ABSTRACT
The kinetics and mechanism of the electron transfer reactions of a
mixed-valence manganese complex (MnIII O2 MnIV) with four different
reductants ( Ascorbic acid (H2A), 1, 3 dihroxybenzene (H2R), SO3
2-and
S2O3
2-) have been studied in acid medium at 28 ± 10C.The reactions were
found to obey second order kinetics with the following general rate law
dt
[Mn O Mn ]
d
2
3 IV
2
III
= k2 [MnIIIO2MnIV] [Reductant].
The experimental data for three of the reaction systems were found to
exhibit first order dependence on acid concentration with this general rate
law.
dt
[Mn O Mn ]
d
2
3 IV
2
III
= (a +b [H+] ) [MnIIIO2MnIV] [Reductant].
For MnIIIO2MnIV / H2A system
a = 3.50 x 10-4 dm3 mol-1 s-1 ,b = 2.13 x 10-3 dm-3 mol-1 s-1 at constant
[H+] 0.50 mol dm-3, I = 0.5 mol dm3 (NaCl), T =30oC and λ = 580nm.
For MnIIIO2MnIV / H2R system
a = 0.70 x 10-3 dm3 mol-1 s-1 ,b = 0.02 x 10-2 dm6 mol-2 s-1 at [H+] =
0.50 mol dm-3, I= 1.0 mol dm-3 (NaCl), T= 28 ± 1oC and λmax = 580nm..
For MnIIIO2MnIV / SO3
2- system
a = 1.22 x 10-6 dm3 mol-1 s-1 ,b = 7.83 x 10-7 dm6 mol-2 s-1at [H+] 5 x 10-2 mol
dm-3, I= 0.5 mol dm-3 (NaCl), T= 28 ± 1oC and λmax = 580nm .
while the rate law for the MnIIIO2MnIV / S2O3
2- system is
dt
[Mn O Mn ]
d
8
3 IV
2
III
= (a +b [H+] ) [MnIIIO2MnIV] [S2O3
2-]
where a = 7.01 x 10-3 dm3 mol-1 s-1 and b = 0.92 x 10-3 dm6 mol-2 s-1at 28 ±
1oC [H+] = 2 x 10-3 mol dm-3 ,I= 0.5 mol dm-3 (NaNO3).
– 8 –
There were no spectroscopic and kinetic evidences for the formation
of an intermediate complex in any of the reaction systems. This is not in
favour of inner sphere mechanism. Therefore, the electron transfer must
probably occur by the outer sphere mechanism in these systems.
In MnIIIO2MnIV/H2A and MnIIIO2MnIV/SO3
2- systems, the reactions
were catalysed by cation and anion species .This is in support of the outer
sphere mechanism but in the MnIIIO2MnIV/H2R and MnIIIO2MnIV/S2O3
2-
systems, no catalysis was observed and free radicals were also not important
in the redox processes. Therefore, the results obtained are in support of
proton coupled electron transfer (PCET) mechanism and is hereby proposed
for the two reaction systems.
The kinetics and mechanism of the electron transfer reaction of
rosaniline hydrochloride (referred to as Ros) with reductants (OH-, NO3
-,
IO4
-, SO3
2- and S2O6
2-) have also been studied in aqueous medium at 30oC, I
=1.0 mol dm-3(NaCl, CH3COONa or NaClO4), [H+] = 1 x 10-4 mol dm-3
except for OH- system that does not involve [H+].
The stoichiometry is 1:1 in all the five systems investigated. The
reaction is first order in both the [oxidant] and [reductant] for OH-, NO3
– and
S2O6
2- systems respectively therefore the overall order for the system is
second order with the following rate law
-d/dt[Ros] = k2[Ros][reductant]
In the IO4
– and SO3
2- systems, the reaction is first order in the
[oxidant] and zero order in the [reductants] hence the overall order is first
order with the rate law
-d/dt[Ros] = ko [Ros]
– 9 –
The rates of the redox reactions showed direct dependence on acid
concentrations for Ros/ NO3
– system and an inverse dependence on acid
concentration for Ros/SO3
2- and Ros/S2O6
2- systems while in Ros/IO4
–
system, acid concentration have no effect on the rate of the reaction. The
overall rate equation for the reactions can be given as
-d/dt[Ros] = (a+b [H+] ) [Ros][NO3
-] for Ros/NO3
-system
-d/dt[Ros] = (a+b [H+]-1 ) [Ros][reductant] for Ros/SO3
2- and Ros/S2O6
2-
systems
The rate of the reaction displayed negative salt effect for Ros /OHsystem
and a positive salt effect for Ros /NO3
– system and the rate also was
sensitive to changes in the di-electric constants.
