Etude à pH physiologique, des mécanismes de ... · (Hemanth) for his continuous encouragement...

1

Université de REIMS CHAMPAGNE-ARDENNE

ICMR, Groupe Chimie de Coordination

CNRS UMR 7312

Sciences Technologie Santé

THÈSE

Pour obtenir le grade de

Docteur de l’Université de Reims Champagne-Ardenne

Discipline : CHIMIE

par

Vijetha MOGILIREDDY

Soutenue publiquement le 16 décembre 2013

Etude à pH physiologique, des mécanismes de transmétallation de complexes linéaires et macrocycliques de gadolinium utilisés en IRM

Jury

M. Stephen James ARCHIBALD Rapporteur Mme Luce VANDER ELST Rapporteur M. Stéphane ROUX Examinateur M. Laurent DUPONT Président Mme Françoise CHUBURU Directeur de thèse Mme Isabelle DECHAMPS Directeur de thèse

© N° attribué par la bibliothèque | | | R | E | I | | | | | |

3

Résumé

L’objectif de ce travail est l’analyse de la stabilité thermodynamique et de l’inertie

chimique de complexes métalliques avec des ligands ou des nanoparticules conçus pour des

applications en IRM. Deux types de ligands polyaminocarboxylates ont été étudiés, ligands

pour lesquels les unités complexantes sont soit linéaires soit macrocycliques.

Les ligands macrocycliques étudiés sont des ligands basés sur des squelettes DO3A,

substitués par des entités benzimidazole (L1H4) ou p-nitrophenylbenzimidazole (L2H3). Les

données thermodynamiques indiquent que les affinités de ces ligands vis-à-vis des ions de la

première série de transition (Cu(II) et Zn(II)) ou vis-à-vis des lanthanides (Gd(III) et Eu(III))

sont plus élevées que celles des complexes correspondants avec le ligand DO3A. Ce

renforcement d’affinité est corrélé avec la participation des groupements benzimidazole à la

sphère de coordination de chacun des métaux. L’inertie chimique du complexe Gd(III)- L1H4

a ensuite été évaluée par relaxométrie en tampon phosphate, en présence d’une quantité

équimolaire de Zn(II). Dans cette expérience, le Zn(II) joue le rôle d’un compétiteur du

Gd(III) c’est-à-dire qu’il peut si le complexe Gd(III)-L1H4 n’est pas inerte chimiquement,

induire une libération de l’ion gadolinium. Pour Gd(III)-L1H4, aucune réaction de ce type n’a

été détectée, ce qui plaide en faveur de l’inertie chimique de ce complexe.

Les ligands linéaires étudiés sont des dérivés dithiolés de ligands DTPA bisamide

L@1H5. Ces ligands ont été conçus pour être greffés sur des nanoparticules d’or. La stabilité

thermodynamique des complexes de Cu(II), Zn(II) et Gd(III) utilisant les ligands L@1H5 et

L@1H5 greffé sur nanoparticule d’or (autrement appelé L@

2H3) suit l’ordre de stabilité

croissant Zn(II) < Cu(II) < Gd(III). Par ailleurs, les résultats montrent que le complexe

Gd(III)-L@1H5 est moins stable d’au moins deux ordres de grandeur que le complexe

Gd(III)-L@2H3. Ceci suggère qu’une fois greffé sur la nanoparticule, le complexe de

gadolinium correspondant gagne en stabilité. Par ailleurs, des études comparatives d’inertie

chimique montre que le complexe Gd(III)-L@1H5 greffé sur la nanoparticule a une inertie

chimique comparable à celle de l’agent de contraste commercial Gd-DTPA. En revanche

lorsque ce complexe est seul, sa vitesse de démétallation est rapide. Le greffage du ligand

L@1H5 à la surface de la nanoparticule est donc au bénéfice de la stabilité et de l’inertie

chimique de son complexe de Gd(III). Ce gain de stabilité peut être attribué à l’’effet de

ballast’ de la nanoparticule qui rigidifie la structure du complexe et limite sa démétallation.

5

Université de REIMS CHAMPAGNE-ARDENNE

ICMR, Groupe Chimie de Coordination

CNRS UMR 7312

Sciences Technologie Santé

THÈSE

Pour obtenir le grade de

Docteur de l’Université de Reims Champagne-Ardenne

Discipline : CHIMIE

par

Vijetha MOGILIREDDY

Soutenue publiquement le 16 décembre 2013

Study of transmetallation mechanisms of macrocyclic and linear gadolinium complexes at physiological pH for MRI

Jury M. Stephen James ARCHIBALD Rapporteur Mme Luce VANDER ELST Rapporteur M. Stéphane ROUX Examinateur M. Laurent DUPONT Président Mme Françoise CHUBURU Directeur de thèse Mme Isabelle DECHAMPS Directeur de thèse

© N° attribué par la bibliothèque | | | R | E | I | | | | | |

7

Abstract

The aim of this work is to analyse the stability of metal complexes with ligands or

nanoparticles of interest in MRI and to study their transmetallation mechanisms in the

presence of endogenous cations near physiological pH. Two types of polyaminocarboxylate

ligands were studied for which the binding unit was either linear or macrocyclic.

Macrocyclic ligands are constituted of a DO3A backbone functionalized with a

benzimidazole (L1H4) or a p-nitrophenylbenzimidazole unit (L2H3). Thermodynamic data

indicated that the affinities of these ligands towards first row transition metal ions (Cu(II) and

Zn(II) or lanthanide ions (Gd(III) and Eu(III)) are increased compared to the corresponding

ones with DO3A. This enhancement is correlated to the involvement of the benzimidazole

moiety to each metal coordination sphere. For gadolinium complex Gd(III)-L1H4, its kinetic

inertness was evaluated in phosphate buffer by relaxometry, in the presence of equimolar

quantities of Zn(II) as a competitor. In these conditions, if the complex is not chemically

inert, it would be subjected to a transmetallation reaction, that is to say that at least,

gadolinium would be released. For Gd(III)-L1H4, no such reaction was detected which is in

favour of kinetic inertness of Gd(III)-L1H4.

Linear ligand, dithiolated DTPA bisamide L@1H5 was designed with an aim of

grafting it onto gold nanoparticles. L@1H5 and the ligand grafted into gold nanoparticle,

namely L@2H3, were analysed for their thermodynamic stability towards mainly Cu(II),

Zn(II) and Gd(III). Whatever the system, L@1H5 or L@

2H3, the general trend of increasing

complex stability was Zn(II) < Cu(II) < Gd(III). Furthermore, Gd(III)-L@1H5 complex was

less stable than Gd(III)-L@2H3, this latter being 2 orders of magnitude more stable at

physiological pH. This suggested that the gadolinium complex stability is enhanced when the

ligand is grafted onto the nanoparticle. Moreover, comparative kinetic inertness studies

showed that the gadolinium complex Gd(III)-L@1H5 is not chemically inert and demetallates

rapidly while the gadolinium complex grafted onto the nanoparticle exhibit almost equal

kinetic inertness as Gd-DTPA (Magnevist). The bulky nanoparticle probably rigidifies the

structure of the complex and prevents Gd(III)-L@2H3 from an extensive demetallation, which

was a good point for the possible use of these nanoparticles in living organisms for imaging

applications.

9

Mots – clés: Ligands DO3A, methyl benzimidazole, ligands DTPA bisamide,

nanoparticules d’or, Stabilité thermodynamique, inertie chimique, transmétallation,

relaxométrie, spectroscopie UV-Visible.

Keywords: DO3A ligands, methyl benzimidazole, DTPA bisamide ligands, gold

nanoparticles, thermodynamic stability, kinetic inertness, transmetallation, relaxometry, UV-

Visible spectroscopy.

Adresse du laboratoire et de l’unité:

Université de Reims Champagne Ardenne

Institut de Chimie Moléculaire de Reims – UMR CNRS 7312

Groupe Chimie de Coordination

UFR Sciences Exactes et Naturelles – Moulin de la Housse - Batiment 18 – BP 1039

51687 Reims Cedex 2

11

To my whole Family

13

Acknowledgements

After three years, and a physico-chemical study of macrocyclic and linear ligands, hideous

and jittery equations for transmetallation experiments, I want to express here my grateful

reconnaissance to important persons who helped me doing this work.

Firstly, I would like to thank Dr. Stephen James ARCHIBALD, Dr. Luce VANDER ELST,

Dr. Stéphane ROUX and Dr. Laurent DUPONT for accepting to be the jury members of my

thesis. I would like to express my gratitude to my supervisors Dr. Françoise CHUBURU and

Dr. Isabelle DECHAMPS for their continuous and affectionate support in lab work as well as

during writing up this thesis. I would like to strongly thank our collaborators Dr. S.J.

ARCHIBALD and Dr. Stephane ROUX for their express fund of ligands when ever asked

for. I could never forget the surprise one day journey to Mons, Belgium for the learning of

relaxometry handling; Dr. Luce VANDER ELST and Dr. Sophie LAURENT are greatly

acknowledged for their valuable advices. For technical support, Dr. Dominic HARAKAT,

Sylvie LANTHONY, Christophe PETERMANN, Agate MARTINEZ and Antony ROBERT

are also thanked for the mass, elemental analysis and NMR experiments. I would like to

thank Dr. Cyril CADIOU for geometrical structures in the thesis and his help in the

laboratory very often. Our group members (Juliette, Stephanie, Laurent, Aminou, Emanuel,

and José) are also thanked for their warm welcome into the group, which made me feel

comfortable in the lab.

I could not imagine my stay in France without my wonderful friends i have acquired. A very

big thanks to Gaelle, Guillaume, Arnaud, Loic, Sai sai, Kun, Ailing, Khoa, Jomy, Audrey,

Hela, Mareen, Anais, Berengere, Axelle, Krupakar, Surendra, Bapuji for their beautiful

hearts. I would like to thank Sylviane for her support during my beginning days in France.

I would like to thank once again my co-supervisor Dr. Isabelle DECHAMPS, who has been

there always for me at all times, bearing me with patience, and ofcourse with my terrible

french. It has been a wonderful experience to work together in the lab.

At the end my special thanks will go for my whole family and especially to my husband

(Hemanth) for his continuous encouragement throughout my stay in France. French Ministry

for Education and the Region Champagne-Ardenne are thanked for their financial support of

my thesis.

15

Résumé substantiel des chapitres

de la thèse

17

Introduction générale

L’utilisation d’agents de contraste à base de gadolinium pour l’imagerie IRM justifie

les nombreux efforts effectués pour augmenter leur innocuité et leur efficacité. Dans ce

contexte, l’utilisation de ligands polyaminocarboxylates pour la complexation de lanthanides

en général, et du gadolinium en particulier, repose sur l’affinité remarquable de ces ligands

pour ces ions. Cette affinité se décline d’une part par une forte stabilité thermodynamique

vis-à-vis des lanthanides et d’autre part par une bonne inertie chimique vis-à-vis de la

démétallation pour les complexes correspondants. L’intérêt de ces ligands réside aussi dans la

possibilité de les fonctionnaliser sélectivement. Par exemple, il est possible de rendre ces

ligands (et donc ces complexes) intelligents en leur permettant via une fonctionnalisation

adaptée, d’atteindre spécifiquement certains types cellulaires. Il est aussi possible par

ingénierie moléculaire du ligand, de coupler plusieurs modalités d’imagerie. En matière de

multimodalité, l’association de ligands polyaminocarboxylates avec des nanoparticules

métalliques telles que les nanoparticules d’or est un bon exemple. Il n’en reste pas moins

qu’avant toute application in vivo, la stabilité thermodynamique et l’inertie chimique de ces

nouveaux systèmes doit être testée.

Dans le contexte de cette thèse quatre composés ont été étudiés, deux d’entre eux

étant basés sur des ligands macrocycliques dérivés du DO3A (L1H4 and L2H3). Les

groupements fonctionnels additionnels sont des dérivés de benzimidazole. Les deux autres

systèmes sont basés sur des ligands linéaires dérivés du DTPA bisamide (L@1H5 et L@

2H3)

pour lesquels les fonctions amide terminales portent des fonctions thiol. Le ligand de base

L@1H5 greffé sur nanoparticule d’or conduit au système L@

2H3. Afin de statuer sur l’affinité

de ces quatre systèmes vis-à-vis des lanthanides et en particulier du gadolinium, notre

démarche a consisté à comparer pour chaque système son affinité pour Gd(III) et pour des

ions potentiellement compétiteurs en milieu biologique.

Le premier chapitre de cette thèse décrira les propriétés physiques des lanthanides qui

peuvent être pertinentes en imagerie, particulièrement en imagerie IRM. Ce chapitre montrera

aussi la nécessité de fortement chélater ces ions pour que les complexes correspondants

puissent être candidats à une mise sur le marché comme agents de contraste IRM.

18

Le second chapitre s’intéressera à la mise en évidence de la stabilité

thermodynamique de complexes issus des ligands L1H4 and L2H3. Ces ligands

macrocycliques ont été synthétisés par l’équipe du Pr S. J. Archibald (Université de Hull).

L’objectif sera d’évaluer par pH-métrie, la stabilité thermodynamique des complexes de

Cu(II), Zn(II), Gd(III) et Eu(III) obtenus à partir des ligands précités. Des études en

spectroscopie UV, 1H NMR, RPE et fluorescence viendront conforter si nécessaire les

hypothèses structurales émises. L’inertie chimique du complexe de gadolinium Gd(III)-L1H4

sera ensuite été évaluée par relaxométrie et comparée à celle d’un agent de contraste

commercial (DOTAREM®). Nous montrerons qu’aucune démétallation n’est détectée pour

ce complexe Gd(III)-L1H4, ce qui est un bon premier point dans l’optique de son utilisation

comme agent de contraste IRM.

Le troisième chapitre de la thèse s’intéressera à la mise en évidence de la stabilité

thermodynamique de complexes issus du ligand L@1H5 et des nanoparticules d’or associées

(L@2H3). Ces systèmes ont été synthétisés par l’équipe du Pr S. Roux (Université de

Besançon). L’approche développée dans le Chapitre II sera transposée dans ce chapitre.

Après l’évaluation de la stabilité thermodynamique des complexes de gadolinium Gd(III)-

L@1H5 et Gd(III)-L@

2H3, leur inertie chimique sera evaluée par relaxométrie et une

proposition de mécanisme de transmétallation sera faite. Nous montrerons en particulier que

le greffage du ligand sur la nanoparticule se traduit pour le complexe correspondant par un

gain de stabilité thermodynamique et un renforcement conséquent de son inertie chimique.

Ceci est aussi un bon point pour l’utilisation de ces nanoparticules d’or greffées par des

complexes de gadolinium, dans le cadre d’applications en imagerie multimodale.

19

A. Chapitre - I

Ce chapitre est dédié à des rappels bibliographiques autour de la chimie des

lanthanides et de leurs applications biomédicales.

La première section de ce chapitre rappelle les propriétés électroniques fondamentales

des éléments lanthanides, en particulier leurs diagrammes d’énergie électronique qui

serviront de base à l’interprétation de leurs spectres d’absorption et d’émission. L’accent sera

d’ailleurs mis sur les mécanismes d’émission des lanthanides et en particulier sur

l’amplification de ce phénomène par l’effet d’antenne que peuvent exercer des ligands

absorbeurs de lumière. Une mention sera aussi faite quant aux propriétés magnétiques de ces

ions.

La seconde section de ce chapitre s’attache à décrire les applications envisageables

pour les ions lanthanides et ce, du fait des propriétés électroniques précédemment rappelées.

La troisième section de ce chapitre s’attache à décrire plus longuement les

applications biomédicales des lanthanides, que ce soit des applications comme sondes

luminescentes ou comme sondes magnétiques.

Les deux dernières sections de ce chapitre bibliographique s’attachent à montrer

l’intérêt de complexes de Gadolinium comme agents de contraste pour l’imagerie IRM ainsi

qu’à établir un cahier des charges pour que l’utilisation de ces complexes permettent de

s’assurer de leur innocuité pour des utilisations chez les patients. Seront à ce niveau précisé

l’importance de la stabilité thermodynamique de ces complexes, ainsi que celle de leur inertie

chimique vis-à-vis des réactions de transmétallation éventuelles en milieu biologique.

B. Chapitre – II

Dans ce chapitre, la complexation de deux ligands basés sur un squelette DO3A et

fonctionnalisés par des groupements méthyl-benzimidazole a été décrite. Ces ligands sont

nommés L1H4 et L2H3 et leur synthèse a été effectuée par le groupe du Pr. S. Archibald à

l’université de Hull (Schéma 1).

20

N N

N N

CO2HHO2C

HO2C

N

N

R

R = H, NO2

Schéma 1: Ligands benzimidazole DO3A étudiés dans ce chapitre

(R = H L1H4 , p-nitrophenyl group L2H3)

Les métaux utilisés pour la complexation sont des ions de la première série de

transition (Cu(II) et Zn(II) ainsi que des lanthanides tels que Gd(III) et Eu(III)).

Le premier but du travail était de décrire les différents schémas de coordination pour les

différents systèmes métal-ligand et de comparer les stabilités thermodynamiques respectives

de ceux-ci. Pour les systèmes basés sur le ligand L1H4, quel que soit le métal, le groupement

benzimidazole participe à la sphère de coordination du métal par son atome d’azote de type

imine, dès que la forme benzimidazolium est déprotonée. Dans le Schéma 2 sont présentés

les résultats pour les complexes de Cu(II) et Zn(II) du ligand L1H4.

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

[ML1H]- [ML1]2-

N N

N NCO2

-HO2C

-O2CHN

NH

M

[ML1H3]+

logK :Cu

Zn

4.3 4.5 9.2

4.2 5.1 9.7 Schéma 2: Schéma de Complexation pour Cu et Zn avec le ligand L1H4

Dans le Schéma 3 sont présentés les résultats pour les complexes de Gd(III) et Eu(III)

du ligand L1H4.

21

Gd

Eu

3.0

4.1

8.4

9.3

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]+

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

OH2

[ML1H] [ML1]-

OH2H2O OH2

Gd

Eu

3.0

4.1

8.4

9.3

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]+

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

OH2

[ML1H] [ML1]-

OH2H2O OH2

Schéma 3: Schéma de Complexation pour Gd et Eu avec le ligand L1H4

Pour les systèmes basés sur le ligand L2H3, la séquence de complexation est différente

puisque le groupement fonctionnel est coordiné même à faible pH. Le complexe se

déprotonant à partir de pH 6.4 est probablement un complexe dans lequel un atome d’azote

du macrocycle est protoné et donc non-coordiné (Schéma 4).

Gd

Eu

7.62

6.5

N N

NH NCO2

--O2C

-O2CN

N

R

M

[ML2H]+

N N

N NCO2

--O2C

-O2CN

N

R

M

[ML2]

H2OOH2H2O

Gd

Eu

7.62

6.5

N N

NH NCO2

--O2C

-O2CN

N

R

M

[ML2H]+

N N

N NCO2

--O2C

-O2CN

N

R

M

[ML2]

H2OOH2H2O

Schéma 4: Schéma de Complexation pour Gd et Eu avec le ligand L2H3

De plus, pour les complexes (Cu(II) et Zn(II)) issus du ligand L2H3, des espèces

hydroxylées sont proposées à des pH plus élevés (Schéma 5 et 6).

N N

NH NCO2

--O2C

-O2CN

N

R

Cu pKa = 6.6

[CuL2H]

N N

N NCO2

--O2C

-O2CN

N

R

CuN N

N NCO2

--O2C

-O2CN

N

R

Cu

OH

[CuL2]- [CuL2H-2]

pKa = 8.9

Schéma 5: Schéma de Complexation pour Cu avec le ligand L2H3

22

N HN

NH NCO2

--O2C

-O2CN

N

R

Zn pKa = 6.4

[ZnL2H2]+

N N

N NCO2

--O2C

-O2CN

N

R

Zn

N N

N NCO2

--O2C

-O2CN

N

R

Zn

OH

[ZnL2]-

[ZnL2H-2]

pKa = 10

N N

N NCO2

--O2C

-O2CN

N

R

Zn

HO OH

[ZnL2H-3]

pKa = 9.6

N N

NH NCO2

--O2C

-O2CN

N

R

Zn

ZnL2H

Schéma 6: Schéma de Complexation pour Zn avec le ligand L2H3

La comparaison des stabilités des différents complexes a été ensuite effectuée sur la

base du calcul du pourcentage de métal libre et ce, sur toute la gamme de pH

(log([M]libre/[M]total) = f(pH)) pour L1H4 et L2H3. Quel que soit le cation, les complexes basés

sur le ligand L1H4 sont les plus stables. Ces complexes sont toutefois moins stables que ceux

obtenus à partir du ligand DOTA. Ceci indique que l’atome d’azote de type imine du groupe

benzimidazole est un moins bon atome donneur qu’un groupement carboxylate du DOTA

(Schéma 7).

-12

-10

-8

-6

-4

-2

0

L5H3

L2H3

L4H4

L1H4

pH

log(

[Gd]

free/[

Gd]

tota

l)

2 4 6 8 10 12

Schéma 7: Comparaison des pouvoirs complexants de L1H4, L2H3, L4H4 et L5H3 vis-à-vis de

Gd(III) ; [L] = [Gd(III)] = 2×10-3 mol.L-1.

23

Le second objectif de ce travail était d’évaluer l’inertie chimique du complexe de

gadolinium Gd(III)-L1H4. Pour cela, des expériences de transmétallation de ce complexe en

présence de Zn(II) ont été menées en tampon phosphate (pH 7.0) et suivies par relaxométrie.

Si une réaction de transmétallation intervient entre le Gd(III) complexé et le Zn(II)

échangeur, un déclin de la relaxivité doit être obtenu. Pour Gd(III)-L1H4, sous ces conditions

opératoires, aucune chute de relaxivité n’a été mesurée (Schéma 8).

0 1000 2000 3000 4000 50000,0

0,2

0,4

0,6

0,8

1,0

1,2

R1 t /

R1 0

t (min)

Gd-L1H4

Gd-L4H4

Schéma 8: Transmétallation de Gd(III)-L1H4 suivie par relaxométrie.

Ce comportement est typique de celui de complexes macrocycliques pour lesquels, la

préorganisation et le nombre d’atomes donneurs du ligand sont parfaitement adaptées aux

contraintes stéréoélectroniques des lanthaanides. En conséquence, aucune réaction de

transmétallation n’intervient en présence de cations potentiellement échangeurs. Ce résultat

est encourageant dans l’optique de l’utilisation du complexe Gd(III)-L1H4 comme agent de

contraste pour des applications en IRM.

C. Chapitre - III

Dans ce chapitre, sont décrits la complexation du ligand DTDTPA (ligand L@1H5) qui

est un ligand de type DTPA bisamide, porteur de fonctions ethylène-thiol, et de ce ligand

greffé sur nanoparticule d’or (système L@2H3) (Schéma 9).

24

a.

N N N

COOHHOOC

O

NH

O

HNHS SH

COOH

b.

Schéma 9: a. DTDTPA et b. AuNPs greffées par GdDTDTPA

Ces composés ont été synthétisés dans le groupe du Pr S. Roux (Université de

Franche-Comté) et leur complexation a été envisagée avec les cations Cu(II), Zn(II), Ca(II),

Na(I) et Gd(III). Dans le ligand L@1H5, les fonctions thiol sont libres (SH) tandis que pour le

système L@2H3, les groupements thiols sont mobilisés pour le greffage du ligand sur la

nanoparticule via la formation de liaisons disulfure.

Quel que soit le système, L@1H5 ou L@

2H3, l’ordre croissant de stabilité des complexes

Ca(II) < Zn(II) < Cu(II) < Gd(III). Pour le ligand L@1H5, l’analyse de chaque diagramme de

spéciation indique que le complexe present à pH est du type [ML@1H2] pour lequel, les

fonctions thiols ne sont pas coordinées et non déprotonées. A pH plus élevé, ces fonctions se

déprotonent successivement.

Pour le système L@2H3 étant donné la densité de ligands L@

1H5 présente à la surface

de la nanoparticule et les probables interactions entre eux, il est peu réaliste de proposer une

structure pour les complexes associés. Néanmoins le diagramme d’existence des différentes

espèces présentes en solution et contenant du Gd(III) ne fait apparaitre qu’une stoechiométrie

de complexe, ce qui est une bonne chose pour l’utilisation de ces systèmes en milieu

biologique. Par ailleurs, la comparaison des stabilités des systèmes Gd(III)-L@1H5 et Gd(III)-

L@2H3 souligne que le système Gd(III)- L@

2H3 est plus stable de deux ordres de grandeurs

que son homologue non greffé sur nanoparticule Gd(III)-L@1H5 (Schéma 10).

25

-10

-8

-6

-4

-2

0pH

log

([Gd]

free/[

Gd]

tota

l)

L@1H5

L@2H3

L@4H5

2 4 6 8 10 12

Schéma 10: Comparaison des pouvoirs complexants de L@

1H5, L@2H3 et L@

4H5 vis-à-vis de

Gd(III) ; [L] = [Gd(III)] = 2×10-3 mol.L-1.

De plus, la comparaison des affinités de chaque système pour Cu(II), Zn(II) et Gd(II)

indique que les meilleures affinités sont obtenues pour Gd(III), ce qui est un point crucial

pour les applications de ces systèmes en imagerie IRM.

L’autre objectif du travail était d’évaluer l’inertie chimique des systèmes Gd(III)-

L@1H5 et Gd(III)- L@

2H3. Les expériences de transmétallation utilisées au chapitre précédent

ont été appliquées pour ces systèmes. Les résultats montrent que la réaction de

transmétallation entre le Gd(III) et le Zn(II) intervient beaucoup plus rapidement lorsque le

complexe n’est pas greffé sur la nanoparticule (Schéma 11).

0 1000 2000 3000 4000 50000.0

0.2

0.4

0.6

0.8

1.0

Gd(III)-L3@H5

Gd(III)-L2@H3

Gd(III)-L1@H5

R1/

R1 0

t (mn)

4

0 1000 2000 3000 4000 50000.0

0.2

0.4

0.6

0.8

1.0

Gd(III)-L3@H5

Gd(III)-L2@H3

Gd(III)-L1@H5

R1/

R1 0

t (mn)

4

Schéma 11: Transmétallation de Gd(III)-L@

1H5, Gd(III)-L@2H3 et Gd(III)-L@

4H5 vis-à-vis du

Zn(II) suivie par relaxométrie.

En d’autres termes, ces résultats indiquent que lorsque le complexe est greffé sur la

nanoparticule, la vitesse de la réaction de transmétallarion se voit réduite. On peut penser que

26

le greffage des complexes à la surface de la nanoparticule tend à les rigidifier ou du moins à

minimiser la capacité de pénétration des ions échangeurs au sein du réseau de complexes

greffés. De fait, leur inertie chimique s’en trouverait augmentée ce qui est un bon point pour

l’utilisation de ces nanoparticules greffées par des complexes de Gd en imagerie IRM. Enfin,

des expériences de transmétallation menées en excès de Zn(II) échangeur suggèrent que d’un

point de vue mécanistique, la force motrice de la réaction de transmétallation est la formation

de complexes hétéro-dinucléaires de type Gd(III)-L@1H5-Zn(II), protonés ou non.

D. Chapitre – IV

Dans ce chapitre sont décrits tous les protocoles expérimentaux utilisés au cours du travail.

La première partie de ce chapitre décrit la synthèse des ligands L1H4, L2H3, L@1H5 et des

nanoparticules L@2H3. Les synthèses ont été effectuées dans les équipes du Pr S. Archibald et

du Pr. S Roux.

L’analyse élémentaire des ligands nécessaire à toute étude préalable en potentiométrie est

reportée dans la deuxième partie de ce chapitre.

La troisième partie décrit la mise en œuvre des expériences de potentiométrie, ainsi que le

traitement des données qui en découle (logiciels PROTAF et HYPERQUAD). Les différents

modes de préparation des solutions sont aussi détaillés. Les constantes de déprotonation des

complexes étudiés dans ce travail sont rassemblées dans cette partie car complémentaires

mais non nécessaires à l’analyse des données effectuées dans le chapitre 3.

La quatrième partie décrit le protocole expérimental suivi en RMN 1H pour le suivi des

signaux 1H des différents ligands et complexes de Zn(II), en fonction du pH.

La cinquième partie décrit le protocole analytique suivi en spectroscopie RPE.

La sixième partie décrit le protocole expérimental suivi en relaxométrie (appareillage,

préparation des solutions) dans le cadre des études de transmétallation des complexes de

Gd(III) avec Zn(II). Cette partie détaille aussi le modèle chimique utilisé, permettant

d’accéder aux constantes de vitesse de démétallation des différents complexes de Gd(III).

27

La septième partie décrit le protocole expérimental suivi en spectroscopie UV

(appareillage, préparation des solutions), dans le cadre des études de transmétallation des

complexes de Gd(III) avec Zn(II).

Conclusion générale

Lorsque les complexes de coordination sont développés dans un but de diagnostic, la

question de leur stabilité thermodynamique et de leur inertie chimique vis-à-vis de réactions

de transmétallation est une question majeure, qui conditionne le développement et

l’utilisation de ces systèmes. Dans notre cas, l’objectif du travail était d’évaluer la stabilité

thermodynamique et de l’inertie chimique de complexes de gadolinium pour l’imagerie IRM.

Les ligands utilisés sont des ligands polyaminocarboxylates macrocycliques ou linéaires

fonctionnalisés. Dans le cas des ligands macrocycliques, la cavité est de type DO3A et les

groupements fonctionnels sont des groupements de type méthyl-benzimidazole. Pour les

ligands linéaires, le ligand est un DTPA bisamide et les groupements fonctionnels sont des

fonctions thiol, permettant d’envisager le greffage de ces ligands sur des nanoparticules d’or.

La démarche utilisée a consisté à étudier dans une première étape le comportement

acido-basique de tous les systèmes puis, sur la base des constantes de protonation obtenues,

d’évaluer l’affinité des différents ligands vis-à-vis d’un ensemble de cations métalliques. Si le

cation cible de l’étude est le gadolinium Gd(III), l’affinité des différents ligands a été

déterminée par rapport à ce cation et par rapport à des cations potentiellement compétiteurs

(Cu(II), Zn(II)). Lorsque cela a été possible, les différentes espèces complexes présentes en

solution ont été caractérisées en s’appuyant sur l’apport de résultats en spectroscopie UV, 1H

RMN, RPE et fluorescence. La seconde étape du travail a consisté à évaluer par relaxométrie

l’inertie chimique des complexes de Gd(III) en présence d’un ion compétiteur, le Zn(II).

Pour les systèmes macrocycliques L1H4 et L2H3, quel que soit le cation, les

complexes basés sur le ligand L1H4 sont les plus stables. De plus, la stabilité de ces

complexes suit l’ordre d’affinité croissant M-DO3A<M-L1H4<M-DOTA. Ceci indique que

l’adjonction du groupe benzimidazole renforce les capacités de complexation du macrocycle

DO3A bien que l’atome d’azote imine coordinant du groupe benzimidazole soit un moins

bon atome donneur qu’un groupement carboxylate du DOTA. Du point de vue inertie

chimique, le complexe Gd(III)-L1H4 présente une très bonne inertie chimique puisqu’elle est

comparable à celle de la référence Gd(III)-DOTA. Ce résultat s’explique par le fait que le

28

ligand L1H4 répond parfaitement aux exigences stéréoélectroniques du Gd(III) en termes de

nombre d’atomes coordinants et de préorganisation du squelette organique.

Pour les systèmes L@1H5 et L@

2H3, les études potentiométriques montrent à nouveau

que les complexes de Gd(III) sont plus stables que ceux de Zn(II) et de Cu(II). Le point à

remarquer dans cette étude est que les propriétés acido-basiques et les propriétés de

complexation du ligand L@1H5 sont modifiées lorsque celui-ci est greffé à la surface de la

nanoparticule. Ainsi, la basicité du ligand est renforcée, la stabilité des complexes s’en

trouvant de fait améliorée. Pour expliquer ce résultat, on peut suggérer que la structure de la

couche organique de ligands à la surface de la nanoparticule favorise la stabilisation des

charges introduites par un réseau de liaisons hydrogène et des réorganisations

conformationnelles. Ce gain de stabilité thermodynamique s’accompagne pour le complexe

greffé à la surface de la nanoparticule par un gain important d’inertie chimique puisque les

études relaxométriques montrent que, la demi-vie du complexe greffé est deux fois plus

importante que celle du complexe libre. Du point de vue de l’utilisation en imagerie IRM de

ces complexes de Gd greffés sur nanoparticules, ces gains en stabilité thermodynamique et en

inertie chimique sont certainement des atouts importants. Du point de vue analytique, ces

résultats contribuent aussi à montrer qu’on ne peut pas se baser sur les propriétés d’un

complexe en solution pour extrapoler ces propriétés au complexe engagé dans un système

plus organisé tel que le réseau étudié ici.

29

Sommaire

31

Introduction générale p 39

General Introduction p 43

Glossaire p 47

Chapitre-I: Introduction p 51

A. Introduction p 53

B. Lanthanides p 54

1. Propriétés fondamentales des lanthanides p 54

a) Configurations électroniques p 54

b) Lanthanides et Orbitales f p 55

c) Niveaux d'énergie des ions lanthanides et spectres électroniques p 57

2. Spectres d'absorption des ions lanthanides p 59

3. Spectres d'émission des ions lanthanides p 60

a) Mécanismes de luminescence des Ln (III) p 61

b) Sensibilisation des lanthanides, effet d'antenne p 62

4. Propriétés magnétiques des ions lanthanides p 63

5. Applications de complexes de lanthanides p 63

6. Applications des complexes de Ln(III) en luminescence p 63

a) Les sondes luminescentes p 64

b) Sondes immuno-luminescentes p 65

c) Imagerie de luminescence p 66

C. Imagerie par résonance magnétique nucléaire p 67

1. Contexte général p 67

2. Principe de l’IRM p 67

D. Les agents de contraste en IRM p 70

1. Chélates des gadolinium et agents de contraste en IRM p 70

a) Agents de contraste commerciaux p 71

b) Relaxivité des complexes de gadolinium p 72

2. Toxicité des agents de contraste à base de gadolinium p 73

E. Fibrose Néphrogène Systémique (FSN) et complexes

de Gadolinium p 74

32

1. Stabilité thermodynamique des complexes de gadolinium

Complexes p 75

2. Inertie chimique p 79

a) Conditions stoechimétriques p 80

b) Conditions de pseudo premier ordre p 82

i. Inertie chimique des complexes linéaires p 84

ii. Inertie chimique des complexes macrocycliques p 92

F. Objectifs de la thèse p 94

Chapitre-II: Etude physicochimique des ligands macrocycliques

L1H4, L2H3 et de leurs complexes p 101

A. Inétrêt des ligands de type azole comme substituants

des ligands DO3A p 103

B. Synthèse et études physicochimiques des ligands

L1H4 and L2H3 p 108

1. Synthèse du méthyl-benzimidazole-DO3A (L1H4) et du p-nitrophenyl

benzimidazole-DO3A (L2H3) p 108

a) Synthèse de L1H4 et de L2H3 p 108

b) Détermination de la formule brute des ligands L1H4 et L2H3 p 110

2. Propriétés Acido-basic properties de L1H4 et de L2H3 p 111

a) Etude potentiométrique des ligands p 111

b) Détermination des constantes de protonation p 113

c) Etude spectroscopique des ligands p 115

i. Etudes UV-Visible p 115

ii. Spectroscopie RMN p 118

C. Complexation de L1H4 et de L2H3 p 123

1. Etudes physicochimiques avec Cu (II) et Zn (II) p 123

a) Etudes potentiométriques p 123

b) Stabilité thermodynamique des complexes de Cu(II) et de Zn(II) p 126

c) Etude spectroscopique des complexes de Cu (II) et de Zn (II) p 131

33

i. Etudes UV – visible p 131

ii. Spectroscopie RMN p 133

iii. Spectroscopie RPE p 136

d) Hypothèses structurales pour M-L1H4 et M-L2H3, où

M = Cu(II) et Zn(II) p 141

2. Etudes physicochimiques avec Gd (III) et Eu (III) p 143


b) Stabilité thermodynamique des complexes de lanthanide(III)

avec L1H4 et L2H3 p 144

c) Etudes spectroscopiques des complexes de Gd (III) et Eu(III)

avec L1H4 et L2H3 p 147

i. Etudes UV – visible p 147

ii. Fluorescence p 149

d) Hypotheses structurales pour les complexes Ln-L1H4 and Ln-L2H3 où

Ln = Gd(III) and Eu(III) p 152

e) Transmétallation avec Zn(II) – Mesures relaxométriques p 153

D. Conclusion p 155

Chapitre-III: Etude physicochimique du ligand linéaire L@1H5 et

de ce ligand greffé sur nanoparticule d’or : mécanismes de

transmétallation de leurs complexes de Gd(III) pour un pH voisin

du pH physiologique p 159

A. Intérêt des nanoparticules comme agents de contraste p 161 1. Agents de contraste nanoparticulaires en IRM p 161 2. Nanoparticules d’or et chélates de gadolinium p 162

B. Synthèse, études physicochimiques des ligands L@1H5 et L@

2H3 de leurs

complexes de Cu(II), Zn(II), Ca(II), Na(I) et Gd(III) p 168

1. Synthèse et études physicochimiques des ligands p 168

a) Synthèse de L@1H5 et L@2H3 p 168

b) Etudes potentiométriques p 170

c) Détermination des constantes de protonation pour

34

L@1H5 et L@2H3 p 171

2. Etudes physicochimiques des complexes métalliques basés sur les

ligands L@1H5 et L@2H3 p 175


b) Détermination des constantes de stabilité des complexes de Cu(II), Zn(II),

Ca(II), Na(I) et Gd(III) p 176

3. Etudes spectroscopiques des complexes basés sur le ligand L@1H5 et le

système L@4H5 p 182

C. Etudes de transmétallation des systèmes Gd(III)-L@1H5 et Gd(III)-

L@2H3 p 187

1. Démétallation suivie par relaxométrie p 187

2. Transmétallation suivie par spectroscopie UV dans des conditions de

pseudo-premier ordre p 189

D. Conclusion p 193

Chapitre-IV: Partie expérimentale p 199 A. Synthèse des ligands L1H4, L2H3, L@1H5 et L@2H3 p 201

B. Analyse élémentaire p 202

C. Etudes potentiométriques p 203

D. Spectroscopie RMN p 208

E. Spectroscopie RPE p 209

F. Relaxométrie p 209

G. Spectroscopie UV-visible p 210

Conclusion p 219

Conclusion générale p 223

35

Table of contents

Introduction générale p 39

General Introduction p 43

Glossary p 47

Chapter-I: Introduction p 51

A. Introduction p 53

B. Lanthanides p 54

1. Fundamental properties of lanthanides p 54

a) Electronic configuration p 54

b) Lanthanides and f Orbitals p 55

c) Energy level diagrams for lanthanide Ions, and their

electronic Spectra p 57

2. Absorption spectra of lanthanide Ions p 59

3. Emission spectra of lanthanide Ions p 60

a) Mechanisms of Ln(III) luminescence p 61

b) Luminescence sensitization or antenna effect p 62

4. Magnetic properties of lanthanide Ions p 63

5. Applications of lanthanide complexes p 63

6. Applications of Ln(III) complexes as luminescent probes p 63

a) Luminescent probes p 64

b) Immunoassays p 65

c) Imagery p 66

C. Magnetic Resonance Imaging p 67

1. General background p 67

2. Principle of MRI and MRI devices p 67

D. Contrast agents in MRI p 70

1. Gadolinium chelates as contrast agents for MRI p 70

36

a) Clinical contrast agents p 71

b) Relaxivity of Gadolinium complexes p 72

2. Toxicity of gadolinium contrast agents p 73

E. Nephrogenic systemic fibrosis (NSF) and Gadolinium

contrast agents p 74

1. Thermodynamic stability of gadolinium complexes p 75

2. Kinetic inertness p 79

a) Stoichiometric conditions p 80

b) Pseudo first order conditions p 82

i. Kinetic inertness of linear complexes p 84

ii. Kinetic inertness of macrocyclic complexes p 92

F. Scope of the thesis p 94

Chapter-II: Physico-chemical studies of macrocyclic ligands and

their metal complexes p 101

A. Interest of azoles as substituents for DO3A derivatives. p 103

B. Synthesis and physicochemical studies of ligands

L1H4 and L2H3 p 108

1. Synthesis of methyl benzimidazole-DO3A(L1H4) and p-nitrophenyl

substituted benzimidazole-DO3A (L2H3) p 108 a) Synthesis of L1H4 and L2H3 p 108

b) Determination of the empirical formula of

ligands L1H4 and L2H3 p 110

2. Acido-basic properties of L1H4 and L2H3 p 111

a) Potentiometric study of Ligands p 111

b) Determination of protonation constants p 113

c) Spectroscopic studies of ligands p 115

i. UV-Visible studies p 115

ii. NMR spectroscopy p 118

C. Coordination chemistry of L1H4 and L2H3 p 123

1. Physicochemical studies with Cu (II) and Zn (II) p 123

37

a) Potentiometric study p 123

b) Thermodynamic stability of transition metal (II) complexes p 126

c) Spectroscopic study of Cu(II) and Zn(II) complexes p 131

i. UV – visible studies p 131

ii. NMR spectroscopy p 133

iii. EPR spectroscopy p 136

d) Structural hypotheses for M-L1H4 and M-L2H3, where M = Cu(II) and

Zn(II) p 141

2. Physicochemical studies with (Gd (III) and Eu (III)) complexes p 143


b) Thermodynamic stability of lanthanide(III) complexes p 144

c) Spectroscopic study of Gd (III) and Eu(III) complexes p 147

i. UV – visible studies p 147

ii. Fluorescence measurements p 149

d) Structural hypotheses for Ln-L1H4 and Ln-L2H3 complexes, where Ln =

Gd(III) and Eu(III) p 152

e) Transmetallation with Zn(II) – Relaxometric measurements p 153

D. Conclusion p 155

Chapter-III: Physico-chemical studies of linear ligands and their

metal complexes; investigation of transmetallation mechanisms

near physiological pH p 159

A. Interest of nanoparticles(Np) as contrast agents p 161 1. Nanoparticle-based contrast agents in context of MRI p 161 2. Gold Nanoparticles and Gd chelates p 162

B. Synthesis, physicochemical studies of ligands (L@1H5 and L@

2H3) and

their metal complexes (Cu(II), Zn(II), Ca(II), Na(I) and Gd(III)) p 168

1. Synthesis and physicochemical studies of ligands p 168

a) Synthesis of L@1H5 and L@2H3 p 168

b) Potentiometric study p 170

c) Determination of protonation constants of L@1H5 and L@2H3 p 171

38

2. Physicochemical studies of metal complexes with ligands L@1H5 and

L@2H3 p 175


b) Determination of stability constants of Cu(II), Zn(II), Ca(II), Na(I) and

Gd(III) complexes p 176

3. Spectroscopic study of complexes with ligands L@1H5

and L@4H5 p 182

C. Transmetallation studies of Gd(III)-L@1H5 and

Gd(III)-L@2H3 p 187

1. Kinetics of demetallation followed by relaxometric

measurements p 187

2. Kinetics of transmetallation followed by UV spectroscopy p 189

C. Conclusion p 193

Chapter-IV: Experimental section p 199 A. Synthesis of ligands L1H4, L2H3, L@1H5 and L@2H3 p 201

B. Elemental analysis p 202

C. Potentiometric study p 203

D. NMR Spectroscopy p 208

E. EPR spectroscopy p 209

F. Relaxometry p 209

G. UV visible spectroscopy p 210

Conclusion p 219

Conclusion générale p 223

39

Introduction générale

40

L’utilisation d’agents de contraste à base de gadolinium pour l’imagerie IRM justifie

les nombreux efforts effectués pour augmenter leur innocuité et leur efficacité. Dans ce

contexte, l’utilisation de ligands polyaminocarboxylates pour la complexation de lanthanides

en général, et du gadolinium en particulier, repose sur l’affinité remarquable de ces ligands

pour ces ions. Cette affinité se décline d’une part par une forte stabilité thermodynamique

vis-à-vis des lanthanides et d’autre part par une bonne inertie chimique vis-à-vis de la

démétallation pour les complexes correspondants. L’intérêt de ces ligands réside aussi dans la

possibilité de les fonctionnaliser sélectivement. Par exemple, il est possible de rendre ces

ligands (et donc ces complexes) intelligents en leur permettant via une fonctionnalisation

adaptée, d’atteindre spécifiquement certains types cellulaires. Il est aussi possible par

ingénierie moléculaire du ligand, de coupler plusieurs modalités d’imagerie. En matière de

multimodalité, l’association de ligands polyaminocarboxylates avec des nanoparticules

métalliques telles que les nanoparticules d’or est un bon exemple. Il n’en reste pas moins

qu’avant toute application in vivo, la stabilité thermodynamique et l’inertie chimique de ces

nouveaux systèmes doit être testée.

