-

assembly of multicomponent systems, and we are currently

investigating the recognition properties of this structurally

diverse class of macrocycle.

Experimental

General synthetic methods

The solvents and reagents were reagent grade unless otherwise

stated and were purchased from Acros Organics, Alfa Aesar,

Fisher Scientific UK, Frontier Scientific Europe Ltd., TCI

Europe or Sigma-Aldrich Company Ltd., and were used

without further purification. Infra-red spectra were recorded

on a Perkin Elmer SpectrumRX spectrometer between 4400 cm�1

and 450 cm�1. Samples were evaporated from CHCl3 on

to a NaCl disc (film). Nuclear magnetic resonance spectra

were run in either chloroform-d. A Bruker AVANCE 300 was

used to acquire 1H-NMR spectra and recorded at 300 MHz,11B-NMR spectra at 100 MHz and 13C{1H} NMR spectra at

75 MHz. Chemical shifts (d) are expressed in parts per million

and are reported relative to the residual solvent peak or to

tetramethylsilane as an internal standard in 1H and 13C{1H}

NMR spectra. Boron trifluoride diethyl etherate was used as

an external standard in 11B NMR spectra. Mass spectra were

acquired with a micrOTOFQ electrospray time-of-flight

(ESI-TOF) mass spectrometer (Bruker Daltonik GmbH).

General procedure for the preparation of boracycles 8a-f and

10a–e

2-Formyl-phenyl-boronic acid 2 (60 mg, 0.4 mmol) or 3-formyl-

furanyl-2-boronic acid 9 (56 mg, 0.4 mmol) was stirred with a

chiral 1,2-amino alcohol 6a–f or 6a–e (0.4 mmol) in chloro-

form (5 mL) for 10 min. The solvent was then removed

under reduced pressure to afford boracycles 8a–f or 10a–e in

84–96% yield.

(R,R)-8a. Yellow oil (70 mg, 84%); [a]20D +22.0 (c 1.0,

CH2Cl2); vmax (film) 1628 (CQN); dH (300 MHz; CDCl3)

8.08 (2H, s, CHQN), 7.51 (2H, d, J 7.4, ArH), 7.35–7.27

(4H, m, ArH), 7.11 (2H, dt, J 7.4 and 1.1, ArH), 4.26 (2H, dd,

J 12.2 and 1.3, CHAHB(O)), 3.98 (2H, dd, J 12.2 and 1.3,

CHAHB(O)), 3.13 (2H, m, CH(iPr)–N), 2.96–2.83 (2H, m,

CH(CH3)2), 1.02 (6H, d, J 6.8, C(CH3)(CH3)) and 0.88

(6H, d, J 6.8, C(CH3)(CH3)); dC (75 MHz; CDCl3) 167.3

(CQN), 136.8, 133.6, 129.4, 127.2, 127.0, 126.2, 76.0, 60.9,

27.0, 21.1 and 19.4; dB (100 MHz; CDCl3) 10.7; m/z LRMS

(ESI+) 418 [(M + H)+, 13%], 283.2 (100), 200.1 (2); HRMS

(ESI+) found 417.2531 ([M + H]+ C24H30B2N2O3 requires

417.2515).

(S,S)-8b. Yellow solid (79 mg, 89%); m.p. 206–210 1C (dec);

[a]20D �26.1 (c 1.0, CH2Cl2); vmax (film) 1628 (CQN);

dH (300 MHz; CDCl3) 8.16 (2H, s, CHQN), 7.51 (2H, d,

J 7.4, ArH), 7.36–7.27 (4H, m, ArH), 7.11 (2H, dt, J 7.4 and 1.1,

Fig. 2 Crystal structure of macrocycle 8h. (a) Viewed along the

boron–boron axis. (b) Viewed perpendicular to the boron–boron axis.

Scheme 4 Preparation of boracycles 10a–e.


ArH), 4.35 (2H, dd, J 11.9 and 1.7, CHAHB(O)), 3.81–3.73

(4H, m, CHAHB(O) and CH–N), 2.16–2.06 (2H, m,

CHAHBC(N)), 2.02–1.92 (2H, m, CHAHBC(N)), 1.72–1.63

(2H, m, CH(CH3)2) and 0.90 (12H, app t, J 6.8, C(CH3)2);

dC (75 MHz; CDCl3) 166.9 (CQN), 136.9, 133.5, 133.1, 129.4,

127.2, 126.1, 66.7, 62.5, 40.4, 24.8, 23.0 and 22.9; dB (100 MHz;

CDCl3) 11.5; m/z LRMS (ESI+) 445 [(M + H)+, 14%], 412.4

(100), 292.2 (66), 227.2 (15); HRMS (ESI+) found 445.2873

([M+H]+ C26H34B2N2O3 requires 445.2828).

(R,R)-8c. Yellow oil (69 mg, 88%); [a]20D +14.7 (c 1.0,


8.13 (2H, s, CHQN), 7.50 (2H, t, J 7.4, ArH), 7.3–7.26

(4H, m, ArH), 7.12–7.07 (2H, m, ArH), 4.33 (2H, dd, J 12.0 and

1.7, CHAHB(O)), 3.78 (2H, dd, J 12.0 and 1.7, CHAHB(O)),

3.51 (2H, m, CH(Et)–N), 2.20–2.09 (4H, m, CHAHBMe) and

0.93 (6H, t, J 7.6, CH3); dC (75 MHz; CDCl3) 167.3 (CQN),

136.9, 133.5, 133.3, 129.3, 127.2, 126.2, 70.7, 62.4, 24.8

and 11.5; dB (100 MHz; CDCl3) 11.2; m/z LRMS (CI+) 389

[(M + H)+, 6%], 188.2 (50), 106.0 (46), 72.0 (100); HRMS

(EI+) found 388.2126 (2 � 11B) (M+� C22H26B2N2O3 requires

388.2124).

(R,R)-8d. Yellow oil (88 mg, 91%); [a]20D +21.1 (c 1.0,


7.66–7.63 (4H, m, CHQN and ArH), 7.43–7.30 (12H, m,

ArH), 7.20 (2H, br t, J 7.4, ArH), 7.09 (2H, dt, J 7.4 and

0.8, ArH), 5.25 (2H, m, CHAHB(O)), 4.65 (2H, dd, J 11.9 and

10.4, CH(Ph)-N), 3.95 (2H, m, CHAHB(O)); dC (75 MHz;

CDCl3) 166.1 (CQN), 137.0, 135.7, 133.9, 133.5, 132.3, 130.0,

129.9, 129.7, 129.5, 127.2, 126.8, 126.7, 71.4, and 69.1; dB (100

MHz; CDCl3) 11.3; m/z LRMS (ESI+) 485 [(M + H)+, 9%],

368.2 (10), 312.1 (100), 278.2 (16); HRMS (ESI+) found

485.2230 ([M + H]+ C30H26B2N2O3 requires 485.2202).

(R,R)-8e. Yellow solid (140 mg, 95%); m.p. 125–129 1C

(dec); [a]20D +19.4 (c 1.0, CH2Cl2); vmax (film) 1635 (CQN); dH(300 MHz; CDCl3) 8.25 (2H, s, CHQN), 7.64 (2H, d, J 7.0,

ArH), 7.48–7.14 (16H, m, ArH), 5.46 (2H, br d, J 9.8,

CHAHB(N)), 4.50–4.42 (2H, m, CH(Ph)(O)) and 4.03 (2H,

br d, J 9.8, CHAHB(N)); dC (75 MHz; CDCl3) 166.0, 137.0,

135.6, 133.9, 133.5, 132.3, 130.0, 129.9, 129.7, 129.4, 128.0,

127.3, 126.8, 71.3 and 65.8; dB (100 MHz; CDCl3) 10.5; m/z

LRMS (ESI+) 485 [(M + H)+, 100%], 312.1 (99); HRMS

(ESI+) found 485.2219 ([M + H]+C30H26B2N2O3 requires

485.2202).

(rac)-8f. Yellow solid (85 mg, 96%); m.p. 142–144 1C (dec);

vmax (film) 1625 (CQN); dH (300 MHz; CDCl3) 8.20 (2H, d, J

3.0, CHQN), 7.47 (2H, d, J 6.8, ArH), 7.35 (2H, d, J 7.4,

ArH), 7.28 (2H, app dt, J 7.5 and 1.1, ArH), 7.10 (2H, app dt,

J 7.5 and 1.1, ArH), 3.96–3.88 (2H, m, CH(N)), 3.79–3.70 (2H,

m, CH(O)), 2.26 (2H, br d, J 12.0, CHAHBC–O), 1.88–1.81

(4H, m, CHAHBC–O and CHAHBC–N), 1.71–1.66 (2H, m,

CHAHBC–N) and 1.50–1.12 (8H, m, 2�(CH2)2); dC (75 MHz;

CDCl3) 164.0 (CQN), 137.2, 133.2, 129.9, 129.1, 126.9, 126.3,

65.6, 36.3, 29.8, 27.3, 24.9 and 24.8; dB (100 MHz; CDCl3)

10.8; m/z LRMS (ESI+) 440 [(M + H)+, 100%], 290.2 (15);

HRMS (ESI+) found 441.2549 ([M + H]+ C26H30B2N2O3

requires 441.2515).

8g. 2-Aminophenol 6g (200 mg, 1.83 mmol) and 2-formyl-

phenylboronic acid 2 (275 mg, 1.83 mmol) were dissolved in

95:5 ethanol–benzene (25 mL) in a round bottom flask fitted

with a Dean–Stark condenser and stirred at reflux for 4 h. The

reaction mixture was cooled and the solvent removed under

reduced pressure. Washing with a little cold methanol afforded

8g as a yellow powder (709 mg, 90%): m.p. 182–183 1C (dec.)

(Lit. 180 1C (dec.)36); dH (300 MHz, CDCl3) 8.65 (s, 2H), 7.49

(2H, d, J= 7.5 Hz), 7.40–7.37 (2H, m, Ar), 7.29–7.12 (8H, m),

6.92 (4H, m); dC (75 MHz, CDCl3) d = 160.9, 155.5, 135.1,

134.2, 134.13, 133.1, 132.9, 131.5, 127.9, 118.7, 115.8, 113.7; dB(96.3 MHz, CDCl3) 9.6; m/z HRMS (ESI+) found 429.1571.

([M + H]+ C26H19B2N2O3 (M + H+) requires 429.1582).

8h. 2-Hydroxy-5-methylaniline 6h (123 mg, 1.0 mmol) and

2-formyl-phenyl-boronic acid 1 (150 mg, 1.0 mmol) were

dissolved in 95:5 ethanol–benzene (20 mL) in a round bottom

flask fitted with a Dean–Stark condenser and stirred at reflux

for 4 h. The reaction mixture was cooled and the solvent

removed under reduced pressure. Washing with a little cold

methanol afforded 8h as a orange powder (374 mg, 82%): m.p.

231–232 1C (dec.); dH (300 MHz, CDCl3) 8.64 (2H, s,

CHQN), 7.42 (2H, m, ArH), 7.36 (2H, s, ArH), 7.29–7.16

(6H, m, ArH), 7.11 (2H, d, J 8.4 ArH), 6.85 (2H, d, J 8.4), 2.35

(6H, s, CH3); dC (75 MHz, CDCl3) 158.24, 154.9, 148.8, 134.9,

134.1, 133.9, 132.9, 131.2, 128.2, 127.8, 115.3, 113.8, 21.5; dB(100 MHz, CDCl3) 8.9; m/z HRMS (ESI+) found 457.2011.

([M + H]+ C28H23B2N2O3 (M + H+) requires 457.1889).

(S,S)-10a.Dark brown solid (74 mg, 93%); m.p. 131–140 1C

(dec); [a]20D �36.8 (c 1.0, CH2Cl2); vmax (film) 1649 (CQN);

dH (300 MHz; CDCl3) 8.13 (2H, d, J 3.0 CHQN), 7.35 (2H, d,

J 1.9, ArH), 6.33 (2H, d, J 1.9, ArH), 4.39 (2H, dd, J 9.4 and

6.2, CHAHB(O)), 4.15 (2H, dd, J 9.4 and 4.0, CHAHB(O)),

3.90–3.84 (2H, m, CHQN), 2.28–2.17 (2H, m, CH(CH3)2) and

1.06 (12H, app dd, J 6.8 and 7.2, C(CH3)2); dC (75 MHz;

CDCl3) 157.2, 144.5, 132.6, 123.3, 110.2, 69.3, 63.5, 32.6, 20.0

and 17.4; dB (100 MHz; CDCl3) 4.7; m/z LRMS (ESI+) 397

[(M + H)+, 9%], 345.2 (100), 283.2 (36); HRMS (ESI+)

found 397.2122 ([M+H]+C20H26B2N2O5 requires 397.2100).

(S,S)-10b.Red oil (76 mg, 90%); [a]20D �34.0 (c 1.0, CH2Cl2);

vmax (film) 1649 (CQN); dH (300 MHz; CDCl3) 8.13 (2H, d, J

3.0, CHQN), 7.36 (2H, d, J 1.9, ArH), 6.32 (2H, d, J 1.9,

ArH), 4.39 (2H, dd, J 8.9 and 6.0, CHAHB(O)), 4.24–4.15

(2H, m, CH–N), 3.99 (2H, dd, J 8.9 and 7.4, CHAHB(O)),

1.77–1.65 (6H, m, CHAHBCH(CH3)2) and 1.00 (12H, app t, J

5.5, C(CH3)2); dC (75 MHz; CDCl3) 157.2, 144.7, 132.7, 122.8,

110.0, 71.9, 52.2, 32.6, 20.5, 8.9 and 8.1; dB (100 MHz; CDCl3)

4.6; m/z LRMS (ESI+) 425 [(M + H)+, 32%], 412.4 (27),

389.3 (35), 375.2 (100); HRMS (ESI+) found 425.2452

([M + H]+ C22H30B2N2O5 requires 425.2419).

(R,R)-10c. Red oil (134 mg, 91%); [a]20D +33.0 (c 1.0,


8.13 (2H, s, CHQN), 7.34 (2H, d, J 1.9, ArH), 6.31 (2H, d,

J 1.9, ArH), 4.42–4.39 (2H, m, CHAHB(O)), 4.09–3.99 (4H, m,

CHAHB(O) and CH–N), 2.05–1.90 (4H, m, CH2Me) and 1.07

(6H, t, J 7.0, CH3); dC (75 MHz; CDCl3) 156.3, 144.5, 132.6,

123.3 110.2, 67.7, 65.7, 26.4 and 10.4; dB (100 MHz; CDCl3)


4.8; m/z LRMS (ESI+) 369 [(M+H)+, 65%], 288.2 (100),

201.1 (34); HRMS (ESI+) found 369.1791 ([M+H]+

C18H22B2N2O5 requires 369.1785).

(R,R)-10d. Red solid (79 mg, 85%); m.p. 115–118 1C (dec);

[a]20D +39.6 (c 1.0, CH2Cl2); vmax (film) 1657 (CQN); dH(300 MHz; CDCl3) 7.83 (2H, d, J 3.0, CHQN), 7.53–7.44

(10H, m, ArH), 7.42 (2H, d, J 1.9, ArH), 6.24 (2H, d, J 1.9,

ArH), 5.35–5.28 (2H, m, CH(Ph)–N) and 4.57–4.45 (4H, m,

CHAHB(O)); dC (75 MHz; CDCl3) 158.3, 144.9, 136.6, 131.4,

129.8, 129.60, 129.59, 129.55, 129.50, 124.6, 110.3, 71.3

and 70.4; dB (100 MHz; CDCl3) 5.4; m/z LRMS (ESI+) 465

[(M + H)+, 100%], 415.2 (33), 335.2 (36), 292.1 (32), 215.1 (10);

HRMS (ESI+) found 465.1833 ([M + H]+C26H22B2N2O5

requires 465.1787).

(R,R)-10e. Dark brown solid (85 mg, 92%); m.p. 124–126 1C

(dec); [a]20D +18.5 (c 1.0, CH2Cl2); vmax (film) 1662 (CQN); dH(300 MHz; CDCl3) 8.26 (2H, br s, CHQN), 7.56 (4H, br d,

ArH), 7.45 (2H, d, J 1.9, ArH), 7.39–7.26 (6H, m, ArH), 6.34

(2H, d, J 1.9, ArH), 5.56–5.51 (2H, m, CHAHB(N)) and

4.11–4.08 (4H, m, CHAHB(N)) and CH(Ph)(O)); dC(75 MHz; CDCl3) 157.7, 145.1, 142.1, 129.7, 128.8, 128.6,

128.1, 126.7, 124.5, 110.4, 76.2 and 63.7; dB (100 MHz; CDCl3)

5.0; m/z LRMS (ESI+) 465 [(M + H)+, 65%], 415.2 (100),

323.2 (83); HRMS (ESI+) found 465.1841 ([M + H]+

C26H22B2N2O5 requires 465.1787).

Acknowledgements

We would like to acknowledge the EPSRC, Royal Society, the

Leverhulme Trust, Beckman-Coulter and University of Bath

for funding.

