Thermodynamic studies of halogen bonding in …...ii Thermodynamic studies of halogen bonding in...

Thermodynamic studies of halogen bonding in solution and application to anion recognition

by

Mohammed Golam Sarwar

A thesis submitted in conformity with the requirements For the degree of Doctor of Philosophy

Department of Chemistry University of Toronto

©Copyright by Mohammed Golam Sarwar 2012

ii

Thermodynamic studies of halogen bonding in solution and application to anion recognition

Mohammed Golam Sarwar Doctor of Philosophy

Department of Chemistry University of Toronto

2012 Abstract Halogen bonding (XB), the interaction between electron deficient halogen compounds and

electron donors, is an established non-covalent interaction in the solid and gaseous phases.

Understanding of XB in the solution phase is limited. This thesis describes experimental

studies of XB interactions in solution, and the application of XB interactions in anion

recognition.

Chapter 1 is a brief review of current understanding of XB interaction: theoretical

models, studies of XB in solid and gaseous phases and examples in biological systems are

discussed. At the end of this chapter, halogen bonding in the solution phase is discussed,

along with applications of halogen bonding in organic syntheses.

In chapter 2, linear free energy relationships involving the thermodynamics of

halogen bonding of substituted iodoaromatics are studied. The utility of substituent constants

and calculated molecular electrostatic potential values as metrics of halogen bond donor

ability are discussed. Density Functional Theory (DFT) calculations are shown to have useful

predictive values for trends in halogen bond strength for a range of donor-acceptor pairs.

Chapter 3 describes the development of new multidentate anion receptors based on

halogen bonding. Bidentate and tridentate receptors were found to exhibit significantly

higher binding constants than simple monodentate donors. These receptors show selectivity

for halide anions over oxyanions. Using 19F NMR spectra at different temperature, the

enthalpies and entropies of anion bindings for monodentate and tridentate receptors were

iii

determined. The results indicate a positive entropy contribution to anion binding for both

mono and tridentate receptors in acetone solvent.

Finally in chapter 4, some mesitylene based receptors with 3-halopyridinium and 2-

iodobenzimidazolium donors are introduced. The receptors perform halide anion recognition

in aqueous solvent system through charge-assisted XB interactions. These findings can allude

to utility in organic synthesis, supramolecular chemistry and drug design.

iv

ACKNOWLWDGEMENTS

First I owe many thank to my supervisor, Professor Mark S. Taylor for giving me

opportunity to work with him throughout my study in this department. His indispensable

guidance, constant encouragement, excellent supervision and thoughtful suggestions

immensely contributed to the successful completion of this work. I cannot imagine any better

environment for encouraging discussion and discovery.

My appreciation goes to group members in this group for their friendship and help. I would

like to thank Sunny Lai, Ali Rostami, Doris Lee, Elena Dimitrijevic, Corey McClay, Mike

Chudzinski, . Many thanks to Pragma Roy, Bojan Dragisic, and Simon Chan, undergraduate

students for their help in syntheses and titration experiments. I wish to thank the Department

of Chemistry and the University of Toronto for financial support.

At last but not the least, I am grateful to my grandfather, parents and siblings for teir

unconditional support for me over the years.

v

Table of Contents

ABSTRACT……………………………………………………………………………

ii

ACKNOWLEDGMENTS……………………………………………………………..

iv

LIST OF TABLES….…………………………………………………………………

viii

LIST OF SCHEMES……………………………………………………………………

x

LIST OF FIGURES ………………………………………………………………….

xii

LIST OF ABBREVIATIONS………………………………………………………….

xvii

1. CHAPTER 1……………………………………………………………………..

1

1. Halogen bonding: A noncovalent interaction………………………………………

1

1.1. Introduction………………………………………………………………………..

1

1.2. History of the halogen bonding interaction………………………………………..

1

1.3. Halogen bonding in the gas phase…………………………………………………

6

1.3.1. Complex formation with ammonia, water and hydrogen sulfide………………

7

1.3.2. Complex formation with non-aromatic π–electron donors………………………

9

1.4. Halogen bonding in the solid state…………………………………………………

11

1.4.1. Halogen bonding in crystal engineering………………………………………..

12

1.4.2. Liquid crystals…………………………………………………………………..

14

1.4.3. Metal-organic frameworks………………………………………………………

17

1.4.4. Nanomaterials…………………………………………………………………...

19

1.5. Halogen bonding in medicinal chemistry…………………………………………

21

1.6. Halogen bonding in the solution phase……………………………………………

24

1.6.1. Halogen bonding in supramolecular chemistry…………………………………

29

1.6.2. Syntheses induced by halogen bonding……………………………………….. 32

vi

1.7. Conclusions………………………………………………………………………... 34

1.8. Bibliography………………………………………………………………………

36

2. CHAPTER 2……………………………………………………………………….

42

2. Thermodynamics of halogen bonding in solution: substituent, structural, and solvent effects…………………………………………………………………………

42

2.1. Introduction……………………………………………………………………….

42

2.1.1. Early studies of XB in solution phase…………………………………………..

42

2.1.2. XB studies in solution using IR spectroscopy………………………………….

44

2.1.3. XB studies in solution using 19F NMR pectroscopy…………………………..

46

2.1.4. XB studies in solution using 13C NMR spectroscopy………………………….

47

2.2. Results and discussion……………………………………………………………

50

2.2.1. Experimental determinations of Ka values for halogen bonding interactions…..

50

2.2.2. Substituent effects on the halogen-bond donor ability of iodoperfluoro- benzenes

51

2.2.3. Structural Effects on halogen bonding: a combined experimental and computational approach………………………………………………………………

58

2.3 Conclusions………………………………………………………………………...

62

2.4 Experimental details………………………………………………………………

64

2.5 Bibliography………………………………………………………………………

71

3. CHAPTER 3………………………………………………………………………

76

3. Development of anion receptors using halogen bonding interactions………………

76

3.1. Introduction………………………………………………………………………..

76

3.2. Results and discussion……………………………………………………………

80

vii

3.2.1 Association constants and thermodynamic data for interactions of C8F17I and C6F5I with anions………………………………………………………………………

80

3.2.2. Development of multidentate halogen bond donors……………………………

84

3.2.3. Solvent effects on halogen bonding……………………………………………

96

3.2.4. Anion receptor based on hydrogen and halogen bonding interactions…………

100

3.2.5. Assembly of ternary complex through hydrogen and halogen bonding interactions…………………………………………………………………………….

101

3.3. Conclusions and outlook………………………………………………………….

102

3.4. Experimental details………………………………………………………………

105

3.5. Bibliography………………………………………………………………………

125

4. CHAPTER 4………………………………………………………………………..

130

4. Development of cationic halogen bonding receptors based on halopyridinium and haloimidazolium groups….……………………………………………………………

130

4.1. Introduction……………………………………………………………………….

130

4.2. Result and discussion…………………………………………………………….

132

4.3. Conclusions………………………………………………………………………

136

4.4. Experimental……………………………………………………………………..

137

4.5. Bibliography………………………………………………………………………

145

5. CHAPTER 5………………………………………………………………………..

130

5. Lesson, hurdles and future directions.……………………………………………….

146

viii

List of Tables

Table 1.01. Inhibition of hCatL influence by halogen bonding (XB)

22

Table 1.02. Binding constants (M-1) of complexes between pyridine derivatives and THF with halogen bonds donors (CCl4, CBr4, and I2).

25

Table 2.01. Equilibrium constant (Ka) for the complex formation of inorganic, perfluorinated haloalkenes (e.g., C6F13I) and the non fluorinated R−C≡C−I halogen bond donors with base pyridine in CCl4

42

Table 2.02. Equilibrium constants (Ka) for the complex formation of inorganic halogen bond donors with pyridine in CCl4 at 30 °C.

43

Table 2.03. Equilibrium constant (Kc) for halogen bonding of XB with Et2NCOMe in CCl4.

45

Table 2.04. Equilibrium constant (Kapp) for the complex of I−C≡CCN with Et2NCOMe in C6H6-CCl4 solvent mixtures.

46

Table 2.05. Halogen bonding association constants of para-substituted iodotetrafluorobenzenes (tri-n-butylphosphine oxide acceptor, cyclohexane solvent), with substituent constants (σ para and σ meta) and computed electrostatic potential at the iodine atom.

53

Table 2.06. Experimental and predicted binding energies of C6F5I and C8F17I with representative halogen bond acceptors.

59

Table 2.07. Computed Halogen Bonding Energies for Comparison to Experimental Data From Table 2.06.

61

Table 3.01. Binding constant of C6F5I and C8F17I with n-Bu4N+X− in acetone.

84

Table 3.02. Calculated geometries of 3.11·X− and 3.12·X−.

86

Table 3.03. Binding constant of 3.11 and 3.12 with n-Bu4N+X− in acetone.

89

Table 3.04. Binding constant of C6F5I, C6F4ICOOH, C6F4ICOOBn, and C6F4ICO(NH)Bn with n-Bu4N+X− in acetone

90

Table 3.05. Binding constant of receptors 3.11, 3.12, 3.13, and 3.14 with n-Bu4N+X− in acetone

92

Table 3.06. Binding constant of 3.21, 3.22 and 3.23 with n-Bu4N+X− in acetone

95

Table 3.07. Binding constant of C6F5I, C8F17I and 3.21 with n-Bu4N+Cl− in different solvent system.

97

ix

Table 3.08. Binding constant of C6F5I with n-Bu4N+Cl− in acetone with the presence of trace amount of water.

98

Table 3.09. Association constant 3.13, 3.14, 3.26, 3.27, and Cl− in acetone at 295 K.

103

Table 3.10. Calculated halogen-bonding complex (3.21–X– ) mass for comparison to actual mass (ESI)

123

Table 4.01. Binding constants (M-1) of receptor 4.07 (X = Cl), 4.08 (X = Br), and 4.09 (X = H) with n-Bu4N+X− in CD3CN:D2O (9:1) and DMSO.

136

Table 4.02. Binding constants of 4.13 and 4.15 with n-Bu4N+X− in CD3CN:D2O (8:2).

137

Table 4.03. Binding constants of 4.14 and 4.16 with n-Bu4N+X− in CD3CN:D2O (8:2).