Dans le contexte de cette thèse quatre composés ont été étudiés, deux d’entre eux

étant basés sur des ligands macrocycliques dérivés du DO3A (L1H4 and L2H3). Les

groupements fonctionnels additionnels sont des dérivés de benzimidazole. Les deux autres

systèmes sont basés sur des ligands linéaires dérivés du DTPA bisamide (L@1H5 et L@

2H3)

pour lesquels les fonctions amide terminales portent des fonctions thiol. Le ligand de base

L@1H5 greffé sur nanoparticule d’or conduit au système L@

2H3. Afin de statuer sur l’affinité

de ces quatre systèmes vis-à-vis des lanthanides et en particulier du gadolinium, notre

démarche a consisté à comparer pour chaque système son affinité pour Gd(III) et pour des

ions potentiellement compétiteurs en milieu biologique.

Le premier chapitre de cette thèse décrira les propriétés physiques des lanthanides qui

peuvent être pertinentes en imagerie, particulièrement en imagerie IRM. Ce chapitre montrera

aussi la nécessité de fortement chélater ces ions pour que les complexes correspondants

puissent être candidats à une mise sur le marché comme agents de contraste IRM.

Le second chapitre s’intéressera à la mise en évidence de la stabilité

thermodynamique de complexes issus des ligands L1H4 and L2H3. Ces ligands

macrocycliques ont été synthétisés par l’équipe du Pr S. J. Archibald (Université de Hull).

41

L’objectif sera d’évaluer par pH-métrie, la stabilité thermodynamique des complexes de

Cu(II), Zn(II), Gd(III) et Eu(III) obtenus à partir des ligands précités. Des études en

spectroscopie UV, 1H NMR, RPE et fluorescence viendront conforter si nécessaire les

hypothèses structurales émises. L’inertie chimique du complexe de gadolinium Gd(III)-L1H4

sera ensuite été évaluée par relaxométrie et comparée à celle d’un agent de contraste

commercial (DOTAREM®). Nous montrerons qu’aucune démétallation n’est détectée pour

ce complexe Gd(III)-L1H4, ce qui est un bon premier point dans l’optique de son utilisation

comme agent de contraste IRM.

Le troisième chapitre de la thèse s’intéressera à la mise en évidence de la stabilité

thermodynamique de complexes issus du ligand L@1H5 et des nanoparticules d’or associées

(L@2H3). Ces systèmes ont été synthétisés par l’équipe du Pr S. Roux (Université de Franche

Comté). L’approche développée dans le Chapitre II sera transposée dans ce chapitre. Après

l’évaluation de la stabilité thermodynamique des complexes de gadolinium Gd(III)-L@1H5

etGd(III)- L@2H3, leur inertie chimique sera evaluée par relaxométrie et une proposition de

mécanisme de transmétallation sera faite. Nous montrerons en particulier que le greffage du

ligand sur la nanoparticule se traduit pour le complexe correspondant par un gain de stabilité

thermodynamique et un renforcement conséquent de son inertie chimique. Ceci est aussi un

bon point pour l’utilisation de ces nanoparticules d’or greffées par des complexes de

gadolinium, dans le cadre d’applications en imagerie multimodale.

42

General Introduction

44

Since gadolinium contrast agents are used in magnetic resonance imaging (MRI),

numerous efforts have been done to increase their safety and obviously efficiency. In such a

case, the interest for polyaminocarboxylate ligands and among them those who are based on

tetraazamacrocyclic frameworks, undoubtedly lies in their remarkable lanthanide complexing

ability. These capabilities are reflected in the complex by a large thermodynamic stability,

but also by kinetic stability towards demetallation. The functionalization ability to change

these ligands helped to develop the chemistry of these compounds and to consider for

instance, the possibility to target these complexes towards specific cells or to improve their

imaging efficiency by enhancing their sensitivity or by combining imaging facilities. In that

respect, functionalization of these ligands in order to associate them to nanoparticles is a

constantly evolving research field. Nevertheless, before any in vivo applications the

thermodynamic stability and the kinetic inertness of each new candidate must be tested.

In this context, four compounds have been studied in this thesis, in which two of them

are macrocyclic based on a DO3A backbone functionalized by benzimidazole derivatives

(L1H4 and L2H3). The two others are constituted of linear ligands derived from DTPA

bisamide framework. L@1H5 is a dithiolated DTPA bisamide ligand, which was further

grafted onto a gold nanoparticle (the resulting grafted nanoparticle being L@2H3). In order to

status on the ability of these four ligands to complex efficiently lanthanides and particularly

gadolinium, our approach was to determine for each ligand the affinity for this ion and to

compare it with the affinity of each ligand for ions that potentially can act as competitors in a

biological context.

In the first chapter of this thesis, is dedicated to lanthanide properties and the

applications of lanthanide complexes in imaging, particularly in Magnetic Resonance

Imaging. This chapter will also focus on the necessity to strongly entrap lanthanide ions and

particularly gadolinium for being able to market gadolinium contrast agents fro MRI.

The second chapter of this thesis includes physicochemical studies of L1H4 and L2H3.

These macrocyclic ligands were synthesized by Pr. S.J. Archibald group. The objective was

to evaluate the thermodynamic stability of metal complexes where the metal can either be

first row metal ions (mainly Cu(II) and Zn(II) or a lanthanide (Gd(III) and Eu(III).

Spectroscopic studies were developed to have an insight of the metal coordination sphere,

when necessary. The kinetic stability towards demetallation of the gadolinium complex of

45

DO3A-methylbenzimidazole (Gd(III)-L1H4) was further evaluated by means of relaxometry

and compared to the one of a commercially available contrast agent (DOTAREM®). We will

demonstrate that no demetallation is detected for Gd(III)-L1H4 which is a good thing, in the

view of Gd(III)-L1H4 utilization as a contrast agent in MRI applications.

The third chapter of this thesis includes physicochemical studies of a dithiolated

DTPA bisamide ligand (L@1H5) and of gold nanoparticles grafted with this ligand (L@

2H3).

These systems were synthesized by Pr. S.Roux group. The approach developed in the

previous chapter was transposed for these systems. After evaluation of thermodynamic

stability of the corresponding gadolinium complexes, their kinetic inertness was evaluated by

relaxometry and a proposition of a transmetallation mechanism was attempted. Comparative

transmetallation experiments of Gd(III)-L@1H5 and Gd(III)-L@

2H3 highlighted that when the

ligand is grafted onto the nanoparticule the stability and the kinetic inertness of the

corresponding gadolinium complex are greatly enhanced. This is also a good point for the

possible use of these nanoparticles in living organisms for multimodal imaging applications.

47

Glossary

49

NH N

NH HN

HN

N

L3H

N N

NN

CO2H

HO2CHO2C

CO2H

L4H4

N N

NN

CO2H

HO2CHO2C

OHOH

HO

L5H3

N N

N

N

CO2H

HO2C

HO2C

L6H3

N N N

COOHHOOC

COOH

O

NH

CH3

O

HN

H3C

L@3H3

N N N

COOH

COOHHOOC

HOOC

COOHL@

4H5

51

Chapter-1

Introduction

53

A. Introduction

This chapter is dedicated to bibliographic review around the chemistry of lanthanides

and their biomedical applications.

The first section of this chapter reviews the fundamental electronic properties of lanthanide

elements, especially their electronic energy level diagrams as a basis for interpretation of

their absorption and emission spectra. Focus will be on the emission mechanisms of the

lanthanides and in particular the amplification of this phenomenon by the antenna effect that

can have light-absorbing ligands. A statement will also be made about the magnetic

properties of these ions.

The second section of this chapter is to describe the possible applications for

lanthanide ions basing on their electronic properties previously recalled.

The third section of this chapter is to describe in more detail the biomedical

applications of lanthanides, such as luminescent probes and magnetic probes.

The last two sections of this bibliographic chapter endeavor to show the interest of

gadolinium complexes as contrast agents for MRI and to establish specifications for the use

of these complexes in imaging and ensure their safety for utilisation in patients. To be

specified at this level, the importance of the thermodynamic stability of these complexes, as

well as their chemical inertness towards any transmetallation reactions in biological media.

54

B. Lanthanides

1. Fundamental properties of lanthanides

Lanthanides occupy unique position in the periodic table, which correspond to the

first period of f-block elements.

a) Electronic configuration

The elements from cerium (Z = 58) to lutetium (Z = 71) constitutes the lanthanide

series. Due to similar chemical properties, lanthanum (Z = 57) is also considered as a member

of the lanthanide series (Table I-1).

Table I-1: Mendeleev table and lanthanide series

Lanthanides have similarities in their electronic configuration, which explains most of

their physical properties. These elements are different from the main group elements due to

the fact that they have electrons in their f orbital. The ground state electronic configurations

of lanthanides are gathered in Table I-2.

55

Table I-2: Ground state electronic configurations across the lanthanide series

Symbol Electronic configuration Symbol Electronic configuration

La [Xe] 5d1 6s2 Tb [Xe] 5d0 6s2 4f9

Ce [Xe] 5d1 6s2 4f1 Dy [Xe] 5d0 6s2 4f10

Pr [Xe] 5d0 6s2 4f3 Ho [Xe] 5d0 6s2 4f11

Nd [Xe] 5d0 6s2 4f4 Er [Xe] 5d0 6s2 4f12

Pm [Xe] 5d0 6s2 4f5 Tm [Xe] 5d0 6s2 4f13

Sm [Xe] 5d0 6s2 4f6 Yb [Xe] 5d0 6s2 4f14

Eu [Xe] 5d0 6s2 4f7 Lu [Xe] 5d1 6s2 4f14

Gd [Xe] 5d1 6s2 4f7

After lanthanum, the energy of the 4f orbitals falls below that of the 5d one (E4f (La) =

-0.95 eV while E4f(Nd) = -5 eV).[1] Therefore, this leads to the electron filling of the 4f

orbitals before the 5d one.

b) Lanthanides and f-orbitals

The f-orbitals affect the properties of lanthanides. As shown in Figure I-1, the radial

f-wavefunction overlaps appreciably the radial part of the Xe core wavefunction.

Figure I-1: Radial wavefunction of the three 4f electrons of Nd(III) compared with the radial

wavefunction of the xenon core (a.u. = atomic units).[2]

This renders the valence 4f ‘inner orbitals’. Therefore, the 4f orbitals are not

particularly effective at shielding the outer shell electrons (n=5 and n=6). 1 Because of this,

across the Ln series, there is a decrease greater than expected in both the atomic radii and in

1 Thus, the shielding effect is less able to counter the decrease in radius caused by increasing nuclear charge.

56

the radii of the Ln(III) ions (decrease from 102 pm for Ce(III) to 86.1 pm for Lu(III)). This

decrease corresponds to the so-called lanthanide contraction.

As a result of the different degrees of stabilization experienced by the 4f, 5d and 6s

orbitals upon ionization of the neutral metal, the lanthanides from La to Lu exist almost

exclusively in their trivalent state Ln(III) ([Xe] 4fn ; n = 0-14).[3]

Another consequence is that in coordination complexes, the ‘inner’ 4f orbitals cannot

overlap with ligand orbitals. Therefore covalence plays a minor role in Ln-ligand dative

bonds, leading to a poor stereochemical control in lanthanide containing edifices. The

coordination number (CN) adopted by a particular complex is determined by how many

ligands can be packed round the central metal ion. Variable coordination numbers (6 < CN <

12) and geometries are thus observed in lanthanide complexes. They are difficult to predict

because the hard Ln(III) ion will complete its co-ordination sphere by binding small

molecules or anions (water, chloride, hydroxide, etc.) if the number of available sites offered

by the ligand(s) is too low. Nevertheless, in the solid state, trivalent lanthanides display a

tendency to adopt nine-coordinate tricapped trigonal prismatic (TTP) arrangements around

the metal ion. (Figure I-2)

TTP SAP

Figure I-2: Various Ln(III) coordination spheres

In solution, large Ln(III) ions at the beginning of the series (La-Nd) adopt TTP

geometries, which are gradually transformed into eight-coordinate square antiprismatic (SAP)

arrangements for small Ln(III) ions (Tb-Lu), equilibria between CN = 8 and CN = 9 being

observed for Ln = Nd-Tb.[4] Since 4f orbitals in Ln coordination complexes are shielded from

the effects of surrounding ligands, crystal-field (CF) effects are not found in transition metal

57

chemistry (CF splitting the order of ~ 100 cm-1 in lanthanide complexes against 15,000-

25,000 cm-1 for d-transition metal complexes).[5]

CF effects are weaker than Russell-Saunders coupling (between orbital angular L and

spin angular S momenta of electrons2) and should only be considered as a second perturbation

of the free-ion levels, LS perturbation being the first one. The corresponding Hamiltonian is

written as below: [6,7]

(1)

where the second term corresponds to electronic repulsion effects, the third one to Russell-

Saunders coupling and the last one to ligand field effects.

c) Energy level diagrams for lanthanide ions and their electronic spectra

The ground state electronic configuration of Ln(III) ions is [Xe] 4fn (n = 0-14). The

excited states are well separated from the ground state, due to electronic repulsion between

electrons. For instance the [Xe] 4fn-1 5d1 configuration is separated from the ground state by

E > 32,000 cm-1. [6] As an example, electronic levels of the ground state and the first excited

state of Eu(III) are given in Figure I-3.

Figure I-3: Ground state (7F) and first excited state (5D) of Eu(III)

2 The spins (s) of the individual electrons are coupled together (added vectorially) to give the spin quantum number for the ion (S). The orbital angular momenta (ℓ) of the individual electrons are coupled similarly to give the angular quantum number of the ion (L).

58

The Russell-Saunders coupling (LS coupling) introduces a new quantum number J, associated

with the total angular momentum with values ranging from (L+S) to |L-S|. As a

consequence each state associated to one value of L and S (or spectroscopic term) is further

split into a number of spectroscopic levels 2S+1𝜞J each with a (2J+1) multiplicity (Table I-3).

Table I-3: Electronic levels of Ln(III) resulting from L-S coupling on the free ions.[5]

The CF effects are weak in Ln(III) complexes and the corresponding perturbation has

to be treated last. The (2J+1) degeneracy of the electronic levels split under the influence of

ligand-field potential. The number of ligand-field sublevels depends on the value of J and are

collected in Table I-4.[8]

59

Table I-4: Number of ligand-field sublevels versus site symmetry and the value of the J

quantum number.[8]

It should be noted that the electrostatic ligand-field effects do not completely lift the

degeneracy of the J levels of odd-numbered electronic configurations. In some situations, all

ligand-field sublevels are at least doubly degenerated (Kramer’s doublets).

2. Absorption spectra of lanthanide ions

The interaction between light and matter is ruled by operators linked to the nature of

light. Three operators are concerned: the odd-parity electric dipole (ED) operator P, the even-

parity magnetic dipole (MD) M and the electric quadrupole (EQ) Q. The expression of these

operators is given below:

(2)

Most lanthanide ions absorb electromagnetic radiation, particularly in the visible

region of the spectrum. They are involved in three types of electronic transitions:

intraconfigurational 4f-4f transitions, 4f-5d transitions and charge-transfer transitions (metal-

to-ligand MLCT or ligand-to-metal LMCT).

Laporte’s parity selection rules indicate that states with the same parity cannot be

connected by the ED interaction. As a consequence, f-f transitions are forbidden.

60

Nevertheless, for lanthanide under the influence of a ligand field, non centrosymmetric

interactions allow the mixing of electronic states of opposite parity into the 4f functions,

which loosen the selection rules.[6] Consequently, the transition probability occurrence

increases. Magnetic dipole transitions MD are allowed by parity but their intensity is weak.

The f-f transitions are also excited by the MD mechanism even though their amplitude is

often as the same order of magnitude as those induced by the ED mechanism. Quadrupolar

transitions QD are also parity allowed but they are much weaker than MD tranisitons and are

generally difficult to observe. The intensities of f-f transitions are modeled by the Judd-Ofelt

theory. Some Ln(III) ions have weak intensities f-f transitions, leading to colourless

solutions.

The electronic spectra of lanthanide compounds resemble those of the free ions, due

to the crystal-field influence being weak. This induces sharp f-f transitions in spectra (Figure

I-4).

Figure I-4: Excitation spectrum of at 77K of Eu(NO3)3·15-crown-5 (λanal = 618 nm).[5]

Because of the quasi negligible influence of the ligand field, electronic spectra of

lanthanide complexes can rarely be used to get information about metal coordination

geometry. Some transitions constitute an exception since they can be hypersensitive to

changes in the symmetry and strength of the ligand field. As a result, they display shifts of

the absorption bands, usually to longer wavelength, as well as band splitting and intensity

variation. They are called ‘hypersensitive’ bands and are encountered for Ho(III), Er(III) and

Nd(III).[5]

3. Emission spectra of lanthanide ions

All Ln(III) are luminescent, excepted La(III) and Lu(III). Their emitted light is

constituted of sharp lines characteristic of f-f transitions of Ln(III) ion. For the Ln series,

these lines cover the entire spectrum,[6] from UV (Gd(III)) to visible (Pr(III), Sm(III), Eu(III),

Tb(III), Dy(III), Tm(III)) and near-infrared (Pr(III), Nd(III), Ho(III), Er(III), Yb(III)) (Figure

I-5).

61

Figure I-5: Normalized emission spectra of luminescent lanthanide complexes in solution

from http://www.pitt.edu/~dave/Lanthanide_spec.html

Some ions are fluorescent (S = 0), others are phosphorescent (S = 1) and some are both.

a) Mechanisms of Ln(III) luminescence

The Jablonski-Perrin diagram gives a description of the Ln(III) emission mechanism (Figure

I- 6).

Figure I-6: Luminescence in Lanthanide complexes – Perrin Jablonski diagram

When a photon is absorbed, an electron is promoted to an excited singlet state. Direct

de-excitation of this singlet excited state to the ground state is accompanied by photon

emission. This corresponds to fluorescence. When conversion of the singlet excited state into

a triplet excited state occurs via an intersystem crossing (IC), de-excitation of the triplet state

to the ground state is also accompanied by photon emission. This corresponds to

phosphorescence. In the lanthanide series, some ions are fluorescent, others are

phosphorescent and some are both. In the excited states, the internuclear distances remain

http://www.pitt.edu/~dave/Lanthanide_spec.html

62

almost the same. Consequently, the f-f emission lines are sharp and very small Stokes’ shifts

are measured.

b) Luminescence sensitization or Antenna effect

The dipole strength of f-f transitions are very small and direct excitation into the 4f

excited levels rarely yields to high luminescence. To circumvent this drawback, an alternative

path has been explored to improve lanthanide luminescence. This mechanism corresponds to

luminescence sensitization or antenna effect (Figure I-7).

Figure I-7: Lanthanide Sensitization by a light harvester (ligand) – de-excitation pathways

are omitted for clarity

In this case, the Ln(III) excited state sensitization is reached by intramolecular energy

migrations from the excited states of the ligand used for the metal complexation. Certain

Ln(III) ions have excited states lying slightly lower in energy than the triplet states of their

surrounding ligands. For Eu(III) and Tb(III), their most likely acceptor levels are 17,200 and

20,400 cm−1 respectively, so the triplet level in the acceptor ligand needs to be above 22,000

cm−1.[5] By resonant energy transfer mechanisms between excited states of the ligand and the

metal ion, lanthanide fluorescence is greatly improved. Then, Eu(III) and Tb(III) luminesce

with green and red colours, respectively. For Tb(III), the main emissions are 5D4 →7Fn (n=6–

0) with 5D4 → 7F5 the strongest. For Eu(III), 5D0 →7Fn are seen (n=4–0) with the main ones 7F0, 7F1, and 7F2 the most useful.

63

4. Magnetic properties of lanthanide ions

Except La(III) and Lu(III), all lanthanide ions are paramagnetic. Their magnetic

properties are determined by their ground state, as their excited states are separated from the

ground state by about 20,000 cm-1 (excepted for Sm(III) and Eu(III) which possess thermally

accessible excited states able to contribute to the magnetic properties). The magnetic moment

of the Ln(III) ions is essentially independent of environment, so that one cannot distinguish

between coordination geometries, as is sometimes possible for transition metals.[5] The

corresponding magnetic moments are given as below:

where gJ corresponds to the Landé factor (3)

This expression takes into account again, the coupling between orbital and spin components.

5. Applications of lanthanide complexes

Over the last twenty years, lanthanide study is being motivated not only by many

medical imaging applications,[9, 10] but also in medicine (analgesic treatment with radioactive

isotopes),[11] biology (catalysing hydrolysis of ADN),[12, 13] materials chemistry (converters of

light),[14, 15] and in the nuclear reprocessing (actinides / lanthanides separation).[16, 17]

Tb(III) and Eu(III) ions have emission in the visible region (green λem = 550 nm and

λem = 620 nm red respectively). These are the most commonly used ions for their application

as luminescent probes for analysis and biomedical diagnostics. To a lesser extent, Sm(III)

(orange λem = 590 nm) and Dy(III) (yellow-orange λem = 570 nm) are also used. The use of

ion emission in infrared (Nd(III): λem = 1065 nm, Er(III): λem = 1509 nm and Yb(III): λem =

978 nm) is booming. These are generally used in the development of optical fibers, lasers and

amplifiers for telecommunications, but now are acquiring considerable interest in the

development of luminescent probes for biomedical analysis. The Gd(III) ion, however, is not

used because of its emission in ultraviolet interferes with the absorption or emission of

organic and biological compounds. Nevertheless, this ion presents good characteristics to

assist diagnostic in MRI applications. In the following section we will focus on biomedical

devices in which Ln(III) are involved.

6. Applications of Ln(III) complexes as luminescent probes

There are three main applications of luminescent complexes in the biomedical field:

luminescent chemical probes, immunoassays and hybridization tests, and medical imaging.

64

a) Luminescent probes

As far as concerned with the probe, detection of ion or molecular species (analyte) is

performed by presence or absence of a light signal ("on-off" system). The modulation of the

luminescence can occur by three main mechanisms (Figure I-8).

Figure I-8: Modulation of the luminescence of Ln(III) by the reversible association of

analyte (an): a) direct influence on the lanthanide luminescence b) influence on the

photophysical properties of ligand, c) addition of a sensitizing analyte onto a low luminescent

Ln- containing sensor.[15]

First, the analyte can interact directly with the coordination sphere of lanthanide. This

mechanism is exploited for the detection of anions (Figure I-8.a). We know that

luminescence of the metal is "quenched" by the presence of water molecules bound to the

metal center. The interaction with these anions helps in restoring luminescence by replacing

the water molecules. Important work in this regard was made by D. Parker et al .[19] Second,

the luminescence of the complex can be influenced by the photophysical properties of ligand

for example by changing the energy level of its excited state (Figure I-8.b and c). The non-

radiative processes disturb the bands of charge transfer or photo-induced electron transfer of

a receiver in the lanthanide complex. This process is often used for the detection of metallic

cations and development of the pH sensitive systems.[20] A cation complexation by the

receiver removes the electron transfer and luminescence is the main relaxation process. Such

systems have been developed for the detection of major endogenous cations: Cu(II), Zn(II),

Mn(II), Ca(II), Mg(II), K(I) and Na(I).[20-23]

65

Finally, systems similar to those used for the development of MRI contrast agents

sensitive to pH were also used. For these systems, therefore the number of donor atoms and

the number of water molecules in the coordination sphere of the metal varies with the pH and

cause modulation of vibrational de-excitations.[24]

b) Immunoassays

Immunoassays are based on a biochemical reaction between an antigen (analyte) and

a specific antibody labeled with a fluorescent probe. Several technologies are used and

marketed with lanthanide complexes: the heterogeneous and homogeneous tests (Figure I-9).

Figure I-9: Principle of Immunoassays.[15]

The DELFIA heterogeneous system (Dissociation Enhanced Lanthanide

FluoroImmunoAssays) was developed by the group of I. Hemmilä in the 1980s.[25] A

luminescent Ln(III) complex is grafted to an specific antibody of interest (green one in

Figure I-9.a). The analyte-labeled antibody entity is put in the presence of a second antibody

attached to a solid support (blue one in Figure I-9.a). After washing the excess labeled

antibodys, Ln(III) ions are released in an acid medium. These ions are then complexed with a

photosensitizer (β-diketone) and the luminescence of this new complex is measured.

Homogeneous tests involve the reaction of an antigen with two different antibodies:

the first is labeled with the Ln(III) chelate (green one in Figure I-9.b), the latter with an

organic acceptor (blue one in Figure I-9.b). After excitation, the light emitted by the

lanthanide is transferred to the acceptor which emits a characteristic wavelength. Transfer

TR-FRET (Time-Resolved Fluorescence Resonance Energy Transfer) occurs only if the

donor and acceptor are close enough, i.e if they are attached to the same antigen. On this

66

principle, G. Mathis and his colleagues have made many immunoassays and hybridization

assays of DNA with cryptands of Eu(III) complexes using cyanine acceptor and other

systems.[26] They are now marketed by Cis Bio International.

c) Imaging

Currently the luminescence imaging technique is applicable only in vitro or for

experiments on small animals. Two types of luminescent markers are used for time-resolved

fluorescence imaging: lanthanide complexes emitting in the visible region (mainly Eu(III)

and Tb(III)) and those emitting in the infrared (mainly Yb(III)). The emission in infrared is a

considerable advantage since it allows overcoming the absorption of incident radiation by

water and biological tissues, and ultimately to explore deeper tissues. In all cases, research

efforts are directed towards the development of specific markers similar to those developed

as contrast agents for MRI strategies. Efforts were mainly focused on the detection of cancer

cells.[15] Ligands used are similar to those used for the complexation of Gd(III). Most are

derived from polyaminocarboxylate ligands, DOTA or DTPA-type, in which one (or more)

acetate (or phosphonate) arms was functionalized with a chromophore.[15] For example, D. J.

Bornhop et al. developed a cyclen derivative (Figure I-10) in which the luminescence of the

Tb(III) complex was observed in vivo in hamsters.[27]

N

N

N

N

N

P

P

P

OC4H9

O

O

O

C4H9O

OC4H9

O

O

O

Figure I-10: Cyclen derivative developped for Tb(III) luminescence and in vivo imaging in

hamsters.

67

C. Magnetic Resonance Imaging (MRI)

1. General background

Medical imaging has become a crucial part of the biomedical sciences, not only for

diagnosis within clinical medicine, but also for providing unique insights into evaluation of

disease, pathophysiology and the translation of novel treatments from the laboratory into

patients. There are excellent imaging facilities including Computed Tomography (CT)

involving the use of ionizing radiation for the tomographic analysis, Nuclear Medicine (NM)

involving the application of radioactive substances in the diagnosis and treatment of disease,

Positron Emission Tomography (PET) using short-lived radiolabeled substances to produce

powerful images of the body's biological function. Unlike CT, NM and PET, Magnetic

Resonance Imaging (MRI) does not involve the use of ionizing radiation. MRI is based on

the principle of nuclear magnetic resonance (NMR) and uses the magnetic properties of

hydrogen and its interaction with both external magnetic field and radio-frequencies (RF),

producing highly detailed images of the body. So, this non-invasive technique has become

one of the most important techniques in medical diagnosis and biomedical research.

2. Principle of MRI and MRI devices

The basic principle of MRI follows that of NMR. MRI measurement is generally

based on the detection of the water proton in human body tissues (water constitutes about

60% of the human body). Water protons respond when a gradient magnetic field is applied

and align either parallel or anti-parallel to the applied magnetic field (Figure I-11).

a.

b.

Figure I-11: a) In the absence of any externally applied magnetic field, individual dipoles

are randomly oriented. b) In the presence of external magnetic field (Bo), nuclei align and a

net magnetization is produced.

http://en.wikipedia.org/wiki/Radioactivity

http://en.wikipedia.org/wiki/Disease

68

Under these orientations, protons precess around the axis of the external magnetic

field. The rate at which the protons precess is given by an equation called Larmor equation.

(4)

where ɷo is the ‘Larmor frequency’, γ is the gyromagnetic ratio and B is the strength of the

external magnetic field.

The Larmor equation states that the precessional freqency ɷo is equal to the strength

of the external magnetic field Bo multiplied by the gyromagnetic ratio γ.

When a pulse of RF energy is applied at the Larmor frequency, the protons are said to

be resonant, and as well the net magnetization vector (Figure I-12). If the RF pulse

continues, some of the low energy state spins absorb energy from the RF field and make a

transition into the higher energy state.

Figure I-12: Upon the application of RF field, the net magnetic moment is disturbed and

begins to precess with the magnetic field Bo.

This transition has the effect of flipping the net magnetization towards the transverse

plane. A pulse that is sufficient to flip the net magnetization to MXY plane is called a 90°

pulse. When the pulse is off, the spins of protons become out of phase and rotate freely in the

transverse plane and tend to align in their natural way giving out excess energy. This

relaxation of protons occurs in two different ways namely the spin-lattice relaxation (T1) and

spin-spin relaxation (T2) seen in Figure I-13.

a. b.

Figure I-13: a. T1 and b. T2 relaxation measurements.

69

This process is known as relaxation and the produced signals are recorded and built as

MRI images (Figure I-14).

a. b.

Figure I-14: T1 MRI Image a. without and b. with injection of a contrast agent (from

Eurorad website)

To perform such examinations, MRI scanners include (i) a superconducting magnet

capable of producing a strong and homogenous magnetic field (ii) a radio frequency (RF)

transmitter and receiver system (iii) a gradient coil system, all three surrounding the patient

and (iv) a computer receiving the signals from the receiver coil and processing the signals

into images (Figure I-15).

Figure I-15: MRI scanner with superconducting magnet and the radio-transceiver which

causes the water protons to flip on their axes.[28]

Contrast between the tissues, in particular between the healthy and diseased tissues,

might be not sufficient to ensure the early detection of a tumor. Thus, in many cases, the

successful use of MRI would not have been possible without a class of pharmacological

products, called contrast agents. The role of the contrast agent is to accelerate the relaxation

of the surrounding protons, allowing reducing the examination time and improving the

contrast of the image.

70

D. Contrast agents in MRI

The expansion of medical MRI has given rise to the development of a new set of

pharmaceutics called contrast agents (CAs). These agents shorten the relaxation time of

nearby water molecules, thereby enhancing the contrast between the area of localization of

contrast agents and background. CA is composed of a paramagnetic metal center surrounded

by organic chelates.[29] The reduction in the relaxation time could be achieved by

paramagnetic metal centers such as Fe(III), Mn(II), Dy(III) and Gd(III). The unpaired

electrons in the paramagnetic center create a fluctuating magnetic field which is responsible

for increasing the relaxation rate of near by water protons.

1. Gadolinium chelates as contrast agents for MRI.

Various CAs based on Mn(II), Fe(III) and Cu(II) were studied, but about 25-30% of

MRI scans that involve CAs utilize gadolinium contrast agents (GdCAs) in their scans.[28]

The importance gained is explained by the fact that Gd(III) has seven unpaired electrons and

also has a relatively slow electronic relaxation (owing to its symmetric S-state), thereby

enhancing the relaxation of surrounding protons. Lanthanides such as Dy(III) and Ho(III)

have also higher magnetic moments, but have relatively short electronic relaxation when

compared to Gd(III). Thus Gd(III) gained its importance ruling out other paramagnetic metal

ions.

Table I-5: Paramagnetic ions that could be useful for MRI applications

Free Gd(III) ion is toxic for the human body. The ionic radius of Gd(III) (0.99Å) is

very nearly equal to that of Ca(II) and hence can compete with this cation in all biological

systems that require Ca(II) for proper function. Therefore, Gd(III) can alter the biological

71

process leading to complexities in human body.[30] It also can form hydroxo complexes and

precipitate at physiological pH. So Gd(III) metal centers should be encapsulated by strong

chelates forming stable complexes in order to reduce the intrinsic Gd(III) toxicity.

a) Clinical contrast agents

Presently, there are nine different commercially available gadolinium based contrast

agents (Figure I-16). The ligands are poly-aminocarboxylates with hard donor atoms such

anionic oxygen atoms and nitrogen atoms. These ligands are either linear, such as DTPA

(DTPA based ligands) or macrocyclic, such as DOTA (DOTA based ligands). They form

stable complexes with Gd(III) and these complexes are highly hydrophilic. All Gd complexes

are nine-coordinated, eight coordinating atoms belonging to the polyamino carboxylate

backbone and one bound water molecule.

N

N

N

N

-OOC

-OOCCOO-

COO-

Gd3+

NH2+

OH

OH

HO

OH

HO

Gd-DOTA, Dotarem ®

N

N

COO-

N

COO-

-OOC

COO-

COO-

Gd3+

NH2+

OH

OH

HO

OH

HO

2

N

N N

COO-

-OOC

COO-

COO-

Gd3+

NH2+

OH

OH

HO

OH

HO

2

COO-

O

Gd-DTPA, Magnevist ® BOPTA, Multihance ®

N

N

COO-

N

COO-

-OOC

COO-

COO-

Gd3+

O

SNa+

Na+

N

N

COO-

N

COO-

-OOC

COO-

COO-

Gd3+

SNa+

Na+

OP

O

O-

O

Na+

Gd-EOB-DTPA,

Primovist ®

MS 325,

Vasovist ®

N

N

N

N

-OOC

COO-

COO-

Gd3+

OH

N

N

N

N

-OOC

COO-

COO-

Gd3+

OH

OH

HO

N

N

C

N

C

-OOC

-OOC

COO-

Gd3+O

HN

O

HN

N

N

C

N

C

-OOC

-OOC

COO-

Gd3+O

HN

O

HN

O

O

Gd-HPDO3A,

Prohance ®

Gd-BTDO3A,

Gadovist ®

Gd-DTPA-BMA,

Omniscan ®

Gd-DTPA-BMEA,

OptiMARK ®

72

Figure I-16: Structures of clinically approved Gadolinium contrast agents.[31]

b) Relaxivity of gadolinium complexes

The aim of contrast agents is to increase the relaxation rate of water protons in the

surrounding tissue and thereby creating a contrast between the pathological area and normal

tissues inside the body. All contrast agents work by reducing the T1 and/or T2 relaxation times

of the target tissue. The contrast enhancement of the contrast agent is directly proportional to

its relaxation of neighbouring water molecules by a paramagnetic ion (Gd(III)). The

relaxation of solvent nuclei around a paramagnetic center has been described by Solomon,

Bloembergen and others.[32-37] The observed relaxation rate (1/Ti, obs) is the sum of

diamagnetic relaxation rate (1/Ti,d), corresponding to the relaxation rate without the

paramagnetic agent and the paramagnetic relaxation rate (1/Ti,p), corresponding to the

relaxation rate enhancement caused by the paramagnetic agent.

1/Ti,obs = 1/Ti,d + 1/Ti,p (5)

The paramagnetic contribution is linearly propotional to the concentration of the

paramagnetic species, [Gd].

1/Ti,p = ri [Gd] where i = 1.2 (6)

1/Ti,obs = 1/Ti,d + ri [Gd] where i = 1.2 (7)

Therefore, Proton relaxivity ri refers to the efficiency of a paramagnetic substance to

enhance the relaxation of water protons. The dipole-dipole interactions between protons

nuclear spins and the fluctuating local magnetic field caused by the unpaired electron spins

contribute mainly to the paramagnetic relaxation of the water protons. Relaxivity ri is the

resultant of inner sphere and outer sphere water molecules contributions (Figure I-17).

73

Figure I-17: Gd(III) complex with one inner sphere water molecule, surrounded by bulk water. τR stands for the rotational correlation time of the molecule, kex for the water proton

exchange rate and 1/T1,2e for the electron spin relaxation rates of Gd(III).[38]

The inner-sphere relaxivity (IS) is due to interaction between protons of the water

molecule bound in the first coordination sphere of the Gd(III) complex and the unpaired

electron spin of the paramagnetic ion. The outer-sphere relaxivity (OS) is more a

consequence of through-space interaction with protons outside the first coordination shell. IS

and OS mechanisms can be summarized by the equation as:

(1/Ti,p) = (1/Ti,p)IS + (1/Ti,p)OS (8)

ri = riIS + ri

OS (9)

Considerations for improving the relaxivity of Gd complexes were based on

modulation of Gd complexes hydration number, on their rotational correlation time and

residence lifetime of coordinated water molecule(s).[39]

2. Toxicity of gadolinium contrast agents After intravenous injection, gadolinium contrast agents are distributed in the

extracellular and intravascular spaces. Doses are required generally from 0.1 to 0.3 mmol.kg-1

and the concentration of the injected solution is important (about 0.5 mol L-1). The complex

should definitely be soluble at these concentrations. The second crucial point is the

thermodynamic and kinetic stabilities of the complex injected. The Gd(III) aqua-ion is very

toxic. Its toxicity is mainly due to the similarity of its ionic radius with Ca(II). It thus replaces

Ca(II) in the human body, sometimes blocks calcium channels and it is also known to interact

74

with serum proteins in such as irreversible process and settles down in the form of insoluble

salts in the bones, liver and spleen. Recently, a scleroderma was discovered in patients with

improper kidney function, who were undergone MRI imaging.[40]

E. Nephrogenic Systemic Fibrosis (NSF) and Gadolinium contrast agents

NSF is a highly debilitating disease occuring in patients suffering from renal failure.

NSF was first recognized in 1997 in 15 dialyzed patients. This is a recently described

scleroderma characterized by extensive thickening and hardening of the skin associated with

cutaneous papules and coalescing plaques with a ‘peau d’orange’ appearance (Figure

I-18).[41, 42]

a.

b.

Figure I-18: a. Fibrosis on the upper arm, b. fibrosis causing deformities on finger nails.

It causes disabling contractures and damages other internal organs which may

sometimes even lead to death. A causal link between NSF and gadolinium chelates used as

contrast agents for MRI has been proposed, on the basis of retrospective analyses.[43, 44] These

free Gd(III) ions that deposit in the tissue can attract circulating fibrocytes, and eventually

initiates the process of fibrosis. Even though gadolinium is found in the biopsy samples of

suffered patients thus supporting the link between NSF and Gd-CAs,[45, 46] the cause and

mechanism of NSF has not been fully elucidated to have efficient therapeutic measures.

Many hypotheses were drawn out of which dechelation of gadolinium chelates, followed by

the release of free Gd(III) ions inside the body (causing health hazard) is the most accepted

hypothesis so far.

In normal renal function, free gadolinium is removed with a half-life less than 2 hours

whereas improper renal function extends this half-life significantly longer. The vast majority

of published cases were associated with linear and non-ionic gadolinium chelate Gd-DTPA-

BMA (Omniscan), then followed by Gd-DTPA (Magnevist) and Gd-DTPA-BMEA

75

(Optimark) according to the reports of US Food and Drug Administration Medwatch

reporting system.[47] One case was reported with the macrocyclic agent Gd-HP-DO3A

(Prohance). Out of all these clinically available contrast agents, Gd-DOTA (Dotarem) is

known to be the safest MRI contrast agent so far.[48, 49]

Non-ionic preparations are less stable in comparison with the ionic ones as the

binding between Gd(III) with negatively charged carboxyl groups is stronger in comparison

with that of amides and alcohols in non-ionic ones. Macrocyclic chelates form a rigid cage

strongly binding the Gd(III) when compared to the linear ones whose chains are flexible and

do not offer a strong binding to Gd(III). As the improper renal function extends the presence

of Gd chelates significantly longer inside the body, their dissociation half-life also plays a

major role. As longer the dissociation half-life, more safer the gadolinium chelates are from

undergoing transmetallation by endogenous ions such as Cu(II), Zn(II), Ca(II).[50]

Two physicochemical parameters have been introduced to predict the inertness of a

gadolinium chelate: its thermodynamic and kinetic stability.[51]

1. Thermodynamic stability – definition

Thermodynamic stability refers to a thermodynamic equilibrium that exists between

the metal [M], the ligand (often multidentate) [L], [H] ions and the complex [MmLℓHh] (for

simplification the charges are omitted). The global equation is:

(10)

The stability of hm HLM complex is expressed as log hmβ where it is defined as:

(11)

where [M], [L], [H] and ]HLM[ hm are the concentrations of the free metal ion, ligand, H+

ions and the complex at equilibrium respectively.

