TABLE OF CONTENTS
Title page — — — — — — — — — — i
Declaration — — — — — — — — — ii
Certification — — — — — — — — — iii
Dedication — — — — — — — — — iv
Acknowledgment — — — — — — — — v
Abstract — — — — — — — — — — vi
Abbreviations — — — — — — — — — vii
Table of contents — — — — — — — — ix
List of tables — — — — — — — — — xi
List of figures — — — — — — — — — xii
Abbreviations — — — — — — — — — xiv
CHAPTER ONE
INTRODUCTION
1.0 Introduction — — — — — — — — 1
1.2 Methods of Monitoring Reaction Rates– — — — 1
1.2.1 Conventional method (slow technique) — — — — 2
1.2.2 Monitoring the rates of fast reactions — — — — 2
1.3 Theories of Reaction Rate — — — — — — 4
1.3.1 Arrhenius Theory– — — — — — — 5
1.3.2 Collision Theory — — — — — — — 5
1.3.3 Theory of Absolute Reaction Rate — — — — — 6
1.4 Theories of Electron Transfer Processes — — — 7
1.4.1 Marcus Theory — — — — — — — — 7
1.4.2 Electron tunneling theory — — — — — — 8
1.4.3 Frank Condon Principle — — — — — — 9
1.5. Electron Transfer Reactions– — — — — — 11
1.5.1 Classes /Types of Electron Transfer Reactions — — — 12
1.5.1.1 Homonuclear or Isotopic Exchange Reactions– — 12
viii
1.5.1.2 Heteronuclear or cross reaction — — — — — 13
1.6 Proton-coupled electron transfer (PCET) — — — — 13
1.7 Mechanism of electron transfer reactions — — — 15
1.7.1 Outer-sphere mechanism — — — — — — 16
1.7.2 Inner-sphere mechanisms — — — — — — 17
1.7.3. Distinction between the Outer-Sphere and
Inner-sphere Reactions– — — — — — 19
1.8 Determination of the mechanisms of redox reaction — — 20
1.8.1 Identification of binuclear intermediate — — — — 20
1.8.2 Reactivity Patterns — — — — — — — 21
1.8.3 kred versus ksub — — — — — — — — 23
1.8.4 Effect of added ions — — — — — — — 23
1.8.5 Activation Parameters — — — — — — — 24
1.8.6 Product Identification — — — — — — — 24
1.8.7 Michaelis-Menten Plots — — — — — — 25
1.9 Objectives of the Project — — — — — 27
CHAPTER TWO:
2.0 Literature Review — — — — — — 29
2.1 Electron transfer reaction of μ-oxo bridged Fe(III) complexes 29
2.2 Kinetics and mechanism of the reduction of μ-adi-di(N,N/- bis
{salicylideneethylenediaminatoiron (III)} by dithionate ion — 30
2.3 Kinetics and mechanism of electron transfer reactions of thiols
(L-cysteine, thiourea, thioglycolic acid, 2-mercaptobenzothiazole
and benzyl mercaptan) — — — — — — — 30
CHAPTER THREE:
3.0 Experimental — — — — — — — — 35
Equipment — — — — — — — — 35
x
3.1 Synthesis of complexes — — — — — — — 35
3.1.1 Ferric hydroxide, [Fe(OH)3] — — — — — — 35
3.1.2 N,N/ bis(Salicylidene)ethylenediamine, (H2Salen) — 35
3.1.3 μ-oxo-di(N,Ni-bis(salicylidene)ethylenediaminatoiron
(III), [(Fe-Salen)2O] — — — — — — — 36
3.1.4 μ-adipato-di (N,NI-bis (salicylidene) ethylene
diaminatoiron (III), [(Fe-Salen)2 adi ] — — — — 37
3.2 Preparation of reagents — — — — — — 38
3.2.1 Sodium Perchlorate salt (NaClO4) — — — — — 38
3.2.2 Perchloric acid (HClO4) — — — — — — 38
3.2.3 Sodium acetate (CH3COONa) — — — — — 38
3.2.4 Magnesium Chloride (MgCl2) — — — — — 38
3.2.5 L-cysteine (LSH) — — — — — — — 38
3.2.6 Thiourea (USH) — — — — — — — — 38
3.2.7 Thioglycolic acid (GSH) — — — — — — 39
3.2.8 2-mercaptobenzothiazole — — — — — — 39
3.2.9 Benzyl mercaptan — — — — — — — 39
3.3. Stoichiometric Studies — — — — — — — 39
3.4. Kinetic measurements — — — — — — — 40
3.5 Test for free radical — — — — — — — 40
3.6. Product Analysis — — — — — — — — 41
3.7 UV Analysis — — — — — — — — 41
3.8.0 UV and IR Analysis — — — — — — — 41
CHAPTER FOUR:
RESULTS AND DISCUSSION
4.0. Result — — — — — — — — — 42
4.1 Stoichiometry — — — — — — — — 42
4.2 Determination of order of reaction — — — — — 48
4.3 Effect of hydrogen ion concentration on the rates of reaction — 65
ix
xi
4.4 The effect of ionic strength — — — — — — 72
4.5 Effect of added ions — — — — — — — 79
4.6 Effect of dielectric constant — — — — — — 79
4.7 Temperature dependence of rates of reaction — — — 86
4.8 Test for formation of intermediate complex and products — 94
4.8.1 Michaelis-Menten plots — — — — — — 94
4.8.2 Product Analysis —- — — — — — — 94
4.8.3 Test for free radicals — — — — — — — 100
4.8.4 UV and IR Analysis — — — — — — — 100
4.9 Summary and conclusion — — — — — — 109
References
Appendix
x
CHAPTER ONE
1.0 INTRODUCTION
The electron transfer reactions of binuclear iron (III) complexes have
attracted a lot of interest in recent time due to their application as models for the
investigation of the physiological role played by iron in biochemical processes 2,
such as hemerythrin 2,3,4.6 and ferric porphyrin7,27,28 47. Previously, the dynamics
of electron transfer reactions of dinuclear oxo bindged iron(III) complexes of the
form [Fe2O]4+ with ascorbic acid 4, b – mercapto acetic acid5 and b –
mercaptoethylamine 6 have been investigated. Most of these reactions followed
outer sphere electron transfer route with intervening ion-pair complexes and free
radicals..