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N-Inversion in 2-azabicyclopentane derivatives: model simulations

for a laser controlled molecular switch

Bastian Klaumunzer* and Dominik Kroner

Received (in Montpellier, France) 17th July 2008, Accepted 18th September 2008


DOI: 10.1039/b812319e

We report model quantum simulations for the nitrogen inversion in 2-azabicyclo[1.1.1]pentane

derivates controlled by laser pulses proposing to use this class of molecules as molecular switches.

The derivatives trans-5-fluoro-2-methyl-2-azabicyclo[1.1.1]pentane and cis-5-fluoro-2-methyl-2-

azabicyclo[1.1.1]pentane are investigated by means of density functional theory and quantum

wave packet dynamics. The molecules have two stable, i.e. energetically well-separated,

conformers along the N-inversion coordinate. In 1D model simulations the transformation from

one conformer to the other is accomplished in the electronic ground state by using two

overlapping chirped linearly polarized IR laser pulses for the trans- and cis-isomer or alternatively

via an electronic excited state employing a pump-dump sequence of ultrashort UV laser pulses.

1. Introduction

Currently molecular switches are of interest in the field of

nanotechnology, e.g. for application in molecular electronics.1,2

In addition, they are also important in biology since

many biological functions are based on them, for instance,

allosteric regulation and vision. In general, theoretical and

experimental research on photo-switchable compounds has

mainly focused on cis-trans isomerization or photocyclic

reactions.3–5 Examples are chiroptical switches based on steri-

cally overcrowded alkenes,6 azobenzenes used as surface

mounted molecular switches7,8 or the laser controlled rever-

sible ring-opening of cyclohexadiene.9

Conformational transformations in molecules without

affecting the bond order have been, however, of rather less

interest for the design of molecular switches. The reason is

obvious since the barrier separating conformers is often in the

order of 1–10 kJ mol�1 making differentiation and, hence,

detection of the switchable molecular property, at room

temperature difficult if not impossible. Nevertheless, energy

barriers between conformers can be increased by sterically

demanding substituents making those molecules more attrac-

tive for controlled conformational switching. For instance,

Umeda et al. presented quantum simulations for the optical

isomerization of helical difluorobenzo[c]phenanthrene10 and

Hoki et al. performed quantum simulations for the change of

axial chiral 1,10-binaphthyl from its P- to M-form by laser

induced torsion around a single bond.11 Recently, we reported

a laser controlled axial chiral molecular switch, which allows

for the selective transformation between the achiral and either

the left- or right-handed form of an F-substituted styrene

derivative by torsion around a C–C single bond.12

A particular type of conformational change is the nitrogen

inversion (N-inversion).13,14 A nitrogen compound like

ammonia in a trigonal pyramid geometry (tertiary amine)

undergoes rapid nitrogen inversion. This interconversion is

very fast at room temperature because the energy barrier

(24.2 kJ mol�1) is relatively small.15 However, if the nitrogen

has sterically demanding substituents or is part of a rigid ring

system, it cannot easily invert around the lone electron pair

making the two conformers separable at room temperature.

Here we report quantum dynamical simulations of laser

controlled N-inversion of two 2-azabicycles. This class of

azabicyclic molecules has a particularly high inversion barrier

due the bicyclic effect, which has been of great experimental

and theoretical interest.13,16,17 We propose that derivatives of

5-X/Y-2-azabicyclo[1.1.1]pentane could serve as laser pulse

controlled molecular switches, which change according to

their conformation the size and direction of their dipole

moment mainly originating from an electronegative substitu-

ent X/Y, see Fig. 1. For a defined setup of the molecular switch

this system could be immobilized by chemi- or physisorption

on a surface via an adequate linking group R, see Fig. 1.

In this paper we investigate cis-5-fluoro-2-methyl-2-azabi-

cyclo[1.1.1]pentane (X = F, Y = H and R = CH3 in Fig. 1)

and trans-5-fluoro-2-methyl-2-azabicyclo[1.1.1]pentane (X =

H, Y = F and R = CH3 in Fig. 1). These molecules possess

two conformers of different dipole moments separated by a

high N-inversion barrier. In the following we will demonstrate

how these molecular systems can be switched via vibrational or

Fig. 1 Model for a laser controlled molecular switch: N-inversion of

5-X/Y-2-R-2-azabicyclo[1.1.1]pentane with X/Y being an electro-

negative substituent (here: cis X = F/Y = H and trans X = H/Y

= F) and R being e.g. a linker for a surface (here R = methyl).

Universitat Potsdam, Institut fur Chemie, Karl-Liebknecht-Str. 24-25,D-14476 Potsdam, Germany.E-mail: [email protected]



vibronic states. For these purposes control mechanisms

employing ultrafast laser pulses have been developed.

The remainder of the paper is organized as follows: The

model and the applied theoretical methods are explained in

section 2, the results of the quantum chemical and quantum

dynamical calculations including the laser control are pre-

sented in section 3. Section 4 provides a summary.

2. Model and methods

2.1 Quantum chemistry

The geometries of the two N-inversion conformers of both

isomers, namely the trans- and cis-isomer, were optimized with

density functional theory (DFT) employing the B3LYP18,19

functional and the ANO-L-DZ20 basis set as implemented in

the Molcas 6.4 program package.21 The obtained geometries

are denoted trans-Min1 and trans-Min2 for the trans-molecule

or cis-Min1 and cis-Min2 for the cis-molecule, see section

3.1 and Fig. 2 and 4. The transition states, called trans-TS

and cis-TS, were also calculated at the same level of theory.

To simulate the change of conformation the molecules are

assumed to be oriented with their N–C1-bond along the space

fixed z-axis, as shown in Fig. 2. Then, the N-inversion is

approximated by a partial rotation of the methyl group

around the y-axis while keeping the rest of the molecule fixed

in space. The angle a between the N–C1-bond and the x-axis is

used as reaction coordinate. In addition, for the trans-isomer

the free rotation of the methyl group around the N–C1-bond is

simulated by a rotation of the hydrogen atoms of the methyl

group around the N–C1-bond. Here the dihedral b, measured

between the H1–C1-bond and the x-axis, is used as reaction

coordinate, see Fig. 2. (In practice, first, the hydrogen atoms of

the methyl group are rotated clockwise around z-axis by b and

afterwards the whole methyl group is rotated around the

y-axis by a.)For the laser control via the ground state N-inversion states,

see section 3.3, the unrelaxed potential energy surface (PES) of

the electronic ground state along a is calculated by B3LYP/

ANO-L-DZ while keeping the rest of the geometrical para-

meters frozen to the minimum energy geometry trans-Min1 or

cis-Min1. For the trans-switch the calculations are also per-

formed along b obtaining a two-dimensional PES. Accord-

ingly, the permanent dipole moment along a is obtained on the

same level of theory as the PES.

For the control scenario using UV laser pulses for the

cis-isomer, see section 3.4, the first ten singlet electronic excited

states along a are calculated by time-dependent DFT

(TDDFT) with B3LYP and 6-31G(d,p) as implemented in

the GAUSSIAN0322 package. Transition dipole moments

between ground and any of the electronic excited states are

obtained on the same level of theory. As previously cis-Min1 is

used as reference geometry.

2.2 Model Hamiltonian

As the moment of inertia of the methyl group with respect to

the space fixed y-axis is about 100-times smaller than that of

the rest of the molecule, we assume the F-azabicyclo group

being fixed in space with only the methyl group moving,23 see

Fig. 2. To obtain the N-inversion eigenenergies eiv and eigen-

functions fiv of the ith electronic state the time-independent

Schrodinger equation

Himol(a)f

iv(a) = eivf

iv(a) (1)

is solved numerically. The molecular Hamiltonian Himol(a) is

given as

Hi

molðaÞ ¼ ��h2

2Iy

d2

da2þ ViðaÞ: ð2Þ

Vi(a) is the potential energy curve of the ith electronic state

with i = 0 for the electronic ground states and i 4 0 for

electronic excited states, cf. Fig. 5. Iy is the moment of inertia

for the rotation of the methyl group around the y-axis: Iy =PAmA�r2A. The distances rA of the atoms A with mass mA,

namely C1 and the hydrogens attached to it, are obtained from

the minimum energy geometry trans-Min1 or cis-Min1, see e.g.

Fig. 2. We obtain Iy = 247390.32 mea20 for the trans- and

Iy = 246484.74 mea20 for the cis-isomer. Note that, the N-in-

version is here modelled as a partial rotation of the methyl

group around the space-fixed y-axis while keeping the rest of

the molecule fixed in space, as described in section 2.1. Eqn (1)

is, then, solved by the Fourier Grid Hamiltonian method24

using N = 256 grid points in the IR-pulse case and N = 1024

points in the UV-pulse case, i.e. the coordinate a is

expressed as

ai = a0 + iDa, i = 0, . . ., N � 1, (3)

where a0 = 601 and Da = 0.70591 in the IR case (section 3.3)

and Da = 0.15641 (section 3.4) in the UV case.

Fig. 2 Optimized geometries of trans-5-fluoro-2-methyl-2-azabicyclo-

[1.1.1]pentane obtained from B3LYP/ANO-L-DZ: trans-Min1 is

the global minimum with angle a set to 90.01 in the space fixed

coordinate system. trans-Min2 is the optimized geometry of the second

N-inversion conformer at a E1891. trans-TS is the transition state

geometry at a E1651. trans-Min3 is the unrelaxed minimum along awith a E1931 using trans-Min1 as reference.


2.3 Quantum dynamics

To describe the laser-driven quantum dynamics the time-

dependent Schrodinger equation is solved numerically:

i�h@

@tCða; tÞ ¼ Hða; tÞCða; tÞ: ð4Þ

for

Cða; tÞ ¼C0ða; tÞ

..

.

Ciða; tÞ

0

B@

1

CA; ð5Þ

with C0(a,t) being the wave function of the electronic ground

state and Ci(a,t) the wave function of the ith excited state. The

Hamilton operator H(a,t) is given by:

Hða; tÞ ¼H

00. . . H

0i

..

. . .. ..

.

Hi0

. . . Hii

0

B@

1

CA; ð6Þ

within the semiclassical dipole approximation:

Hii ¼ H

i

mol�^m!iiðaÞ � E

!ðtÞ; ð7Þ

Hij ¼ �^m!ijðaÞ � E

!ðtÞ; ð8Þ

where^m!iiðaÞ are the permanent dipole moments of the ith

electronic state and^m!ijðaÞ are the electronic transition dipole

moments. For the dynamical simulations we set the permanent

dipole moments of the ith excited state equal to that of the

electronic ground state ð^m!iiðaÞ ¼ ^m!00ðaÞÞ , the transition dipole

moments are set^m!0iðaÞ ¼ ^m!i0ðaÞ and the transition dipole

moments between all other electronic excited states are set

zero ðj^m!ij j ¼ 0Þ .The electric field ~E(t) of the laser pulses used here is

given by:

E!ðtÞ ¼ e

!j E

0cosðo � ðt� tcÞ þ ZÞsin2 pðt� tcÞ2fwhm

þ p2

� �

; ð9Þ

for |t � tc|r fwhm. Z is the time-independent phase and fwhm

the full width at half maximum (2fwhm equals the pulse

duration). The polarization vector ~ej = ~excos(j) + ~ezsin(j)with polarization angle j, where ~ex/z is the unit vector along

the x/z-axis. Hence, the laser is chosen to propagate in

y-direction. E0 is the electric field amplitude and tc the pulse

center, i.e. the time when the sin2 shape-function reaches its

maximum. The laser pulse frequency o can be linearly chirped

by _o ¼ do=dt :

oðtÞ ¼ o0 þ _o � ðt� tcÞ; ð10Þ

where o0 is the central frequency at t = tc. All quantum

dynamical propagations were performed with the wavepacket

program package25 using the second order splitting26 in grid

representation with a time step of 0.25 fs for the IR case

(section 3.3) and 0.025 fs for the UV case (section 3.4).


3.1 Geometries and PESs

The geometry optimization of the trans-5-fluoro-2-methyl-2-

azabicyclo[1.1.1]pentane with B3LYP/ANO-L-DZ resulted in

two stable minima trans-Min1 and trans-Min2, shown in

Fig. 2, where trans-Min1 is the global minimum. In terms of

the angle a the two minima are at 901 (trans-Min1) and at

189.41 (trans-Min2) according to the space fixed coordinate

system. Thus, the transformation between them is achieved by

a rotation of the methyl group around the y-axis by about

1001. At the transition state trans-TS of the nitrogen inversion

the angle a E140.21. The molecule has CS-symmetry with

respect to the xz-plane in all conformations. The energy

difference between the transition state trans-TS and the abso-

lute minimum trans-Min1 is 6397.9 cm�1 corresponding to

76.5 kJ mol�1 which is more than three times higher than the

inversion barrier of ammonia (24.2 kJ mol�1).15 As one can see

from Fig. 2 the methyl group is rotated around the C1–N bond

by about 1801 while going from trans-Min1 to trans-Min2.

Therefore a two-dimensional PES (Fig. 3) along a and b was

calculated.

The PES shows three minima belonging to three different

molecular structures. We find minima at a = 901/b = 01

(trans-Min1), at a = 1901/b = 01 (trans-Min3) and at a =

1901/b = 1801 (trans-Min4), while the latter corresponds to

the unrelaxed geometry of trans-Min2. Additionally there

are three distinct maxima: trans-Max1 at a = 901/b = 1801,

trans-Max2 at a = 1501/b = 01, which belongs to the

unrelaxed geometry of trans-TS, and trans-Max3 at a =

1501/b = 1801. The energy differences between trans-Min1

and trans-Max1 and the barrier height between trans-Min4

and trans-Min3 are approximately of the same size, namely

1000 cm�1 (12 kJ mol�1). This barrier, resulting from the free

rotation of the methyl group around the C1–N bond, is, as

expected, of the same height as the rotational barrier of ethane

(12 kJ mol�1).27 The PES shows that the N-inversion should

not significantly be affected by the free rotation of the methyl

group, see Fig. 3. Yet, the methyl group used here represents

merely a placeholder for an arbitrary substituent R, for

instance, a linker to a surface. The free rotation of the methyl

Fig. 3 Unrelaxed potential energy surface along a and b for trans-5-

fluoro-2-methyl-2-azabicyclo[1.1.1]pentane (B3LYP/ANO-L-DZ).

trans-Min1 denotes the minimum energy geometry which was used

as reference geometry.


group around the C1–N bond is neglected in the following, i.e.

the unrelaxed PES only along a with trans-Min1 as reference

geometry is used for all dynamical simulations.

It should be noted, as an aside, that a normal mode analysis

of the relaxed minimum geometry trans-Min1 employing

B3LYP/6-31G(d,p) reveals three modes that should be con-

sidered important for the dynamical simulations: (i) The

rotation of the methyl group around the C1–N bond (b) at245 cm�1, (ii) the bending of the C1–N-bicylo-angle describing

the rotation of the methyl group around the y-axis at 251 cm�1

(a), and (iii) the torsion of the C1–N-bicylo-dihedral describ-

ing the rotation of the methyl group around the x-axis at

286 cm�1 (all frequencies scaled according to ref. 29). Hence, a

coupling of mode (ii), which characterizes our model reaction

coordinate a, to modes (i) and (iii) cannot completely be ruled

out for higher excited states, because their energies lie within

the range of the inversion excitation energies, see section 3.2.

Due to the unrelaxed geometry the inversion barrier is

approx. 1600 cm�1 higher compared to the relaxed one. To

get an idea whether the barrier can be crossed thermally

we calculated N-inversion rates according to the theory of

Eyring.28 The necessary thermodynamic quantities were ob-

tained with the GAUSSIAN03 program package employing

B3LYP/6-31G(d,p). At 298 K we obtain an inversion rate

from trans-Min1 to trans-Min2 of 2.20 s�1. So at room

temperature we find a rather small rate for spontaneous

N-inversion compared to ammonia (about 109 s�1 without

tunneling). As the backward reaction rate is also fairly small at

room temperature (3.54 s�1) the here investigated conformers

are considered thermally sufficiently stable to monitor the

change of the dipole moment, see discussion below. For a

more detailed discussion the reader is referred to ref. 23.

The geometry optimization of the cis-5-fluoro-2-methyl-2-

azabicyclo[1.1.1]pentane with B3LYP/ANO-L-DZ resulted in

two stable minima, denoted cis-Min1 and cis-Min2. The

corresponding structures are shown in Fig. 4. Cis-Min1 is

the global minimum, however, the energy difference between

the two minima is only 7 cm�1. In terms of the inversion angle

a the two minima are found at 901 (cis-Min1) by definition and

at 179.81 (cis-Min2), i.e. the transformation between them is

achieved by flipping the methyl group by about 901. As the

steric interactions of the X-substituent (X = F) with

the methyl group (R) is stronger than for the trans-isomer

(X = H), the change in a going from one conformer to the

other (E901) is smaller than for the trans-isomer (E1001).

At the transition state cis-TS of the nitrogen inversion

the angle a E164.71. The molecule has a mirror plane in the

xz-plane in all conformations. The energetic difference

between the transition state cis-TS and the absolute minimum

cis-Min1 is now 6358.6 cm�1 corresponding to 76.1 kJ mol�1.