Spectroscopic investigation showed the presence of short-lived
intermediate complex formation in the Ros/IO4
-and Ros/ SO3
2- systems. This
suggested that both systems are proposed to take place by inner sphere
mechanism. In Ros/ OH-, Ros/NO3
-, and Ros/S2O6
2- systems, there was no
evidence for the formation of an intermediate complex of significant
stability and free radicals are absent therefore the mechanisms of their
reactions are discussed in terms of outer sphere mechanism.
On the basis of the experimental results obtained, the outer sphere
mechanism has been proposed for Ros/ OH-, Ros/NO3
– and Ros/S2O6
2-
systems respectively and inner sphere mechanism for Ros/IO4
– and Ros/SO3
2- systems.
– 10 –
TABLE OF CONTENTS
Title page ….i
Declaration ….ii
Certification ….iii
Dedication ….iv
Acknowledgment ….v
Abstract ….vii
Table of Contents ….x
List of Tables ….xiv
List of Figures ….xvi
List of Appendices ….xxi
Abbreviations
….xxii
Chapter One.
1.0 Introduction. ….1
1.1 Oxidation – Reduction (redox) reaction. ….3
1.1:1 Electron transfer reactions. ….3
1.1.1:1 Homonuclear (or Isotopic Exchanges) reactions ….4
1.1.1:2 Heteronuclear electron exchange or cross-reactions ….5
1.2 Theories of electron transfer. ….6
1.2:1 Franck-Condon principle. ….6
1.2:2 Electron tunnelling theory ….7
1.2:3 Marcus theory. ….9
1.3 Importance of electron transfer reactions. ….15
1.4 Mechanism of electron transfer reactions. ….15
1.4:1 Inner sphere mechanism. ….16
1.4:2 Outer sphere mechanism. ….23
1.4:2.1 Proton coupled electron transfer (PCET) mechanism ….26
1.4:3 Solvated electron theory. ….28
1.5 Factors affecting outer-sphere reactivity ….29
1.5.1 Intramolecular vibration trapping ….29
1.5.3 Solvent trapping ….30
1.5.4 Electron coupling. ….31
1.6 Criteria for assigning the mechanism of electron
transfer reactions ….32
1.6.1 Rate of redox process versus rate of substitution . ….33
1.6.2 Reactivity pattern. ….34
1.6.3 Identification of binuclear intermediates. ….36
– 11 –
1.6.4 Product analysis ….37
1.6.5 Activation parameters. ….38
1.6.6 Marcus theory correlation ….39
1.7 Intervalence transfer (I.T ) ….39
1.7.1 Optical electron transfer process. ….39
1.7.2 Thermal electron transfer process. ….40
1.8 Mixed valence compounds ….42
1.9 Photosynthesis process. ….46
1.9.1 Photosynthetic reaction centre ….46
1.9.2 Photosystem II ….49
1.10. Dyes ….53
1.10.1. Chemistry of dyes (Artificial Dyes). ….53
1.10.2 . Classification of dyes. ….55
1.10.2.1. Natural dyes. ….55
1.10.2.2 Synthetic dyes. ….55
1.10.2.3 Reactive dyes. ….55
1.10.2.4. Fluorescent brightness dyes. ….56
1.10.2.5 Disperse dyes. ….56
1.10.2.6 Direct dyes ….56
1.10.2.7 Vat dyes ….56
1.10.2.8. Sulphur dyes ….56
1.10.2.9 Acid dyes ….57
1.10.2.10 Basic dyes/fuchsin. ….57
1.10.3 Uses of dyes as stains ….60
1..11 Objective of this research work. ….62
Chapter Two.