The behaviour of transition metal ions with respect to their electron
transfer and the roles played by bridging ligands in the course of redox reaction
formed the bed rock of this study. 37,39 The main advantage of this research is that
the results provide additional insight into the complexities attending reactions of
bridged iron(III) complexes and the extent of influence of the bridging ligand on
the rate of electron transfer. It is therefore hoped that this research will enhance
the knowledge of the kinetics and mechanisms of electron transfer reactions of
binuclear iron (III) complexes and other transition metal complexes with these set
of thiols.
1.2 Methods of Monitoring Reaction Rates
The first step in kinetic analysis of a given reaction is to ascertain the
stoichiometry of the reaction and to identify any side reaction. The fundamental
data of chemical kinetics are the concentrations of the reactants and products at
different times after a reaction has been initiated.1 The rates of most chemical
reactions are sensitive to the temperature aid. In conventional experiments, the
temperature of the reaction mixture must be held constant throughout the course
of the reaction.
1
2
The method employed in monitoring the rate of a reaction depends on the
concentration of the species involved and on how fast the concentrations change.
Reactions may take seconds, minutes or hours before they can reach equilibrium.
The techniques used to monitor the change in concentration are as follows:
1.2.1 Conventional Method (Slow Technique)
Conventional methods involve the measurement of the concentration or any
physical property of one or more of the reactants or products as a function of
time. For instance, in some reactions absorbance of any of the reactants or
products could be measured and related directly to the concentration.
In kinetic analysis, the composition of the system is examined while the reaction
is in progress by either withdrawing a small sample or the bulk and the reactants
are mixed as they flow together in a reaction container. At different level in the
observation tube, the mixtures are examined at different time of mixing and by
doing so, the rate of the reaction is obtained.
The conventional method is difficult for rapid reactions due to the fact that:
(i) The time it takes to mix reactants or to bring them to a specified temperature
may be significant in comparison with the half life of the reactants.
(ii) Also, the time that it takes to make measurement of concentration is
significant compared with the half life.
1.2.2 Monitoring of the Rates of Fast Reactions
The rates of fast reactions can be monitored effectively by the following
methods:
3
(i) Flow Techniques:
Flow techniques were developed in an effort to monitor the rates of a very
fast reactions at the shortest possible time.3 Different flow techniques exist
depending on the treatment given to the reaction after mixing. They include
continuous flow technique, quenched flow method and stopped flow technique.
In continuous flow technique, the reaction solution is allowed to flow along an
observation tube where the changes in the reaction mixture is monitored at
different points along the tube or at a fixed point in the tube.
Quenched flow method involves quenching a reaction in progress after it
has been allowed to proceed for a certain period of time. In this way, a reaction
mixture which has reaction time scale on the order of milliseconds can be studied
with ease. Once the reaction has been quenched, the mixtures comprising the
concentration of reactants, intermediates and products can be measured by
chromatographic (slow technique) or spectroscopic method.
In stopped flow technique, the reaction mixture is put to the reaction
cuvette, where the reactants are brought into a complete contact in less than 10-3
second.1 The technique allows for the study of reactions that take place on the
time scale of millisecond. This technique is efficient in monitoring many
biochemical reactions like the enzymatic action of some proteins. Spectroscopic
method is used effectively in this technique.
(ii) Relaxation Method: (Temperature Jump Method)
Relaxation method is used to analyze a very fast reaction. 1,2 When an electric
spark is passed through the solution, the spark causes a very large, but brief rise
in temperature. This upsets the solution in equilibrium such that it relaxes to
another equilibrium state. In this way the concentration of the solution can be
measured spectrophotometrically. This is popularly known as temperature jump
method.
4
(iii) Resonance Techniques:
Rates of reaction could be monitored by using nuclear magnetic resonance
technique 1. Resonance absorption line is related to the
2
1 t of the
nucleus in a given energy state. If the life-time of these states is shortened by a
chemical interaction, it results into line broadening. 1H n.m.r line broadening has
been used to measure the rate of change of various mono and bidentate nitrogen
and oxygen donor ligands coordinated to Mn(II),Fe(II), Co(II),Ni(II) and Cu(II).
(iv) Flash Photolysis
This technique can measure rates of reactions that are extremely fast. In
this case, a very short but intense flash of light passes through the mixture. After
a brief period of time, another flash of light passes through the mixture. The
molecules produced in the reaction absorb light from the second flash.3 By taking
a photograph, the spectrum of the molecules can be recorded and the intensity of
the lines in the spectrum gives a measure of the concentrations of the molecules.
If the time interval between the first and second flashes is changed, the intensity
of the lines changes. In this way, a series of experiments allow the way the
concentration of the molecules changes with time to be found. An example is the
light induced dissociation of chlorine gas. Other methods of monitoring rates of
reactions are titrations, colour changes, volume changes, and pressure changes.
1.3 The Theories of Reaction Rate
The general goal of theoretical chemical kinetics is to rationalize many of
the empirical (or observed) facts of chemical kinetics in terms of molecular
properties. Prominent among these facts are the effects of concentration and
temperature on reaction rates. Indeed, the ultimate goal of theoretical chemical
kinetics is the calculation of the rate of any reaction from a knowledge of the
fundamental properties of the reacting molecules, namely, their masses,
diameters, moments of inertia, vibrational frequencies, binding energies etc. The
main theories describing the rates of reaction are highlighted below.
5
1.3.1 Arrhenius Theory
Arrhenius theory states that the rates of a chemical reaction always
increases with increase in temperature to a marked extent. It has been observed
that as a rule, the specific rate constant of a homogeneous reaction is usually
increased by a factor of about two or three for every 1 degree rise in
temperature.9,38 An expression relating rate constant with temperature was
derived by Arrhenius in 1889. According to him,
k = Ae – RT
Ea
k = Ae- ————————————————————— 1.10
Where k is rate constant
A is called pre-exponential factor or frequency factor.
Ea is the activation energy
R is the universal gas constant
A and Ea are collectively known as the Arrhenius parameters.