Ligands are often present in their different protonated forms, so the role of H+ ions

should be taken into account. The equation is given as:

(12)

)1h(m HLM + H hm HLM

76

And the successive stability constant of hm HLM complex is defined as:

(13)

As the ligand and/or their complexes exist in various protonated forms according to the

pH, the overall stability constants are mostly determined by pH-potentiometric titrations. In

this case, the experiments are conducted in solution, when the equilibrium state is reached.

Depending upon the type of the ligand used for complexation, the time that is required for the

complexes to reach equilibrium state varies:

- Usually, in the case of linear ligands the equilibrium state is reached rapidly.

Technique used in such cases is the direct titration, where ligand and metal solutions

are added directly into the cell followed by the titration with the base (‘in-cell

titrations).

- In the case of macrocyclic complexes, the complex formation is very slow. For

example, in DOTA, the complex formation is very slow and the technique employed

is called the “out-of-cell” method.[52,53] In this technique, the overall titration curve is

reconstituted from a series of pH measurements determined in individual flasks, each

flask corresponding to a given mixture of ligand, metal and base. From a flask to the

other, the volume of the base varies, and then the pH. These flasks are kept at 37°C

for 15-30 days. The solution is then titrated against the base in a normal

potentiometric method.

The determination of these formation constants are required to discuss the stability of the

complex in solution. The direct comparision of global constants of different ligands is not

appropriate, because the latter does not take into account the different protonated forms of

free ligands or complexes in solution, as they vary for different ligands. Moreover, in

biological systems such as body fluids, many reactions are possible between the free ligand,

the free metal and also with the complex. General conditional stability constant is calculated

taking into account all the important side reactions possible. On one hand, the free ligand can

interact with endogenous ions such as Mg(II), Ca(II), Zn(II), Cu(II) and Fe(III). On the other

hand, free metal ion can interact with number of biological ligands such as citrate, phosphate,

bicarbonate and oxalic acids. Gadolinium complexes could be in their protonated forms,

which then could form ternary complexes with other small ligands such as carbonate,

77

phosphate and dicarboxylic acids. By considering all the above possible side reactions, and to

have an insight of the complex stability at physiological pH, a conditional thermodynamic

stability constant cond, or Kcond, stemming from can be defined as: [54]

(14)

where [Gd’] = [Gd] + [GdA] + [GdB] +.....

[L’] = [L] + [LH] + [LHn] +.....+ [M’L] + [M’’L] +....

[GdL’] = [GdL] + [GdLH] + [GdLX] + [GdLY] +.....

and M’, M’’ correspond to endogenous metal ions.

A, B correspond to biological competitor ligands.

GdLH represents protonated gadolinium complex

GdLX, GdLY complexes correspond to ternary complexes.

To discuss and compare the stability of complexes in solution, the comparision could

be done either by calculating the formation constant for a given system or by the direct

comparision of these conditional constants, which permits to know the concentration of the

free metal ions present in the solution. The latter should be done only when the same

experimental conditions are practiced, such as pH and total concentration of the metal and

ligand solution. For example, studies comparing the complexation of Fe(III) in the mimicked

living environment using transferrin are performed at pH = 7.4, a total iron concentration of

10-3 M and concentration of ligand 10 times more the concentration of metal,[55] while for

metals other than iron, the reference conditions used are: pH=7.4; CM=1×10-5 mol.L-1,

CL=2×10-5 mol.L-1.[56] An alternative is also possible to compare the stability of different

systems by plotting the logarithm of the free metal ratio versus pH for the systems.[57]

Gd contrast agent structure and thermodynamic relationships

The stability of gadolinium contrast agents chelates is derived from the electrostatic

interactions of gadolinium ion with the donor groups of the chelating ligand. Chemical bonds

are predominantly ionic in gadolinium chelates. Three main structural factors that are said to

be influencing the thermodynamic stability of these chelates are:

1) the basicity of the polyaza-carboxylate scaffold,

78

2) the number of five membered rings formed by the chelate between metal and

various donor atoms of the ligand,

3) the macrocyclic effect that is related to the cavity size, rigidity and the

conformation of the ligand.

The stability constants of the current Gd chelates are reported in Table I-6.[58] In this

context, conditional stability constants can be correlated to their ionic or non-ionic chemical

structure.

Table I-6: Stability constants of commercially available contrast agents.[58]

Gd-CAs Thermodynamic stability constant

Gd-DOTA 25.8

Gd-HP-DO3A 23.8

Gd-BT-DO3A 21.8

Gd-EOB-DTPA 23.5

Gd-BOPTA 22.6

Gd-DTPA 22.1

MS-325 22.06

Gd-DTPA-BMA 16.9

Gd-DTPA-BMEA 16.6

The main differences can be explained by the ionic/ non-ioninc character of the

chemicals bonds between the metal and the ligand and by the macrocyclic/ linear character of

the complexes.

From the chemical bond nature point of view, anionic oxygen atoms in carboxylate

substituent are more basic than neutral oxygen atoms in alcohol or amide substituents.

Anionic oxygen atoms of the carboxylic substituent form a stronger bond with the metal

compared to the metal-alcoholic interaction in other Gd complexes. Therefore, electrostatic

attractions would drastically decrease if the ionic pendant arms are replaced by the non-ionic

pendant arms and carboxylate substituents would form stronger bonds with the metal

compared to the metal-alcoholic interaction. This is illustrated by the Gd-DOTA stability

constant (25.8) which is 2-4 orders of magnitude higher than those of Gd-HP-DO3A (23.8)

and Gd-BT-DO3A (21.8). The same trend is observed in linear complexes between ionic

(Gd-DTPA (22.1), Gd-EOB-DTPA (23.5), Gd-BOPTA (22.6), MS325 (22.06)) and non-ionic

ones (Gd-DTPA-BMA (16.9), Gd-DTPA-BMEA (16.6)). From the ligand preorganization

79

point of view, even if with linear ligands the chelate rings are five-membered ring chelates,

the global macrocylic structure is lacking. This structural point is translated into the values of

the stability constants since stability constants of macrocylic complexes are 2-4 orders of

magnitude higher than those of linear complexes.

2. Kinetic inertness

Thermodynamic stability constants measured in water are neither neccessary nor

sufficient to have an outlook of in vivo stability.[59]

Kinetic stability of a complex is defined as the relative ability of this complex to stay

stable in the presence of various endogenous cations and anions. Therefore, kinetic stability is

also one of the important physico-chemical parameters to understand the relative in vivo

dissociation. This is particularly important in the context of Gd chelates and NSF in which

injected Gd chelate should remain chemically inert during it span inside the body. Exchange

reactions of Gd-complex with endogenous cations such as Cu(II) and Zn(II) ions are studied

by following the absorbance[60] using UV-Visible spectroscopy and relaxometry[61]

respectively. Exchange reactions of Gd-complex with endogenous anions are studied by

following their longitudinal relaxation rates using relaxometry.

Dissociation of a gadolinium complex can occur via different pathways (Scheme I-1),

which include spontaneous dissociation (0), proton assisted dissociation (I and II), ligand

assisted dissociation (III) and metal-ion catalyzed dissociation (IV).[62]

Scheme I-1: Dissociation mechanisms of Gd(III) complexes.[62]

Spontaneous dissociation (0) is explained by the reaction in which the dissociation

occurs individually without any other interferent.

80

Proton-assisted dissociation (I) and (II) includes the protonated species GdLHn

contribution in the complex dissociation (mono and diprotonated complexes, KGdHL and

KGdH2L).

Ligand–assisted dissociation (III) occurs by the involvement of an endogenous ligand

which triggers the Gd complex dissociation. It implies the formation of an intermediate

LGdL* where L* is the endogenous ligand, its corresponding stability constant being KLGdL*.

Metal-ion-catalyzed dissociation (IV) is characterized by the formation of

heterodinuclear complexes with endogenous metal ions (Mg(II), Ca(II), Zn(II), Cu(II),

Fe(III)) eventually releasing Gd(III) as a free metal ion. These metals were presumed to be

the most abundant metals in plasma and best understood in their biological roles.[63] The

stability constant of the corresponding dinuclear complex is given by KGdLM.

Kinetic inertness of a complex can be studied either by relaxometric measurements

using stoichiometric conditions (Zn(II) as competitive metal ion) or by UV-spectroscopy

(Eu(III) and/or Cu(II) as competitive metal ions) in which the reaction is reduced to pseudo

first order by using excess of competitive metal ion.

a) Stoichiometric conditions

Transmetallation of gadolinium complexes can be assessed through the evolution of

the paramagnetic longitudinal relaxation rate of water protons (R1). These experiments are

raised in stoichiometric conditions at 37°C in pH = 7 phosphate-buffered solutions.

Gadolinium complex and zinc chloride are mixed in equal ratios. When the

transmetallation of a gadolinium complex by diamagnetic Zn(II) ions occurs in such a

buffered solution, the released Gd(III) ions react with PO43- ions to form GdPO4 (Scheme I-

2). Consequently a decrease in the proton relaxation rate is observed.

GdLn- L(n+3)- + Gd3+

L(n+3)- + Zn2+ ZnL(n+1)-

Gd3+ + PO43- Gd(PO4) insoluble

(15)

(16)

(17)

Scheme I-2: Transmetallation of Gd(III) complexes assisted by Zn(II)

GdPO4 salt has very low solubility. So the influence of this salt on the longitudinal

relaxation rate of water is negligible. Basing on the evolution of R1, the extent of

81

transmetallation process can be estimated.

The measurements were performed with both commercially available macrocyclic

and linear contrast agents (Figure I-19).

Figure I-19: Evolution of R1(t) / R1(t = 0) vs time.[64]

For macrocyclic complexes Gd-DOTA and Gd-HP-DO3A, no evolutions are

observed during the experiment which indicated that no transmetallation occurred, illustrating

the very high kinetic inertia of these complexes. At the end of the observation period, more

than 98% of the paramagnetic relaxation rate is retained for these macrocyclic complexes.

For linear complexes Gd-DTPA, Gd-DTPA-BMA and Gd-EOB-DTPA, a marked

difference is seen in the evolution of the normalized paramagnetic longitudinal relaxation

rates for the early times of the experiment. A decrease in the rate is observed, with a steep

slope for Gd-DTPA-BMA which retains only 50% of its initial relaxation rate after 200

minutes. Therefore, all these linear complexes show more extensive transmetallation than the

macrocyclic ones. At the end of the observation period, less than 70% was seen for these

linear complexes.

A theoritical description has been attempted to account for these experimental data.

For that, a ‘kinetic index’ and a ‘thermodynamic index’ have been arbitrarily defined.[65]

Kinetic index is defined by the time required to reach 80% of the initial R1 value and the

thermodynamic index by the ratio of R1(t) / R1(t=0) measured after 3 days of starting the

experiment. The values obtained for this experiment were given in Table I-7.

82

Table I-7: Time required to reach [R1(t) / R1(t = 0)] = 0.80 and ratio R1(t=4320) / R1(t=0)

for gadolinium complexes

Complexes Kinetic Index : time

(min) required to reach

R1 (t) / R1 (t = 0) = 0.80

Thermodynamic Index : R1

(t=4320) / R1 (t = 0)

Gd-DOTA >5000 0.99

Gd-HP-DO3A >5000 0.99

Gd-(S)EOB-DTPA 1500 0.69

Gd-DTPA 250 0.49

Gd-DTPA-BMA 70 0.10

All these data illustrate the high kinetic and thermodynamic stabilities reported for

macrocyclic gadolinium chelates.

Relaxometric measurements using this protocol requiring a simple low-resolution

NMR system allows the study of the transmetallation process of the gadolinium complexes of

our interest.

b) Pseudo first order conditions

Kinetic measurements of the gadolinium complex (GdL) transmetallation in the

presence of endogenous cations (M) are also studied by UV- spectroscopy. Let’s consider the

reaction:

GdL + M Gd + ML for which the reaction rate is: k[GdL][M]r (18)

The reaction rate depends on both the concentration of Gd complex and endogenous

metal ion concentration. Measuring a second-order reaction rate with two reactants can be

problematic: either the concentrations of two reactants must be followed simultaneously,

which is more difficult, or the concentration of one of them must be measured and the other

calculated as a difference, which is less precise.

Hence, a common solution for that problem is the pseudo first order approximation.

To reduce the reaction order to pseudo first order, the competitive metal ion is added in

excess. Therefore, one of the reactants remains constant and the expression of the rate of the

reaction becomes:

[GdL]k' k[GdL][M] r (19)

83

Whereas the concentration of excess of reactant is absorbed within the rate constant obtaining

a pseudo first order rate constant according to:

k[M]k'

where k[M]k' (k' or kobs)

(20)

Hence, the order degeneracy leads then to a rate constant which is only dependent on

the total concentration of the complex:

[GdL]tkobsdtd[GdL]t

where kobs is the pseudo first order rate constant.

(21)

Taking into account all possible dissociation pathways described in Scheme 1, the

total concentration of the complex [GdL]t can be expressed as:

GdL][LGdLM][L][GdH[GdHL][GdL][GdL]t*

2 (22)

Experimentally, the rates of exchange reactions were studied at different

concentrations of exchanging metal ions and at various pH values. When the analytical

technique is UV-spectroscopy, the pseudo-first-order rate constants of each kinetics are

calculated by fitting the absorbance data with the use of equation:

tk

e0etobse)AA(AA

(23)

where At, A0 and Ae are the absorbance at time t, at the start and at equilibrium of the

reactions, respectively.

The determination of kobs under different experimental conditions (pH, concentration

of endogenous metal M) allows then to have an insight of the mechanisms of transmetallation

for a given Gd complex and then of their in vivo stability.[51] It is found that the macrocyclic

and linear Gd chelates differ not only in the rate of dissociation, but also in the mechanism of

the reactions that occur.[51,66]

84

i. Kinetic inertness of linear gadolinium chelates

The prototype reaction is the dissociation of GdDTPA, which is probably the most

widely used MRI contrast agent. The evaluation of its kinetic inertness can be characterized

by rate constants of exchange reactions that take place in the plasma. The determination of

the in vivo rate of dissociation reactions of Gd complexes in fluids that mimic the plasma

would be difficult. The results of in vitro studies may however provide important information

and from these results, some extrapolation could be done to predict the kinetic behaviour of

these complexes under in vivo conditions.

The main possibility is the dissociation of Gd(III) ion from its complex by the endogenous

metal cations M in a metal exchange reaction according to the already mentioned reaction: [67]

GdL + M Gd + ML (18)

These metal-exchange reactions can occur via associative or dissociative mechanisms,

like substitution reactions.

On the one hand, exchange can take place by direct attack of exchanging metal ion on

the complex forming a dinuclear intermediate (GdLM), followed by the release of Gd(III) ion

according to:

GdL + M Gd + MLGdLM (24)

Thus, competitive metal ions M, can attack the gadolinium complex forming a hetero

dinuclear complex:

GdL + M GdLMΚGdLM

where [GdL][M]

GdLMGdLM Κ is the stability constant of the hetero dinuclear

complex3

(25)

and then slowly dissociate gadolinium ion from its complex:

GdL M Gd + MLkGdLM

where kGdLM is the rate constant of metal ion catalyzed dissociation reaction

(26)

3 KGdLM can be determined by pH titrations and used for kinetic calculations.

85

The other possibility can be the attack of the exchanging metal ions onto the

protonated gadolinium complex at lower pH ranges, where the weakly coordinated functional

groups to the gadolinium ion can slowly transfer from Gd(III) to the exchanging metal ion M:

GdHL + M Gd + ML + HkH

GdLM

and

GdHL + M GdLM + H Gd + MLkGdLMkM

GdHL

where kH

GdLM is the rate constant of the proton-metal-assisted reaction,

kMGdHL is the metal assisted dissociation of protonated gadolinium complex.

(27)

(28)

On the other hand, exchange can take place slowly by any spontaneous (S) or proton-

assisted dissociation (PA) of the gadolinium complex. In a second step, rapid reaction occurs

between the free ligand or a protonated species (Scheme I-3).[68]

Gd L

Gd H L

Gd + L

ML + Gd

S

PA M

M

H+

Scheme I-3: Spontaneous (S) or proton-assisted dissociation (PA) mechanisms.[68]

For the spontaneous dissociation reaction, the reaction is:

GdL Gd + Lk0

where k0 is the rate constant of spontaneous dissociation reaction.

(29)

For the proton-assisted dissociation mechanisms, the protons can help in the

dissociation of gadolinium ion from its complex according to:

GdHL Gd + LHkGdHL

(30)

GdHL + H+ Gd + LH2

kHGdHL

(31)

where kGdHL and kHGdHL are rate constants of protonated dissociation reactions

KGdLH being the protonation constant of the protonated complex GdLH4 4 KGdLH can be determined by pH titrations and used for kinetic calculations.

86

By taking into account all these pathways, the pseudo first order rate equation is given as:

[GdHL][M]k[GdLM]k]H[GdHL][k[GdHL]kdtd[GdL] M

GdHLGdLMHGdHLGdHL

t

(32)

Taking into account the total concentration of the complex:

[GdLM][GdHL][GdL][GdL]t

and stability constants of protonated and dinuclear complexes,

the rate constant kobs is expressed as:

(33)

which can be written as:

[M]]H[1

]H[M][k[M]k]H[k]H[kkGdLMGdLH

M4

M3

221

obsKK

(34)

where k1, k2 are the rate constants for proton assisted dissociation, M3k is the rate constant for

metal proton assisted dissociation and kM4 for metal assisted dissociation.

The kobs values obtained by fitting experimental data to this equation allows the

determination of the rate constants k1, k2, M3k and kM

4 respectively, together with the

thermodynamic constant KGdLM (when they are not experimentally determined).

It is said that the mechanisms differ according to the exchanging metal ions used in these

exchange reactions.[67] An evolution in the design of the experiments can be found since the

pioneering works used experimental conditions which are far from being physiological (for

instance with exchanging cations such as Eu(III)and Cu(II) for the nature of the ion or its

concentration by comparison with biological conditions). More recent studies aimed to mimic

physiological conditions as best as possible by performing experiments at physiological pH in

the presence of exchanging ligands to envisage the possible role of small endogenous ligands

in the Gd contrast agents demetallation.

MH1

Mk MHkHkHkkMGdLH

MGdLMGdHLMGdHL

2GdHL

HGdHLGdHLGdHL

obs KKKKKK

87

i.a. Exchange reactions with competitive metal cations.

In this section, exchange reactions with Eu(III) and Cu(II) are presented.

i.a-1. Exchange reactions with Europium.

Exchange reactions of Gd-DTPA with Eu(III) ions were studied in the presence of

excess of Eu(III).[67] The rate constant kobs is found to be increased with increasing [H+]

concentrations, particularly at low Eu(III) concentrations, whereas at higher concentrations of

Eu(III) there is a gradual decrease in the rate constant.

Figure I-20: Plots of log kobs versus Eu(III) concentration for the reaction between Gd-DTPA

and Eu(III). Concentration of [Gd-DTPA] = 5×10-4M; pH = 3.67 (○), 3.86 (×), 4.30 (◊),

4.33 (∆), 4.75 (■), 5.08 (▲), 5.38 (*). (25 °C, 1.0M KCl).

The rate of the reaction is directly proportional to [H+], which indicates the exchange

of the metal ions by proton-independent and proton-assisted pathways, presumably by the

formation and dissociation of mono or diprotonated complexes.

Moreover, kobs increase with an increase in the concentration of Eu(III), which

indicates the characteristic of direct attack of Eu(III) on the complex, by the formation of

dinuclear complexes. By taking all the possible reactions into account, the rate of exchange

reaction between Gd-DTPA and Eu(III) (M) is expressed as:

[GdHL][M]k[GdLM]k]H[GdHL][k[GdHL]kdtd[GdL] M

GdHLGdLMHGdHLGdHL

t (35)

88

[M]]H[1

]H[M][k[M]k]H[k]H[kkGdLMGdLH

M4

M3

221

obsKK

(36)

The kobs values obtained by fitting experimental data to this equation, [67] allows the

determination of the rate constants k1, k2, M3k and kM

4 respectively, together with the

thermodynamic constant KGdLM.

This indicates that with Eu(III), exchange can take place:

- via proton-assisted dissociation of GdDTPA (characterized by k1 and k2), followed by

a fast reaction between the free ligand and Eu(III) (because mono- and diprotonated

complexes dissociate much faster then the non-protonated ones)

- by direct attack of Eu(III) on the non-protonated complex GdL ( M3k ) and on the

mono-protonated complex, GdHL ( kM4 ). Once these dinuclear reaction intermediates

are formed, the functional groups of the ligand are then gradually transferred from the

gadolinium to the attacking ion.

Same behaviour was found in the exchange reactions with Eu(III) ion in other linear

Gd chelates, Gd-DTPA-BMA[68-69] and its derivatives. Even though there is a decrease in the

stability constants of Gd-DTPA-BMA and its derivatives compared to Gd-DTPA (because of

the replacement of carboxylic groups with amide groups), no change in the kinetic stability

towards Eu(III) is seen.

i.a-2. Exchange reactions with Copper and Zinc.

Rates of the exchange reactions between Gd-DTPA and Cu(II) or Zn(II) occur much

faster than with Eu(III).[67]

The rate constant and the stability constant of dinuclear complex are calculated using

the equation below:

[M]1

[M]kkkGdLM

M30

obsK

(37)

89

Rates of the reactions increase by increasing the concentration of the exchanging

metal ions and the values are practically pH independent.

Figure I-21: a). Plots of kobs versus Cu(II) concentration for the reaction between Gd-DTPA

and Cu(II). Concentration of [Gd-DTPA] = 5×10-4M; pH = 5.20 (○), 4.92 (∆) (25 °C, 1.0M

KCl); pH 5.21 (◊) (37 °C, 1.0M KCl).

b). Plots of kobs versus Zn(II) concentration for the reaction between Gd-DTPA and Zn(II).

Concentration of [Gd-DTPA] = 5×10-4 M; pH 4.62 (○), 4.80 (×), 5.51 (◊) (25 °C, 1.0M KCl);

pH 5.21 (∆) (37 °C, 1.0M KCl).

This indicate that exchange reactions occur via the direct attack of the Cu(II) or Zn(II)

ions on the complex, forming an hetero-dinuclear complex. The proton assisted dissociation

mechanism is slow and is considered as negligible in the range of investigated pH.

Various DTPA derivatives have been studied on the basis of this methodology which

allows classifying them according to their rate of dissociation in the presence of Cu(II)or

Zn(II).[70]

The exchange reactions of the DTPA derived Gd complexes with Cu(II) and Zn(II)

show the same pattern of transmetallation, where the direct attack of the Cu(II) or Zn(II) ions

on the complex lead to the formation of an hetero dinuclear complex. kobs obtained were

fitted (Figure I-22) using the equation below:

]M[1]M[kk

GdLM

M3

obsK

(38)

90

Figure I-22: a). Plots of kobs versus Cu(II) concentration for the reaction between Gd-

DTPA-BMA (▲ and +), Gd-DTPA-N’-MA (∆ and ◊) and Gd-DTPA-N-MA (○ and ×) and

Cu(II). pH = 4.91 and 5.211 (25°C, 1.0M KCl).

b). Plots of kobs versus Zn(II) concentration for the reaction between Gd-DTPA-BMA (▲ and

+), Gd-DTPA-N’-MA (▲) and Gd-DTPA-N-MA (•) and Zn(II). pH = 5.21(25°C, 1.0M KCl).

i.b. Exchange reactions with endogenous anions - Ligand assisted dissociation studies

The ligand assisted dissociations are also possible inside the body, due to the presence

of endogenous anions (L*) such as citrates (cit), phosphates, hydrogeno-carbonates,

aminoacids or proteins. These exchanging substrates can trigger the dissociation of GdL,

thereby complexing gadolinium and leaving the organic chelate free according to:

GdL + L* L*GdL L + GdL* (39)

The small anions can form ternary complexes with Gd-complexes by replacing the

water molecule forming a L*GdL species, where L*= CO32-, cit3-.

Exchange reactions were performed between linear complexes (Gd-DTPA, GdDTPA-

BMA) and citrate, phosphate and carbonate ions.[71] In these cases, 1H relaxation rates of

ternary complexes formed during the reaction were determined according to the pH, in order

to determine the influence of these complexes and the pH on the mechanism.

91

Figure I-23: The extent of ternary complex formation in the presence of (1) 25 mM carbonate, and

(2) 0.5 mM phosphate ligands. The concentrations of Gd-DTPA and Gd-DTPA-BMA were each 0.2

mM (25 °C, 0.1 M NaCI).

The ternary complex formation is seen only at pH greater than 8. So, the relaxation

effect of Gd-complexes is not influenced by the ternary complex formation at physiological

pH.

Recently, a more complete study was performed to understand the mechanisms of the

ligand exchange reactions inbetween linear Gd contrast agents GdL, in the presence of an

exchanging ligand (TriethyleneTetraamine-HexaAcetate, TTHA as a model) and small

endogenous anions in a stepwise methodology.[72]

First, in order to obtain the rate law for the ligand exchange reactions, the pseudo-first-order

rate constants have been determined at different TTHA and hydrogen ion concentrations. All

the protonated foms of the exchanging ligand were considered according to:

GdL + HiTTHA(6-i)- Gd(TTHA)3- +HjL + (i-j)H+

(40)

The rates of the reactions are directly proportional to the concentration of TTHA,

indicating that the reactions take place with the direct attack of the various protonated

HiTTHA species (0 ≤ i ≤ 3) on the Gd(III) complexes, through the formation of ternary

intermediates. The results also indicated that the less protonated HiTTHA species can more

efficiently attack the Gd(III) complex, improving the formation of the ternary complex.

92

Therefore in a second step, the influence of small endogenous ligands was evaluated.

The exchange reactions take place more quickly in the presence of endogenous citrate,

phosphate and carbonate ions at a pH of 7.4. The increase in the reaction rates may be the

result of the formation of low stability ternary complexes between the endogenous ligands

and the GdL species (mostly with CO32- ions since the concentration of the carbonate in the

blood plasma is larger (25 mM), which accelerates the intramolecular rearrangements and, so,

the dissociation of the complexes).

Dissociation kinetics of linear gadolinium complexes (Gd-DTPA, Gd-DTPA-BMA)

were studied with Cu(II) in the presence of endogenous ligands such as citrate, phosphate,

carbonate and histidinate ligands under near to the physiological pH at 25°C in NaCl to

investigate the catalytic effect of endogenous ligands.[73]

GdL + Cu(X) Gd(X) + CuL where X = citrate, carbonate, phosphate ions.

(41)

Dissociation of Gd-DTPA and Gd-DTPA-BMA is assisted mainly by the carbonate,

citrate and phosphate ions, where free aminoacids showed no role in the dissociation of

gadolinium complexes. Considering all the possible reaction pathways, the reaction rate is

expressed as:

])[[][][][][][(][

]42323

231 4233

GdLPOHkHCOkCOkHCitkCitkHkdt

GdLPOHHCOCOHCitCit

t

From the reaction pathway consisdered above, an equation was derived, where the

kobs was fitted giving the rate constants of Gd-DTPA and Gd-DTPA-BMA. Dissociation of

Gd-DTPA-BMA in the presence of citrate, carbonate and phosphate ions occur much faster

than the dissociation of Gd-DTPA.

ii. Kinetic inertness of macrocyclic complexes

Macrocyclic ligand is generally characterized by a rigid cavity that can enclose the

metal ion and stay inert. The clinically available macrocyclic contrast agents are Gd-DOTA,

Gd-BT-DO3A, Gd-HP-DO3A. Gd-DOTA is considered as the highly stable complex so far,

amongst all the clinically available contrast agents. At physiological pH the dissociation is

even lower or unrealizable when compared to the acidic medium where the dissociation is

93

rapid.[74] Nevertheless, to have an insight about the reaction, these experiments were raised at

excess of [H+] concentration, since the complexes are unstable under pH = 2.

Dissociation kinetics (metal exchange reactions) of Gd-DOTA were studied with

Eu(III).[75] Reactions were studied in the pH range 3.2-5.0, and the concentration of Gd-

DOTA (0.1M) was kept high compared to the Eu(III) (aq) (0.01M) solution to have

measurable rates. Under these conditions the reaction follows a pseudo-first order reaction.

The exchange rate is considered proportional to Gd-DOTA concentration and given by:

[GdDOTA]kdtd[GdDOTA]

obs

where kobs is a pseudo first order rate constant

(42)

.

The kobs values obtained do not differ according to the varying concentrations of Eu3+

indicating that the exchange rates are independent of Eu(III) concentration and are dependent

on [H+] concentration. The expression of kobs is given by:

]H[kkk 10obs (43)

where the rate constants k0 and k1 are characteristic of the spontaneous and proton-assisted

dissociation of Gd-DOTA.

To neglect the spontaneous dissociation process, dissociation studies were carried out

in acidic solutions (0.02-0.23 mol.L-1 HCl).[74, 75] By fitting the experimental data to the

previous equation (where k0 was fixed to zero), the proton assisted rate constant k1 was

determined. The linear dependence of kobs on [H+] concentration indicates the proton

catalyzed dissociation of the GdL complex (L = DOTA, DO3A derivative) (Scheme I-4).

Scheme I-4: Proton-assisted dissociative mechanism.[75]

Upon stepwise protonation, several Gd species can be involved in the dissociation

mechanism. The rate constants are conditioned by the respective thermodynamic stability and

kinetic inertness of each species. In the mechanism, the weakest complex is GdH2L because

94

of intrinsic destabilizing electrostatic repulsions spreading between two protonated nitrogen

atoms of the macrocyclic cavity and the partially demetallated Gd(III) ion.

From the above reaction (Scheme 4), the rate of dissociation can be given as:

][GdLkGdL]H[k[GdHL]kdt]d[GdL

tobs2LHGdGdHLt

2

(44)

Since: GdL]H[[HGdL][GdL]][GdL 2t

and: [HGdL][H]GdL][H;[GdL][H]

[HGdL] 2LHGdHGdL 2 ΚK

The pseudo-first order rate constant, kobs, which describes the relationship between the

dissociation rates and [H+] in a broad range of concentrations, is then given by:

][HKK[H]K1][HKkk[H]kk

k 2GdLHHGdLHGdL

2GdLHHGdLLHGdHGdLGdHL

obs

2

22

(45)

The dissociation of Gd-HP-DO3A was studied in similar conditions,[74] where

monoprotonated complex predominates. The linear dependence of the dissociation rate on

proton concentration was interpreted by proton-assisted dissociation of the monoprotonated

species, which is also the dissociation of diprotonated complexes.

All these calculations indicate that the pathway involving monoprotonated species is

considered important for dissociation near physiological pH values of macrocyclic Gd

chelates.[74, 76]

F. Scope of the thesis

Many advances were done in the last years improving contrast agents for imaging.

Design of CAs combining both the optical and MRI imaging, introduction of nanoparticles

for both diagnosis and therapy etc. are nowadays gaining importance because of their

multimodal abilities. Since contrast agents are used in MRI, numerous efforts have been done

to improve these agents to minimize the toxic effects of Gd(III) ions. No efficient therapy for

NSF could be proposed due to the unknown mechanisms of the Gd chelates when injected

95

into the body. So, study of the physico-chemical parameters became primary and an

important step before introducing any Gd chelates as contrast agents.

In the first part of my thesis, the physico-chemical parameters of newly designed

ligands which could be used for both optical and MRI imaging are studied (Figure I-24).

These macrocyclic ligands are designed in such a way to ensure the stability and relaxivity

when complexed with Gd(III) ion and luminescence when complexed with Eu(III) ion. The

ligands are synthesized by Prof. Stephen J ARCHIBALD et al. (University of Hull, United

Kingdom).

N N

N N

CO2HHO2C

HO2C

N

N

R

R = H, NO2

Figure I-24: Benzimidazole DO3A chelate (R = H, p-nitrophenyl group)

In the second part of my thesis, the physicochemical properties of a linear ligand

(DTDTPA) designed for grafting onto gold nanoparticles and finally incorporating Gd(III)

will be studied (Figure I-25). These gold nanoparticles are synthesized by Prof. Stéphane

ROUX et al. (Université de Franche-Comté, France).[77, 78]

a.

N N N

COOHHOOC

O

NH

O

HNHS SH

COOH

b.

Figure I-25: a) DTDTPA and b) multilayered GdDTDTPA-loaded AuNPs. [77]

The challenge will be to study if when complexed with lanthanide ions, the resulting

complexes will be thermodynamically stable and chemically inert.

96

References:

1. T. Moeller, J. Chem. Educ. 1970, 47, 417-423.

2. G. Liu, B. Jacquier, Spectroscopic properties of rare earth in optical materials,

Tsinghua University Press & Springer: Beijing & Heidelberg.

3. J. C. Bunzli, J. C. Piguet, Chem. Rev. 2002, 102, 1897-1928.

4. E. N. Rizkalla, G. R. Choppin, in Handbook on the Physics and Chemistry of Rare

Earths, K. A., Jr. Gschneidner, L. Eyring, Eds., Elsevier Science Publ. B.V.:

Amsterdam, 1991.

5. S. A. Cotton in Lanthanide and Actinide Chemistry, John Wiley & Sons Ltd, The

Atrium: Southern Gate, Chichester, 2006.

6. J. C. Bunzli, S. V. Eliseeva, Basics of Lanthanide Photophysics in Lanthanide

Luminescence: Photophysical, Analytical and Biological Aspects.

7. P. Hanninen, H. Harma Eds., Springer Ser Fluoresc, 2010, Springer-Verlag: Berlin

Heidelberg.

8. J. C. Bunzli, S. V. Eliseeva, Chem. Sci. 2013, 4, 1939-1949.

9. V. Yam, K. Lo, Coord. Chem. Rev. 1999, 184, 157–240.

10. P. David, J. A. Gareth Williams, J. Chem. Soc. Dalton Trans. 1996, 3613–3628.

11. J. H.Turner, P. G. Claringbold, E. L. Hetherington, P. Sorby, A. A. Martindale,

Journal of Clinical Oncology. 1989, 7, 1926–1931.

12. L. E. Hegg, N. B. Judith, Coord. Chem. Rev. 1998, 173, 133–165.

13. H. Schneider, J. Rammo, R. Hettich, Angew. Chem. Int. Ed. Engl. 1993, 32, 1716–

1719.

14. N.Sabbatini, M. Guardigli, J. M. Lehn, Coord. Chem. Rev. 1993, 123, 201–228.

15. J. C. Bunzli, J. C. Piguet, Chem. Soc. Rev. 2005, 34, 1048–1077.

16. H. H. Dam, N. R. David, W. Verboom, Chem. Soc. Rev. 2007, 36, 367–377.

17. R. Wietzke, M. Mazzanti, J. M. Latour, J. Pécaut, P. Y. Cordier, C. Madic, Inorg.

Chem. 1998, 37, 6690–6697.

18. D. Parker, J. A. Gareth Williams, Metal Ions in Biological Systems. 2003, 40, 233.

19. T. Gunnlaugsson, A. M. D. Dónall, D. Parker, J. Am. Chem. Soc. 2001, 123, 12866–

12876.

20. T. Gunnlaugsson, P. L. Joseph, K. Sénéchal, J. H. Andrew, Chem. Commun. 2004,

782–783.

97

21. O. Reany, T. Gunnlaugsson, D. Parker, J. Chem. Soc. Perkin Trans 2. 2000, 1819–

1831.

22. A.P.De Silva, H.Q.N. Gunaratne, T.E. Rice, S. Stewart, Chem. Commun. 1997, 1891–

1892.

23. C. Li, W. T. Wong, Chem. Commun. 2002, 2034–2035.

24. M.P. Lowe, D. Parker, O. Reany, S. Aime, M. Botta, G. Castellano, E. Gianolio,

R. Pagliarin, J. Am. Chem. Soc. 2001, 123, 7601–7609.

25. I. Hemmilä, S. Dakubu, M. Veli-Matti, H. Siitari, T. Lövgren. Anal. Biochem. 1984,

137, 335–343.

26. G. Mathis, Clin. Chem. 1995, 41, 1391–1397.

27. D. J. Bornhop, J. M. M. Griffin, S. G.Timothy, R. S. Mark, B. Bell, M. Motamedi,

Appl. Spectrosc. 2003, 57, 1216–1222.

28. P. Caravan, J.J. Ellison, T.J. McMurry, R.B. Lauffer, Chem. Rev. 1999, 99, 2293–

2352.

29. M. F. Tweedle, Lanthanide Probes in Life. In Chemical and Earth Sciences; J. C. G.

Bunzli, G. R. Choppin, Eds.; Elsevier: Amsterdam, 1989, 127.

30. D. W. Darnall, R. B. Edward, Biochemistry. 1973, 12, 3489–3491.

31. M. Port, J. M. Idée, C. Medina, C. Robic, M. Sabatou, C. Corot, BioMetals. 2008, 21,

469–490.

32. N. Bloembergen, E. M. Purcell, R. V. Pound, Phys Rev. 1948, 73, 678.

33. I. Solomon, Phys. Rev. 1955, 99, 559.

34. I. Solomon, N. Bloembergen, J. Chem. Phys. 1957, 23, 261.

35. N. Bloembergen, J. Chem. Phys. 1957, 27, 572.

36. N. Bloembergen, L. O. Morgan, J. Chem. Phys. 1961, 34, 842.

37. R. E. Connick, D. J. Fiat, J. Chem. Phys. 1966, 44, 4103.

38. E. Tóth, L. Burai, A.E. Merbach, Coord. Chem. Rev. 2001, 216–217, 363–382.

39. A.D. Sherry, Z. Shanrong & W. Mark, Med. Inorg. Chem. 2005, 903, 151–165, ACS

Symposium Series, American Chemical Society.

40. J. M. Idée, M. Port, C. Medina, E. Lancelot, E. Fayoux, S. Ballet, C. Corot,

Toxicology. 2008, 248, 77–88.

41. S. E. Cowper, H. S. Robin, M. S. Steven, D. S. Lyndon, S. Gupta, P. E. LeBoit, The

Lancet. 2000, 356, 1000–1001.

42. S. E. Cowper, L. D. Su, J. Bhawan, H. S. Robin, P. E LeBoit, Am. J. Dermatopathol.

98

2001, 23, 383–393.

43. T. Grobner, Nephrol. Dial. Transplant. 2006, 21, 1104–1108.

44. P. Marckmann, S. Lone, K. Rossen, A. Dupont, M. B. Damholt, J. G. Heaf, H. S.

Thomsen, J. Am. Soc. Nephrolog. 2006, 17, 2359–2362.

45. A. S. Boyd, J. A. Zic, J. L. Abraham, J. Am. Acad. Dermatol. 2007, 56, 27–30.

46. W. A. High, R. A. Ayers, S. E. Cowper, J. Am. Acad. Dermatol. 2007, 56, 710–712.

47. E. Kanal, D. R. Broome, D. R. Martin, H. S. Thomsen, Radiology. 2008, 246, 11–14.

48. T. Ishiguchi, S. Takahashi, Drugs in R&D. 2012, 10, 133–145.

49. G. Deray, O. Rouviere, L. Bacigalupo, B. Maes, T. Hannedouche, F. Vrtovsnik, C.

Rigothier, Eur. Radiol. 2013, 36, 387-408.

50. W.P. Cacheris, C.Q. Steven, S.M. Rocklage, Magnetic Resonance Imaging. 1990, 8,

467–481.

51. E. Brucher, A.D. Sherry, 2001. Stability and toxicity of contrast agents. In: Merbach,

A.E., Toth, E. (Eds.), The Chemistry of Contrast Agents in Medical Magnetic

Resonance Imaging. John Wiley & Sons Ltd., Chichester, pp. 243–279.

52. Ė.Tóth, E. Brücher, Inorg. Chim. Acta. 1994, 221, 165–167.

53. E.T. Clarke, A.E. Martell, Inorg. Chim. Acta. 1991, 190, 37–46.

54. J. M. Idée, M. Port, I. Raynal, M. Schaefer, S. L. Greneur, C. Corot, Fundam. Clin.

Pharmacol. 2006, 20, 563–576.

55. J. M.Harrington, G. Winkelmann, K. Haselwandter, A. L. Crumbliss, J. Inorg.

Biochem. 2011, 105, 1670–1674.

56. B. Achour, J. Costa, R. Delgado, E. Garrigues, C. F. G. C. Geraldes, N. Korber, F.

Nepveu, M. I. Prata, Inorg. Chem. 1998, 37, 2729–2740.

57. J. Moreau, E. Guillon, P. Aplincourt, J. C. Pierrard, J. Rimbault, M. Port, M.

Aplincourt, Eur. J. Inorg. Chem. 2003, 3007–3020.

58. J. M. Idée, M. Port, C. Robic, C. Medina, M. Sabatou, C. Corot, J. Magn. Reson.

Imaging. 2009, 30, 1249–1258.

59. M. F. Tweedle, J. J. Hagan, K. Kumar, S. Mantha, C.A. Chang, Magnetic Resonance

Imaging. 1991, 9, 409–415.

60. E. Balogh, R. Tripier, P. Fousková, F. Reviriego, H. Handel, E. Tóth, Dalton Trans.

2007, 32, 3572–3581.

61. R. B. Lauffer, Chem. Rev. 1987, 87, 901-927.

62. G. Tircsó, Z. Kovács, A. D. Sherry, Inorg. Chem. 2006, 45, 9269 – 9280.

99

63. P. M. May, P. W. Linder, D. R. Williams, J. Chem. Soc., Dalton Trans. 1977, 588–

595.

64. S. Laurent, L. Vander Elst, C. Henoumont, R. N. Muller. Contrast Media Mol.

Imaging. 2010, 5, 305–308.

65. S. Laurent, L. Vander Elst, F. Copoix, R.N. Muller, Invest. Radiol. 2001, 36, 115–

122.

66. E. Balogh, R. Tripier, P. Fousková, F. Reviriego, H. Handel, E. Tóth, Dalton

Trans. 2007, 32, 3572–3581.