2.0 Literature review ….64
2.1 Kinetics and mechanism of oxy-anion reaction ….64
2.2 Redox reactions of sulphite ions ….66
2.3 Redox reactions of thiosulphate ions ….72
2.4 Redox reactions of periodate ions ….73
2.5 Redox reactions of 1,3 dihydroxybenzene ….75
2.6 Redox reactions of L-ascorbic acid ….77
2.6.1 Oxidation of ascorbic acid with group 5 complexes ….78
2.6.2 Oxidation of ascorbic acid with group 6 complexes
….79
2.6.3 Oxidation of ascorbic acid with group 7 complexes
….79
– 12 –
2.6.4 Oxidation of ascorbic acid with group 8 complexes
….80
2.6.5 Oxidation of ascorbic acid with group 9 complexes
….82
2.6.6 Oxidation of ascorbic acid by group 10 complexes
….83
2.6.7 Oxidation of ascorbic acid by group 11 complexes ….85
2.6.8 Oxidation of ascorbic acid by lanthanides and actinides .…85
2.7 Reactions of basic fuschin.
….86
2.8 Reduction of dioxo- bridge manganese complexes.
….88
Chapter Three
3.0. Experimental.
….91
3.1 Materials and reagents
….91
3.2 Preparation of reagents.
….91
3.2.1 Rosaniline hydrochloride.
….91
3.2.2 Perchloric acid solution.
….91
3.2.3 Standard sodium perchlorate solution
….93
3.2.4 Standard sodium periodates solution
….93
3.2.5 Standard sodium sulphite solution.
….93
3.2.6 Standard sodium dithionate solution
….94
3.2.7 Standard hydrochloric acid and nitric acid solution
….94
3.2.8 Standard solution of sodium thiosulphate.
….94
3.2.9 Manganese sulphate
….95
– 13 –
3.2.10 Potassium chloride solution
….95
3.2.11 Sodium acetate solution
….95
3.2.12 Sodium nitrate solutions
….95
3.2.13 Potassium iodate solution
….96
3.2.14 Standard L-ascorbic acid solution.
….96
3.2.15 Synthesis and characterisation of di-μ – oxo bridged
phenanthroline complex of manganese
[(Phen)2 MnIIIO2MnIV(Phen)2] (ClO4)3 .
…..96
3.3 Stoichiometric studies.
….97
3.4 Kinetic measurements.
….110
3.5 Test for intermediate complex
….110
3.5.1 Spectrophotometric test
….110
3.5.2 Test for free radical.
….111
3.6. Product analysis.
….111
Chapter Four.
4.0 Results.
….112
4.1 Stoichiometry
….112
– 14 –
4.2. Order of reaction
….114
4.3. Effect of hydrogen ion concentration
on the rate of reaction
….145
4.4. Effect of ionic strength.`
….154
4.5. Effect of added cations and anions
on the rate of reactions.
….156
4.6 Effect of di electric constant.
….171
4.7 Test for intermediate complex
….171
4.7.1. Spectrophotometric tests.
….171
4.7.2. Michael – Mentens plots.
….172
4.8 Test for free radicals
….183
4.9 Product analysis.
….183
Chapter Five
….185
5.0 Discussion.
….185
5.1. Rosaniline hydrochloride – hydroxyl ion reactions
….185
5.2. Rosaniline hydrochloride – periodate ion reactions
….187
5.3. Rosaniline hydrochloride –sulphite ion reactions
….190
5.4. Rosaniline hydrochloride –dithionate ion reactions
….195
5.5. Rosaniline hydrochloride –nitrate ion reactions.
….196
5.6. MnIIIO2MnIV – L-ascorbic acid reaction
….198
– 15 –
5.7. MnIIIO2MnIV – 1, 3 – dihydroxybenzene reaction
….201
5.8. MnIIIO2MnIV –thiosulphite ion reaction.
….204
5.9. MnIIIO2MnIV –sulphite ion reaction
….207
5.10. Summary and conclusion
….209
References
….213
Appendix
….231
– 16 –
CHAPTER ONE
1.0 INTRODUCTION
After a period of quiescence in the early part of the 20th century, inorganic chemistry has again become an exciting area of research.
The most interesting developments in inorganic chemistry bridge the gap with other disciplines. These include organometallic chemistry with
bias towards catalysis, coordination chemistry and bioorganic chemistry (Purcell and Kotz, 1977)
Many organometallic compounds play important roles in industrial chemistry as catalysts. Some success has been achieved in the
use of such catalysts for converting natural gas to related but more useful chemical substances.
Chemists have also synthesized large inorganic molecules that contain a core of metal atoms, such as manganese, surrounded by a
shell of different chemical units. Some of these compounds, referred to as metal clusters, have characteristics of metals, while others react in
ways similar to biological systems. Although organic molecules were once thought to be the distinguishing chemical feature of living creatures,
it is now known that inorganic chemistry plays a vital role as well. Trace amounts of metals such as manganese in biological systems are
essential for processes such as photosynthesis, respiration, nerve function, and cell metabolism. Processes of this kind form the object of study
of bioinorganic chemistry.(Cotzias,1962)
A great deal of work in inorganic reaction mechanism came from the research efforts of Taube and co-workers (Taube et
al; 1953; Taube and Meyers 1954). They started with the study of reduction of substitutionally inert amine complexes of CoIII by
substitutionally labile Cr(H2O)6
2- and established the ultimate details of how electron transfer reactions occur. (Taube et al 1953;
Taube, 1970). Using isotopic labelling techniques, they were able to suggest that the reaction can be represented by equation (1.1)
[Cr(H2O)6]2++[Co(NH3)5Cl]2++ 5H+ + 5H2O [Co(H2O)6]2+ + [Cr(H2O)5Cl]2+ + 5NH4
+…(1.1)
Significant advances have been recorded in the synthesis and characterization of transition metal complexes of mono-, diand
poly nuclear centers. (Arulsamy et al, 1994; Cooper and Calvin, 1977; Dave and Czernuszewicz; 1994 (a and b)). The reactions
of these complexes are being studied continuously, particularly the electron transfer processes (Monzyk and Holwerda, 1992. Ghosh et
al, 1994).