Hence, the barrier height is similar to the one of the trans-

isomer (76.5 kJ mol�1).

For the cis-isomer we also observe that the methyl group is

rotated around the C1–N bond by about 1801 while going

from cis-Min1 to cis-Min2. As the coupling of the N-inversion

to the rotation of the methyl group is, as discussed above,

rather weak, only the one-dimensional PES along a starting

from cis-Min1 is considered. The unrelaxed electronic ground

state potential along a (B3LYP/ANO-L-DZ) is shown in

Fig. 5. The inversion barrier height is 8000 cm�1 and due to

the unrelaxed geometry approx. 1600 cm�1 higher than in the

relaxed case (cis-TS). As noted previously, due to the frozen

geometry the inversion angles at the top of the inversion

barrier (cis-Max2) and for cis-Min3 differ from those of the

optimized geometries, i.e. a(cis-Max2) E1481 and a(cis-Min3)

E1911. Here we can denote that the 1D cut of the potential

energy surface of the trans- and cis-isomer, Fig. 5, are quanti-

tatively similar, so that the curves overlap in the figure.

Fig. 4 Optimized geometries of cis-5-fluoro-2-methyl-2-azabicyclo-

[1.1.1]pentane obtained from B3LYP/ANO-L-DZ: cis-Min1 is the

global minimum with angle a set to 90.01 in the space fixed coordinate

system. cis-Min2 is the optimized geometry of the second N-inversion

conformer at a E1791. cis-TS is the transition state geometry

at a E1651. cis-Min3 is the unrelaxed minimum along a with aE1911 using cis-Min1 as reference.

Fig. 5 Potential energy surface of 5-fluoro-2-methyl-2-azabicyclo-

[1.1.1]pentane along a for the electronic ground state S0 (B3LYP/

ANO-L-DZ) (fitted from 33 single point calculations with a cubic

spline) and first three electronic excited states, S1 to S3 (TD-B3LYP/

6-31G(d,p)). Min1 denotes cis- and trans-Min1, Min3 cis- and trans-

Min1 and Max2 denotes cis- and trans-Max2, since the potentials for

the cis- and trans-isomers overlap on the scale of the picture.


For the UV laser pulse control of the cis-isomer electronic

excited states were calculated as explained in section 2.1. The

first three excited states S1, S2 and S3 (TD-B3LYP/6-31G(d,p))

of the cis-isomer are depicted in Fig. 5. The shape of the

excited state potentials are similar to those of other aliphatic

tertiary amines calculated by Solling et al. with TD-B3LYP/

6-31++G(2df).30 We observe a vertical excitation energy

from S0 to S1 at a = 901 of approx. 7.2 eV (58071.9 cm�1).

Compared to the amines in ref. 30 the here computed excita-

tion energies are similar, so that we consider the level of theory

used for the calculation of the electronic excited states to be

sufficient for our needs, although Rydberg states might not be

described precisely due to the restrictions of the basis set used

here, i.e. the lack of diffuse functions. An orbital analysis

shows that the transitions are predominantly of n - p and

n - s character in accordance to the amines in ref. 30.

The topology of the excited states differ from the ground

state. Instead of two minima along a there is only one single

minimum. The topology of the S1 potential is suited well for

switching the molecule via this excited state since no barrier

has to be crossed on V1 by switching from one conformer to

the other, see section 3.4.

The components of the permanent dipole moment^m!00 along

a are shown in Fig. 6(a) for the trans-isomer. Because of the

Cs-symmetry in the xz-plane the permanent dipole moment

along y is zero at all a. The laser will, therefore, be chosen to

propagate in y-direction, see eqn (9). The major change in

m00x occurs in range of 150–2101, while m00z has its major change

in the range of 70–1501.

For the cis-isomer the components of the permanent dipole

moment,^m!00ðaÞ , and of the transition dipole moments from S0

to S1,^m!01ðaÞ , are plotted in Fig. 6(b) and (c). For reasons of

symmetry again m00y = 0 for all a. In contrast to the trans-

isomer the major change in m00x is in range of 70–1201, while for

m00z the major change occurs in the range of 120–2101 now. For

an efficient pump–dumpmechanism31,32 for both the trans- and

the cis-isomer the laser pulses will be polarized in accordance

to the regions of largest change in the dipole components, see

section 3.3. For the electronic transition to the S1 state of the

cis-isomer, UV laser pulses will be xz-polarized as well, as

transitions in y polarization are forbidden by symmetry.

3.2 Inversion eigenstates

The inversion eigenstates of the ground state were calculated

as described in section 2.2. There are 36 eigenstates (trans-

isomer)/35 eigenstates (cis-isomer) below the barrier whose

eigenfunctions f0v are localized in the left and 21 eigenstates

(both isomers) whose eigenfunctions f0v are localized in the

right potential well. All eigenfunctions f0v which satisfy

PK�1i=1

|f0v(ai)|

2Da Z 0.999 are called ‘‘left localized’’ eigenfunctions.

Correspondingly, an eigenfunction is called ‘‘right localized’’ ifPN

K+1|f0v(ai)|

2Da Z 0.999 is fulfilled, where aK is the grid

point defined by the maximum of the inversion barrier V0(aK)(trans-Max2/cis-Max2). All states in the left quantum well are

termed ‘‘L’’ and those in the right quantum well are termed

‘‘R’’. The associated wave functions f0v are denoted fuL with

f0L to f35L/f34L (trans/cis) for the left well and fuR with f0R to

f20R for the right well. (The superscript 0 is omitted for the

‘‘localized’’ eigenfunctions since the concept applies only for the

ground state.) In addition, there are two more eigenstates below

the barrier (both isomers) which are considered ‘‘delocalized’’

in these terms. Table 1 lists the two lowest N-inversion

eigenenergies e0v in each ground state minimum, their energy

difference De = e0v0 � e0v and the corresponding dipole matrix

elements hf0v0|m

00x/z|f

0vi = hmx/zi for trans- and cis-isomer.

3.3 Switching via ladder climbing

For the quantum dynamical simulations the system is

initially assumed to be in the inversion ground state 0L, i.e.

Fig. 6 x-, y- and z-component along a for the permanent dipole

moment ~m 00 (B3LYP/ANO-L-DZ) of (a) trans- and (b) cis-5-fluoro-2-

methyl-2-azabicyclo[1.1.1]pentane, and (c) the transition dipole

moments^m!01 to S1 (TD-B3LYP/6-31G(d,p)) for the cis-isomer.

Table 1 Selection of eigen-energies e0v in cm�1 of the electronicground state for the trans- and cis-isomer, energy differences De andtransition dipole matrix elements hmx/zi in Debye

Isomer

trans cis

f0L f1L f0R f1R f0L f1L f0R f1R

e0v 126.2 379.0 3397.2 3649.3 124.5 389.7 3383.9 3642.0De 252.8 252.1 265.2 258.1hmxi �0.0095 �0.029 �0.078 �0.019hmzi �0.044 �0.0082 0.013 0.051


C(t=0)= f0L. The goal of the laser control is, to transfer the

population to the R states. This is achieved, first, by climbing

up the vibrational ladder until the wave packet is above the

barrier. Then, a dump laser pulse will be used to induce a

downward ladder climbing in the right potential well to trap

the wave packet there. To achieve a most efficient ladder

climbing the laser pulses will be linearly chirped33–35 according

to eqn (10) to compensate for the anharmonicity of the

potential at higher energies.

Fig. 7(a)/(d) show the time evolution of the electric field of

the laser pulse sequences for switching the trans- (a) or cis-

isomer (d). The resulting time evolution of the expectation

value of the angle a and the population of the L (PL) and R

(PR) states are shown in Fig. 7(b)/(e) and (c)/(f) for trans (b, c)

or cis (e, f), respectively. The populations in the L and R states

and the population in all other states (PD) are calculated as

follows:

PLðtÞ ¼X35=34

u¼1jhCðtÞjfuLij

2; ð11Þ

PRðtÞ ¼X20

u¼1jhCðtÞjfuRij

2; ð12Þ

PD(t) = 1 � PL(t) � PR(t). (13)

3.3.1 Switching the trans-isomer. The electric field consists

of an overlapping pump–dump sequence, see Fig. 7(a), with

optimal parameters as given in Table 2. All pulse parameters

were tuned manually to obtain the best possible result. In

general, the laser parameters are chosen in accordance with the

molecular properties, i.e., transition dipole moment elements

and energy differences. This approach allows for a deeper

understanding and more flexible control of the underlying

switching mechanism. As initial guess we set the frequency

o0 of the pump/dump pulse to the transition frequency of

0L - 1L (252.8 cm�1)/1R - 0R (252.1 cm�1). Further fine

tuning then lead to a frequency close to a transition frequency

between higher R states (202.00 cm�1). Initially the polariza-

tion angles j were estimated by tanj ¼ hmzihmxi36 for the transition

between 0L and 1L (77.801) and 1R and 0R (15.801). Further

optimization of the pulse parameter resulted in almost the

same values for the pump (77.801) and dump pulse (15.951),

see Table 2. A non-overlapping sequence of pump and dump

pulse sequence was found less efficient. At first, the overall

pulse is more z-polarized and has a negative chirp; after

450–500 fs it changes its polarization towards x-direction

and the chirp becomes positive. One can see the change in

polarization direction as well as the frequency chirp more

clearly in the Husimi probability distributions in Fig. 8. The

Husimi distribution37 is obtained from:

PHðt; eÞ ¼Z

dt0Z

1

�hde0e�

ðt�t0Þ2k e�

kðe�e0Þ2�h PW ðt0; e0Þ ð14Þ

with PW being the Wigner probability distribution:38

PW ðt; eÞ ¼1

p

Z1

�1

dt0E?ðt� t0ÞEðtþ t0Þe�2it0e

�h : ð15Þ

where E is the x- or z-component of the electric field, and

k = 2s2 the parameter of the gaussian distribution with

s ¼ 4000�h=Eh , e the energy, and t the time.

From Fig. 7(b) one can see that once the propagation is

started the wave packet begins to oscillate in the left quantum

well until it crosses the barrier after 500 fs and is dumped into

the right quantum well (a = 1851). The final population is

spread over several R states such that the expectation value of

a still oscillates around 1851 as the mainly R-localized wave

packet evolves in time. Nevertheless, the switching of the

molecule was successful as at final time 92.5% of the popula-

tion has been transferred from the left potential well (a = 901)

to the right potential well (a = 1911). States 0R to 6R are the

most populated eigenstates after the laser pulse sequence. The

missing 7.5% of population remains in theD-states, i.e.mainly

above the barrier.

It should be noted that the mean peak intensity

(I = 0.5e0c|E0|2) of the IR pulse is rather high: I =

16.5 TW cm�2 due to the high N-inversion barrier and the

comparatively small change of the dipole moment components

along a. The high laser amplitudes could be decreased by using

longer pulses. But longer pulses could cause a decrease of

efficiency in transferring population from L to R for effects as

wave packet broadening. For very long times even intramole-

cular vibrational redistribution (IVR) cannot be neglected any

more. Therefore, we considered the cis-isomer in the next

Fig. 7 Laser pulse sequence for the N-inversion via IR ladder

climbing from trans-Min1 to trans-Min3 (a)–(c) and cis-Min1 to

cis-Min3 (d)–(f); for the laser pulse parameters see Table 2. Time

evolution of (a)/(d) the x- and z- component of the electric field, (b)/(e)

the expectation value of the inversion angle hai, and (c)/(f) the

population of the L, R and D-states according to eqns (11)–(13).


section which has greater transition dipole matrix elements, cf.

Table 1, between the eigenstates and such should lead to a

decrease of the laser pulse intensities.

3.3.2 Switching the cis-isomer. For switching the cis-isomer

the electric field consists as in the previous case of an over-

lapping pump–dump sequence with optimal parameters as

given in Table 2. Again a non-overlapping sequence of pump

and dump pulse sequence was found less efficient. However,

now the pump pulse is more x-polarized while the dump pulse

is more z-polarized, according to the regions where major

changes in the dipole moment components occur, cf. Fig. 6

and discussion in section 3.1. Here the polarization angles

were initially determined as above for the 0L–1L transition as

j = �9.51 and for the 1R–0R transition as j = �69.61.Hence, the optimized polarization angle for the pump pulse

(�10.01) is in a good agreement with the calculated one, for

the dump pulse the optimized angle (�751) differs slightly.For getting the wave packet above the barrier a pump pulse

with a slight positive chirp was found beneficial. The second

pulse, then, has a notable positive chirp and dumps the wave

packet into the right potential well. The wave packet propaga-

tion shows almost the same behaviour as for the trans-isomer

concerning the expectation value of a and the population

dynamics, see Fig. 7(d–f). The pulse sequence is found with

90% PR at final time almost as efficient as what we achieved

for the trans-isomer. The population is mostly transferred to

the states 3R–9R. As expected the intensity of the laser pulse

(I = 11.5 TW cm�2) was lowered by 5 TW cm�2, however, it is

still high. For that reason a switching mechanism via the

electronic excited states will be investigated in the following

section.

3.4 Switching the cis-isomer via S1

For the transformation of the cis-isomer of the azabicycle via

the excited state S1 the initial state is the inversion ground state

0L as in the previous case. Now the wave packet is to be

excited to S1 employing a UV pump laser pulse, and after some

short time the wave packet is dumped into the right potential

well of the electronic ground state. The criterion for a success-

ful propagation is again that a large part of population is

transferred to the twenty R states below the central barrier of

the ground state potential. Higher states are omitted at this

point, see discussion below.

Fig. 9 shows (a) the electric field of the UV pump dump

pulse sequence, (b) the expectation value of angle a, (c) and the

population (P) of S0 (P(S0)), S1 (P(S1)) and of the R states (PR).

We obtain a sequence of pulses where the pump and dump

pulse do not overlap at all. The UV pulse sequence is polarized

in the xz-direction. The pulse parameters for the initial guess

were obtained in analogy to the procedure discussed above:

We set the frequency o0 to the energy difference V1–V0 at a =

901 for the pump pulse and at a = 1901 for the dump pulse.

From the transition dipole moments hf00Lj

^m!01jf170i and

hf170j

^m!01jf00Ri we computed the initial polarization angles

j of the pump pulse (�83.11) and of the dump pulse (7.41),

see 3.2. For the pump pulse we obtained j = �83.21 and for

the dump pulse j = 10.51 after further manual optimization.

During the control sequence the pump pulse transfers 99%

of the population from S0 to S1 (t = 360 fs), see Fig. 9(c). The

center of the wave packet then travels on S1 back and forth

until the dump pulse transfers 91.5% of the population back

to the ground state (t = 600 fs) where the wave packet is then

mostly trapped in the right potential well of the ground state

(89.5%), see Fig. 9(c). All the population transferred from S1to the R-states of S0 is found in the inversion states 14R to

17R; so the wave packet is still highly vibrationally excited and

it is therefore oscillating between a = 1631 and a = 1881, see

Fig. 9(b). The missing 2% of the ground state population

remains above the barrier (Max2). 8.5% of the electronically

excited population remains in S1.

The goal of reducing the laser pulse intensity is reached. The

intensity was brought down to 0.83 TW cm�2 for the pump

Table 2 Laser pulse parameters for the IR laser pulse sequences, depicted in Fig. 7(a) and (d), for the trans- and cis-isomer

Isomer Pulse type j (1) fwhm/fs tc/fs E 0/GV m�1 o0/cm�1 _o/cm�1 fs�1 Z/rad

trans Pump 77.80 640 640 16.5 243.25 �0.365 �3.25Dump 15.95 955 1060 16.5 202.00 0.178 0.125

cis Pump �10.00 400 400 8.0 252.25 0.036 �0.25Dump �75.00 550 650 11.5 220.25 0.180 0.75

Fig. 8 Husimi probability distributions of the (a) z- and (b) x-com-

ponent of the electric field of the IR switching pulse sequence for the

trans-isomer.


pulse and 2.3 TW cm�2 for the dump pulse. As the population

is transferred to high N-inversion states of S1, namely eigen-

states 67–72, whereas 70 is the most populated, the wave

packet is not in a compact form any more on the excited state

and hence it turned out to be very difficult to dump the

population effectively to the ground state. The transition

frequency of the pump pulse o0 = 58170 cm�1 is close to

the energy difference of the eigenenergies of the eigenstates f170

and f0L (58169.65 cm�1). Neither shortening or prolonging

the laser pulse durations nor lowering or increasing their

intensity increased the population of the R states.