1.3.2 The Collision Theory of Reaction Rate
This theory makes the basic assumption that for a chemical reaction to
occur, particles must collide. 9,38 In the reaction
A + B ® AB ……………………………………………………………(1.11)
The particles A, be the molecules, ions or atoms must collide with particles B. In
collision, chemical bonds in atoms and electrons are always rearranged and as a
result, new species are produced. According to the collision theory, the rate of
any step in a reaction is directly proportional to,
(i) The number of collisions per second between the reacting particles
involved in that step and
(ii) The fraction of these collisions that are effective
Actually, not all collisions lead to reaction, otherwise every bimolecular reaction
occurring at the same temperature and concentration would occur at the same
rate. Besides, since the frequency of binary collision is proportional to
6
2
1
T an increase in temperature say from 500K to 510K will increase the
collision frequency by a factor of 2
1
500
510
÷ø
ö
çè
æ = 1.01 or 1 percent. The rate of
chemical reaction on the other hand, may have increased by 200% or more.
1.3.3 The Theory of Absolute Reaction Rates
The theory of absolute reaction rate is also called the transition state theory
9. The theory as developed by Eyring (1935), postulates the existence of a
transitory molecular species known as the activated complex which is in
equilibrium with the reactants. The activated complex is the configuration of the
atoms which corresponds energetically to the top of the energy barrier separating
the reactants from the products. This region of high energy defines the transition
state or the activated complex. The energy difference between the stable reactants
and products is the heat of reaction, which is a thermodynamic quantity.9 On
the other hand, this theory postulates a state of equilibrium between reactant and
the activated complex. The theory asserts that, if the reactants progress along the
path of products, an intermediate complex or transition state prevails. The
transition state –complex exists in equilibrium with reactants. The rate of reaction
is then assumed to depend on the concentration of the activated complex and the
rate with which it break up to give the products. Thus, for a reaction between A
and B molecules, we can write
A + B [AB]# ® Products …………………………………….. (1.12)
The concentration of the activated complex is obtained from the equilibrium law
since it is assumed to be a thermodynamic entity. It is stated as follows:
K# = [ ] [ ][ ]
[ ][ ]
[ ] # #
#
or AB K A B
A B
AB = …………………………………… (1.13)
The activated complex is an unstable species and is held together by loose bonds.
A suitable vibration of frequency v will cause its dissociation into products. The
rate at which the products are formed is then given by,
7
rate = v [AB]# = vK# [A] [B] ……………………………………. ….. (1.14)
The transition state theory suggests that the structure of the activated
complex is necessary for the calculation of the entropy of activation. The
uncertainties about the structure of the activated complex and the assumptions
involved in computing it’s thermodynamic properties seriously limit the practical
value of the theory. However, it does provide qualitative interpretation of how
molecules react and a reassuring foundation for the empirical rate expressions
inferred from experimental data.
1.4 Theories of Electron Transfer Processes
The first and accepted theory of electron transfer was proposed in 1965 by
Rudolph A Marcus. The theory was meant to address the issue concerning
outer electron transfer and was based on transition state theory approach. This
theory was extended to include inner-sphere electron transfer by Noel Hush.
Other theories like electron tunneling theory and Franck-condon principle have
also been developed through extensive studies by chemists and physicists.1 The
three outstanding theories are therefore discussed below.
1.4.1 Marcus Theory
Calculation of electron transfer rates using such parameters like interatomic
distance, dielectric constants, force constant, e.t.c is difficult. However for
reactions occurring by outer-sphere mechanism, the weak interaction between
reactants during electron transfer makes it possible that kinetics and
thermodynamic parameters can be related.47
Marcus calculated the minimum energy needed for electron transfer to
occur. According to Marcus theory, the rate constant for outer-sphere electron
transfer is a product of four factors as related in the equation below:
÷ø
ö
çè
æ -D
= – RT
WR G
k ZK
*
*exp ……………………………………………………(1.15)
8
From the equation
(1) Z represents the collision frequency between two neutral molecules in
solution. It is not the diffusion limited rate constant since it also includes
encounters between reactants in a solvent cage. For water at 25oC, Z =
1011 cm3 s-1.
(2) K* is the transmission coefficient. It is related to the probability that
electron transfer will occur once the intersection between the potential
coordinate modes of the redox couple is reached. K* have values close to
unity in most simple outer-sphere electron transfer reactions.
(3) WR is the free energy change associated with bridging together of the
reactants and is unfavourable for unlike charged reactants since they have
mutual attraction.
(4) DG * is the minimum free energy increase above the back ground thermal
energy. R and T are the universal gas constant and absolute temperature
respectively. RT is required in the vibration and solvent trapping modes in
order for electron transfer to occur with energy conservation. DG* is also
related to the inner sphere and outer sphere reorganization energies for
self exchange reaction.
1.4.2 Electron Tunneling Theory
In comparing the classical potential energy barrier to electron migration
between complexes, the electron tunneling theory sees the electronic energy as
being low in both, reactants and product activated complexes. The theory
explains that the electron migrates by passing through the potential energy barrier
rather than over it. 38 This implies that the electron will be able to travel distances
much greater than would correspond to the actual collision of reactants.
9
Theoretically, this theory gives a relationship between the transmission
coefficient and the rate constant for electron transfer as
k = ÷
÷
ø
ö
ç ç
è
æ D
–
D
–
RT
Ge
RT
k xp G
h
TK r
o * *
1 e …………………………………………(1.16)
Where
k1 = Electron transmission coefficient
k = Rate constant
Ko = Boltzmann constant
*
e DG = Activation energy
*
r DG = Hydration energy for inner coordination shell arrangement
T = Absolute temperature
R = Universal gas constant
h = Planck’s constant
The value of the transmission coefficient is less than unity and increases as the
exchanging partners come close together. Electrostatic repulsion ensures that
activation energy also increases. As a result of that the rate of the reaction tends
to decrease. At an optimum distance, a maximum exchange rate is obtained.
Electron tunneling theory is viewed as being involved in most electron transfer
reactions but might not be the rate determining step in most cases.
1.4.3 Franck Condon Principle
Electron transfer reactions which occur either by inner sphere or outer
sphere mechanisms are subjected to restrictions which was defined by the Franck
Condon Principle. Franck Condon Principle states that, the motion of the nuclei
is slow (10-13s) compared to that of the electron (10-15s), and electron transfer
occurs without significant movement of atoms.38 Since electron transfer reactions
involve bond breaking and formation, this principle must come into play. The
atomic distances between ligand and metal ions alter the oxidation state of the
10
metal ion. Therefore, the reorganization of metal-ligand distances for the
reactants and products occur before electron transfer takes place.