138

x

List of Schemes

Scheme 1.01. Mechanisms for the reaction of halogen molecules with alkenes (outer and inner complexes) in polar solvents in the dark.

2

Scheme 1.02. Halogen bonding and hydrogen bonding interactions: A) halogen bond donor (D−X) with acceptor A, B) hydrogen bond donor (D−H) with acceptor A, and C) hydrogen bond donor (D−H) with acceptor (D−X)

3

Scheme1.03. The formation of assembly consisting of functionalized AuNP with organic cross-linkers (BPEB) in solution.

20

Scheme 1.04. Stepwise generation of assemblies consisting of functionalized gold nanoparticles (AuNP-1) and organic cross-linkers (BPEB) on silicon organic monolayers (M1, M2).

20

Scheme 1.05. Halogen bond catalyzed reduction of 2-phenylquinoline (1.43).

33

Scheme 1.06. Reaction of benzhydryl bromide (1.45) with wet CD3CN catalyzed by 1.48 or 1.49.

33

Scheme 1.07. Baylis-Hillman reaction between aromatic aldehydes and Michael acceptors.

34

Scheme 1.08. Synthesis of N-arylaziridine (1.53) using halogen bonding interaction.

34

Scheme 3.01. formation of polyiodide( m and n integers > 0, n = 1 to 4)

78

Scheme 3.02. Complex formation of tetraiodide with 1,6-bis(trimethylammonium)-hexadiiodide (3.06).

79

Scheme 3.03. Synthesis of 4,4’-bis(2,3,4,5-tetrafluoro-6-iodophenyl)buta-1,3-diyne (3.11) and 3.12

87

Scheme 3.04. Synthesis of 4,4’-bis(2,3,4,5-tetrafluoro-6-iodophenyl)buta-1,3-diyne (3.11) and 3.12

89

Scheme 3.05. Synthesis of 2,3,4,5-tetrafluoro-6-iodobenzaldehyde (3.16)

91

Scheme 3.06. Synthesis of hexadentate receptor 3.25.

96

Scheme 3.07. Synthesis of calix[4]arene based intramolecular hydrogen and halogen bonding receptor (3.26).

100

xi

Scheme 3.08. Synthesis of ester 3.27 from 6-methylpyridin-2-amine.

101

Scheme 4.01. Syntheses of receptors 4.07, 4.08 and 4.09 when R = Cl, Br, and H respectively.

133

Scheme 4.02. Syntheses of benzimidazolium derived receptors 4.13 and 4.14 when R = CH3 and CH2Ph respectively

134

Scheme 4.03. Syntheses of benzimidazolium derived receptors 4.15 and 4.16 when R = CH3 and CH2Ph respectively

135

xii

List of Figures Figure 1.01. Halogen bonding in Br2 −1,4-dioxane, Br2 –hexamethylenetetramine and Br2 –acetone complexes

2

Figure 1.02. The computed B3LYP/6-31+G(d,p) electrostatic potential, on the 0.001 electrons/bohr surface, of C6F5I .

5

Figure 1.03. Molecular electrostatic potential surfaces (B3LYP/6-31+G**- LANLdp, Gaussian 03) of trifluoro-halomethane species: CF4 (left), CF3Cl (middle left), CF3Br (middle right), and CF3I (left).

6

Figure 1.04. Geometry of both XB and HB complexes a) NH3 contains lone pairs of electrons and the observed C3v geometries in b) H3N···Br2, c) H3N···BrCl, and d) H3N···HBr complexes.

7

Figure 1.05. Nonbonding pair electron density model of H2O and the observed geometries of a) H2O, b) H2O···HF, and c) H2O···ClF. The equilibrium geometry for H2O···HF is shown, for which the angle φ = 46(5)° and For H2O···ClF, the estimated value is φ ≈ 20°, (effective value for the zero-point state).

8

Figure 1.06. Geometries of a) Nonbonding pair electron density model of H2S and the observed geometries of b) H2S···HCl (φ = 93.8°), and d) H2S···ClF (φ = 95.8°).

8

Figure 1.07. The geometry of H2O···F2 , which is effectively planar in the zero-point state and geometry of H2S···F2 . The angle φ = 113°, and the internal rotation of H2S about its local C2 axis (angle y) gave a low-energy vibrational satellite in the rotational spectrum.

9

Figure 1.08. a) π-Bonding electron density model of ethene and ethyne observed geometries (perpendicular or planar, C2v) of a) ethene···HCl b) ethene···ClF c) ethyne···HCl and d) ethyne···ClF.

9

Figure 1.09. a) pseudo-π-electron density model of cyclopropane and observed geometries b) cyclopropane···ClF c) cyclopropane···HCl.

10

Figure 1.10 Observed geometries of a) benzene···ClF and b) benzene···HCl.

10

Figure 1.11: 1D supramolecular network of cocrystal of two layers top: 1,6-diiodoper-fluorohexane with trans-4,4’-azobis (pyridine) and bottom:1,8-diiodoperfluorooctane with trans-4,4’-azobis (pyridine).

12

Figure 1.12. 1D chain produced through combination of hydrogen and halogen bond.

12

Figure 1.13. Matching and mismatching various I–(CF2)m–I with bis-

xiii

(trimethylammonium) alkane to separate α,ω-diiodoperfluoroalkane from their mixtures.

13

Figure 1.14. a) 2D sheet generated through a combination of self-complementary hydrogen bonds and I···I interactions.

14

Figure 1.15. The molecular structure of the XB complexes between alkoxystilbazole and 4-substituted iodotetrafluorobenzene.

15

Figure 1.16. The molecular structure of the XB complexes between 4-substituted pyridine and 4-substituted iodotetrafluorobenzene.

16

Figure 1.17. The bent-core halogen bonding complexs between stilbezole and 1,3-disubstituted perfluorobenzene.

17

Figure 1.18: Halogen bonded polymer.

17

Figure 1.19. Geometry of type I and II halogen-halogen interactions. Type I: θ1 = θ2; and type II: θ 1 = 90° involving the halogen bond acceptor role, θ 2 = 180° involving the halogen donor role.

18

Figure 1.20. Structure of the (2-haloimidazolium)2(TeX2)salt (1.14) and (DIETSe)2FeCl4 (1.15).

19

Figure 1.21. Binding mode of covalent inhibitors 1.16 and 1.17 at the active site of hCat L with its three pockets. Pocket S1 binds with covalent bond, ArSO2- group accommodates in pocket pocket S2 contains and the substituent at position 4 of the phenyl ring approaches the C=O group of Gly61 in the S3 pocket.

22

Figure 1.22. X-ray structure bound to Inhibitors KH-CB19 (1.18) and CLK3 reveals the halogen interaction with the hinge CO.

23

Figure 1.23. The chemical structures of adenosine triphosphate and 5,6-dichlorobenzimidazone-1-β-D-ribofuranoside (DRB, 1.19) bound to a Cdk9 hinge region

24

Figure 1.24. L-thyroxine, 3,5,3',5'-tetraiodothyronine (1.20) is the precursor of thyroid hormone.

24

Figure 1.25. . Equilibrium geometries of CF3X···N(Me)3 complexes (X = Cl, Br,

or I).

26

Figure 1.26. Equilibrium geometries of CF3X···O(Me)2 complexes where ((X = Cl, Br, or I) and (CF3Br)2···O(Me)2

27

Figure 1.27. The computed B3LYP/LANLdp electrostatic potential, on the 0.001 electrons/bohr surface of I2.

28

xiv

Figure 1.28. Hydrogen bond and halogen assisted interaction between heterocyclic amines (2-aminopyridine 1.21, N,N-dimethylaminopyridine 1.22, N,-phenylamino pyridine 1.23, 1H-pyrrolo[2,3-b]pyridine 1.24) and diiodine

28

Figure 1.29. The symmetric (a) and asymmetric (b) geometries of the systems bis(pyridine) bromonium and iodonium triflates (1.25, X = Br, for 1.26, X = I and for 1.27, X = H).

29

Figure 1.30. Perylene diimide based C-shaped receptors (1.28), S-shaped receptors (1.29) and phenazine (1.30).

30

Figure 1.31. Interactions between the nitropyrrole containing bisaniline derivatives (1.31, 1.32, 1.33 and 1.34) with isothaloyl derivative (1.35). (φ = 30 ≈ 37°)

31

Figure 1.32. Nickel fluoride complex trans-[NiF(C5F4N)(PEt3)2]·C6F5I (1.36), trans-[NiF(C5F4N) (PEt3)2]·2C6F5I (1.37) and nickel chloride complexes trans-[NiCl(C4F2N)(PEt3)2] (1.38).

32

Figure 1.33. Ni, Pd, and Pt fluoride complex

32

Figure 2.01. Halogen bonding interactions between iodoalkynes and DMSO.

48

Figure 2.02. Chemical shift of C-1 of n-C4H9−C≡C−I (2.03) versus donor number (DN) of the solvent.

48

Figure 2.03. Chemical shift of C-1 of n-C4H9−C≡C−I (2.03) with respect to β.

49

Figure 2.04. Chemical shift of C-1 of n-C4H9−C≡C−I (2.03) verses solvent soft basicity Bsoft.

49

Figure 2.05. The computed B3LYP/6-31+G(d,p) electrostatic potential, on the 0.001 electrons/ bohr surface, of C6F5I. The iodine atom is at the right.

51

Figure 2.06. Structures of the substituted iodoperfluoroarenes employed in the substituent effect study.

52

Figure 2.07. Molecular electrostatic potential surfaces (B3LYP/6-31+G**-LANLdp; Gaussian ’03) of (left to right) iodopentafluorobenzene, 2.03a and 2.03b.

54

Figure 2.08. Correlation of 4-X-C6F4I···Bu3PO halogen bond strength log (Ka) with (a) σ para and (b) σmeta substituent constants of X.

55

Figure 2.09. Correlation of 4-X-C6F4I···Bu3PO halogen bond strength log(Ka) with the electrostatic potential density at the iodine atom, calculated with (a) AM1 and

56

xv

(b) DFT (B3LYP/6-31+G**/LANLdp) computational methods. Figure 2.10. Molecular electrostatic potential surfaces (B3LYP/6-31+G**-LANLdp; Gaussian ’03) of (iodoethynyl)benzene (left); 5-iodouracil.