Comparing the ladder climbing mechanism to the excited

state mechanism one can say that almost the same amount of

the population is transferred to the twenty R states below the

barrier. But in the ladder climbing case lower vibrational R

states have been populated, which leads to smaller oscillations

in a. In the UV case the laser pulse durations are much shorter

(690 fs) than in the IR case (1200 fs). However, upon including

the excited state S2 in the calculations the population trans-

ferred back to S0 is excited to S2 by the dump pulse such that at

the end of the propagation only 50% of the population is

transferred to the R states of S0 (not shown). Hence, we have a

loss of efficiency, when S2 is included. When including more

states up to S10 there is only additional population transfer to

S6 by 2%. There will be, however, no population transfer to

S3, S4, S7 and S8 as the transitions from S0 are symmetry

forbidden in x and z direction. For energetic reasons there is

also no population transfer to S9 or S10. Furthermore, due to

the different sign of the transition dipole moment of S5 in the z

direction the coupling to the x–z polarized dump pulse is

rather weak due to the improper polarization. Therefore, no

significant population transfer to S5 takes place. Here we also

note, that a more precise description of the electronic excited

state potentials might be necessary in order to correctly

account for avoided crossings in the regions of a = 1651 to

2151 which might lead to nonadiabatic transitions. In sum-

mary, while the laser induced switching via electronic excited

states is faster and allows for more realistic laser parameters—

at least within the framework of our model—the efficiency is

reduced due to undesired electronic excitations mainly to S2.

4. Summary and conclusions

Quantum simulations for a laser driven model system were

presented based on the approximate description of the nitro-

gen inversion in two azabicycles. Each of the two proposed

molecules, cis- and trans-5-X-2-R-2-azabicyclo[1.1.1]pentane,

possesses two stable conformers due to the sterically hindered

N-inversion. The molecules change the size and direction of

their dipole moment upon N-inversion where the electronega-

tive substituent X/Y mainly determines this property. The

substituent R can be used to immobilize the system, for

instance, by mounting it via a linker group R on a surface.

To investigate the possibilities of laser control we chose here

for a toy model X/Y = F and R= CH3. For trans-5-fluoro-2-

methyl-2-azabicyclo[1.1.1]pentane a two-dimensional poten-

tial along a and b was calculated, where a models the

N-inversion and b describes the free rotation of the methyl

group around its single bond to N. For the cis-isomer the

potential energy curve only along the nitrogen inversion

coordinate (a) was computed. In addition, the potential energy

curve was evaluated along a for several electronic excited

states. The designed laser pulse sequences allow to transfer

the molecules from their energetically more stable conformer

(Min1) to the less stable conformer (Min3) along the model

reaction coordinate (N-inversion). Linearly polarized laser

pulses were used to switch the molecule by either IR induced

ladder climbing or alternatively using the electronic excited

state (S1) as intermediate state in case of the cis-isomer.

In the case of vibrational ladder climbing two overlapping

laser pulses with chirped frequencies in the IR range were used

to switch the molecules. Thereby excitation and de-excitation

were mainly controlled by changing the chirp and the polari-

zation of the laser pulse. But the obtained laser intensities are

rather high.

Hence, we applied a control mechanism for the cis-isomer via

the first electronic excited S1. In this case the molecule is initially

excited by a UV pump laser pulse to a highly vibrational state

Fig. 9 Laser pulse sequence for the transformation of the cis-isomer

via the excited state S1 from cis-Min1 to cis-Min3; for the laser pulse

parameters see Table 3. Time evolution of (a) the x- and z-component

of the envelope function of the electric field, (b) the expectation value

of the inversion angle hai, and (c) the population in S0, S1 and R-states.

Table 3 Laser pulse parameters for the UV laser pulse sequencedepicted in Fig. 9(a)

Pulsetype j (1)

fwhm/fs tc/fs

E0/GVm�1

o0/cm�1

_o/cm�1

fs�1 Z/rad

Pump �83.20 180 180 2.5 58170.0 0.0 0.0Dump 10.50 155 535 4.5 51387.0 0.86 0.0


of S1. Afterwards a second UV laser pulse brings it down to the

electronic ground state in the right potential well.

In all cases the molecule was switched effectively as at least

90% of population was found in the target potential well.

However, if the second excited state S2 is considered as well

there is only population transfer of 50% to the R states in case

of UV pulses.

The efficiency of the population transfer could be dimin-

ished due to dissipative effects as intramolecular vibrational

redistribution (IVR). However, due to the overlapped pump-

dump control scheme the system will most probably return to

its thermal equilibrium geometry from where it can be

switched again, cf. ref. 14. The inversion mode may also

couple to other bending modes of the methyl group making

the proposed control strategy more complicated. Here, simu-

lations including more degrees of freedom could help to

quantify the effect. In this sense the mechanism via electronic

excited states could be more efficient as the switching process is

faster than in the case of vibrational ladder climbing. The

switching process via electronic excited states contains, how-

ever, the possible risk of undesired photochemical pathways

which could lead to a diminution of the desired control.

The degree of orientation of the molecule with respect to the

polarization vector of the laser field determines the efficiency

of the control mechanism. There are, however, theoretical and

experimental methods for orienting or aligning molecules, e.g.

in strong electric fields,39,40 using elliptically polarized lasers41

or applying optimal control theory.42

For the molecular system to be used in electronic devices, it

should be immobilized e.g. by adsorption to a surface. For this

one has to find a suitable linking group. By immobilizing the

molecule on a surface the switching mechanism will be different

from those presented here since the azabicycle will flip instead

of the R-group. Investigations along this thread are on the way.

Still, different fixed molecular orientations (with potentially

restricted rotations with respect to the surface normal) are

possible upon chemisorption depending on the symmetry of

the surface and the linker group. For surface mounted mole-

cules with different orientation along the surface normal sto-

chastically optimized elliptically polarized laser pulses were

found efficient for control of molecular isomerization.43 Note

that the coupling of the vibrational and electronic degrees of

freedom of the molecule to the surface degrees of freedom

(phonons, electron-hole pairs) may intensify energy dissipation

depending on the nature of the solid and the linking groups.

Nevertheless, the model molecules presented here could be a

good supplement to the model molecular switches which are

based on cis–trans isomerization or photocyclization reaction.

Acknowledgements

We thank P. Saalfrank for stimulating discussions. Financial

support by the Deutsche Forschungsgemeinschaft, project KR

2942/1 is gratefully acknowledged.

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How does non-covalent Se� � �SeQO interaction stabilize selenoxides at

naphthalene 1,8-positions: structural and theoretical investigationsw

Satoko Hayashi,a Waro Nakanishi,*a Atsushi Furuta,a Jozef Drabowicz,b

Takahiro Sasamoricand Norihiro Tokitoh

c

Received (in Montpellier, France) 9th June 2008, Accepted 25th September 2008


DOI: 10.1039/b809763a

Bis-selenides (LL), such as 8-[MeSe(X)]-1-[MeSe(Z)]C10H6 (1 (LL)), 8-[EtSe(X)]-1-[EtSe(Z)]C10H6

(2 (LL)), 8-[p-YC6H4Se(X)]-1-[MeSe(Z)]C10H6 (3 (LL)) and 8-[p-YC6H4Se(X)]-1-[p-YC6H4Se(Z)]C10H6

(4 (LL)) were oxidized with ozone at 0 1C, where (X, Z) = (lone pair, lone pair) for LL.

Bis-selenoxides, 1 (OO), 3 (OO) and 4 (OO) where (X, Z) = (oxygen, oxygen), were obtained in

the oxidation of 1 (LL), 3 (LL) and 4 (LL), respectively, via corresponding selenide-selenoxides,

1 (LO), 3 (LO) and 4 (LO), respectively. A facile Se–C bond cleavage was observed in 2 (LL).

The structures of 1 (LO) and 1 (OO) were determined by the X-ray analysis. Three Se� � �SeQO

atoms in 1 (LO) and four OQSe� � �SeQO atoms in 1 (OO) align linearly. While the non-covalent

Se� � �SeQO 3c–4e interaction operates to stabilize 1 (LO), the non-covalent OQSe� � �SeQO 4c–4e

interaction would not stabilize 1 (OO). The 3c–4e interaction must play an important role to

control the stereochemistry of selenoxides. The 8-G-1-[MeSe(OH)2]C10H6 (n (OH�OH)) are the key

intermediates in the racemization of 8-G-1-[MeSe(O)]C10H6 (n (O)) in solutions, where G = SeMe

(1), H (5), F (6), Cl (7) and Br (8). Energies of n (OH�OH), relative to n (O), are evaluated based

on the theoretical calculations. G of SeMe is demonstrated to operate most effectively to protect

from racemization of selenoxides among n = 1 and 5–8, since the relative energies

for G of cis- and trans-SeMe are largest.

Introduction

Selonoxides1–4 [RSe(O)R0] afford optically active enantiomers,

as well as sulfoxides,5,6 if R and R0 are not the same, since Se

in each selenoxide is three-coordinated containing a lone pair.

However, it is usually difficult to utilize optical active selen-

oxides to introduce the optical activity in a target mole-

cule,1,2,4,7 since the racemization of optical active selenoxides

is usually very fast. Nevertheless, some efforts have been made

to stabilize the stereochemistry of selenoxides, by taking

advantage of non-covalent coordination by the neighboring

groups (G) of the G� � �SeQO type.2,4,7

Naphthalene 1,8-positions supply a good system to investi-

gate such interactions, since the non-bonded distances between

heteroatoms at the positions are close to the sum of the van

der Waals radii minus 1 A.8,9 Various types of non-covalent

interactions are detected in naphthalene 1,8-positions.8–11 The

s-type three center-four electron interactions (s(3c–4e)),12–14

s(2c–4e),12 p(2c–4e),12 distorted p(2c–4e),12 and Z4 4c–6e13

are typical examples. Such non-covalent interactions are

demonstrated to control the fine structures of molecules.15

Recently, we investigated fine structures of 8-G-1-(arylseleninyl)-

naphthalene with G = H, F, Cl and Br, together with

the factors to control the structures, as the first step to

control the stereochemistry of selenoxides.16 The factors are

called G, O and Y dependences, which originate from the

np(G)� � �s*(Se–O), np(O)� � �p(Nap) and np(O)� � �p(Ar) interac-

tions, respectively.16

We paid much attention to G = MeSe and ArSe in 8-G-1-

(arylseleninyl)naphthalenes, since many conformers are plau-

sible around the two Se–CNap bonds, relative to the case of

G = H and halogens. Scheme 1 shows the orbitals taking part

in the non-covalent Se� � �SeQO interaction. A bis-selenide

Scheme 1 Orbitals taking part in the non-bonded Se� � �SeQO inter-

actions in naphthalene 1,8-positions.

aDepartment of Material Science and Chemistry, Faculty of SystemsEngineering, Wakayama University, 930 Sakaedani, Wakayama640-8510, Japan. E-mail: [email protected];Fax: +81 73 457 8253; Tel: +81 73 457 8253

bCenter of Molecular and Macromolecular Studies, Polish Academyof Science, Sienkiewicza, 112, 90-363 Lodz, Poland

c Institute for Chemical Research, Kyoto University, Gokasho, Uji,Kyoto 611-0011, Japanw Electronic supplementary information (ESI) available: Energies andrelative energies for 8-G-1-[MeSe(X)]C10H6 [G = MeSe (1), H (5), F(6), Cl (7) and Br (8) with X = lone pair (L), O (O), OH+ (OH+) andO2H2 (OH�OH)], the packing structures of 1 (OO), counter map for1 (OO), Cartesian coordinates for optimized structures of 1 and 5–8

with X = lone pair (L), O (O), OH+ (OH+) and O2H2 (OH�OH)].CCDC reference numbers 688690 (1 (LO)) and 688691 (1 (OO)). ForESI and crystallographic data in CIF or other electronic format seeDOI: 10.1039/b809763a



contains double ns(Se), np(Se), s(Se–C) and s*(Se–C) orbitals.However, ns(O), np(O), np0(O), s(Se–O) and s*(Se–O) appear

newly with the quit of an np(Se), when a selenide-selenoxide is

formed from the bis-selenide.

The oxidation and formation of 8-[2RSe(X)]-1-

[1RSe(Z)]C10H6 (1 (1R = 2R = Me), 2 (1R = 2R = Et), 3

(1R =Me, 2R = p-YC6H4: Y = H (a), MeO (b) and NO2 (d))

and 4 (1R = 2R = p-YC6H4: Y = H (a) and tBu (c)) are

investigated for LL where (X, Z) = (lone pair, lone pair), LO

(lone pair, oxygen) and OO (oxygen, oxygen) (Chart 1). The

reactions are easily controlled and each process is followed by

the spectroscopic method. Non-bonded OQSe� � �SeQO

interactions are also the subject of interest.

The structures around the naphthyl group (Nap) in 8-G-1-

RSeC10H6 are well explained by three types, type A (A), B and

C.8c,d,f–h,17,18 The combined notation are used to specify the

structures of 1–4 with G = RSe, where the notation, such as

AA, BA or CA, shows the conformers around the two CNap–Se

bonds. Scheme 2 draws the notations employed in this work,

exemplified by 1 (LO).

The structures of 1 (LO) and 1 (OO) are determined by

X-ray crystallographic analysis. Quantum chemical (QC) cal-

culations are performed on 1 (LO) and 1 (OO), to elucidate the

role of the Se� � �SeQO interaction in 1 (LO) and the

OQSe� � �SeQO interaction in 1 (OO) as the factor to control

the fine structures. Orbitals of two Se atoms in 1 (LO) and

1 (OO) must overlap directly with each other, which would

stabilize the fine structures. QC calculations are also

performed on 8-G-1-[MeSe(X)]C10H6 [G = MeSe (1), H (5),

F (6), Cl (7) and Br (8) with X = lone pair (L), O (O), OH+

(OH+) and O2H2 (OH�OH)], where OH�OH must be the key

intermediate in the racemization of 1 and 5–8, in the presence

of a trace of water. The relative energy [DE = E(n (OH�OH))

� (E(n (O)) + E(H2O)) (n = 1 and 5–8)] is evaluated: that for

G = MeSe is largest among them. The larger value must

correspond to a selenoxide with the stronger resistance for

racemization, although n (OH�OH) is not the transition state.

The G� � �SeQO interactions containing the Se� � �SeQO

and OQSe� � �SeQO interactions are also analyzed with the

natural orbital (NBO)19,20 and atoms-in-molecules (AIM)21,22

analyses.

Oxidation of 1,8-bis(selanyl)naphthalenes (LL) with ozone

is well controlled and monitored, which gives 1,8-bis(seleninyl)-

naphthalenes (OO) via 8-selanyl-1-seleninylnaphthalenes (LO).

Factors to control the fine structures of 1 (LO) and 1 (OO) are

clarified based on QC calculations, after determination of the

structures. The Se� � �SeQO interaction is demonstrated to

control the fine structure of 1 (LO), whereas the role of the

OQSe� � �SeQO interaction in 1 (OO) is critically discussed.

The role of G in 1 and 5–8 in the racemization process is also

evaluated.


Survey of oxidation

Bis-selenides (n (LL): n= 1–4) were oxidized with ozone in the

methylene dichloride solution of each bis-selenide at 0 1C. The

bis-selenides (n (LL)) gave corresponding bis-selenoxides

(n (O)) via corresponding selenide-selenoxides (n (LO)), except

for 2 (LL). While 1 (LL) gave 1 (LO), followed by the quanti-

tative formation of 1 (OO), a facile Se–C bond cleavage

occurred on the oxidation of 2 (LL), resulting in the formation

of naphtho-1,8-[c,d]-1,2-diselenole (9).23 b-Elimination of the

selenoxide may be responsible for the facile Se–C bond

cleavage. In the case of 3 (LL), the methylselanyl Se atoms

were attacked exclusively. 3 (LO) were consumed to produce

the corresponding 3 (OO) with more ozone. 4 (LO) were also

produced from the corresponding 4 (LL) with ozone, followed

by the formation of the corresponding 4 (OO), respectively.

The results are summarized in Chart 1. The reactions are well

followed by NMR.

Structures of 1 (LO) and 1 (OO)

Single crystals of 1 (LO) and 1 (OO) were obtained via slow

evaporation of methylene dichloride-hexane solutions and one

of suitable crystals was subjected to X-ray crystallographic

analysis for each compound.24 Only one type of structure

corresponds to each of 1 (LO) and 1 (OO) in the crystals.

Table 1 shows the crystallographic data of 1 (LO) and

1 (OO). Fig. 1 shows the structures of 1 (LO) and 1 (OO).25

The packing structure of 1 (OO) is shown in Fig. S1 of the

Electronic Supplementary Information (ESIw). Selected inter-

atomic distances, angles and torsional angles of the com-

pounds 1 (LO) and 1 (OO) are collected in Table 2, together

with those of 1 (LL), which contains two types, 1 (LL)Aand 1 (LL)B.

26

Chart 1 Bis(selanyl)naphthalenes, 1–4, together with 5–8.

Scheme 2 Structures around naphthyl group in 8-G-1-[RSe(X)]C10H6,

exemplified by 1 (LO).


