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STUDIES ON CHARGE TRANSFER COMPLEXES

BETWEEN ORGANIC COMPOUNDS

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DISSERTATION 4/ .V,

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SUBMITTED FOR THE AWARD OF THE DEGREE OF

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ifSafifter of $l)tlosiop{)p

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BY LAL MIYAN '-' ^

UNDER THE GUIDANCE OF

PROF. AFAQ AHMAD

DEPARTMENT OF CHEMISTRY LL ALIGARH MUSLIM UNIVERSITY

ALIGARH ( I N D I A ) 2014

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A NOV 2014 DS4397

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CANDIDATE'S DECLARATION

I, Lai Miyan, Department of Chemistry certify that the work embodied in this M.Phil

dissertation is my own bonafide work carried out by me under the supervision of Prof. Afaq

Ahmad at AUgarh Muslim University, Aligarh. The matter embodied in this M.Phil

dissertation has not been submitted for the award of any other degree.

I declare that I have faithfully acknowledged, given credit to and referred to the research

workers wherever their works have been cited in the text and the body of the dissertation. I

further certify that I have not will fully lifted up some other's work, para, text, data, result,

etc. reported in the journals, books, magazines, reports, dissertations, theses, etc., or available

at web-sites and included them in this M. Phil dissertation and cited as my own work.

Date: l.^ AL/.l^n/M

(Signat c candidate)

LAL MlyAN (Name of the candidate)

Certificate from the Supervisor

This is to certify that the above statement made by the candidate is correct to the best of my knowledge.

Signature of the Supervisor ^ :p-f'

Name & Designation:

Department

Prof. Afaq Ahmad (Supervisor)

Chemistry

(Signature p e r

n of the Department with seal)

COURSE/COMPREHENSIVE EXAMINATION COMPLETION CERTIFICATE

This is to certify that Mr. Lai Miyan Department of Chemistry has satisfactorily

completed the course work/ comprehensive examination requirement which is part of

his M.Phil, programme.

Date:- (Signati^,

^ 1 i<iy

airma^ of the Department)

ACKNOWLEDGEMENTS

All (hanks lo Allah who is mosl henejkcnt and merciful. Xoihin^^ is possible Miliiota

his blessings including this work. I owe my deepest gratitude to Prof. Afaq Ahmad.

my worthy supervisor. I express my sincere thanks lo him /or his vahiuhle guiJunL i /.•.•

carrying out work under his supervision, encouragement and cooperation. His

visiomiry thoughts hare influenced me greatly. His dynamical allinnle has enipowc'-a!

me with zeal of energy to conquer the minor details of my research work. Under his

supervision I really fell freedom and curiosity of learning new things and gene'-ailny

ideas. I wish to express my thanks for proposing this investigation and for the many

helpful discussions and suggestions concerning the problems arising out during die

course of the research.

/ also wish the deepest sense of gratitude to the Prof. Zafar A. Siddiqi, Chair num.

Department of Chemistry, Aligarh Muslim University, Aligarh, for his constant

support and for providing facdities to carry out experimental work This thanks u>

University Sophisticated Instrument Facility (USIF), Instrumentation center fin-

providing research facilities and Department of Applied f^hysics. Aligarh Mw-iini

University, Aligarh, for providing the SEM andXRD facilities. Thanks are also due to

SAIF. Punjab University Chandigarh for providing '11 \'\IR and "C XMR spcclra

I am also sincerely thankful to Dr. Ishaat Mohammad Khan, for his keen guidance

and support for carrying out experimental works. His great hnmledge and iii-.v/t/i.,••;/;,••/

attitude help me tremendously; his kindness, patience is much appreciable. 1 could

not finish my study without his help and encouragcmcnl. I would like lo pay special

thanks to my teachers. Prof S.A. Nabi, PofM. Shakir.Profi Sartaj Tabassum. Prof

Rqfiuddin, Prof Riyazuddin, Dr. Mohd. Akram, Prof .Jawcdd Iqhul. Prof Abdul Ruuf

and Prof (Mrs.) Farrukh Arjmand for their immense help and encouragement. A

special word of thanks to my parents, laboratory colleague'^ and seniors Dr

Noorusaba, Dr. Neeti Singh, Ms. Neelam Singh, Mr. Zulkar Main, Mr. Mohd.

Shoeb, Mr. Mohd. .4sif, Mr. .4bad .4Ii, Mr. Khairujjaman, Mr. Nayeem .Ahmad, for

their constant help in carrying out experimental works and finalizing the report.

There are no special words to express my gratitude lo my parents and brolher<< ,/v ,i

source of inspiration and encouragement. Above all I am thankful to Allah, the

ultimate .source of knowledge: with his blessings in completing the work.

(Lai Mivan)

CONTENTS

Chapter-I Page No.

GENERAL INTRODUCTION 1-46

1.1. Theory of Charge Transfer Complexes 3-7

1.2. Mulliken's Theory 7-12

1.3. Dewar's Theory 12 15

1.4. Ligand-Metal and Metal- Ligand Charge Transfer 15-19

Complexes

1.5. Spectrophotometric Determination of Equilibrium Constant and

Molar Absorptivity 19-2!

1.6. Hydrogen Bonding 21-28

] .7. Interatomic interaction in Charge Transfer Complexes 28-36

1.8. Recently Research Work on Charge Transfer Complexes 36-38

References 39-46

Chapter-II 47-56

SYNTHESIS, SPECTROPHOTOMETRIC AND SPECTROSCOPIC

STUDIES OF 1, 2-DIMETHYLIMIDAZOLE (DMI) AS AN ELECTRON

DONAR WITH 2, 4-DINITRO-l-NAPHTHOL (DNN) AS AN

ELECTRON ACCEPTOR IN DIFFERENT POLAR SOLVENTS AT

ROOM TEMPERATURE.

2. Introduction 48-49

2.1. Experimental Study 49

2.1.1 Material and methods 49

Procedure 49

2.2.1 Synthesis of solid charge transfer (CT) complex 49

2.2.2 Preparation of standard stock solutions 49 50

0 O

2.3. Spectral Measurement and Determination of formation 50-51

constant

2.4. Result and Discussion 51 -53

2.4.1 Observation of CT electronic spectra 51 -53

2.5. Conclusions 54

References 55-56

CHAPTER-1 GENERAL INTRODUCTION

1. INTRODUCTION

Charge transfer complexes (CTC) are an electron donor/ electron acceptor association

for which an intermolecular electronic charge transfer transition is observed. An

electron-donor-electron acceptor-complex, characterized by electronic transition to an

excited state in which there is a partial transfer of electronic charge from the donor to

the acceptor moiety. The nature of this transition is made apparent in theoretical

discussion to follow, but experimentally, a charge transfer complex is typically

identified spectrally. By the combination of two compounds, absor- ption maximum

appear that are not characteristic of either of the compound alone. The current view is

that the electronic transition is associated with the transfer of electron from the donor

to the acceptor.

The attraction in a charge transfer complex is much weaker than covalent forces;

rather it can be better characterized as a weak electron resonance. The formation of

charge transfer complexes also depend upon the polarity of the solvent and hydrogen

bonding between the donor and acceptor. In many instances there appear charge

transfer bands although no complexes are formed. The nature of these types of charge

transfer complexes is discussed in connection with the theory of spectral transition

but it may be mentioned here that the molecular interaction occurs when random

collision of pair permit an overlap between the lowest virtual orbital of the acceptor

and donor molecular orbital. Since these pairs are not associated with each other for

any long time, they do not form a stable complex, and therefore there is no minimum

in the potential energy surface describing ground state. Charge transfer (CT)

complexes play an important role in hormone action, in material science, in drug

design and their carcinogenic activity (with possible formation of CTC). They are also

important in the field of analytical chemistry and organic semiconductors and recent

work have provided that appear as intermediates in various organic and inorganic

reactions. The charge transfer complexes can be characterize with the help of most

powerful optical spectroscopic technique, NMR, FTIR, TGA/DTA etc. Up to now the

factors governing the behavior, stability etc. of these complexes have not been clearly

established. Among the different theory proposed, one of the best known was

Mullikan Theory, Dewar's Theory Ligand- Metal and Metal-Ligand Charge Transfer

Theory.

1.1. Theory of Charge Transfer Complexes

Benesi and Hildebrand [1-2] studied the effect of various solvents on the absorption

spectra of molecular iodine. They noted that a mixture of aromatic hydrocarbons (e.g.

benzene) and iodine possessed the charge absorption maximum not present in the

spectra of either benzene or iodine. They attributed this new band to the formation of

an adduct between the two components and began to examine the nature of this

complex by altering substituents on benzene. From the spectral changes resulting

from the addition of electron withdrawing or releasing groups to the benzene, it was

concluded that these complexes were the result of an acid- base interaction in the

Lewis sense.

The charge transfer complexes (also called donor-acceptor complexes) are generally

coloured. It has been known for a long time that hydrocarbons such as quinone is

yellow in colour and quinhydrone crystals possess a highly green metallic colour. In

the quinhydrone complex, the donor and acceptor orbitals are delocalized 7i-molecular

orbitals and the chemical bond between the two components of the complex is a

delocalized bond. The theory of charge transfer complex could explain that the

aromatic carbonyl compounds have absorption bands which can be identified as

Ti 'Ti* and n^-Ti* and charge transfer (CT) transitions. Of these transitions, the CT

transition has the highest intensity, which means that they have the highest value of

the extinction coefficient and of the oscillator strength. Theoretical model of charge

transfer complexes basically depend upon the nature of donors and acceptors and also

the polarity of solvents. This is practically true when one component is a good

electron donor (has a high electron affinity or has a low ionization potential) and other

is a good electron acceptor (has a low electron affinity or has a high ionization

potential). Charge transfer complexes are as formed in biological systems, e.g., charge

transfer takes place from donor to acceptor [3-4] during drug action, enzyme catalysis,

hydrogen bonding ion transfer through lipophilic membranes. Various aromatic

molecules can behave as electron donors and form molecular complexes with electron

acceptor molecules such as halogens, nitro compounds and quinines [5-6]. Extensive

works have been carried out to elucidate the nature of intermolecular interactions in

these molecular complexes. Mulliken has developed the theory of the intermolecular

CT interactions, which has been applied successfully to the interpretation of the

absorption bands characteristic of molecular complexes in various systems [7].

The nature of the attraction in a charge-transfer complex is not a stable chemical

bond, and is thus much weaker than covalent forces [8-11]. Many such complexes can

undergo an electronic transition into an excited electronic state. The excitation energ>'

of this transition occurs very frequently in the visible region of the electro-magnetic

spectrum, which produces the characteristic intense colour for these complexes. These

optical absorption bands are often referred to as charge transfer bands (CT bands).

Optical spectroscopy is a powerful technique to characterize charge-transfer.