57

Figure 2.11. Experimental binding data (–∆Gbinding) against values according to the model of Hunter and coworkers.

60

Figure 2.12. Experimental binding data (–∆Gbinding, Table 2.05) against DFT-calculated (–∆ECPB3LYP, Table 2.06) gas phase binding energy.

62

Figure 3.01. Bidentate coordination of nitrates by 2-bromo-1,3-diisopropyl-4,5-dimethyl imidazolium nitrate (3.01), infinite chains formed by: 3,5,3´-tri-iodo-L-thyronine (3.02), chain formation on sulfonate coordination in 4-chloropyridine-3-sulfonic acid (3.03)

77

Figure 3.02. Bidentate coordination of Br− with 2,5 dibromoaniline (3.04), Cl− with 2,5 dichloroaniline (3.05).

78

Figure 3.03. Heteroditopic receptor (3.09) and (3.09a).

79

Figure 3.04. Top: 19F NMR titration of C8F17I (acetone, 295 K; change inchemical shift |∆δ| (in ppm) vs [n-Bu4N+Cl−] (in mm)).

81

Figure 3.05. Van’t Hoff plot (ln(Ka) vs 1000/T) for the interaction of C8F17I with Cl− anion in acetone solvent.

82


83

Figure 3.07. Hydrogen bonded molecular aggregation among 2-pyridones and dipyridones (3.10a, 3.10b, 3.10c).

85

Figure 3.08. Complex formation by propose receptors with halide anions. Complexes are 3.11·X− and 3.12·X−.

86

Figure 3.09. Solid-state structures of the halogen-bonded complex of a) 3.11···I− and n-Bu4N+ b) 2 × (3.11···I−).

88

Figure 3.10. Bidentate receptors.

91

Figure 3.11. Tridentate receptors of mesityle (3.21), amine (3.22) and benzene scaffold (3.23); spatial arrangement of ethyl groups in hexaethylbenzene (3.24).

93

Figure 3.12. Left: Calculated structure of the 3.21–Cl− complex (HF/6-31+G**-LANL2DZdp: H white; C gray; O red; F blue; I purple; Cl green). Right: Electrostatic potential surface of 3.21 (chloride-bound conformation, B3LYP/6-

94

xvi

31G*-LANL2DZdp: blue indicates sites of partial positive charge and red sites of partial negative charge). Figure 3.13. Job plot ([3.21+n-Bu4N+Cl−] = 11 mm, |χ∆δ| vs mole fraction χ).

94

Figure 3.14. Cu+2 coordinated hexadentate receptor 3.25.

95

Figure 3.15. Van’t Hoff plot (ln(Ka) vs 1000/T) for the interaction of receptor 3.21 with Cl− anion in acetone solvent.

99

Figure 3.16. Propose ternary complex employed by both hydrogen and halogen bonding. The donors are the molecules 3.26 and 3.13.

100

Figure 3.17. Tridentate receptors (3.21a, 3.21b, and 3.21c) and bidentate receptors syn bromoimidazoliophane (3.29.2PF6−).

103


119


119

Figure 3.20. Van’t Hoff plot (ln(Ka) vs 1000/T) for the interaction of receptor 3.21 with Cl− anion in acetone solvent.

120

Figure 3.21. Calibration curve of probe temperature (°C) with methanol solvent.

Figure 4.01. Macrocyclic receptor 4.01·2PF6− having the interlocked nature of catenane.

130

Figure 4.02. Imidazolium based receptors 4.03 and 4.04 using anthracene as spacer or scaffold.

131

Figure 4.03. Imidazolium based receptors 4.05 and 4.06 using 1,3,5-trimethylbenzene-based scaffold.

132

xvii

List of Abbreviations

CDCl3 deuterated chloroform

B3LYP Becke, three-parameter, Lee-Yang-Parr

CH2Cl2 methylene chloride

CH3CN acetonitrile

DCE 1,2-dichloroethane

DCM dichloromethane

DFT density functional theory

DIC N,N’-diisopropylcarbodiimide

DMAP 4-N,N-(dimethylamino)pyridine

DMF N,N-dimethylformamide

DMSO dimethyl sulfoxide

EI electron impact

equiv. equivalents

ESI electrospray ionization

Et ethyl

Et3N triethylamine

EtOAc ethyl acetate

FID free induction delay

g gram

GC gas chromatography

HRMS high-resolution mass spectrometry

Hz Hertz

IR infrared

J coupling constant

LDA lithium diisopropylamide

LANL2DZ Los Alamos National Laboratory 2-double-z (density functional theory) M molarity

Me methyl

MgSO4 magnesium sulfate

MHz megahertz

xviii

MP2 Møller-Plesset Perturbation theory

δ chemical shift in parts per million

m/z mass-to-charge ratio

NaHCO3 sodium bicarbonate

Na2SO4 sodium sulfate

NaHSO4 sodium bisulfate

NaS2O3 sodium thiosulfate

NBS N-bromosuccinimide

NH4Cl ammonium chloride

NMR nuclear magnetic resonance

mmol millimole

mL milliter

µL microliter

MS mass spectrometry

n-Bu n-butyl

OAc acetate

Piv pivaloyl

ppm parts per million

py pyridine

Rf perfluoroalkane/ perfluoroarene

TBAF tetrabutylammonium fluoride

TBACl tetrabutylammonium chloride

TBABr tetrabutylammonium bromide

TBAI tetrabutylammonium iodide

t-Bu tert-butyl

THF tetrahydrofuran

TMS tetramethylsilane

TMSOTf trimethylsilyl trifluoromethanesulfonate

TOF time of flight

UV ultraviolet

Vis visible

1

Chapter 1

Halogen Bonding: A Noncovalent interaction

1.1. Introduction Halogenated compounds play key roles in polymer chemistry, crystal engineering, medicinal

chemistry, and environmental science. The interaction between electron deficient halogen

compounds and electron donors is known as halogen bonding. Halogen bonding has seen

applications in areas such as crystal design, molecular recognition, topochemical reactions,

molecular conductors, and liquid crystals; its role in structural biology has also been

examined. The aims of this review is not to summarize all recent work related to halogen

bonding, but rather to provide an overview of studies of the interaction in different phases.

1.2. History of the halogen bonding interaction Although halogen bonding interactions were first observed by Guthrie in 1863,1 the term

“halogen bond” was not coined until late 1970’s by Dumas.2 In Guthrie’s experiment, iodine

powder was added to an ammonia solution to obtain an iodoammonium complex. Half a

century later, the group of Odd Hassel used X-ray crystallography to probe the structure of

the bromine−dioxane complex3 and the bromine−hexamethylene tetramine complex.4 The

Br−Br bond length had elongated upon complex formation, changing from 2.28 Å in the free

molecule (Br2) to 2.31 Å whereas Br···O distance in the complex was 2.71 Å, shorter than the

sum of van der Waals (VDW) radii (3.35 Å). Similarly, in the Br2-hexamethylenetetramine

complex, the Br−Br bond length was 2.3 Å and the Br···N distance 2.4 Å. The noncovalent

O···Br−Br and N···Br−Br angles were found to be linear in both crystals. Hassel showed that

“halogen molecule bridges” of the dioxane−halogen addition compounds were not only

confined to halogen compounds formed with ethers and amines but had also been found in

the 1:1 complex formed by acetone and bromine; the 2:1 ratio in bromine−

hexamethylenetetramine complex as well. In the case of bromine−acetone complexes, the 1:1

complex was formed by alternating acetone and bromine end chains, indicating that both

electron pairs were capable of forming such bonds simultaneously (Figure 1.01).

2

O

O

O

OBr Br

N N

NN

BrBrO

H3C CH3

BrBrBrBr

Br Br Figure 1.01. Halogen bonding in Br2 −1,4-dioxane, Br2 –hexamethylenetetramine and Br2 –acetone complexes.

According to Hassel, charge transfer from electron donors (Lewis bases) to electron

acceptors (iodine or bromine) was responsible for these interactions.5 In 1949 Benesi and

Hildebrand investigated solutions of iodine in benzene and mesitylene, and reported the

presence of absorption bands in the ultraviolet region. Their quantitative studies were able to

determine the equilibrium constants 1.72 M-1 and 1.15 M-1 for the iodine−benzene complex

in CCl4 and n-heptane solvent respectively, whereas equilibrium constants of the iodine-

mesitylene complex in CCl4 and n-heptane solvent were 7.2 M-1 and 5.3 M-1, respectively.6

At the same time, Mulliken worked out a quantum mechanical theory of complex resonance

that explained halogen bonding as a charge-transfer interaction. The direction of a bond

linking between a lone pair donor atom to a halogen atom was determined by the orientation

of the orbitals of non-bonding electron pairs in the donor atom.7 He also proposed a

mechanism for halogenation of alkenes involving an outer complex where there was little

charge transfer between components and an inner complex where there was extensive charge

distribution (Scheme 1.01).8

Scheme 1.01. Mechanism for the reaction of halogen molecules with alkenes (outer and inner complexes) in polar solvents in the dark.

C C XY C C

XY

C C

X

Youter complex inner complex

3

According to their Nobel lectures in 1966 and 1970, both Mulliken and Hassel respectively

described interactions involving negative charge transfer from electron donors to electron

acceptor halogens along with the properties of these systems elicited by X-ray

crystallography.9 In his work, Dumas defined the “halogen bond” (XB), analogous with the

hydrogen bond, by determining the association constants of pyridine derivatives with CX4 (X

= Cl, and Br) and comparing them with those between pyridine derivatives and pyrrole. The

attractive interaction between a halogen atom X from R–X (e,g. X = I, Cl, and Br), in which

R is an electron withdrawing group and an atom or a group of atoms A, and the acceptor

species is termed halogen bonding. Similarities between halogen bonding and hydrogen

bonding interaction are depicted in Scheme 1.02, where A is a halogen bonding while B and

C are hydrogen bonding interactions. The details of these interactions will be discussed later.

