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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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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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