Charge-transfer complexes exist in many types of molecules, inorganic as well as

organic, and in solids, liquids, and solutions. A well-known example is the complex

formed by iodine when combined with starch, which exhibits an intense blue charge-

transfer bands Charge-transfer complexes are formed by weak association of

molecules or molecular subgroups, one acting as an electron donor and another as an

electron acceptor. Methods have been developed to determine the equilibrium

constant for these complexes in solution by measuring the intensity of absorption

bands as a function of the concentration of donor and acceptor components in

solution. The methods were first described for the association of iodine dissolved in

aromatic hydrocarbons. This procedure is called the Benesi- Hildebrand method [12].

The absorption wavelength of charge-transfer bands, i.e., the charge-transfer

transition energy, is characteristic of the specific type of donor and acceptor entities.

The formation of electron donor-acceptor (EDA) charge-transfer complexes between

two molecules involves the transfer of an electron from the highest occupied

molecular orbital (HOMO) of the donor to the lowest unoccupied molecular orbital

(LUMO) of the acceptor [13]. The donor may donate an unshared pair (an n-donor) or

a pair of electrons in a TT orbital of a double bond or aromatic system (a n donor). One

of the tests for the presence of an EDA complex is the electronic spectrum. These

complexes generally exhibit a spectrum (called a charge transfer spectrum), which is

not the same as the sum of the spectra of the two individual molecules. Because the

first excited state of the complex is relatively close in energy to the ground state, there

is usually a peak in the visible or near-UV region, and electron donor-acceptor

complexes are often coloured. These charge transfer complexes have unique

absorption bands in the ultraviolet-visible region. The compositions of donor-acceptor

complexes could not be isolated at ordinary temperatures in pure state since they are

mostly unstable. However, they exist only in solutions in equilibrium with their

components. The rates of formafion of complexes in solution are generally so rapid

that kinetic studies of the reactions cannot be made, at least by ordinary procedures.

The values of heat of interaction are generally smaller than the forces of coordination

and are much feebler than those established in the formation of covalent bonds [14].

That is, the degree to which electron transfer from the donor component to the

acceptor component takes place is much less than that ordinarily occurs when new

compounds are formed. Aromatic hydrocarbons are rich in K- electron and favor the

formation of charge-transfer complexes with electron deficient molecules. This may

be the due to the intermolecular attractive forces or through dipole-dipole interaction

with electron deficient molecules. Charge-transfer complexes exist in two states: a

ground state and an excited state. In the ground state, the two molecules composing

the complex undergo the normal physical forces expected between two molecules

which are in close proximity to each other. These include London dispersion forces

and any electrostatic interactions, such as between dipole moments. In addition, a

small amount of charge is transferred from the donor to the acceptor. This contributes

some additional binding energy to the complex [15]. The excited state of the complex

occurs when the ground state complex absorbs a photon of light having appropriate

frequency. In the excited state, the electron which had only been slightly shifted

towards the acceptor in the ground state is almost totally transferred. Depending on

the structural features of both the donor and acceptor, the wavelength of light

absorbed may be in the visible range of the electromagnetic spectrum. In many cases,

therefore, charge-transfer complexes are colored substances. Examples of such

charge-transfer complexes include the complex formed between a metal ion and a n

orbital of a double bond or an aromatic system [16-17], the complex formed between

polynitro aromatics, such as picric acid [18] as well as n-orbital-containing molecules,

and complexes of h and Br2 with amines, ketones, aromatics, etc. Phenols and

quinones also form charge-transfer complexes [19].

The ground state of the complex is described by the wave function, *FN, which is the

hybrid of two wave functions, \\i (A, D) and \|/ (A'D"") [20] .The wave function \\i (A,D)

is termed as the no-bond function and represents the wave function of the two

molecules in close proximity to each other but with no charge-transfer between them.

4 (A, D) may include, however, the normal electrical interactions between molecules.

Consequently, the ground state wave function for a weak complex can be described

as:

^N = a»F(A,D) + b4^(A"D^)

Here a » b . The wave function FCA'D" ) is called the dative function and represents

two molecules held together by total transfer of an electron from the donor, D, to the

acceptor, A. The excited state of the complex can then be described by:

^E = b* 4 (A"D*) - a* ^ (A,D)

Hereb*»a*.

Charge transfer complexes may conveniently be classified according to the types of

orbitals in the donor and acceptor molecules which are undergoing the interaction.

Donor and acceptor molecules may each be divided into three classes [21], as shown

in the Table-1 below. The v-acceptors refer to metal atoms possessing a low-lying

vacant valence-orbital. Hypothetically, there are nine possible types of complexes.

However in practice, the n-donors do not form complexes with metal ions but form

covalent bonds instead [22].

Table 1: Charge-Transfer Complexes: Donor and Acceptor Types

Donor type

a

n

71

Acceptor type

n

a

71

Examples

R-X, cyclopropane etc.

R2O, R3N, pyridine, dioxane etc.

Aromatic and unsaturated hydrocarbons.

especially if containing electron releasing

substituent (hexamethyl-benzene or phenols)

etc.

Examples

Ag+, certain organometalics etc.

I2, Br2, ICI etc.

Aromatic and unsaturated hydrocarbons.

especially if containing electron withdrawing

substituents (TCNE, halogenated quinones)

etc.

-The occurrence of a charge-transfer usually requires some amount of overlap between

the molecular orbitals of the donor and acceptor. Normally, the interaction is between

the highest occupied molecular orbital (HOMO) of the donor with the lowest

unoccupied molecular orbital (LUMO) of the acceptor. This overlap principle is

certainly true for intermolecular complexes. However, for intramolecular complexes,

examples are also known of indirect, through-bond interactions, besides the direct,

through-space interactions [23-24]. The amount of orbital overlap between the donor

and acceptor plays a critical role in the magnitude of the charge-transfer interaction

observed. Constraints resulting from steric hindrances are major factors in this regard.

For instance, the binding energies for the charge-transfer complexes between phenol

and hydroquinone with p-benzoquinone were found to be larger than those for anisol

and hydroquinone dimethyl ether with p-benzoquinone [25]. This was attributed to the

steric influence of the methyl groups. For highly substituted molecules, steric

hindrance may well prevent the close approach necessary in order for charge-transfer

interactions to take place. For most complexes, however, the charge transfer

contribution to the complex stability appears to be minor or negligible, and these

complexes may be considered to be non-covalently bound, despite their possession of

intense electronic absorption spectra. The electron donating power of a donor

molecule is measured by its ionization potential, which is the energy required to

remove an electron from the highest occupied molecular orbital. The electron

accepting power of the electron acceptor is determined by its electron affinity, which

is the energy released when filling the lowest unoccupied molecular orbital. As a

result of Mulliken's and Deware's theory [26-27], there has been a great stimulus to

the developments in the study of the charge transfer (CT) complexes. A brief history

of their theory are presented here.

1.2. Mulliken's Theory

The theory of charge transfer CT) complex was proposed by R. S. Mulliken who won

the 1966 Chemistry novel prize [28-30]. Mulliken consider the complex as a hybrid

resonating between the non-polar structure and polar one resulting, from the transfer

of one electron [31]. According to Mulliken, when two molecules which form loose

complexes are mixed together, one of them acts as an electron-donor while other as an

electron-acceptor and they give rise to a donor-acceptor complex, neglecting other

types of interactions.

Mulliken's charge- transfer (hereafter abbreviated CT) has been widely and

successfully applied to the interpretation of various properties of electron donor-

acceptor (hereafter abbreviated to EDA) complexes such as their stabilities,

geometrical structures and spectroscopic, electric and magnetic properties [32-43].

The existence of CT states characteristic of EDA complexes has been demonstrated

by the measurement of dipole moments and absorption and electron spin resonance

spectra of their excited states. The theory of donor acceptor complexes and their

spectra as presented by MuUiken is a vapor state theory, except for the omission of the

London dispersion attraction terms. This theory essentially valid for solution in inert

solvents. Mulliken had been previously involved in the interpretation of band spectra

of diatomic molecules and decided to treat the problem of molecular complexes in

similar manner. As a result his theoretical treatment of the complex was very similar

to the valence bond treatment of diatomic molecules

What Mulliken did in his treatment of molecular complexes was to consider each

member of the complex as an 'atom' and the overall pair as a diatomic molecule of

sorts. He then wrote a very simple diatomic-like bond wave function

4^(DA) = a4^o(D,A) +b4'i(D^A~) (1)

Where, D refers to the donor and A the acceptor. To be more general, the equation

should include terms of higher ionic character (which Mulliken later did) but the

theory and its ramifications are best explained using the simple form above. Equation

(1) states simply that the complex may be considered as a mixture of two states, a

non-ionic pair \|/o(D,A) which, in addition to describing the nonbonding pair, includes

modifying terms due to polarization effects, and an ionic pair v|/| (D' A") which

describes a weak covalent bond between the pair and also includes some modifying

terms. That this can be done is stipulated within the rules of quantum mechanics since

we are regarding *Po and NKi as our basis functions with which we are to describe our

system. The beauty of this assumption in describing the wave function of the complex

is that from it all the properties of charge-transfer spectra can be derived even though

we have no idea as to the form of *Po and 4*1. Although it has been stated above that

*Fo and *Fi are unknown, this is true in the absolute sense only. The physical nature

can be inferred by constructing these states from the individual wave functions

describing the donor and acceptor ^{D) and T(A) respectively. One can therefore

express *Fo as a product wave function

4 0 = A [^ (D) H'(A)] (2)

Where, each component wave function describes the donor or acceptor with its full

component of electrons without actually exchanging electrons. However, despite their

being separate entities, the two electronic distributions on the donor and acceptor are

mutually influenced by one another, and subject to exchange repulsion forces,

dispersion, and classical electrostatic forces. It follows then that *F(D) is not identical

to the wave function which describes the donor in uucuo, but possesses some

modification due to the nearby molecule; these modifications naturally would be more

extreme as the strength of the interaction increases. The dative basis function may be

likewise written as a product of two ionic wave functions ¥(0"^), Y(A"). In this case,

however, an electron is exchanged from the donor to acceptor and may be regarded as

being delocalized over the entire donor-acceptor moiety. It is in this sense that the

ionic form of the complex's basis is considered to contain some covalent character. It

must be realized that for two molecules to exchange an electron, they must approach

one another to the extent that orbitals localized on the individual molecules overlap.

From equation (1) the ground state wave function can be written as sum of the

functions 4 0 and 4 1. The excited state is then

Pv (DA) = c4 o (D,A) + d^'i (D^ A") (3)

By applying the condition of orthonormality, ^'N and 4'v provides information

regarding the nature of the charge transfer transition and gives the three relations in

terms of the coefficients and integral over ^o and Ti:

(4 N I 'N ) = 1 = a + b^ + ab (4'i|4'o) (4)

(4/v|4^v) = 1 = c + d + cd(4^i|To) (5)

(4>v|4 v) = 0 = ac + bd + ac (4^i|To) + bd (4^i|To) (6)

Which imply that c = - b and d = a is during the excitation process the character of the

complex changes. If in the ground state the complex is predominantly non-ionic (a »

b), in the excited state it becomes ionic. The excitation process has associated with it a

transfer of an electron from the donor to the acceptor. If we want to the represents the

interaction between the donor-acceptor pair in terms of a potential energy profile for

the ground and excited states, which can be shown in fig. 1 that at infinite separation

dissociation products D and A are obtained in the ground state and ionic forms are the

dissociation product of the excited state.

w A"D*

Figure 1. Potential surface representation of the charge transfer transition and its dependence

upon ID and t A.