D X A D H A D X

DH

··· ··· ···

A B C Scheme 1.02. Halogen bonding and hydrogen bonding interactions: A) halogen bond donor (D−X) with acceptor A, B) hydrogen bond donor (D−H) with acceptor A, and C) hydrogen bond donor (D−H) with acceptor (D−X).

In 1979, Bertran and Rodriguez first reported the interaction of halogen bond donors (e.g.,

chloroform, bromoform and iodoform) with different amines, ketones, ethers and esters by

measuring the chemical shift (δ) change of the haloformic proton in different solvents

relative to that in hexane.10 Analyzing the environment around halogen centers in hundreds

of crystal structures, Parthasarathy et al. in 1986 studied the angular preferences of

intermolecular forces around halogen centers; nucleophiles approach C−X along the backside

of the bond whereas electrophiles approach nearly perpendicular to C−X.11 From the late

1990’s there was a rapid increase in the number of publications involving halogen bonding,

including applications in the solid state, studies in the gas phase, and computational work.

The Milan group successfully used halogen bonding in crystal engineering, and to solve

problems in supramolecular chemistry.12 The Legon group contributed to the research

advancement by determining the symmetries, radial and angular geometries, as well as

4

intermolecular force constants Kσ of weakly bound halogen bonded complexes in the gas

phase.13

Trends in halogen bond donor ability have been investigated through both theoretical and

experimental approaches, and the order I > Br > Cl >> F has emerged from many different

studies in solid, liquid and gas phases, correlated positively with the polarizabilities of the

halogen atoms and correlated negatively with their electronegativities. The donor abilities of

these halogen atoms are also influenced by the hybridization of adjacent carbon atoms as the

order of donor strength increases in the order X−X > X−C (sp) > X−C (sp2) > X−C (sp3).14

Paradoxically, halogen atoms having partial negative charge in R−X compounds can act as

electrophiles and accept electron density from Lewis bases.

The group of Politzer has advanced an explanation for this phenomenon that is distinct from

the charge-transfer hypothesis discussed above: it relies on a concept named ‘sigma-hole’.

The σ-hole represents a region of positive electrostatic potential that develops along the

covalent carbon−halogen bond as a result of anisotropic distribution of electron density

around the halogen nucleus. As a result, the atomic radius of halogen atom along the covalent

bond axis is smaller than that along the perpendicular direction.15

5

δ− δ+

Figure 1.02. The computed B3LYP/6-31+G(d,p) electrostatic potential, on the 0.001 electrons/bohr surface, of C6F5I . The iodine atom is at the right. Red indicates negative charge density and blue positive charge density The molecular electrostatic potential surface of any ground state may be defined as the three-

dimensional contour of the electronic density envelope.16,17 The electrostatic potential of any

ground-state spherically symmetric atom is positive everywhere. On the contrary, when

atoms are combined to form a molecule, the accompanying rearrangements of electronic

charge normally produce one or more regions of negative potential. Figure 1.02 depicts the

molecular electrostatic potential for C6F5I where a positive potential is developed on the

outer most portion of the iodine surface along the C−I axis. This portion of electrostatic

potential is represented by blue color and is called the σ-hole. Previously it was also observed

by Allen group that the effective atomic radius of atom I along the extended C−I bond axis is

smaller than the radius perpendicular to this axis, a phenomenon known as “polar

flattening”.18 The σ-hole of a halogen bond donor can create the possibility of an attractive

electrostatic interaction with a nucleophile to form a halogen bond which is along the C−I

axis. The size of the σ-hole depends on the hybridization state and electronegativity of the

halogen atom X (e.g. X = I, Cl, and Br) along with the electron withdrawing ability of the

group attached with it.

6

Figure 1.03. Molecular electrostatic potential surfaces (B3LYP/6-31+G**- LANLdp, Gaussian 03) of trifluoro-halomethane species: CF4 (left), CF3Cl (middle left), CF3Br (middle right), and CF3I (left). Red indicates negative charge density and blue positive charge density: the plots have been set to the same color scale so that a visual comparison can be made. One view of bonding in CF3X places two unshared electron pairs for X in two p-orbitals and

one unshared electron pairs in one s-orbital. The amount of s orbital contribution to the

hybridization of the halogen atom influences the size of the σ-hole; the s-contribution to the

bonding orbital of X decreases according to the following sequence F, Cl, Br, and I (25%,

12%, 9% and 9%, respectively). Using natural bond order (NBO) population analysis,

Politzer et al. showed that the share of bonding electrons by X in CF3X was 71%, 53%, 49%

and 46% for F, Cl, Br, and I, respectively.15 The combined effects of hybridization and

electronegativity make the fluorine atom the poorest halogen bond donor and make the

iodine atom the best halogen bond donor. Moreover, incorporation of strongly electron

withdrawing functionalities on the halogen donor increases the size of the σ-hole and thus

increases the halogen bond donor ability. The electrostatic potential maps of CF4 (left),

CF3Cl, CF3Br, CF3I and CF4 (right) illustrate trends in the magnitude of the σ-hole as a result

of hybridization, electron withdrawing effects, and electronegativity (Figure 1.03).

1.3. Halogen bonding in the gas phase The nature of the halogen bond in the gas phase was studied by Legon et al. using microwave

rotational spectra of complexes B···XY, where B is a simple Lewis base and XY a dihalogen

molecule. The radial and angular geometries as well as the strengths of the intermolecular

binding of complexes B···XY were found closely parallel to those of the corresponding

hydrogen-bonded systems B···HX. To carry out these measurements, a supersonic jet or

beam of gas mixture was formed by premixing the components and then immediately

diluting them with argon followed by expanding from a relatively high pressure vessel

7

through a small circular nozzle into a vacuum. The weakly bound complexes emerged as

collisionless complexes that were frozen in their lowest rotational and vibrational energy

states and were then interrogated by microwave radiation in a high electric field to record

their rotational spectra. By analyzing the rotational spectra, different spectroscopic constants

(e.g. moment of inertia Ia, intermolecular stretching force constant k and bond angles θ, φ)

that contain much information about the complexes, were determined.13 The Legon group

found that B···XY complexes were best described as Mulliken outer type complexes with

very small charge redistribution on complex formation.8 On the other hand, almost negligible

bond strength for B···F2 complex indicated the effective absence of halogen bonding to

fluorine atom.17

1.3.1. Complex formation with ammonia, water and hydrogen

sulfide

The angular geometries of halogen bonded or hydrogen bonded complexes of the type

B···XY and B···HY were determined based on the molecular axis of Base B. The molecular

NH

HH N

H

HH Br Br N

H

HH Br Cl N

H

HH H Br

Figure 1.04. Geometries of both XB and HB complexes ( H3N···Br2, H3N···BrCl, and H3N···HBr). axes of bases lie along the axis of a nonbonding (n) electron pair of acceptor B, along the

local symmetry axis of a π (pseudo-π orbital) or along the axis of one of the n pairs when B

carries both n- and π-pairs. The complexes H3N···Br2 and H3N···BrCl both have C3v geometry

as shown in Figure 1.04 and are consistent with the above rules, where the electrophilic

region of the halogen interacted with the n-pair on N so that XY lay along the C3 axis.19

8

H FOHH

Cl FOHH

OHH

a b c

ϕϕ

Figure 1.05. Nonbonding pair electron density model of H2O and the observed geometries of a) H2O, b) H2O···HF, and c) H2O···ClF. The equilibrium geometry for H2O···HF is shown, for which the angle φ = 46(5)° and For H2O···ClF, the estimated value is φ ≈ 20°, (effective value for the zero-point state). The complexes of H2O with HF and ClF were expected to be pyramidal, but experimental

evidence was in favor of effectively planar structures, indicating that the energy barrier

between pyramidal and the planar conformation is low enough that the vibrational wave

functions can be classified as the C2v point group; even in the zero-point state, the molecule

tunnels rapidly between the pyramidal conformers. In Figure 1.05, the angle φ for both HB

and XB complexes is shown.

S

H H

S

H H

HCl S

H H

ClF

a b c Figure 1.06. Geometries of a) Nonbonding pair electron density model of H2S and the observed geometries of b) H2S···HCl (φ = 93.8°), and d) H2S···ClF (φ = 95.8°).

To determine the molecular angular geometry of halogen bonding complexes different

isotopomers, H232S···35ClF, H232S···37ClF, H234S···35ClF, HD32S···35ClF, D232S···35ClF,

DH32S···37ClF, and D232S···37ClF, of a complex formed by H2S with ClF was observed by

pulsed-nozzle, Fourier-transform microwave spectroscopy. From these rotational spectra,

vibrational satellites and Cl-nuclear quadrupole hyperfine components were obtained. By

using Hamiltonian H operator, these ground state rotational spectra were analyzed and

centrifugal distortion constants Djk, along with centrifugal distortion term χk, χD were

obtained. From the value of Djk, χk, χD.as well as moment of inertia Ia, Ib, and Ic, the angle φ

and ө were extracted.20 One of the rules for determining the equilibrium angular geometry of

halogen bonded complex is the axis of the complex lies along the axis of non bonding

electron paired carried by halogen bond acceptor. In terms of halogen bonding complex

formation the donor-acceptor interactions are between halogen bond donor and non bonding

9

(n) lone pair of electron such as H2O···ClF. In case of H2S···ClF complex, the two non

bonding electron pairs occupy two sp hybrids orbitals. The axes of these two orbitals are

180° to each other and 90° to the H2S molecular plane. As a result, the axis of complex

formed with halogen bond donor acceptor interaction is different from that of with H2S···ClF

complex in Figure 1.06. Paradoxically, the F2 molecule showed halogen bond donor ability

in the gas phase. The rotational spectrum of H2O···F2 complex indicated weak binding at zero

point energy level with a planar geometry, whereas a pyramidal geometry was observed for

the H2S···F2 complex with φ = 113° along with a vibrational satellite, suggesting that a low

potential energy barrier separates two equivalent pyramidal conformations (Figure 1.07).13, 21 φ

OH

HF F S

H H

F F

Figure 1.07. The geometries of H2O···F2, which is effectively planar in the zero-point state, and of H2S···F2. The angle φ = 113°, and the internal rotation of H2S about its local C2 axis (angle Y) gave a low-energy vibrational satellite in the rotational spectrum.