These electron transfer reactions gave insight into the actual process of transfer. They also have found application in
chemical synthesis, histological techniques, biological systems and electron transfer catalysis. (Bugress, 1978, Wilkin, 1974)
Various reactions in inorganic and biological systems (biological reactions
that involve organic catalysts or enzymes) involve transfer of electrons at one
stage or the other. In addition, the need for industries to have a new set of stable
polymers particularly those that could withstand extremes in thermal and physical
stress aroused great interest in inorganic polymers with skeletons consisting of
boron, aluminium silicon and phosphorus atoms. The formation of such inorganic
polymers demands good knowledge of inorganic reactions kinetics. (Armtage,
1972)
– 28 –
Iyun in 1982 also highlighted the fact that the knowledge gained from these
reactions will constitute an inevitable prerequisite to the understanding, development and
eventual effective control of a wide area of science and technology.
1.1 OXIDATION – REDUCTION (REDOX) REACTIONS
Redox reactions are among the most common types of reactions in chemical and biological processes (Yoshihana et al, 1995).
They are usually spontaneous and often accompanied by changes in oxidation state of at least two of the reactants. (Purcell and
Kotz, 1977). These reactions are basically of two types;
i) Reactions involving electron transfer (ET) which play important roles in chemical,
biological and technological processes.
ii) Reactions involving atom transfer with or without electron transfer.
The simplest of all these redox reactions are those that involve the transfer
of electrons. An example is the reaction between MnO4
2– and MnO4
–
which was
studied by the use of an isotopic tracer (Horning et al, 1950) and is represented by
equation (1.2)
*MnO4
– + MnO4
2– *MnO4
2– + MnO4
– …(1.2)
1.1.1 Electron Transfer Reactions
Electron transfer reaction mechanisms are inherently of interest because of
the insight they give into the actual process of transfer. (Sheldon and Kochi, 1981;
Meyer and Taube, 1987).
Electron transfer is very important in different types of molecular and
bimolecular systems. It is also important in polymerization reaction, photography,
electrochemistry, photosynthesis, metabolism and many other processes of
individual and industrial applications. (Yoshihana et al, 1995). Electron transfer
reactions can be classified in two main types on the basis of thermodynamic
parameters. These are homonuclear electron exchange reactions and heteronuclear
electron exchange reactions (Lawal, 1997)
1.1.1.1 Homonuclear (or Isotopic Exchange) Reactions
In homonuclear electron exchange reactions, electron transfer takes place
between two ions of the same element existing in different oxidation states.
– 29 –
Mn(III) + *Mn(II) Mn(II) + *Mn(III) …(1.3)
[*Co(NH3)6]2++ [Co(NH3)6]3+ *[Co(NH3)6]3+ + [Co(NH3)6]2+ …(1.4)
[*Fe(phen)3]2+ + [Fe(phen)3]3+ *[Fe(phen)3]3+ + [Fe(phen)3]2+ …(1.5)
*Cr2+ + CrCl2
2+ *CrCl2
2+ + Cr2+ …(1.6)
where * is an isotopically labelled species.
In all of the above reactions, there is no net chemical change and the rate
constant for the forward and backward reactions are equal. The reactant and
product concentrations are the same. Therefore the equilibrium constant =1. The
only change in free energy is that due to mixing and therefore the overall free
energy change ΔG ≈ 0
1.1:1.2 Heteronuclear Electron Exchange or Cross Reactions
These reactions involve transfer of electrons between different metal ion
centres and the products are chemically distinct from the reactants. The rate can be
measured by chemical methods. Examples are equations (1.7–1.11).