Alternatively, electron transfer can occur before the reorganization. For
this route, the intermediate product possesses non equilibrium configuration and
therefore, reorganization of the coordination shell must take place. This gives rise
to a highly endothermic and exceedingly low reaction.38 Electron transfer only
takes place when ions approaches each other. If the electron transfer step is fast,
the overall rate is that at which the ions diffuse together to form an ion pair.
Reactions of this type which is studied by temperature jump techniques, had rate
of the order of magnitude of the diffusion limited value. Reorganization is
undergone by the reactants before electron transfer takes place in such a way that
their transition state energy becomes almost identical and energy change on
electron transfer is minimized.
The scheme for the electron transfer is shown below:
M N approach and reorganisation M – – – – – – – – N
Electron transfer
M separation and reorganisation M – – – – – – – – N
Alternatively it can occur as follows:
M M
m
2+ + n
3+
o 2
+ 3+
o
ion pair
m. + Nn 3+ 2+.
o
3+
o
2+
m
2+ + Nn
3+ Electron transfer 3+
m n
2+
+ N
reorganisation
Mm. Nn . 3+
where
2+
11
Subscripts m and n are equilibrium configuration of the coordinate shell for
metals M2+ and N3+ respectively. Subscript 0 = intermediate configuration. The
total energy change DG* involved in the process can be represented as follows:
* * * * …………………………………………………………………………(1.17)
DG = DG a + DG i + DG o
Where DG*a = the association free energy.
G i D * = the inner sphere reorganization energy.
and DG*o = the outer sphere reorganization energy.
The principle also assume that no angular momentum is transferred to or from the
transition state during the electron transfer and a restriction is also imposed on
the change in spin angular momentum.55 For the reaction;
[Co (phen)3 ]2+ + [*Co (phen)3 ]3+ [Co (phen)3 ]3+ + [*Co (phen)3 ]2+ … (1.18)
It involves only electron transfer and so has a rate of 1.1 dm3mol-1 s-1 at 25oC. On
the other hand, for the reaction,
[Co (NH3)6 ]2+ + [*Co(NH3)6 ]3+ ¾¾® [Co(NH3)6 ]3+ +[*Co(NH3)6 ]2+ … (1.19).
It involves both electron transfer and change in spin multiplicity and so, it is slow
with a rate of 10-9 dm3mol-1 s-1 at 25oC.38, 42
1.5 Electron Transfer Reactions
Electron transfer is the process whereby an electron moves from one atom
or molecule to another. Electron transfer is a mechanistic description of the
thermodynamic concept of redox reaction where oxidation state of both reaction
partners change. Electron transfer reactions also known as oxidation-reduction
(Redox) reactions are usually studied in aqueous solution because most ions are
inert in non-aqueous solution. Oxidation of a particular species involves electron
loss and reduction involves electron gain, implying that the rate at which a redox
reactions occurs is qualitatively related to the redox potential. Each ion in
aqueous media has its standard electrode potential Eo measured in volts
12
which is determined in comparison to the standard hydrogen electrode which
is assigned zero potential.
The electrode potential of an ion gives an indication of its readiness to be
oxidized or reduced by another ion. So, ions with higher negative values of
standard reduction potentials are good reducing agent while those with less
negative values or those with positive values function as good oxidizing agent.
Numerous processes in biology like oxygen binding, photosynthesis,
respiration and detoxification routes involve electron transfer reactions. In most
cases, electron transfer reactions involve transition metals complexes, but many
examples of electron transfer reaction abounds in organic chemistry.
1.5.1 Classes/Types of Electron Transfer Reactions
Electron transfer reactions can be divided into two broad classes. They
include homonuclear or isotopic or self exchange reactions popularly known as
outer-sphere electron transfer reactions and heteronuclear or cross reactions
popularly called inner-sphere electron transfer reactions.
1.5.1.1 Homonuclear or Isotopic Exchange Reactions
This is a type of electron transfer which involves the exchange of electrons
between two identical metal ion centres in different oxidation states. The
participating redox centres are not linked through any bridge during the electron
transfer, rather the electron “hops” through space from reducing centre to the
acceptor. 15,39 The reactants and the products are the same and identical. As a
result of that they have the same concentrations. The free energy change for such
reaction is mainly due to mixing and so, it is approximately zero. Under this type
of electron transfer, there is no net chemical change and as a result, the
equilibrium constant is one since the rate constant for the forward and reverse
reactions are equal. The reactions below represents examples of homonuclear or
isotopic exchange reactions.
13
[Fe(H2O)6]2+ + [Fe* (H2O)6]3+ ® [Fe(H2O)6]3+ + [Fe* (H2O)6]2+ ……(1.20)
[*Fe(phen)3]2+ + [Fe(phen)3]3+ ® [*Fe(phen)3]3+ + [Fe (phen)3]2+ ………(1.21)
1.5.1.2 Heteronuclear or Cross Reactions
Heteronuclear reaction is a class of reaction that involve the electron
transfer between different metal ion centres. The products of the reaction are
chemically different from the reactants and so, the over all free energy change is
not equal to zero.15,39 In this type of electron transfer reaction, the participating
redox centres are linked through a bridge during the course of electron transfer
although not in all cases. The reaction can be complementary if the oxidant and
reductant undergo equal changes in oxidation states. The stoichiometry for such
reaction is 1:1. The equation for the reaction is shown below.
[Co(en)3]3++[Ru(NH3)6]2+®[Co (en)3]2+ + [Ru (NH3)6]3+………………….(1.22)
Heteronuclear reaction could also be non- complementary whereby the oxidant
and reductant undergo unequal changes in their oxidation states. The
stoichiometry for such reaction is not equal to 1:1 and it is shown in the equation
below.
Sn
2++ 2Fe3+ ®Sn
4+ + 2Fe2+…………………………………………….(1.23)
1.6 Proton-Coupled Electron Transfer (PCET)
Proton- coupled electron transfer (PCET) is a reaction mechanism that is
thought to be common in redox reactions. It involves the concerted transfer of an
electron and proton to or from a substrate.40,41 In PCET, the proton and the
electron (i) start from different orbitals and (ii) are transferred to different
orbitals. They transfer in a concerted elementary step. PCET contrast to step-wise
mechanisms in which the electron and proton are transferred sequentially.