67. L. Sarka, L. Burai, E. Brücher. Chem. Eur. J. 2000, 6, 719–724.

68. L. Sarka, L. Burai, R. Király, L. Zékány, E. Brücher, J. Inorg. Biochem. 2002, 91,

320–326.

69. L. Sarka, I. Bányai, E. Brücher, R. Király, J. Platzek, B. Radüchel, H. Schmitt-

Willich, J. Chem. Soc. Dalton Trans. 2000, 3699–3703.

70. Z. Jászberényi, I. Bányai, E. Brücher, R. Király, K. Hideg, T. Kálai, Dalton Trans.

2006, 1082–1091.

71. L. Burai, V. Hietapelto, R. Király, É. Tóth, E. Brücher, Magn. Reson. Med. 1997, 38,

146–150.

72. Z. Palinkas, Z. Baranyai, E. Brücher, B. Rózsa, Inorg. Chem. 2011, 50, 3471–3478.

73. Z. Baranyai, Z. Palinkas, F. Uggeri, A. Maiocchi, S. Aime, E. Brücher, Chem. Eur. J.

2012, 18, 16426–16435.

74. K. Kumar, A. Chang, M. F. Tweedle, Inorg. Chem. 1993, 32, 587–593.

75. E. Toth, E. Brucher, I. Lazar, I. Toth, Inorg. Chem. 1994, 33, 4070–4076.

76. X. Wang, T. Jin, V. Comblin, A. Lopez-Mut, E. Merciny, J. F. Desreux., Inorg.

Chem. 1992, 31, 1095–1099.

77. P.J. Debouttière, S. Roux, F. Vocanson, C. Billotey, O. Beuf, A. Favre-Réguillon, Y.

Lin, et al, Adv. Funct. Mater. 2006, 16, 2330–2339.

78. C. Alric, J. Taleb, G. L. Duc, C. Mandon, C. Billotey, A. L. Meur-Herland, T.

Brochard, et al, J. Am. Chem. Soc. 2008, 130, 5908–5915.

101

Chapter - II Physico-chemical studies of

macrocyclic ligands and their metal complexes

103

In the bibliographic introduction, we have indicated that in this second chapter we

will focus on the physicochemical properties of newly designed tetrazamacrocyclic ligands

which could be used for MRI imaging. These ligands5 will be based on a parent backbone

constituted of a DO3A (1,4,7,10-tetraazacyclododecane-1,4,7-triacetic acid) ligand for which

the fourth backbone nitrogen atom will be substituted by benzimidazole derivatives (Figure

II-1).

N N

N N

CO2HHO2C

HO2C

N

N

R

NO2

L1H4

L2H3

HR =

Figure II-1: Benzimidazole DO3A chelate (R = H, L1H4 and R = p-nitrophenyl group, L2H3)

After a rapid survey of azole derivatives interest as substituents for

tetraazamacrocyclic ligands, this chapter will deal with coordination properties of L1H4 and

L2H3 ligands towards first row metal ions (namely Cu(II) and Zn(II)) and lanthanide ions

(namely Gd(III) and Eu(III)). The thermodynamic stability and kinetic inertness of lanthanide

complexes able to act as MRI contrast agents will also be evaluated.

A. Interest of azoles as substituents for DO3A derivatives

Aromatic N-donor ligands such as imidazole, triazole or benzimidazole are interesting

substituents for tetraazamacrocyclic derivatives. Indeed, the N-functionalization of the

macrocyclic backbone by such ligands can offer additional coordination sites for d- and f-

metals and then can improve their coordination. Besides this structural aspect, these ligands

were also used to design luminescent structures, mainly for biological applications. At the

light of some recent examples, we will illustrate this interest.

The first example highlights the role of imidazole as substituents of a cyclen

backbone. Depending on the number of substituent groups on the macrocyclic framework, as

5 These ligands were synthesized by Prof. Stephen J ARCHIBALD et al. (University of Hull, United Kingdom).

104

well as on the requirements of the metal ions employed, the substituted cyclen ligands easily

adapt to coordination numbers six, seven or eight, and thus they might coordinate mono-, di-

and trivalent metal ions.

For the ligand described in Figure II-2, three lanthanide complexes were synthesized

(La(III), Eu(III) and Lu(III)) and their structure established in solution at the light of VT-

NMR investigations.

Figure II-2: 1,7-bis(carboxymethyl)-4,10-bis(1-methylimidazol-2-ylmethyl)-1,4,7,10-

tetraazacyclododecane for the complexation of lanthanides

These results confirm the ligand ability to behave as an octa-coordinating ligand and

for the La(III) complex, allow the identification of two isomers in solution, exchanging via an

enantiomerization process.[1] The results point out that the major and the minor isomers have

a prismatic and an antiprismatic geometry respectively, highlighting that the prismatic

geometry is the most stable for the three [LnL]+ cations. In this regard the present complexes

differ from those with DOTA. Similar geometries were also observed in solution for divalent

ions (Cd(II), Pb(II), Hg(II)).[2] The stabilization of the nearly prismatic structure in the [LnL]+

complexes is likely to be related to the two heterocyclic imidazole rings.

The second example concerns the synthesis by Watkinson et al. of cyclen or cyclam

backbones functionalized by a triazole moiety which bears a fluorophore. The triazole linker

was obtained by a Cu(I)-mediated Huisgen [3+2] ‘click’ cycloaddition between a propargyl-

functionalized cyclen or cyclam ligand and an azide-functionalized naphthalimide

fluorophore. [3,4,5]

105

Figure II-3: a. Cyclen (n = 0) and b. cyclam (n = 1) fluorescent clickates for Zn(II)

complexation

This ‘click’ reaction can also be used to generate di-functionnalised cyclen or cyclam

ligands.[4] The triazole is not only used as a junction in the cyclam derivatives but also as an

additional coordinating ligand. The clickate cyclen or cyclams were used to sense selectively

Zn(II) by fluorescence (Figure II-4).

a. b.

Figure II-4: Fluorescence emission of cyclam clickate a. upon titration with Zn(II) at pH 7

b. in the presence of competitive metals [3]

Their fluorescence emission was also tested in the presence of other biologically relevant

metal ions. It was shown that both cyclen and cyclam ligands have excellent switch-on

selectivity for zinc, with a significant enhancement in fluorescence at physiologically pH.

The following examples concern the synthesis and the photophysical properties of ligands

functionalised with benzimidazole.

Tetraaza macrocycles with functionalized benzimidazole moieties as pendant arms

and their complexes with divalent ions such as Cu(II) and Zn(II) were synthesized to test

their use as chemosensors for analytical purpose based on fluorescence spectroscopy.[6, 7, 8]

106

Figure II-5: a. Cyclen (n = 0) and cyclam (n = 1) mono-N- and di-N-substituted by

benzimidazole groups – Zn(II) complex of cyclen-benzimidazole

Benzimidazole is involved in metal complexation and this involvement plays a role in

the photophysics of the ligand in the presence of metal ions. Without metal, fluorescence of

the ligand is quenched by an intramolecular photoinduced electron transfer between the lone

pairs of the macrocyclic nitrogen atoms and the benzimidazole (PET effect). In the presence

of divalent ions and particularly in the presence of Zn(II) the benzimidazole fluorescence is

restored (Figure II-6).

a. b.

Figure II-6: Fluorescence emission of cyclan benzimidazole a. upon titration with Zn(II) at

pH 10.4 b. in the presence of competitive metals [8]

With increasing amounts of Zn(II) ion, the fluorescence emission increased linearly.

On the other hand, on addition of successive amounts of Cu(II) ion, the fluorescence emission

decreased linearly. Therefore these ligands are selective in their fluorescent response towards

Zn(II) which makes them selective sensors towards this divalent ion.

Finally, benzimidazole was also involved in lanthanide coordination, its role consisting in

allowing an effective energy transfer to the emitting levels of the Ln(III) ions, and on the

107

other hand, shielding the central ion against the solvent in order to avoid non-radiative

deactivation processes.

For example, Pan et al. used tripodal benzimidazole ligands L to form homoleptic

ML2 complexes with lanthanide ions.[9] The alkylation of N-1 position in benzimidazole was

used to tune the singlet and triplet energy levels of the ligand, to influence the ligand antenna

effect and therefore the luminescent properties of Ln(III).

Figure II-7: Tripodal benzimidazole ligands for Ln(III) complexation and some relevant

complexes

Similar studies were carried out with tridentate benzimidazole-substituted pyridine-2-

carboxylates that readily form 9-coordinate neutral homoleptic anhydrous lanthanide

complexes. The ligands sensitise very efficiently the luminescence of europium, in solid

state, in thin film and in dichloromethane.[10]

R’ = CH3 R = H, F, Cl, Br, OCH3C4H9C8H17

R’ = CH3 R = H, F, Cl, Br, OCH3C4H9C8H17

Figure II-8: Tridentate benzimidazole-substituted pyridine-2-carboxylic acids and

sensitization of Eu(III)

The remarkable photophysical properties of the europium complexes result from the

good protection of the metal ion by the ligands from non-radiative deactivation provided by

the N6O3 coordination environment. Thus, the authors concluded that benzimidazole-

substituted pyridine-2-carboxylates are promising building blocks for the design of

luminescent materials.

108

To sum up, the previous examples have shown that azole ligands are useful to act as

donors for metal complexation. They can be easily introduced in pre-existant ligands

(tetraazamacrocycles for instance) and then they allowed to adapt macrocyclic ligands to

coordination numbers superior to four. In this chapter, the design of the studied ligands rely

upon these possibilities. As previously said at the beginning of the chapter, these ligands will

be based on a parent backbone constituted of a DO3A (1,4,7,10-tetraazacyclododecane-1,4,7-

triacetic acid) ligand for which the fourth backbone nitrogen atom will be substituted by

benzimidazole derivatives.

B. Synthesis and physicochemical studies of ligands L1H4 and L2H3

1. Synthesis of methyl benzimidazole-DO3A (L1H4) and p-nitrophenyl

substituted benzimidazole-DO3A (L2H3)

a) Synthesis of L1H4 and L2H3

Methylbenzimidazole-DO3A consists of a DOTA parent molecule in which one

carboxylic arm is replaced by a benzimidazole group. Therefore, the corresponding ligand

possesses four ionisable protons in its neutral form (Scheme II-1, protons indicated in

red) which leads to write the neutral ligand as L1H4. L1H4 can be synthesized by two ways.

First way is followed by directly utilising tris-tert-butylacetate cyclen derivative

(synthesized by following a reported procedure [11,12]) followed by the introduction of a

single methylbenzimidazole substituent to produce DO3A(t-Bu)3 benzimidazole (Scheme

II-1).

N N

N N

NH

NOtBu

tBuO

tBuO

O

OO

N N

N N

HN

NOH

HO

HO

O

OO

N HN

N N

OtBu

tBuO

tBuO

O

OO

HN

NCl

L1H4A B

MeCN, Cs2CO3

HCl

6 mol. L-1, n HCl

Scheme II-1: Synthesis of DO3A-benzimidazole in the form of its hydrochloride salt

In A, the macrocyclic secondary amine function was reacted with 2-

chloromethylbenzimidazole to give 1,4,7-tris(tert-Butoxycarboxymethyl),10-(1H-

benzimidazole)-1,4,7,10 tetraazacyclododecane B as a light brown solid. Deprotection of

109

the ester groups of B was carried out in HCl to give DO3A-benzimidazole L1H4 as an

hydrochloride salt.

An alternative procedure utilises a bis-aminal cyclen derivative, which has been

used previously to produce cyclen benzimidazole L3H.[13] A condensation reaction

between cyclen and glyoxal leads to the formation of an azamacrocyclic bis-aminal C.[14]

The alkylation of the bis-aminal in THF with 2-chloromethyl benzimidazole, leads to the

quaternized methyl benzimidazole substituted bis-aminal cyclen D.[13,15,16] The aminal

bridge is then removed by hydrazinolysis to give cyclen methylbenzimidazole E.

N N

N N

N N

N N

HN

N

NH N

NH HN

HN

N

N N

N N

HN

NOtBu

tBuO

tBuO

O

OO

N N

N N

HN

NOH

HO

HO

O

OO

NH

NCl

anhydrous THF

H2N-NH2,H2O

BrO

O

HCl

EtOH

5 mol. L-1

L1H4

C D

E

B Scheme II-2: Synthesis of DO3A-benzimidazole in the form of its hydrochloride salt

This compound was further alkylated by reaction with tert-butylbromoacetate to

form 1, 4, 7-tris (tert-Butoxycarboxymethyl), 10-(1H methyl benzimidazole)-1,4,7,10

tetraazacyclododecane B whose deprotection by HCl gives DO3A-methylbenzimidazole

L1H4 as an hydrochloride salt. (Scheme II-2)

p-Nitrophenylmethylbenzimidazole-DO3A (L2H3) consists of a DOTA parent

molecule in which benzimidazole is replaced by a p-nitrophenylmethylbenzimidazole

group. Therefore, the corresponding ligand possesses three ionisable protons in its neutral

form (Scheme II-3, protons indicated in red) which leads to write the neutral ligand as

L2H3.

L2H3 was synthesized by directly utilising tris-tert-butylacetate cyclen derivative

(1,4,7-Tris(tert-butoxycarboxymethyl)-1,4,7,10-tetraazadodecane[11,12]) followed by the

110

introduction of a single 1-(4-nitrobenzyl)-2-chloromethyl benzimidazole [17] to produce

DO3A(t-Bu)3 benzimidazole (Scheme II-3). Deprotection of the ester groups of 1,4,7-

tris(tert-butoxycarbonylmethyl)-10-(1-(4-nitrobenzyl)-2-methylbenzimidazole)-1,4,7,10-

tetraazacyclododecane was carried out in HCl to give DO3A- p

nitrophenylmethylbenzimidazole L2H3 as a hydrochloride salt.

N N

N NN

NOtBu

tBuO

tBuO

O

OO

N HN

N N

OtBu

tBuO

tBuO

O

OO

L2H3A

MeCN, Cs2CO3

HCl

6 mol.L-1

, n HCl

N

N

Cl

NO2 NO2

N N

N NN

NOH

HO

HO

O

OO

NO2

B Scheme II-3: Synthesis of DO3A-p-nitrophenylmethylbenzimidazole in the form of its

hydrochloride salt

b) Determination of the empirical formula of ligands L1H4 and L2H3

As the deprotection of ester groups in the last step of synthesis was carried out in HCl,

ligands obtained are in the form of their hydrochloride salts. Microanalysis is not sufficient to

determine the exact number of HCl molecules that accompanied the ligand, as the molecular

mass of two water molecules is equal to the molecular mass of one HCl molecule. To precise

this aspect argentometric titrations were performed for both ligands.

Argentometry involves the titration of silver nitrate of known concentration against

the ligand solution. Silver nitrate reacts with the chloride ions forming silver chloride as a

white precipitate, according to:

Cl− (aq) + Ag+ (aq) → AgCl (s) (1)

To avoid the possible complexation of silver ion by the macrocyclic cavity, (CuNO3)2

was first added to ensure the ligand complexation. Therefore, the added silver ions are

available to determine the chloride ions brought by the ligands.

The amount of silver nitrate used to reach the equilibrium point is sufficient to

determine the percentage of chloride ions present in the ligand solution of given

concentration (Figure II-9). These titrations indicated that L1H4 is obtained as a

pentahydrochloride salt while L2H3 is obtained as a tetrahydrochloride salt.

111

N N

N NCO2HHO2C

HO2C

N

N

R

NO2

L1H4; R = H

L2H3; R =

L1H4, 5HCl,

4H2O

L2H3, 4HCl,

3H2O

C = 36.15% C = 41.1%

H = 6.21% H = 6.06%

N = 11.50% N = 11.57%

Cl = 24.25% Cl = 16.73%

Figure II-9 : L1H4, L2H3 and their compositions

2. Acido-basic properties of L1H4 and L2H3 To evaluate the affinity of L1H4 and L2H3 towards any metal ion it is necessary to

determine their protonation constants, since on the course of the complexation process,

protons and metal ions could be in competition.

a) Potentiometric study of Ligands

Protonation constants are determined by potentiometry. Potentiometric studies are

performed at a constant temperature of 25°C with ionic strength I = 0.1 (NMe4Cl) and ligands

are titrated with tetramethyl ammonium hydroxide (NMe4OH, 5×10-2 mol.L-1) in between

pH = 2 and 12. For example, potentiometric titration of L1H4 solution is reported below

(Figure II-10).

0,0 0,2 0,4 0,6 0,8 1,0 1,22

4

6

8

10

12

pH

VOH- / mL

HClL1H4

Figure II-10: Neutralization curve of L1H4 by NMe4OH 5×10-2 mol.L-1, 25°C:

(∆) 4mL of L1H4 (c = 7.76×10-4 mol.L-1 in 5×10-3 mol.L-1 HCl) and (○) 4mL of 10-2 mol.L-1

HCl, I = 0.1 (NMe4Cl)

112

The difference observed between the ligand titration curve and the HCl one, on either

side of the equivalent point (Figure II-10), evidenced for the presence of weak acidities for

L1H4. Furthermore these differences occurring between and after pH = 7, one can expect to

determine acidity constant values inferior and superior to 7 for L1H4. In order to precise the

number of these constants, a preliminary analysis of L1H4 titration curve based on the

knowledge of the equivalent volume, consists in determining the number of protons released

by L1H4 at pH = 7.

For the curve depicted in Figure II-10, L1H4 as a hydrochloride salt (L1H4, 5HCl) is

solubilized in hydrochloric acid (HCl 5×10-3 mol.L-1). Therefore at the equivalence, the

number of moles of base added neutralise all the protons released by the ligand (HCl5,HL 4

1n )

and those from the acid in excess (excess,Hn ) according to:

excess,HHLexcess,HHCl5,HLOH nn*)5x(nnn4

14

1

where x represents the number of protons released by L1H4 (in its neutral form) before pH = 7

and 4

1HLn corresponds to the number of moles of L1H4 in solution.

In the current conditions (Figure II-10):

mol10835.3n,mol101.3n,mol102n 5OH

6HL

5excess,H 4

1

which leads to a value of x = 0.92 which can be rounded to 1.

One proton is therefore released from ligand L1H4 in its neutral form, leading to the

formation of L1H3- species at pH 7. As previously said, L1H4 possesses four ionisable protons

in its neutral form. This result implicates that three acidity constant values will be superior to

7. However, since L1H4 titration begins at acidic pH (pH = 2), one cannot exclude

reprotonation of the ligand, then the existence of protonated species in solution for L1H4. The

calculation will then allow the determination of supplemental acidity constants for these

cationic species.

A similar analysis is done for L2H3 (x~1). One proton is released from ligand L2H3 in

its neutral form, leading to the formation of L2H2- species at pH 7. As previously said, L2H3

possesses three ionisable protons in its neutral form. The previous result implicates then that

two acidity constant values will be superior to 7. As highlighted for L1H4, titration of L2H3

begins at acidic pH (pH = 2). Calculations will then allow the determination of supplemental

acidity constants for the cationic species of L2H3.

113

b) Determination of protonation constants

Calculation of the protonation constants is done using PROTAF[18] from various

titrations of the ligand, varying the concentrations (for L1H4: 7.76×10-4 mol.L-1 to 4.0×10-4

mol.L-1 and 8.5×10-5 mol.L-1 to 4.25×10-5 mol.L-1 for L2H3) against NMe4OH

(5×10-2 mol.L-1). As said before, ligand is solubilized in hydrochloric acid, which permits the

reprotonation of the ligand at low pH values.

In these conditions, six protonation constants (log K01h) were determined for L1H4 and

four protonation constants for L2H3 (Table II-1). These constants were analysed by

comparison with protonation constants known for the relevant ligands L3H, L4H4 and L5H3.

(See supplementary page given for ligand structures L3H, L4H4 and L5H3).

Table II-1: Protonation constants log K0lh of L1H4 and L2H3, T = 25°C, (I = 0.1(NMe4Cl))

Equilibrium reactions log K0lh(a) L1H4 L2H3 L3H[13] L4H4

[19] L5H3[20]

L + H+ LH log K011 12.5 (1) 9.73 (4) 11.05 11.74 11.75

LH + H+ LH2 log K012 9.92 (3) 8.82 (2) 10.21 9.76 9.23

LH2 + H+ LH3 log K013 8.32(2) 5.34 (2) 9.01 4. 86 4.13

LH3 + H+ LH4

log K014 4.67 (3) 3.48 (3) 4.55 4.11 2.97

LH4 + H+ LH5 log K015 4.10 (3) <2 2.37

LH5 + H+ LH6

log K016 2.6 (1) <2

a - K01h corresponds to the reaction equilibrium

and is defined by equation ]H][LH[

]LH[K 1)(h

1h

hh

01h

Six protonation constants (log K0lh) were determined for L1H4, in which three

constants are greater than 7. By comparison with protonation constants of L3H,[13] two

protonation constants 9.9 (log K012) and 8.3 (log K013) could be associated to the acid-base

equilibria of two nitrogen atoms in the macrocycle. Protonation constants 4.1 (log K015) and

2.6 (log K016) are attributed to the acetate group functions by comparison with the protonation

constants of acetate functions in L4H4.[19] The two latter constants 12.5 (log K011) and 4.67

(log K014) could be possibly associated to the benzimidazole moiety, as first these constants

HLH 1)(h1h

LHh

h

114

are very near to the ones of ligand cyclen benzimidazole L3H (11.05 (log K011) and 4.55 (log

K014), respectively) and second, these constants are very similar to the those determined for

the protonation events in benzimidazole (Scheme II-4).[21,22,23]

NH

N

NH

HN

N

Nlog K011 log K012

log K011 (25°C) = 12.75 [21,22] log K012 (25°C) = 5.58 [23]BIM

Scheme II-4 : Protonation constants of benzimidazole (BIM)

Therefore, in ligand L1H4, protonation constant log K011 (12.5) corresponds to the

benzimidazole substituent, in which the benzimidazole (L1H3-) is deprotonated to form

benzimidazolate ([L1]4-). The constant log K012 (4.7) is associated to the protonation of the

imine nitrogen atom in benzimidazole and the formation of a benzimidazolium cation. To

check these assumptions the acid-base behaviour of L1H4 can be investigated by following

the UV absorption spectrum of the benzimidazole moiety in the range 250–290 nm, which

will be discussed in the coming section.

A similar trend is followed in the ligand L2H3, where the first and second protonation

constants (9.73 (log K011) and 8.82 (log K011)) could be associated to the protonation

constants of two nitrogen atoms in the macrocycle. The protonation constant log K014 (3.48)

could be attributed to the acidity constant of acetate group, whereas the latter constant log

K013 (5.34) could either be assigned to the re-protonation of benzimidazole (imine group) or

to the acetate function. To confirm either of the above propositions, UV absorbance

spectrum of the benzimidazole moiety in the whole pH range (2 – 12) is discussed in the

coming section.

Another way to represent these data is to determine, by using HYSS software, [24] the

domain of existence of the protonated and deprotonated species of L1H4 and L2H3 (Figure II-

14).

115

a. b.

2 4 6 8 10 120

20

40

60

80

100

[L1]4-

L1H3-L1H2

2-L1H3-

L1H4

L1H5+L1H6

2+

% o

f pro

tona

ted

spec

ies

of L

1 H4

pH 2 4 6 8 10 12

0

20

40

60

80

100

[L2]3-

L2H2-

L2H2-L2H3

L2H4

+

% o

f pro

tona

ted

spec

ies

of L

1 H3

pH

Figure II-14: Species distribution curves of ligands in function with pH.

a. [L1H4] = 7.76×10-4 mol.L-1, b. [L2H3] = 8.5×10-5 mol.L-1, T = 25°C, I = 0.1 (NMe4Cl)

Species differentiation curves of L1H4 and L2H3 show different protonated species in

the pH range 2-12. The neutral form of each ligand is different and hence the protonated

species at respective pH. In the distribution curve with L1H4, at pH = 5 to 8, L1H3- is present

in majority. In this domain of pH, in the presence of a metal ion, the complexation occurs

with this species. In L2H3 distribution curve with, at pH = 5 to 9, L2H2- is present in majority.

In this domain of pH, in the presence of a metal ion, the complexation occurs with this

species

c) Spectroscopic study of ligands

i. UV-Visible studies

Some aspects of the protonation patterns of L1H4 and L2H3 can be followed by UV-

Visible spectroscopy monitoring. Indeed, benzimidazole substituent has absorbance in the

UV region, and its signal is pH-sensitive (Figure II-15).[25] In Figure II-15, bathochromic

(red shift) and hypochromic shifts (decrease in the intensity of absorbance) are seen in the

spectrum when pH moves from acidic to basic.

116

in 0.01N HClin 0.01N NaOH

λ / nm

ε/ L

mol

-1. c

m-1


ε/ L

mol

-1. c

m-1


λ / nm

ε/ L

mol

-1. c

m-1

in 0.01N HClin 0.01N HClin 0.01N HClin 0.01N NaOH

ε/ L

mol

-1. c

m-1

Figure II-15: Absorbance of benzimidazole in UV region at acidic and basic pH.[25]

Therefore, since the UV spectrum of benzimidazole species can constitute good

witness of the protonation state of this chromophore, the evolution of L1H4 UV signal with

the pH was monitored in the range 220-350nm. The results obtained are shown in Figure II-

16.

a. b.

250 300 3500,0

0,5

1,0

% o

f pro

tona

ted

spec

ies o

f L1 H

4

pH

pH=2.58 pH=4.05 pH=6.08 pH=9.09 pH=11.45

2 4 6 8 10 120

20

40

60

80

100

LH62+

LH5+

LH4

LH3-

LH22- LH3-

L4-

% L

pH

4000

5000

6000

7000

8000

27

8 (m

ol-1 L

cm

-1)

Figure II-16: a. Evolution of UV signal of L1H4 in function of pH.

b. Superimposition of 278 (mol-1 L cm-1) = f (pH) to the speciation diagram

At pH 6, L1H3- species predominates, the electronic spectrum registered shows a

signal separated widely into two peaks at 274 and 280 nm. This signal is a characteristic

signal of benzimidazole chromophore.[26] By analogy with benzene derivatives, these bands

correspond to the π*π,[27a, b] S1(1La)S0 and S1(1Lb)S0 transitions where S1(1La) and

S1(1Lb) correspond to the two first excited states of benzimidazole.[27c]

117

When the ligand solution switches from neutral to acidic medium, protonated species

such as L1H62+, L1H5

+ are predominated in-between pH 2 and 4. From pH 4, neutral species

L1H4 is predominant with a slight hypsochromic shift of the two bands (∆ = 2 nm). In

parallel an hyperchromic shift is seen ( = 272 and 278 nm), with an increase of the molar

absorption coefficient of almost 50% compared to the initial value determined at pH = 6. This

phenomenon, already observed in cyclen benzimidazole L3H[13] showcases the protonation of

nitrogen atom of imine in benzimidazole.[26]

No difference in the absorbance is seen between pH 6 to pH 11. So, it is shown that

the benzimidazole probe has no involvement in the protonation constants between these pHs.

Above pH 11, an hypochromic shift is observed for the benzimidazole signal and at the same

time a new band raised at λ = 300nm. These two phenomenons, already observed for the

benzimidazole chromophore, [22,25,26] correspond to the deprotonation of the nitrogen atom of

the pyrrole in benzimidazole and then to the formation of a benzimidazolate species.

For L2H3, the evolution of the benzimidazole UV signal upon pH variation of probe

was monitored. Figure II-17.b shows this evolution. A little decrease in the evolution is seen

around pH 4. However, due to its diluted concentration, the significance in the changes of

evolution is tricky to consider.

a. b.

200 300 4000,0

0,5

1,0

1,5

Abso

rban

ce

nm

pH = 3.03 pH = 4.23 pH = 5.27 pH = 6.27 pH = 7.25 pH = 8.34 pH = 9.28 pH = 10.17 pH = 11.25

2 4 6 8 10 12

4000

4500

5000

5500

6000

6500

7000

7500

8000

278(m

ol-1 L

cm

-1)

pH

L1H4

L2H3

Figure II-17: a. Evolution of UV signal of L2H3 in function of pH. b. Comparison with the

epsilon (at = 278 nm) of L1H4

To confirm and assign this evolution to the deprotonation of benzimidazolate ion,

NMR spectroscopic studies are followed (next subsection).

To resume, the absorbance of benzimidazole probe in UV region permitted us to

monitor its protonation or deprotonation varying the pH. The passage of L1H3- species

118

(major species at pH = 6) to L1H4 (major species at pH = 4) correspond to the protonation of

the imine nitrogen atom of benzimidazole probe (log K014 = 4.7). The passage of L1H3- species

(major species at pH = 10) into [L1]4- (major species at pH = 11) correspond to the

deprotonation of the nitrogen atom of the pyrrole in benzimidazole probe (log K011 = 12.5).

For L2H3, the deprotonation of benzimidazole moiety and the attribution of log K could be

confirmed following the NMR study.

ii. NMR spectroscopy

The similar type of study could be followed using 1H NMR spectroscopy by

monitoring L1H4 and particularly the benzimidazole moiety ones according to pH. 1H NMR of L1H4 and peak assignments are reported in Figure II-18.

Hα and Hβ N-CH2-imidazole Aliphatic DO3A

(22) H(2) H(2) H (2) H

N N

N NCO2HHO2C

HO2C

HN

N


(22) H(2) H(2) H (2) H


(22) H(2) H(2) H (2) H

N N

N NCO2HHO2C

HO2C

HN

N

Figure II-18: 1H NMR spectra of L1H4 in D2O.

The peaks present in upfield around δ = 3.0-3.32 ppm represent aliphatic protons of

methyl groups in DO3A backbone. The peak around δ = 4.2 represent the protons of

methylene group that link DO3A backbone and benzimidazole (N-CH2-benzimidazole).

Peaks toward downfield around δ = 7.4-7.7 represent the aromatic protons of benzimidazole.

The corresponding signals are present in the form of two doublets of doublets. The shape of

119

this signal agrees with a -type local symmetry making the aromatic Hα protons and the Hβ

protons as almost equivalent pairs.

N N

N NCO2HHO2C

HO2C

HN

Nα

α

β

β

N N

N NCO2HHO2C

HO2C

HN

Nα

α

β

βN N

N NCO2HHO2C

HO2C

HN

Nα

α

β

β

Figure II-19: 1H NMR spectra of benzimidazole probe in L1H4, D2O, c = 10-2 mol.L-1,

300MHz

Changes in proton chemical shifts as a function of pH show that the progression to the

acidic pH is accompanied by a deshielding of the proton signals of benzimidazole probe. This

phenomenon is similar to that observed in L3H,[13] is due to an increase in the positive charge

density on the aromatic ring. The existence of a more pronounced deshielding around pH =

4.7 can be correlated to the protonation of the imine function of the benzimidazole moiety

and confirms the attribution of constant (log K014 = 4.67) to the protonation of benzimidazole.

7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8

Aromatic zone Aliphatic zone ppm

pH = 11.0

pH = 6.5

pH = 1.9pH = 1.0

pH = 3.0pH = 3.7

pH = 4.7

pH = 8.8

pH = 7.7

pH = 10.0

7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8


7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8


7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8

Aromatic zone Aliphatic zone

7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8


pH = 11.0

pH = 6.5

pH = 1.9pH = 1.0

pH = 3.0pH = 3.7

pH = 4.7

pH = 8.8

pH = 7.7

pH = 10.0

7.8 7.6 7.4 7.2 4.4 4.2 4.0 3.8


Figure II-20: Evolution of δ (Haromatics and Haliphatic) of L1H4 (c =

10-2 mol.L-1) according to the pH (D2O, 300 MHz)

120

The protons of methylene group (labelled as *, δ ~ 4.4 ppm) that binds the parent

DO3A and benzimidazole group also undergo a downfield shift simultaneously with the

modification in benzimidazole moiety on varying the pH.

For L2H3, 1H NMR spectra and peak assignments are reported in Figure II-21

Aromatic (Benzimidazole;Nitrobenzene) N-CH2-Nitrobenzene N-CH2-Imidazole Aliphatic DO3A

(2) H (22) H(2) H(8) H

N N

N NCO2HHO2C

HO2C

N

N

O2N


(2) H (22) H(2) H(8) H


(2) H (22) H(2) H(8) H

N N

N NCO2HHO2C

HO2C

N

N

O2N

Figure II-21: 1H NMR spectra of L2H3

The peaks present in upfield around δ = 2.8-3.96 ppm represent aliphatic protons of

methylene groups in DO3A backbone. Their assignement, given in Figure II-21 is

established by comparison with L1H4 signals. The main differences are due the nitrobenzene

substituent that is to say at δ = 6.0 ppm for the methylene protons that connect benzimidazole

with nitrobenzene (N-CH2-Nitrobenzene) and signals around δ = 7.4-8.3 that represent the

aromatic protons of both benzimidazole and nitrobenzene. The nitrobenzene group has a

symmetrical plane and protons are present as equivalent pairs. The corresponding signals at δ

= 7.4 ppm and 8.28 ppm are present in the form of two doublets which could correspond to

the protons of nitrobenzene. Other signals at 7.65-7.77 ppm, present in the form of multiplets

and at 7.94 ppm in the form of doublet could correspond to that of benzimidazole moiety,

which lost their symmetry compared to L1H4.

121

Changes in proton chemical shifts as a function of pH show the progression of

deprotonation of ligand (Figure II-22).

pH = 1.7

pH = 2.6

pH = 4.1

pH = 5.0

pH = 5.9

pH = 6.8

pH = 8.1

pH = 9.0

pH = 10.2

pH = 11.1

8.4 8.2 8.0 7.8 7.6 7.4 6.0 5.8 4.46.2 4.6 4.4

Aromatic Zone Aliphatic Zone

pH = 1.7

pH = 2.6

pH = 4.1

pH = 5.0

pH = 5.9

pH = 6.8

pH = 8.1

pH = 9.0

pH = 10.2

pH = 11.1

8.4 8.2 8.0 7.8 7.6 7.4 6.0 5.8 4.46.2 4.6 4.4


Figure II-22: Evolution of δ (Haromatics and Haliphatic) of L2H3 (c = 8.5×10-5 mol.L-1) as a

function of pH ( D2O, 500 MHz)

For 1.7<pH<4.1, an upfield shift is seen in benzimidazole group (δ = 7.7 ppm), in

aliphatic protons of methylene groups linking benzimidazole to nitrobenzene group (δ = 6.0

ppm) and DO3A parent backbone to benzimidazole (δ = 4.5 ppm). This indicates the

deprotonation of benzimidazolium ion and therefore helps to remove any doubt on the

deprotonation event that corresponds to a constant value of log K = 3.48. Beyond pH = 5.0,

no chemical shift is seen indicating the absence of role of benzimidazole moiety. The latter

deprotonations correspond to the modification of one acetate function (log K013= 5.34) and

finally to macrocyclic amine functions (Scheme II-6).

Therefore, basing on the above results, the following protonation schemes could be

proposed for L1H4 and L2H3.

122

N N

NN

-OOC COO-

-OOCN

N

N N

NN

-OOC COO-

-OOCN

HN

N N

NN

-OOC COO-

-OOCN

HN

H+

N N

NN

-OOC COO-

-OOCN

HN

2H+N N

NN

-OOC COO-

-OOCHN

HN

N N

NN

-OOC COOH

-OOCHN

HN

N N

NN

-OOC COOH

HOOCHN

HN

log K011 = 12.5

[L1]4- L1H3- L1H22-

L1H3-L1H4L1H5

+

L1H62+

2H+2H+

2H+

log K012 = 9.9

log K013 = 8.3

log K016 = 2.6

log K014 = 4.7log K015 =4.1

Scheme II-5 : Protonation sequence of ligand L1H4

N N

N NCO2

--O2C

-O2CN

N

R

N N

N NCO2

--O2C

-O2CN

N

R

N N

N NCO2

--O2C

-O2CN

N

R

H+log K011 = 9.73 log K012 = 8.82 2H+

L2H2-L2H2-

[L2]3-

N N

N NCO2

--O2C

HO2CN

N

R

log K013 = 5.34

log K014 = 3.48 2H+

L2H4+ L2H3

N N

N NCO2

--O2C

HO2CN

NH

R

2H+

Scheme II-6 : Protonation sequence of ligand L2H3

123

C. Coordination chemistry of L1H4 and L2H3 This section is devoted to the study of coordination chemistry of ligands L1H4 and

L2H3 towards transition metal ions (Cu(II) and Zn(II)) and lanthanides (Gd(III) and Eu(III)).

1. Physicochemical studies with Cu(II) and Zn(II) Studies with transition metal(II) ions such as Cu(II) and Zn(II) give primary

information about their stabilities, which could be used further in the investigation of kinetic

inertness of Gd complex in the presence of endogenous metal cations.

a) Potentiometric study

Generally the stability constants of a complex including their protonated forms are

determined by potentiometric method. For this kind of experiments, acidic solution

containing the equivalent amount of the metal ion and the ligand is titrated directly by a base

(NMe4OH). As the ligands are macrocyclic, complex formation is slow in acidic medium

and this restricts to use the direct titration method. To overcome this problem, the technique

called ‘Out-of-cell’ method is used.[8,13,28] In this method, preparation of solutions is done

separately by fixing the amount of ligand and metal to be utilized. The ligand is dissolved in

acid to start the experiments from acidic pH (pH ~ 2). To each of these prepared solutions, a

known quantity of base NMe4OH (5×10-2 mol.L-1) is added consequently varying the pH

(until pH ~ 7). These solutions are stored for about 30 days under argon at 40°C to speed up

the complex formation and to ensure that the thermodynamic equilibrium is reached. After

this period of time, pHs are measured and plotted according to the varying volume. An

example of such a titration is presented in the Figure II-23.a.

a. b.

0,0 0,1 0,2 0,3 0,4 0,5 0,6 0,72,0

2,5

3,0

3,5

4,0

4,5

pH

VOH- /mL 0,0 0,1 0,2 0,3 0,4 0,5 0,6

4

5

6

7

8

9

10

11

12

pH

VOH- /mL

Figure II-23 a. Out-of-cell method, b. Normal titration continued; ([L1H4] = [M-L1H4, (M

= Cu, Zn)] = 7.76×10-4 mol.L-1, [OH-] = 5×10-2 mol.L-1 , T = 25°C, I = 0.1 (NMe4Cl)

124

As these solutions are stored for 30 days, possibility of fungal growth or above pH~7

CO2 dissolution in each cell restricts the method applicability. Nevertheless, to have pH

information above pH = 7, a direct titration of individual flasks, for which pH was previously

adjusted in-between pH 3 ~ 4 was performed with NMe4OH (5×10-2 mol.L-1) as a base and

this, until pH = 12 (Figure II-23.b).

The formation constants of copper and zinc complexes ( β hm ) are determined by

potentiometric titrations at 25°C and with ionic strength I = 0.1 (NMe4Cl). These

determinations are done by experimenting the ligand solutions at known concentration

(7.7610-4 mol.L-1 for L1H4 and 8.510-5 mol.L-1 for L2H3) dissolved in HCl 5×10-3 mol.L-1,

and by varying the ratios of metal for each set of solutions. The ratios R = [Ligand]/ [Metal]

used are 1.06 and 1.24 respectively.

As an example in Figure II-24, two neutralization curves are represented, in which of

the first one corresponds to the neutralization of the ligand L1H4 (initial volume = 3mL,

7.7610-4 mol.L-1 in HCl 5×10-3 mol.L-1) and the other corresponds to Cu(II)-L1H4 titration

in ratio R = [L] / [M] = 1.24.

0,0 0,2 0,4 0,6 0,8 1,02

4

6

8

10

12

pH

VOH- / mL

L1H4 R = 1.24

Figure II-24: Neutralization curves of Cu(II)-L1H4 solutions by NMe4OH 510-2 mol.L-1

At the beginning of the titrations, the neutralization curve of Cu(II)-L1H4 shows

decrease in pH when compared to the pH of the ligand alone. This decrease in pH

corresponds to the liberation of protons due to complexation.

In Figure II-25, the difference in-between the equivalent points of ligand alone and

Cu(II)-L1H4 corresponds to the supplementary protons released from the ligand due to

complexation.

125

∆V = 0.066

0,0 0,2 0,4 0,6 0,8 1,02

4

6

8

10

120.571 0.637

pH

VOH- / mL

L1H4

R = 1.24

∆V = 0.066

0,0 0,2 0,4 0,6 0,8 1,02

4

6

8

10

120.571 0.637

pH

VOH- / mL

L1H4

R = 1.24

Figure II-25: Neutralization curves for L1H4 and L1H4-Cu(II) system (R = [L]/[M] = 1.24)

[NMe4OH] = 5×10-2 mol.L-1, [L1H4] = 7.76×10-4 mol.L-1, [HCl] = 5×10-3 mol.L-1, I = 0.1

(NMe4Cl) at 25°C.