This information follows from the orthogonality relations above. At infinite

separation the difference in energy between the ground and excited states is ID - EA,

where ID is the ionization potential of the donor, and EA is the electron affinity of the

acceptor. Therefore the transition energy is approximated by hv = ID + EA + A which

is very nearly identical to the equation proposed by Rabinowitch for Electron Affinity

Spectra; the difference being that the signs are reversed. At the time Mulliken

reported his theory of charge transfer, the modern theory of quantum mechanics,

including the method of molecular orbitals [44] was well developed. With the

molecular orbital theory Mulliken was able to approximate A by more sophisticated

means. Since he was able to write the wave function in a meaningful (albeit not

calculable) form, an application of the variational principle of quantum mechanics

was possible. By substitufing the approximate wave function Tn (DA) = a4^o (D,A)

+ b Ti in to the Schrodinger equation and requiring that 5W/ca and 6W/ob (where

10

W designates energy) be zero, a secular determinant of the form was obtained, which

when solved, gives a quadratic equation in W with two roots as solutions.

Wo - W HoiWSoi

= 0

H o i - W S o i VViW

One of these roots corresponding to the ground state energy

(Hoi - SoiWo)' WN = Wo -

(Wi - Wo)

(')

(8)

Where the another root obtain the excited state

Wv = Wx + (Hoi - SoiWi )-

(9)

In the above equations Ho is the expaction value <4'o|H|^i> and Soi is the overlap

integral <4'o|4'i>.

To obtain hv in the form of equation one starts from its definition

Which, following substitution from equations (8 ) and (9), yields

1>VCT = \ \ , Wo ( Hoi - SoAVo )• - ( H O I - SoiWi )"

(Wi - \Ao) (10)

The form of equation above is obtained if the expressions for W are written as

energies at infinhe separation plus a correction term. For example, the ground-state

energy is expressed as Wo = Woo - G and from this the excited-state energy can be

written as Wi = Woo + IQ - EA - G'. Substitution in to equation (10) gives

(Hoi - SoiWi)-

(11) livci = Ijj + E A - G ' + G -

(W, Wo)

11

Which is the desired form expressing the functional dependencies upon ID and EA.

Mulliken an offered an explanation as to the possible nature of the terms G and G'.

He suggested that G was a non-bonding stabilization term which largely consisted of

contributions from London dispersion forces. G' was a term which had to include

major contributions from coulombic and exchange forces due to the nature of the

excited state. What is important is that not only has this valence bond approach agreed

with the diatomic molecule analogue (and could be directly applied to the ionic

molecule problem), but it successfully predicts experimental trends in the variation of

hv, as a function of the ionization potential and electron affinity of the donor and

acceptor respectively. In the variational treatment of the charge-transfer problem the

coefficients a and b were the parameters with respect to which the energy was

minimized. The results of this procedure can be used to obtain a ratio of the

coefficients, b/a, by the application of second-order perturbation theory [45]. In the

application of the perturbation methods to the complexes the dative wave function, is

considered as a perturbation to the non-bonding pair, which ultimately gives for b/a:

(Hoi - SoiWo) b;) = - (12)

(Wi - Wo)

The importance of this equation has been discussed in detail by Mulliken [46].

1.3. Dewar Theory

The factors governing the behavior, stability, etc. of CT complexes have not been

clearly established. Among the different theories proposed, one of the best known was

derived by Mulliken, [47] who considers the complex as a hybrid resonating between

a non-polar structure and a polar one, resulting from the transfer of one electron from

donor (D) to acceptor (A) .In general, the charge transfer complexation occurs as an

ionic band in the simple ion-radical pair interaction [48]. The appearance of a new

band spectrum of this type of complex is ascribed to a transition from a ground state

which is mostly (I) mixed with little (II) to an excited state which is mostly (II) mixed

with little (I) this type of CT transition termed as CT complexes.

(I) DA (II) D^A"

12

The above problem has been modified in terms of molecular orbital treatment by

Dewar and Lepley [49]. In Dewar's theory the transfer is supposed to take place

between the highest occupied orbital of the donor and the lowest empty one of the

acceptor.

Here the complex DA is represented as a 7t-complex formed by interaction of the 71-

orbitals of D and A; science the interaction is known to be small it can be

conveniently treated by using perturbation theory [50]. Consider the orbitals of D and

A (Fig. 2). Interactions between the filled bonding orbitals of and of D and A lead to

no change in their total energy and to no net transfer of charge between D and A.

Interactions of the filled orbitals of D with the empty anti-bonding orbitals of A

depress the former and raise the latter, leading to a net stabilization with a

simultaneous transfer of negative charge from D to A; interactions of the filled

orbitals of A with the empty orbitals of D likewise lead to stabilization with a net

charge transfer in the opposite direction. These interactions are inversely proportional

to the difference in energy between the interacting orbitals. In complexes of this kind

one component is normally a molecule of donor type (i.e., with filled orbitals of

relatively high energy), the other an acceptor (i.e., with empty orbitals of relatively

low energy).The main interaction is therefore between the filled orbitals of the donor

and the empty orbitals of the acceptor, as indicated in Fig. 2; this leads to a net

transfer of negative charge from the donor D to the acceptor A. The heats of

formation of complexes of this kind are at least an order of magnitude less than their

lowest transition energies; this suggests

that the changes in energy of the orbitals in forming the complex are small compared

with the spacing between the filled (bonding) and empty (anti-bonding) orbitals. The

energies of the orbitals in the complex should therefore be little different from those

in the separate components; all the possible transitions observed in A and B should

therefore appear in the spectrum of the complex AB, and this is commonly the case.

Transitions of this type are described as locally excited [51]. There should also be

charge transfer transitions of electrons from a filled orbital of D into an empty orbital

of A, and from a filled orbital of A into an empty orbital of D. Fig.2 indicates that

transitions of the former kind may occur at lower energies than the locally excited

transitions and so lead to the appearance of new absorption bands at lower

frequencies. This accounts for the new bands commonly observed in the spectra of

such complexes and responsible for their color.

13

This treatment leads to conclusions similar to those given by the valence bond

approach, but it seems preferable for two reasons. First, there are cases when more

than one new charge transfer band appears in the complex DA; this can be explained

at once in terms of the molecular orbital approach since there should be bands

corresponding to transitions between any of the occupied orbitals of D and empty

orbitals of A. Secondly, the term "charge transfer complex" is misleading in that very

little charge is transferred in the ground states of such complexes and in that an

appreciable part of their stability may be due to back-coordination involving

interactions between the filled orbitals of the acceptor and the empty orbitals of the

donor, The term "71-complex" seems preferable for compounds of this type. If the

interactions between donor and acceptor are small, the transition energy AEo for the

first charge transfer band should by either treatment by

AEo = 1 D - E A + Constant (13)

Where ID is the ionization potential of D (equal to the energy of the highest, occupied

MO in a naive molecular-orbital approach) and EA is the electron affinity of A

(likewise equal to the energy of its lowest unoccupied orbital). If then the acceptor is

kept constant, AEo should vary linearly with the ionization potential of the donor; this

relation has been observed in a number of cases.3 In the molecular orbital approach

equation (13 ) is replaced by the more general relation which is given below

AEij = Di - Aj + Constant (14)

Where AEy is the transition energy for CT band involving the field orbital i of D

(energy Di) and the empty orbital j of A (energy Aj). This is equivalent to equation

(14) in the case of the first charge transfer band, for the ionization potential of the

donor should be equal to the energy of its highest occupied molecular orbital.

If equation (14) is valid, the energies of the charge transfer transitions should be

predictable from simple molecular orbital theory. Thus, the energies of the first charge

transfer transitions for a variety of donors with a given acceptor should be a linear

function of the energies of the highest occupied orbitals of the donors. If the

arguments given above are correct, and if equation (14) were calibrated by using data

for a variety of hydrocarbons, the charge transfer spectra of compounds would

immediately provide estimates of the energies of their highest occupied molecular

orbitals.

14

anti-bondmg MO's

A\

locally excited transitions

cliarge transfer transitions

A\ locally excited transitions

bonding MO's

^ n u u

^

ft

VJU

D A

Figure 2. Orbital energies and transitions in a molecular complex formed by a donor

and acceptor.

1.4. Ligand to metal and metal to ligand charge transfer complexes

Ligands possess a, a*, TT, TI*, and nonbonding (n) molecular orbitals. If the ligand

molecular orbitals are full, charge transfer may occur from the ligand molecular

orbitals to the empty or partially filled metal d-orbitals. The absorptions that arise

from this process are called ligand-to-metal charge-transfer bands (LMCT) (Figure 3)

[52]. LMCT transitions result in intense bands. Forbidden d-d transitions may also

take place giving rise to weak absorptions. Ligand to metal charge transfer results in

the reduction of the metal.

If the metal is in a low oxidation state (electron rich) and the ligand possesses low-

lying empty orbitals (for example CO or CN-) then a metal-to-ligand charge transfer

(MLCT) transition may occur. LMCT transitions are common for coordination

compounds having 7i-acceptor ligands. Upon the absorption of light, electrons in the

15

metal orbitals are excited to the ligand n* orbitals [52] Figure 4 illustrates the metal to

ligaiid charge transfer in a d' octahedral complex. MLCT Iraiisitions result in inlcnsc

bands. Forbidden d - d transitions may also occur. This transition results in the

oxidation of the metal.

-tii_-L-L-l d^ uncoordinated metal

eg

fi JJl t2g Ligand to Metal Charge

Transfer (LMCT)

/ ligand sigma orbitals

iiiiiliiii/ octahedral complex

Figure 3. Ligand to metal charge transfer (LMCT) involving an octahedral d*

complex [53].

i_i_i_± d uncoordinated metal

Ugand Tc* orbitals

eg

Metal to Ligand Ckarge Transfer (MLCT)

octahedral complex

Figure 4. Metal to ligand charge transfer (MLCT) involving an octahedral d'

complex [53].

Carbon monoxide bonds to transition metals using "synergistic n* back-bonding."

The bonding has three components, giving rise lo a partial triple bond. A sigma hcMid

arises from overlap of nonbonding sp-hybridized electron pair on carbon with a blend

of d-. S-. and p-orbita!s on the metal. A pair of TT bonds arises from overlap of tilled d-

16

orbitals on the metal with a pair of TI anti-bonding orbitals projecting from the carbon

of the CO. The latter kind of binding requires that the metal have d-electrons, and that

the metal is in a relatively low oxidation state (<+2) which makes the back donation

process favorable. As electrons from the metal fill the 7i-anti-bonding orbital of CO,

they weaken the carbon-oxygen bond compared with free carbon monoxide, while the

metal-carbon bond is strengthened. Because of the multiple bond character of the M-

CO linkage, the distance between the metal and carbon is relatively short, often < 1.8

A. about 0.2 A. shorter than a metal-alkyl bond. Several canonical forms can be drawn

to describe (he approximate metal carbonyl bonding modes.