The structures of 1 (LO) and 1 (OO) are all AA for two

methyl groups (Fig. 1 and Table 2).15 The planarity of the

naphthyl (Nap) planes is very good. All Se–O bonds are placed

in the naphthyl plane. The superior tendency of the Se–O

bonds to stay on the naphthyl plane (O dependence)16 must be

the driving force for the structures of 1 (LO) and 1 (OO). Three

Se� � �Se–O atoms align linearly (+SeSeO = 173.31(15)1) and

the Se–O bond is almost perpendicular to another CNapSeCMe

plane in 1 (LO). The non-covalent np(Se)� � �s*(Se–O) 3c–4e

interaction operates effectively to keep the Se–O bond on the

naphthyl plane in 1 (LO) (G dependence).16 These results show

that the structure of 1 (LO) is well stabilized by the O and G

dependences observed in 1-naphthyl selenoxides.16

On the other hand, there is no np(Se) in 1 (OO). Therefore,

the G dependence of the np(Se)� � �s*(Se–O) type cannot

operate in 1 (OO). Consequently, the driving force for the

structure must come from the O dependence for both Se–O

bonds. Namely, the non-covalent O–Se� � �Se–O s(4c–4e) in-

teraction must be carefully examined as a factor to stabilize the

fine structure of 1 (OO), although the non-bonded Se� � �Sedistances are less than the sum of van der Waals radii by

ca. 0.65 A.27 The s(4c–4e) interaction seems not so important.

How does G of MeSe control the fine structure and the

behavior? QC calculations are performed on 1 and 5–8.

QC calculations

QC calculations were performed on 1 (LO) with the B3LYP/

6-311+G(d) method of the Gaussian 98 program.28–30 QC

calculations revealed energy profiles of the compounds.31

Table 3 collects the results of the QC calculations. The NBO

analysis19,20 were performed on 1 (LO) and 1 (OO) with the

B3LYP/6-311+G(d) method. The results are shown in

Table 4. The AIM parameters21,22 are calculated for 1 (LO)

and 1 (OO) with the Gaussian 03 program32 employing the

6-311+G(3df) basis sets for Se with the 6-311+G(3d,2p) basis

sets for C and H at the B3LYP level. They are analyzed

employing the AIM 2000 program.33 Table 5 collects the

results of AIM calculations.

Indeed, the results of QC calculations essentially correspond

to those in the gas phase, but the factors to control and/or

stabilize the structures in gas phase must also operate in solid

states and in solutions. Therefore, it must be instructive to

consider those predicted by QC calculations, although we

must be careful for the crystal packing effect in crystals and

the solvent effect in solutions, since such effects often larger

than the predicted factors.

The effect of G to stabilize 8-G-1-[MeSe(X)]C10H6 [G =

MeSe (1), H (5), F (6), Cl (7) and Br (8) with X = lone pair

(L), O (O), OH+ (OH+) and O2H2 (OH�OH)] will be dis-

cussed in detail, here. The results clarified the factors for the

racemization of selenoxides. n (OH�OH) (n = 1 and 5–8) must

be the key intermediates in the racemization of n (O), in the

presence of (a trace of) water in solutions.

Effect of G in 1 and 5–8

Racemization of an optically active selenoxide is believed to

proceed via a selenide dihydroxide (n (OH�OH)).4a–d Scheme 3

shows a hypothetical racemization process of optically active

n (O*) via n (OH�OH).

Protonation of n (O*) occurs at O of an optically active

isomer of n (O*: R) to give n (O*H+: R) at the initial stage of

the reaction. n (OH�OH) will form in the reaction of n (O*H+:

R) with water followed by the deprotonation to yield n (OH�OH). Elimination of water from n (OH�OH) results in the

racemization, since of n (OH�OH) is not optically active as a

whole. Similar reactions occur starting from n (O*: S) to yield

n (OH�OH) via n (O*H+: S), which also leads to racemization.

Table 1 Crystallographic data for 1 (LO) and 1 (OO)

1 (LO) 1 (OO)

Empirical formula C12H12OSe2 C12H12O2Se2�2.5H2OFormula weight 330.14 391.18Temperature/K 298(2) 103(2)Crystal system Monoclinic MonoclinicSpace group P21/n (#14) C2/c (#15)a/A 5.8460(19) 25.549(9)b/A 14.473(3) 5.8653(18)c/A 14.1490(16) 20.850(8)b/1 97.660(17) 117.329(4)V/A3 1186.5(5) 2775.6(16)Z 4 8Dc/g cm�3 1.848 1.872F(000) 640 1544Reflections observed [I 4 2s(I)] 2200 2435Parameters 136 190R1 [I 4 2s(I)] 0.032 0.021R1 [all data] 0.082 0.022oR2 [I 4 2s(I)] 0.065 0.053oR2 [all data] 0.077 0.054Goodness-of-fit on F2 1.029 1.109

Fig. 1 Structures of 1 (LO) (a) and 1 (OO) (b) with atomic numbering

scheme for selected atoms (thermal ellipsoids are shown at the 50%

probability level).


Table 2 Selected interatomic distances (A), angles (1) and torsional angles (1) around Se atom in 1 (LO) and 1 (OO), together with those of 1 (LL)

1 (LL)Aa

1 (LL)Ba

1 (LO) 1 (OO)

Interatomic distancesSe1–Se2 3.051(4) 3.064(4) 3.1587(10) 3.1512(8)Se1–C1 1.929(4) 1.932(3) 1.983(5) 1.959(2)Se1–C11 1.944(4) 1.953(4) 1.954(5) 1.940(2)Se1–O1 1.653(4) 1.6771(15)Se2–C9 1.926(4) 1.932(4) 1.928(5) 1.970(2)Se2–C12 1.944(4) 1.949(4) 1.938(6) 1.934(2)Se2–O2 1.680(15)

AnglesSe2–Se1–C11 164.47(3) 146.46(3) 85.93(16) 88.26(7)Se2–Se1–O1 173.31(15) 167.54(5)Se1–Se2–C12 150.34(3) 159.73(3) 85.65(18) 89.41(7)Se1–Se2–O2 167.51(5)Se1–C1–C10 122.9(3) 123.9(3) 126.9(4) 124.33(16)C1–Se1–C11 99.29(16) 98.41(16) 96.0(2) 94.73(9)C1–Se1–O1 101.1(2) 102.69(8)C11–Se1–O1 100.7(2) 102.85(9)Se2–C9–C10 123.9(3) 122.9(3) 124.1(4) 124.67(16)C9–Se2–C12 99.27(16) 98.50(16) 98.1(2) 93.38(9)C9–Se2–O2 102.21(8)C12–Se2–O2 102.29(9)C1–C10–C9 126.4(3) 127.2(3) 127.0(4) 128.1(2)

Torsional anglesSe1–C1–C10–C5 173.5(2) �176.0(2) �177.6(4) 179.19(15)C10–C1–Se1C11 �154.1(3) 136.8(3) 82.8(4) �86.49(19)C10–C1–Se1–O1 �175.0(4) 169.19(17)Se2–C9–C10–C5 172.2(2) �170.2(2) 178.8(4) 178.90(15)C10–C9–Se2–C12 �138.8(3) 148.0(3) 84.6(4) �87.10(19)C10–C9–Se2–O2 169.54(18)O1–Se1–Se2–O2 140.3(3)

a Ref. 26.

Table 3 Energies and relative energies for 8-G-1-[MeSe(OiHj)]C10H6

(i, j = 0, 1 and 2)a

Form O: A/AAb OH+: A/AAb OH�OH: AC

5 (G = H) �2902.0303 �2902.4017 �2978.4686Qn(Se) 1.309 1.307 1.324Qn(O) �0.968 �0.837 �0.996, �0.993Qn(H) 0.497 0.433, 0.433+Wc �2978.4741 �2978.2198 �2978.4686Dd,e as 0.0 667.7 (as 0.0) 14.4 (as 0.0)6 (G = F) �3001.3000 �3001.6744 �3077.7371+Wc �3077.7438 �3077.4925 �3077.7371Dd,e as 0.0 659.8 (�7.9) 17.6 (3.2)7 (G = Cl) �3361.6478 �3362.0250 �3438.0833+Wc �3438.0916 �3437.8431 �3438.0833Dd,e as 0.0 652.4 (�15.2) 21.8 (7.4)8 (G = Br) �5475.5651 �5475.9442 �5552.0007+Wc �5552.0089 �5551.7623 �5552.0007Dd,e as 0.0 647.4 (�20.2) 21.5 (7.1)1 (G = trans-MeSe) �5342.8896 �5343.2854 �5419.3241+Wc �5419.3334 �5419.1035 �5419.3241Dd,e as 0.0 603.6 (�64.1) 24.4 (10.0)1 (G = cis-MeSe) �5342.8869 �5343.2821 �5419.3214+Wc �5419.3307 �5419.1002 �5419.3214Dd,e as 0.0f 612.3 (�55.4) 31.5 (17.1)

a Calculated with the B3LYP/6-311+G(d) method. bA for 5–8 and

AA for 1. c Evaluated based on the values of E(H2O2) = �151.5891 au,

E(H2O) = �76.4438 au and E(OH�) = �75.8181 au calculated with

the same method. d Relative to that of the corresponding n (O): A.e Relative to the same structure derived from 5 (G = H) being

given in parenthesis. f 7.1 kJ mol�1 from the corresponding species of

1 (G = trans-MeSe; O): AA.

Table 4 Second order perturbation energies in the donor (D)–accep-tor (A) interactions of the n(G)� � �s*(Se–O) type in 8-G-1-[MeSe(O)]-C10H6 and 8-G-1-[MeSe+(OH)]C10H6, calculated with the NBOmethodab

D; A np(G); s*(Se–O) np(G): s*(Se+–OH)

G = F 1.44 9.15 (0.87)c

G = Cl 3.29 13.65 (1.09)c

G = Br 3.73 27.95 (1.19)c

G = cis-SeMe 4.77d 34.99 (1.76)c

G = trans-SeMe 5.52 41.86 (2.69)c

G = trans-SeMee 5.86G = trans-Se(O)Mee 1.53 (�2)fa The 6-311+G(d) basis sets being employed. b In kcal mol�1. c Corres-

ponding to the ns(G)� � �s*(Se+–OH) interaction. d 0.76 kcal mol�1 for

the ns(Se)� � �s*(Se–O) type interaction. e The 6-311+G(3df) basis sets

being employed for Se with the 6-311+G(3d,2p) basis sets for C and

H. f Corresponding to the ns(Se)� � �s*(Se–O) interactions.

Table 5 Second order perturbation energies in the donor–acceptorinteractions of the n(G)� � �s*(Se–O) type at the naphthalene 1,8-positions in 1 (LO) and 1 (OO), calculated with the NBO methoda

Compound ro(Se, Se)/A rb(rc)/eao�3 Drb(rc)/eao

�5 Hb(rc)/au

1 (LO) 3.2521 0.0195 0.0420 �0.00051 (OO) 3.2851 0.0160 0.0393 0.0002

a The 6-311+G(3df) basis sets being employed for Se and the

6-311+G(3d,2p) basis sets for C and H.


Water may originate from the solvent and the racemization

would proceed under the neutral conditions. The stability of

n (OH�OH) must affect on the rates of racemization for the

optical active selenoxides.

The effect of G on the stability of 8-G-1-[MeSe(OiHj)]C10H6

[1 and 5–8: L (i = j = 0), O (i = 1, j = 0), OH+ (i = j = 1)

and OH�OH (i = j = 2)] are examined based on the QC

calculations. The results of QC calculations performed with

the B3LYP/6-311+G(d) method are collected in Table 3.

Table 3 also contains natural charges (Qn) of Se and O

calculated employing the natural population analysis.20

Scheme 4 shows optimized structures of the global minimum

for each of 1 (LL), 1 (LO) and 1 (LOH+), together with the

three types, AA0, BB and AC, for 1 (LOH�OH). The values for

AC of n (LOH�OH) are given in Table 3, since AC is most

stable among the three for each.34

Energy differences of the reactions in Scheme 3 are exam-

ined based on the values shown in Table 3. The energy of n (O) +

H2O (E(n (O) + H2O)) is taken as the standard for each, for

convenience of comparison. How are the selenoxides stabilized

by G at the 8-position? The effect of G on the stabilization of

selenoxides is examined before discussion the energy profile

shown in Scheme 3.

Eqn (1) shows the energies of n (L) + H2O2 (E(n (L) +

H2O2)) relative to E(n (O) + H2O) [DE(n (LO) = E(n (L) +

H2O2) � E(n (O) + H2O)) (see Table S1 in the ESIw). Simi-

larly, eqn (2) and (3) exhibit DE(n (OH+)) and DE(n (OH�OH)),35

respectively, which are defined as [E(n (OH+) + HO�) �E(n (O) + H2O)] and [E(n (OH�OH)) � E(n (O) + H2O)],

respectively.36

DE(n (LO)) = E(n (L) + H2O2) � E(n (O) + H2O)

G = H (121.8 kJ mol�1) o F (131.3) o cis-MeSe (133.9)

o Cl (136.8) r Br (137.6) o trans-MeSe (141.0) (1)

DE(n (OH+)) = E(n (OH+) + HO�) � E(n (O) + H2O)

G = H (667.7 kJ mol�1) 4 F (659.8) 4 Cl (652.4) 4 Br

(647.4) c cis-MeSe (605.2) 4 trans-MeSe (603.6) (2)

DE(n (OH�OH)) = E(n (OH�OH)) � E(n (O) + H2O)

G = H (14.4 kJ mol�1) o F (17.6) o Cl (21.8) E Br (21.5)

o trans-MeSe (24.4) o cis-MeSe (31.5) (3)

The order in eqn (1) corresponds the energy lowering effect

by the G� � �Se–O interactions in the formation selenoxides

relative to the G� � �Se–C interactions in selenides. However, we

must be careful to examine the values for G = cis-MeSe and

trans-MeSe, since the structure of the corresponding selenide is

commonly CC (see Table S1 in the ESIw).Eqn (2) exhibits that the protonation on the seleninyl O

atom occurs more easily in the order of G = H o F o Cl oBr { cis-MeSe o trans-MeSe. The results show that the

protonation occurs more easily when G become better donors,

especially for G = MeSe. The evaluated DE(n (OH+)) values

are very large in magnitudes, however, they do not mean that

the process is very difficult to occur. The large magnitudes are

the results of the calculations for the charge separated species

of the n (OH+) + HO� type. Only the relative values are

important, since protonation will occur easily in solutions.

Resulting hypervalent np(G)� � �s*(Se–OH+) interactions sta-

bilize further the species in the order shown in eqn (2), relative

to the case of the selenoxides.

The activation energies for the racemization of optically

active selenoxides are closely related to the values shown in

eqn (3), although they are the energies for the intermediates,

n (OH�OH). The activation energies are expected to increase in

this order. The activation energy for G = cis-MeSe is pre-

dicted to be larger than that with trans-MeSe. However, cis-

MeSe and trans-MeSe isomers interconvert with each other.

Therefore, it may be better to evaluate the value by G= trans-

MeSe under the experimental conditions: The activation

energy of 1 (LO) with G = MeSe is estimated to be about

10 kJ mol�1 larger than that of 5 (L) with G = H and the

former is also larger than the case of G = Br by ca. 3 kJ mol�1.

Fig. 2 summarizes the effect of G given in eqn (2).

G at the 8-position will protect sterically from the racemiza-

tion of an optical active n (O*). G must stabilize the optical

active n (O*) and the protonated n (O*H+) whereas G would

destabilize n (OH�OH). The steric congestion at the backside

of the Se+–OH bond in n (O*H+) by G will block the space

for H2O to attack to produce n (OH�OH) (Scheme 4). We must

be careful, since G could also stabilize n (OH�OH) in some

Scheme 3 Mechanism for racemization of n (O*) via n (OH�OH) (n= 1

and 5–8).

Scheme 4 Optimized structures for 1 (G = SeMe) and the deriva-

tives.


cases. The calculated values might correspond to the total

effects of the electronic and steric effects. Energy profiles for

the racemization evaluated by above calculations must contain

main factors. The energies for the transition states must be

close to those of the intermediates, n (OH�OH).

NBO analysis for n(G)� � �r*(Se–O) interactions

Table 4 summarizes the second order perturbation energies

(E(2)) for the charge transfer (CT) interactions of the

n(G)� � �s*(Se–O) type in 8-G-1-[MeSe(O)]C10H6 and 8-G-1-

[MeSe+(OH)]C10H6 evaluated with the NBO method.37 The

B3LYP/6-311+G(d) method is employed for the calculations.

The E(2) values becomes larger in an order shown in eqn (4).

E(2): G = F (1.44) o Cl (3.29) o Br (3.73) o cis-MeSe

(4.77) o trans-MeSe (5.52) (4)

The E(2) values are also evaluated for 1 (LO) and 1 (OO),

employing the 6-311+G(3df) basis sets for Se and the

6-311+G(3d,2p) basis sets for C and H at the B3LYP level.38

Table 4 also contains the values. The np(G)� � �s*(Se–O) inter-

action are evaluated to be 5.9 kcal mol�1 for 1 (LO)39 and as

1.5 (� 2) kcal mol�1 for the ns(G)� � �s*(Se–O) interactions in

1 (OO). The larger value for 1 (LO) relative to 1 (OO) implies

the more effective interaction of the np(G)� � �s*(Se–O) type in

1 (LO). The contribution of the 4c–4e interaction of the

O–Se� � �Se–O type was not detected by the NBO analysis.

Fig. 3 summarizes the interactions.

The nature of the np(G)� � �s*(Se–O) interaction in 1 (LO)

and the ns(G)� � �s*(Se–O) interactions in 1 (OO) are evaluated

based on the AIM analysis, next.