1.3.2. Complex formation with non-aromatic π–electron donors

Complexes between hydrogen bond donors and π-pair Lewis bases B showed that the HX

stayed along the local symmetry axis of a π-pair orbital in the absence of a n-pair orbital.

Both ethylene and ethyne bind with HX or XY perpendicular to their C2 axis. Angular

geometries of the complexes of ethylene and ethyne with XY (Cl2, BrCl, ClF) in Figure 1.08

showed a T shape where Xδ+ of XY interacts with the π-electrons as shown. Two π-orbitals in

ethyne are perpendicular to each other, giving it a cylindrical shape; the interaction with HX

or XY (HCl, HF, HBr, Cl2, Br2, ICl) resulted in a planar and T-shaped geometry.

Figure 1.08. Observed geometries (perpendicular or planar, C2v) of ethene···HCl, ethene···ClF, ethyne···HCl, and ethyne···ClF.

10

Due to similar behaviors of cyclopropanes and alkenes, Coulson and Moffitt introduced the

idea of a pseudo-π carbon−carbon bond generated by overlapping a pair of sp3 hybrid orbitals

on adjacent carbons.22 The symmetry axis of the pseudo-π-orbital coincides with a median of

the cyclopropane equilateral triangle. Hence the angular geometry of cyclopropane···ClF and

cyclopropane···HCl lie along the extension of the median.

H

HH

H

H

H

Cl

F

H

HH

H

H

H

H

Cl

H

H

H

H

H

H

a b c

Figure 1.09. a) Pseudo-π-electron density model of cyclopropane and observed geometries of b) cyclopropane···ClF c) cyclopropane···HCl. In case of the interaction between benzene with ClF or HCl, the aromatic π-orbitals act as the

halogen bond or hydrogen bond acceptor. The geometry of the benzene···ClF complex is

shown in Figure 1.10. The Cl end of the ClF molecule approximately circulated the cyclic

path in the potential energy minimum and sampled the π-electron density on one face of the

aromatic molecule. This motion is therefore governed by a “Mexican-hat” type of potential

function. Due to the lighter weight of the hydrogen atom and the absence of vibrational

satellites associated with the intermolecular modes, the vibrational wave function for

benzene···HCl complex was either strictly C6v equilibrium symmetry or would have preferred

a specific carbon−carbon bond. However, in either case, the vibrational wavefunctions will

have C6v symmetry.23

H

Cl

a b

φ = 23°

H H

HH

H H

φ = 14.4°Cl

F

H H

HH

H H

Figure 1.10. Observed geometries of a) benzene···ClF and b) benzene···HCl.

11

In contrast with general perception, fluorine that is bound to a strong electron withdrawing

group such as N or O may act as a halogen bond donor. By using DFT calculations at the

B3PW91/6-31G(d,p) level of theory, Metrangolo et al. proposed the presence of a positive σ-

hole on fluorine atom in CF4, CF3COF, FOCF2OF, HC=CF, O=CF2, FOC(O)F, F3CC(O)OF,

F3CSO2)2NF, and FSO2CF3.24 This proposal was supported by gas phase studies on the

H2S···F2 complex, H2O···F2 complex as well as the measurement of gas phase F2 elimination

from F3− (98.4±10.6 KJmol-1).25 The experimental XB bond strength of F3− is comparable

with that of the other trihalide anions (99 ± 5, 127 ± 7, and 126 ± 6 kJmol-1 for Cl3−, Br3−, and

I3−, respectively).

1.4. Halogen bonding in the solid state Due to advances in electronics and optics, single crystals are becoming increasingly

important tools in industrial areas and play crucial roles in semiconductors, optical crystals,

laser crystals, jewelry and watch industries. Crystal engineering is relevant to several

important applications, including gas storage in metal-organic frameworks and the

understanding of polymorphism in pharmaceuticals and explosives. Although hydrogen

bonds and coordination bonds remain the most preferable options for crystal engineering

strategies, halogen bonds have received increasing attention over the last decade. By

analyzing crystal structures from the Cambridge Structural Database (CSD), Allen et al.

found that the intermolecular contacts between carbon-bonded halogen (C−X, X = Cl, Br, or

I), and electronegative atoms (O or N) were highly directional. They also proposed that the

attractive nature of the interaction was due to electrostatic interactions although polarization,

charge transfer, and dispersion contributions may also be important.18 In further studies on

X-ray crystal structures, scatter plots illustrated that the angles between the covalent and non-

covalent bonds around the halogen in C−I···N interaction were close to 175°.12 Similar

behavior was observed for iodide ion interactions C−I···I−, but the C−I···O interaction angle

showed a larger spread, the mean being ca. 165°. The interaction angle for hypervalent iodine

showed a bimodal interaction geometry where two major interaction angles (ca. 175° and

80°) were observed with the same interaction distance.26

12

1.4.1. Halogen bonding in crystal Engineering

Halogen bonding was applied for synthesis of 1D, 2D and 3D networks in the crystalline

state. Linear and zig zag chains were assembled with the use of bidentate 4,4’-bipyridine and

different (1,2- or 1,3- or 1,4-) dibromotetrafluorobenzenes.27 Similarly, 1,6-diiodoper-

fluorohexane and 1,8-diiodoperfluorooctane (electron-acceptor tectons) were assembled with

the XB acceptor trans-4,4’-azobis(pyridine) to form halogen bonded 1D cocrystals as shown

in Figure 1.11.28

I

I

FF

F

FF

FF

FF

FFF

I

I

FF

F

FF

FF

FF

FFF

NN

NN

I

FF

F

FF

FF

FF

FFF

I

FF

F

FF

FF

FF

FFF

NN

NN

I

F F

F F

I

FF

F F

Figure 1.11. 1D supramolecular crystalline networks Top: 1,6-diiodoperfluorohexane with trans-4,4’-azobis(pyridine) and bottom:1,8-diiodoperfluorooctane with trans-4,4’-azobis(pyridine).

Akeröy et al. have demonstrated the simultaneous action of both hydrogen bonds and

halogen bonds for construction of cocrystal networks. Mixing hydrogen bond donor 2-

aminopyrazine and 4-iodo-2,3,5,6-tetrafluorobenzoic acid gave a cocrystal in which the Py···I

bond directed the assembly of the cocrystal and the amine−acid dimeric synthon was

responsible for propagating the inherent geometry in the resulting supramolecules into

infinite 1D chains.29

O

O

F F

I

FF

NN

NH

H

HO

OF F

I

FF

H

NN

NH

H

Figure 1.12. 1D chain produced through a combination of hydrogen and halogen bonding.

13

Recently, halogen bonded cocrystal 1D networks have been used for resolving mixtures of

α,ω-diiodoperfluoalkanes, and for isolation and characterization of both stable and unstable

polyanions. Mixtures of α,ω-diiodoperfluoalkanes were synthesized by telomerization

reaction of tetrafluoroethylene with iodine and purified by fractional distillation which left a

mixture of heavier inseparable diiodoperfluoalkanes. Treating these inseparable

diiodoperfluoalkanes I–(CF2)m–I, where m = 2, 4, 6, 8, 10 or 12 with a bis-

(trimethylammonium) alkane diiodide [I−(CH3)3N+–(CH2)m+6–N+(CH3)3I−] formed

supramolecular anions I−···I(CF2)mI···I−. The dications formed a crystal framework with the

suitable superanions. The CH2 groups of the dications adopted the all-trans conformation,

and the carbon chains were parallel and interdigitated, thereby forming the required crystal

framework. When there are (m +6) CH2 and m CF2 groups in the bis-(trimethylammonium)

alkane diiodide and the α,ω-diiodoperfluoroalkane respectively, there is a match between the

length of the anion I−···I(CF2)mI···I− formed and the size of the cavities in the dication

framework. When this rule is not maintained, there is a mismatch, the crystal packing is

different and there is no cavity to trap the α,ω-diiodoperfluoroalkane. Based on matching and

mismatching, various I–(CF2)m–I were separated from their mixtures. This was achieved by

mixing solutions of I–(CF2)m–I (m = 2, 4, 6 and 8) with a solution of a given bis-

(trimethylammonium) alkane diiodide (say, m = 4). Only the size-matched (m = 4) adduct

crystallized from the solution and pure (m = 4) α,ω-diiodoperfluoroalkane was then readily

recovered from the solid by raising the temperature (Figure 1.13).30

N Nm

I- I-(CF2)I I

N Nm

-I I-(CF2)I Im m

m = 2, 4, 6, 8, 10, and 12 Figure 1.13. Matching and mismatching various I–(CF2)m–I with bis-(trimethylammonium) alkane to separate α,ω-diiodoperfluoroalkane from their mixtures.

Bis(trimethylammonium)hexane diiodide was also used to encapsulate I2 from solution

whereas bipyridinium-derived cations were used to encapsule I3− and I4−2. 31 In case of

bipyridinium-derived cations, the synergistic effect of size and space matching were used as

an effective strategy to develop a general and reliable protocol where the crystal

stoichiometry is independent of the stoichiometry in the crystallization solution.32

14

Generally, construction of two dimensional 2D networks involves the use of tridentate

halogen bond donors or XB acceptor tectons. The ability of halogen atoms to act as either

donors or acceptors can also be used to generate 2D networks. For example using the

bifurcated bonding ability of XB donor atom, a 2D sheet was created from 2,3,5,6-

tetrafluoro-4-iodobenzaldehyde oxime (Figure 1.14).33

N

H

HO

FF

F F

IN

H

OH

F F

FF

IN

H

HO

F F

FF

I

N

H

HO

F F

FF

I

Figure 1.14. A 2D sheet generated through a combination of self-complementary hydrogen bonds and I···I interactions.