In a first intention, data exploitation consists in analysing the value of the equivalent

point as following: in the zone of equivalence, the species distribution curve (Figure II-19.a,

page 27) shows that the ligand is present in the form of L1H3- species. At pH = 7, the addition

of copper ions to L1H3- species leads to the formation of protonated [CuL1Hn] (0<n<2). The

equation of the reaction and the species formed could be represented as below:

Cu2+ + L1H3- [CuL1Hn] + (3-n) H+

ninitial(mol) 1/R 1

nt(mol) (1-1/R) 1/R (3-n)×1/R

For instance, for a ratio R = [L]/[M] = 1.24 and when the copper species formed at pH

= 7 is [CuL1H2], the previous balance sheet will become as following:

Cu2+ + L1H3- [CuL1H2] + H+ .

ninitial(mol) 0.8066 1

nt(mol) 0.1934 0.8066 0.8066

Now, for the same ratio [L]/[M] and and when the copper species formed at pH = 7 is

[CuL1H]-, the previous balance sheet will become as following:

126

Cu2+ + L1H3- [CuL1H]- + 2H+ .


nt(mol) 0.1934 0.8066 1.6130

Finally, for the same ratio [L]/[M] and and when the copper species formed at pH = 7

is [CuL1]2-, the previous balance sheet will become as following:

Cu2+ + L1H3- [CuL1]2- + 3H+ .


nt(mol) 0.1934 0.8066 2.4198

For the ratio R = 1.24, the formation of [CuL1H2] theoretically releases 0.8066 moles

of protons from the ligand introduced, which needs 0.037 mL of base for neutralization. The

formation of [CuL1H]- releases 1.613 moles, which needs 0.073 mL and the formation of

[CuL1]2- needs 0.112 mL of base for neutralization. Experimentally, the volume of the base

required for neutralization for the ratio R = [L]/[M] = 1.24 is 0.066 mL (Figure II-25), which

is close to the value of 0.073 mL. This comparison validates the hypothesis of the formation

of [CuL1H]- species at pH 7.

b) Thermodynamic stability of Cu (II) and Zn(II) complexes

PROTAF software permits to determine the best chemical model in solution able to

account for the experimental potentiometric curves. From this fit, the overall stability

constants ( β hm ) of the transition metal complexes of L1H4 and L2H3 could therefore be

obtained. These constants correspond to the equilibrium:

m M + L + h H hm HLM

with hmhm

hm ]H[]L[]M[]HLM[

β

For L1H4, the computed chemical model suggests the formation of [MLH4], [MLH3],

[MLH2], [MLH] and [ML] species while for L2H3 it implies the formation of [MLH], [ML],

[MLH-1] and [MLH-2] species with metal M(where M = Cu(II) or Zn(II)). The results are

presented in Table II-2. These results are compared with other ligands L4H4 and L5H3 (Table

127

II-3) which were determined under similar experimental conditions.[19, 20] (these ligand

structures are gathered in the loose sheet)

Table II-2: Stability constants of M/L1H4 and M/L2H3 complexes (where M = Cu(II) or

Zn(II)), T = 25°C, I = 0.1 (NMe4Cl) ) – Standard deviations are given in parenthesis

L1H4 L2H3

Cu(II) Zn(II) Cu(II) Zn(II)

log mlh M+ L+4H+ = MLH4

41.1 (1)

M+L+3H+ = MLH3 39.8 (4) 37.8 (1)

M+L+2H+ = MLH2 35.5 (5) 33.6 (1)

M+L+H+ = MLH 31.0 (4) 28.5 (1) 22.6 (2) 22.3 (2)

M+L = ML 21.8 (4) 18.8 (2) 16.0 (5) 15.9 (4)

M+L = MLH-1 + H+ 7.1 (5) 5.9 (5)

M+L = MLH-2 + H+ -3.9 (5)

Table II-3: Stability constants of M/L4H4 and M/L5H3 complexes (for comparison with those

of M/L1H4 and M/L2H3 complexes.)

L4H4 L5H3

Cu(II) Zn(II) Cu(II) Zn(II)

log mlh M+ L+4H+ = MLH4

M+L+3H+ = MLH3

M+L+2H+ = MLH2 30.27 28.36 27.4

M+L+H+ = MLH 26.75 24.87 25.0 23.2

M+L = ML 22.44 20.52 21.1 19.0

The obtained stability constant values show that the complexes are

thermodynamically stable. If we compare the thermodynamic stability of different complexes

formed, the direct comparison with their overall stability constants is not relevant because of

the different protonation constants for their ligands. Nevertheless for the same ligand, one can

compare the stability of copper and zinc complexes. The data show that copper complexes

128

are more stable than the zinc ones (especially for L1H4 ligand) and this result follows the

known Irving - Williams order for the first row transition metal ions.[8]

From the logmlh values, the deprotonation constants of both Cu(II) and Zn(II)

complexes are calculated.

Table II-4: Deprotonation constants of M/L1H4 and M/L2H3 complexes (where M = Cu(II) or

Zn(II)), compared with other metal complexes.

L1H4 L2H3 L4H4 L5H3

Log K Cu(II) Zn(II) Cu(II) Zn(II) Cu(II) Zn(II) Cu(II) Zn(II)

MLH4 = MLH3 + H 3.3

MLH3 = MLH2 + H 4.3 4.2

MLH2 = MLH + H 4.5 5.1 3.52 3.49 2.4

MLH = ML + H 9.2 9.7 6.6 6.4 4.31 4.35 3.9 4.2

ML = MLH-1 + H 8.9 10.0

ML = MLH-2 + H 9.8

In metal complexes with L1H4, if we assign the deprotonation constants (log K = 9.2

of Cu(II) complex and log K = 9.7 of Zn(II) complex) to the pyrrole nitrogen atom of the

benzimidazole, the high acidic nature of the deprotonation constants than that observed in the

ligand (log K (L1H4) = -12.5) could explain the involvement of benzimidazole in complex

formation. This means that in the complex, the acidity of the hydrogen of pyrrole nitrogen

atom of the benzimidazole is enhanced. This type of deprotonation has already been observed

in the case of the ternary complex [Zn(cyclen)]2+-imidazole where the imidazole

complexation by [Zn (cyclen)]2+ assists imidazole deprotonation.[29]

N

N

N

N

N

NH

Zn

H

HH

H

2+

N

N

N

N

N

N

Zn

H

HH

H

+

log K = 9.5

Scheme II-7: Effect of Lewis acidity of Zn(II) on the deprotonation of imidazole.[29]

Nevertheless, these log K values could also be assigned to the deprotonation of water

molecule. To decide between these two hypotheses, M-L1H4 and M-L2H4 systems (M =

Cu(II) or Zn(II)) will be studied by UV spectroscopy in the forthcoming section. In metal

complexes with L2H3, the deprotonation constants (log K = 8.9 of Cu(II) complex and

129

log K = 9.8 and 10.0 of Zn(II) complex) could be due to the deprotonation of a water

molecule.

Basing on the potentiometric values, speciation diagrams of M-L1H4 and M-L2H3

systems, with M = Cu(II) and Zn(II), are traced with the help of HYSS (Figure II-26)

a. b.

2 4 6 8 10 120

20

40

60

80

100

L1H62+

Cu2+

% Cu

[CuL1]2-[CuL1H]-

[CuL1H2]

[CuL1H3]+

pH 2 4 6 8 10 12

0

20

40

60

80

100[CuL2OH]-[CuL2]-[CuL2H]Cu2+

%C

u

pH Figure II-26: Species distribution curves of a. Cu(II)-L1H4 and b. Cu(II)-L2H3 systems.

This diagram shows different species for two ligands (L1H4 and L2H3) at a specific

point of pH. This can be explained by the difference in the number of protons of the ligands

in their neutral forms. [CuL1H3]+ species forms at pH = 2 and the other species [CuL1H2] and

[CuL1H]- prevails between pH 2 and pH 10. Beyond pH 9, the deprotonated [CuL1]2-

complex becomes predominant.

For Cu(II)-L2H3 system, protonated species [CuL2H] begins to form at pH = 2. The

deprotonated [CuL2]- complex prevails from pH 5 to 10. Beyond pH = 9, the formation of

hydroxylated complexes could be proposed.

a. b.

2 4 6 8 10 120

20

40

60

80

100

LH5+

LH62+

Zn2+

[ZnL1H4]2+

[ZnL1H3]+

[ZnL1H2]

[ZnL1H]-

[ZnL1]2-

% Z

n

pH

2 4 6 8 10 120

20

40

60

80

100

[ZnL2(OH)2]-

[ZnL2(OH)]-

[ZnL2]-[ZnL2H]Zn2+

%Zn

pH

Figure II-27: Species distribution curves of a. Zn(II)-L1H4 and b. Zn(II)-L2H3 systems.

130

This diagram (Figure II-27) shows that species [ZnL1H4]2+ forms before pH 2, while

[ZnL1H3]+ and [ZnL1H2] are prevalent from pH = 2 to 5. [ZnL1H2]- prevails between pH 4 and

pH = 10. Beyond pH 9, the deprotonated [ZnL1]2- complex becomes predominant, similar to

that of [CuL1]2-.

In the case of Zn(II)-L2H3, protonated species [ZnL2H] is present before pH = 4. The

deprotonated species [ZnL2]- prevail from pH = 4. Deprotonation of water molecules were

determined in [ZnL2]- above pH 10.

As previously said, the direct comparison of the overall stability constants of the

complexes to judge their thermodynamic stability is not relevant because of the different

protonation constants for their corresponding ligands. One way to do so is to calculate the

percentage of free metal ion present between pH 2 and 12 for each complex. The logarithm of

the percentage of free metal ion (log([M]free/[M]total)) in function of pH is traced for

M-L1H4 and M-L1H4 systems (M = Cu(II) or Zn(II)). The obtained curves are compared with

those of L4H4 and L5H3.

a. b.

-10

-8

-6

-4

-2

0

L4H4

L2H3

L1H4,L5H3

pH

log(

[Cu]

free/[

Cu]

tota

l)

2 4 6 8 10 12

-10

-8

-6

-4

-2

0

L5H3

L4H4

L2H3

L1H4

pH

log(

[Zn]

free/[

Zn] to

tal)

2 4 6 8 10 12

Figure II-28: Logarithm of percentage of free metal ion concentration of L1H4, L2H3, L3H

and L5H3 in relation to M(II), [ligand] = [M = Cu/Zn] = 2.0×10-3 mol.L-1

a. Cu complexes and b. Zn complexes

The lower is the curve the more stable is the complex. From the Figure II-28.a, it is

therfore clear that Cu(II) complex with L1H4 is more stable than Cu(II)-L2H3 complex. The

stabilities of Cu(II)-L1H4 and Cu(II)-L5H3 are almost the same, but nevertheless less stable

than Cu(II)-L4H4.

In the case of zinc complexes (Figure II-28.b), Zn(II)-L1H4 and Zn(II)-L2H3 are

equally stable and are slightly lower in stability when compared with Zn(II)-L5H3 and Zn(II)-

L4H4 systems.

131

c) Spectroscopic study of Cu (II) and Zn (II) complexes

i. UV – visible studies

The evolution of UV signal of L1H4-metal complexes in function with pH is

monitored using UV spectroscopy basing on the benzimidazole probe (Figure II-29).

a. b.

260 280 300 320 3400

1

2

3

4

Abso

rban

ce

/ nm

pH = 3.09 pH = 4.25 pH = 5.06 pH = 5.89

260 280 300 320 340

0

1

2

3

Asor

banc

e

/ nm

pH = 8.94 pH = 9.76 pH = 10.25 pH = 10.83 pH = 11.36

Figure II-29: Evolution of UV signal of Cu(II)-L1H4 by addition a. of HCl (1×10-2 mol.L-1)

b. NMe4OH (5×10-2 mol L-1); T = 25°C; I = 0.1 (NMe4Cl) ; l = 1cm

The speciation diagram of Cu complex shows (Figure II-26.a) that below pH = 3, the

[CuL1H3]+ complex is not yet completely formed. At this pH, the UV spectrum resembles to

that of the ligand alone (Figure II-14.a, Page 82). Above pH = 3, [CuL1H2] species is

formed where a slight decrease in the intensity of the bands = 272 nm and = 278 nm is

seen in the spectrum.

Figure II-29.b represents the evolution of UV signal of Cu(II) complex at basic pH.

After pH = 9, the important and significant changes are seen in the spectrum where two

supplementary bands are formed ( = 276 nm and = 282 nm). We can deduce that the

variation observed in the bands is due to the modification of the electron density around the

metal, consecutive to the deprotonation of benzimidazole.[30]

The superimposition of UV signal variation on the speciation diagram of Cu(II)

complexes was reported (Figure II-30). The evolution of 278 indicates the diminution of

epsilon during the formation of [CuL1H2] and the evolution of 282 superposed quite well with

the formation of [CuL1]2- species. The band at = 282 nm indicate the deprotonation of

nitrogen atom of the pyrrole in benzimidazole forming benzimidazolate.[26]

132

a. b.

2 4 6 8 10 120

20

40

60

80

100

LH73+

LH62+

[CuL]2-

[CuLH]-

[CuLH2]

[CuLH3]+

Cu2+

%C

u

pH10500

11000

11500

12000

12500

13000

13500

14000

14500

15000

15500

27

8 (L

mol

-1 c

m-1)

2 4 6 8 10 120

20

40

60

80

100

L1H73+

L1H62+

[CuL1]2-

[CuL1H]-

[CuL1H2]

[CuL1H3]+

Cu2+

%C

u

pH

7500

8000

8500

9000

9500

10000

10500

11000

11500

28

2 (L

mol

-1 c

m-1)

Figure II-30: Evolution of a. 278 (mol-1 L cm-1) = f(pH) and b. 282 (mol-1 L cm-1) = f(pH)

followed on species differentiation diagram. ([Cu(II)-L1H4] = 5×10-4 mol.L-1)

Zinc complex exhibits almost the same behaviour as copper complex. Formation of

benzimidazolate ion is seen after pH = 9 (supplementary bands at = 276 nm and = 282

nm) (Figure II-31).

a. b.

3000,0

0,5

Abso

rban

ce

nm

pH=3.415 pH=4.089 pH=5.106 pH=6.347 pH=7.15 pH=8.249 pH=9.2 pH=10.059 pH=11.46

2 4 6 8 10 12

0

20

40

60

80

100

[ZnL1H4]2+

[ZnL1H]-

[ZnL1]2-

[ZnL1H2]

[ZnL1H3]+

L1H5+

L1H62+

Zn2+

% Z

n

pH7500

8000

8500

9000

9500

10000

10500

27

8 (L

mol

-1 c

m-1)

Figure II-31: a. Evolution of UV signal of Zn(II)-L1H4 in function to pH. b. 278 (mol-1 L cm-

1) = f (pH) followed on species differentiation diagram. ([Zn(II)-L1H4] = 5×10-4 mol.L-1)

These results give us the confirmation of the deprotonation of the benzimidazole in

the complex. The deprotonation of benzimidazole in ligand alone (L1H4) is at pKa 12.5,

whereas in the metal complexes pKa values are decreased to pKa ~ 9 because the

deprotonation is made easy in the complexes than in the ligand.[29]

Similar UV spectroscopic studies were reproduced with M-L2H3 complexes (where

M = Cu(II) and Zn(II)).

133

a. b.

200 300 4000,0

0,5

1,0

1,5

Abso

rban

ce

nm

pH = 2.99 pH = 4.03 pH = 5.10 pH = 6.21 pH = 7.10 pH = 8.48 pH = 9.45 pH = 10.34 pH = 11.42

200 300 4000,0

0,5

1,0

1,5

Abso

rban

ce

(nm)

pH = 3.34 pH = 4.23 pH = 5.33 pH = 6.91 pH = 8.35 pH = 9.80 pH = 11.20

Figure II-32: a. Evolution of UV signal of Cu(II)-L2H3 in function with pH. b. Evolution of

UV signal of Zn(II)-L2H3 in function with pH

Figure II-32.a represents the evolution of copper complex according to pH. No

significant changes are seen in the spectrum, indicating the absence of benzimidazolate ion

formation. Same results were obtained for Zn(II)-L2H3 (Figure II-32.b). Therefore, the

deprotonation events calculated on the basis of potentiometric results could be attributed to

the deprotonation of water molecule.

ii. NMR spectroscopy

The similar type of study could be performed by following the evolution of chemical

shifts of L1H4 protons in presence of Zn(II) as a function of pH(Figure II-33). For that, the

synthesis of the zinc complex was first performed to lead to the [ZnL1H]- species and then the

pH was adjusted in the NMR tube with DCl or NaOD.

134

7.8 7.6 7.4 7.4 4.3 4.2 4.1 4.0


ppm

pH = 6.7

pH = 5.7

pH = 4.7

pH = 3.8pH = 2.8

pH = 1.7

pH = 9.1

pH = 10.2

pH = 11.4

pH = 8.1

7.8 7.6 7.4 7.4 4.3 4.2 4.1 4.07.8 7.6 7.4 7.4 4.3 4.2 4.1 4.0


ppm

pH = 6.7

pH = 5.7

pH = 4.7

pH = 3.8pH = 2.8

pH = 1.7

pH = 9.1

pH = 10.2

pH = 11.4

pH = 8.1

Figure II-33: Evolution of δ (Haromatics and Haliphatic) of [ZnL1H]- complex as a function to pH,

D2O, c = 5×10-3 mol.L-1, 300 MHz

For 2.8 < pH < 5.7, the signals associated to the benzimidazole probes exhibit an

upfield shift, when the pH increased. In the aliphatic zone, an upfield shift of δ = 4.26 ppm to

δ = 4.12 ppm is seen in the methylene group proton signal. This indicates an increase of the

electronic charge on the benzimidazole ring, due to the deprotonation of the benzimidazolium

moiety.

For 8.1 < pH < 11.4, the aromatic signals of the benzimidazole probe undergo another

upfield shift when the pH increased. This evolution signifies a second deprotonation on the

benzimidazole ring. This result provides a confirmation of the deprotonation of secondary

amine function in benzimidazole moiety in the complex, with a constant near log K110 = 9.7.

Changes in proton chemical shifts as a function of pH are followed for Zn(II)-L2H3

system, and are reported in Figure II-34.

135

pH = 1.7

pH = 2.9

pH = 3.7

pH = 4.8

pH = 6.1

pH = 7.0

pH = 8.1

pH = 9.3

pH = 10.4

pH = 11.3

8.5 8.0 7.5 6.0 4.5 3.5 3.0 2.5 2.0

Aromatic Zone Aliphatic ZoneAromatic Zone Aliphatic ZonepH = 1.7

pH = 2.9

pH = 3.7

pH = 4.8

pH = 6.1

pH = 7.0

pH = 8.1

pH = 9.3

pH = 10.4

pH = 11.3

8.5 8.0 7.5 6.0 4.5 3.5 3.0 2.5 2.0


pH = 1.7

pH = 2.9

pH = 3.7

pH = 4.8

pH = 6.1

pH = 7.0

pH = 8.1

pH = 9.3

pH = 10.4

pH = 11.3

8.5 8.0 7.5 6.0 4.5 3.5 3.0 2.5 2.0

Aromatic Zone Aliphatic ZoneAromatic Zone Aliphatic Zone

Figure II-34: Evolution of δ (Haromatics and Haliphatic) of [ZnL2H2]+ as a function of pH, D2O,

c = 8.5×10-5 mol.L-1, 500 MHz

For Zn(II)-L2H3 system, at pH = 1.7, the spectrum is different compared to the

spectrum of the ligand alone (at the same pH). This could indicate an early complexation of

Zn, in acidic pH conditions. Furthermore, no chemical shift in the signals of methylene group

linkers is seen in the whole pH range 2-12, indicating the involvement of benzimidazole

moiety in the metal coordination in the whole pH range. For pH>4, since no other spectrum

modification can be seen, one can deduce that deprotonation events determined by

potentiometry would not involve the benzimidazole moiety. Consequently, the deprotonation

constant of [ZnL2H] complex at logK= 6.4 could correspond to the proton release from a

protonated macrocyclic nitrogen atom. The two latter ones (logK= 9.8 and 10) could be

attributed to the deprotonation of water molecules in the Zn(II) coordination sphere, probably

accompanied by the decoordination of a carboxylate function.

In the aliphatic zone, at 1.7<pH<3.7, slight modification of DO3A backbone signals

(δ = 2.8-3.3 ppm) is seen, and in the aromatic zone, a slender downfield shift is seen in one of

the signals of benzimidazole moiety. This downfield shift could be due to the influence of

supplementary acid base equilibrium in the ligand backbone. Basing on the NMR

experiments, at acidic pH conditions, the presence of a [ZnL2H2] species could be proposed

in which the benzimidazole group participate at the Zn(II) coordination sphere.

136

iii. EPR spectroscopy

To have an additional insight in the coordination sphere of copper complexes,

Electron Paramagnetic Resonance (EPR) of Cu(II)-L1H4 and Cu(II)-L2H3 systems were

recorded at pH~7 (Figure II-35).

Copper complexes have an absorption around 3200G, characteristic

of the ∆ms = 1 transition for Cu(II) ion. Each spectrum is axial with a perpendicular signal

centered at 3200G and a parallel component at 2900G. The parallel component has a

hyperfine structure consisting of four equidistant lines. The lines are due to the coupling

between the electron spin of the unpaired electron in Cu(II) and the nuclear spin of copper

(I = 3/2).

a. b.

Figure II-35: EPR spectra a. [CuL1H]- and b. [CuL2]- ( pH 7, water/ethanol 10% - 150K)

Table II-5 represents the parameters of the corresponding EPR spectra and those of

relevant copper complexes with ligands cyclen, DTPA, L4H4 and L3H respectively. [CuL1H]-

and [CuL2]- parameters were simulated with the help of software Xsophe.[31]

137

Table II-5: EPR data for Cu(II)-L1H4, Cu(II)-L2H3 systems and other similar complexes

(pH~7, water/ethanol 10% - 150K)

g// g┴ A// (10-4 cm-1) A┴ (10-4 cm-1) Ref

[CuL1H]- 2.302 2.085 133.73 21.26 This work

[CuL2]- 2.302 2.085 133.73 21.26 This work

CuL4H2 2.300 2.062 150.3 -- [32 ]

Cu(cyclen) 2.198 2.057 184 24 [33]

CuL3H 2.207 2.049 173.0 13.1 [13]

Cu(DTPA) 2.30 2.10 140.0 10.0 [34]

The spectra obtained at pH~7 for both copper complexes are similar

(with A// = 133.73×10-4 cm-1, g// = 2.302 and g┴ = 2.085). This indicates the presence of a

similar geometry in both systems at pH = 7. For each complex, the g// value is superior to g┴

and are >2. This is typical of axially symmetric d9 copper(II) complexes in a ground-state

doublet with the unpaired electron in the dx2–y2 orbital.[35] The values (A// and g//) obtained for

these complexes are entirely different from those corresponding to pentacoordinated

Cu(cyclen) (A = 18410-4 cm-1, with a N4O coordination sphere for Cu(II))[33] or to cyclen

benzimidazole (A// = 173×10-4 cm-1, with a N5 coordination sphere for Cu(II))[13]. The values

obtained for [CuL1H]- and [CuL2]- complexes are closer to that obtained for copper

complexes with DTPA (A// = 14010-4 cm-1[34]) where the copper is hexacoordinated with a

N2O4 or NO5 coordination sphere.

Therefore, this could indicate that for [CuL1H]- and [CuL2]- complexes present at

pH 7, copper is hexacoordinated with less than four nitrogen atoms in the coordination

sphere.

Further for each copper system, several species are involved in the whole pH range

and their formation could be followed by EPR spectroscopy. The evolution of EPR spectra

of Cu(II)-L1H4 system according to pH is given in (Figure II-36).

138

2500 2700 2900 3100 3300 3500Gauss

CuNO3

pH=1.7

pH=4.2

pH=5.2

pH=6.1

pH=7.6

pH=9.1

pH=11.2

Figure II-36: EPR spectra evolution with pH for Cu(II)-L1H4

For pH < 5, the signal evidence in solution the presence two copper species, in which

one corresponds to unreacted copper ion (Figure II-37).

2500 2750 3000 3250 3500

experimental simulation

****

Gauss

Figure II-37: EPR spectrum at pH 2: ‘*’ for Cu(H2O)6 and ‘’ for the second species

(solid line: experimental spectrum – dotted lines simulated spectrum with Xsophe)

139

Xsophe simulation allows to estimate the EPR parameters of the second copper

species, for which g// = 2.21, g┴ = 2.05, A// = 165×10-4 cm-1 and A┴ = 20×10-4 cm-1. These

values are close to the ones determined for CuL4H2 (Cu-DOTA) which seems to indicate that

this species, which corresponds to [CuL1H3]+ (see speciation diagram (Figure II-26.a,

page 129), could be hexacoordinated with a N4O2 coordination sphere.

For 5<pH<8, the signal shows a single system of four equidistant lines which

correspond to the formation of a single complex in solution. This signal corresponds to the

one shown in Figure II-35.a and is characteristic of a hexacoordinated copper ion (see

before).

At pH > 9, the spectrum is more complicated since the parallel system shows two

systems of four equidistant lines (Figure II-38, first set indexed by Δ, second set indexed by

), attributable to two species in solution.

2400 2600 2800 3000 3200 3400 3600 3800

Experimental Simulation

Gauss Figure II-38: EPR spectra of [CuL1]2- – pH 11.2) solid line: experimental spectrum and

dotted lines: simulated spectrum

On the other hand, Cu(II)-L1H4 potentiometric results of at pH > 9 indicate the

existence of a single stoichiometry for the species in solution. One can conclude that the

existence of two isomers for [CuL1]2- complex could account for the EPR signal. Similar

formation of isomers was previously observed with CuL6 system.[36]

140

For Cu(II)-L1H4 isomers, since EPR parameters are close to those of [CuL1H]-

complex, one can propose that both isomers are in an octahedral distorted ligand field. One

can propose that these isomers are configurational isomers for which the Cu(II) coordination

sphere N3O3 is differently distorted, due to the hindrance imposed by the functionalized

cyclen ring.

For Cu(II)-L2H3, the evolution of the spectra according to the pH (Figure II-39)

shows no change in the pH range 2-12.

2500 2700 2900 3100 3300 3500

Gauss

CuNO3

pH=3.3

pH=4.1

pH=6

pH=7.1

pH=9

pH=11

Figure II-39: EPR spectra evolution with pH for Cu(II)-L2H3 system

Therefore, it can be proposed that the complex exhibits a single geometry in solution

whatever the pH. This geometry as previously described for [CuL2]- (Figure II-35.b and

Table II-5, line 2) is octahedral, the copper coordination sphere being constituted of three

nitrogen atoms and three oxygen atoms (N3O3).

Table II-6: EPR data for Cu (II)-L1H4 system at pH > 9 (Xsophe simulation)

g// g┴ A// (10-4 cm-1) A┴ (10-4 cm-1) Ref

Cu(II)-L1H4 (isomer I) 2.31 2.08 135.0 20.0 This work

Cu(II)-L1H4 (isomer II) 2.26 2.08 130.0 20.0 This work

141

To resume, identical values obtained for Cu(II)-L1H4 spectrum at pH ~ 7 with that of

Cu(II)-L2H3 shows that the geometry of the complexes are similar. Hence, [CuL1H]- complex

is formed by N3O3 with imine nitrogen of benzimidazole involved in the complexation at

pH ~ 7 for Cu(II)-L1H4.

d) Structural hypotheses for M-L1H4 and M-L2H3 complexes, where M = Cu(II) and

Zn(II)

Deprotonation schemes of the transition metal complexes were proposed basing on

the above spectroscopic studies.

In Scheme II-8 are gathered deprotonation sequence of M-L1H4 systems

(with M = Cu(II) and Zn(II)).

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

[ML1H]- [ML1]2-

N N

N NCO2

-HO2C

-O2CHN

NH

M

[ML1H3]+

logK :Cu

Zn

4.3 4.5 9.2

4.2 5.1 9.7 Scheme II-8: Structural hypotheses for the deprotonation steps of Cu(II)-L1H4 and Zn(II)-

L1H4 complexes.

For Cu complexes, EPR spectra suggest an evolution of the coordination sphere from

N4O2 for [ML1H2] species to N3O3 for [ML1H]- species. This implies a decoordination of one

macrocyclic nitrogen atom in favour of one carboxylic oxygen atom. This ligand exchange

was already observed in cyclen substituted by a picolinate moiety.[37] UV-visible spectra

indicate that the final deprotonation step (between [ML1H]- species and [ML1]2- species)

implicates the deprotonation of the benzimidazole substituent. Therefore, the coordination

sphere between [ML1H]- species and [ML1]2- species is maintained to N3O3. This proposal is

corroborated by EPR spectroscopy since no major evolution is observed in the corresponding

pH range. Nevertheless, it must be mentioned that EPR spectra highlight the presence of two

isomers in solution for [ML1]2- species.

142

For the homologous zinc system Zn(II)-L1H4, we propose a similar deprotonation

sequence, the final step being highlighted by UV-visible spectroscopy.

In Schemes II-9 and 10, are reported deprotonation sequence of M-L2H3 systems

(with M = Cu(II) and Zn(II)).

N N

NH NCO2

--O2C

-O2CN

N

R

Cu pKa = 6.6

[CuL2H]

N N

N NCO2

--O2C

-O2CN

N

R

CuN N

N NCO2

--O2C

-O2CN

N

R

Cu

OH

[CuL2]- [CuL2H-2]

pKa = 8.9

Scheme II-9: Structural hypothesis for the deprotonation steps of Cu(II)-L2H3 complex

N HN

NH NCO2

--O2C

-O2CN

N

R

Zn pKa = 6.4

[ZnL2H2]+

N N

N NCO2

--O2C

-O2CN

N

R

Zn

N N

N NCO2

--O2C

-O2CN

N

R

Zn

OH

[ZnL2]-

[ZnL2H-2]

pKa = 10

N N

N NCO2

--O2C

-O2CN

N

R

Zn

HO OH

[ZnL2H-3]

pKa = 9.6

N N

NH NCO2

--O2C

-O2CN

N

R

Zn

ZnL2H

Scheme II-10: Structural hypothesis for the deprotonation steps of Zn(II)-L2H3 complex

Globally, the deprotonation sequences are similar for Cu(II) and Zn(II). UV-visible

spectroscopic studies of M-L2H3 systems, highlight that benzimidazole is not involved in the

deprotonation sequence and is coordinated to the metal for all the species. Furthermore, no

evolution with pH is observed for EPR spectra which indicate no evolution of the metal

coordination sphere on the whole pH range. The coordination sphere that can be deduced

143

from the EPR spectra indicate that the metal is hexacoordinated (MN3O3 species). Since the

deprotonation do not concern benzimidazole moiety one can propose that the first

deprotonation constant ([ML2H] [ML2]-) is associated to the deprotonation of one

macrocyclic nitrogen atom that is non coordinated to the metal. The next deprotonation

constants are associated to the deprotonation of one (for copper) or two (for zinc) water

molecule(s).

2. Physicochemical studies of lanthanide (Gd(III) and Eu(III)) complexes

In order to quantify the stability of Ln(III) complexes (Ln(III) = Gd(III), Eu(III)) with

L1H4 and L2H3, the overall complexation constants of these complexes and their protonated

forms were determined. In this section, overall stability constants of the metal complexes are determined by

potentiometry in the pH range 2-12, and compared with the constants of reference metal

complexes, M-L4H4 and M-L5H3 (M = Gd(III) and Eu(III)).


Potentiometric study of these lanthanide complexes are followed using the same

conditions of the transition metal complexes. The lanthanide(III) complexes are prepared at

various ratios of Ligand/Metal. For example, neutralization curves of Gd(III)-L1H4 are shown

in Figure II-40, where the variation in equivalence points is seen according to the ratio

[Ligand/Metal].

0,0 0,2 0,4 0,6 0,82

4

6

8

10

12

pH

V OH- / mL

L1H4 R = 1.15 R = 1.24

Figure II-40: Neutralization curves of Gd(III)-L1H4 solutions by NMe4OH (5×10-2 mol.L-1)

144

In Figure II-40, the equivalent points vary according to the different ratios (R = [L] /

[M]). The difference in-between the equivalent points correspond to the ratio of metal

solution added in each set of solutions.

b) Thermodynamic stability of lanthanide(III) complexes

Analysis of potentiometric data with PROTAF allows to propose the best chemical

model able to account for the experimental data. A good superposition of theoretical and

experimental curves is obtained, in which the formation in solution of [MLH2], [MLH] and

[ML] complexes with ligands L1H4 and L2H3 respectively are considered. The corresponding

overall stability constants and the deprotonation constants of lanthanide complexes are

gathered in Table II-7.

Table II-7: Stability constants of M/L1H4 and M/L2H3 complexes, (M = Gd(III) or Eu(III)),

compared with other metal complexes, T = 25°C, I = 0.1 (NMe4Cl) – Standard deviations are

given in parenthesis

L1H4 L2H3 L4H4 L5H3

Gd(III) Eu(III) Gd(III) Eu(III) Gd(III) Eu(III) Gd(III) Eu(III)

log mlh

M+L+2H+= MLH2 35.0

(5)

35.0

(4)

M+L+H+= MLH 32.1

(3)

30.9

(2)

23.6

(2)

24.5

(3)

M+L= ML 23.7

(4)

21.6

(2)

16.0

(1)

18.0

(2)

25.3 23.7 20.8 21.2

From the logmlh values, the deprotonation constants of both Gd(III) and Eu(III)

complexes are calculated.

Table II-8: Deprotonation constants of M/L1H4 and M/L2H3 complexes, ((M = Gd(III) or

Eu(III)).

log K L1H4 L2H3

Gd(III) Eu(III) Gd(III) Eu(III)

MLH2 = MLH + H 2.9 4.1 - -

MLH = ML + H 8.4 9.3 7.62 6.5

145

Basing on the above stability constants, the species distribution diagrams for L1H4 and

L2H3 with Gd(III) and Eu(III) were traced using the HYSS software. [24]

a. b.

2 4 6 8 10 120

20

40

60

80

100

L1H4

L1H5+

L1H73+

L1H62+

[GdL1]-[GdL1H]

[GdL1H2]+

Gd3+

%G

d

pH 2 4 6 8 10 120

20

40

60

80

100[GdL2][GdL2H]+

Gd3+

%G

d

pH

Figure II-41: Species distribution curves of a. Gd(III)-L1H4 and b. Gd(III)-L2H3 systems.

[GdL1H2]+ species begins to form at pH = 2 and the other protonated species [GdL1H]

prevails between pH 3 and pH = 10. Beyond pH 9, the deprotonated [GdL1]- complex

becomes predominant. For Gd(III)-L2H3 system, protonated species [GdL2H]+ forms at

pH = 2. The deprotonated [GdL2] complex prevails from pH 6 to 12.

Europium complexes Eu(III)-L1H4 and Eu(III)-L2H3 exhibit similarities in species

distribution according to pH with Gd(III)-L1H4 and Gd(III)-L2H3 complexes (Figure II-42).

a. b.

2 4 6 8 10 120

20

40

60

80

100

[EuL1]-[EuL1H][EuL1H2]

+

L1H5+

L1H6

2+

Eu3+

% E

u

pH 2 4 6 8 10 12

0

20

40

60

80

100[EuL2][EuL2H]+

Eu3+

%Eu

pH

Figure II-42: Species distribution curves of a. Eu(III)-L1H4 and b. Eu (III)-L2H3 systems.

[EuL1H2]+ species begins to form at pH = 2 and the other protonated species [EuL1H]

prevails between pH 2 and pH = 11. Beyond pH 9, the deprotonated [EuL1]- complex

146

becomes predominant. For Eu(III)-L2H3 system, protonated species [EuL2H]+ prevails from

pH = 2. The deprotonated [EuL2] complex prevails from pH 6 to 12.

To compare L1H4 and L2H3 affinities for Gd(III) and Eu(III), the logarithm of the

percentage of free metal ion (log ([M]free/[M]total)) in function with pH was traced (Figure II-

43).

a. b.

-12

-10

-8

-6

-4

-2

0

L5H3

L2H3

L4H4

L1H4

pH

log(

[Gd]

free/[

Gd]

tota

l)

2 4 6 8 10 12

-12

-10

-8

-6

-4

-2

0pH

log(

[Eu]

free/[

Eu] to

tal)

2 4 6 8 10 12

L5H3

L2H3

L4H4

L1H4

Figure II-43: Logarithm of percentage of free metal ion concentration of L1H4, L4H4 and

L5 H3 in relation to Ln(III) equal to a. Gd(III), b. Eu(III) ([ligand] = [Ln] = 2.010-3

mol.L-1)

Gd(III)-L1H4 curve is enclosed between Gd(III)-L4H4 one and Gd(III)-L5H3, and its

position indicates that the affinity of L4H4 for Gd(III) is intermediate between L4H4 (DOTA

ligand, which possesses the higher affinity) and L5H3 (DO3A derivative, which possesses a

lower affinity). In consequence for Gd(III)-L1H4 system, the involvement of nitrogen of

imidazole group in Gd(III) coordination counterbalance the lack of one carboxylic arm (by

comparison to DOTA) and improves the complex stability (by comparison to the DO3A

backbone in Gd(III)-L5H3). Moreover, affinity of L1H4 for Gd(III) is better to the one of L2H3

especially for pH superior to 6.

Similar behaviour follows in the Eu(III)-L1H4 complex, which was expected due to

the analogy between Gd (III) and Eu (III) ions.

147

c) Spectroscopic study of Gd (III) and Eu(III) complexes

i. UV – visible studies

The hypotheses of deprotonation of complexes can be monitored using UV

spectroscopy as a function of pH basing on the presence of benzimidazole probe (Figure II-

44).

a. b.

260 280 300 320 3400,0

0,1

0,2

0,3

0,4

0,5

Abso

rban

ce

/ nm

pH = 3.05 pH = 6.42 pH = 8.53 pH = 10.20 pH = 11.58

2 4 6 8 10 120

20

40

60

80

100

LH5+

LH62+

[GdL]-[GdLH]

[GdLH2]+

Gd3+

%G

d

pH7800

8000

8200

8400

8600

8800

9000

(m

ol-1 L

cm

-1)

Figure II-44: a. Evolution of UV signal of Gd(III)-L1H4 in function of pH. b. 278 (mol-1 L

cm-1) = f (pH) followed on the species differentiation diagram. ([Gd(III)-L1H4] = 6.25×10-5

mol.L-1)

Significant changes are seen in the spectrum around pH = 9, where a supplementary

band at = 284 nm is formed (Figure II-44.a). As seen earlier in the case of copper and zinc

complexes (see in paragraph 2, page 131), formation of a new band around = 282 nm

indicate the formation of benzimidazolate ion. Besides, This proposition was verified by

superimposing the evolution of the epsilon values at = 278 nm according to the pH on the

speciation diagram with existence of [GdL1H] and [GdL1]- species in majority. The

absorbance superposed quite well with the formation of [GdL1]- species.

Europium complex exhibits similar behaviour as gadolinium complex. Formation of

benzimidazolate ion is seen after pH = 9 (supplementary band at = 284 nm) (Figure II-45).

148

a. b.

260 280 300 320 3400,0

0,1

0,2

0,3

0,4

0,5A

bsor

banc

e

/ nm

pH = 3.60 pH = 6.75 pH = 8.91 pH = 9.80 pH = 10.94 pH = 11.64

2 4 6 8 10 12

0

20

40

60

80

100[EuL1H] [EuL1]-

[EuL1H2]+

Eu3+

L1H62+

L1H5+

%E

u

pH8400

8600

8800

9000

9200

9400

(m

ol-1 L

cm

-1)

Figure II-45: a. Evolution of UV signal of Eu(III)-L1H4 in function to pH. b. 278 (mol-1 L

cm-1) = f(pH) followed on speciation diagram. ([Eu(III)-L1H4] = 6.25×10-5 mol.L-1)

Evolution of metal complexes (M-L2H3, where M = Gd(III) and Eu(III)) with pH are

monitored in the UV range using UV spectroscopy (Figure II-46).

a. b.

200 300 4000,0

0,5

1,0

1,5

Abso

rban

ce

(nm)

pH = 2.94 pH = 3.28 pH = 4.23 pH = 5.76 pH = 6.98 pH = 7.72 pH = 8.72 pH = 9.76 pH = 10.98 pH = 11.47

200 300 4000,0

0,5

1,0

1,5

Abso

rban

ce

(nm)

pH = 3.31 pH = 4.43 pH = 6.21 pH = 7.38 pH = 8.48 pH = 9.77 pH = 11.07

Figure II-46: a. Evolution of UV signal of Gd(III)-L2H3 in function of pH. b. Evolution of UV

signal of Eu(III)-L2H3 in function of pH

Similar results were obtained for Eu(III)-L2H3 complex. As depicted by the Figure II-

46, no significant change in absorbance is seen which provides evidence that no

benzimidazolate formation is reached in the relevant L2H3 metal complexes.

ii. Fluorescence measurements

Fluorescence experiments were conducted at University of Hull for europium

complexes, to determine the number of water molecules in the lanthanide coordination

149

sphere. Luminescence spectra of europium complexes with L1H4 were taken at either side of

the pKa 9.3. The lifetimes of europium complexes and their associated q values were

calculated.

Emission spectra of the complexes were recorded in aqueous solutions with

excitation at 273 nm via the benzimidazole-centred triplet state.

a. b.

Figure II-47: Luminescence spectra of Eu(III) complex at pH 11.86 in H2O. a.

Excitation spectrum at 616 nm emission. b. Emission spectrum at 265 nm excitation with

a 10 nm slit width (the red line represents the emission spectrum at 265 nm excitation

with a 5 nm slit width).

Europium(III) has a 5D0 excited state and the complex showed a typical emission

spectrum corresponding to the transitions )4n0(FD n7

05 , where the transition

07

05 FD corresponds to λem = 575, 1

70

5 FD to 592 nm, 27

05 FD to 616 nm, 3

70

5 FD

to 655 nm and 47

05 FD to 699 nm.

150

Figure II-48: Emission spectra of Eu(III) complex of L2H3

A better knowledge of the number of water molecules coordinated to the metal ion is

important in interpreting the reactivity of metal complexes in solution. For that, the

comparison of the Eu complex lifetime in water and in D2O was undertaken. The number of

water molecules bound to the hydrated lanthanide ion was then calculated using the initial

expression of Horrocks and Sudnick:[38]

)kk(Aq ODobs

OHobsLn

22 (2)

where OHobs

2k and ODobs

2k are the rate constants of deexcitation of europium in H2O and in D2O,

ALn is the proportionality factor for a given lanthanide.

In this case, the ALn value of europium (AEu = 1.05) was taken into consideration basing on

the literature review.[38, 39]

The luminescence intensity measured for various H2O/D2O solutions where their ratio

ranging from 0-1, decreases exponentially with time. The obtained curves were fitted using

the equation (3) to determine the observed rate constant kobs.

)t.kexp(.II obs0 (3)

Table II-9 represents the rate constants of deexcitation of lanthanide in H2O and in

D2O.

151

Table II-9: Lifetimes and their associated q values of Eu(III)-L1H4 and Eu(III)-L2H3

pH Predominant

complex

H2O / ms D2O / ms q

7.34 [EuL1H] 0.62 2.15 1.1

11.86 [EuL1]- 0.56 1.96 1.2

>7.0 [EuL2] 2.28 1.04 1.30

The above q values show that above pH 7, [EuL1H] and [EuL1]- complexes

possess both one water molecule in their inner coordination sphere. Therefore, the

deprotonation event occurring between [EuL1H] and [EuL1]- could not be due to water

deprotonation. For [EuL2] complex, one water molecule is present in the inner

coordination sphere.

d) Structural hypotheses for Ln-L1H4 and Ln-L2H3 complexes(where Ln = Gd(III)

and Eu(III))

Taking into account all the previous results, structural hypotheses could be formulated

for Ln-L1H4 and Ln-L2H3 complexes.