AIM analysis of 1 (LO) and 1 (OO)

The AIM analysis are carried out on 1 (LO) and 1 (OO). The

6-311+G(3df) basis sets are employed for Se and the

6-311+G(3d,2p) basis sets for C and H at the B3LYP level.

Table 5 collects the AIM parameters of 1 (LO) and 1 (OO) for

the bond critical points (BCPs: rc) on the interaction lines

between non-bonded Se atoms.

The low values of electron densities at BCPs (rb(rc)) in

1 (LO) and 1 (OO) (0.016–0.020 eao�3) show that the interac-

tions are ionic in nature. Laplacian values of rb(rc) (Drb(rc))are both positive, whereas the total electron energy densities at

BCPs (Hb(rc)) for 1 (LO) is negative but it is positive for

1 (OO). The results strongly suggest that the np(G)� � �s*(Se–O)

interaction in 1 (LO) is the CT interaction in nature similarly

to the case of R2Se� � �Br2 (MC) but the ns(G)� � �s*(Se–O)

interactions in 1 (OO) seems weaker than such CT interac-

tions.40

Fig. 4 shows the counter map of rb(rc) in the SeSeC9 plane

for 1 (LO), together with BCPs ( ), ring critical points ( ),

bond paths and the interaction lines. BCP are detected on the

Se� � �Se and O� � �2H interaction lines. The BCP on the Se� � �Seinteraction line well visualize the np(Se)� � �s*(Se–O) interac-

tion in 1 (LO). While BCP is also detected on the O� � �2Hinteraction line, the interaction is very small. A similar counter

map is also drawn for 1 (OO), which is shown in Fig. S2 of the

ESI.w

Conclusion

X-Ray crystallographic analysis of 8-methylselanyl-1-(methyl-

seleninyl)naphthalene (1 (LO)) and 1,8-bis(methylseleninyl)-

naphthalene 1 (OO) revealed that the three Se� � �SeQO

atoms in 1 (LO) and the four OQSe� � �SeQO atoms in

1 (OO) align linearly. All Se–O bonds are placed in the

naphthyl plane. The superior tendency for the Se–O bonds

to stay on the naphthyl plane (O dependence) must be the

Fig. 2 Energies of n (OH�OH), relative to n (O) for n = 1 and 5–8.

Fig. 3 The np(G)� � �s*(Se–O) interaction in 1 (LO) and the

ns(G)� � �s*(Se–O) interactions in 1 (OO) evaluated by the NBOmethod.

Fig. 4 Contour map of rb(rc) for 1 (LO) in the SeSeC9 plane, together

with BCPs ( ), ring critical points ( ) and bond paths. The contours

[eao�3] are at 2l (l = �8, �7,. . .0) and 0.0047 (heavy line). Two Me

groups are located upside and downside of the SeSeC9 plane. The C2,

C3, C6 and C7 atoms with the C–H bonds deviate substantially from

the plane.


driving force for the fine structures of 1 (LO) and 1 (OO). The

noncovalent np(Se)� � �s(Se–O) 3c–4e interactions (G depen-

dence) operate effectively to stabilize the structure of 1 (LO).

On the other hand, the driving force for the structure of 1 (OO)

must mainly come from the O dependence for each Se–O bond

in 1 (OO), since the G dependence cannot operate without

np(Se).

QC calculations clarify the factors that protect from race-

mization of selenoxides. The energies of 8-G-1-[MeSe(OH)2]-

C10H6 from (8-G-1-[MeSe(O)]C10H6 + H2O) are shown to be

in an order of G = H (14.4 kJ mol�1) o F (17.6) o Cl (21.8)

E Br (21.5) o trans-SeMe (24.4) o cis-SeMe (31.5). The

activation energies for the racemization should increase in this

order, since 8-G-1-[MeSe(OH)2]C10H6 must be the key inter-

mediates. The activation energy of 1 (LO: G = MeSe) is

evaluated to be larger than that of 5 (L: G = H) and 8

(L: G = Br) by 10 and 3 kJ mol�1, respectively. The results

will help to design the optically stable selenoxides. The NBO

and AIM analyses support the discussion and visualize the

interactions.

Investigations on the chiral 3a (LO), prepared in the oxida-

tion of 3a (LL) with chiral reagents, are in progress. Details

will be reported elsewhere.

Experimental

General considerations

Manipulations were performed under an argon atmosphere

with standard vacuum-line techniques. Glassware was dried at

130 1C overnight. Solvents and reagents were purified by

standard procedures if necessary. Melting points were deter-

mined on a Yanaco MP-S3 melting point apparatus and

uncorrected. NMR spectra were recorded at room tempera-

ture on a JEOL AL-300 spectrometer (1H, 300 MHz; 13C,

75 MHz) and on a JEOL Lambda-400 spectrometer (1H, 400

MHz; 77Se, 76 MHz). The 1H, 13C and 77Se NMR spectra were

recorded in CDCl3. Chemical shifts are given in ppm relative

to Me4Si for the 1H and 13C NMR spectra and relative to

reference compound MeSeMe for the 77Se NMR spectra.

Column chromatography was performed by using silica gel

(Fujishilysia PSQ-100B) and basic alumina (E. Merck) and

analytical thin layer chromatography was performed on pre-

coated silica gel plates (60F-254) with the systems (v/v)

indicated.

Syntheses

Bis(methylselanyl 1,8-bis(methylselanyl)naphthalene (1 (LL)).

To a solution of the dianion of naphtho[1,8-c,d]-1,2-diselenole,

which was prepared by reduction of the diselenole 923 (1.03 g,

3.64 mmol) with NaBH4 in an aqueous THF solution, was

added methyl iodide (1.29 g, 9.06 mmol) at room temperature.

After a usual workup, the crude was purified by column

chromatography (flash column, SiO2, hexane). Recrystalli-

zation of the chromatographed product from hexane gave

1 (LL) as colorless prisms in 98% yield, mp 85.0–85.5 1C, 1H

NMR (300MHz, CDCl3, d, ppm, TMS): 2.33 (s, 6H), 7.32 (t, 2H,

J = 7.7 Hz), 7.70 (dd, 2H, J = 1.2 and 8.2 Hz), 7.73 (dd, 2H,

J = 1.2 and 7.5 Hz); 13C NMR (75 MHz, CDCl3, d, ppm,

TMS): 13.3, 125.7, 128.3, 131.9, 132.3, 135.3, 135.6; 77Se NMR

(76 MHz, CDCl3, d, ppm, Me2Se): 231.4. Anal. Calc. for

1 (LL) (C12H12Se2): C, 45.88; H, 3.85%. Found: C, 45.73;

H, 3.77%.

8-Methylselanyl-1-(methylseleninyl)naphthalene (1 (LO)).

1 (LL) (0.98 mg, 3.12 mmol) was dissolved in 20 mL of CH2Cl2and the solution was bubbling with the ozone for 5 min. TLC

was checked for the completion of the reaction (rf = 0.07

(chloroform)). Then the solution was evaporated and dried

in vacuo. The crude product was purified by column chromato-

graphy (flash column, Al2O3, CH2Cl2). 1 (LO) gave 85% yield

as colorless powder, mp 129.8–130.1 1C; 1H NMR

(400 MHz, CDCl3, d, ppm, TMS): 2.29 (s, 3H), 2.78 (s, 3H),

7.48 (t, J 7.6 Hz, 2H), 7.76 (t, J 7.7 Hz, 2H), 7.98–8.05 (m, 2H),

8.10 (dd, J 1.1 and 7.2 Hz, 1H), 8.88 (dd, J 1.2 and 7.4 Hz,

1H); 13C NMR (75 MHz, CDCl3, d, ppm, TMS): 13.87, 41.12,

125.73, 126.28, 126.35, 126.57, 131.01, 132.44, 133.06, 136.13,

138.93, 141.34; 77Se NMR (76 MHz, CDCl3, d, ppm, Me2Se):

210.8, 833.0. Anal. Calc. for 1 (LO) (C12H12OSe2): C, 43.66; H,

3.66%. Found: C, 43.61; H, 3.60%.

1,8-Bis(methylseleninyl)naphthalene (1 (OO)). 1 (LL) (0.58 g,

0.30 mmol) was dissolved in 20 mL of CH2Cl2 and the solution

was bubbling with the ozone for 15 min. TLC was checked for

the completion of the reaction (rf = 0.00 (chloroform)). Then

the solution was evaporated and dried in vacuo. The crude

product was purified by column chromatography (flash

column, Al2O3, CH2Cl2). 1 (OO) gave 59% yield as colorless

powder, mp 154.8–155.2 1C; 1H NMR (400 MHz, CDCl3, d,ppm, TMS): 2.71 (s, 6H), 7.84 (t, J 7.7 Hz, 2H), 8.18 (dd, J 1.2

and 6.9 Hz, 2H), 8.71 (dd, J 1.4 and 6.9 Hz, 2H); 77Se NMR


1 (OO) (C12H12O2Se2): C, 41.64; H, 3.49%. Found: C, 41.55;

H, 3.45%. Anal. Calc. for 1 (OO)�2.5H2O (C24H24O4Se4�5H2O): C, 36.84; H, 4.38%. Found: C, 36.87; H, 4.41%.

1,8-Bis(ethylselanyl)naphthalene (2 (LL)). Following the

similar method to that used for 1 (LL), 2 (LL) gave 80% yield

as colorless powder, mp 52.3–52.8 1C; 1H NMR (400 MHz,

CDCl3, d, ppm, TMS): 1.35 (t, J 7.4 Hz, 6H), 2.89 (q, J 7.5 Hz,

4H), 7.32 (t, J 7.6 Hz, 2H), 7.70 (dd, J 1.2 and 8.1 Hz, 2H),

7.76 (dd, J 1.1 and 7.2 Hz, 2H); 77Se NMR (76 MHz, CDCl3,

d, ppm, Me2Se): 341.7. Anal. Calc. for 2 (LL) (C14H16Se2): C,

49.14; H, 4.71%. Found: C, 49.23; H, 4.72%.

8-Phenylselanyl-1-(methylseleninyl)naphthalene (3a (LO)).

Following the similar method to that used for 1 (LO),

3a (LO) gave 80% yield as colorless needles, mp 129.8–130.2

1C; 1H NMR (400 MHz, CDCl3, d, ppm, TMS): 2.72 (s, 3H),

6.98–7.02 (m, 2H), 7.11–7.16 (m, 3H), 7.56 (t, J 7.6 Hz, 1H),

7.76 (t, J 7.7 Hz, 1H), 8.05 (dd, J 1.2 and 8.0 Hz, 1H), 8.10

(dd, J 1.3 and 8.1 Hz, 1H), 8.15 (dd, J 1.3 and 7.2 Hz, 1H),

8.82 (dd, J 1.3 and 7.3 Hz, 1H); 13C NMR (75 MHz, CDCl3, d,ppm, TMS) 40.56, 123.19, 126.51, 126.57, 126.75, 126.88,

128.42 (2J(Se,C) 5.9 Hz), 129.63, 131.95, 132.42, 133.03,

133.50, 136.29, 140.89, 141.37; 77Se NMR (76 MHz, CDCl3,

d, ppm, Me2Se): 398.2, 831.4. Anal. Calc. for 3a (LO)

(C17H14OSe2): C, 52.06; H, 3.60%. Found: C, 52.11;

H, 3.66%.


8-Phenylseleninyl-1-(methylseleninyl)naphthalene (3a (OO)).

Following the similar method to that used for 1 (OO),

3a (OO) gave 63% yield as colorless needles, mp 148.0–148.8 1C;1H NMR (400 MHz, CDCl3, d, ppm, TMS): 2.74 (s, 3H),

7.33–7.50 (m, 3H), 7.51–7.58 (m, 2H), 7.78 (t, J 7.7 Hz, 1H),

7.81 (t, J 7.7 Hz, 1H), 8.13 (dd, J 1.1 and 8.1 Hz, 1H), 8.14 (dd,

J 1.1 and 8.1 Hz, 1H), 8.63 (dd, J 1.4 and 7.4 Hz, 1H), 8.72

(dd, J 1.3 and 7.3 Hz, 1H); 13C NMR (75 MHz, CDCl3, d,ppm, TMS): 38.25, 126.72, 126.80, 126.85, 127.01, 127.64,

127.92, 130.02, 131.70, 133.33, 133.71, 135.55, 138.88,


820.0, 832.5. Anal. Calc. for 3a (OO) (C17H14O2Se2): C, 50.02;

H, 3.46%. Found: C, 50.07; H, 3.57%.

8-p-Anisylselanyl-1-(methylseleninyl)naphthalene (3b (LO)).


3b (LO) gave 88% yield as colorless needles, mp 129.6–130.4 1C;1H NMR (400 MHz, CDCl3, d, ppm, TMS): 2.70 (s, 3H),

3.82 (s, 3H), 6.70 (d, J 8.8 Hz, 2H), 7.01 (d, J 8.8 Hz, 2H), 7.51

(t, J 7.2 Hz, 1H), 7.74 (t, J 7.2 Hz, 1H), 7.88 (dd, J 1.1 and

6.8 Hz, 1H), 8.01 (dd, J 1.1 and 6.8 Hz, 1H), 8.03 (dd, J 1.1

and 6.8 Hz, 1H), 8.10 (dd, J 1.1 and 6.8 Hz, 1H); 77Se NMR

(76 MHz, CDCl3, d, ppm, Me2Se): 385.9, 833.9. Anal. Calc.

for 3b (LO) (C18H16O2Se2): C, 51.20; H, 3.82%. Found: C,

50.98; H, 3.83%.

8-p-Anisylseleninyl-1-(methylseleninyl)naphthalene

(3b (OO)). Following the similar method to that used for

1 (OO), 3b (OO) gave 43% yield as colorless powder, mp

144.5–145.0 1C; 1H NMR (400 MHz, CDCl3, d, ppm, TMS):

2.75 (s, 3H), 3.76 (s, 3H), 6.87 (d, J 8.9 Hz, 2H), 7.45 (d, J 8.9

Hz, 2H), 7.83 (t, J 7.7 Hz, 2H), 8.15 (dd, J 1.0, 8.2 Hz, 1H),

8.17 (dd, J 1.0, 8.2 Hz, 1H), 8.69 (dd, J 1.3, 9.1 Hz, 1H), 8.71

(dd, J 1.2, 9.1 Hz, 1H); 77Se NMR (76 MHz, CDCl3, d, ppm,

Me2Se): 821.6, 846.4. Anal. Calc. for 3b (OO) (C18H16O3Se2):

C, 49.33; H, 3.68%. Found: C, 49.30; H, 3.73%.

8-p-Nitrophenylselanyl-1-(methylseleninyl)naphthalene

(3d (LO)). Following the similar method to that used for

1 (LO), 3d (LO) gave 61% yield as colorless powder, mp


2.67 (s, 3H), 7.07 (dt, J 2.1 and 9.0 Hz, 2H), 7.64 (t, J 7.5 Hz,

1H), 7.83 (t, J 7.5 Hz, 1H), 7.99 (dt, J 2.4 and 9.0 Hz, 2H), 8.11

(dd, J 1.2 and 6.9 Hz, 2H), 8.18 (dd, J 1.2 and 4.2 Hz, 1H),


1H); 77Se NMR (76 MHz, CDCl3, d, ppm, Me2Se): 426.4,

835.6. Anal. Calc. for 3d (LO) (C17H13NO3Se2): C, 46.70; H,

3.00; N, 3.20%. Found: C, 46.75; H, 3.03; N, 3.22%.

8-p-Nitrophenylseleninyl-1-(methylseleninyl)naphthalene

(3d (OO)). Following the similar method to that used for

1 (OO), 3d (OO) gave 82% yield as colorless powder, mp


2.83 (s, 3H), 7.72–7.92 (m, 4H), 8.12–8.27 (m, 4H), 8.57 (dd,

J 1.1 and 6.2 Hz, 1H), 8.74 (dd, J 1.3 and 6.1 Hz, 1H); 77Se

NMR (76 MHz, CDCl3, d, ppm, Me2Se): 821.4, 849.0. Anal.

Calc. for 3d (OO) (C17H13NO4Se2): C, 45.05; H, 2.89; N,

3.09%. Found: C, 45.12; H, 2.83; N, 3.12%.

1,8-Bis(phenylselanyl)naphthalene (4a (LL)).Under an argon

atmosphere, 1,8-diiodonaphthalene (4.33 g, 11.40 mmol) was

dissolved in 100 mL of dry THF and the solution was added to

nBuLi (15.0 mL, 23.94 mmol, 1.6 N) at �78 1C. After 20 min,

a THF solution of phenylselenobromide (22.80 mmol) was

added to the above solution at �78 1C. Then the reaction

mixture was stirring for 2 h and warmed up room temperature.

Then, 20 mL of 5% acetone hydrochloric acid and 100 mL of

benzene were added. The organic layer was separated, washed

with brine, 10% aqueous solution of sodium hydroxide,

saturated aqueous solution of sodium bicarbonate and brine.