1.4.2. Liquid crystals

Although most liquid crystals are covalent materials, non-covalent interactions such as

hydrogen bonding, quadrupolar and charge transfer interactions, have been used to generate

mesogenic species.34 Due to similarities with hydrogen bonding, halogen bonding has also

been explored for the synthesis of liquid crystal materials. The liquid crystal state can be

reached from the solid state either by the action of temperature (thermotropic liquid crystals)

or solvent (lyotropic liquid crystals). The halogen bonding contribution to thermotropic

liquid crystals can be divided into two portions: a) halogen bonded low molar mass liquid

crystals b) supramolecular liquid-crystalline polymers. The first examples of low molar mass

halogen bonded liquid crystals were obtained by self-assembly; the evaporation of equimolar

solutions of iodopentafluorobenzene and 4-alkoxystilbazoles ((E)-4-(4-alkoxystyryl)

pyridines) afforded the halogen bonded dimers 1.01 (Figure 1.15). This complex, formed by

I of the iodopentafluorobenzene with the N of the alkoxystilbazole, was liquid crystalline in

nature even though each of the components was non-mesomorphic. X-Ray diffraction of the

complex of iodopentafluorobenzene and n-octyloxystilbazole confirmed the absence of

quadrupolar phenyl and perfluorophenyl interactions, and the presence of halogen bond was

clearly evident with an observed N···I distance of 2.811 Å and N···I−C angle of 168.4°.

15

Comparison with the transition temperatures of hydrogen-bonded systems (phenol or benzoic

acid derivatives) suggested that halogen bond and hydrogen bond have effects of similar

magnitudes in the fluid, liquid crystal phase.35

.

H2m+1Cm I F

FF

F F

I

FF

F FOCnH2n+1

1.01

I

FF

F F

OCnH2n+1

1.02

1.03

NO

H2m+1Cm NO

H2m+1Cm NO

H2m+1Cm NO

m = 4, 6, 8, 10, 12n = 4, 6, 8, 10, 12

FF

F F

OI

OOCnH2n+1

1.04

H2m+1Cm NO

FF

F F

IO OCnH2n+1

1.05

H2m+1Cm NO

FF

F F

OI

OOCnH2n+1

F F

FF

1.06

O

NO = NO

Figure 1.15. The molecular structure of the XB complexes between alkoxystilbazole and 4-substituted iodotetrafluorobenzene. Extensions of this work were reported by Bruce et al. who prepared new complexes from

alkoxystilbazole derivatives, 4-substituted pyridine derivatives and 4-substituted

iodotetrafluorobenzene derivatives. In supramolecular dimer 1.02 (m = 10, n = 6), all four

phenyl rings are coplanar and the alkyl chains also adopt a linear trans arrangement coplanar

with the phenyl rings. In the crystal structure, two different dimers of 1.02 (m = 10, n = 6)

were held together in an antiparallel fashion in the plane of the aromatic rings by

intermolecular H···F contacts (2.49 Ǻ). In the case of tetramer 1.03, the mean square plane of

16

the stilbazole rings rotates with respect to the XB donor ring by 31.1° and 13.3°, but in case

of 1.02 the angle was 8.6°. The alkyl chains in 1.03 were not coplanar with the aromatic

rings. However, transition temperatures of these complexes primarily depended on the

stilbazole chain length, and on the alkyl chain length of the iodobenzene to a lesser extent.

For complexes between octylpyridine and esters 1.07-1.10 (Figure 1.16), the transition

temperatures were significantly lower than expected in spite of the presence of three aromatic

rings.36

n = 4, 6, 8, 10, 12

FF

F F

OI

OOCnH2n+1

1.07

NC8H17

FF

F F

IOCnH2n+1

NC8H17

FF

F F

OI

OOCnH2n+1

F F

FF

NC8H17

1.09

1.10

1.08

FF

F F

IO OCnH2n+1

NC8H17O

Figure 1.16. The molecular structure of the XB complexes between 4-substituted pyridine and 4-substituted iodotetrafluorobenzene.

The telechelic donors 1,3-diiodoperfluorobenzene or para-iodoperfluorophenol self-

assembled with the stilbazole to give rise to bent core systems with 2:1 stoichiometry. Once

again, both starting materials were nonmesogenic, but 1.11 and 1.12 complexes showed

monotropic liquid crystalline phases on cooling (Figure 1.17).37

.

17

CmH2m+1

O

FF

I

FF H

NO

N

O

H2m+1Cm

F

I

FF

F

1.121.11

NO = NO

H2m+1Cm

N

O

I

CmH2m+1

N

O

Figure 1.17. The bent-core halogen bonding complexs between stilbazole and 1,3-disubstituted perfluorobenzene.

A layer-by-layer polymer 1.13 was assembled on the basis of halogen bonding between

poly(4-(4-iodo-2,3,5,6-tetrafluorophenoxy)-butyl acrylate) and poly(4-vinylpyridine) shown

in Figure 1.18. This multilayer water insoluble polymer was synthesized in methanol.38

OO

O

In

FF

FF

Nn

1.13 Figure 1.18. Halogen bonded polymer.

1.4.3. Metal-organic frameworks

Due to their redox, magnetic, and optical properties, metal ions can play an interesting role in

the area of supramolecular chemistry. The semi-occupied d and f orbitals may give rise to

such properties as strong absorption, high quantum yields, luminescence, tunable redox

states, and suitable excited state lifetimes. The splitting of d orbitals, in response to ligand

field strength, can influence spin crossover phenomena along with physical properties.39 In

organometallic networks, M−C bonds are longer than C−C bonds, leading to increases in the

sizes of cavities and channel sizes of the crystal networks. Metal-organic compounds that

contain halogen−halogen interactions are divided into two groups: type I, where 140° < θ1 =

18

θ2 < 160° and type II where θ1 is close to 180° and θ2 is close to 90°. The later type is halogen

bonding.40

C X

X CC X X

C

Type I Type II

θ1

θ1

θ2

θ2

Figure 1.19. Geometries of type I and II halogen−halogen interactions. Type I: θ1 = θ2; and type II: θ 1 = 90° involving the halogen bond acceptor role, θ 2 = 180° involving the halogen donor role.

Halogenated and positive charged heterocyclic systems such as halopyridinium derivatives

have been commonly used for maximizing halogen bond donor ability in complex

formation.41 In M−X···X−C interactions involving halopyridinium derivatives, the halogen

bonding order is I > Br > Cl and the bond angles are close to 90° and 180° for θ1 and θ2

respectively. Complexes trans-[MCl2(3-X-py)2] were synthesized by the reaction of ML4-2

(M = Pd, Pt) with 3-X-pyridine (X = Cl, Br, I) in acidic aqueous solution. After

crystallization, M−X···X−C interactions were observed in the crystal. Similarly,

halopyridinium derivatives also formed complexes with Cu, Au, Co, Pd, Ni and Ru salts by

halogen bonding interactions. Halogen substituted imidazolium, imidazolinium,

benzodithiazolium, and anilium derivatives also showed halogen bond donor properties.41

The design of conducting materials was achieved relying on interactions of

diiodo(ethylenedithio)tetraselenafulvalene (DIETSe) (donor molecules) with anionic

acceptors. Here stronger π-d exchange was proposed to occur when intermolecular

halogen···halogen contacts were present, promoting exchange interactions between the

conducting π-electrons of the donors and the localized d-electrons of the anions.

(DIETSe)2FeCl4 showed a metal–semiconductor transition at 11 K (Figure 1.20).42

19

N N

X

NN

X

TeX

XX

X

X

X

S

S

Se

Se

Se

Se I

I

FeCl Cl

Cl Cl

S

S

Se

Se

Se

Se I

I

S

S

Se

Se

Se

SeI

I

S

S

Se

Se

Se

SeI

I

1.14 1.15 Figure 1.20. Structure of the (2-haloimidazolium)2(TeX2)salt (1.14) and (DIETSe)2FeCl4 (1.15).

1.4.4. Nanomaterials

Shirman et al. have employed halogen bonding as a tool for stepwise assembly of gold

nanoparticles onto a planar surface.43 Previously they demonstrated the formation of highly

crystalline halogen bonded thin films by vapor deposition on silicon substrates, where the

solid state structure was independent of the substrate surface chemical functionalization.44

They also demonstrated the supramolecular assembly of gold nanoparticles (AuNPs) by

activating with an XB donor ligand. To form the AuNP-1, they substituted

tetraoctylammonium bromide (TOAB) from AuNP−TOAB with the polar N-oxide moiety of

(E)-4 (2,4,6-trifluoro-3,5-diiodostyryl)-pyridine-1-oxide (FIPO). The AuNP- FIPO particle

was formed by coordination of the polar N-oxide moiety of FIPO to the gold surface and

addition of a bifunctional halogen bond linker (BPEB) to the solvent mixture. The final

aggregates AuNP−FIPO /BPEB which were formed could be controlled by time during the

first assembly steps and by varying the concentration of the linker (BPEB). The aggregation

process (scheme 1.03) was monitored by UV/Vis spectroscopy and TEM.45

20

Scheme1.03. The formation of assembly consisting of functionalized AuNP with organic cross-linkers (BPEB) in solution. Reprinted with permission from reference 42. Copyright 2011, American Chemical Society.

The gold nanoparticles (AuNPs) were assembled on silicon and quartz substrates which were

functionalized by p-chloro benzyl derivatives and reacted with (E)-1,2-di(pyridin-4-yl)ethane

or (E)-4-(2-(pyridin-4-yl)vinyl)phenol to construct M1 or M2. The reactions of AuNP-1 with

M1/ M2 in presence of BPEB assembled AuPN-1 on M1/ M2 planar surfaces. The progress

of the assembly of M1-AuNP-1- BPEBB could be monitored by color change, from red to

deep blue after 10 AuNP-1 deposition steps (scheme 1.04).

N

N

=

Scheme 1.04. Stepwise generation of assemblies consisting of functionalized gold nanoparticles (AuNP-1) and organic cross-linkers (BPEB) on silicon organic monolayers (M1, M2). Adapted and reprinted with permission from reference 42. Copyright 2011, American Chemical Society.