M~^ C^(f< > M=^C^=0 < > M " = C >0~

(Resonance structures of a metal carbonyl, from left to right the contributions of the

right-hand-side canonical forms increase as the back bonding power of M to CO

increases).

CO pi aniibt>nding orbitals (empty)

mcJal l;y orbital (filled)

Figure 5. Charge transfer from metal to ligand through n- back bonding.

Marcus-Hush theory relates kinetic and thermodynamic data for two self-exchange

reactions with data for the cross-reaction between the two self-exchange partners.

This theory determines whether an outer sphere mechanism has taken place. This

theory is illustrated in the following reactions

Self exchange 1: [ ML6 ] ^ + [MU] ^^ ^ [ MLg ] ^^ + [ MLa ] ^ AG" = 0

Self exchange 2: [ MLs ] ^' + [ MU] " [ MLs ] ^ + [ MU ] ^' AG"= 0

Cross Reaction: [ ML6] " + [ MU] * -> [ ML6 ] ^ + [ ML^ ] ' '

The Gibbs free energy of activation A G ' is represented by the following equation:

AG^ - AwG + AoCJ + i\oQ^ + Rl in (k'T/hZ)

17

Where:

T Icmperulure in K, R molar gas CDnsianl. k" Buil/.mann coiislani. li Planck^

constant

Z = etfccti\c trcqiicnc}' collision in solution ~ lO um' mo! s . \u C; " the cncr^}

associated with bringing the reactants together, includes the work done to counter an\

repulsion

AQG^ = energy associated with bond distance changes. As AG^= energy associated u ith

the rearrantiements takinc place in the solvent spheres

In ( k'T / hZ) = accounts for the energy lost in the formation of the encounter complex

The rate constant for the self-exchange is calculated nsing The followini' reaction

K = kZe"^^^'^^^

where K is the transmission coefficient ^1

The Marcus-Hush equation is given by the following expression

k,. = ('k!,k.2K,.f,2)"-

where,

Z is the collision frequency ,T

:''^:nr>n.'l ' M •;:V kii and \G n correspond to self exchange !. k:: and \G 22 corresponc

exchange 2, k|2 and AG 12 correspond to the cross-reaction. K12 = cross reaction

equiiiunum constant, AG 12 = standard Gibbs free energ}' of the rcactiwii.

The following equation is an approxinialc iVom ofihc Marcus-llush cqualion:

logki2 ~ 0.5 logki, + 0.5 logk22 + 0.5 logK^

since

f^l

and

logfs^O

How is the Marcus-Hush equation used to determine if an outer sphere mechanism is

taking place?

values of kn. ]i22 Kj2. and k,2 are obtained experimeniallv k.. .nnd k . are

theoretically values

Ki2is obtained frotn F,.,.ii

If an outer sphere mechanism is taking place the calculated values of ki2 will match or

agree with the cxperimeiilal values. If these values do iioL agree, this would indicate

that another mechanism is taking place [54].

1.5. Spcctrophotomctric Dctcrminnatio of Equilibrium Constant and

Molar Absoptivity

Theoretically. K| (formation constant) can range from zero (implying no complex

formation) to infinity (implying complete conversion of stoichiometric quantities of

the donor and the acceptor into the complex). Generally, complexes with formation

constants less than unity are classified as weak, while for strong complexes K>1. In

many siluations. these complexes are non-iso!ahIe in the pure state, thus ruling out the

possibility of the direct measurement of their molar extinction coefficients at the

charge-transfer maximum and. therefore, the direct evaluation of their formation

constants. In 1949, however, Benesi and Hildebrand [55] reported a graphical method

for the simultaneous evaluation of the formation constant (Kf) and the molar

extinction coefficient (e) of 1:1 EDA (electron donor-acceptor) complexes formed

between some aromatic hydrocarbons (acting as electron donors) and iodine (an

electron acceptor) in inert solvents like carbon tetrachloride and heptanes:

The equilibrium constant Kf (formation constant) for DA association and molar

absorpti\it}' (s) for DA CT absorption can be defined by using the equation

K( D + A '^ [ D - A ] C T

Ihe Benesi-IIildebrand analysis of Kf involves the measurement of the D—A CI

absorbance (Abs) as a function of varied [A] when [A]»[D]. A plot of x = l/[A]o vs

y = [D]o/Abs gives a y-intercept = I/E and slope = (1/Kfe) as defined by the Benesi-

Hildebrand equation:

[D]o/Abs = (1 /[A]o) (1 /K, 8) + I /£ (15)

Where, [D]o = total of donor (fixed), Abs = CT absorption of DA complex at

wavelength X, [A]o - total cone, of acceptor (varied), Kf = equilibrium constant for

DA complex formation

8 = molar absorptivity of DA complex at X.

19

1 he derivation is as follows:

Kf = [DA]/[D][A] (16)

define

[D]o = total D (uncomplexed and complexed) = [D] + [DA]

|A],i ~ total A (uncoinplexed and complexed) ^ [A] ^ jD \ |

substitution into equation (16) gives

K, = [DA]/JfDlo-fDA]} ![A]o-[DA!; (H)

i f [Aln » [nio, then f lAlo - rnA]>= fAln. so

K, = [DA] /{[D]o - [DA]} [A]o (18)

rearranging gives.

[nA] = K,rA]ornio/(i +KirA]o) (i9i

CT absorbance by DA according to Beer's law is:

Abs= cUDA] - cl(Ki[A]orDlo)/(n KrjAjo) (20)

Where, 1 = sample path length in cm (typical 1 cm)

The value of Absorbance [A] must increase as the concentraiion c-rj"Di incrca.-,e.

which has been sliovvn in rigure (6).

zu

[*] • >

Figure 6. The concentration of of complex as a function of varied [A] for a fixed total

concentration [D]o, Region 1: DA is approximately a linear function of [A]o. Region

3: Saturation has been reached, and [DA] is constant and equal to [D]o.

when 1 ~ 1, rearranging equation (20) gives the B-H equation

[D] c)/Abs = (1 / [A]()) (1 /eKf) +1/8 (21)

The equation (21) known as Benesi-Hildebrand equation. The values of formation

constant (Kf) and the molar absorptivity (e) can be obtained by using the method of

Benesi- Hildebrand [77].

The Benesi-Hildebrand analysis starts from the assumption that only one equilibrium

esist in the solution.

1.6. Hydrogen Bonding

Hydrogen Bonding is a complex process that plays an important role in biophysics.

Hydrogen bonds, although much weaker than many other forms of chemical bonds,

have major effects on the structure and function of the proteins, polypeptide chains,

and strands of DNA that are essential to biological systems. Hydrogen bonding will

only occur under very specific conditions and only between particles with very

particular properties. Hydrogen bonds form between proton donor molecules that

contain a hydrogen atom covalently bonded to one or more electronegative atoms and

a proton acceptor molecule. The electronegative atoms in the proton donor molecule

attracts the electron of the bonded hydrogen to such an extent that the attached

hydrogen gains a net positive charge that allows it to interact with the more negatively

charged proton acceptor molecule. In fact, there are several other factors that must be

21

taken into account to fully explain the nature of hydrogen bonds (H bonds), but a

simple model of a H- bond is a proton being shared by two different molecules [56].

A more complicated model of hydrogen bonds involves contributions from several

different interactions. The most prevalent of these interactions are: electrostatic,

delocalization, dispersive, and repulsive interactions. As previously noted, the proton

donor molecule consists of a hydrogen atom bonded to one or more electronegative

atoms which attract the electron on the attached hydrogen, essentially reducing the

negative charge distribution in the region around the hydrogen proton. This

delocalization of charge produces a columbic attraction between the donor molecule's

hydrogen atom and the electrons of the proton acceptor molecule. Furthermore, the

motion of electrons in both the donor and acceptor molecules behaves much like

fluctuating electric dipoles resulting in additional attractive interactions. Finally,

overlaps in the electron clouds of the proton -donor and -acceptor molecules add a

repulsive interaction to the nature of H-bonds [78]. There are several different

methods which attempt to model the complex interactions that produce hydrogen

bonds. Among the most successful are the molecular orbits (MO) method and the

Hartree-Fock (HF) (or self consistent field (SCF) method) [57-58].

Hydrogen bonds often provide the strongest intermolecular forces between molecules

in organic molecular crystals and hence often dictate the preferred packing

arrangement. The general principles underlying hydrogen- bond formation are

reasonably well understood and the structures of hydrogen-bonded crystals can often

be rationalized in preferred combinations of hydrogen-bond donors and acceptors

(Etter, 1990; Etter, McDonald & Bernstein, 1990; Etter & Reutzel, 1991) [59-61]. In

general, the strongest hydrogen-bond donors pair off with the strongest hydrogen-

bond acceptors. Similar pairing processes are repeated until all the hydrogen-bond

donors and acceptors have been utilized. However, when a system contains excess

donors or acceptors, at least two hydrogen-bonding strategies are available to

accommodate the mismatch (Hanton, Hunter & Purvis, 1992) [62]: (i) change in

hybridization or (ii) the formation of hydrogen bonds involving the n system of an

aromatic group as the acceptor. Several examples of the formation of intermolecular

X~H...;i (arene) bonds for X = O or N have been observed where there is a deficiency

of sterically accessible acceptor sites of the conventional type (Hanton, Hunter &

Purvis, 1992; Atwood, Hamada, Robinson, Orr & Vincent, 1991; Rzepa, Webb,

Slawin & Williams, 1991) [62-64]. The strongest bonds are formed between the most

22

electronegative atoms such as fluorine nitrogen and oxygen whicii interact with atoms

having electronegativity greater than that of hydrogen. The weakest of hydrogen

bonds are formed by the acidic protons of CH groups, as in chloroform acetylene and

by olefmic and aromatic 7r-electrons. The stability of the formation of charge transfer

complex depends upon the strength of hydrogen bond that is the weaker the hydrogen

bond, the shorter the lifetime of the complex formed. The intermolecular hydrogen

bonding between the organic compounds are formed very weak in comparison of

covalent bonding i.e., covalent bonding typical near about twenty times stronger than

that of intermolecular hydrogen bonding. These bonds mainly are formed between the

molecules (i.e., intermolecular hydrogen bonding) or within different parts of a single

molecule (i.e., intramolecular hydrogen bonding) [65]. The hydrogen bond is a very

strong fixed dipole-dipole van der waals force, but weaker than covalent and ionic

bonds. The hydrogen bond is somewhere between a covalent bond and an electrostatic

intermolecular attraction.