Then the solution was dried over sodium sulfate, evaporated

and dried in vacuo. The crude product was purified by column

chromatography (flash column, SiO2, hexane). 4a (LL) gave

89% yield as yellow prisms, mp 64.0–64.8 1C; 1H NMR

(300 MHz, CDCl3, d, ppm, TMS): 7.22–7.28 (m, 8H),

7.39–7.45 (m, 4H), 7.64 (dd, J 1.1 and 7.3 Hz, 2H), 7.74

(dd, J 1.1 and 8.3 Hz, 2H); 13C NMR (75 MHz, CDCl3,

d, ppm, TMS): 126.0, 127.4, 129.2, 129.4, 131.4, 133.4, 135.18,

135.19, 135.5, 135.9; 77Se NMR (76 MHz, CDCl3, d, ppm,

Me2Se): 435.4. Anal. Calc. for 4a (LL) (C22H16Se2): C, 60.29;

H, 3.68%. Found: C, 60.21; H, 3.75%.

8-Phenylselanyl-1-(phenylseleninyl)naphthalene (4a (LO)).


4a (LO) gave 65% yield as colorless prisms, mp 155.5–156.3 1C;1H NMR (400 MHz, CDCl3, d, ppm, TMS): 6.90–6.95

(m, 4H), 7.10–7.13 (m, 6H), 7.22–7.26 (m, 6H), 7.48–7.53

(m, 4H), 7.52 (t, J 8.2 Hz, 1H), 7.83 (d, J 7.7 Hz, 1H), 8.07

(dd, J 1.3 and 7.2 Hz, 1H), 8.08 (dd, J 1.3 and 8.2 Hz, 1H), 9.02

(dd, J 1.3 and 7.3 Hz, 1H); 13C NMR (75 MHz, CDCl3, d, ppm,

TMS): 123.90, 126.45, 126.51, 126.79, 127.86, 127.91, 128.70,

129.17, 129.49, 130.11, 131.73, 132.76, 133.27, 133.78, 136.31,

140.17, 140.71, 146.21; 77Se NMR (76 MHz, CDCl3, d, ppm,

Me2Se): 400.1, 863.7. Anal. Calc. for 4a (LO) (C22H16OSe2):

C, 58.17; H, 3.55%. Found: C, 58.11; H, 3.65%.

1,8-Bis(phenylseleninyl)naphthalene (4a (OO)). Following

the similar method to that used for 1 (OO), 4a (OO) gave

78% yield as colorless prisms, mp. 187.5–188.3 1C; 1H NMR

(400 MHz, CDCl3, d, ppm, TMS): 7.21–7.30 (m, 8H), 7.37 (tt,

J 1.5 and 6.8 Hz, 2H), 7.76 (t, J 7.7 Hz, 2H), 8.15 (dd, J 0.9 and

7.5 Hz, 2H), 8.47 (dd, J 1.1 and 6.2 Hz, 2H); 77Se NMR


4a (OO) (C22H16O2Se2): C, 56.19; H, 3.43%. Found: C, 56.22;

H, 3.53%.

1,8-Bis[(p-tert-butylphenyl)selanyl]naphthalene (4c (LL)).

Following the similar method to that used for 4a (LL),

4c (LL) gave 87% yield as yellow prisms, mp 97.8–98.3 1C;1H NMR (400 MHz, CDCl3, d, ppm, TMS): 1.30 (s 18H), 7.24

(t, J 7.7 Hz, 2H), 7.27 (d, J 8.6 Hz, 4H), 7.38 (d, J 8.6 Hz, 4H),


2H); 77Se NMR (76 MHz, CDCl3, d, ppm, Me2Se): 424.6.

Anal. Calc. for 4c (LL) (C30H32Se2): C, 65.45; H, 5.86%.

Found: C, 65.41; H, 5.88%.

8-[(p-tert-Butylphenyl)selanyl]-1-[(p-tert-butylphenyl)seleninyl]-

naphthalene (4c (LO)). Following the similar method to that

used for 1 (LO), 4c (LO) gave 86% yield as colorless powder,

mp 179.5–180.2 1C; 1H NMR (400 MHz, CDCl3, d,ppm, TMS): 1.22 (s, 9H), 1.27 (s 9H), 6.91 (d, J 8.3 Hz, 2H),


7.15 (d, J 8.1 Hz, 2H), 7.28 (d, J 8.8 Hz, 2H), 7.43 (d, J 8.6 Hz,

2H), 7.51 (t, J 7.9 Hz, 1H), 7.82 (t, J 7.7 Hz, 1H), 8.07 (d, J 8.3

Hz, 3H), 9.01 (dd, J 1.3 and 7.5 Hz, 1H); 77Se NMR (76 MHz,

CDCl3, d, ppm, Me2Se): 393.2, 861.2. Anal. Calc. for 4c (LO)

(C30H32OSe2): C, 63.61; H, 5.69%. Found: C, 63.55;

H, 5.58%.

1,8-Bis[(p-tert-butylphenyl)seleninyl]naphthalene (4c (OO)).

Following the similar method to that used for 1 (OO),

4c (OO) gave 87% yield as colorless powder, mp 172.5–173.2 1C;1H NMR (400 MHz, CDCl3, d, ppm, TMS): 1.63 (s, 18H),

7.39 (d, J 8.6 Hz, 4H), 7.47 (d, J 8.3 Hz, 4H), 7.83 (d, J 7.6 Hz,

2H), 8.16 (dd, J 0.8 and 7.3 Hz, 2H), 8.73 (dd, J 1.0 and 6.2

Hz, 2H); 77Se NMR (76 MHz, CDCl3, d, ppm, Me2Se): 843.7.

Anal. Calc. for 4c (OO) (C30H32O2Se2): C, 61.86; H, 5.54%.

Found: C, 61.93; H, 5.58%.

1-(Methylselanyl)naphthalene (5 (L)). Following the similar

method to that used for 1 (LL), 5 (L) gave 99% yield as pale

yellow oil; 1H NMR (400 MHz, CDCl3, d, ppm, TMS): 2.37

(s, 2JSe,H 11.7 Hz, 3H), 7.35 (dd, J 7.3 and 8.1 Hz, 1H), 7.47

(ddd, J 1.6, 6.9 and 8.2 Hz, 1H), 7.53 (ddd, J 1.6, 6.9 and 8.3

Hz, 1H), 7.66 (dd, J 1.1 and 7.3 Hz, 1H), 7.71 (d, J 8.2 Hz,

1H), 7.80 (dd, J 1.7 and 7.9 Hz, 1H), 8.24 (ddd, J 0.7, 1.6 and

8.1 Hz, 1H); 13C NMR (75 MHz, CDCl3, d, ppm, TMS):

36.54, 122.08 (J 14.9 Hz), 124.03 (J 6.2 Hz), 126.11, 126.79,

127.56, 129.35, 130.27, 131.40, 133.88, 138.68; 77Se NMR

(76 MHz, CDCl3, d, ppm, Me2Se): 158.6. Anal. Calc. for 5 (L)

(C11H10Se): C, 59.74; H, 4.56%. Found: C, 59.90; H, 4.49%.

1-(Methylseleninyl)naphthalene (5 (O)). Following the simi-

lar method to that used for 1 (LO), 5 (O) gave 67% yield as

colorless needles, mp 97.2–97.8 1C; 1H NMR (400 MHz,

CDCl3, d, ppm, TMS): 2.71 (s, 2JSe,H 12.3 Hz, 3H),

7.56–7.65 (m, 2H), 7.70 (dd, J 7.4 and 8.2 Hz, 1H),

7.81–7.87 (m, 1H), 7.95–8.02 (m, 2H), 8.29 (dd, J 1.1 and 7.3

Hz, 1H); 13C NMR (75 MHz, CDCl3, d, ppm, TMS): 7.52,

125.83, 126.14, 126.39, 126.63, 127.16, 128.58, 128.67, 131.10,


809.3. Anal. Calc. for 5 (O) (C11H10OSe): C, 55.71; H, 4.25%.

Found: C, 55.88; H, 4.18%.

X-Ray crystal structure determination. The colorless crystals

of 1 (LO) and 1 (OO) were grown by slow evaporation of

methylene dichloride-hexane solutions at room temperature.

A crystal of 1 (LO) was measured on a Rigaku AFC5R

diffractometer with graphite monochromated Mo-Ka radia-

tion source (l= 0.71069 A) and a rotating anode generator at

298(2) K. That of 1 (OO) was measured on a Rigaku/MSC

Mercury CCD diffractometer equipped with a graphite-mono-

chromatedMo-Ka radiation source (l=0.71070 A) at 103(2) K.

The structures of 1 (LO) and 1 (OO) were solved by direct

method (SHELXS-97)41 and refined by full-matrix least-

square method on F2 for all reflections (SHELXL-97).42 All

the non-hydrogen atoms were refined anisotropically.

QC calculations

QC calculations are performed on 1 (LO) and 1 (OO), as the

models of n (LO) and n (OO) (n = 1, 3 and 4), respectively,

employing the 6-311+G(d) basis sets of the Gaussian 98

program.27 Calculations are performed at the density func-

tional theory (DFT) level of the Becke three parameter hybrid

functionals combined with the Lee-Yang-Parr correlation

functional (B3LYP).28,29 QC calculations are also performed

on 8-G-1-[MeSe(X)]C10H6 [G = MeSe (1), H (5), F (6), Cl (7)

and Br (8) with X = lone pair (L), O (O), OH+ (OH+) and

O2H2 (OH�OH)], employing the B3LYP/6-311+G(d) method.

The NBO19,20 analysis were performed with the B3LYP/

6-311+G(d) method. The AIM21,22 analysis are performed on

1 (LO) and 1 (OO) with the Gaussian 03 program employing the

6-311+G(3df) basis sets for Se with the 6-311+G(3d,2p) basis

sets for C and H at the B3LYP level. They are analyzed

employing the AIM 2000 program.21,22 NBO analysis are also

performed on 1 (LO) and 1 (OO) with the same method for the

AIM analysis. Optimized structures and the molecular orbitals

are drawn using MolStudio R3.2 (Rev 1.0).43

Acknowledgements

This work was partially supported by a Grant-in-Aid for

Scientific Research (Nos. 16550038, 19550041 and 20550042)

from the Ministry of Education, Culture, Sports, Science and

Technology, Japan.

References

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2 T. Shimizu andN.Kamigata, Synthesis and Stereochemistry of OpticallyActive Chalcogen Compounds, in Handbook of Chalcogen ChemistryNew Perspectives in Sulfur Selenium and Tellurium, ed. F. A.Devillanova, Royal Society of Chemistry, Cambridge, 2006, ch. 10.1.

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13 (a) R. A. Hayes and J. C. Martin, Sulfurane Chemistry, in OrganicSulfur Chemistry: Theoretical and Experimental Advances, eds.F. Bernardi, I. G. Csizmadia and A. Mangini, Elsevier Scientific,Amsterdam, 1985, ch. 8; (b) J. Bergman, L. Engman and J. Siden,Tetra- and Higher-Valent (Hypervalent) Derivatives of Seleniumand Tellurium, in The Chemistry of Organic Selenium and TelluriumCompounds, eds. S. Patai and Z. Rappoport, John Wiley & Sons,New York, 1986, vol. 1, ch. 14; (c) Chemistry of HypervalentCompounds, ed. K.-y. Akiba, Wiley-VCH, New York, 1999;(d) W. Nakanishi, Hypervalent Chalcogen Compounds, in Hand-book of Chalcogen Chemistry: New Perspectives in Sulfur, Seleniumand Tellurium, ed. F. A. Devillanova, Royal Society of Chemistry,London, 2006, ch. 10.3.

14 For the 3c–4e treatment of the CT complexes, see: W. Nakanishi,S. Hayashi and H. Kihara, J. Org. Chem., 1999, 64, 2630–2637.

15 For the weak interactions, see, (a)Molecular Interactions. From vander Waals to Strongly Bound Complexes, ed. S. Scheiner, Wiley,New York, 1997; (b) K. D. Asmus, Acc. Chem. Res., 1979, 12,436–442; (c) W. K. Musker, Acc. Chem. Res., 1980, 13, 200–206.

16 S. Hayashi, H. Wada, T. Ueno and W. Nakanishi, J. Org. Chem.,2006, 71, 5574–5585.

17 W. Nakanishi, S. Hayashi and T. Uehara, Eur. J. Org. Chem.,2001, 3933–3943.

18 The structure is A if the Se–CAr bond is placed almost perpendi-cular to the naphthyl plane, it is B when the bond is located on theplane and C is the intermediate between A and B.

19 A. E. Reed, R. B. Weinstock and F. Weinhold, J. Chem. Phys.,1985, 83, 735–746; J. E. Carpenter and F. Weinhold, J. Mol.Struct. (THEOCHEM), 1988, 169, 41–62.

20 E. D. Glendening, A. E. Reed, J. E. Carpenter and F. Weinhold,NBO Ver. 3.1.

21 Atoms in Molecules. A Quantum Theory, ed. R. F. W. Bader, OxfordUniversity Press, Oxford, 1990; The Quantum Theory of Atoms inMolecules: From Solid State to DNA and Drug Design, eds.C. F. Matta and R. J. Boyd, Wiley-VCH, Weinheim, 2007, ch. 1.

22 (a) R. F. W. Bader, T. S. Slee, D. Cremer and E. Kraka, J. Am.Chem. Soc., 1983, 105, 5061–5068; (b) R. F. W. Bader, Chem. Rev.,1991, 91, 893–926; (c) R. F. W. Bader, J. Phys. Chem. A, 1998, 102,7314–7323; (d) F. Biegler-Konig, R. F. W. Bader and T. H. Tang,J. Comput. Chem., 1982, 3, 317–328; (e) R. F. W. Bader, Acc.

Chem. Res., 1985, 18, 9–15; (f) T. H. Tang, R. F. W. Bader andP. MacDougall, Inorg. Chem., 1985, 24, 2047–2053; (g) F. Biegler-Konig, J. Schonbohm and D. Bayles, J. Comput. Chem., 2001, 22,545–559; (h) F. Biegler-Konig and J. Schonbohm, J. Comput.Chem., 2002, 23, 1489–1494. See also: W. Nakanishi,T. Nakamoto, S. Hayashi, T. Sasamori and N. Tokitoh, Chem.Eur. J., 2007, 13, 255–268.

23 J. Meinwald, D. Dauplaise and J. Clardy, J. Am. Chem. Soc., 1977,99, 7743–7744.

24 The structures of 3b (LO) and 3a (OO) are also determined by theX-ray crystallographic analysis. The results are essentially the sameas those of 1 (LO) and 1 (OO), respectively, which will be reportedelsewhere.

25 Water molecules in 1 (OO) are omitted for clarity.26 S. Hayashi and W. Nakanishi, Bull. Chem. Soc. Jpn., 2008, 12, in

press (CCDC 640537 for 1 (LL)).27 A. Bondi, J. Phys. Chem., 1964, 68, 441–451.28 M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria,

M. A. Robb, J. R. Cheeseman, V. G. Zakrzewski, J. A.Montgomery, Jr., R. E. Stratmann, J. C. Burant, S. Dapprich,J. M. Millam, A. D. Daniels, K. N. Kudin, M. C. Strain,O. Farkas, J. Tomasi, V. Barone, M. Cossi, R. Cammi,B. Mennucci, C. Pomelli, C. Adamo, S. Clifford, J. Ochterski,G. A. Petersson, P. Y. Ayala, Q. Cui, K. Morokuma, P. Salvador,J. J. Dannenberg, D. K. Malick, A. D. Rabuck, K. Raghavachari,J. B. Foresman, J. Cioslowski, J. V. Ortiz, A. G. Baboul,B. B. Stefanov, G. Liu, A. Liashenko, P. Piskorz, I. Komaromi,R. Gomperts, R. L. Martin, D. J. Fox, T. Keith, M. A. Al-Laham,C. Y. Peng, A. Nanayakkara, M. Challacombe, P. M. W. Gill,B. G. Johnson, W. Chen, M. W. Wong, J. L. Andres, C. Gonzalez,M. Head-Gordon, E. S. Replogle and J. A. Pople, GAUSSIAN 98(Revision A.11), Gaussian, Inc., Pittsburgh, PA, 2001.

29 C. Lee, W. Yang and R. G. Parr, Phys. Rev. B, 1988, 37, 785–789;B. Miehlich, A. Savin, H. Stoll and H. Preuss, Chem. Phys. Lett.,1989, 157, 200–2006.