21

1.5. Halogen bonding in Medicinal chemistry Although halogens are not generally prevalent atoms in biological molecules, it is clear that

they can play important roles in some biological systems. In medicinal chemistry, one of the

common strategies to develop new drugs is insertion of halogen atom /atoms on hit or lead

compounds to exploit their steric effect in the binding sites of molecular targets having

deeper pockets. Other than steric effects, halogen atom incorporation may increase

membrane permeability and/or blood-brain barrier (BBB) permeability or block sites of

metabolism.46 Roughly one fourth of total number of papers and patents related to medicinal

chemistry involve halogen atom insertion to the final compounds as per Scifinder, Web of

Science, Science Direct and ACS. The majority of the halogenated drugs are fluorine

containing drugs followed by chlorine, bromine and iodine. In structure-activity relationship

(SAR) discussions the effect of halogenation is largely confined to steric effects or changes

in lipophilicity. The role of halogen bonds in protein-ligand complexes was pointed out by

Auffinger et al. (2004) who studied halogen bonding interactions between halogenated

ligands and cyclin-dependent kinase 2 (CDK 2).47

Recently, Hadegger et al. studied halogen bonding in human Cathepsin L (hCat L)

inhibitors.48 The nitrile group was responsible for binding with the hCat L in S1 pocket,

forming a covalent thioimidate adduct while ArSO2- interacts with S2 by lipophilic

interaction and the other aryl group interacts with the oxyanion hole pocket S3 through a

halogen bonding interaction involving p-X−Ar (X = Cl, Br, I, and CF3) (Figure 1.21).

Experimental results were consistent with a halogen bonding interaction between Gly61 in

pocket S3 with the inhibitors 1.16 (X = Cl, Br, I, and CF3). Methyl substitution at the 4 X-

aryl position, which occupied the S3 receptor site, also showed moderate activity probably

due to the flexibility and reorganization of S3. X- ray crystal structures showed the O···Cl−C

angle was 174° (close to 180°) and the O···Cl distance was 3.10Å consistent with a halogen

bonding definition. The free energy gain with this halogen bond was as much as ∆∆G = −2.6

kcal mol-1 from X = H to X = Cl (van der Waals radii of Cl and CH3 are1.75 and 2.00 Å

respectively).

22

N

S

Cl

O

X

ONHN

OO

N

S

Cl

O

X

ONHN

OOO

F3C

SH

HN

O

HN

OCys25

gly61

S3

S1

S2

1.16

1.17

Figure 1.21. Binding mode of covalent inhibitors 1.16 and 1.17 at the active site of hCat L with its three pockets.

Table 1.01. Activity data for hCatL inhibitors.

Inhibitor

IC50 Log D X Inhibitor IC50 Log D

1.16a 0.29 2.11 H 1.17a 0.59 n.d

1.16b 0.13 2.57 Me 1.17b 0.81 n.d

1.16c 0.34 2.36 F 1.17c 0.93 n.d

1.16d 0.022 2.73 Cl 1.17d 0.036 2.19

1.16e 0.012 2.96 Br 1.17e 0.0081 >3.0

1.16f 0.0065 3.23 I 1.17f 0.020 n.d

1.16g 0.095 3.12 CF3 1.17g 0.22 n.d

Fedorov et al. (2011) introduced dichloroindoyl enaminonitrile KH-CB19, a potent and

highly specific inhibitor of the CDC2-like kinase isoforms 1 and 4 (CLK1/CLK4). The ATP

binding site in the protein kinase contains a hinge connection between the two protein lobes

and forms H-bonds to bound ATP through NH and CO, leaving a hinge CO as an available

accepting site. The X-ray crystal structure of KH-CB19 with CLK1 and CLK3 revealed a

non-ATP-mimetic binding mode involving conformational changes in helix αC and the

23

phosphate binding loop and halogen bonding to the kinase hinge region. KHCB19 effectively

suppressed phosphorylation of SR (serine/arginine) proteins in cells (Figure 22). 49, 50

NH3C

Cl

Cl

O

O

CN

H2N

18

Figure 1.22. X-ray cocrystal structure bound to inhibitors KH-CB19 (1.18) and CLK3. Reprinted with permission from reference 49. Copyright 2011, Elsevier.

As protein kinase Cdk9 inhibition contributes to anticancer activity, inhibitors have been

under clinical investigation, including 5,6-dichlorobenzimidazone-1-β-D-ribofuranoside

(DRB), studied by Baumli et al. (2010). DRB, an adenosine analog, binds to the ATP binding

site via halogen bonds from the chlorine to the CO in the hinge region that connects the N-

and C-terminal kinase lobes. It also induced conformational changes in the glycine-rich loop

and the β3-αC loop that enclose the inhibitor as shown in Figure 1.23.51

24

NN

N

N

O

HOHO

O

N

PO

O

P OO O

O

NH

OHN

OHN

OHN

HHPhe105

Asp104

Cys106

NH

OHN

ONH

OHN

Phe105

Asp104

Cys106

N

N

OHO

HOOH

Cl

Cl

ATPDRB (1.19)

Figure 1.23. The chemical structures of adenosine triphosphate and 5,6-dichlorobenzimidazone-1-β-D-ribofuranoside (DRB, 1.19) bound to a Cdk9 Hinge Region An important iodinated compound in human physiology is L-thyroxine, 3,5,3',5'-tetra-

iodothyronine (1.20), present in thyroid hormones. Halogen bonding is thought to play a role

in thyroxine binding to its transport protein transthyretin, and this is also supported by short

I···O contacts that result when bound. Thyroxine binding to RNA sequences has also been

attributed to halogen bonding, and the T3 (thyroxine having three iodine atoms) and T4

(thyroxine having four iodine atoms) forms have been identified as the active forms of the

protein.27

O

OH

O

NH2

I

II

IHO

1.20 Figure 1.24. L-thyroxine, 3,5,3',5'-tetraiodothyronine (1.20) is the precursor of thyroid

hormone.

1.6. Halogen bonding in the solution phase In spite of enormous success in explorations of XB in computational studies, the gas phase,

metal organic frameworks, and polymer chemistry, there were limited studies of halogen

bonding in solution before 2010. The first such work was done by Hildebrand, Glascock

(1909), and Benesi, Hildebrand (1949) to determine the binding constant of the iodine-

25

benzene complex in CCl4 and n-heptane solvent as well as the iodine-mesitylene complex in

CCl4 and n-heptane solvent.6 Mulliken et al. (1959) studied the formation of iodine

complexes with benzene and pyridine in heptane followed by binding constant determination

using both UV and IR spectroscopy. The binding constant of Py·I2 at 26 °C was 200 M-1

which roughly two orders of magnitude stronger than that of the benzene·I2 complex. The

observed frequencies for Py·I2 and Bn·I2, 180 cm-1 and 205 cm-1 respectively, were slightly

lower than that of I2 vapor (213 cm-1) because of partial loosening of intrahalogen bond as

well as an increase in effective mass due to partial formation of a donor-halogen bond.52 The

binding constants of pyridine and pyridine derivatives (4-MePy, 2,6-Me2Py, and 2,4,6-

Me3Py) were determined in CCl4 and CBr4 and gave significantly lower values. Similar

studies with I2 and pyridine derivatives were helpful for understanding steric and substituent

effects on halogen bonding. In the same study, THF showed much weaker binding strength

as shown in Table 1.02.53

Table 1.02. Binding constants (M-1) of complexes between pyridine derivatives and THF with halogen bond donors (CCl4, CBr4, and I2).

CCl4 CBr4 I2

Py 0.14 ± 0.04 0.83 ± 0.16 117 ± 10

4-MePy, , and 0.32 ± 0.07 1.2± 0.22 260 ± 10

2,6-Me2Py 0.10 ± 0.02 0.41 ± 0.08 74 ± 7

2,4,6-Me3Py 0.09 ± 0.02 0.64 ± 0.18 125 ± 10

THF 0.08 ± 0.05 0.46 ± 0.10 1.9 ± 0.2

To understand the halogen bonding interactions of haloforms with different amines, ketones,

and esters, NMR spectroscopy was used based on the chemical shift (δ) change with respect

to the same haloform proton chemical shift in hexane. On the other hand, chemical shift (∆δ)

changes were poorly correlated for anisotropic solvents e.g benzene, pyridine derivatives,

and amides.10 The 1:1 complex formed between trifluoroiodomethane and 2,4,6-trimethyl-

pyridine in cyclopentane solution was studied, and the enthalpy and entropy of formation

were −20.9 × 0.4 kJ mol-1 and −52.7 × 1.6 JK-1mol-1, respectively.54 Laurence and co-

workers reported that the association constant of 1-cyano-2-iodoacetylene with Ph3PO was

higher in CCl4 than in benzene. Kochi and co-worker also studied solvent effects on halogen

26

bonding in the context of the quinuclidine···CBr4 and CBr4···DABCO complexes. The

association constants for CBr4···DABCO decreased as a function of solvent, in the order

CH3CN > CHCl3 > CH3OH. Experimental investigation of n-C3F7I···toluene and 2-

C3F7I···toluene complexes determined the enthalpies and entropies −2.9 kJmol-1 and −19.1

JK-1mol-1 for 2-C3F7I···toluene-d8 complex and −2.7 kJmol-1 and −16.0 JK-1mol-1 1-

C3F7I···toluene-d8 complex. These were correlated with the predicted results.55

Cl

FFF

173.9ο Br/I

FFF

180ο

NH

H

H

H HHH

H

HN

HH

H

H HHH

H

H

Figure 1.25. Equilibrium geometries of CF3X···N(Me)3 complexes (X = Cl, Br, or I)

The halogen bonding complexes between trimethylamine and trifluorohalomethane (CF3I,

CF3Br, and CF3Cl) were investigated using FT-IR spectroscopy. Using MP2/6-311++G(d,p)

+LanL2DZ basis set, the complexes of CF3X (X = Cl, Br, or I) with NMe3 showed C−X···N

(180°) with C3v symmetry for CF3I···NMe3 and CF3Br···NMe3 complexes but Cs symmetry

was observed for CF3Cl···NMe3 complex (angle C−Cl···N = 173.9°). The tilt of the

CF3Cl···NMe3 complex was caused by additional stabilization of van dar Waals contact

between one of the fluorine atoms and a hydrogen atom of one of the methyl groups in

trimethylamine. The CF3 stretching frequencies of CF3I···NMe3, CF3Br···NMe3, and

CF3Cl···NMe3 complexes in xenon, kryton and argon respectively were compared with the

spectra of respective monomers and assigned to 1:1 complexes. Van’t Hoff plots of these

complexes gave complexation enthalpies of −28.7(1) kJ mol-1for CF3I···NMe3, −18.3(1) kJ

mol-1 for CF3Br···NMe3, and −8.9(2) kJ mol-1for CF3Cl···NMe3.56 On the other hand, the

weakly bound molecular complexes between Me2O and CF3I, CF3Br, and CF3Cl in liquid

argon and krypton were investigated, using both IR and Raman spectroscopy. Ab initio

calculated geometries of CF3I···OMe2, CF3Br···OMe2, and CF3Cl···OMe2 complexes showed

that the bond angles for O···X−C were 180 ± 0.5° in all cases. The experimental standard

complexation enthalpies for 1:1 complexes were derived using the Van’t Hoff equation and

found to be −6.8(3) kJ mol-1 for Me2O···CF3Cl, −10.2(1) kJ mol-1 for Me2O···BrCF3, and

27

−15.5(1) kJ mol-1for Me2O···ICF3. At higher concentrations and low temperatures for the

Me2O / CF3Br mixture, the presence of an additional band near 916.8 cm-1 indicated the 1:2

Me2O:(CF3Br)2 complex, whereas the band at 923.8 cm-1 corresponded to the 1:1 complex

between Me2O / CF3Br at low concentration. The complexation enthalpy of the 1:2 complex

is −17.8(5) kJ mol-1, which is slightly less than twice the experimental enthalpy −10.2(1) kJ

mol-1 for the corresponding 1:1 complex (Figure 1.26).