Mainly two types hydrogen bonding are involves in organic compounds are

intermolecular and intramolecular hydrogen bonding. The intermolecular hydrogen

bonding takes place between the molecules and intramolecular hydrogen bonding

within the molecule. Distinction between inter-and intramolecular hydrogen bonding

can be thus made by effect of dilution. Intramolecular hydrogen bonds remain

unaffected and as a result the absorption band also remains unaffected. Intermolecular

hydrogen bonds are however, broken on dilution and as a result there is a decrease in

the bonded O—H absorption and an increase in or the appearance of free O—H

absorption. Hydrogen bonding in chelates and enols is very strong as shown in fig.7.

Since these bonds are not easily broken on dilution by an inert solvent, free O—H

stretching may not be seen at low concentration.

enol oi

(CH3COCH2COCH3) chelate oi

(methyl salicylate)

Figure 7. Intramolecular hydrogen bonding in chelate of methyl salicylate and enol of

CH3COCH2COCH3.

23

Hydrogen bonding (H-bonding) has recently been defined by lUPAC as "an attractive

interaction between a hydrogen atom from a molecule or a molecular fragment X-H

in which X is more electronegative than H, and an atom or a group of atoms in the

same or a different molecule, in which there is evidence of bond formation". In most

cases, the strength of an H-bond increases with the increase of the electronegativity

value of the acceptor atom (Pauling, 1960) [66]. This is exactly the case for oxygen

and nitrogen atoms. The H-bonds formed between them and the NH and OH groups

are usually strong, which play essential roles in studies in supramolecular, crystal

engineering, materials, and life sciences (Scheiner, 1997; Jeffrey, 1997) [67-68]. As a

result of their growing applications in supramolecular chemistry and crystal

engineering, in the past two decades, the critical assessment of the weaker H-bonds

has also become an important topic (Desiraju & Steiner, 2001) [69]. In this context,

organic halogen and sulfur atoms, C-X (X = F, CI, Br, I, S), have all been

demonstrated to be weak H-bonding acceptors (Dunitz & Taylor, 1997) [70], although

their electronegativities (Pauling scale: 3.98, 3.16, 2.96, 2.66, and 2.58, respectively)

are all higher than that of hydrogen (2.20). Indeed, over years it has been accepted

that organic fluorine "hardly ever accepts hydrogen bonds (Dunitz, 2004) [71],"

presumably due to its low polarizability and tightly contracted lone pairs. For other

organic heteroatoms, the increased van der Waals radius and decreased

electronegativities may also weaken their capacity of forming the intramolecular

electrostatic interaction, i.e., H-bonding, with the amide hydrogen and lose the

competition with the amide oxygen of another molecule which forms the

intermolecular N-H---0=C H-bonding. In contrast, the halogen anions are capable of

forming strong intermolecular H-bonding with NH, OH or even CH protons (Harrell

& McDaniel, 1964; Simonov et al., 1996; Del Bene & Jordan, 2001) [72-74]. This

chapter summarizes recent progresses in the assessment of the weak intramolecular

six- and five-membered H-bonding patterns formed by aromatic amides bearing the

above five atoms. Theoretical investigations show that similar intermolecular H-

bonding patterns can be formed by fluorine in DNA or RNA base analogues (Frey et

al., 2006; Roller et al., 2010; Manjunatha et al., 2010) [75-77], although they are

difficult to be confirmed in solution experimentally. The crystal structures of many

organic halogen or sulfur (ether) compounds exhibit such intermolecular short

contacts, which may be mainly driven by the intrinsic preference of these atoms in

forming the H-bonding or formed due to the assistance of the intermolecular stacking

24

and van der Waals force (Toth et al., 2007) [78] and other intra- and intermolecular

interactions. Due to tlie increased conformational tlexibility of the backbones and the

decreased acidity of the amide proton, the H-bonding in aliphatic amide derivatives is

expected to be even weaker. However, five-membered intramolecular N-H---F (F:

Hughes & Small, 1972; O'Hagan et al., 2006) [79-80], N-H---C1 (De Sousa et al.,

2007; Kalyanaraman et al., 1978) [81-82] and N-H--I (Savinkina et al., 2008) [83] H-

bonding patterns have been observed in aliphatic amides. To the best of our

knowledge, the six-membered one is not available yet in simple organic molecules.

One consideration for exploiting the intramolecular N-H--X (X = F, CI, Br, I, S) H-

bonding of the aromatic amides is that the new patterns may find applications in

designing new preorganized building blocks for crystal and supramolecular

engineering (Biradha, 2003; Desiraju, 2005) [84-85]. Furthermore, new H-bonding

motifs may also be useful in building foldamers (Zhu et al., 2011; Zhao & Li, 2010;

Saraogi & Hamilton, 2009; Li et al„ 2008; Li et al., 2006; Hue, 2004; Sanford &

Gong, 2003) [86-92], the artificial secondary structures, and for designing

biologically or medicinally useful structures (Tew et al., 2010; Li et al., 2008;

Bautista et al., 2007) [93-95]. For doing this, the more competitive intermolecular

N-H"-0=C H-bonding of the amide unit has to be suppressed. There are two

approaches for realizing this purpose. The first one concerns the introduction of a

strong intramolecular H-bond to "lock" the amide proton. The second one is to

introduce one or more bulky groups to impede the contact of the amides. In these

ways, the very weak intramolecular N-H---I hydrogen bonding can be observed.

There are several techniques for investigating the formation of the weak

intramolecular H-bonding. The NMR spectroscopy is promising for studies in

solution (Manjunatha et al., 2010) [96], and the infrared spectroscopy can be used to

detect samples in both the solution and solid state (Legon, 1990) [97], while the

computational modeling can provide useful information about the effects of discrete

factors on the stability of the H-bonds (Dunitz, 2004; Liu et al., 2009) [98-99], which

are particularly valuable when experimental evidences are not available. In view of

the feature of this book, we will focus on the investigations by the X-ray

crystallography. The crystal structure of an aromatic amide molecule is affected by

many factors, including the stacking pattern, van der Waals force, intra- and

intermolecular hydrogen and halogen bonding, and shape matching of the molecule.

25

The entrapped solvent molecules, particularly those containing heteroatoms, may also

play an important role because they are able to form hydrogen or halogen bonding

with the molecule and thus affect the stacking pattern to suppress or promote the

formation of the intramolecular H-bonding. Concerning the criterion for the formation

of the weak intramolecular H-bonding, we simply check the distance between the

heteroatom and the amide hydrogen in the crystal structure. If it is shorter than the

sum of the radius of the two atoms, we consider that an H-bonding is formed

(Desiraju & Steiner, 2001) [100]. Although in X-ray structures the proton/hydrogen is

not located accurately and may bend toward or away from the acceptor, for clarity we

simply use the reported distances between the two concerned atoms as the criteria.

Fluorine atom has the highest electronegativity. In 1996, Howard et al. carried out a

review on the short F--H contacts from all of the organofluorine compounds

deposited in the Cambridge Structural Database System (CSDS) and concluded that

organic fluorine is at best only a weak H-bonding acceptor (Howard et al., 2006)

[101]. In 1997, Dunitz and Taylor also executed an intensive search of the CSDS and

confirmed that organic fluorine accepts hydrogen bonds only in the absence of a

better acceptor (Dunitz & Taylor, 1997) [102]. They also examined the evidence for

H-bonding to organic fluorine in protein-ligand complexes and found that it is

unconvincing. They thus proposed that, due to its low polarizability and tightly

contracted lone pairs, organic fluorine does not compete with stronger H-bond

acceptors such as oxygen or nitrogen, and only when other better acceptor atoms are

sterically hindered that the 0-H---F or N-H---F H-bonding can be formed (Barbarich

etal., 1999) [103].

26

OMe

O7N

Figure 9. Intramolecular N-H---F hydrogen bonding in aromatic amide derivatives.

In 1982, Kato et al. reported the crystal structure of 2-fluorobenzamide (Kato &

Sakurai, 1982) [104]. Although the positions of hydrogen atoms were not determined,

the N---F distance is 2.80 A, which corresponded to an NH---F distance of 2.15 A by

molecular modeling. Clearly, an intramolecular six-membered N-H---F hydrogen

bond exists in the crystal. In 2003, Li et al. found that 2-fluorobenzamide derivatives

might promote the stability of hydrazidebased quadruply hydrogen-bonded

heterodimers by forming six-membered Intramolecular N-H---F hydrogen bonding

(Zhao et al, 2003) [105]. A number of model compounds were then designed and

prepared (Li et al., 2005) [106]. The structures of compounds (a)-(c), which bear one

triphenylmethyl or two nitro groups to increase their cystallinity (Corbin et al, 2003;

Yin et al., 2003) [107-108], were obtained (Figure 9). All the three compounds adopt

a well-defmed planar conformation rigidified by the intramolecular N-H---F H-bonds.

The F---H (amide) distance of compound (a) is 2.23 A, and the N-H---F angle is 106°.

27

The fluorine atoms of both (b) and (c) are located to the proximity of the amide

hydrogen due to the formation of the three centered H-bonds, which is common for

similar alkoxyl-substituted aromatic amide (Gong, 2001) [109]. The F---H (amide)

distance of the six- and five-membered H-bonds is 1.94 and 2.18 A in (b), and 1.97

and 2.18 A in (c), respectively. The corresponding F---H-N angle is 136 and 108° for

(b), and 136 and 111° for (c). All these values fall into the range of the criterion for

the judgment of a F---H-N H-bond the F---HN distance < 2.3 A and the N-H---F angle

> 90° proposed by Dunitz and Taylor [110] The NH---F distance of the amino group

of (a) is 2.39 A, which is larger than that of the amide, also reflecting the preference

of the amide proton to form the intramolecular hydrogen bond. 'H N M R experiments

also support that the five- and six-membered and three-center H-bonds are formed in

solution. Recently, the crystal structures of more N-aryl 2-fluorobenzamides have

been reported, most of which display the six-membered N-H---F H-bonding motif

We can also display the Intramolecular hydrogen bonding with other electronegative

atoms like F, CI, Br, I, and S etc.

1.7. Interatomic interaction in charge transfer complexes

Understanding strong and weak interatomic interactions enables the design and

manipulation of molecular systems whose physical properties depend on crystal

packing. The importance of such a study is shown by many examples, such as organic

charge-transfer (CT) complexes which are built by co-crystallization of organic planar

electron donor (D) and acceptor (A) molecules, often aromatic rings. The properties

of these complexes depend on their crystal packing which drives the interactions

between D and A to control the charge transfer. Most studies concern 1:1 charge-

transfer complexes in which D and A form segregated (...-D-D-D- -A-A-A-...) or

mixed (...-D-A-D-...) stacks. The former generally present a high electric conductivity

in the direction of stacking [111]. The latter are usually insulators or semiconductors

at ambient conditions. Some of them undergo an unusual phase transition, called a

neutral-ionic phase transition, related to the variation of the partial degree of charge

transfer [112-114]. This phase transition is accompanied by structural modifications

of the stacking, most often with symmetry breaking, with the formation of D-A pairs.