The spectroscopic results give the confirmation of benzimidazole deprotonation in

Ln-L1H4 complex. The deprotonation of benzimidazole of ligand alone (L1H4) is at

log K = 12.5, whereas in the lanthanide complexes it’s decreased log K values (inbetween

8~9) or increased acidic nature is due to the involvement of benzimidazole in complex

formation.

Gd

Eu

3.0

4.1

8.4

9.3

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]+

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

OH2

[ML1H] [ML1]-

OH2H2O OH2

Gd

Eu

3.0

4.1

8.4

9.3

N N

N NCO2

--O2C

-O2CHN

NH

M

[ML1H2]+

N N

N NCO2

--O2C

-O2CHN

NM

N N

N NCO2

--O2C

-O2CN

NM

OH2

[ML1H] [ML1]-

OH2H2O OH2

Scheme II-11: Structural hypotheses for the deprotonation steps of the complexes Gd(III)-

L1H4 and Eu(III)-L1H4.

152

For L2H3, we proposed a hypothesis (Scheme II-12) in which log K= 7.62 for

Gd(III)-L2H3 and log K = 6.5 for Eu(III)-L2H3 are assigned to the deprotonation of the

nitrogen atom of macrocycle.

Gd

Eu

7.62

6.5

N N

NH NCO2

--O2C

-O2CN

N

R

M

[ML2H]+

N N

N NCO2

--O2C

-O2CN

N

R

M

[ML2]

H2OOH2H2O

Gd

Eu

7.62

6.5

N N

NH NCO2

--O2C

-O2CN

N

R

M

[ML2H]+

N N

N NCO2

--O2C

-O2CN

N

R

M

[ML2]

H2OOH2H2O

Scheme II-12: Structural hypotheses for the deprotonation steps of the complexes Gd(III)-

L2H3 and Eu(III)-L2H3.

e) Transmetallation with Zn(II) – Relaxometric measurements

Zn(II) is one of the most abundant endogenous metal ions with a concentration of ~32

μM in the human plasma.[40] Therefore, this ion can be an exchanging ion for Gd(III)

involved in gadolinium contrast agents (GdL). The corresponding exchanging reaction can be

considered as follows:

GdL + Zn = ZnL + Gd (4)

In this case, from the thermodynamic point of view, the percentage of free metal ion

(log ([M]free/[M]total)) in function of pH for Gd(III)-L1H4 and Zn(II)-L1H4 systems were

calculated (Figure II-49).

-12

-10

-8

-6

-4

-2

0

log(

[M] fre

e/[M] to

tal)

Gd(III)-L1H4

Zn(II)-L1H4

2 4 6 8 10 12

Figure II-49: Logarithm of percentage of free metal ion concentration of L1H4 for Gd(III)

and Zn(II) ([ligand] = [Gd] = [Zn] = 2.010-3 mol.L-1)

153

Effectively, the curve shows that L1H4 has a better affinity for Gd(III) on the whole

pH range. However, the difference between the stability of these two systems is not immense.

Hence, to determine the chemical inertness of Gd(III)-L1H4 system in the presence of Zn(II)

ions, a transmetallation experiment was carried out in phosphate buffer (pH 7.4).

Replacement of Gd(III) in the complex with Zn(II) leads to free gadolinium ions, which

precipitate in the presence of phosphate as GdPO4. On the course of the experiment, the total

amount of paramagnetic species in solution should decrease because the precipitated GdPO4

could no more be involved in the relaxation process. Therefore the decrease in relaxation

allows following the transmetallation reaction.

In Figure II-50, the evolution with time of the ratio of relaxation rates R1(t) / R1(t =

0), where R1 is the relaxation rate at time t and R1(t = 0) is the relaxation rate at time zero

(just before addition of Zn(II)), is shown for Gd(III)-L1H4 together with the evolution of

Gd(III)-L4H4 for comparison.

0 1000 2000 3000 4000 50000,0

0,2

0,4

0,6

0,8

1,0

1,2

R1 t /

R1 0

t (min)

Gd-L1H4

Gd-L4H4

Figure II-50: Evolution of R1(t)/ R1(t = 0) (T= 310 K; pH 7.0) versus time for (■) Gd(III)-

L1H4 and (●) Gd(III)-L4H4 (Gd-DOTA) in the presence of equimolar amounts of Zn(II) ions

in phosphate buffer solution.

No decrease in the relative relaxation rates is seen for Gd(III)-L1H4.

Ligand L2H3 having some solubility issues in water for concentrations required for

transmetallation experiments, the comparison of its thermodynamic and kinetic index with

the other metal complexes could not be done.

For Gd(III)-L1H4, according to Laurent et al.[41] two characteristic values can be used

to describe the behavior of a Gd(III) chelate in a transmetallation experiment, i.e. the time to

reach R1(t)/R1(t = 0) = 0.8 (ratio index) which gives information about the kinetics of the

154

reaction, and the R1(t = 4320) / R1(t = 0) value at very long time, (considered after 3days =

4320 min) reflecting the thermodynamic aspect of the system.

Table II-10: Transmetallation with Zn(II): time required to reach the ratio R1(t)/ R1(t = 0) =

0.8 and value of R1(t) (t = 4320) / R1(t = 0) for Gd(III)-L1H4 and Gd(III)-L1H4 (37° C, pH 7).

t for R1(t) / R1(t = 0) = 0.8 [min]

R1(t) (t = 4320 min) / R1(t = 0)

Gd-L1H4 > 5000 0.99 This work

Gd-DOTA > 5000 0.99 [42]

From the values in Table II-10, it can be concluded that no transmetallation reaction

occured in Gd(III)-L1H4 complex with Zn(II) ions.

155

D. Conclusion In this chapter we described the complexation of two DO3A ligands functionalised by

methyl-benzimidazole groups (named L1H4 and L2H3 and synthesized in Pr. S. Archibald

group), by a set of metal ions (both transition metal ions such as Cu(II) and Zn(II) and

lanthanide ions such as Gd(III) and Eu(III)).

The goal of this work was first to propose a coordination scheme for the different

L1H4 and L2H3 systems and to compare their respective stabilities. For L1H4 complexes,

whatever the metal ion, the benzimidazole group participates to the metal ion coordination

sphere as soon as the benzimidazolium was deprotonated. For L2H3 complexes, the

complexation sequence is not similar since the benzimidazole moiety is coordinated even at

low pH, the deprotonation observed in this range being attributable to deprotonation of a

nitrogen atom un-coordinated to the metal ion. Furthermore for these systems, hydroxylated

complexes were proposed at higher pHs. A comparison of complex stabilities was also

proposed by means of the calculation of the logarithm of the percentage of free metal ion

(log([M]free/[M]total)) in function of pH for L1H4 and L2H3. Whatever the metal ion,

complexes based on L1H4 backbone were the most stable. It should however be noticed that

these complexes are generally less stable than their analogues based of DOTA backbone.

This indicated that the nitrogen of imidazole group (from benzimidazole) is a less good donor

atom than anionic oxygen (of carboxylate in DOTA).

The goal of this work was secondly to evaluate the kinetic inertness of L1H4 Gd(III)

complex. Transmetallation experiments of Gd(III)-L1H4 complex in the presence of Zn(II)

was carried out in phosphate buffer (pH 7.0) and followed by means of relaxometry. It was

expected that if transmetallation occurred, a decrease in relaxivity would have been detected.

For Gd(III)-L1H4, under these conditions, no decrease in the relative relaxation rates was

seen. This behaviour, similar to the one obtained for Gd(III)-L1H4 complex is typical of

macrocyclic complexes for which, the ligand pre-organisation and the number of ligand

donor atoms are perfectly fitted for the Ln(III) coordination sphere. Consequently, no

transmetallation reaction occurred in Gd(III)-L1H4 complex in the presence of Zn(II) ions.

This encouraging result is important and interesting, in the view of Gd(III)-L1H4 utilization as

a contrast agent in MRI applications.

156

References

1. F.Mani, R. Morassi, P. Stoppioni, A. Vacca, J. Chem. Soc. Dalton Trans. 2001, 2116–

2120.

2. M. Di Vaira, F. Mani, S. S. Costantini, P. Stoppioni, A. Vacca, Eur. J. Inorg. Chem.

2003, 3185–3192.

3. E. Tamanini, A. Katewa, L. M. Sedger, M. H. Todd, M. Watkinson, Inorg. Chem.

2009, 48, 319–324.

4. E. Tamanini, K. Flavin, M. Motevalli, S. Piperno, L. A. Gheber, M. H. Todd, M.

Watkinson, Inorg. Chem. 2010, 49, 3789–3800.

5. K. Jobe, C. H. Brennan, M. Motevalli, S. M. Goldup, M. Watkinson, Chem. Commun.

2011, 47, 6036–6038.

6. A. E. Majzoub, C. Cadiou, I. Déchamps-Olivier, F. Chuburu, M. Aplincourt, B.

Tinant, Inorg. Chim. Acta. 2009, 362, 1169–1178.

7. A. E. Majzoub, C. Cadiou, I. Déchamps-Olivier, F. Chuburu, M. Aplincourt, Eur. J.

Inorg. Chem. 2007, 5087–5097.

8. A. E. Majzoub, C. Cadiou, I. Déchamps-Olivier, B. Tinant, F. Chuburu, Inorg. Chem.

2011, 50, 4029–4038.

9. M. Pan, X. L. Zheng, Y. Liu, W. S. Liu, C. Y. Su, Dalton Trans. 2009, 2157–2169.

10. N. M. Shavaleev, S. V. Eliseeva, R. Scopelliti, J. C. G. Bünzli, Chem. Eur. J. 2009,

15, 10790–10802.

11. A. Beeby, L. M. Bushby, D. Maffeo, J. A. Gareth Williams, J. Chem. Soc. Dalton

Trans. 2. 2002, 48–54.

12. J. Bhumasamudram, G. L. Brickert-Albrecht, G. S. Nichol, A. E. Mash, R. Natarajan,

Tetrahedron Lett. 2011, 52, 2058–2061.

13. A. El Majzoub, C. Cadiou, I. Déchamps-Olivier, F. Chuburu, M. Aplincourt.

Eur. J. Inorg. Chem. 2007, 5087–5097.

14. G. R. Weisman, S. C. H. Ho, V. Johnson, Tetrahedron Letters. 1980, 21, 335–338.

15. A. Khan, G. Nicholson, J. Greenman, L. Madden, G. McRobbie, C. Pannecouque, E.

DeClercq, S. J. Archibald, et al, J. Am. Chem. Soc. 2009, 131, 3416–3417.

16. R. Smith, D. Huskens, D. Daelemans, R. E. Mewis, C. D. Garcia, A. N. Cain, T. N. C.

Freeman, et al, Dalton Trans. 2012, 41, 11369–11377.

17. A. E. Sparke, C. M. Fisher, R. E. Mewis, S. J. Archibald, Tetrahedron Lett. 2010, 51,

4723.

http://www.sciencedirect.com/science/article/pii/S0020169308004271






157

18. R. Fournaise, C. Petitfaux, Talanta., 1987, 34, 385–395.

19. A. Bianchi, L. Calabi, C. Giorgi, P. Losi, P. Mariani, P. Paoli, P. Rossi, B. Valtancoli,

M. Virtuani, J. Chem. Soc. Dalton Trans. 2000, 697–705.

20. E. Tóth, R. Király, J. Platzek, B. Radüchel, E. Brücher, Inorg. Chim. Acta. 1996, 249,

191–199.

21. J. Catalan, J. L. G. De Paz, M. Yanez, F. Amat-Guerri, R. Houriet, E. Rolli, R.

Zehringer, P. Oelhafen, R. W. Taft, F. Anvia, J. H. Qian, J. Am. Chem. Soc. 1988,

110, 2699–2705.

22. H. Walba, R.W. Isensee, J. Org. Chem. 1961, 26, 2789–2791.

23. L. J. Thomas, K. P. Quinlan, J. Am. Chem. Soc. 1960, 82, 2994–2997.

24. P. Gans, A. Sabbatini, A. Vacca, Hyperquad simulation and Speciation. 1998-2000,

Protonics software

25. P. N. Preston, ed. Chemistry of Heterocyclic Compounds: Benzimidazoles and

Cogeneric Tricyclic Compounds, Part 1, Volume 40. Chemistry of Heterocyclic

Compounds: A Series Of Monographs, 2008.

26. M. Krishnamurthy, P. Phaniraj, S.K. Dogra, J. Chem. Soc. Perkin Trans. 2. 1986,

1917–1925.

27. a. M. Regueiro-Figueroa, D. Esteban-Gómez, C. Platas-Iglesias, A. de Blas, T.

Rodríguez-Blas, Eur. J. Inorg. Chem. 2007, 2198–2207. b. G. Berden, W. L. Meerts,

E. Jalviste, J. Chem. Phys. 1995, 103, 9596. c. J. R.Platt, J. Chem. Phys. 1951, 19,

101-118.

28. J. Moreau, E. Guillon, J. C. Pierrard, J. Rimbault, M. Port, M. Aplincourt, Chem. Eur.

J. 2004, 10, 5218–5232.

29. M. Kruppa, D. Frank, H. Leffler-Schuster, B. König, Inorg. Chim. Acta. 2006, 359,

1159–1168.

30. I. Sóvágó, T. Kiss, A. Gergely, J. Chem. Soc. Dalton Trans. 1978, 964–968.

31. M. Griffin, A. Muys, C. Noble, D. Wang, C. Eldershaw, K. E. Gates, K. Burrage, G.

R. Hanson, Mol. Phys. Rep. 1999, 26, 60.

32. R. Delgado, J. J. R. Fraústo da Silva, M. C. T. A. Vaz, Talanta. 1986, 33, 285–287.

33. K. Miyoshi, H. Tanaka, E. Kimura, S. Tsuboyama, S. Murata, H. Shimizu, K. Ishizu,

Inorg. Chim. Acta. 1983, 78, 23–30.

34. T. D. Smith, A. E. Martell, J. Am. Chem. Soc. 1972, 4123-4129.

35. H. Kurosaki, C. Bucher, E. Espinosa, J. M. Barbe, R. Guilard, Inorg. Chim. Acta.

158

2001, 322, 145–149.

36. R. Delgado, S. Quintino, M. Teixeira, A. Zhang, J. Chem. Soc. Dalton Trans. 1997,

55–64.

37. L. M. P. Lima, D. Esteban-Gómez, R. Delgado, C. Platas-Iglesias, R. Tripier, Inorg.

Chem. 2012, 51, 6916-6927.

38. W. D. Horrocks, R. D. Sudnick, Acc. Chem. Res. 1981, 14, 384–392.

39. A. Beeby, M. I. Clarkson, S. R. Dickins, S. Faulkner, D. Parker, L. Royle, S. A. de

Sousa, A. J. Gareth Williams, M. Woods, J. Chem. Soc. Perkin Trans. 2. 1999, 493–

504.

40. S. Laurent, L. Vander Elst, F. Copoix, R. N. Muller, Invest. Radiol. 2001, 36, 115-

122.

41. S. Laurent, L. Vander Elst, R. N. Muller, Contrast Med. Mol. Imaging. 2006, 1, 128-

137.

42. R. S. Dickins, S. Aime, A. Batsanov, A. Beeby, M. Botta, J. I. Bruce, J. A. K.

Howard, C. S. Love, D. Parker, R. D. Peacock and H. Puschmann, J. Am. Chem. Soc.,

2002, 124, 12697-12705.

159

Chapter-III Physico-chemical studies of linear ligands and their metal complexes; investigation of transmetallation

mechanisms near physiological pH

161

A. Interest of nanoparticles(Np) as contrast agents

Linear polyamino carboxylic acids gained their importance in MRI by being used as

contrast agents. DTPA (Diethylene Triamine Pentaacetic Acid named as L@4H5 in this

chapter) is used as a parent molecule mostly for developing novel linear ligands in the

context of MRI. Despite of its use as a contrast agent in MRI, its non-ability to detect the very

small tumors is explained by a lack of sensitivity that is inherent to the MRI process, which

detects the difference between malignant and normal tissue. Novel contrast agents based on

nanoparticles are being actively investigated as next generation of smart contrast agents for

magnetic resonance imaging due to their ability to function at the cellular and even at the

molecular level of biological interactions.[1] Furthermore, these magnetic nanoparticles are

very attractive platforms as contrast agents, besides as carriers for drug delivery.[2, 3, 4] In

addition, they offer the opportunity to incorporate in a single nano-object, at least two

synergistic imaging functionalities.[5] The resulting multimodal nanoparticles are very

attractive because they provide a more complementary, effective, and accurate information

about the physical, anatomical structure, and the physiological function for diagnosis and

clinical settings.

1. Nanoparticle based contrast agents in the context of MRI.

Gd chelates associated with macromolecular species or nanosystems, offering a

number of advantages compared to conventional contrast agents secured their importance in

MRI.[5] As examples, Gd complexes covalently attached to macromolecules or

macromolecular species such as dendrimers, micelles, liposomes, nano-emulsion, carbon

nanotubes can be cited (FigureIII-1),[5,6-8] Those systems have the advantage to ensure a

high gadolinium payload and to increase the rotational correlation time of metal chelates.

Consequently, relaxivity enhancement can be reached provided that the link between Gd

chelate and the macromolecule or the nano-object is rigid. This point is of particular

importance for Gd chelates grafted onto linear polymers or at dendrimer surface,[6b] to avoid

individual chelate motion that break away from the overall system motion and be deleterious

to the relaxivity.

162

Figure III-1: Schematic representation of Gd-complexed species (i.e dendrimers, liposomes,

carbon nanotubes etc). [6]

The nanoparticle-based contrast agents can be classified into T1-weighted MRI

contrast agents and T2-weighted MRI contrast agents. The ability of nanoparticles to enhance

proton relaxation of specific tissues and serve as magnetic resonance imaging contrast agents

have been actively investigated mainly as T2-weighted contrast agents in the form of

superparamagnetic iron oxides (SPIO) for over two decades, with applications such as

liver/spleen imaging (i.e., Endorem® and Feridex IV®).[9-11] Gadolinium based nanoparticles

are being investigated mainly as T1-weighted contrast agents. Metal nanoparticles such as

gold nanoparticles as templates for MRI contrast agents are of interest due to their highest

stability among metal nanoparticles, low toxicity and strong resistance against oxidation. In

this section, the discussion will focus on recent advances in the field of gold nanoparticle

based contrast agents in diagnosis and therapy (or theranostics).

2. Gold Nanoparticles and Gd chelates Gold nanoparticles have a wide range of applications in medical science and in

catalysis because of their non-toxic behaviour and good resistance to oxidation. They are easy

to synthesize, mainly by reduction of AuCl4- in the presence of thiols as stabilizers.[12] Their

synthesis can also be performed in aqueous solution in the presence of reducing agents and

water-soluble polymers, surfactants or capping agents with the aid of externally supplied

energy.[13] These methods allow for adequate control of the size and concentration of the

dispersed particles.

163

The contrast enhancement studies revealed that gold due to high atomic number (Z)

and greater absorption coefficient than iodine can be used as contrast agent in CT imaging.

Study of Gold nanoparticles for its therapeutic and theranostic applications[14] is focussed

presently in the field of medicine. We will focus our discussion on the recent reports of gold

nanoparticle based multimodal contrast agents.

S.Roux et al.[15] pioneered the design of gold nanoparticles for the purpose of both

MRI and CT (Computed Tomography) contrast agents. They synthesized dithiolated-DTPA

multilayered nanoparticles that could accommodate 150 gadolinium ions per nanoparticle

(Figure III-2). These multilayered nanoparticles are with diameter 2-2.5 nm. The relaxivity

values for protons in the presence of these Gd-coated gold nanoparticles (r1 = 3.9 mM -1 s-1)

are similar to the one determined in the presence of DTPA:Gd chelates (r1 = 3.0 mM-1s-1).

This limited enhancement was explained by the presence of relatively mobile Gd-chelates at

gold nanoparticle surface. Besides, the multilayered nanoparticles were not sufficiently stable

at room temperature. In their second report[16] regarding Gd-loaded gold nanoparticles, the

authors explained that the stability of these nanoparticles can be enhanced by decreasing the

number of gadolinium ions per particle from 150 to 50.

Figure III-2: Multilayered Gd-loaded gold nanoparticles.[15]

From the X-ray imaging point of view, the enhancement of contrast in CT imaging is

observed for lower gold concentration (10 mg.mL-1). Besides, the contrast enhancement is

greater than for iodinated compounds, highlighting the interest of gold nanoparticles as

contrast agents for CT.

164

Park et al. reported a similar approach and proposed AuNPs coated with Gd-complex

of DTPA-bis(amide) conjugate of gluthathione as MRI contrast agents (Figure III-3).

Figure III-3: Gold nanoparticles coated with gluthathione or cysteine conjugates of Gd-

DTPA coating.[17 , 18]

The corresponding particles have a mean size of 5-7 nm and a total number of

gadolinium complexes per nanoparticle to be about 1.34×104.[17] The relaxivity of these

nanoparticles is (10.2 mM-1s-1 at 293 K and 1.5 T) three fold higher than for the commercially

available MRI contrast agent (Omniscan : 3.30 mM-1s-1). Such a high relaxivity may be

explained in terms of of slower tumbling motion of Gd complexes grafted on gold

nanoparticles.

Bigger gold nanoparticles (~ 14 nm) coated by Gd-DTPA conjugates of cysteine as

multimodal contrast agents were reported by the same authors (Figure III-3).[18] A total

number of gadolinium complexes per nanoparticle of about 2.9×103 was reported with a

relaxivity per Gd center of 17.9 mM-1s-1 (293 K, 1.5 T). This relaxivity, five fold higher than

for the commercially available MRI and isostructural contrast agent Omniscan (3.30 mM-1s-

1), may be explained in terms of of slower tumbling motion of Gd complexes grafted on gold

nanoparticles.

The cytotoxic studies showed that these nanoparticles are non-toxic and hence are

suitable for practical applications.

Recently, Moriggi and coworkers[19] reported the synthesis of ~ 5.0 nm DTTA thiol-

functionalized gold nanoparticles. The ligand grafted at the surface of the nanoparticle is a

DTPA derivative (Figure III-4.a).

165

a. b.

Figure III-4: a. Gd-DTTA thiol functionalized gold nanoparticles and b. 1H NMRD profile

(298K) [19]

The contribution of paramagnetism from the gold core to the magnetic moment of

Gd(III) was explored by these authors. They reported that gold nanoparticles do not

contribute significantly to the magnetic moment of Gd(III) ions grafted at the gold

nanoparticle surface and then to the overall relaxivity of surrounding water protons. The 1H

nuclear magnetic relaxation dispersion profile (NMRD) of these nanoparticles (Figure III-

4.a) showed a very high relaxivity maximum of ~ 60 mM-1 s-1 at 30 MHz. The NMRD fitting

highlighted that this remarkable enhancement has to be correlated with a slow rotational

motion of the Gd chelates. Moreover, the dense packing of the Gd chelates on the surface of

the nanoparticle (Figure III-4.a) rendered them rigid, which was beneficial for the relaxivity

enhancement.

Modification of the hydration sphere of gold nanoparticles can also be beneficial to

rhe relaxivity enhancement. This modification can be addressed by post-coating with sugar

conjugates of β-glucose (glycC2S and glycC5S), β-galactcose (galC5S), and β-lactose

(lacC5S).[20] (Figure III-5)

166

Figure III-5: Thiol-ending sugar conjugates and the corresponding paramagnetic

glyconanoparticles.[20]

The relaxivities of most of these paramagnetic glyconanoparticles were higher than

for the corresponding small molecule Gd-chelate (10< r1 < 25 mM-1 s-1).

Coating Gd-functionalised gold nanopaticles with polyelectrolyte by electrostatic

interactions were also used by Warsi et al[21] to reduce the tumbling of Gd-chelates bound to

nanoparticle, thereby enhancing the relaxivity. The overall relaxivity increase obtained for

polyelectrolyte-coated nanoparticles was well over 80% as compared to Gd-DTPA.

When the nanoparticles are functionalized with a targeting agent they can be

addressed to specific cells and thereby their efficiency for diagnostic can be improved. Such

kind of nanoparticles designed to target epidermal growth factor receptors (EFGR)

overexpressed on the surfaces of breast cancer cells were reported by Lim et al.[22] The

principle of this recognition is to graft at the surface of the gold nanoparticle a monoclonal

antibody that specifically recognizes receptors at the cell surface. Here monoclonal anti-

HER2 was chemically conjugated with Gd-DTPA and the combination grafted on the

nanoparticle surface (Figure III-6).

167

Figure III-6: Targeted paramagnetic gold nanostructures.[21]

These nanoparticles with Gd(III) and cancer cell targeting units were evaluated for both MRI

and optical imaging purpose. The relaxivity values of r1 and r2 for paramagnetic for these

nanoparticles were 23.7 mM-1 s-1 and 89.5 mM-1 s-1, respectively. One more time the

conjugation of GdDTPA chelate on antibody and its strong anchorage at the gold

nanoparticle surface is in favour of gain in relaxivity. This relaxivity enhancement is

translated in MR images of SKBR3 breast cancer that over-expressed EFGR receptors.

Furthermore, these gold nanostructures due to their absorption in the NIR region were used to

kill breast cancer cells by conversion of the absorbed radiation into thermal energy.

Recently, gold nanoparticles coated with a Gd-DTPA conjugate of penicillamine-

chelate were reported[23] (Au@GdL, Figure III-7). Average size of Au@GdL is 12 nm with

loading of GdL reaching up to 3.0×103 per particle. They exhibit a r1 relaxivity per Gd of

20.1 mM–1 s–1 and good X-ray attenuation. Due to their small size, these nanoparticles also

exhibited a contrast enhancement in the liver, blood pool and lymph node.[24]

Figure III-7: Structure of Au@GdL (L = DTPA–penicillamine conjugate).[23]

Due to their small size, these nanoparticles also exhibited a contrast enhancement in

the liver, blood pool and lymph node.[24]

168

To resume, in all reports of gold nanoparticle based contrast agents, it has been found

that several gadolinium ions can be loaded per particle and these contrast agents gave better

relaxivity and a good T1 contrast in MRI images. In many reports[15-18] multilayers of Gd-

chelates were formed at nanoparticles surface, which might not favour the stability issues of

these nanostructures. Hence, we are interested in studying the physicochemical properties of

the dithiolated-DTPA (named L@1H5) grafted multilayered gold nanoparticles (named

L@2H3)[15] to understand their thermodynamic stability and kinetic inertness over competitive

endogenous cations.

This chapter mainly deals with the determination of thermodynamic stabilities of

dithiolated-DTPA (L@1H5), dithiolated-DTPA (L@

1H5) grafted multilayered gold

nanoparticles (L@2H3) along with the commercially available DTPA (L@

4H5) to compare and

understand their stabilities. Kinetic inertness of Gd-L@1H5 grafted gold nanoparticles will be

evaluated in the presence of endogenous Zn(II) ions. The investigation of transmetallation

mechanism near physiological pH will also be presented.

B. Synthesis, physicochemical studies of ligands (L@1H5 and L@

2H3) and their metal complexes (Cu(II), Zn(II), Ca(II), Na(I) and Gd(III))

1. Synthesis and physicochemical studies of ligands

a) Synthesis of L@1H5 and L@2H3

L@1H5 is a derivative of DTPA bisamide derivative (L@

3H3) where amide group at

each terminal nitrogen atom is replaced by amidothiol group. From a synthetic point of view,

this ligand was obtained by reaction of diethylenetriamine pentaceticacid bis-anhydride and

aminoethanethiol in the presence of triethylamine in DMF (Scheme III-1).6 After filtration,

the filtrate was then reacted with chloroform which leads to L@1H5 as a white powder.[15]

6 This ligand was synthesized by Pr. S. Roux’s group (Université de Franche-Comté, France).

169

N N NOH

OO

O

O O

O

OH2N SH

DMF; 70°CN N N

OH

O O

O

HN SHH

NHS

OH

O

HO

OCHCl3

N+

Scheme III-1: Synthesis of L@

1H5 (thiolated ligand of DTPA)

The ligand possesses five ionisable protons in its neutral form. So for the further

potentiometric titrations it will be named L@1H5 (protons indicated in red in Scheme III-1)).

In the second step, L@1H5 was grafted onto a gold nanoparticle.7 From a synthetic

point of view, the nanoparticles are synthesized by reducing a gold salt (HAuCl4·3H2O) with

NaBH4 in a mixture of methanol and water in the presence of L@1H5.[16] Basing on the FT-IR

and XPS analysis, it was revealed that ligands (L@1H5) are grafted forming multilayers at the

surface of gold nanoparticle with inter and intra layer disulphide bonds and are stabilized by

the electrostatic repulsions between charged particles; the zeta potentials measured in

solution show that the gold nanoparticle is always in its anionic form inbetween pH 2 and

12.[15]

Due to the nanoparticle structure that is to say, due to the L@1H5 packing at the

nanoparticle surface, all the protonation events that could be determined by potentiometry

experiments will be the result of overall proton exchanges. Therefore, these equilibria will be

studied by measuring the average number of protons released for a fixed amount of grafted

nanoparticles. To simplify the analysis and notation, the grafted gold nanoparticles will now

be considered as an entire entity named L@2H3 (see potentiometric results for this notation)

(Scheme III-2).

7 The gold nanoparticles were synthesized by Pr. S. Roux’s group.

170

N

N

N

COOH

HOOC

O

NH

OHN S

S

HOOC

NpAu

N

N

NCOOH

COOH

NH

OHN

S

S

HOOC

N NN

COOHCOOH

O

NH

O

NH

S S

COOH Scheme III-2: Model unit (L@

2H3) of ligand grafted on a gold nanoparticle

As references, DTPA-BMA, which has three ionisable protons in its neutral form

(named L@3H3: protons indicated in red (Figure III-8.a)[25], and DTPA, which has five

ionisable protons in its neutral form (named L@4H5: protons indicated in red (Figure III-8.b)

were used.

a. b.

Nt Nc Nt

C

COOHHOOC

C

COOH O

NHCH3

O

H3CHN

Nt Nc Nt

COOH

COOHHOOC

HOOC

COOH

Figure III-8: Structures of L@3H3 (DTPA-BMA) and L@

4H5 (DTPA)

b) Potentiometric study

Potentiometric studies are done to determine protonation constants of ligands at a

constant temperature of 25°C with ionic strength I = 0.1 (NMe4Cl). Ligands are solubilised in

in NMe4Cl and then titrated with tetra methyl ammonium hydroxide (NMe4OH, 5×10-2

mol.L-1) in between pH 2 and 12.

Potentiometric titrations could be depicted as curves of h versus pH, where h is the

average number of protons bounded to ligand (Figure III-9).

171

2 4 6 8 10 120

1

2

3

4

5

6

pH

L@1H5

L@2H3

L@4H5

h

2 4 6 8 10 120

1

2

3

4

5

6

pH

L@1H5

L@2H3

L@4H5

h

Figure III-9. Average number h of protons bound per mole of L@1H5 (), L@

2H3 (●) and

L@4H5 () as a function of pH.

For L@1H5, the curve begins at h >5 indicating the formation of a reprotonated

cationic species L@1H6

+ in acidic medium. The curve shows a plateau in the pH range 6-8 for

a h value of 3. This indicates that in this range, a sole triprotonated species exists, which then

could be deprotonated stepwise upon increasing the pH. For L@2H3, the curve starts at h

value ~ 4 which gives the evidence of the reprotonation of ligand and formation of protonated

species L@2H4. At pH range 8-9, appearance of a slight plateau for h =2 is observed which

indicate the coexistance of more than one species in this pH range and the formation of

L@2H2 as a major species. For L@

4H5, the curve begin at h =4 which indicates the first

deprotonation of ligand that is relatively strong and the non existence of cationic species. At

pH range 6-8, the curve shows the presence of a single diprotonated species for a h value of

2.

c) Determination of protonation constants of ligands L@1H5 and L@2H3

Calculation of the protonation constants is done using PROTAF software[26] from

several titrations of the ligand, varying their concentrations against NMe4OH (5×10-2

mol.L-1). Ligand is solubilised in hydrochloric acid, which permits the reprotonation of ligand

at low pH values.

172

Table III-1: Protonation constants log K0lh of L@1H5, L@

2H3, L@3H5 and L@

4H5, T = 25°C, I

= 0.1 (NMe4Cl) ) – Standard deviations are given in parenthesis

Equilibrium reactions log K01h(a) L@

1H5 L@2H3 L@

3H3[25] L@

4H5

L+H+ = LH log K011 10.37 (2) 11.26 (3) 9.4 10.61 (1)

LH+H+ = LH2 log K012 9.77 (1) 10.12 (2) 4.4 8.63 (1)

LH2+H+ = LH3 log K013 8.96 (2) 7.27 (3) 3.1 4.34 (1)

LH3+H+ = LH4 log K014 4.79 (1) 5.75 (2) 2.80 (1)

LH4+H+ = LH5 log K015 3.43 (1) 3.78 (1) 2.13 (1)

LH5+H+ = LH6 log K016 2.34 (1)

(a) - where K01h corresponds to the reaction equilibrium

And is defined by equation ]H][LH[

]LH[K 1)(h

1h

hh

01h

For L@1H5, six macroscopic protonation constants were determined. The two first

values (log K011 = 10.37 and log K012 = 9.77, Table III-1) could be assigned to successive

ionic equilibria of the two pendant RSH functions, successively deprotonated in two RS-

functions. Similar deprotonation constant values for the thiol groups were determined in the

case of 1,2-ethanedithiol (log K = 10.43 and 9.00), and dimercaptosuccinic acid (log K = 9.54

and 12.05).[27, 28] The four following constants could then involve the protonable sites of the

DTPA bisamide backbone. For DTPA bisamides derivatives,[25] it is well established that

protonation equilibria measured in the pH range 9.5-3, take place at the backbone nitrogen

atoms of the ligands. Thus for DTPA-BMA (L@3H3), the first proton (log K011 = 9.4, Table

III-1) is added at the central nitrogen atom Nc (Figure III-8.a) while the second (log K012 =

4.4) and the third (log K013 = 3.1) are added at the terminal nitrogen atoms (Nt, Figure III-

8.a). This led for L@3H3 a large ΔlogK12 (logK011-logK012 = 5) value, this difference being

greatly superior to the one in the parent DTPA (L@4H5) ligand (ΔlogK12 2). This difference

indicates that it is more difficult to add a second proton in the L@3H3 backbone than in L@

4H5.

This difference was interpreted by Geraldes et al. by the poorer ability of amide functions to

stabilize this second proton addition by electrostatic interactions. Indeed, for L@3H3 only one

carboxylate anion can exert an electrostatic field, while for L@4H5, two carboxylate ions can

play this role (Figure III-10).

HLH 1)(h1h

LHh

h

173

a. b.

O

HNt

O

N

O

R

R'H

O

HNt

O

O

O

R

Figure III-10: Stabilisation of the proton added on terminal nitrogen atom (Nt) a. L@3H3

(DTPA-BMA) and b. L@4H5 (DTPA) by carboxylate electrostatic field.[25]

Therefore, the second proton added in L@3H3 backbone, less stabilized, is more acidic

leading to a large difference between the two successive protonation constants log K011 and

log K012. This criterion was used for L@1H5 between protonation constants K013 and K014,

ΔlogK34 4.2. This difference although inferior to the one determined for L@3H3, can be

considered similar. This result is also similar to the one measured for a DTPA bisamide

conjugate of penicillamine (Δlog K34 3.2)[29] Therefore for L@1H5, log K013 and log K014

protonation events could be assigned to the central backbone nitrogen atom Nc and to a

terminal backbone nitrogen atom Nt protonations respectively, the two last constants being

attributed to the protonation of the second terminal backbone nitrogen atom Nt and to a

carboxylate oxygen respectively.

On the basis of these propositions, the acid-base behaviour of L@1H5 was summarized in

Scheme III-3.

2.34 3.43

8.96

10.37 9.77

H+N-OOC NH+

COOH

H+N COO-

HN

O

HS

NH

SH

O

H+N-OOC NH+

COO-

H+N COO-

HN

O

HS

NH

SH

O

H+N-OOC NH+

COO-

N COO-

HN

O

HS

NH

SH

ON

-OOC NH+

COO-

N COO-

HN

O

HS

NH

SH

O

N-OOC N

COO-

N COO-

HN

O

-S

NH

S-

ON

-OOC N

COO-

N COO-

HN

O

-S

NH

SH

ON

-OOC N

COO-

N COO-

HN

O

HS

NH

SH

O

4.79

log K011 log K012

log K013

log K014log K015log K016

Scheme III-3: Protonation scheme of L@1H5

174

In the case of L@2H3, five macroscopic constants are determined. In a simplified

outlook, which would consist of ligands grafted at the nanoparticle surface as indicated in

Scheme III-4, one can consider that the three first constants can be associated to the

deprotonation of the model unit (log K = 11.2, 10.12, 7.27, hence the system name L@2H3)

and the two last ones to its re-protonation (log K = 5.75, 3.78).

N

N

N

COOH

HOOC

O

NH

OHN S

S

HOOC

NpAu

N

N

NCOOH

COOH

NH

OHN

S

S

HOOC

N NN

COOHCOOH

O

NH

O

NH

S S

COOH Scheme III-4: Model unit (L@

2H3) of ligand grafted on a gold nanoparticle

The comparison of these values to the ones determined for the ligand L@1H5 alone

shows that in a general trend, the basicity of the ionisable protons in L@2H3 increase. This

hike in basicity can be interpreted as the result of H-bonds between the added protons.

Indeed, all the acido-basic sites, due to the ligand packing at the nanoparticle surface, are

spread all around the nano-object and can stabilize these protons. These results are similar to

the ones reported for polyaminocarboxylate ligands embedded on the surface of

macromolecules, for which the existence of an extended hydrogen bond network alters the

overall charge distribution at the vicinity of these ligands and modifies their basicity.[30] As a

consequence, the identification of the protonation sites at the surface of the functionalized

nanoparticle, that is to say a site-specific description of the nanoparticle protonation is rather

speculative since first accessibility of acido-basic functions is not probably homogenous at

the surface and second, successive additions of protons are probably followed by ligand re-

organisation at the surface.[31] One should note that one argument for this re-organisation is

the fact that there is a general increase of the ligand protonation constant values when it is

grafted at the surface. Indeed, without any reorganisation, the addition of protons at low pH

values would result8, due to electrostatic repulsion, in the basicity reduction of the

corresponding acido-basic functions,[32] which is not observed. With these macroscopic

8 For a given ionic strength

175

deprotonation constants in hand, the stability constant values of metal complexes will be

determined.

Another way to represent these data is to determine, by using HYSS software,[33] the

domain of existence of the protonated and deprotonated species of L@1H5 and L@

2H3 (Figure

III-11).

a. b.

2 4 6 8 10 120

20

40

60

80

100

[L@

1]

5-

L@

1H

4-

L@

1H

2

3-

L@

1H

3

2-

L@

1H

4

-

L@

1H

5

L@

1H

6

+

% o

f pro

tona

ted

spec

ies o

f L@

1 H5

pH 2 4 6 8 10 12

0

20

40

60

80

100

[L@2]L@

2H

L@2H2

L@2H3

L@2H4

L@2H5

% p

roto

nate

d sp

ecie

s of L

@2 H

3

pH

Figure III-11: Species distribution curves of ligands as a function of pH.

a. [L@1H5] = 2×10-3 mol.L-1, b. [L@

2H3] = 2×10-3 mol.L-1

In the distribution curve with L@1H5, from pH = 5 to 8, L@

1H32- is present in majority.

Similar to DTPA or EDTA, the neutral form L@1H5 is present in acidic solution (pH<4) and

doesn’t exceed 60% of the total species in solution. In L@2H3, distribution curve show the

formation of successive protonated forms in the total pH range. At pH ~ 7 this curve displays

the presence of L@2H3 as a major species. Hence, at physiological pH and in the presence of a

metal ion, the complexation will occur with L@1H3

2- and L@2H3 species.

2. Physicochemical studies of complexes with ligands L@1H3 and L@2H3


To evaluate the thermodynamic stability of Gd(III) complexes with L@1H5 and L@

2H3,

their overall stability constants were determined by potentiometry. Potentiometric studies are

done at a constant temperature of 25°C with ionic strength I = 0.1 (NMe4Cl). Ligands are

solubilised in NMe4Cl and then titrated, between pH = 2 and 12, with tetra methyl

ammonium hydroxide (NMe4OH, 5×10-2 mol.L-1). Direct titration of ligand-metal solutions is

achievable, since ligands are linear and they attain equilibrium rapidly.[34]

176

For all titration curves, it is possible to depict them as curves of h versus pH.

where h is the average number of protons bounded per mole of ligand. As an example, for

Gd(III)-L@1H5 system, the comparison of h versus pH curves for the ligand alone and in the

presence of Gd(III) ions (metal-to-ligand ratio R~1) are represented in Figure III-12.

2 4 6 8 10 120

1

2

3

4

5

pH

L@1H5

Gd(III)-L@1H5

h

2 4 6 8 10 120

1

2

3

4

5

pH

L@1H5

Gd(III)-L@1H5

h

Figure III-12. Average number h of protons bound per mole of Gd(III)-L@1H5 (), L@

1H5

(○) as a function of pH.

In the presence of Gd(III), the titration curve was depreciated relative to the titration

curve of the ligand alone. This deviation revealed the complex formation from pH 2.5. The

Gd(III)-L@1H5 curve represent a plateau in the pH range of 5<pH<8 and pH>11 for h =2 and

0, respectively. The first plateau correspond to the existence of a protonated complex in

solution while the second one could be correlated to the existence of a deprotonated species.

b) Determination of stability constants of Gd(III), Cu(II), Zn(II), Ca(II) and Na(I)

For L@1H5 and L@

2H3 the computed chemical model PROTAF[26] suggests the

formation of [MLH2], [MLH] and [ML] species for L@1H5 and L@

2H3. The results are

presented in Table III-2 and compared with those of L@4H5 ones.