X

FFF

H

HH

O

H

HHO

HHH

H HH

Br F

FF

Br

F

FF

Figure 1.26. Equilibrium geometries of CF3X···O(Me)2 complexes where ((X = Cl, Br, or I) and (CF3Br)2···O(Me)2

One of the most fascinating uses of halogen bonding in solution phase is Lewis acid-base

scales (pKBI2) introduced by Laurence group.57 Unlike Bronsted acid and bases, there is no

single reference Lewis acidity or basicity scale. Existing Lewis acid-base scales were

constructed against a particular acid such as BF3, 4-FC6H4OH, I2, Li+, Na+, K+, Al+, Mn+,

CpNi+, and CH3NH3+. These scales are based on dative bond, hydrogen bond, halogen bond,

and cation-molecule interactions.58 The new diiodine basicity scale pKBI2 is quasi-orthogonal

to the hydrogen bond, dative bond, and cation bonds scale. This scale covers an extended free

energy range of 9 kcalmol-1 and the order of diiodine basicity of bases followed the sequence

R2NH > R2PH ≈ R2Se > R2S > RI ≈ R2O > RBr > RCl > RF. According to the HSAB (Hard-

Soft Acid-Base) theory, I2 is a soft acid. The diiodine basicity of halogen, chalcogen and

pnictogen bases follow the sequences, I > Br > Cl > F, and Se > S > O, and N > P

respectively. Contradictory to predictions based on HSAB, diiodine preferred to bind to

amines than phosphanes. The diiodine basicities of oxygen functionalities occur in the

following order NO ≈ AsO > SeO > PO > SO > C=O > -O- > SO2. In these series, the

chalcogen and pnictogen oxides basicity appeared to be correlated with the electrostatic

potential on oxygen atom. The diiodine basicity order also found in sulfur and nitrogen bases

and these are PS >> -S- ~ C=S >> N=C=S and amine > pyridine > nitrile. The high accepts

ability of the PS moiety can be explained in terms of the zwitterionic contribution of P+−S-,

28

and the order of nitrogen bases can be explained by the decreasing p character of the lone

pair.

Figure 1.27. The computed B3LYP/LANLdp electrostatic potential, on the 0.001 electrons/bohr surface of I2.

The electrostatic potential surface of I2 reveals the σ-hole as well as a belt of negative

electrostatic potential surrounding the σ-hole as shown in Figure 1.27. Exploiting the σ-hole

of diiodine as a halogen bond donor and the surrounding belt of negative charge as a

hydrogen bond acceptor, pKBI2 was determined for heterocyclic amines. The value of pKBI2

of 2-aminopyridine (2.45) is higher than the pKBI2 of N,N-dimethylaminopyridine (0.95). The

difference of 1.50 pKBI2 units is due partially to the steric effect of the NMe2 group compared

to NH2 group and partially by the existence of an intramolecular NH···I hydrogen bond,

assisting to the formation of N···I halogen bond. Though the diiodine scale is orthogonal to

the manganese cation, lithium cation, aluminium cation, trimethylboron, and 4-fluorophenol

basicity scales, it correlated significantly with cyclopentadienylnickel cation and copper

cation basicity scales which can be explained due to the soft Lewis acidic character of I2,

CpNi+, and Cu+.57

N

NH H

II

N

NH3C CH3

II

N

NH

N

NPh H

II

II

1.21 1.22 1.23 1.24pKBI2 = 2.45 pKBI2 = 2.08pKBI2 = 0.95 pKBI2 = 2.10

Figure 1.28. Hydrogen bond and halogen assisted interaction between heterocyclic amines (2-aminopyridine 1.21, N,N-dimethylaminopyridine 1.22, N,-phenylamino pyridine 1.23, 1H-pyrrolo[2,3-b]pyridine 1.24) and diiodine.

29

Recently, Erdélyi et al. (2012) successfully applied an NMR spectroscopic method called

isotopic perturbation of equilibrium (IPE) to probe the symmetry of bis(pyridine) halonium

halogen bonded systems (Figure 1.29). Perrin had previously shown that N···H···N-type HBs

was asymmetric in solution, but symmetric in crystals. Analogous iodous 1.26 and bromous

1.25 halogen bonded (XB) systems preferred symmetric arrangements in solution whereas

1.26 preferred symmetric but 1.25 asymmetric arrangements in the solid state.59

N

X

N

N

XN

a b

NX

N

b

(D)H (D)H (D)H

Figure 1.29. The symmetric (a) and asymmetric (b) geometries of the systems bis(pyridine) bromonium and iodonium triflates (1.25, X = Br, for 1.26, X = I and for 1.27, X = H).

1.6.1. Halogen bonding in supramolecular chemistry

The polyaromatic perylene diimide core of C- and S-shaped receptor clefts was designed to

probe interactions between phenazine (1.30) nitrogens and halogens. Using UV/vis

spectroscopy, the determined binding constants of halogen-containing clefts receptor 1.28a,

1.28b and 1.28c with phenazine (1.30) were 480 ± 50, 1100 ± 100 and 2200 ± 200 M-1

respectively in benzene solution whereas the binding constants of receptors 1.29a, 1.29b and

1.29c were 780 ± 80, 900 ± 90 and 760 ± 80 M-1. The cleft-like shapes were well suited for

the investigation of weak intermolecular forces because they provide convergent functional

groups: the acid for hydrogen bonding, π-stacking, and halogen bonding interaction with the

substrate held within. The energy differences were 0.5 and 0.9 kcal/ mol for nitrogen-

chlorine (1.28b and 1.28a) and nitrogen−bromine interactions (1.28c and 1.28a). The energy

differences were 0.3, 0.1, 0.6 kcal/ mol for among the C- and S-shaped (1.28a and 1.29a),

(1.28b and 1.29b) (1.28c and 1.29c) receptors, respectively.60

30

N

N

N N

O

OO

O

N N

O

OO

O

OO

O

O

XO

HO

X

O

HO

a = H, b = Cl, C = Br

C - shape

S - shape

1.28

1.29

1.30

Figure 1.30. Perylene diimide based C-shaped receptors (1.28), S-shaped receptors (1.29) and phenazine (1.30). Adams et al. determined the free energy of arene-halogen interactions in CHCl3 solvent. The

interactions between nitropyrrole-containing bisaniline derivatives (1.31, 1.32, 1.33, and

1.34) with isothaloyl derivative (1.35) were studied with double-mutant-cycles. As the

geometry of the zipper complexes is locked, the architecture of the system from one side

allows them to determine the thermodynamic contribution of relatively weak interactions.

The measured halogen−aromatic interaction (repulsive) free energies (∆∆G) were 3.0, 3.2

and 1.2 kJmol-1 for 1.32, 1.33 and 1.34 where X= F, Cl and Br respectively, compared to the

free energy with 1.31.61

31

N NH

OX

X

X

O

NO2N

O

O

NH

N

O2NNO2

H

H

H

1.35

1.31 = H1.32 = F1.33 = Cl1.34 = Br

φ

Figure 1.31. Interactions between the nitropyrrole containing bisaniline derivatives (1.31, 1.32, 1.33 and 1.34) with isothaloyl derivative (1.35). (φ = 30 ≈ 37°)

Brammer et al. (2008) reported experimental studies of the interaction enthalpies and

entropies of C−F···I−M halogen bonds, determined by titration of a nickel fluoride complex

trans-[NiF(C5F4N)(PEt3)2] with the halogen bond donor C6F5I. Thermodynamic data in

heptane solvent for trans-[NiF(C5F4N)(PEt3)2]·C6F5I (∆H° = −26.1 kJmol−1, ∆S° = 63.4

Jmol−1K−1) (1.36) and for trans-[NiF(C5F4N)(PEt3)2]·2C6F5I (∆H° = −21.1 kJmol−1, ∆S° =

83.5 Jmol−1K−1) (1.37) indicated a higher entropic penalty for the trimolecular complex, as

expected.62 For the synthesis of trans-[NiCl(C5F4N)(PEt3)2] (1.38), a halogen bonding

interaction (Ni−F···Cl−C) might be involved in the substitution of the metal-bound fluorine

by Me3SiCl in THF solvent (Figure 1.32).63 Although the nickel fluoride complex trans-

[NiF(C5F4N)(PEt3)2] interacted with pentafluoroiodobenzene (C6F5I) in 1:1 and 1:2 ratios,

the complex formation ratio in toluene was 1:1. In an extension of this work, Perutz et al.

(2011) studied Ni, Pd, and Pt fluoride complexes in toluene (d8). Using 19F NMR titrations at

various temperatures, the enthalpies and entropies of interaction were determined. The trend

for halogen bonding interactions based on −∆H° was Pt > Pd > Ni. In terms of ligands,

Thermodynamic studies of halogen bonding in …...ii Thermodynamic studies of halogen bonding in...

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