ET
[HX] + [M] [HX]+ + [M]———————————— (1.24)
PT
14
[HX] + [M] [X]- + [HM]+———————————– (1.25)
PCET
[HX] + [M] [X] + [HM]———————————– (1.26)
PCET is thought to be pervasive in redox reactions that appear to be net
hydrogenations and dehydrogenations. Relevant examples include water
oxidation in photosynthesis, nitrogen fixation and oxygen reduction in many
pathways for respiration. Inorganic chemists often study simple reactions to test
this mechanism, one example being the comproportionation of a Ru(II) aquo and
a Ru(IV) oxo reactants
cis-[(bipy)2 (py) RuIV (O)]2+ + cis-[(bipy)2 (py) RuII (OH2)]2+
2cis- [(bipy)2 (py) RuIII (OH)]2+ — —————————-(1.27)
PCET is also often invoked in electrochemical reactions where reduction is
coupled to protonation or where oxidation is coupled to deprotonation. 40,43
Although it is relatively simple to demonstrate that the electron and proton begin
and end in different orbitals, it is more difficult to prove that they do not move
sequentially. General sequential pathways are lower in energy than concerted
pathways. The main evidence that PCET exists is that a number of reactions
occur faster than expected for the sequential pathways. In the initial electron
transfer (ET) mechanism, the initial redox event has a minimum thermodynamics
barrier associated with the first step. Similarly, the initial proton transfer (PT)
mechanism has a minimum barrier associated with the protons initial PKa.
Variations on these minimum barriers are also considered. The important finding
is that there are a number of reactions with rates greater than these minimum
barriers would permit. This suggests a third mechanism lower in energy; the
15
concerted PCET has been offered as this third mechanism. This assertion has
also been supported by the observation of unsually large kinetic isotope effects
(KIE).
A typical method for establishing PCET pathway is to show that the
individual ET pathways operate at higher activation energy than the concerted
pathway. 40,41 In some literature, the definition of PCET has been extended to
include the sequential mechanisms listed above. This confusion in the definition
of PCET has led to the proposal of alternate names including electron transferproton
transfer (ETPT), electron-proton transfer (EPT), and concerted protonelectron
transfer (CPET).
Also distinct is hydrogen atom transfer (HAT), in which the proton and electron
start in the same orbitals and move together to the final orbital. HAT is
recognized as a radical pathway, although the stoichiometry is similar to that for
PCET.
1.7 Mechanisms of Electron Transfer Reactions
It might be assumed that there would be little to study in the mechanism of
electron transfer; that the reducing agent and the oxidizing agent would simply
bump into each other and electron transfer would take place. Reactions in
solutions are complicated, however, by the fact that metal ions are often
surrounded by shields of ligands and solvating molecules.
The kinetics of electron transfer reactions and their mechanistic importance
revolves around finding answers to the following questions:
(i) What is the stoichiometry of the reaction and the composition of the
activated complex?
(ii) Whether the transfer of electrons, atoms or other species are involved.
(iii) What is the relative rate of electron transfer as compared to the rate of
substitution?
16
(iv) How many electrons are transferred in a single step for multivalent
reactants?
(v) For reactions that are not feasible thermodynamically, what provides
the driving force?
(vi) Are the products isolable and identifiable
(vii) Can intermediate formed before electron transfer be identified?
(viii) What is the importance of acid-base catalysis obtained in the rate law?
(ix) Could it be rationalized in terms of reactants, products or transition
state?
Electron transfer reactions involving transition metal complexes have been
divided into two possible broad mechanistic class called the outer sphere and
inner sphere electron transfers. In this section, these mechanisms and factors
which influence them are examined.
1.7.1 Outer-Sphere Mechanisms
Outer sphere mechanism is a type of reaction whereby bonds are neither
formed nor broken during the electron transfer. 53. For example, in the reaction
below:
[Fe(CN)6]4- + [Mo(CN)8]3- ®[Fe(CN)6]3- +[Mo(CN)8]4-…………………..(1.28)
There is an electron transfer from the reductant to the oxidant, with the
coordination spheres of each remaining intact. Such reaction may be considered
to approximate a simple collision model. The rate of electron transfer for such
reaction is faster than the rate of cyanide substitution for either reactant. So, the
process is considered to consist of electron transfer from one stable complex to
another without the breaking of Fe-CN or Mo-CN bonds. Outer sphere reaction
pathways may be represented stepwise as follows for the reaction between two
metal ions MII and NIII.
(a) formation of a precursor complex
[MII(H2O)6]2+ + NIII (NH3)5L]2+ [(H2O)6MII //NIII(NH3)5L]4+ ……(1.29)
17
(b) Activation of the precursor complex
[(H2O)6MII//NIII (NH3)L]4+ [(H2O)6MII //NIII (NH3)5L4+]# …………(1.30)
(c) Electron transfer and formation of a successor complex (rate determining
step)
[(H2O) 6MII//NIII (NH3)L4+]# ®[(H2O)6MIII //NII (NH3)5L]4+ ……..….. (1.31)
(d) Dissociation of the successor complex to give the final products.
[(H2O)6MIII//NII (NH3)5L]4+ ® [MIII(H2O)6]3+ +[NII (NH3)5L]+ …… (1.32)
It is crucial to note here that, according to the Franck-Condon principle,
the energies of the participating electronic orbitals must be the same for electron
transfer to occur. The little difference in energy observed is as result of
vibrational stretching and compression along the metal-ligand bonds in order to
achieve the required configuration. So, the actual process occurs with the
shortening of the bonds in the MII complex and lengthening of the bonds in NIII.
1.7.2 Inner-Sphere Mechanisms
Inner – sphere reactions are more complicated than outer-sphere reactions
because, in addition to electron transfer, bonds are broken and made 53. A ligand
which bridges two metals is intimately involved in the electron transfer. This type
of mechanism involves penetration into the inner-coordination sphere of reactants
with the formation of a bridged activated intermediate. Substitution occurs at one
of the metal centres to give a ligand-bridged binuclear complex before electron
transfer. The two metal centres participating in the reaction are linked by at least
one bridging ligand common to their inner coordination shells.