Table III-2: Stability constants of M/L@1H5, M/L@

2H3 and M/L@4H5 complexes (where M

= Gd(III)), T = 25°C, I = 0.1 (NMe4Cl) ) – Standard deviations are given in parenthesis

L@1H5 L@

2H3 L@4H5

Gd(III) Gd(III) Gd(III)

log mlh M+L+2H+ = MLH2 34.94 (4) 35.6 (2)

M+L+H+ = MLH 26.41 (5) 31.7 (2)

M+L = ML 16.46 (5) 21.2 (2) 21.57 (3)

177

Table III-3: Deprotonation constants of M/L@1H5, M/L@

2H3 and M/L@4H5 complexes

(where M = Gd(III)), T = 25°C, I = 0.1 (NMe4Cl) )

log K

L@1H5 L@

2H3 L@4H5

MLH+H+ = MLH2 8.53 3.9

ML+H+ = MLH 9.95 10.5

The two successive deprotonation constant values of [GdL@1H2] (log K = 9.95 and

8.57) are almost similar to the two first deprotonation constants of L@1H5 alone (log K011 =

10.37 and log K012 = 9.77). This suggests that the successive deprotonations of [GdL@1H2]

into [GdL@1H]- and [GdL@

1]2- involve the thiol functions. The two deprotonation values

being at lower pH in Gd(III)-L@1H5 than in the ligand alone could be interpreted as an

assistance of the metal to the complex deprotonation. For Gd(III)-L@2H3 system, as L@

2H3 is

a unit model, a proposition for the structure of GdL@2 complex and its protonated forms

could only be speculative and the more reliable way to analyse this system is to describe the

distribution diagram of Gd(III)-L@2H3 system and to determine the percentage of free metal

ion present in the entire pH range.

The species distribution diagrams of Gd(III)-L@1H5 and Gd(III)-L@

2H3 systems were

reported in Figure III-13.

a. b.

2 4 6 8 10 120

20

40

60

80

100

[GdL@1]2-

[GdL@1H]-

[GdL@1H2]

Gd3+

%G

d

pH

2 4 6 8 10 120

20

40

60

80

100

[GdL@2]

[GdL@2H]

[GdL@2H2]

Gd3+

% G

d

pH

Figure III-13: Species distribution curves of a. Gd(III)-L@1H5 and b. Gd(III)-L@

2H3

systems.

178

It is interesting to note that for Gd(III)-L@1H5 system, only one species exist at

physiological pH ([GdL@1H2]). This protonated Gd complex forms at pH 2, and is

deprotonated successively to form [GdL@1]2- predominantly at pH>10. In the case of Gd(III)-

L@2H3 system, the complex is completed at about pH 3. At lower pH, two species (GdL@

2H2

and GdL@2H) are present in solution. The GdL@

2H species is in majority in the pH range of 4-

10 and deprotonates after pH 10 into GdL@2.

The stability of these Gd complexes is judged by calculating the percentage of free

metal ion present between pH 2 and 12 for each complex. The logarithm of the percentage of

free metal ion (log([M]free/[M]total)) calculation as a function of pH is one of the ways for

comparison, as the direct comparison of the overall stability constants of the complexes is

inappropriate because of the different protonation constants for their corresponding ligands.

The logarithm of the percentage of free metal ion (log([M]free/[M]total)) in function of pH is

traced for Gd(III)-L@1H5 system and Gd(III)-L@

2H3 system. The obtained curves are

compared with that of Gd(III)-L@4H5 system. The lower is the curve the more stable is the

complex.

-10

-8

-6

-4

-2

0pH

log

([Gd]

free/[

Gd]

tota

l)

L@1H5

L@2H3

L@4H5

2 4 6 8 10 12

Figure III-14: Logarithm of percentage of free metal ion concentration of L@1H5, L@

2H3,

and L@4H5 in relation to Gd(III), [ligand] = [Gd] = 2.0×10-3 mol.L-1

From the Figure III-14, it is clear, before pH>4, that stability of Gd(III)-L@1H5

system is lower compared to that of other Gd complexes. The stabilities of Gd(III)-L@2H3

and Gd(III)-L@4H5 are similar. This result highlights an increase of the Gd complex stability

when these complexes are anchored at the surface of the nanoparticle. This improvement in

the complex thermodynamic stability, which is very satisfactory, could be due to the positive

179

influence of the ligand packing at the surface that stabilise probably by cooperative effects,

even the protonated forms of these complexes.

It is equitably important to compare the stability of Gd complexes to the ones of

potentially competitive endogenous ions such as Cu(II), Zn(II), and Ca(II). It is indeed well

known that Zn(II) ions could behave as exchanging metal towards Gd(III).[35, 36] Besides,

Ca(II) and Gd(III) possess similar ionic radii, leading to potential exchanges in vivo.[37] The

overall complexation constants of M-L@1H5, M-L@

2H3 and M-L@4H5 (where M = Cu(II),

Zn(II) and Ca(II)) are reported in Tables III-4.

Table III-4: Stability constants of M/L@1H5, M/L@

2H3 and M/L@4H5 complexes (where M =

Cu(II)/Zn(II)/Ca(II)), T = 25°C, I = 0.1 (NMe4Cl)) – Standard deviations are given in

parenthesis L@

1H5 L@2H3 L@

4H5

Cu(II) Zn(II) Ca(II) Cu(II) Zn(II) Ca(II) Cu(II) Zn(II) Ca(II)

Log mlh

M+L+4H+=MLH4 36.7 (1)

M+L+3H+=MLH3 37.2 (2) 37.1 (1) 32.3 (1)

M+L+2H+=MLH2 33.7 (1) 33.0 (1) 27.1 (1) 31.3 (1) 31.2 (1) 27.9 (1)

M+L+H+=MLH 28.4 (1) 24.8 (1) 18.2 (1) 25.96 (1) 25.3 (1) 20.8 (1) 22.87 (2) 23.0 (3) 16.6 (1)

M+L=ML 19.8 (1) 15.2 (1) 8.0 (1) 16.7 (1) 14.8 (2) 10.0 (1) 18.19 (3) 17.42

(4)

10.31

(6)

M+L=MLH-1 + H+ 6.36 (4) -2.4 (2)

2M+L+H+=M2LH 36.8 (2)

2M+L=M2L 32.8 (2) 4.5 (1) 24.9 (1) 24.9 (1)

2M+L=M2LH-1+H+ 23.6 (3)

2M+L=M2LH-2+2H+ 13.2 (2) 6.7 (1) 6.7 (1)

M+2L+4H+ =

ML2H4

53.8 (2)

M+2L+3H+ =

ML2H3

45.4 (1)

M+2L+2H+ =

ML2H2

36.7 (2)

M+2L+H+ = ML2H 25.8 (1)

M+2L = ML2 15.9 (2)

From the logmlh values, the deprotonation constants of Cu(II), Zn(II) and Ca(II)

complexes are calculated. (see experimental section, Table 1, page 208)

Basing on the above stability constant values, speciation diagrams for M-L@1H5 and

M-L@2H3 complexes (where M = Cu(II), Zn(II) and Ca(II)) are reported in Figure III-15, 16,

and 17. All constants values are taken into consideration to plot speciation curves.

180

a. b.

2 4 6 8 10 120

20

40

60

80

100

[Cu2L@1(OH)2]

3-

[CuL@1]3-[CuL@

1H]2-

[Cu2L@1]-

[CuL@1H2]

-

[Cu2L@1H]

Cu2+

[CuL@1H3]

% C

u

pH

2 4 6 8 10 120

20

40

60

80

100[CuL@

2H2][CuL@

2H]

[CuL@2]

Cu2+

%C

u

pH

Figure III-15: Species distribution curves of a. Cu(II)-L@1H5 and b. Cu(II)-L@

2H3 systems.

This diagram shows different species for two ligands (L@1H5 and L@

2H3) at a specific

point of pH. Dinuclear [Cu2L@1H] and mononuclear [CuL@

1H3]+ species form

simultaneously at pH = 2. Deprotonated [Cu2L@1]- species and other mononuclear protonated

species [CuL@1H2]+ and [CuL@

1H]2- prevail between pH 2 and 9. Beyond pH 9, the

deprotonated [CuL@1]3- complex becomes predominant. Formation of hydroxylated species

could be noticed in a very minute quantity. For Cu(II)-L@2H3 system, complex formation

begins at pH >3. Protonated complexes ([CuL@2H2] and [CuL@

2H]) exist until pH 9 in

solution. Beyond pH 9, the deprotonated [CuL@2] complex becomes predominant.

Speciation diagrams of Zn systems are reported in Figure III-16

a. b.

2 4 6 8 10 120

20

40

60

80

100

[ZnL@1H]2-

[ZnL@1]3-

[ZnL@1H2]

-[ZnL@

1H3]

Zn

%Zn

pH 2 4 6 8 10 12

0

20

40

60

80

100

[ZnL@2]

[ZnLH@2]

[ZnL@2H2]

Zn2+

%Zn

pH

Figure III-16: Species distribution curves of a. Zn(II)-L@1H5 and b. Zn(II)-L@

2H3 systems.

For Zn(II)-L@1H5 system, stability constant values for dinuclear complex are

determined in the various conditions. However, in the 1:1 stoichiometric conditions, the

formation of dinuclear zinc species in L@1H5 is not seen in the species distribution curve

181

indicating its formation in an insignificant quantity. This could be considered due to the

lower stability of dinuclear zinc complex compared to copper complex, as the ligand could

not provide sufficient coordination sites for the second Zn to form two octahedral complexes.

Similar to that of Cu(II)-L@2H3, Zn complex doesn’t involve in dinuclear complex formation.

Beyond pH 10, deprotonated zinc complexes ([ZnL@1]2- and [ZnL@

2]) exist in solution.

Speciation diagram for Ca systems are represented in Figure III-17.

a. b.

2 4 6 8 10 120

20

40

60

80

100

Ca2+

[CaL@1H4]

+ [CaL@1H3]

[CaL@1H2]

-

[CaL@1H]2-

[CaL@1]3-

%C

a

pH 2 4 6 8 10 12

0

20

40

60

80

100

[CaL@22H4]

[CaL@2H2]

[CaL@2H]

[CaL@2]

Ca2+

%C

a

pH

Figure III-17: Species distribution curves of a. Ca(II)-L@1H5 and b. Ca(II)-L@

2H3 systems.

In both the systems, Ca(II) is not complexed effectively by the ligand in acidic pH. For

Ca(II)-L@1H5 system, presence of protonated species even in highly basic pH indicate that the

thiol groups are not involved in complex formation. This can also be interpreted basing on the

log K values (10.2 and 8.9) which are relatively close to the log K011 = 10.37 and log K012 =

9.77 determined for the ligand itself (Table III-1). For Ca(II)-L@2H3 system, successively

protonated [CaL@2Hn] species are present in solution. The [CaL@

22H3] species doesn’t exist

and [CaL@2

2H4] species prevail around less than 10% in the total composition of species.

As previously said, the overall stability complexes of L@1H5 and L@

2H3 could not be

compared. Nevertheless, the stabilities of different metal complexes of the same ligand could

be compared. The general trend of increasing complex stability is Ca(II) < Zn(II) < Cu(II) <

Gd(III) for all systems. The data show that copper complexes are more stable than the zinc

ones according to the Irving - Williams order. To circumvent the different protonated states

of the complexes, an efficient way to compare the sequestering ability of L@1H5 and L@

2H3

for all these ions is to determine the logarithmic amount of free metal relative to the metal

coordinated on the whole pH range.

182

a. b. c.

-12

-10

-8

-6

-4

-2

0

log(

[Cu]

free/[

Cu]

tota

l)

L@1H5

L@2H3

L@4H5

2 4 6 8 10 12pH

-12

-10

-8

-6

-4

-2

0pH

log(

[Zn]

free/[

Zn] to

tal)

L@1H5

L@2H3

L@4H5

2 4 6 8 10 12

-5

-4

-3

-2

-1

0pH

log(

[Ca]

free/[

Ca]

tota

l)

L@1H5

L@2H3

L@4H5

2 4 6 8 10 12

Figure III-18: Logarithm of percentage of free metal ion concentration of L@1H5, L@

2H3, and

L@4H5 in relation to M equal to a. Cu(II), b. Zn(II), c. Ca(II), [ligand] = [M] = 2×10-3 mol.L-1

In the Figure III-18.a, Cu(II)-L@1H5 display more affinity for copper than Cu(II)-

L@2H3 and Cu(II)-L@

4H5. The stability of Cu(II)-L@1H5 could be explained by the affinity of

copper for thiol groups and due to the formation of dinuclear species(see in Species

distribution Figure III-15, page 180).

For zinc complexes (Figure III-18.b), the affinity of both L@1H5 and L@

2H3 ligands

towards Zn(II) is similar. However, Zn(II)-L@4H5 system exhibits higher stability compared

to other complexes. This can be argued by the ability of acetate groups in DTPA to stabilize

the complex than the bisamide-DTPA ligand.

In the case of Ca complexes (Figure III-18.c), all the systems almost exhibit similar

affinity towards Ca(II) ion at physiological pH.

3. Spectroscopic study of metal complexes with ligands L@1H5 and L@4H5 EPR spectroscopic measurements are realised for copper complexes with

ligands L@1H5 (and L@

4H5 for comparison) at pH ~ 6 using frozen solution (water/ethanol) of

complexes (T=150K). EPR spectra were taken at various ratios of Cu(II) and ligand, by

adding Cu(II) just up to 2:1 mole ratio (R = [Metal]/[Ligand]). The interest in measuring

these complexes at various ratios is to favour the formation of polynuclear complexes.

Indeed, in the distribution curve of Cu(II)-L@1H5 (see speciation diagram Figure III-15, page

180), at pH ~ 6 and at 1:1 mole ratio, [Cu2L@1]- complex exist around < 50% in total species

composition. Dinuclear complexes were also described and observed by EPR spectra [38] for

Cu(II)-L@4H5 system, even at 1:1 ratio. However, an EPR spectrum of a dinuclear complex

depends on many considerations. In dicopper complexes, where S=1, the hamiltonian

expression becomes complicated as it depends on coupling of electrons, the distance of the

nuclide and their position compared to each other.[39] In agreement with this consideration,

183

the EPR spectra of dinuclear complex is centered at 3200 G or at 1500G and the Hamiltonian

spin is taken as

H = BBgASA + IAAASA + BBgBSB + IBABSB + JABSASB+ SDS

where B is the Bohr magneton, gA, gB, AA, AB correspond to the individual site g and A

tensors, SA and SB are spins on the two copper centres (SA = SB = ½) and S is the total

angular momentum of spin, JAB accounts for the isotropic exchange interaction and D is the

zero field splitting tensor.

With L@4H5, the dinuclear complex have a distance of Cu-Cu at 5.5Å.[37] In this case the

isotropic exchange interaction (JAB) is negligible when compared to zero field splitting

perturbation (D). This results in two series of seven lines, which is theoretically expected.[40]

For each allowed ΔMS = 1 transition; they correspond to the hyperfine coupling between the

two copper d9 ions (ICu = 3/2). According to the zero-field splitting amplitude, the two septets

can either overlap (weak D factor) or could be shifted, or even in some cases it is difficult to

observe all the lines.[41, 42]

The dinuclear formation in distribution curves at stoichiometric conditions prompted

us to follow the formation of these dinuclear complexes for Cu(II)-L@1H5 system by EPR

spectroscopy. Figure III-19.a represents the spectra of copper complexes with L@1H5 and for

comparison, copper complexes with L@4H5 recorded at various ratios of metal and ligand

(upto 2:1 mole ratio of R = [M]/[L]) are represented in Figure III-19.b.

184

a. b.

2600 2800 3000 3200 3400 3600Gauss

0.75 Cu 1.0 Cu 1.5 Cu 2.0 Cu 2.5 Cu

2600 2800 3000 3200 3400 3600

Gauss

0.75 Cu 1.0 Cu 1.5 Cu 2.0 Cu

Figure III-19: EPR spectra evolution of a. Cu(II)-L@1H5 b. Cu(II)-L@

4H5 according to

addition of Cu(II) ([L@1H5] = 10-3 mol L-1, [L@

4H5] =2×10-3 mol L-1 in water/ethanol, T=

150K)

For Cu(II)-L@1H5 system, at R<1 shows an axial spectra centered at 3200G with the

resolution of hyperfine structure of four equidistant peaks in the parallel region. On further

addition of Cu(II), broadening of the perpendicular region and disappearance of the peaks in

parallel region is clearly visible. For R=2, the absence of peaks in parallel region show that

copper complex is no more mononuclear but probably complexed with ligand forming a

dinuclear species with an EPR signal at 3200G constantly observed. However, the signal is

very broad, the hyperfine structure is not resolved, and a new peak is noticed at 3400G. In an

excess of copper condition (i.e. R > 2), the appearance of peaks in parallel region in EPR

signal gives the evidence of excess copper which could, before R=2, the total coordination of

copper by the ligand and the formation of dinuclear complex. In this case the hyperfine

structures are not sufficiently resolved to observe the fourteen expected peaks.

In Cu(III)-L@4H5 system, at R<1 the hyperfine signal in the parallel region is

conspicuous and with the addition of copper, similar characteristics as Cu(II)-L@1H5 system

are observed. This evolution is similar to the one earlier observed by Martell et al.[38] in

polyamino linear ligands, upon addition of more than one equivalent of metal ion to ligand.

185

It is interesting to give a commentary for the spectrum of Cu(II)-L@4H5 system at 1

equivalent of Cu (Figure III-20).

2600 2800 3000 3200 3400 3600

AN//

ACu//

Gauss

Experimental Simulation

Figure III-20: Experimental and simulated spectra of Cu(II)-L@4H5 system.

In the parallel region, the spectrum shows three out of four set of equidistant peaks, in

which a set of peaks belong to the hyperfine structure of Cu(II) ion (ICu = 3/2). The other

peaks are considered to be the super hyperfine structures of one of the nitrogen atoms which

is coordinated to copper ion. These lines are due to the super-hyperfine coupling between the

nuclear spin of copper (ICu = 3/2) and the nuclear spin of a single nitrogen atom in one of the

axial positions (IN = 1). The coupling constant obtained by the simulation is

AN = 45×10-4 cm-1. This result is comparable with that of polypropylene imine copper

complex (AN = 36.8×10-4 cm-1).[43] The unusually high value of nitrogen super-hyperfine

constant could be due to the nature of the copper-ligand bonding.[43]The other values obtained

by simulating the experimental spectrum gave us (A// = 148×10-4 cm-1; g// = 2.288; g┴ = 2.06)

are similar to that of Cu(II)-DTPA in 3000G region (A// = 140×10-4 cm-1; g// = 2.30; g┴ =

2.10)[38], thereby possibly possessing a N3O3 geometry.

Comparison of Cu(II)-L@1H5 spectrum with Cu(II)-L@

4H5 indicate that the spectra are

identical. In this case, the super-hyperfine coupling should also be seen in Cu(II)-L@1H5

system. Unfortunately due the high noise in the peaks obstruct the clear visibility.

186

Figure III-21: EPR spectra of copper complexes with L@

1H5 and L@4H5 at 1:1 mole ratio of

metal (Cu(II)) and ligand.

The above comparison of mononuclear complex spectra leads us in proposing an

identical geometry (N3O3) for both copper complexes.

To resume, dinuclear copper formation for Cu(II)-L@1H5 system under conditions

where R = [M]/[L] > 1, could not be resolved in the EPR spectra in 3200G region. However,

in the ratio range 1 < R < 2, the modification of spectra, comparable to the one described for

Cu(II)-L@4H5 system and ascribed in this case to the formation of a dinuclear complex,[38]

allowed us to propose the formation of dinuclear complexes for Cu(II)-L@1H5 system.

2600 2800 3000 3200 3400 3600

Gauss

L @ 1 H 5

L @ 3 H 5

2600 2800 3000 3200 3400 3600

Gauss

L @ 1 H 5

L @ 4 H 5

187

C. Transmetallation studies of Gd(III)-L@1H5 and Gd(III)-L@

2H3

In order to get a deeper insight on this point, kinetic measurements were performed for

Gd(III)-L@1H5 and Gd(III)-L@

2H3 in the presence of copper and zinc. Unfortunately, all the

solutions involving Cu(II) have precipitated which did not allow experiments with this ion.

With Zn(II), the exchanging reactions (Equation (1)) were followed by means of

relaxometry for Gd(III)-L@1H5 and Gd(III)-L@

2H3 in the conditions used for DO3A

benzimidazole derivative (Chapter 2) and by UV spectroscopy for only Gd(III)-L@1H5 (since

the experiment was not possible with the dark nano-suspension).

GdLT + Zn2+ Gd3+ + ZnLT

with L = Gd(III)-L@1H5, Gd(III)-L@

2H3

n

nT GdLHGdL

(1)

where GdLT and ZnLT are complexes of Gd(III) and Zn(II). The experimental conditions

were different by these two techniques since by relaxometry, Gd(III)-L@1H5 and Zn(II) were

mixed in 1:1 conditions while for UV experiments, Zn(II) was in excess towards Gd(III)-

L@1H5. Therefore, formally these experiments could not be compared but analyzed in

combination, to state about the kinetic inertness of the Gd complexes (relaxometry) and to

have an insight of the mechanisms (UV experiments).

1. Kinetics of demetallation followed by relaxometric measurements

The kinetic stabilities of Gd(III)-L@1H5 and Gd(III)-L@

2H3 were characterized by their

normalized paramagnetic relaxation rates R1(t) / R1(t = 0) and compared them to those of

Gd(III)-L@4H5 (Figure III-22).

188

a. b.

0 1000 2000 3000 4000 50000.0

0.2

0.4

0.6

0.8

1.0

Gd(III)-L3@H5

Gd(III)-L2@H3

Gd(III)-L1@H5

R1/

R1 0

t (mn)

4

0 1000 2000 3000 4000 50000.0

0.2

0.4

0.6

0.8

1.0

Gd(III)-L3@H5

Gd(III)-L2@H3

Gd(III)-L1@H5

R1/

R1 0

t (mn)

4

0 100 200 300 400 5000.0

0.2

0.4

0.6

0.8

1.0

Gd(III) - L3@H5

Gd(III) - L2@H3

Gd(III) - L1@H5

R1/

R1 0

t (mn)

4

0 100 200 300 400 5000.0

0.2

0.4

0.6

0.8

1.0

Gd(III) - L3@H5

Gd(III) - L2@H3

Gd(III) - L1@H5

R1/

R1 0

t (mn)

4

Figure III-22: a. Evolution of R1(t) / R1(t = 0) (T= 310 K; pH 7.0) versus time for (■) Gd(III)-

L@1H5 complex and (●) Gd(III)-L@

2H3 ( for Gd-DTPA) in the presence of equimolar amounts

of Zn(II) ions in phosphate buffer solution. b. Enlargement of R1(t)/ R1(t = 0) for

0< t<500 min

A semi-quantitative evaluation of the transmetallation kinetics was performed by

comparing the time needed to reach 80% of the initial R1(t = 0) value. The shortest delay was

obtained for Gd(III)-L@1H5 (110 min) while for Gd(III)-L@

2H3 the delay was similar to the

one measured for Gd(III)-L@4H5 (220 mn). These results were in favour of the kinetic

inertness of Gd(III)-L@2H3 system.

A theoretical description was attempted to evaluate in the current experimental

conditions the rate of transmetallation of the Gd species and the Gd species lifetimes

(Equations (2) and (3)).

The rate of the reactions was expressed as given in Equation (2):

Tobs

T GdLkdt

GdLdv (2)

The value kobs is a pseudo first order constant and [GdL]t is the total concentration of

the complex (L = L@1H5, L@

2H3, L@4H5 (DTPA) respectively). Under our experimental

conditions (phosphate buffer, pH 7), the concentration of free Zn(II) and Gd(III) are very low

owing to their low solubilities and were assumed to be approximately constant.[44]

Consequently, the rate of transmetallation kobs were determined with the use of Equation (2),

where R1(t0), R1(t) and R1(tq) are the relaxation rate values at the start, at time t and at

equilibrium of the reaction.

R1(t) = R1(te) + (R1(t0) - R1(te)) exp (-kobs t) (3)

189

The kobs values obtained by fitting the experimental data were equal to 3.910-3 min-1

and 2.710-3 min-1 for Gd(III)-L@1H5 and Gd(III)-L@

2H3, respectively, while for the

reference DTPA:Gd this constant was equal to 2.510-3 min-1. Therefore, Gd(III)-L@1H5 and

Gd(III)- L@2H3 t1/2 values were equal to 177 min and 257 min respectively, while for the

reference Gd(III)-L@4H5 the half-life was equal to 277 min. At the end of the observation

period (~ 3 days) the ratio R1(t) / R1(t = 0) was about 10% for Gd(III)-L@1H5 compared with

50% for the reference Gd(III)-L@4H5. This indicated that the bisamide complex Gd(III)-

L@1H5 show more extensive transmetallation than the parent compound. This behaviour was

consistent with what was already determined in Gd(III)-DTPA bisamide series.[44, 45] The

reason lies probably in the substitution of two carboxylate groups by two amide groups in the

first coordination sphere of Gd(III). For Gd(III)-L@2H3 the ratio of R1(t) / R1(t = 0) was

remarkably higher (33%) and comparable to the one of the reference Gd(III)-L@4H5. It

indicated that when Gd(III)-L@1H5 is grafted onto the Au nanoparticle, its transmetallation

kinetics tended to be reduced. This behaviour was also already reported in bisamide DTPA

series for which amide functions were functionalized with bulky groups. The similar result

obtained here suggested that the nanoparticle with the ligand packing at its surface, rigidifies

the structure of the complex, probably limiting competitive ion accessibility and then

preventing Gd(III)-L@1H5 from an extensive demetallation.

2. Kinetics of transmetallation followed by UV spectroscopy

Pathways involving the catalysis by endogenously available metal ions, such as Zn(II)

are very important in the dissociation of acyclic complexes. The results of earlier kinetic

studies on the transmetallation reactions of Gd(III)-DTPA with Cu(II), Zn(II), and Eu(III)

have shown that the reactions take place with direct attack of the exchanging metal ions on

the complex and also with proton-assisted dissociation of the complex (followed by fast

reaction between the metal ions and the free ligands).[46, 47]

The rates of the exchange reactions (Equation 4) were studied at different concentrations of

the exchanging Zn(II)ı (410-3 < [Zn(II)] < 10-2 mol L-1) and at different pH values in the

5.8 – 6.5 range.

Gd(III)-L@1H5 + Zn(II) Gd(III) + Zn(II)-L@

1H5 (4)

Under such conditions, the reactions can be regarded as pseudo-first order and the

rates of the reactions can be expressed as in Equation (5).

190

Tobs

T GdLkdt

GdLdv

with : GdLZnGdLZnHGdLHGdLHGdLGdL 2T

(5)

In Figure III-23 are shown the dependency of kobs values towards pH and Zn(II)

concentration.

a. b.

10

15

20

25

30

2.0E-07 4.0E-07 6.0E-07 8.0E-07 1.0E-06 1.2E-06 1.4E-06 1.6E-06 1.8E-06

[H+] (mol L

-1)

k ob

s (1

0-4

s-1

)

[Zn2+

] = 4 mM

[Zn2+

] = 6 mM

[Zn2+

] = 8 mM

[Zn2+] = 10 mM

10

15

20

25

30

3.00E-03 5.00E-03 7.00E-03 9.00E-03 1.10E-02

[Zn2+] (mol L-1)

k ob

s (1

0-4

s-1

)

pH = 5.8

pH = 6.0

pH = 6.2

pH = 6.5

Figure III-23: a. Plots of kobs versus H+ concentration and b. plots of the kobs versus Zn(II)

concentration for the reaction between Gd(III)-L@1H5 and Zn(II).

First, the increase in the H+ concentration resulted in an increase of the kobs values.

This result can be explained by the formation and faster dissociation of the protonated

[GdL@1H]-.[46, 47] From a structural point of view and as already proposed for pH assisted

Gd(III)-DTPA demetallation,[47] protonation can occur at a carboxylate group of the L@1H5

backbone, followed by the transfer of this proton to a neighbouring nitrogen atom. Then, or

the resulting free glycinate group can be re-coordinated or the L@1H5 donor atoms can be

stepwise de-coordinated to lead to the complex dissociation. The resulting protonated ligand

is then able to react with the exchanging Zn(II) ion.[46] The corresponding dissociation

reactions can be summarized in the reaction scheme shown in Scheme III-5.

GdL GdLHn Gd + HnLKGdLHn kGdLH

(6)

Scheme III-5: Proton-assisted demetallation (L = L@1H5)

Second, the increase in the kobs values with an increase in the concentration Zn(II)

reported in this study was already reported for DTPA bisamide derivatives.[46, 47] The reaction

191

mechanism proposed implicated the direct reactions between Gd(III)-L@1H5 and the

exchanging metal ion, which occurred via the formation of the hetero-dinuclear intermediate

GdLM. This hypothesis is supported by the fact that for DTPA derivatives, the formation of

dinuclear complexes [Ln2(DTPA)]+ has been already detected.[46-48] From a structural point

of view, in the first step of the formation of the hetero-dinuclear intermediate, a carboxylate

group is probably coordinated to the attacking Zn(II) metal ion, and in the course of the

reaction, the functional groups of L@1H5 are slowly transferred to the metal step by step to

form the ZnL complex (Scheme III-6).

GdL + M GdLM Gd + ZnLKGdLM kM

GdL

(7)

Scheme III-6: Zn-assisted demetallation (L = L@1H5)

Finally, for interpreting the dependence of kobs values on the concentration of H+ and

Zn(II), a reaction pathway that occurs with the direct attack of the exchanging metal ion on

the protonated complex has to be assumed (Equation 8)

GdLH + M GdLHM Gd + ZnLHKGdLHM kM

GdLH

(8)

Scheme III-7: Proton and Zn-assisted demetallation (L = L@1H5)

By taking all the possible reaction pathways into account, the rate of exchange

between Gd(III)-L@1H5 and the exchanging metal ion Zn(II) can be expressed as shown in

Equation (9), where [GdHnL] and [GdLZnHn] are the concentrations of the protonated and

dinuclear complexes, respectively.

Tobsi i GdLkvv (9)

If we take into account the total concentration of the complex

GdLZnGdLZnHGdLHGdLHGdLGdL 2T

the equations defining the overall stability constants of the protonated and hetero-dinuclear

complexes

LGd

GdLGdL

3110 , HLGd

GdLH3111GdLH

, 23

2112GdLH HLGd

GdLH2

and

192

231110GdLZn ZnLGdGdLZn

, HLZnGd

GdLZnH231111GdLZnH

and Equation (5),

the rate constant kobs can be expressed as follows in Equation (10)

21110

21111111

2112110

229

28

227

265

2224

223

22

31

ZnZnHHH

ZnAZnAZnHAZnHAHAZnHAZnHAHAHAkobs

(see Experimental Section for the demonstration)

The kobs values obtained for the reaction between the complex and Zn2+ were fitted and the

result of the fit correspond to the simplified Equation (10 b).

21110

21111111

2112110

229

28

227

26

ZnZnHHH

ZnAZnAZnHAZnHAkobs

(10 b)

where: A6 = k1× 1111, A7 = k2×1111, A8 = k3×1110, A9 = k4×1110

and are associated to the following equations:

ZnGdLZnkvGdLZnkv

ZnGdLZnHkvGdLZnHkv

44

33

222

11

The rate constants calculated were then as follows:

k1 = 38 10-4 M-1 s-1, k2 = 1007 10-4 M-2 s-1, k3 = 0.08 10-4 M-1 s-1, k4 = 654 10-4 M-2 s-1

These results indicated that the exchange reaction with Zn(II) predominantly occur

with direct attack of Zn(II) on the complex present in solution either in its monoprotonated

and non-protonated hetero-dinuclear form.

D. Conclusion

193

In this chapter we described the complexation of thiolated ligands DTDTPA and

Au@DTDTPA which corresponded to the previous ligand anchored at the surface of a gold

nanoparticle (named L@1H5 and L@

2H3 synthesized in Pr. S. Roux group), by Cu(II), Zn(II),

Ca(II), Na(I) and Gd(III). In L@1H5 the thiol groups SH were free while in L@

2H3, the thiol

groups were mobilized for the grafting onto the nanoparticle and the formation of disulide

bonds.

Whatever the system, L@1H5 or L@

2H3, the general trend of increasing complex

stability was Ca(II) < Zn(II) < Cu(II) < Gd(III). For L@1H5, the analysis of each speciation

diagram indicated that at physiological pH only one complex is formed whose stoichiometry

was [ML@1H2] for which the thiol functions are in SH form. Therefore, the pH increase

provoked the successive deprotonation of these functions. For L@2H3 given the density of

ligands L@1H5 present on the surface of the nanoparticle and the likely interactions between

them, it is unrealistic to propose a structure for the associated complex. Nevertheless, the

existence of different species in solution containing Gd(III) did not appear, which is a good

thing for the use of these systems in biological media. Furthermore, the speciation diagram of

L@1H5 and L@

2H3 in the presence of Gd(III) ions indicated at physiological pH the formation

of a sole complex, which was relevant for MRI application. The comparison of L@1H5 and

Gd(III)-L@2H3 system stabilities highlighted that Gd(III)-L@

1H5 was less stable than Gd(III)-

L@2H3, this latter being 2 orders of magnitude more stable at physiological pH. Moreover, the

comparison of L@1H5 and L@

2H3 affinities for the three metal ions indicated that for each of

them the best affinity was obtained for Gd(III), which is a crucial point for the use of these

systems as MRI contrast agents.

The goal of this work was secondly to evaluate the kinetic inertness of gadolinium

complexes with L@1H5 and L@

2H3. Transmetallation experiments of Gd(III)-L@1H5 and

Gd(III)- L@2H3 complexes in the presence of stoichiometric amounts of Zn(II) in phosphate

buffer (pH 7.0) and followed indicated that transmetallation occurred more rapidly for

Gd(III)-L@1H5 than for Gd(III)- L@

2H3 (whose kinetic inertness was similar to the one of Gd-

DTPA). In other words, these results indicated that when Gd(III)-L@1H5 is grafted onto the

Au nanoparticle, its transmetallation kinetics tended to be reduced. Therefore one can

propose that the bulky nanoparticle rigidified the structure of the complex and prevented

Gd(III)-L@1H5 from an extensive demetallation, which was a good point for the possible use

of these nanoparticles in living organisms for imaging applications. Transmetallation

experiments in the presence of an excess of Zn(II) suggested that for Gd(III)-L@1H5, from a

194

mechanistical point of view, the driven force of the transmetallation could be the formation of

hetero-dinuclear Gd(III)-L@1H5-Zn(II) complex, protonated or not.

References

195

1. C. Sun, J. S. H. Lee, M. Zhang, Adv. Drug Delivery Rev. 2008, 60, 1252-1265.

2. J. Dobson, Drug Dev. Res. 2006, 67, 55-60.

3. M. Ferrari, Nat. Rev. Cancer. 2005, 5, 161-171.

4. S. A. Wickline, A. M. Neubauer, P. M. Winter, S. D. Caruthers, G. M. Lanza, J.

Magn. Reson. Imaging. 2007, 25, 667-680.

5. Z. Derong, L. Fuyao, N. Lina, L. Dianjun, W. Zhenxin, Int. J. Mol. Sci. 2013, 14,

10591-10607.

6. a- S. Aime, D. D. Castelli, S. G. Crich, E. Gianolio, E. Terreno, Acc. Chem. Res.

2009, 42, 822-831. b- A. J. L. Villaraza, A. Bumb, M. W. Brechbiel, Chem. Rev.

2010, 110, 2921–2959.

7. M. Brinkley, Bioconjugate Chem. 1992, 3, 2-13.

8. P. Caravan, J. J. Ellison, T. J. McMurry, R. B. Lauffer, Chem. Rev. 1999, 99, 2293-

2352.

9. K. C. Briley-Saebo, V. Mani, F. Hyafil, J. C. Cornily, Z. A. Fayad, Magnetic

Resonance in Medicine, 2008, 59, 721-730.

10. F. Liu, S. Laurent, H. Fattahi, L. Vander Elst, R. N. Muller, Nanomedicine. 2011, 6,

519–528.

11. R.Weissleder, Science. 2006, 312, 1168–1171.

12. M. Brust, M. Walker, D. Bethell, D. J. Schiffrin, R. Whyman, J. Chem. Soc. Chem.

Commun. 1994, 801.

13. P. Alexandridis, Chem. Eng. Technol. 2011, 34, 15–28.

14. L. Vigderman, E. R. Zubarev, Adv. Drug. Deliv. Rev. 2013, 65, 663–676. K. Saha, S.

S. Agasti, C. Kim, X. Li, V. M. Rotello, Chem. Rev. 2012, 112, 2739−2779. M. C.

Daniel, D. Astruc, Chem. Rev. 2003, 104, 293-346.

15. P. J. Debouttière, S. Roux, F. Vocanson, C. Billotey, O. Beuf, A. Favre-Réguillon, Y.

Lin, S. Pellet-Rostaing, R. Lamartine, P. Perriat, O. Tillement, Adv. Funct. Mater.

2006, 16, 2330- 2339.

16. C. Alric, J. Taleb, G. L. Duc, C. Mandon, C. Billotey, A. L. Meur-Herland, T.

Brochard, F. Vocanson, M. Janier, P. Perriat, S. Roux, O. Tillement, J. Am. Chem.

Soc. 2008, 130, 5908–5915.

17. J. A. Park, P. A. N. Reddy, H. K. Kim, I. S. Kim, G. C. Kim, Y. Chang, T. J. Kim,

Bioorg. Med. Chem. Lett. 2008, 18, 6135-6137.

18. J. A. Park, H. K. Kim, J. H. Kim, S. W. Jeong, J. C. Jung, G. H. Lee, J. Lee, Y.

196

Chang, T. J. Kim, Bioorg. Med. Chem. Lett. 2010, 20, 2287-2291.

19. L. Moriggi, C. Cannizzo, E. Dumas, C. Mayer, A. Ulianov, L. Helm, J. Am. Chem.

Soc. 2009, 131, 10828-10829.

20. M. Marradi, D. Alcantara, J. M. d. l. Fuente, M. L. Garcia-Martin, S. Cerdan, S.

Penades, Chem. Comm. 2009, 3922-3924.

21. M. F. Warsi, R. W. Adams, S. B. Duckett, V. Chechik, Chem. Commun. 2010, 46,

451–453.

22. Y. T. Lim, M. Y. Cho, B. S. Choi, J. M. Lee, B. H. Chung, Chem. Comm. 2008, 4930-

4932.

23. H. K. Kim, H. Y. Jung, J. A. Park, M. I. Huh, J. C. Jung, Y. Chang, T. J. Kim, J.

Mater. Chem. 2011, 20, 5411–5417.

24. G. H. Lee, Y. Chang, T. J. Kim, Eur. J. Inorg. Chem. 2012, 1924–1933.

25. C. F. G. C. Geraldes, A. M. Urbano, M. C. Alpoim, A. D. Sherry, K. T. Kuan, R.

Rajagopalan, F. Maton, R. N. Muller, Magn. Reson. Imaging. 1995, 13, 401-420.

26. R. Fournaise, C. Petitfaux, Talanta. 1987, 34, 385–395.

27. D. L. Leussing, G. S. Albert, J. Am. Chem. Soc. 1960, 82, 4458-4461.

28. G. Crisponi, A. Diaz, V. M. Nurchi, T. Pivetta, M. J. Tapia Estevez, Polyhedron.

2002, 21, 1319-1327.

29. H. K. Kim, H. Y. Jung, J. A. Park, M. I. Huh, J. C. Jung, Y. Chang, T. J. Kim, J.

Mater. Chem. 2010, 20, 5411–5417. K. Kimura, S. Takashima, S. Ohshima. J. Phys.

Chem. B. 2002, 106, 7260-7266.

30. Z. Baranyai, E. Gianolio, K. Ramalingam, R. Swenson, R. Ranganathan, E. Brucher,

S. Aime, Contrast Media Mol. Imaging. 2007, 2, 94–102.

31. L. J. Garcés, C. R. Castro, C. David, S. Madurga, F. Mas, I. Pastor, J. Puy, J. Phys.

Chem. B. 2009, 113, 15145–15155.

32. C. David, E. Companys, J. Galceran, J. L. Garcés, F. Mas, C. R. Castro, J. Salvador, J.

Puy, J. Phys. Chem. B. 2007, 111, 10421–10430.

33. P. Gans, A. Sabbatini, A. Vacca, Hyperquad simulation and Speciation. 1998-2000,

Protonics software.

34. T. Koike, S. Kajitani, I. Nakamura, E. Kimura, M. Shiro, J. Am. Chem. Soc. 1995,

117, 1210-1219.

35. S. Laurent, L. Vander Elst, C. Henoumont, R. N. Muller, Contrast Media Mol.

Imaging. 2010, 5, 305-308.

197

36. L. Sarka, L. Burai, E. Brucher, Chem. Eur. J. 2000, 6, 719-724.

37. J. M. Idée, M. Port, I. Raynal, M. Schaefer, S. Le Greneur, C. Corot, Fundam. Clin.

Pharmacol. 2006, 20, 563–576.

38. T. D. Smith, A. E. Martell, J. Am. Chem. Soc. 1972, 94, 4123–4129.

39. C. Peter, D. L. Yaroslav, I. P. Alexander, G. Rajaraman, Inorg. Chem. 2006, 45,

3632-3638.

40. E. F. Hasty, L. J. Wilson, D. N. Henrickson, Inorg. Chem. 1978, 17, 1834.

41. M. Lachkar, R. Guilard, A. Atmani, A. De Cian, J. Fischer, R. Weiss, Inorg. Chem.,

1998, 37, 1575.

42. R. Schneider, A. Riesen, T. A. Kaden, Helv. Chim. Acta, 1986, 69, 53.

43. N. Domracheva, A. Mirea, M. Schwoerer, L. T. Lorente, G. Lattermann,

ChemPhysChem. 2005, 6, 110–119.

44. S. Laurent, L. Vander Elst, F. Copoix, R. N. Muller, Invest. Radiol. 2001, 36, 115-

122.

45. S. Gu, H. K. Kim, G. H. Lee, B. S. Kang, Y. Chang, T. J. Kim, J. Med. Chem. 2011,

54, 143-152.

46. Z. Baranyai, Z. Pálinkás, F. Uggeri, E. Brücher, Eur. J. Inorg. Chem. 2010, 1948–

1956.