We have recently shown the possibility of characterizing the two microscopic control

parameters (dimerization and charge transfer) via experimental charge-density studies

[114].

28

By applying temperature, pressure and light excitation charge transfer may also be

induced in 2:1 molecular complexes, but no structural evidence has been published

yet. Previous studies concerned pressure evolution of the ionicity of the D and A

molecules from spectroscopic studies: it has been proposed that ionicity increases

with pressure [115-116], and a non-uniform charge distribution between two moieties

of the D-A-D trimer was deduced from electron-molecular vibration coupling [117-

118]. Moreover, this non-uniform charge distribution has also been observed on A

sites, suggesting that coupled trimers may behave cooperatively [119]. The crystal

structure of the 2:1 charge-transfer complex of tetrathiafulvalene [2,2'-bis(l,3-

dithiolylidene)] and bromanil (tetrabromo-1,4-benzoquinone) [(TTF)(2)-BA,

(C(6)H(4)S(4))(2)-C(6)Br(4)0(2)] has been determined by X-ray diffraction at room

temperature, 100 and 25 K. No structural phase transition occurs in the temperature

range studied. The crystal is made of TTF-BA-TTF sandwich trimers. A charge-

transfer estimation between donor and acceptor molecules is proposed in comparison

to the molecular geometries of TTF-BA and TTF and BA isolated molecules.

Displacement parameters of the molecules have been modeled with the TLS

formalism. Donor (TTF) and acceptor (BA) molecules are represented below (fig. 19).

Rr fir

During the later part of the 1940's the interest in complexes formed by donor and

acceptor molecules was stimulated by quantitative spectroscopic work dealing first

with solutions of iodine in benzene, later with a considerable number of liquid

systems containing different combinations of donor and acceptor molecules dissolved

in solvents regarded as "inert" with respect to the interacting species. From the

spectroscopic data equilibrium constants and thermodynamic values associated with

the formation of 1:1 complexes were evaluated. A quantum mechanical theory of the

"complex resonance" was worked out by Mulliken which is of a very general nature

and explains spectroscopical observations, but has not yet made possible reliable

predictions regarding the atomic arrangements within the complexes. Direct

interferometric structure determinations in the vapour phase are regarded virtually

29

impossible in most cases because of the extremely low concentration of the complex

which may be expected to be present. X-Ray crystallographic structure determinations

have made it possible, however, to draw conclusions regarding atomic arrangements

not only in solids, but even in isolated I : I complexes, conclusions which should be of

value for theoretical workers trying to establish a more elaborate theory of charge-

transfer interaction. Particular importance may be attributed to complexes in which

direct bonding exists between one atom belonging to the donor molecule and another

atom belonging to the acceptor molecule. Complexes of this kind are above all those

formed by donor molecules containing atoms possessing "lone pair electrons" and

halogen or halide molecules. A presentation and discussion of the general results

obtained by X-ray analysis of solid adducts exhibiting charge-transfer bonding

between such atoms might therefore be of some interest. The considerations were

based on the assumption that halogen atoms are directly linked to donor atoms with a

bond direction roughly coinciding with the axes of the orbitals of the lone pairs in the

non-complexed donor molecule. The oxygen atom of an ether or ketone molecule was

supposed to form bonds with both halogen atoms in an isolated 1 : 1 complex, which

would require the halogen molecule axis to run orthogonal to the COC plane in an

ether complex, and to be situated in the plane in a ketone complex.

>=0

X-ray crystallographic investigations of halogen adducts were started in Oslo,

beginning with the solid 1:11,4-dioxan-bromine compound. The most striking feature

of the resulting crystal structure [120], the endless chains of alternating dioxan and

bromine molecules depending on linear 0-Br-Br-O arrangements running in a

direction roughly equal to the "equatorial" direction in cyclohexane (Fig. 10) was

rather unexpected. It proved that both atoms belonging to a particular halogen

molecule may simultaneously be bonded to oxygen atoms, although probably not to

the same oxygen atom The existence of halogen molecule bridging between donor

atoms contradicts previous assumptions according to which a charge transfer bond

formed by one of the atoms belonging to a particular halogen molecule creates a

marked negative charge on the partner atom.

30

'A

Figure 10. Chains in the 1:1 adduct of 1,4-dioxan and bromine.

The oxygen-bromine distance in the dioxan-bromine adduct is 2.7 A and thus

considerably larger than the sum of the covalent radii of oxygen and bromine, but at

the same time definitively shorter than the sum of the corresponding van der Waals

radii. The type of "polymerisation" of simple complexes into endless chains observed

in the crystalline dioxan-bromine compound has also been observed in the crystalline

1:1 adducts of dioxan and chlorine, resp. iodine. A comparison between the oxygen-

halogen separations and of the interhalogen bond lengths in the three 1,4-dioxan

adducts leads to the conclusion that the former increases rather slowly from chlorine

to iodine, indicating a certain degree of compensation of the effect of larger halogen

radius by the increase in charge-transfer bond strength. On the other hand, the

halogen-halogen bond length increases, although slowly, from chlorine to iodine

compared with that observed in "free" halogen molecules, an observation which must

also find its explanation in the stronger charge transfer interaction between oxygen

and halogen.

From the structures of the crystalline adducts of 1,4-dioxan it became clear that both

atoms of a particular halogen molecule are able to form bonds to donor atoms,

although apparently not to the same donor atom. The question then arose whether or

not a particular donor atom may be involved in more than one bond to halogen. This

would appear possible for n donor atoms like oxygen possessing two lone electron

pairs. In the crystal structure of the 1 : 1 acetone-bromine adduct it was actually found

that each keto oxygen atom is linked in a symmetrical way to two neighbouring

bromine atoms, thus serving as a starting point for two bromine molecule bridges,

both with a linear 0-Br-Br-O arrangement and with an angle between the two bond

directions of 110".

In the case of amine adducts it would not be expected that a nitrogen atom might be

capable of forming more than one single bond to halogen. The correctness of this

anticipation has been borne out by the results of a considerable number of crystal

structure determinations of addition compounds, usually choosing iodine or an iodine

31

monohalide as the acceptor partner. Here again, the bond direction corresponds to that

expected from simple considerations regarding the orbital of the lone electron pair on

the amine nitrogen atom in the donor molecule. The nitrogen-halogen-halogen

arrangement has always been found to be nearly linear; the bonds between the

nitrogen atom and the carbon, resp. the iodine atom are tetrahedrally arranged in the

case of aliphatic amines, essentially co-planar if the donor molecule is a

heteroaromatic amine. The strength of the charge-transfer bond may be inferred from

the short nitrogen-halogen bond distance which is only 2.3 A in all complexes formed

by tertiary amines and iodine or iodine monohalides, a value only about 0.25 A larger

than the sum of the covalent radii of nitrogen and iodine. A lengthening of the

interhalogen bond by approximately 0.2 A observed in these complexes is therefore

not surprising. The fact that halogen molecule bridges have never been observed

between amino nitrogen atoms also indicates that the nitrogen-halogen bond is rather

strong. This does not imply, however, that such bridges cannot be stable between

other kinds of nitrogen atoms. Thus, in the isolated complex containing two molecules

of acetonitrile and one bromine molecule such bridges are present, which appears

very natural because nitriles are known from spectroscopical measurements to be

weaker donors than are the amines. The only crystal structure of an amine adduct so

far investigated in which a cyanogen halide acts as the acceptor molecule is that

formed by pyridine and cyanogen iodide [121]. The complex contains a linear

arrangement N-I-C=N which is symmetrically situated in the pyridine plane along the

line drawn between the pyridine nitrogen and the g -carbon atoms. The N-I bond

distance is larger than 2.3 A by about one quarter of an A unit; in agreement with the

spectroscopical finding that cyanogen iodide is a relatively weak acceptor.

In most adducts so far investigated both participants contain more than one atom

capable of taking part in charge-transfer bonds. When each oxygen, sulphur or

selenium atom is linked to only one iodine atom in 1:1 addition compounds of

iodoform with 1,4-dioxan or its analogues, structures exhibiting endless chains of

alternating donor and acceptor molecules would be anticipated. Such chains are

actually present in these adducts, chains analogous to those found in the sulphuric

acid-dioxan compound. Figugers (11) and (12) illustrate the shape of the chains

observed in the sulphuric acid-dioxan, resp. the iodoform-dithiane compound. The

similarity between these chains again affords an example of analogy between

hydrogen and charge-transfer bonding.

32

The possibility that an oxygen, sulphur or selenium atom may be involved in two

charge-transfer bonds with halogen atoms should always be kept in mind, however,

particularly when the acceptor molecule contains a large number of halogen atoms.

Thus, in the 1:1 diselenane-etraiodoethylene adduct [122] every selenium atom is

bonded to two iodine atoms, the bond directions being roughly equatorial resp.axial

and the selenium-iodine bond lengths almost identical. A "cross-linking" of the chains

results, all iodine atoms are linked to selenium and each selenium atom to two iodine

atoms.

\ y K A

/

X y

A /

/

Figure 11. Chains of sulphuric acid and dioxin molecule in 1:1 adduct.

In view of the moderate energies apparently involved in bonding between halide

halogen atoms and n donor atoms it would appear natural to suggest that the Van der

Waals interaction energy between acceptor molecules containing a sufficiently large

number of the heaviest halogen species may contribute substantially to the lattice

energy of a solid addition compound. Crystal structure determinations of tetrabromo -

and tetraiodoethylene and of their 1:1 pyrazine adducts actually give some support to

33

this suggestion. The mutual arrangement of the halide molecules is virtually identical

in the tetrahalogenoethylene crystals and in the corresponding crystalline addition

compounds. The structures of the 1:1 adduct crystals may formally be derived from

those of the tetrahalide crystal by removing one set of equivalent molecules and

replacing them by pyrazine molecules.

Figure 12. Chains of iodoform and dithiane molecules in the 1:1 adduct.

34

For both adducts this resuhs in the formation of endless chains of alternating donor

and acceptor molecules in which each nitrogen atom is bonded to a halogen atom

situated near the plane of the hydrazine ring with a nearly linear nitrogen-halogen-

carbon arrangement and a nitrogen-halogen bond about 3 A long. In the

tetrabromoethylene adduct these chains are all parallel, in the tetraiodoethylene

adduct, however, the chains are running along two different crystallographic

directions [123]. A phenomenon which has not yet apparently been met with great

interest, but should perhaps deserve more attention, is the formation of solid solutions

between donor and acceptor molecules. Until recently, the experimental facts

favouring the suggestion of mixed crystal formation were somewhat meagre,

however, and no attempt of a crystallographic investigation had apparently been made

before an X-ray crystallographic investigation of the system

hexamethylenetetraminine-carbon tetrabromide was carried out [124]. These two

substances actually form mixed crystals, containing from zero to about sixty mole per

cent of the acceptor partner, which could be examined in the form of single crystals.