The ligand bridge acts as the conducting route for electron transfer from one
metal ion to the other. Dissociation of the activated complex after the electron
transfer produces the products of the reaction.
The classic example of this type of mechanism involved the reduction of
cobalt(III) in [Co(NH3)5Cl]2+ by chromium(II) in [Cr(H2O)6]2+, and it was
specifically chosen because (1) Both Co(III) and Cr(III) form inert complexes
18
and (2) the complexes of Co(II) and Cr(II) are labile. 3,4,16 Under these
circumstances the chlorine atom while remaining firmly attached to the inert
Co(III) ion, can displace a water molecule from the labile Cr(II) complexes to
form a bridged intermediate as shown below:
[Co(NH3)5Cl]2+ + [Cr(H2O)6]2+® [(H3N)5Co-Cl-Cr(H2O)5]4+ + H2O… (1.33)
The redox reaction now takes place within this dinuclear complex with the
formation of reduced Co(II) and oxidized Cr(III). The latter species form an inert
chloroaqua complex, but the cobalt (II) is labile, so that the intermediate
dissociates with the chlorine atom remaining with the chromium.
[(H3N)5Co-Cl–Cr(H2O)5]4+®[(H3N)5Co]2++[(ClCr(H2O)5]2+ ……… (1.34)
The five coordinate cobalt (II) species presumably immediately picks up a water
molecule to fill its sixth coordination position and then hydrolyzes rapidly to
[(H3N)5Co(H2O)]2+. Formally, such an inner sphere reaction consists of the
transfer of a chlorine atom from cobalt to chromium thereby decreasing the
oxidation state of the former but increasing that of the latter. In addition to the
self consistency of chlorochromium complex, further evidence for this
mechanism has been obtained by running the reaction in the presence of free
radioisotopes of chloride ion in the solution. Very little of this labeled chloride is
ever found in the product, indicating that the chloride transfer has indeed been
through the bridge rather than indirectly through free chloride.
The following pathways have been identified in most inner-sphere electron
transfer.1
(a) Formation of collision complex.
[L5MIIIX]2+ + [NII (H2O)6]2+ [L5MIII X //NII (H2O)6]4+………….(1.35)
(b) formation of bridged precursor complex
[L5MIII X //NII (H2O)6]4+ [L5MIII- X – NII (H2O)5]4+ + H2O……(1.36)
(c) Activation of precursor complex, electron transfer and formation of successor
complex.
[L5MIII-X-NII(H2O)5]4+®[L5MII-X-NIII(H2O)5]#…………………….(1.37).
19
(d) Deactivation of successor complex and formation of products.
[L5MII-X-NIII(H2O)5
4+]# ¾¾® [L5MII(H2O)]2+ + [XNIII(H2O)5]2+….(1.38).
Any of the steps in this reaction could be rate determining depending on which
one is the slowest step. If the rate of formation of the precursor complex or the
rate of dissociation of the successor complex is slow, then we are dealing with a
substitution controlled reaction. Alternatively, if the rate of electron transfer is
slow, then we have a redox controlled system.4,16
1.7.3 Distinction between the outer-sphere and inner-sphere reactions
It is generally quite difficult to distinguish between the outer-sphere and
inner-sphere reactions.13 A few more or less clear-cut cases that have been
observed between them by scientists are:
a. Electron transfer by outer sphere mechanism occurs by its tunneling through
space between two coordination spheres, but for an inner-sphere, it occurs
by its tunneling through a common bridging ligand.
b. No specific type of ligand is required for an outer sphere but for an innersphere,
a good bridging ligand is needed for an effective redox reaction.
c. The coordination sphere remains intact for an outer-sphere but in the case
of an inner-sphere, substitution reaction must precede electron transfer.
d. A reaction must be outer-sphere if the rate of electron transfer exceeds that
of ligand substitution, for example, when two inert complexes show a fast
redox reaction with each other.
e. If an inert complex rapidly transfer ligands/atoms to a labile complex the
reaction is very likely to be that of an inner-sphere.
f. Inner-sphere rates are dependent on the nature of the bridging ligand either
kinetically or electronically.
H2O
20
1.8 Determination of the Mechanism of Redox Reaction
One of the aims of an inorganic reaction mechanist is to determine the actual
pathway by which a redox reaction occurs. In order to achieve this objective, the
following modalities are considered.
1.8.1 Identification of Binuclear Intermediate
The detection of a binuclear complex, either as a stable product or as a
transient intermediate along the pathway between reactants and products
represents a piece of experimental information that is taken to be very persuasive
evidence in favour of an inner-sphere mechanism. Until relatively recently, the
binuclear complexes that were detected were successor complexes.18 Such
complexes are expected to be produced when an inner-sphere mechanism is
operative and both the reduced form of the oxidant and the oxidized form of the
reductant are inert with respect to substitution.
Under these circumstances, neither metal centre will “let go” of the bridging
ligand, and a binuclear complex is the final product of the reaction or a relatively
long-lived intermediate. It has been observed that d3 and low-spin d5 and d6
octahedral complexes are inert with respect to substitution and therefore, it is not
surprising that most successor complexes that have been detected so far contain
combination of d3, d5 and d6 octahedral metal centers connected by a suitable
bridging ligand.
An example of a system that features a binuclear successor complex and
which has been studied in considerable detail is the IrCl6
2-–Cr(H2O)6
2+ system.18
The reaction, first studied proceeds in two discernible states. The first is the very
rapid (k>106m-1s-1) disappearance of the reddish brown IrCl6
2- and is
accompanied by the formation of a green intermediate. The second stage involves
the disappearance (k=4.2 x10-2m-1s-1) at 25oC of the green intermediate and the
formation of the final products, olive –brown in colour as shown below
21
[Cr(H2O)6]2+ + [IrCl6]2-®[(H2O)5Cr-Cl-IrCl5]+H2O……………………(1.39)
[(H2O)5Cr-Cl-IrCl5] ¾H¾¾2O®[Cr(H2O)6]3+ + [IrCl6]3+………………..….(1.40)
On the basis of its electronic spectrum, it is evident that the binuclear
complex (H2O)5CrClIrCl5 contains chromium (III) (d3 electronic structure) and
Ir(III) (low-spind6 electronic structure) and is therefore a successor complex. This
system therefore substantiated the inner sphere mechanism.