47. L. Sarka, L. Burai, E. Brücher, Chem. Eur. J. 2000, 6, 719–724.

48. J. A. Peters, Inorg. Chem. 1988, 27, 4686–4690.

199

Chapter-IV

Experimental Section

201

A. Synthesis of ligands L1H4, L2H3, L@1H5 and L@

2H3 Macrocyclic ligands L1H4 and L2H3 were synthesized at University of Hull, U.K, by

S.J Archibald group. Linear ligands L@1H5 and L@

2H3 (L@1H5 incorporated gold

nanoparticles) were synthesized at Université de Franche-Comté by Stéphane Roux group.

L@4H5 was obtained commercially from FLUKA.

1. Synthesis of L1H4 1,4,7-tris(tert-butoxycarboxymethyl)-10-(2-methylbenzimidazolyl)

tetraazacyclododecane (1 g, 2.09 mmol) was dissolved in 6M HCl (200ml) and heated

under reflux for 18 hours. The solvent was removed under reduced pressure to give the

crude product which was purified by dissolving impurities into diethyl ether and

decantation (3 x 100 cm3) to leave a light brown solid (3.0g, 98%).

2. Synthesis of L2H3 1,4,7-tris(tert-butoxycarbonylmethyl)-10-(1-(4-nitrobenzyl)-2-methyl benzimidazole)-

1,4,7,10-tetraazacyclododecane (1.3 g, 1.67 mmol) was dissolved in 6M HCl (50 ml) and

heated under reflux for 18 hours. The reaction was then concentrated in vacuo and diethyl

ether (3 x100 ml) was added and decanted off to remove any impurities to yield a light brown

solid (1.3 g, 98%).

3. Synthesis of L@1H5 In a 250 mL round-bottomed flask, 2 g (5.6 × 10–3 mol) of DTPA-BA (DTPA

Bisanhydride) was dissolved in 40 mL of DMF and heated to 70 °C. In another flask, 1.4 g

(1.23 × 10–2 mol) of aminoethanethiol was dissolved in 30 mL of DMF and 1.74 mL of

triethylamine. This solution was added to the round-bottomed flask and stirred magnetically

at 70 °C overnight. Subsequently, the solution was cooled to room temperature and placed in

an ice bath. A white powder (NEt3·HCl) was seen to precipitate out and was filtered. The

filtrate was then concentrated at low pressure. After addition of the filtrate to a chloroform

solution, a white precipitate was formed. After filtration of the solution, washing with 50 mL

chloroform, and drying under vacuum, DTDTPAwas obtained as a white powder (90%

yield).

202

4. Synthesis of L@2H3 The synthesis described by Brust et al.[1] consists of reducing HAuCl4·3H2O with

NaBH4 in the presence of thiols (stabilizers), which, by adsorption on the growing particles,

control the size and the stability of the colloids. In a typical preparation of gold nanoparticles,

200 mg (5.1×10–5 mol) of HAuCl4·3H2O dissolved in 120 mL methanol was placed in a 500

mL round-bottomed flask. 482 mg (9.4×10–5 mol) of DTDTPA in 40 mL of methanol and 2

mL of acetic acid was added to the gold salt solution while continuously stirring the mixture.

The mixture changed color from yellow to orange. After 5 min, 192 mg (5×10–5 mol) of

NaBH4 dissolved in 14 mL water was added to the orange mixture under vigorous stirring at

room temperature. At the beginning of the NaBH4 addition, the solution became dark brown,

followed by the appearance of a black flocculate. The mixture was stirred vigorously for 1 h

before adding 5 mL of 1 M aqueous HCl solution. After partial removal of the solvent under

reduced pressure at a maximum temperature of 40°C, the precipitate was filtered using a

polymer membrane and washed thoroughly and successively with 0.01 N HCl, water, and

diethyl ether. The resulting black powder (L@1H5 grafted on goldnanoparticle) was dried and

either stored as a solid or dispersed in 10 mL of 0.01 M NaOH solution (20 mg of dry

powder).

The salts used for the experiments were of the highest analytical grade diluted in

water solution. The concentration of GdCl3, EuCl3, CuCl2, and ZnCl2 stock solutions was

determined by complexometric titration with standardized Na2H2EDTA solution, with the use

of xylenol orange (GdCl3, EuCl3, ZnCl2) and murexide (CuCl2) indicators. 0.1 mol.L-1

solutions of stock solutions were further diluted to 5×10-3 mol.L-1 (for GdCl3, EuCl3) and

1×10-2 mol.L-l (for CuCl2, ZnCl2) in 0.1 mol.L-1 NMe4Cl for potentiometric studies.

B. Elemental analysis

The centesimal composition C, H and N were obtained using a "Perkin-Elmer 2400"

analyzer.

In macrocyclic ligands, the exact amount of HCl present could not be identified as the

molecular mass of two water molecules is equal to the molecular mass of one HCl molecule.

To precise this aspect, argentometric titrations were performed for both ligands.

203

Argentometric titrations are performed using a Metrohm 713 pH using a combination

electrode Metrohm AG 9101 (Ag / AgCl / KCl). A known concentration of AgNO3 (5×10-3

mol.L-1) is used for the titration.

To avoid complexation of silver with ligand, several argentometric titrations of ligand

solution with excess of Cu(II) were done.

a. titration of 3mL of ligand solution with H2O

b. titration of 3mL ligand solution with excess of Cu

c. titration of 3mL ligand solution with excess of Cu and HNO3

%C %N %H %Cl %S Formula

L1H4 % exp 36.30 11.52 5.77

% theoretical 36.15 11.50 6.21 24.19 C22H32N6O6, 5HCl, 4H2O

L2H3 % exp 43.02 13.07 5.67

% theoretical 42.94 12.08 6.21 13.11 C29H37N7O8, 3HCl, 5H2O

L@1H5 % exp 39.92 12.63 6.88 10.33

% theoretical 39.49 12.8 6.76 11.70 C18H33N5O8S2, 2H2O

C. Potentiometric study

1. Materials used In the study of ligands and complexes by potentiometry, we use computerized

titrations including:

a Metrohm 713 pH meter to measure the pH millesimal using a combination

electrode Metrohm AG 9101 (Ag / AgCl / KCl).

a microburette, microprocessor METROHM Dosimats 665 equipped with a

replaceable unit of 1mL. The end of the tip is immersed in the solution provided

with an anti-diffusion Teflon valve.

a measurement cell connected to a water flow cryostat. The measured temperature

at 0.1°C is set at 25°C using a Pt100 sensor connected to the pH meter.

a scanning device which enables argon to isolate the solution from external

environment.

204

The micro-burette and the pH meter are controlled by a Software developed at

laboratory by J. C. PIERRARD and J. RIMBAULT.This system automates the potentiometric

titrations and acquire pairs of measurements (volume of titrant delivered, pH) for further

treatment.

The measuring system is calibrated at pH = 2.000 by reference with hydrochloric acid

as exactly equal to 10-2 mol.L-1 and ionic strength of 0.1 NMe4Cl solution.

2. Determination of ionic product of water

The ionic product of water at an ionic strength of 0.1 NMe4Cl and at 25°C was determined

with precision. This is indeed necessary because a small variation of this parameter affects

the accuracy of the protonation constants values of ligands or complex formation.

The experimental value of the ion product is obtained by using the titration curves of acetic

acid at various concentrations by tetramethyl ammonium hydroxide. Processing the curves by

PROTAF software allows simultaneous fitting of the concentration and the acidity constant

of acetic acid, and also the ionic product of water. In our conditions, the value of the ionic

product of water is pKe = 13.78 0.01.

3. Description of h versus pH

Protometric titrations could be depicted as curves of h versus pH, where h is the average

number of protons bounded per mole of ligand (Equation (a)).[2]

HBLL

CCOHHCnC

h 1 Equation (1)

where: n = number of protons on the ligand (in its neutral form); CB = concentration of the

base; CH = added acid concentration; CL = ligand concentration

4. Calculation of software PROTAF[3] and HYPERQUAD 2008[4]

The determination of equilibrium constants is performed using multiparameter

PROTAF refinements program based on a method of least squares. This method is used to

minimize the weighted sum of squared residuals on experimental variables (volume of base

205

added, pH measured). Residues on experimental variables are defined by RVi = Vi exp - Vi

refined and RpHi = pHi exp - pHi refined with weighted coefficients WVi and WpHi related to the

accuracy of measurements on the volume and pH.

The weighted sum of squared residuals is defined by the equation:

i

2II

2VV RWRW iiii

S

Simultaneously with equilibrium constants, it is possible to refine the other titration

parameters such as the ionic product of water and the concentrations of the solutions used.

The refined equilibrium constants are the overall formation constants mlh

corresponding to the equilibrium:

m M + ℓ L + h H MmLℓHh

with mlh = hm

hm

HLMHLM1

1

In the particular case of the study of ligand alone (m = 0, ℓ = 1), the protonation

constants 01h are connected to the acidity constants Ka by the equation:

0lh = Kah ii

h

11

1

The precision of the results given in paranthesis in Chapter–II and III for overall

formation constants is a standard deviation which was given by the software. For the stability

constants the standard deviations were calculated by propagation of the variance method.

The computer program HYSS was used to obtain the species distribution curves.[3]

5. Preparation of solutions

To maintain constant activity coefficients of the various species present in solution,

measurements were made an ionic strength fixed by adding a salt background. All solutions

are prepared at an ionic strength of I = 0.1 NMe4Cl. The cyclic ligands being known to have a

slow rate of complexation, ‘out-of-cell’ or ‘batch’ method is required, while for linear ligands

a classic titration is used to determine the overall stability constants with metal complexes.

For ligands L1H4 and L2H3, ‘out-of-cell’ method is required, while for ligands L@1H5,

L@2H3 and L@

3H5 the titration is direct.

206

a) Preparation and titration of L1H4 and L2H3 solutions for the determination of acidity constants

The mother solution of ligand (L1H4) is prepared at a concentration of 2×10-3 mol.L-1 with an

ionic strength of 0.1 (NMe4Cl). The exact concentration of the ligand solution is determined

by titration with a base (NMe4OH, 5×10-2 mol.L-1). Various volumes such as 2, 3 and 4ml of

the mother solution were titrated by NMe4OH.

In the case of L2H3, the concentration of the mother solution was reduced to 10-4 mol.L-1 due

to its solubility issues in water. The solutions are prepared by saturation, and then filteration

of solution before utilisation. Exact concentration of the solution is determined by titration

with base (NMe4OH).

b) Preparation and titration of L@1H5, L@2H3 and L@4H5 solutions

The mother solutions of ligands were prepared at a concentration of 2×10-3 mol.L-1 with an

ionic strength of 0.1 (NMe4Cl). The exact concentration of the ligand solution is determined

by titration with a base (NMe4OH, 5×10-2 mol.L-1). The determination of protonation

constants was done basing on several titrations with NMe4OH by varying the volumes of the

mother solution such as 2, 3 and 4ml.

c) Preparation and titration of transition metal and lanthanide complexes of L1H4 and L2H3

i. Batch methods

500 mL of L1H4 solution is prepared at a concentration of 10-3 mol.L-1 in hydrochloric

acid 5×10-3 mol.L-1. The exact concentration of the solution is determined by a traditional

titration (4ml). With 96 mL of this solution, different volumes of Gadolinium /Europium

chloride (5×10-3 mol.L-1) or Zinc /copper chloride (1×10-2 mol.L-1) are added in respective

flasks in such a way to obtain the ratios [L]/[M] at 1.05, 1.15, 1.24 respectively. The final

volume (128 ml) is obtained by adding NMe4Cl 0.1 mol.L-1. From this mother solution, 26

tubes are prepared (5mL of this solution in each tube). Base NMe4OH is added in each tube

in such a manner that the pH in a tube is superior to the pH in the earlier tube. Four such

207

series were prepared, each with a respective metal ion. The four series are identical in

constitution and in pH. These four series are stocked in argon at 37°C for four weeks.

After the incubation period, the tubes are cooled to room temperature for one day. The

pH of each container of the first series is then measured. The same procedure of measurement

is followed in the other series.

For metal complexes with L2H3, the same procedure is followed with the preparation

of complexes in solution in ratios [L]/[M] at 1.05, 1.15, 1.24 respectively. The concentration

of mother solution of ligand was 8.5×10-5 mol.L-1

ii. Classic titration

After measuring the pH of all tubes in a series, solution with pH ~ 4 was titrated with

base NMe4OH in a classic method to follow the evolution of complexes, as the batch method

is restricted until pH 7.

d) Preparation and titration of transition metal and lanthanide complexes of L@1H5, L@2H3 and L@4H5

A solution of ligand is prepared at a concentration of 2×10-3 mol.L-1, the exact

concentration of the solution is determined by a classic titration. For linear ligands, a classic

titration is done to determine the stability constants of complexes. 2ml of the ligand solution,

volumes Gadolinium /Europium chloride (5×10-3 mol.L-1) or Zinc /copper chloride (1×10-2

mol.L-1) are added in respective titrations in such a way to obtain the ratios [L]/[M] at 1.05,

1.15, 1.24 respectively. The total volume (4ml) for each of these solutions are adjusted by

adding NMe4Cl 0.1 mol.L-1.

208

Tables:

Deprotonation constants of metal complexes (M/L@1H5, M/L@

2H3 and M/L@4H5)

Table 1: Deprotonation constants of M/L@1H5, M/L@

2H3 and M/L@4H5 complexes (where M

= Cu(II)/Zn(II)/Ca(II)), T = 25°C, I = 0.1 (NMe4Cl) )

L@1H5 L@

2H3 L@4H5

Cu(II) Zn(II) Ca(II) Cu(II) Zn(II) Ca(II) Cu(II) Zn(II) Ca(II)

Log K

MLH4= MLH3 + H 4.4

MLH3=MLH2 + H 3.5 4.1 5.2

MLH2=MLH + H 5.3 8.2 8.9 5.34 5.9 7.1

MLH=ML + H 8.6 9.6 10.2 9.26 10.5 10.8 4.68 5.58 6.29

ML=MLH-1+H+ 11.83 12.71

M2LH=M2L + H 4.0

M2L=M2LH-1 + H 9.2

M2LH-1=M2LH-2+H+ 10.2

ML2H4=ML2H3+H 8.4

ML2H3=ML2H2+H 8.7

ML2H2=ML2H+H 10.9

ML2H=ML2+H 9.9

D. NMR Spectroscopy

The 1H NMR spectra were recorded on DPX 300 (300 MHz) at the common NMR

service of University of Reims. The chemical shifts expressed in ppm, are counted positively

downfield and are given relative to tetramethylsilane (TMS).

The evolution of 1H NMR spectra in function to pH of ligand L1H4 or Zinc complexes

was performed from the solutions of ligand or the complex at 10-2 mol.L-1. The ligands and

metal complexes were solubilised in D2O and the pD of these solutions were adjusted by

adding 4% NaOD or 3.5% DCl. These solutions were obtained by dilution of mother

solutions of NaOD 40% or DCl 35% in D2O. pH values are then calculated by the equation

pH = pD – 0.40[5]

209

E. EPR spectroscopy

The EPR spectra were performed using the spectrometer "BRUKER ESP 300e" with a

gauss meter "BRUKER E035 M" and X-band at 9.43 GHz. The spectra were recorded in

solution (water / ethanol) (80/20)at 150 K (variable temperature unit: BRUKER ER4111VT).

The simulation of the spectra was performed using the software XSophe 1.1.4[6] developed by

Prof. G. Hanson of the University of Queensland, Brisbane, Australia.

The Cu complexes of L1H4 (5×10-4 mol.L-1) and L2H3 (7×10-5 mol.L-1) were recorded

using a frozen solution (5 mL) prepared (water/ethanol) at 150K. pH of these solutions were

adjusted by NMe4OH (5×10-2 mol.L-1) and HCl (1×10-2 mol.L-1). In the case of linear ligands

L@1H5 (10-3 mol.L-1) and L@

4H5 (2×10-3 mol.L-1), EPR spectra were recorded at various mole

ratios of metal to ligand (upto 2:1 ratio of M/L). 5 mL of ligand solution was prepared in a

mixture of water and ethanol, 0.5 equivalent of CuCl2 was added and passed through EPR

spectrometer. The addition of CuCl2 upto 2:1 ratio of Cu(II) to ligand was performed and the

spectra were recorded.

F. Relaxometry

The measurements were performed at 40 MHz and 37°C on a spin analyser the

Minispec MQ-40. Gadolinium complex solution in tampon phosphate (MERCK) was added

with Zinc chloride at a ratio of 1:1. The concentration of gadolinium complex is 2.5×10-3

mol.L-1. When Gd complex in phosphate buffer and aqueous solution of ZnCl2 were mixed, a

little turbidity appeared. The mixture was stirred (homogenized) and taken for the

measurements. The samples were contained in 7mm o.d. Pyrex tubes and kept at 37°C in a

dry block between measurements (at least up to 4320 min or 4 days). kobs determination and

their half life calculations are discussed in the forthcoming section.

Preparation of phosphate buffer

Phosphate buffer solution was obtained from normadose Merck. 500 mL of this buffer

contains a concentration of 0.026 mol.L-1 KH2PO4 and 0.041 mol.L-1 of Na2HPO4.

mailto:L@3H5(2.10-3

210

G. UV visible spectroscopy

The spectrometer utilised for recording the spectra is a spectrophotometer

“SHIMADZU UV-2410-PC”. Spectra were recorded at 25°C in the form of solutions.

1. Evolution of absorbance according to pH

The evolution of UV signals of benzimidazole moiety in ligand L1H4 or different

metal complexes(Gd(III), Eu(III), Cu(II) and Zn(II)) were followed with the solutions of

ligand alone or the metal complexes at 5×10-5 mol.L-1 (Ionic strength I = 0.1(NMe4Cl)). The

evolution of UV signal of benzimidazole in ligand and metal complexes according to pH was

adjusted by adding NMe4OH (5×10-2 mol.L-1) and HCl (10-2 mol.L-1). For ligand L2H3, the

concentrations used were under 8.5×10-5 mol.L-1, prepared under similar conditions than

L1H4.

2. Evolution of absorbance according to time (Kinetics experiments)

Linear ligands were analysed for their kinetic inertness under pseudo first order

conditions by adding excess of competitive metal ion (ZnCl2 in our case). The Gd complex

and zinc are mixed and the reaction is followed on UV spectrophotometry, since Gd

complexes show no absorbance in the UV range. Formation of zinc complexes could be seen

in the UV region. The increase in absorbance according to time is monitered. These

experiments were done by varying the concentration of Zinc added, eventually for calculating

the rate of reactions. The calculation of kobs and investigation of transmetallation mechanism

are described in the coming section.

a) Determination of kobs

The following transmetallation reaction has been studied at stoichiometry (1:1) conditions by

relaxometry and in the presence of excess Zn(II) by UV spectroscopy.

211


with :

2ZnGdLkdt

GdLdv T

T

with : GdLZnGdLZnHGdLHGdLHGdLGdL 2T

The measurements were performed by relaxometry at pH 7.4 ([GdLT] = [Zn(NO3)2] =

2.5×10-3 mol.L-1) in the presence of phosphate buffer. In this case, the precipitation of GdPO4

(pKs (GdPO4) = 22.26) and Zn3(PO4)2 (pK (Zn3(PO4)2) = 35) were observed. UV-Visible

spectroscopy measurements were performed at a concentration of [GdLT] = 2×10-4 mol.L-1 in

the presence of ZnCl2 in excess ([ZnCl2] = 4×10-3 mol.L-1, 6×10-3 mol.L-1, 8×10-3 mol.L-1 and

10×10-3 mol.L-1). In both cases, the free metal ion concentration of Zn(II) in solution is kept

constant and the rate law is written in simplified first order kinetics:

Tobs

T GdLkdt

GdLdv

Determination of kobs depends only on the concentration GdLT


ti = 0 C0 constant - -

t C = C0-x x x

te = ∞ Ce = C0-xe xe xe

In such case, the determination of formation constants of GdL and ZnL highlights the

formation of ZnL at the expense of GdL. In this case, at t = ∞, the value of Ce≈ 0 can be

substituted and one can consider C0 = xe

In spectroscopy, A = f(C); in relaxometry R1 = f(C), however both are followed as a function

of time.

212

A0 = C0 l

At = C0 l- x l

Ae = C0 l- xe l

or

0t AA

x

0e

eAA

x

and

tee

AAxx

R10 = f C0

R1t = fC0 - fx

R1e = fC0 - fxe

or

fRR

x 01t1

fRR

x 01e1e

and f

RRxx t1e1

e

where Ae, A0 and At are the experimental measurements obtained by spectroscopy and R1e,

R10, R1t are the experimental measurements obtained by relaxometry. By analogy, these

similar parameters were denoted combinedly as Ae, A0 and At for both UV spectroscopy and

relaxometry.

In the case of a 1st order kinetics

Tobs

T GdLkdt

GdLd

after integration, the kinetic law is written as:

)x()xx(

lntkCC

lne

eobs

0

t

or

)AA()AA(

lntk0e

teobs

where Ae, A0 and At are the experimental measurements obtained by spectroscopy and

relaxometry.

For each experiment, kobs is obtained by plotting the linear regression using Excel software

)AA()AA(

ln0e

te

versus time. (Figure IV-1)

213

a. b.

y = -0.0039x + 0.1965R2 = 0.9985

-1

-0.8

-0.6

-0.4

-0.2

00 50 100 150 200 250 300 350

t (min)

-0.6

-0.4

y = -0.0022x - 0.1157R2 = 0.9922

-3

-2.5

-2

-1.5

-1

-0.5

00 250 500 750 1000 1250 1500

t (s)

-1.5

-1

Figure IV-1 : Determination of kobs a. by relaxometry [Zn2+]=[GdL@

1T]= 2.5×10-3

mol.L-1 , b. by UV spectrometry for [Zn(II)] = 6×10-3 mol.L-1, [GdL@1] = 5×10-4 mol.L-1

pH = 6.0

Since the kinetics of first order is determinedobs

2/1 k2lnlnt

The results obtained by relaxometry at pH = 7.4 in stoichiometry 1:1; kobs = 3.910-3 min-1 or

6510-4 s-1 (t1/2 = 177 min).

By spectroscopy; kobs values are shown in Table IV-2

Table IV-2: Values of kobs (10-4 s-1) obtained by linear regression.

[Zn2+]

(mM)

pH=5.8 pH=6.0 pH=6.2 pH=6.5 pH=6.9 pH=7.0 pH=7.1

4 24 21 16 11 26 24 29

6 26 22 17 11 24 21 25

8 28 23 18 13 14 14 19

10 30 25 20 16 8 9 14

or t1/2 values are represented in Table IV-3

214

Table IV-3: Values of t1/2 obtained by linear regression.

[Zn2+]

(mM)

pH=5.8 pH=6.0 pH=6.2 pH=6.5 pH=6.9 pH=7.0 pH=7.1

4 288.8 330.1 433.2 630.1 266.6 288.8 239.0

6 266.6 315.1 407.7 630.1 288.8 330.1 277.3

8 247.6 301.4 385.1 533.2 495.1 495.1 364.8

10 231.0 277.3 346.6 433.2 866.4 770.2 495.1

b) Determination of rate constants

The calculations were performed by using the computer program Micromath

Scientist, version 3.0. Scientist is a comprehensive modeling, data analysis and curve fitting

(regression) in which linear or non-linear regression minimizes the sum of the squares of the

differences between the observed values and the calculated values of the model using the

“best fit” parameters.

The evolutions of kobs according to the metal concentration and according to pH are

reported in the Figure IV-2.

a. b.

0,004 0,005 0,006 0,007 0,008 0,009 0,010

5

10

15

20

25

30

ko

bs (1

0-4 s

-1)

[Zn2+

] (mol L-1)

pH=5.8

pH=6

pH=6.2

pH=6.5

pH=6.9

pH=7.0

pH=7.1

kobs

= f([Zn2+

])

5,6 5,8 6,0 6,2 6,4 6,6 6,8 7,0 7,2

6

8

10

12

14

16

18

20

22

24

26

28

30

32

ko

bs (

10

-4 s

-1)

pH

c=4 e-3M

c=6e-3M

c=8e-3M

c=10e-3M

kobs

= f(pH)

Figure IV-2: a. kobs as a function of [Zn2+] b. kobs as a function of pH

The plot shows a different tramsmetallation mechanism before and after pH = 6.8.

The evolution of kobs after pH = 6.8 is not successive. The reason might be due to the

improper regulation of hydroxylated species formation. Hence, the transmetallation

mechanism from pH = 5.8 to 6.8 was investigated.

215

According to the species distribution curves of Gd(III)-L@1H5 in the pH range

between 5 and 7.5, the major species prevail are: GdLH2 and GdLH ( noted as L (instead of

L@1) for simplification.

The rate determining reactions can include:

202 GdLHk....GdLH 0 v

HGdLHk.....HGdLH 212

1 v

222

22 ZnGdLHk.....ZnGdLH 2 v

2223

22 ZnGdLHk.....Zn2GdLH 3 v

GdLHk....GdLH 4 4 v

HGdLHk.....HGdLH 5

5 v

26

2 ZnGdLHk.....ZnGdLH 6 v

227

2 ZnGdLHk.....Zn2GdLH 7 v

GdLZnHk....GdLZnH 8 8 v

HGdLZnHk.....HGdLZnH 9

9 v

210

2 ZnGdLZnHk.....ZnGdLZnH 10 v

GdLZnk....GdLZn 11 11 v

HGdLZnk.....HGdLZn 12

12 v

213

2 ZnGdLZnk.....ZnGdLZn 13 v

Each step is considered independent from one another to determine the predominant slow

steps. Tobsi i GdLkvv

then T

i iobs GdL

vk

with


216


with

23

2112GdLH HLGd

GdLH2

, HLGd

GdLH3111GdLH

, LGd

GdL3110GdL

HLZnGd

GdLZnH231111GdLZnH

,

231110GdLZn ZnLGdGdLZn

where

110

2112

2HGdL

GdLH

,

110

111 HGdLGdLH

110

21111 HZnGdL

GdLZnH

,

110

21110 ZnGdL

GdLZn

Therefore, T

i iobs GdL

vk

21110

211111111

2112110

229

28

227

265

2224

223

22

31

ZnZnHHHZnAZnAZnHAZnHAHAZnHAZnHAHAHA

with :

11211 kA 1110121111811166 kkkA

111511202 kkA 11101011177 kkA

1111911223 kkA 1110118 kA

11234 kA 1110139 kA

11145 kA

In the case of the study of kobs as a function of the concentration of Zn2+, the expression of

kobs becomes:

2

54

223

221

ZnBBZnBZnBB

kobs

with:

217

B1=A1×[H]3 + A2×[H]2 + A5×[H] B4= 110+112×[H]2 +111×[H]

B2= A3×[H]2 + A6×[H] + A8 B5= 1111×[H] +1110

B3= A4×[H]2 + A7×[H] + A9

In the case of kobs as a function of pH, the expression of kobs becomes:

76

25

432

23

1obs PHPHP

PHPHPHPk

with:

P1=A1 P5= 112

P2= A2 + A3×[Zn2+] + A4×[Zn2+]2 P6= 111 +1111×[Zn2+]

P3= A5 + A6×[Zn2+] + A7×[Zn2+]2 P7= 1110×[Zn2+] +110

P4= A8×[Zn2+] + A9×[Zn2+]2

218

References:

1. M. Brust, J. Fink, D. Bethell, D. J. Schiffrin, C. Kiely, J. Chem. Soc. Chem. Commun.

1995, 1655.

2. a. J. Bjerrum, Metal-ammine formation in aqueous solution. 1941, Copenhagen:

Haase. b. Ahmed Messadi, Aminou Mohamadou, Isabelle Déchamp-Olivier &

Laurent Dupont, J. Coord. Chem. 2012, 65, 2442-2458.

3. a. R. Fournaise, C. Petitfaux, Talanta 1987, 34, 385-395. b. R. Fournaise, C.

Petitfaux, Analusis. 1990, 18, 242-249.

4. a. P. Gans, A. Sabatini, A. Vacca, Coord. Chem. Rev. 1996, 43, 1739. b. L. Alderighi,

P. Gans, A. Ienco, D. Peters, A. Sabatini, A. Vacca, Coord. Chem. Rev. 1999, 184,

311.

5. A.K. Covington, M. Paabo, R. A. Robingson, R. G. Bates, Anal. Chem. 1968, 40, 700.

6. M. Griffin, A. Muys, C. Noble, D. Wang, C. Eldershaw, K. E. Gates, K. Burrage, G.

R. Hanson, Mol. Phys. Rep. 1999, 26, 60.

http://www.tandfonline.com/action/doSearch?action=runSearch&type=advanced&searchType=journal&result=true&prevSearch=%2Bauthorsfield%3A%28Messadi%2C+A%29

http://www.tandfonline.com/action/doSearch?action=runSearch&type=advanced&searchType=journal&result=true&prevSearch=%2Bauthorsfield%3A%28Mohamadou%2C+A%29

http://www.tandfonline.com/action/doSearch?action=runSearch&type=advanced&searchType=journal&result=true&prevSearch=%2Bauthorsfield%3A%28D%C3%A9champ%5C-Olivier%2C+I%29

http://www.tandfonline.com/action/doSearch?action=runSearch&type=advanced&searchType=journal&result=true&prevSearch=%2Bauthorsfield%3A%28Dupont%2C+L%29

http://www.tandfonline.com/loi/gcoo20?open=65#vol_65

219

Conclusion

221

When coordination complexes are developped for medical applications, their

thermodynamic stability and their kinetic inertness are parameters that must be assessed. In

this work, the goal was to evaluate the thermodynamic stability and kinetic inertness of

gadolinium complexes designed for MRI imaging. The envisaged ligands were

polyaminocarboxylate ligands based either on macrocyclic or on linear backbones.

Macrocyclic ligands were based on a DO3A cavity and the functional groups were methyl-

benzimidazole derivatives. Linear ligands were based on a DTPA bisamide and the functional

groups were thiol functions, in order to ensure ligand grafting at the surface of a gold

nanoparticle.

The approach consisted of few steps, in which the first step is the evaluation of the

acid-base behavior of each system and then, on the basis of the protonation constants

knowledge, the evaluation of their affinities towards a set of metallic cations. If the target

cation in the study is Gd(III), ligand affinities were determined also towards potentially

competitors such as Cu(II) and Zn(II). When possible, species in solution were characterized

with the help of spectroscopic techniques such as UV, 1H NMR, EPR and fluorescence

spectroscopic techniques. The second step consisted of the evaluation of Gd(III) complexes

kinetic inertness in the presence of Zn(II) as a competitor, by means of relaxometry.

For macrocyclic systems L1H4 and L2H3, whatever the cation, L1H4 was the most

stable. Moreover their stability followed the rising order M-DO3A<M-L1H4<M-DOTA. This

indicated that the additional benzimidazole group reinforced the complexation ability of the

DO3A backbone but its coordinating imine nitrogen atom is a less efficient donor than the

carboxylate function in DOTA. From the kinetic inertness point of view, Gd(III)-L1H4

exhibited a very satisfactory inertness since it was similar to the inertness of the reference

Gd(III)-DOTA. This result can be interpreted by the fact that L1H4 is well adapted to the

stereoelectronic demand of Gd(III) in terms of the nature and the number of coordinating

atoms and in terms of ligand preorganization.

For L@1H5 and L@

2H3 systems, potentiometric studies have shown that Gd(III)

complexes are more stable than Zn(II) and Cu(II) ones. It is interesting to notice that acid-

base and complexation properties of L@1H5 were greatly modified when the ligand is grafted

onto the nanoparticle. Indeed, the ligand basicity was enhanced and then the complex

stabilities. To explain this result, one could suggest that this improvement could be due to the

positive influence of the ligand packing at the surface that stabilise probably by cooperative

effects, the protonated forms of the ligand and the complexes. This gain in stability was

222

accompanied by a gain in kinetic inertness for Gd(III)-L@1H5 grafted onto the nanoparticle.

Effectively, relaxometric studies showed that the half-life of the complex when grafted onto

the nanoparticle was two times more important than for the complex alone. These results are

important in the context of the utilization of these nanoparticles in living organisms for

imaging applications. They also call for caution in assuming that a given property defined for

a free complex could be extrapolated when this complex is embedded in a more organized

system such as the network studied here.

223

Conclusion générale

225

Lorsque les complexes de coordination sont développés dans un but de diagnostic, la

question de leur stabilité thermodynamique et de leur inertie chimique vis-à-vis de réactions

de transmétallation est une question majeure, qui conditionne le développement et

l’utilisation de ces systèmes. Dans notre cas, l’objectif du travail était d’évaluer la stabilité

thermodynamique et de l’inertie chimique de complexes de gadolinium pour l’imagerie IRM.

Les ligands utilisés sont des ligands polyaminocarboxylates macrocycliques ou linéaires

fonctionnalisés. Dans le cas des ligands macrocycliques, la cavité est de type DO3A et les

groupements fonctionnels sont des groupements de type méthyl-benzimidazole. Pour les

ligands linéaires, le ligand est un DTPA bisamide et les groupements fonctionnels sont des

fonctions thiol, permettant d’envisager le greffage de ces ligands sur des nanoparticules d’or.

La démarche utilisée a consisté à étudier dans une première étape le comportement

acido-basique de tous les systèmes puis, sur la base des constantes de protonation obtenues,

d’évaluer l’affinité des différents ligands vis-à-vis d’un ensemble de cations métalliques. Si le

cation cible de l’étude est le gadolinium Gd(III), l’affinité des différents ligands a été

déterminée par rapport à ce cation et par rapport à des cations potentiellement compétiteurs

(Cu(II), Zn(II)). Lorsque cela a été possible, les différentes espèces complexes présentes en

solution ont été caractérisées en s’appuyant sur l’apport de résultats en spectroscopie UV, 1H

RMN, RPE et fluorescence. La seconde étape du travail a consisté à évaluer par relaxométrie

l’inertie chimique des complexes de Gd(III) en présence d’un ion compétiteur, le Zn(II).

Pour les systèmes macrocycliques L1H4 et L2H3, quel que soit le cation, les

complexes basés sur le ligand L1H4 sont les plus stables. De plus, la stabilité de ces

complexes suit l’ordre d’affinité croissant M-DO3A<M-L1H4<M-DOTA. Ceci indique que

l’adjonction du groupe benzimidazole renforce les capacités de complexation du macrocycle

DO3A bien que l’atome d’azote imine coordinant du groupe benzimidazole soit un moins

bon atome donneur qu’un groupement carboxylate du DOTA. Du point de vue inertie

chimique, le complexe Gd(III)-L1H4 présente une très bonne inertie chimique puisqu’elle est

comparable à celle de la référence Gd(III)-DOTA. Ce résultat s’explique par le fait que le

ligand L1H4 répond parfaitement aux exigences stéréoélectroniques du Gd(III) en termes de

nombre d’atomes coordinants et de préorganisation du squelette organique.

Pour les systèmes L@1H5 et L@

2H3, les études potentiométriques montrent à nouveau

que les complexes de Gd(III) sont plus stables que ceux de Zn(II) et de Cu(II). Le point à

remarquer dans cette étude est que les propriétés acido-basiques et les propriétés de

complexation du ligand L@1H5 sont modifiées lorsque celui-ci est greffé à la surface de la

226

nanoparticule. Ainsi, la basicité du ligand est renforcée, la stabilité des complexes s’en

trouvant de fait améliorée. Pour expliquer ce résultat, on peut suggérer que la structure de la

couche organique de ligands à la surface de la nanoparticule favorise la stabilisation des

charges introduites par un réseau de liaisons hydrogène et des réorganisations

conformationnelles. Ce gain de stabilité thermodynamique s’accompagne pour le complexe

greffé à la surface de la nanoparticule par un gain important d’inertie chimique puisque les

études relaxométriques montrent que, la demi-vie du complexe greffé est deux fois plus

importante que celle du complexe libre. Du point de vue de l’utilisation en imagerie IRM de

ces complexes de Gd greffés sur nanoparticules, ces gains en stabilité thermodynamique et en

inertie chimique sont certainement des atouts importants. Du point de vue analytique, ces

résultats contribuent aussi à montrer qu’on ne peut pas se baser sur les propriétés d’un

complexe en solution pour extrapoler ces propriétés au complexe engagé dans un système

plus organisé tel que le réseau étudié ici.

227

Etude à pH physiologique, des mécanismes de transmétallation de complexes linéaires et macrocycliques de gadolinium utilisés en IRM

L’objectif de ce travail est l’analyse de la stabilité thermodynamique et de l’inertie chimique de complexes métalliques avec des ligands ou des nanoparticules conçus pour des applications en IRM. Deux types de ligands polyaminocarboxylates ont été étudiés, ligands pour lesquels les unités complexantes sont soit linéaires soit macrocycliques.

Les ligands macrocycliques étudiés sont des ligands basés sur des squelettes DO3A, substitués par des entités benzimidazole (L1H4) ou p-nitrophenylbenzimidazole (L2H3). Les données thermodynamiques indiquent que les affinités de ces ligands vis-à-vis des ions de la première série de transition (Cu(II) et Zn(II)) ou vis-à-vis des lanthanides (Gd(III) et Eu(III)) sont plus élevées que celles des complexes correspondants avec le ligand DO3A. Ce renforcement d’affinité est corrélé avec la participation des groupements benzimidazole à la sphère de coordination de chacun des métaux. L’inertie chimique du complexe Gd(III)- L1H4 a ensuite été évaluée par relaxométrie en tampon phosphate, en présence d’une quantité équimolaire de Zn(II). Pour Gd(III)-L1H4, aucune réaction de ce type n’a été détectée, ce qui plaide en faveur de l’inertie chimique de ce complexe. Les ligands linéaires étudiés sont des dérivés dithiolés de ligands DTPA bisamide L@

1H5. Ces ligands ont été conçus pour être greffés sur des nanoparticules d’or. La stabilité thermodynamique des complexes de Cu(II), Zn(II) et Gd(III) utilisant les ligands L@

1H5 et L@1H5 greffé sur nanoparticule d’or (autrement appelé L@

2H3) suit l’ordre de stabilité croissant Zn(II) < Cu(II) < Gd(III). Par ailleurs, les résultats montrent que le complexe Gd(III)-L@

1H5 est moins stable d’au moins deux ordres de grandeur que le complexe Gd(III)-L@2H3. Ceci suggère

qu’une fois greffé sur la nanoparticule, le complexe de gadolinium correspondant gagne en stabilité. Par ailleurs, des études comparatives d’inertie chimique montre que le complexe Gd(III)-L@

1H5 greffé sur la nanoparticule a une inertie chimique comparable à celle de l’agent de contraste commercial Gd-DTPA. En revanche lorsque ce complexe est seul, sa vitesse de démétallation est rapide. Le greffage du ligand L@

1H5 à la surface de la nanoparticule est donc au bénéfice de la stabilité et de l’inertie chimique de son complexe de Gd(III). Ce gain de stabilité peut être attribué à l’’effet de ballast’ de la nanoparticule qui rigidifie la structure du complexe et limite sa démétallation. Study of transmetallation mechanisms of macrocyclic and linear gadolinium complexes at physiological pH for MRI

The aim of this work is to analyse the stability of metal complexes with ligands or nanoparticles of interest in MRI and to study their transmetallation mechanisms in the presence of endogenous cations near physiological pH. Two types of polyaminocarboxylate ligands were studied for which the binding unit was either linear or macrocyclic.

Macrocyclic ligands are constituted of a DO3A backbone functionalized with a benzimidazole (L1H4) or a p-nitrophenylbenzimidazole unit (L2H3). Thermodynamic data indicated that the affinities of these ligands towards first row transition metal ions (Cu(II) and Zn(II) or lanthanide ions (Gd(III) and Eu(III)) are increased compared to the corresponding ones with DO3A. This enhancement is correlated to the involvement of the benzimidazole moiety to each metal coordination sphere. For gadolinium complex Gd(III)-L1H4, its kinetic inertness was evaluated in phosphate buffer by relaxometry, in the presence of equimolar quantities of Zn(II) as a competitor. For Gd(III)-L1H4, no such reaction was detected which is in favour of kinetic inertness of Gd(III)-L1H4. Linear ligand, dithiolated DTPA bisamide L@

1H5 was designed with an aim of grafting it onto gold nanoparticles. L@

1H5 and the ligand grafted into gold nanoparticle, namely L@2H3, were analysed for their

thermodynamic stability towards mainly Cu(II), Zn(II) and Gd(III). Whatever the system, L@1H5 or L@

2H3, the general trend of increasing complex stability was Zn(II) < Cu(II) < Gd(III). Furthermore, Gd(III)-L@

1H5 complex was less stable than Gd(III)-L@

2H3, this latter being 2 orders of magnitude more stable at physiological pH. This suggested that the gadolinium complex stability is enhanced when the ligand is grafted onto the nanoparticle. Moreover, comparative kinetic inertness studies showed that the gadolinium complex Gd(III)-L@

1H5 is not chemically inert and demetallates rapidly while the gadolinium complex grafted onto the nanoparticle exhibit almost equal kinetic inertness as Gd-DTPA (Magnevist). The bulky nanoparticle probably rigidifies the structure of the complex and prevents Gd(III)-L@

2H3 from an extensive demetallation, which was a good point for the possible use of these nanoparticles in living organisms for imaging applications.

Mots – clés: Ligands DO3A, methyl benzimidazole, ligands DTPA bisamide, nanoparticules d’or, Stabilité thermodynamique, inertie chimique, transmétallation, relaxométrie, spectroscopie UV-Visible.

Keywords: DO3A ligands, methyl benzimidazole, DTPA bisamide ligands, gold nanoparticles, thermodynamic stability, kinetic inertness, transmetallation, relaxometry, UV-Visible spectroscopy.

Adresse du laboratoire et de l’unité: Université de Reims Champagne Ardenne, ICMR– UMR CNRS 7312, Groupe Chimie de Coordination, UFR Sciences Exactes et Naturelles – Moulin de la Housse - Batiment 18 – BP 1039, 51687, Reims Cedex 2

Etude à pH physiologique, des mécanismes de ... · (Hemanth) for his continuous encouragement...

Documents

Transcript of Etude à pH physiologique, des mécanismes de ... · (Hemanth) for his continuous encouragement...