Mn(VII) + 4Mn(II) 5Mn(III) …(1.7)
[Fe(CN)6]4 – + [IrCl]2- [Fe(CN)6]3- + [Ir(Cl)6]3- …(1.8)
Sn(II) + Ti(III) Sn(IV) +Ti(I) …(1.9)
[Co(en)3]3+ + [Ru(NH3)6]2+ [Co(en)3]2- + [Ru(NH3)6]3+ …(1.10)
– 30 –
Sn(II) + 2Fe(III) Sn(IV) + 2Fe(II) …(1.11)
In equations (1.7) and (1.8), the net change in free energy in most cases is
less than zero (G < 0).
In equations (1.8) and (1.10), the stoichiometry is 1:1 and the reactions is
complementary in which case the oxidant and reductant undergo equal changes in
oxidation states. However in equation (1.11), the stoichiometry of these types of
reactions is usually not 1:1 and they are non-complimentary in that the oxidant and
reductant undergo unequal changes in oxidation states. These reactions certainly
involve the formation of reactive intermediates in which unstable oxidation states
of the metal are formed.
1.2 THEORIES OF ELECTRON TRANSFER
1.2.1 Franck – Condon Principle
Franck – Condon principle state that “the motion of nuclei is so slow
(10-13 s-1) when compared to that of electrons (10-15 s-1) and that electron transfer
occurs without an appreciable movement of the nuclei” (Platzmann and Franck,
1954; Sutin, 1966). The principle shows that the position of the nuclei remains
virtually frozen (intact) during the process of electron transfer. There are two
important consequences of the electron transfer processes from the Franck –
Condon principle.
Firstly, no angular momentum can be transferred to or from the transition
state during the act of electron transfer and a restriction is also imposed on the
change in spin angular momentum.
Secondly, the oxidant and the reductant must undergo reorganization before
electron transfer in such a way that their energies in the transition state become
identical thus minimizing the energy change on electron transfer.
– 31 –
The total change in free energy involved in the process can then be
represented by equation (1:12) (Marcus 1956)
G# Gt
# + Gi
# + Go
# …(1.12)
Where
Gt
# = Activation free energy
Gi
# = Inner-sphere reorganization energy
G#
o = Outer-sphere reorganization energy
1.2.2. Electron Tunnelling Theory
Tunneling is a term often associated with electron transfer, which involves
transfer of a particle between electronically coupled chemical sites through a
tunneling process. In that sense, every electron transfer process involves electron
tunneling with tunneling frequency given in the classical limit by equation. (1.13).
2
1
2 2
e λRT
π
h
2πV
V = …(1.13)
Ve = Electron transfer frequency
V = Electronic coupling
h = Boltzman constant
T = Absolute temperature
R = Gas constant
The electron tunneling theory was developed by some workers (Weiss 1954;
Marcus et al, 1954). They concluded that the electronic energy in the reactants and
products is not as high as would ordinarily be expected from the classical point of
view. Under this condition, electron transfer process was then viewed as a
tunneling process in which electron passes through the potential energy barrier
rather than over it (Basolo and Pearson, 1967). The outcome of the tunneling
process is that electron can be moved to a longer distance beyond the distance of
– 32 –
actual collision of the reactants. The potential energy barrier leakage is illustrated
in Figure- 1.1
1 Electron coordinate
Electron Coordinate
Fig.1.1 ELECTRON TRANSFER BY PENETRATION OF A POTENTIAL ENERGY BARRIER.
U1 and U2 = Ground states of electron in cation 1and 2 respectively.
Uo
w d
U1 U2
Energy
– 33 –
d = Width of barrier at height of penetration
w = Kinetic energy of the electron.
Uo = Height of the barrier
A theoretical basis was developed to explain this electron transmission
model (Marcus et al, 1954; Marcus, 1956) and their results were expressed in
terms of transition state theory of chemical kinetics and could be written in the
form
k =
RT
ΔG
RT
ΔG
k exp
h
K *
e
*r
T / …(1.14)
KT = Equilibrium constant at constant temperature.
k = Rate constant
k/ = Electron transmission co-efficient
h = Boltzman constant
T = Absolute temperature
Ge
* = Activation energy
Gr
* = Hydration energy for inner co-ordination shell arrangement
R = Gas constant
As the exchanging partners come close, transmission coefficient increases,
while the energy of activation also increases as a result of electrostatic repulsion
between the reactants. Both effects tend to decrease the rate constant. The rate at
an optimum distance between the reactants, a maximum exchange rate is obtained.
Taube (1959) and Lewis (1980) then suggested that electron tunnelling is probably
involved in most electron transfer processes and that the tunnelling step might not
necessarily be the rate-determining step.