Although the products of the dissociation complex, Cr(H2O)6
3+ and
IrCl6
3+ which signify outer-sphere are dominant reaction products at significant
amounts of Cr(H2O)5Cl2+ and IrCl5(H2O) (24% yield) are also produced, the
yields increase with increasing temperature and hence reaches a value of 45% at
25oC. This observation therefore, supports an inner sphere mechanism.
The spectrum of the intermediate was recorded and it was found that, in order to
assign molar absorbances that did not vary with temperature to the intermediate,
it was necessary to postulate that the amount of binuclear complex produced was
equal to the yield of Cr(H2O)5Cl2+.
This finding therefore accommodated the fact that the reaction occurs by an inner
sphere mechanism as shown below:
[(H2O)5 CrClIrCl5]¾H¾¾2O®[Cr(H2O)5Cl]2+ + [IrCl5(H2O)]2-…….(1.41)
1.8.2 Reactivity Patterns
(A) Hydroxide versus water
Most electron transfer reactions between aqua complexes exhibits a rate law
consisting of the sum of acid –independent term and an inverse –acid term
Rate = (ko + )
[ + ]
–
H
k I [Ox] [Red] ………………………………………(1.42)
The rate terms are given a mechanistic interpretation to enquire whether the ko
term represent a genuine chemical pathway or is the manifestation of a medium
effect. Thus, acid independent terms are observed for the
22
CO(NH3)5OH2
3+ – Cr(OH2)6
2+ and Fe(OH2)6
3+ – Cr(OH2)6
2+ reactions when the
measurements are carried out utilizing sodium perchlorate to maintain ionic
strength. However, when the background electrolyte is lithium perchlorate, the
acid-independent terms varnish. So, LiClO4 – HClO4 mixtures are preferred over
NaClO4 – HClO4 mixtures when carrying out kinetic studies at varying acidity
and constant ionic strength.
By considering first the inverse –acid path, it is usually interpreted on the
basis of an inner-sphere hydroxide –bridged mechanism.14,15 Direct proof for
such mechanism is lacking in most cases because oxygen tracer studies are
precluded by the lability of reactants and or products.
However, the Co(NH3)5OH2+- Cr(H2O)6
2+ reaction for which trace studies are
feasible, is accompanied by quantitative oxygen transfer from cobalt to
chromium and therefore, show an inner-sphere through the activated complex.
(B) Trends for Halides – Relative Stability of Transition States:
The effects of halide ions on the rates of redox reactions have been
investigated extensively. For historical reasons, the reactivity order I > Br >Cl-
>F- is known as “normal”, whereas the opposite trend is called “inverse”. The
inner-sphere reductions of [Co(OH2)6X]2+ by [Cr(OH2)6]2+ and [Co(CN)6]3-, and
of [Fe(OH2)6X]2+ by [Cr(OH2)6]2+ obey the normal order while the reduction of
[Co(NH3)5X]2+ by [Eu(OH2)6]2+and [Fe(OH2)6]3+ and of [Ru(NH3)5X]2+ by
[Cr(OH2)6]2+ conform to the inverse order 10,14.
For complexes of the form [Co(NH3)5X]2+, (X = Cl, F-, Br-, l- or NO3
-) the
formation of the reductant –X bond in the transition state is of most importance
and the strength of the bond follows the sequence M-F > M-Cl > M-Br > M-I (M
= oxidant or reductant).11 If this complex is reacted with another metal ion, rates
of reaction should be sensitive to the nature of X if the reaction is inner-sphere
whereas for outer-sphere reaction, rates will be unaffected irrespective of the
nature of X .
22
23
1.8.3 Rate of reduction ( kred ) versus rate of substitutation (ksub)
If kred >> ksub, such a reaction is likely to occur by the outer-sphere path.
This was observed for the electron exchange reaction between Fe(CN)6
4- and
Fe(CN)6
3-. Also for the reaction,
[Fe(phen)3]2++[*Fe(phen)3]3+ [Fe(phen)3]3++[*Fe(phen)3]2+ ………(1.43)
ksub was determined to be 7.5 x10-5s-1 (*Fe3+) and 5.0 x 10-5 s-1 (Fe3+) while k for
exchange is 105 mol-1 dm-3 s-1 indicating outer-sphere mechanism.1 For a reaction
in which ksub >> kred, and in the presence of a suitable bridging ligand, innersphere
exchange may occur.
1.8.4 Effect of Added ions:
Substitution of anions into the inner-sphere of labile reactants can alter the
rate of electron transfer greatly. This could be as a result of the formation of
different bridging groups. For an electron transfer reaction that follows the outersphere
mechanism, the absence of bond-making/bond-breaking steps makes the
rate of reaction theoretically easier to be determined.
However, for an outer-sphere reaction the reactants must be in sufficiently close
proximity to create an electron interaction which provides basis for the
delocalization of the exchanging electron. This implies that reactions operating
by the outer-sphere mechanism can be catalyzed in the presence of added ions
that can increase the proximity between the oxidant and reductant thereby
shortening the distance within which the electron can be transferred.8,15,16
However, for redox partners that carry opposite charges, added ions could
retard the rate of reaction since coordination to any of the reactants could reduce
the degree of attraction between the reactants. This will increase the distance
between the redox partners and slow down the rate of electron transfer.
Activated complex of the form [(CH3N)5Co-X-Cr(OH2)4Y]# has been
suggested for the reaction between (H3N)5CoIII X and Cr(II) where Y is an added
anion (catalyst). This shows that the anion affects largely the reactivity of the
24
reducing agent and usually appears in the Cr(III) product for both inner and
outer sphere reactions. A typical rate law for the effect of added anion is given
by;
Rate = (ko + k1 (external ion) [oxidant] [reductant] ——————— (1.44)
Where k0 = rate constant for independent of rate on external ion effect.
k1 = rate constant for dedepdent of rate on external ion effect.
1.8.5 Activation Parameters
Activation parameters DH # , DG# and DS # do not seem to have strong
mutual relationship with the type of mechanism operating in a particular redox
process. However, their signs or magnitude could give a clue as to which
mechanism is existing in a reaction. Negative DH # indicates formation of a
precursor complex as in an inner-sphere mechanism.16 For example, despite the
difference in mechanisms; the DS # for the reaction of Cr2+ and V2+ with Ru3+
complexes are almost the same.