The crystals are cubic with an over-structure depending on the composition, but with

a subcell that corresponds to the true unit cell of the donor component, only slightly

decreasing in dimension as the acceptor concentration increases. The experimental

findings seem to prove

that the tendency towards the formation of solid solutions actually depends on the

faculty of the two partners to form nitrogen-bromine charge-transfer bonds. Accurate

thermodynamic measurements of this and of certain related binary systems would

appear to be of considerable interest.

Complex formation due to charge-transfer interaction between n donor atoms and

halogen atoms belonging to halide molecules does not always result in bond distances

significantly shorter than those expected for Van der Waals contacts between the two

atoms. This may partly be due to the somewhat "diffuse" character of Van der Waals

radii, partly also be explained if the bond, in the case of very weak charge-transfer

interactions, has an intermediate character. The presence of charge-transfer interaction

is indicated by the angle between the halogen-donor atom bond direction and the bond

between the halogen atom and the atom in the acceptor molecule which is directly

linked to it. This angle tends to be about 180°. It is readily recognized that the charge-

transfer contribution to the bond is substantially increased when a lighter halogen

atom is replaced by a heavier one. Thus, the nitrogen-halogen distance is actually a

35

little shorter in the 1:1 adduct pyrazine-tetraiodoethylene than in the corresponding

tetrabromoethylene adduct. Even in complexes where "active" hydrogen atoms are

linked to nitrogen or oxygen atoms bond distances are difficult to predict accurately,

and observed values are often insignificantly shorter than those calculated under the

assumption of a Van der Waals interaction. In such cases arguments in favour of a

weak "hydrogen bond" between donor and acceptor molecule must to some extent be

based on the actual geometry of the complex.

1.8. Recently Research Work on Charge Transfer Complexes

Electron transfer or charge transfer (CT) is one of the most important elementary

processes in chemistry. Since Mulliken presented the well-known theory [125-126] of

the charge-transfer interaction between electron donor and acceptor, it has been

successfully and widely applied to many interesting research subjects [127-128]. One

of them is the possible role of CT complexes in chemical reactions [129]. Organic

charge transfer solids offer a wide range of materials from insulator to

superconductors [130-131]. Charge-transfer (CT) complexes of carbazole, N

ethylcarbazole and l,ndi ( N-carbazolyl) alkanes with p-chloranil (p-CHL) have been

investigated by Arslan et al [132]. The synthesis of small-molecular organic

conductors and nanowires of TTF-TCNQ have been carried out by selective

inducement in a two-phase method by n-n stacking interaction [133].

TCNQ (tetracyanoquinodimethane) as 7i-acceptor has been used with different

phenolic donors like p-aminophenol, a-naphthol, 2, 4, 5-trichlorophenol and p-cresol

by Chauhan et al. [134] Different types of radical ion salts have also been synthesized

using TCNQ molecules [135-137], but no attention has been paid to the use of

chloranil as p-accepter for the formation of charge transfer complexes. Braun et al.

[138] studied, by means ultraviolet photoelectron spectroscopy (UPS), the organic

hetero-junctions in multilayered thin film stacks comprised of alternating layers of

TTF and TCNQ. He showed that energy level alignment at the organic-organic

interfaces in the stacks depended solely upon the relative energy structure of the

donor and acceptor molecules. The scheme of charge transfer complex between

tetrathiafulvalene and tetracyanoquinodimethan (TTF-TCNQ CTC) shown in (fig. 13).

36

NC •:H wc CM r\

^ o 'x.

/ • ^ - .

S o

\,.,.j' N

1 J ^

Ij Solvent

i i

; ' CN

r^-. J ^

('

/ ' • • • . .

K "f ^ \ = i ' '

ii t

X ^ "^~-,

— . J ^

N :• \ : N

'ITF TCMQ ITFTCNQ Ch^ge tt^aislc: comsilex

Figure 13. Formation of the TTF-TCNQ charge transfer complex.

A single crystal of the charge transfer complexes of piperazine, N,N-

dimethylpiperazine and hexaethylbenzene with tetracyanoethylene was investigated

for conformational effect and charge transfer transitions [139-141]. In recently so

many compounds has been synthesis for the formation of newly type of charge

transfer complexes. One of them the N-heterocyclic compound bipyridine or 1,10-

phenanthroline possesses a system of TI and n-electrons and has been extensively used

in the synthesis of various coordinate complexes. CT complexes of 4,4'-bipyridine

with benzoquinone derivatives are also reported [142], in which the IR and IH NMR

spectroscopic analyses reveal the migration of a proton from acceptor to donor

followed by intermolecular hydrogen bonding in charge transfer interaction [143].

Charge-transfer complexes (CTC) of metoclopramide with picric acid (PA), 2,3-

dichloro-5,6-dicyano-p-benzoquinon (DDQ), tetracyanoquinodimethane (TCNQ), m-

dinitrobenzene (DNB), p-nitrobenzoic acid (p-NBA) and tetrachloro-p-quinon (p-CL)

have been studied spectrophotometrically in absolute methanol at room temperature.

The stoichiometrics of the complexes were found to be 1:1 ratio by the

spectrophotometric titration between metoclopramide and represented Jt-acceptors.

Metoclopramide (Mcp), chemically known as 4-amino-5-chloro-N-(2-

diethylaminoethyl)-2-methoxybenzamide. It is also used for the prevention of cancer

chemotherapy-induced emesis at much higher doses [144]. The great therapeutic

importance of Mcp in both clinical and experimental medicine has resulted in

extensive literature on its determination in dosage forms and biological fluids. Both

British Pharmacopoeia [145] and United State Pharmacopeia [146] describe acid-base

titration with potentiometric end point detection. Several methods have been reported

37

for the determination of Mcp in pharmaceuticals, biological fluids or mixtures with

other drugs; by HPLC [147], ' H N M R spectrometry [148], differential scanning

calorimetry [149], X-ray powder diffractometry [149], voltammetry [150].

potentiometry [151], flow-injection chemiluminescence spectrometry [152-153],

fluorimetry [154], UV-spectrophotometry [155] or flow-injection spectrophotometry

[156]. Some of the reported procedures are not simple for routine analysis and

required expensive or sophisticated instruments. Literature survey revealed that no

titrimetric assay of Mcp has ever been reported except the official methods. Donor-

acceptor interactions are vital and important subject in biochemical and

bioelectrochemical energy transfer process [157]. The charge transfer complexes have

been widely studied spectrophotometrically in the determination of the drug based on

the CTcomplexes formation with some 7i-acceptors [158-160]. The interactions of the

charge-transfer complexes are well known in many chemical reactions such as

addition, substitution and condensation [158-161]. The molecular interactions

between electron donors and acceptors are generally associated with the formation of

intensely colored charge transfer complexes, which absorb radiation in the visible

region [162]. Electron donor-acceptor CT-interactions are also important in the field

of drug-receptor binding mechanism [163], in solar energy storage [164] and in

surface chemistry [164] as well as in many biological fields. On the other hand, the

CT-reactions of p-acceptors have successfully utilized in pharma-ceutical analysis and

electrochemical properties. Many drugs are easy to be determined by

spectrophotometric methods based on formation of colored charge-transfer (CT)

complexes between electron acceptors, either-or 7i-acceptors and drugs as electron

donors [165-166] or formation of colored compounds with a number of organic acid

dyes.

The electron donor 1,2-dimithylimidazole (DM1) have been selected to study the

interaction for the first time with acceptor 2,4-dinitro-l-naphthol (DNN). The

interaction have been studied spectrophotometrically in different solvents at room

temperature with view to understand the reactivity of donor and acceptor in different

solvents like chloroform, acetonitrile, methanol, methylene chloride etc. and

interpreted this new type of interaction by using instrumental techniques such as

IHNMR, 13NMR, FTIR, TGA-DTA and electronic absorption.

38

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46

CHAPTER-2 Synthesis, spectrophotometric and spectroscopic studies oi charge transfer complex of 1,2-dimethyimidazole as an

71- acceptor 2,4-dinitro- 1-naphthol in different polar

solvents at room temperature

2. INTRODUCTION

In ihe past twcKe >cai"s. much allenlioii has bcei) locuscu on Drgaiiic'clccLRJiii*.

devices, non linear optical material, solar energy storage, organic conductors and

semiconductors al! based on electron donor-acccplor (FDA) complexes [12].

Nowadays the important role of charge transfer (CT) complex in material science |3-

7]. photo catalysis [8]. dendrimers [Q] and also play an imporlanl role in }}Kt<)\

biological systems, e.g., transfer of charge from one molecule to another 110-15]

during drug action, enzyme catalysis and ion transfer through lipophilic mcmhra'ics

all involving complexation. Characterization of new type donor-acceptor complexes

(DAC) is very challenging and for this purpose many kindes of electron donors and

acceptors have been designed synthesized in the last decades [16-17]. The charge

transfer complexes (CTC) are formed, when combining two compounds one is good

electron donor (i.e., has a low ionization potential) and other is a good electron

accepter (i.e.. has a high electron affinity), absorption maxima appear that otherwise

are not characteristic of either compounds alone, it is suspected that a charge transfer

complex (CTC) has formed between the components of mixture, fhis theor\- was Hrst

proposed by Mulliken [18-19] and was very successful in explaining the origin of the

charge transfer absorption band and also the variations in the spectra as d(>!!or and

acceptor properties of the components were varied.

Up to now the factors governing the behavior, stability, etc. of these complexes have

not been clearly established. Among the different theories proposed, one of the best

known was deri\ed b) Mulliken [20-22J. who considers the complex as a h\lniu

resonating between a non-polar structure and a polar one, resulting from the transfer

of one electron. In Dcwar's theory [23], the transfer is supposed lo take place hetwcer;

the highest occupied orbital of the donor and the lowest empty one of the acceptor.

More recently. Flurry [24] studied these complexes as new molecules whcise u-i-.c

function is a linear combination of the HOMO and LUMO, from the donor and the

acceptor, respectively. Using perturbation theory. Murrcll [25] calculated the cnerg;

as the total of the contributions from different types of interactions (van der Waals.

electrostatic, etc.). Chesnut and Moseley [26] consider these complexes as a

"supermolecule" made up by the donor and acceptor molecules. Formation of the

charge transfer (CT) complex has been characterized by an intense, broad, electronic

absorption band in the UV-Visible region by using different type of techniques [27]

(i.e.. TGA-DTA. FTIR. 'H NMR and ' C NMR spectroscopic techniques etc )

48

Ultraviolet -visible spectroscopy (k = 200-800 nm ) studies the changes in electronic

energ) levels within the molecules arising due lu transfer of electrons from n-or non-

bonding orbital's .Basically the formation of charge transfer (CT) complex depend

upon the excitation of electrons from donor to acceptor compounds [28-29]. The

stoichiometry, structure, spectral, thermal, and other physical properties of the

complexes depend upon the naUire of the donor base as well as ii-acceptors. In this

paper we studied formation of charge transfer complex between electron donor 1,2-

dimethyiimida7,oic{DM!) and n-acceptor 2.4- dinitro-!-naphthoI(DNN) in different

polar solvents such as chloroform, acetonitrile .methanol and methylene chloride at

room temperature by using UV-visible spectrophotometer and also studied the effect

of solvents on the formation of CT complex. The molecular association between

electron donor (DMI) and n-acceptor (DNN) absorbed in the visible region [30]. 2,4-

dinitro-1-naphthol forms molecular complexes with aromatic hydrocarbons, aromatic

amines [31 -33] and also with some aniline derivatives.