( )
– 34 –
1.2.3. Marcus Theory
The approach is to calculate the rate of an outer-sphere electron transfer
reaction from the first principle. In the outer-sphere mechanism, the weak
interaction between reactants during electron transfer, enables theoretical
treatment and correlations between the kinetic and the overall thermodynamic
parameters. (Sutin 1966, Diebler and Sutin, 1964). The most widely used of such
treatment was developed by Marcus (Marcus, 1956) and is commonly known as
Marcus theory.
Marcus theory was initially used for the calculation of absolute rate constants
for homonuclear exchange reactions which was later extended to cross reactions
and related processes (Marcus, 1956, 1957 a–c, 1963, 1965 a–c, 1968 a and b
1977). He assumed in his theory that when the work term is small, then little
reorganization is involved prior to electron transfer. Therefore:
i. There is small electronic interaction between the reacting species (which
are treated as rigid spheres of radii a1 and a2) and that there is no change
of inter-atomic distance between the species during the electron transfer
reaction.
ii. The probability of electron transfer within the activated complex is
unity.
iii. The work terms of the self–exchanging and cross-reaction are the same.
iv. The motion of the inner coordination spheres are harmonic with
breathing forces having the reduced value p
r i
p
i r
f + f
2f f
where fi and f p
i are force constants for the symmetrical breathing
vibration of the species in the reactant and product respectively.
It is necessary to estimate the contribution of G* made by the various steps
by which the reaction is thought to occur. Those terms include the free energy
– 35 –
required to bring reactants to within reactant distance and to re-organize bond
distance in each reactant so as to bring each to a common state prior to electron
transfer and finally the free- energy change of the net reaction, G*.
The free energy G* can therefore be represented by equation
G* =
( ) ( )
4λ
ΔG w w
+
2
ΔG w w
+
4
λ
*p p r *p p r 2
…(1.15)
where λ = λ p + λ i … (1.16)
λ p and λ i are given as follows
λ p = (ne)2 ( ) ( )
D
1
D
1
r
1
2a
1
+
2a
1
1 2 op
… (1.17)
λ I = p
i i
p
i i
f + f
f f
( Δ a1)2 + p
2 2
p
2 2
f + f
f f
( Δ a2)2 …(1:18)
where
wr and wp –are work terms required to bring the reactants to their mean
separation distance (a1 + a2) and then remove the products to infinity.
λ p and λ i –are the free energies required to reorganize the solvent molecules
around the reactants (the outer-sphere coordination shell) and to reorganize the
inner coordination shell of the reactants.
G* – is the standard free energy of the reaction at the separation distance.
n – is the number of electron transferred which is based on the assumption that
( ) p
1 1
a
2 1
1
1 2
1 2 = a a and a = a a
2
a + a
is the difference in the radius of the
specie i when it is a reactant and its radius when it is a product.
Dop and Dr are the square of the refractive index and dielectric constant for the
medium respectively.
– 36 –
The free energy G*is related to the free energy of activation as follows.
kT
hZ
ΔG* = ΔG# + RT ln …(1:19)
1
or ΔG* = ΔG# 2.444kcalmol …(1:20)
So that the rate constant becomes
RT
ΔG
k = Z exp
*
…(1.21)
where Z = 1011 dm3 mol-1s–1 ( for homogeneous reaction).
Iyun in 1982 reported that from the above equation the activation
parameters for electron transfer reactions can be predicted. Therefore equation
(1.15) can be re written as
( )
2
1 + α
+ ΔG
2
ΔG + ΔG
ΔG = 12
*
22
*
* 11 …(1.22)
where *
12
*
22
*
11 ΔG ,ΔG and ΔG are the free energy for the exchange reactions
respectively.
This relationship can also be written for entropy and enthalpy.
1 2 …(1.24)
2
H
1 4
2
H H
H
1 …(1.23)
2
1 4 S
2
S S S
*
2 12
*
22
*
* 11
*
2 12
*
22
*
* 11
From the expressions in equations (1:20)–(1:23) above, the activation
parameters can be calculated using the appropriate equations below
kT
hZ
ΔG* = ΔG# + RT ln …(1:25)
10.2256 Jmol-1K-1
– 37 –
2RT … (1: 27)
H H 1
2R … (1: 26)
1
kT
S S RT ln hZ
* #
* #
For a cross reaction (1:28), and (1:29). Marcus theory predicts that the rate of the
electron exchange is given by equation (1.31)
OX1 + Red2 Red1 + OX2 …. (1.28)
*OX1 + Red1 *Red1 + OX1 …(1.29)
*OX2 + Red2 *Red2 + OX2 …(1.30)
2
1
12 11 22 12k k k K f …(1.31)
where
2
11 22
12
Z
k k
4log
log k
log f …(1.32)
12 12 k and K are the rate constant and equilibrium constant respectively for the cross
reaction 11 22 k and K are appropriate rate constants for the isotopic exchange
reactions. Z is the collision frequency for the hypothetically unchanged reaction
ions.