30 A. D. Becke, Phys. Rev. A, 1988, 38, 3098–3100; A. D. Becke,J. Chem. Phys., 1993, 98, 5648–5652.

31 S. Hayashi and W. Nakanishi, J. Mol. Struct. (THEOCHEM),2007, 811, 293–301.

32 M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria,M. A. Robb, J. R. Cheeseman, J. A. Montgomery, Jr, T. Vreven,K. N. Kudin, J. C. Burant, J. M. Millam, S. S. Iyengar, J. Tomasi,V. Barone, B. Mennucci, M. Cossi, G. Scalmani, N. Rega,G. A. Petersson, H. Nakatsuji, M. Hada, M. Ehara, K. Toyota,R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda,O. Kitao, H. Nakai, M. Klene, X. Li, J. E. Knox,H. P. Hratchian, J. B. Cross, V. Bakken, C. Adamo, J. Jaramillo,R. Gomperts, R. E. Stratmann, O. Yazyev, A. J. Austin, R. Cammi,C. Pomelli, J. W. Ochterski, P. Y. Ayala, K. Morokuma,G. A. Voth, P. Salvador, J. J. Dannenberg, V. G. Zakrzewski,S. Dapprich, A. D. Daniels, M. C. Strain, O. Farkas, D. K. Malick,A. D. Rabuck, K. Raghavachari, J. B. Foresman, J. V. Ortiz,Q. Cui, A. G. Baboul, S. Clifford, J. Cioslowski, B. B. Stefanov,G. Liu, A. Liashenko, P. Piskorz, I. Komaromi, R. L. Martin,D. J. Fox, T. Keith, M. A. Al-Laham, C. Y. Peng, A. Nanayakkara,M. Challacombe, P. M. W. Gill, B. Johnson, W. Chen,M. W. Wong, C. Gonzalez and J. A. Pople, GAUSSIAN 03(Revision B.05), Gaussian, Inc., Pittsburgh, PA, 2003.

33 The AIM2000 program (Version 2.0) is employed to analyze andvisualize atoms in molecules: F. J. Biegler-Konig, Comput. Chem.,2000, 21, 1040–1048; see also ref. 22g.

34 Data for A and B, together with LL, are given in the ESIw.35 The type C of 1 (OH�OH) is discussed which is predicted to be most

stable among the three44.36 Eqn (R1) shows the energies of n (L) + H2O2 (E(n (L) + H2O2))

relative to E(n (O) + H2O) [DE(n (LO)) = E(n (L) + H2O2) �E(n (O) + H2O)], although E(n (L)) are not given in Table 3.45

DE(n (LO)) = E(n (L) + H2O2) � E(n (O) + H2O)G = H (121.8 kJ mol�1) o F (131.3) o cis-MeSe (133.9)o Cl (136.8) r Br (137.6) o trans-MeSe (141.0) (R1)

37 NOB analysis were also performed on the AC conformer of 8-G-1-[MeSe(O2H2)]C10H6. However, the corresponding CT interactionswere not detected.


38 The nonbonded Se� � �Se distance in 1 (LO) is predicted to beshorter than that of 1 (OO) by ca. 0.03 A, while the observedvalues are almost equal (see Table 5). The crystal packing effectmight contribute to the results.

39 The value is very close to that evaluated with the B3LYP/6-311+G(d) method.

40 W. Nakanishi, S. Hayashi and K. Narahara, unpublished results.41 G. M. Sheldrick, SHELXS-97, Program for Crystal Structure

Solution, Universitat Gottingen, Germany, 1997.42 G. M. Sheldrick, SHELXL-97, Program for Crystal Structure

Refinement, University of Gottingen, Germany, 1997.43 MolStudio R3.2 (Rev 1.0), NEC Corporation, 1997–2003.

44 Three structures (type A, type B and type C) were optimizedfor each of n (OH�OH). The type C is the global minimum,which is slightly stable than type B and much stable thantype A, although the steric repulsion between OH and G seemslargest.

45 Eqn (R1)36 shows that selenoxides are stabilized in this orderthrough the non-bonded n(G)� � �s*(Se–O) 3c–4e interactions,together with the O dependence.16 While G = trans-MeSe isdemonstrated to be most effective to stabilize in the selenoxiderelative to the corresponding bis-selenide, the effect of G = cis-MeSe places between F and Cl, where the CC form is postulatedfor the bis-selenide.


Mechanistic aspects of nitrate ion reduction on silver electrode:

estimation of O–NO2�bond dissociation energy using

cyclic voltammetry

Mohsin Ahmad Bhat,w Pravin Popinand Ingole, Vijay Raman Chaudhari

and Santosh Krishna Haram*

Received (in Montpellier, France) 27th August 2008, Accepted 9th October 2008


DOI: 10.1039/b814895c

Voltammetric investigations of mechanistic aspects and estimation of thermo-chemical parameters

of nitrate ion reduction at silver electrode, in alkaline medium are reported. The activation barrier

determined from cyclic voltammetry fits a quadratic relation rather than the expected Butler-

Volmer kinetics. Intrinsic barrier calculations show that the reduction of nitrate ion on silver

follows a concerted mechanism, involving electron transfer initiated bond cleavage, followed by

chemical reaction. The bond dissociation energy for the O–NO2� bond was estimated to be

48.40 kcal mol�1, which matches well with the reported value of 47.5 kcal mol�1, determined

from photodissociation experiments.

1. Introduction

The high water solubility of nitrate ions is responsible for its

virtual presence everywhere. Serious clinical symptoms have

been reported to be caused by their ingestion,1–3 which neces-

sitates effective monitoring and development of sensing tools

for this ion. For its detection and estimation, a series of

methods have been proposed in refs. 4–8. Among them, the

electrochemical methods have proved to be advantageous in

terms of reproducibility, accuracy, time response and durability.9

The voltammetric detection is based on the irreversible two-

electron transfer process shown in eqn (1):10

O–NO2� + 2e + H2O - NO2

� + 2OH� (1)

Using conventional voltammetry,11,12 in combination with

electrochemical scanning tunneling microscopy, surface

enhanced Raman spectroscopy,13 and differential electro-

chemical mass spectrometry,14 it is reported that the nitrate

ion reduction is very sensitive to the solution conditions, pH

and the nature of the electrode material. For example, on

polycrystalline platinum electrode, it proceeds through a dis-

sociative adsorption pathway,15 albeit with slow kinetics.

Similar results have been also reported for palladium electro-

des,16 while on Cu, the reduction is found to be very facile

and proceeds through the formation of NO2� in alkaline

medium.13 Dima et al.14 have reported varying activity of

transition and coinage metals and proposed a probable general

scheme for the nitrate ion reduction process. In general, nitrate

reduction is understood to be a multistep process with the first

electron transfer as the rate determining step.12–14 Among

all the metals studied, silver shows highest sensitivity for

nitrate ion reduction, and is thus strongly advocated for

their electrochemical sensing.10,17,18 For better understanding

and development of Ag as a nitrate ion sensor, it is of utmost

importance to have knowledge about the mechanism of this

reduction—especially with respect to the rate determining

step. To our knowledge, these aspects have not been con-

sidered so far. Thus, it was of our immense interest to study the

mechanism of this reaction on Ag-electrode, voltammetrically.

With this aim, the kinetic investigations of nitrate ion

reduction in alkaline media were undertaken through cyclic

voltammetry. Our analysis for the first time revealed that the

nitrate ion reduction on Ag follows a dissociative electron

transfer mechanism. Besides, the related calculations led

to the quantification of thermo-chemical parameters, such

as bond dissociation energy of the O–NO2� bond, which is

otherwise estimated through thermal- and photo-dissociation

measurements.19

2. Experimental

Potassium nitrate and sodium hydroxide were purchased from

Merck. Ag-bulk electrode was prepared by sealing a 2 mm

diameter Ag (99.9%) wire in a glass tube, with the help of

epoxy adhesive. The electrode surface was exposed by grinding

it on emery paper. It was polished with commercially available

silver cleaner, followed by 0.2 mm alumina powder. Ag/AgCl,

KCl (3.0 M), and a Pt rod from Metrohm devices were used as

reference and counter electrodes, respectively. Cyclic voltam-

metric (CV) investigations were performed using Metrohm

PGSTAT 100 POTENTIOSTAT/GALVANOSTAT in a

three-electrode setup. All the measurements were performed

under an argon atmosphere. Prior to measurements, the Ag

electrode was electrochemically activated by cycling the

potential twenty times (scan rate 1 V s�1) in the potential

range �1.3 to 1.0 V, followed by ten potential steps

(of 1 s duration) in increasing order of potential, ranging from

Department of Chemistry, University of Pune, Ganeshkhind, Pune,411007, India. E-mail: [email protected];Fax: +91 20 2569 3981; Tel: +91 20 2560 1394w Permanent address: Department of Chemistry, University of Kashmir,Srinagar-190006, India.



�0.25 to 0.9 V in 2.0 M NaOH. CVs used for analysis were

background corrected for the charging current. In view of the

reported complications observed during nitrate ion reduction

at Ag-electrode,17 every scan used for the analysis was re-

corded on a freshly polished and electrochemically cleaned

electrode surface. All the measurements were carried out in

thermostatted condition at 298 � 0.1 K.


A typical CV recorded using the Ag electrode in 10 mMKNO3

and 2 M NaOH (pH = 12) at the scan rate of 10 mV s�1 is

shown in Fig. 1. High pH helps to shift hydrogen evolution

towards more negative potential in comparison to the nitrate

ion reduction. A cathodic peak at �0.94 V is assigned to the

reduction of nitrate ions. The linear relationship between ipand the square root of scan rate (n1/2) (Fig. 1, inset) indicatesthat the process is diffusion controlled.20 The scan rate depen-

dent shift in peak potential (Ep), as shown in Fig. 2, is

attributed to the irreversibility in the electron transfer process.

The magnitude of Ep � Ep/2 was in the range 47 to 96 mV,

which is an indication of electron transfer being the rate

determining step.21 Interestingly, ip/n1/2 (Fig. 3) and Ep �

Ep/2 (Table 1) were found to be scan rate dependent. Addi-

tionally, peak broadening with increase in scan rate is ob-

served (Fig. 2). Prima facie, both these observations could

be attributed to the uncompensated iR drop and charging

current contributions. However, the electrolyte used was

highly conducting and moreover, we had subtracted back-

ground charging current. Therefore, the contributions of iR

drop and the capacitance are ruled out. Thus, the shift in

the potential and Ep � Ep/2 with the scan rate are attributed

to the potential dependent electron transfer coefficient (a)for the reaction22 and nitrate ion reduction does not follow

a normal Butler–Volmer kinetics. Since the ip vs. n1/2 plot

shows that the process is diffusion controlled, it also suggests

that the potential dependent free energy of activation for the

reaction is a quadratic function of electrode potential, as per

eqn (2):23

DGz ¼ DG zo 1þ DGo

4DG zo

� �2

ð2Þ

Fig. 1 Typical cyclic voltammogram recorded on Ag electrode in

10 mM KNO3 and 2 M NaOH, at a scan rate of 10 mV s�1. Inset

shows a linear fit of the peak current (ip) vs. square root of scan rate

(n1/2), which indicates diffusion controlled reaction.

Fig. 2 Cyclic voltammograms recorded on Ag electrode in 10 mM

KNO3 and 2 M NaOH, at varying scan rates from 10 to 500 mV s�1.

Inset shows peak potential vs. scan rate, which indicates that the

reduction is irreversible.

Fig. 3 ip/n1/2 vs. scan rate for NO3 reduction on Ag electrode in

alkaline medium, indicating a deviation from Butler-Volmer kinetics.

Table 1 Analysis of cyclic voltammetric (CV) data obtained fornitrate ion reduction (10 mM) at an Ag electrode in 2 M NaOH at298 K

Scan rate/V s�1 104Ip/A Ep/V (Ep � Ep/2)/VDGo

z/kcal mol�1

0.01 0.891 �0.944 0.060 19.970.02 1.40 �0.920 0.065 19.560.04 2.01 �0.926 0.072 19.660.08 2.70 �0.932 0.078 19.770.1 2.80 �0.934 0.083 19.800.2 3.74 �0.944 0.085 19.970.3 4.27 �0.948 0.087 20.040.4 4.77 �0.958 0.089 20.220.5 5.03 �0.962 0.094 20.28


where,

DGo = F(E � Eo0) (3)

in which Eo0 is the formal potential for overall reaction and

has a value of 0.01 V24 and DGoz, is the intrinsic energy barrier

for the reduction reaction.

The potential dependent rate constant (k(E)) is given by

eqn (4):

kðEÞ ¼ Z: exp�DGzRT

� �

ð4Þ

where, Z (= (RT/2pM)1/2,M is the molecular mass of NO3�) =

2521.79 cm s�1.

Such dissociative redox reactions initiated with electron

transfer can occur through two plausible mechanisms, namely

stepwise (eqn (5) and (6)) and concerted (eqn (7)):

A–B + e� - A–B�� (5)

A–B�� - A�–B� (6)

A–B + e� - A� + B� (7)

Products formed through both these mechanisms can un-

dergo further electron transfer or chemical reactions, which

affect the thermodynamic and kinetic aspects of the overall

process. Theory as well as experimental predictions associated

with these mechanisms for alkyl halides,25 peroxides22,26 and

other analytes27,28 have been well documented in the literature.

The two mechanisms can be differentiated on the basis of

difference in the value of intrinsic energy barrier, as given

below (eqn (8) and (9)):

DG zo ðstepwiseÞ ¼l0 þ li

4ð8Þ

DG zo ðconcertedÞ ¼l0 þ li þ BDFE

4ð9Þ

where, BDFE is the bond dissociation free energy, and l0, thesolvent reorganization energy, which is calculated through the

Marcus equation (10):

l0 ¼e

8peoao

1

eo� 1

es

� �

ð10Þ

Here, eo and es are the optical and static dielectric constants of

the solvent, respectively and ao, the effective radius of the

analyte, which is calculated using eqn (11):

ao ¼ aB2aAB � aB

aAB

� �

ð11Þ

where aAB = aNO3� = 2.64 A and aB = aO� = 1.76 A are as

reported previously.29 li is regarded as the internal reorgani-

zation energy and can be neglected due to its comparatively

small magnitude. l0 was calculated using reported values for

optical and static dielectric constants for water.30 Use of

convolutive analysis of CVs has been reported for the calcula-

tion of intrinsic barriers and other mechanistic details of such

electron transfer reactions. Though, the method has many

advantages, its use is limited due to the requirement of

information regarding the double layer structure. Except for

mercury, such information is not available for other metal

electrodes. We used another simple approach for calculation

of the barrier using cyclic voltammetry, without any convolu-

tion analysis, which is as follows.

The free energy of activation at peak potentials obtained in

cyclic voltammograms is given by eqn (12):31

DG zP ¼RT

Fln Z

ffiffiffiffiffiffiffiffiffiffiffiffiRT

FanD

r !

� 0:78

" #

ð12Þ

Knowing the value of a calculated from Ep � Ep/2, where Ep

is the peak potential and Ep/2 is the potential where the current

is at half the peak value (Table 1), through eqn (13):20

a ¼ 1:86RT

FjEp � Ep=2j

� �

ð13Þ

and D (1.9 � 10�5 cm2 s�1), DGPz values for the various scan

rates were calculated. These values were substituted in eqn (2)

and the resulting quadratic equation was solved for DGoz with

the help of the FORTRAN program, written specifically for

this purpose. Among the two roots obtained, the negative root

was not considered as the value emerged out to be less than

that obtained from eqn (8), which is the bare minimum value

expected for the overall reaction. The positive root, gives a

value of DGoz much greater than that expected for a stepwise

mechanism (eqn (5) and (6)) and hence negating the possibility

that the reduction follows a stepwise mechanism. Therefore,

the reaction occurring through the concerted mechanism as

shown in eqn (7) is inferred. Preference of a concerted

mechanism over a stepwise one is also realized by considering

the resonance structure of nitrate ion and the nitrite ion as

against the open shell structure of NO3�2�—an intermediate

which would be formed in a stepwise mechanism.

The bond dissociation energy was calculated by substituting

the value of DGoz (from the above procedure) and lo

(from eqn (10)) in eqn (9) and found to be ca. 48.4 kcal mol�1,

which matches well with the value of 47.5 kcal mol�1, reported

from photodissociation measurements.19 The small difference

is attributed to the neglecting the value of li in the calculations

as a first approximation.

Based on our experimental findings and earlier re-

ports,13,14,32,33 we conclude that, electron transfer to the

nitrate ion is the rate determining step, similar to the process

reported for Cu, and following the overall reaction scheme

given by eqn (14) and (15):

NO3� + e� - NO2

� + O�� (rate determining step) (14)

O�� + e� + H2O - 2OH� (15)

Results published recently by Broder et al.34 also suggest the

formation of O�� as an intermediate in reduction of nitrate ion

at Pt electrode in room-temperature ionic liquids.

4. Conclusion

For the first time, we have used a simple procedure for

analyzing the cyclic voltammetric data for nitrate ion reduc-

tion at a silver electrode. Data fits very well in the dissociative

electron transfer concerted mechanism. Besides, the value of


the bond dissociation energy, ca. 48.4 kcal mol�1, calculated

from these investigations, matches with 47.5 kcal mol�1 the

value obtained from photodissociation measurements. This

knowledge of mechanistic and thermo-chemical parameters is

believed to be useful in designing Ag electrodes as nitrate ion

sensors.

Acknowledgements

M. A. B. would like to thank the University authorities,

especially Vice Chancellor, University of Kashmir, and Head,

Department of Chemistry, University of Kashmir, for sanction

of study leave. P. P. I. thanks CSIR—India, for a fellowship.

V. R. C. thanks the BARC–Pune University collaborative PhD

program for financial support. The authors would like to thank

CNQS, University of Pune, for financial support.

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ISSN 1144-0546










Volume

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aterials Chem

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ISSN 1757-9694

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ISSN 1144-0546

Volume 33 | N

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