2.1. Experimental Study:

2.1.1. Materials and methods:

1,2-Dimithylimidazole (DMI) and 2,4-Dinitro-l-naphthol (DNN) was obtained from

Sigma-.Mdrich ( fine chemical society ) and were used without purification. Double

distilled water, chloroform (Merck), acetonitrile (Merck), Methanol (Merck analytical

grade), and Methylene chloride (Merck) were distilled prior to use according to the

standard protocols.

2.2. Procedure:

2.2.1. Synthesis of solid charge transfer (CT) complex:

The isolated CT complex were formed by mixing 2 m. mol. saturated chlorofomic

solution of donor 1,2-dimethylimidazole ( 0.19226 g ) with 2 m. mol. saturated

chlorofomic solution of acceptor 2.4-dinitro-l-naphlhol ( 0.46834 g ). The mixture

was stirred at room temperature for a 5-10 minutes, there very fine solid CT complex

were formed in yellow needles shape. Then after that filtered it through whatmann: 41

grade filter paper to remove the suspended impurities, washed with minimum amount

of chloroform and dried it under vacuum.

2.2.2. Preparation of standard stock solutions:

The standard stock solutions were prepared by using total concentration of acceptor

(DNN) (varied) and the total concentration of acceptor (DMI) (fixed) in different

49

solvents. A standard stock solution of 2,4-dinitro-l-naphthol (DNN) 10" M

(acceptor) was prepared by dissolving 0.23417 g in a 100 ml '^oluniclric ilask uiMng

chloroform as a solvent and the solution of different concentration of acceptor ( 1 >

10~ M. 1.5x 10" M, 2.0>' 10" M. 2.5'' 10^ M and 3.0>^!0^ M ) were prepared in

50 ml individual volumetric flask by diluting 10~ M solution with same solvent. The

standard stock solution of 1.2-dimithylimidazole (donor) 10 " M was prepared h}

dissolving 0.09613 g in a 100 ml volumetric flask in the same solvents. Then after

that another five solutions were prepared by mixing 3 ml volume froin each : 0 mi

individual volumetric flask and 3 ml volume from the fixed standard stock solution of

donor (DMI) in 25 ml individual volumetric flask without further diluting. In similar

way many other solutions were prepared in different those solvents in which the

compound is soluble.

2.3. Spectral measurement and determination of formation constant:

The fine electronic absorption spectra of the donor 1.2-dimethylimida7ole (DM!)

acceptor 2,4-dinitro-l-naphthol (DNN), and the resulting CT complex in solvents

like chloroform, acctonitrile. methanol and methylene chloride were recorded in the

UV region (200-490 nm) using the spectrophotometer modal SONAR Ll-295 IJV

visible speclropholomeler with 1 cm quart/ cell path length in the range (200-3=)f> nm)

and above this range (345-490 nm) with 1 cm glass cell path length. The electronic

spectra of DMI and DNN in chloroform, acctonitrile and methanol. The FTIR spectra

were recorded employing spectroscopic interspec 2020 FTIR spectrometer using the

KBr disk technique. The thermo gravimetric analysis (TCJA) and differential therniai

analysis (DTA) of reactants resultant CT complex were recorded with the heating rate

of 20 °C/ min under the nitrogen atmosphere using Shimad/u model DTG -60! I

Thermal Analyzer and also the melting point ( m. p.) was calculated 142-150 "C by

using Reichert Thermovar. 'H NMR and 13C NMR of resulting CT complex were

recorded in CDCL3 solvent using the Bruker advance II 400 NMR spectrometer. Ihe

Bcnesi-Ilildcbrand [34] equation was used for determination of KCT (formalioii

constant of CT) for DA association and £CT (extinction coefficient) for DA CI

absorption with the help of spectrophotometric data. The Benesi-Hildebrand analvsis

of KcT involves the measurement of D—A CT absorbance (A) as a function of varied

[A] » fD].

A plot of X = 1/ [A]o vs Y = [D]o/A gives a Y-intercept = 1/ 8CT and slope = 1/ Sex

Kci as a defined by Bencsi- Hildebrand equation:

[D]o/ A = (1/KCT£CT) (1/[A]O) + 1/(8CT)

Where [DJo and [AJo are the initial total concentration of donor (tlxed) and acceptor

(varied). A is the CT absorption of DA complex at wavelength X. The theoretical

foundation of the above equation is the assumption thai when either one of the

reactants is present in excess amounts over the reactant, the characteristic electronic

absorption spectra of the other reactant will be transparent in the collective

absorption/emission range of the reaction system [35]. Therefore by measuring the

absorption spectra of the reaction before and after the formation of the product and its

equilibrium, the association constant of the reaction can be determined. This equation

is west applicable to study reactions with 1:1 product complexes.

2.4. Results and discussion:

2.4.1. Observation of CT electronic spectra

The electronic absorption spectra of 1 x 10" M solution of the donor (DMI), 1x10^

M of acceptor (DNN) and their 1:1 (donor : acceptor) CT complex recorded in

chloroform (CHCI3), acetonitrile (CH3CN). methanol (CH3OH) and methylene

chloride (CH2C12)). To obtained the CT bands, the spectrum of the solution of 1 ^ 10"

M 1,2-dimethylimidazole (DMI) and Ix 10"* M 2,4-dinitro-l-naphthol (DNN) in

different solvents was recorded with solvents used as reference, it is observed that

new absorption peak appear in the visible region. In some cases multiple peaks were

obtained, the longest wavelength peak was considered as CT peak [36J where neither

donor nor acceptor has any absorption. It was the solid information for the formation

of the CT complex of donor (DMI) and acceptor (DNN) and absorption band

appeared at 245 nm of donor-acceptor mixture in chloroform (CHC13) at room

temperature. The change of the absorption intensity to higher for all complexes in this

study when adding the donor is reported in TabIe-1.

51

TabIe-1

Absorption data for spectrophotometric determination of stoichiometry, absorption

maxima (kcT), formation constant (Kc r) and molar extinction coefficient (SCT) of the

[(1,2-DMl)^ (2,4-DNN) J complex at room temperature.

Concentration of

Acceptor (M)

In chloroform

1.5x10-*

2.0x10"*

2.5 X 10"*

3.0x10-*

In Methylene

Chloride

1.5 X10^

2.0 X 10-"

2.5 X10^

3.0 X 10^

In Methanol

1.5 X 10-"

2.0 X 10"*

2.5 X IQ-*

3.0 X 10""

In Acetonitrile

1.5 X 10 "

2.0 X 10^

2.5 X 10""

3.0 X 10-"

Concentration of

Donor (M)

1x10-"

1x10^

1x10-"

1 X 10 "

Absorbance at

^CT

(nm)

At 245 (nm)

1.569

1.821

2.043

2.208

At 255 (nm)

1.509

1.801

2.089

2.170

At 270 (nm)

1.644

1.904

2.282

2.584

At 290 (nm)

1.792

2.210

2.760

2.867

Formation

constant

(KcT)/mor'

0.4829

0.3814

0.2595

0.1758

Molar extinction

coefficient

(DCT)/

cm"' mol'

3.7278

4.1635

5.7796

8.6006

Charge transfer (CT) absorption bands of all CT are exhibited by the spectra of the

system mentioned as above and depicted in fig.l.

52

200 250 300 350 400

Wavelength (nm)

450 500

Figurel. Electronic absorption spectra of [ (1,2-DMI)'^ (2,4-DNN) ] complex (1 x

10""* M + 1 X 10"^ M) in (A) chloroform; (B) methylene chloride; (C) methanol (D)

acetonitrile.

The charge transfer (CT) spectra are analyzed by fitting to the Gaussian function y =

y*' + [A/(w V (jr/2))] cxpf-2(x-xc)^/w^] where x and y denote wavelength and

absorbance, respectively.

The wavelength at these new absorption maximum (Xcr = xc) are summarized. It is

observed that the ACT of CTCs increase as the polarity of the solvent is increased due

to change in dative structure. The concentration of the acceptor in each of the reaction

mixture was kept greater then donor and change over the wide range of concentration

from 1x10"^ M to 3.0x10"* M while the concentration of the donor (DMl) was kept

fixed at 1X 10^ M in each of the reaction mixture, which are used to obtain a straight

line diagram for the determination of formation constant (KCT) molar extinction

coefficient (ECT) of the resulting CT complex.

53

2.5. Conclusions:

Charge Irunslcr ct)inple\alion between eleelron donor l,2-Diniilh>liniida/.olc (i)Mli

and acceptor 2,4-Dinitro-l-naphthol (DNN) has been studied spectrophotometricalK

in different polar solvents included chloroform (CllCf;)- acctoniirilc (CllCN;.

methanol (CH3OH) and methylene chloride (CH2CI2) at room temperature. Ihe

complex formation was confirmed from the appearance of a new wavelength h.'nd

associated with colour change from colorless to yellow. The formation constant has

been estimated by using Bcncsi-llildcbrand equation where it reached high \aluc

confirming high stability of the complex. The stability of the formation constant

depends upon the polarity of the solvents i.e.. lower the polarity of the solvent, hiyher

will be the formation constant and hence more stable complex has been formed

because in less polar solvents CT complex very less dissociate in to the ions and this

is due to the low dielectric constant of very less polar solvent. Furthermore, the

formation constant recorded higher values and molar extinction coefficient recorded

lower values in chloroform compared with methylene chloride, methanol and

acctonitrile. This confirmed the strong interaction between the molecular orbital's o!'

donor and acceptor in the ground state in less polar solvent. Some spectroscopic

physical parameters like formation constant (K( r)- molar extinction coefilcien! ir, ;•).

energy of interaction (ECT), ionization potential (ID), resonance energy (RN). free

energy (AG), resonance energy (R,\) oscillator strength (!) and transition dipolc

moment (I^N) are estimated where they showed solvent dependent. The solid complex

was obtained, its elemental analysis confirmed its formation in 1:1 ratio (doiiv»r.

acceptor). The infrared spectrum of the solid complex confirmed the presence of

proton transfer beside charge transfer that adds extra stabilit} to it. The spectra!

analysis indicate 71-%* transitions and N^-H-"0~ type intermolecular H-bonding

between l,2-Dimithylimida/.ole (DM1) and 2,4-Dinitro-l-naphlhol (DNN). The

results confirmed the presence of charge transfer complex with hydrogen bonding in

agreement with infrared, 'll NMR and '' CNMR measurements.

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