However, the theoretical rate constants for electron transfer reactions can be
calculated from equation (1:33) below by considering the solvent reorganization
occurring outside the inner coordination shells
– 38 –
k = Z exp (
RT
ΔG#
) of each of the reactant …(1:33) G#
= Change in free energy of cross reaction
R = Gas constant
T = Absolute temperature
A reasonable agreement between the rate calculated and rate observed is taken
as evidence that the reaction is of the outer-sphere type. This relation has been
used to:-
i) Affirm that the experimental and calculated values of k12 are in accord;
when the values of k11 k12 K12 are known; compilation of observed and
calculated rate constants has been documented (Pennington, 1978) and
comparing the different oxidants with the same series of reductant.
ii) Check for gross deviations from the theory, which signal a change in
mechanism and can be used to estimate the self-exchange rate constant
for one reactant when the value is unknown, based on measured values
for one or more cross reactions. (Pladziwicz and Espenson, 1973).
1.3 IMPORTANCE OF ELECTRON TRANSFER REACTIONS
Electron transfer is important in reactions of metal ion complexes. They
often involve ligand substitution or electron transfer or both. In some redox
reactions, ligand substitution occurs and this provides a low energy pathway for
electron transfer as shown in the oxidation of both metal ion and non- metallic
substrates by chromium (VI) (Beatle and Haight, 1972).
Various reactions in organic and biological systems involve the transfer of
electrons at one stage or the other especially in biological redox reactions
involving organic catalysts
– 39 –
More recently, it has become increasingly obvious that many reactions in
inorganic chemistry once thought to be concerted in nature also occur via
sequential one electron step. (Thompson and Meyer, 1982)
In view of the obvious importance of electron transfer reactions; further
investigations on these reactions are necessary. Since the main approach to
reaction mechanism is the kinetic study, it is necessary that emphasis be placed on
rate measurements and their interpretations.
1.4 MECHANISM OF ELECTRON TRANSFER REACTIONS
The mechanism of a reaction is the detailed stepwise process involving
molecules, atoms, radicals or ions that occur simultaneously or consecutively and
culminates in the observed overall reaction. Therefore in studying a redox reaction
mechanism, one needs to investigate the following parameters and apply it
appropriately.
The first set of parameters to be investigated are the stoichiometry of the
reaction and the nature of the activated complex, whether the given chemical
reaction occurs in a single molecular process or several processes. If the latter
occurs, one should consider whether the molecular steps occur as concurrent
alternatives or in succession along a single pathway.
In addition, the stereochemical arrangement of the activated complex, the
state of maximum potential energy through which the system passes in its key
phase is the rate determining step and the presence or absence of intermediate(s) in
the reaction must also be considered.
The effects of acid or base (pH), anions and or cations, dielectric constant,
ionic strength and activation parameters Ea, Δ H, Δ S e.t.c on the rate of the
reaction must be considered. However, when the reaction is accompanied by
transfer of electrons or atoms, the nature of all these parameters can be
rationalized in terms of two distinct mechanisms that have been established for
– 40 –
reactions i.e. the outer-sphere and the inner-sphere mechanisms These two
mechanisms were first proposed by Taube and his co-workers (Taube et al, 1953,
Taube 1959).
1.4.1 Inner-Sphere Mechanism:
An inner-sphere mechanism is one in which the reductant and oxidant share a
ligand in their inner or primary coordination spheres and the electron is transferred
across a bridging group. That is, the two reactants are linked together by at least
one bridging ligand common to the coordination shell. The essential feature of this
mechanism is that substitution takes place at one of the metal centres to give
binuclear ligand bridged specie prior to the transfer of an electron. The first redox
reaction that demonstrated this type of mechanism was the reaction between
[Co(NH3)5Cl]2+ and [Cr(H2O)6]2+
[CoIII(NH3)5 Cl]2+ + [CrII (H2O)6]2+ Co2+ + 5NH4
+ + CrCl2+ …(1.34)
This reaction involves a binuclear ligand -bridged intermediate of
[CoIII(NH3)5 –Cl- CrII (H2O)5]4+ + H2O
which leads to transfer of electron .
[CoIII(NH3)5 –Cl- CrII (H2O)5]4+
and forming a product of
Co2+ + 5NH4
+ + CrCl2+ (the bridging ligand has been transferred).