1.8.6 Product Identification
Inner-sphere mechanism can be ascertained without ambiguity if the
oxidant and the reducing agent are both substitution inert, and where atom
transfer occurs during redox reaction.
The transferred atom or group is usually the bridging ligand. This is equally
about the single most conclusive evidence that demonstrates the operation of a
bridged complex in the course of the reaction.16
A lot of work has been done with Co(III) and Cr(II) complexes as oxidant
and reductant respectively. It was reported that CrCl2+ was formed as one of the
products which resulted from the formation of the binuclear intermediate,
[(NH3)5–Co-Cl-CrCl5]4+. This shows the inner-sphere nature of the reaction.
However, there are inner-sphere reaction which are not accompanied by atom
25
[(H Co Y Fe (OH2)5] 3N)5
(n+2)+
transfer. For example reductants like Fe2+, V2+, Eu2+ and in such reactions
where easily hydrolysable products are formed, identification of products is
difficult.
Such a situation has been observed in a case like Co(NH3)5SCN2+/V2+ system
where stopped-flow technique has measurements of the volume of activation (
DV # ) for the reduction of various complexes has been applied as diagnostic tool
in reaction kinetic. It has been reported that for the reaction:
[ ] + [ ] + + 2
(H3N)5Co Y Fe (OH2 )6 III n II
I.S O.S
Inner-Sphere (I.S) pathway should be retarded with increasing pressure (volume
of activitation DV # should be positive) if it is assumed that the volume of “free”
H2O is larger than that of coordinated H2O. Obtained results support an innersphere
mechanism.16 However the same trend has not been obtained in some
other redox systems making the application of DV # as a diagnostic tool of limited
scope.21
1.8.7 Michaelis-Menten Plots
For the enzymatic action of the form
[ ] [ o ]
obs k E
dt
d product = ——————————————————– (1.46)
[ ]
k [S]
k S
k
m
obs +
= 1 —————————————————————- (1.47)
[(H3N)5 Co (Y) (H2O) Fe (OH2)5] (n+2)+
……………….(1.45)
26
Michaelis – Menten observed that equation (1.48) can be arranged to give
[S]
k
k
k k
m
obs
÷ ÷ø
ö
ç çè
æ
= +
1 1
1 1 ——————————————- (1.48)
where kobs is the rate constant for the overall reaction and Eo is total
enzyme concentration, km represents Michaelis – Menten rate constant. k1 is rate
constant for the break up of an active intermediate into products while S is the
substrate. For a normal redox reaction, S represents the reductant. Equation 1.46)
shows that a plot of 1/kobs versus 1/[S] gives 1/k1 as intercept.
If however, a linear plot which passes through the origin is obtained, it
shows that the value of 1/k1 is zero. This means that there is no equilibrium
constant for the intermediate in the reaction and so, the reaction is said to follow
outer-sphere pathway. If however a linear plot with intercept 1/k1 is obtained,
then it means that the intermediate in the reaction has an equilibrium constant and
it is said to undergo inner-sphere pathway. These two conditions will always give
positive slope. But there is another condition whereby a negative slope with
interecept is obtained. Under that condition, Michaelis-Menten equation in
equation 1.49 is rearranged to give equation 1.50 known as Eadie Hofstee
equation. 58,59
[ ]
K [S]
V S
m +
n = max ——————————————- (1.49)
[ ] max V
S
Km = – +
n
n ————————————– (1.50)
Where v represents reaction rate, Km is the Michaelis-Menten constant, [S]
is the substrate concentration, and Vmax is the maximum reaction rate. From
equation (1.49), by inverting and multiplying with Vmax we have that
( [ ])
[ ]
[ ]
[S]
K S
V S
V V Km S m +
=
+
=
max
max max
n
————————– (1.51)
Rearangeing equation (1.52) gives
Vmax = [ ]
[ ]
[ ] [ ] n
n n n
+ = +
S
K
S
S
S
K m m —————————– (1.52)
27
Isolating v from equation 1.52 then gives the Eadie-Hofstee equation
shown below:
[ ] max V
S
= -Km +
n
n
So a plot of v against v/[s] will hence yield vmax as the y-intercept, vmax/Km
as the x-intercept, and Km as the negative slope.
1.9 Objectives of the Project
Iron complexes and other transition metal complexes occupy important
positions in chemistry, biochemistry and chemical technology. For instance, in
haemoglobin, iron serves as an oxygen carrier in the blood of mammals, birds
and fish.18 On the other hand, the thiols under study have numerous industrial
applications. For example, the L-cysteine derivative, N-acetyl cysteine (NAC) is
often used in cough medicine as it breaks up the disulphide bonds in the mucus
and thus liquefies it, making it to be coughed up.15 Thioglycolic acid is used as a
chemical depilatory and for detection of iron, molybdenum, silver and tin.
Thiourea is used in textile processing and in the reductive work up of ozonolysis
to give carbonyl compound.20,21 2-mercaptobenzothiazole which is a toxic
pollutant is used as a corrosion inhibitor in petroleum products.22 while
benzylmercaptan which is found in most transformer oil has been implicated as
the main cause of transformer failure due to it’s corrosive sulphur attack on the
metal surfaces.23 In this research, we investigated the dynamics of the electron
transfer reaction of μ-adipato bridged iron (III) complex; μ-adipato-di[N,N/-
bis(salicylideneethylenediaminatoiron(III)] [(Fe-salen)2adi], hereafter also
denoted as [Fe2adi] with thiols (L-cysteine, thiourea, thioglycolic acid, 2-
mercaptobenzothiazole and benzylmercaptan)
Based on the above information, the objectives of this study are :
27
28
a. To investigate the redox reactions of these thiols with m -adipato bridged
iron(III) complex due to the role played by RSH/RSSR couple in
mediating redox potentials at biological sites.
b. To investigate the possibilities of using the iron(III) complexes of the
thiols under study in metal chelation therapy.
c. To determine the dynamics of electron transfer of thiols under study and
the possibility of using them effectively as reductants for some toxic metal
ions.
d. To determine whether the redox reactions of m -adipato bridged iron(III)
complex with the thiols follow outer-sphere or inner sphere mechanism
and to generate rate equations for the various reactions.
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