The reaction product will then indicate that there was transfer of electron
from cobalt to the chromium through the ligand which is usually but not always
transferred from one reactant to another. The bridging ligand X may be F- Cl-, Br-,
I- and group VI etc. The only essential feature for a bridging ligand is that it
– 41 –
should have at least one lone pair of electrons available for bonding to metal
cation that is at least one lone pair electron beyond the electron pair needed to
bond to cobalt in the first place in equation (1.34) above. For example, ligands
such as NH3 and other Group 5 bases have only one lone pair of electron available
for bonding to a cation and they use this to coordinate with the first metal. They
therefore cannot act as bridging ligand in inner-sphere redox reaction.
The prerequisite for inner sphere mechanism includes:
(i) One of the reactants (usually the oxidant) should possess at least one
ligand capable of bonding simultaneously to two reaction centres (metal ions).
Although this bridging ligand (atom) is frequently transferred from the oxidant to
the reductant in the course of electron transfer, it is not always the case.
(ii) One ligand of the reactants (usually the reductant) should be
substitutionally labile i.e. capable of being replaced by a bridging ligand in a facile
substitution process
The basic step in this type of reaction may be summarized by the sequence
below as proposed by Sutin, (1968), Bennet, (1972), and Linck, (1972).
Step 1: – Formation of an Encounter or Collision Complex.
L5MIIIX2++ NII(H2O)6
2+ [L5 MIII X // NII (H2O)6] 4+ …(1.35)
( collision complex)
Step 2: Formation of a bridged precursor complex
[L5MIII X // NII (H2O)6 ]
4+ H2O + [L5MIII X- NII (H2O)5]
4+ (1.36)
Step 3: Formation of activation of the precursor complex
[L5MIII X-NII (H2O)5 ]4+
[L5MIII X- NII (H2O)5]
4+ …(1.37)
Step 4: Formation successor complex and electron transfer.
[L5MIII X-NII (H2O)5] 4+ [ L5MII X-NIII (H2O)5] 4+ …(1.38)
– 42 –
Step 5: Deactivation of the successor complex
[L5MII – X-NIII (H2O)5 ]4+ [L5MII –X- NIII (H2O)5
] 4+ …(1.39)
Step 6: Decomposition of the successor complex
[L5MII – X-NIII (H2O)5 ]4+
H2O L5MII (H2O)2+ + NIII (H2O)5X2+…(1:40)
A typical energy profile for a redox reaction which occurs by the inner -sphere
mechanism is often illustrated as below
C
B intermediate
A
Reaction coordinate
Fig:1.2 Energy profile diagram for a redox reaction which occurs by the innersphere
mechanism
D
Energy
– 43 –
A. = Collision complex
B. = Bridged precursor complex formation
C. = Successor complex formation and act of electron transfer.
D. = Decompostion of successor complex giving the products
Any of the reaction steps could be the rate-determining step but there are
three special steps that are practicable.
i. Formation of successor complex is the rate-determining step. This is
found in substitution controlled redox reactions.
ii. Electron transfer within the bridged complex is the rate-determining
step. This is the most commonly observed result in inner- sphere
mechanism.
iii. The decomposition of the successor complex is the rate-determining
step e.g
[Co(CN)6]3- + [Fe(CN)6]3- [ CN5CoIIICNFe(CN)6 ]6- ..(1:41)
The successor complex has been identified as a solid.
However, in 1972, Linck reported that the reaction usually encountered ratedetermining
step that involves the actual transfer of electron(s) i.e the transition states are
1) systems in which the highest point on the free energy-reaction co-ordinate
diagram corresponds to the acts of electron transfer itself
2) systems in which the precursor complex formation dominates the process.
3) systems in which the successor complex destruction is of importance
He then reported that the composition of all the three transition states are the same
and can be schematically illustrated as in Fig 1:3.
– 44 –
Scheme 1
This involves the formation of the precursor complexes with their rates of
transformation to successor complexes relatively slow. For example, the reduction of
[Co(en)2H2OCl]2– by Fe2+. This reaction proceeds by substitution on the Fe2+ ion
followed by electron transfer with subsequent rupture of the Co–Cl bond to give FeCl2–
and Co2– ( Linck, 1972). The binuclear complex formed must be stable so that this
reaction will dominate.
Scheme II
In Scheme II, the rate of electron transfer occurs rapidly as soon as the precursor
complex is formed. For example, oxidation of CoIII by series of reductants as shown in
Figure 1:3.
Scheme III
Here the reactants and the successor complexes are in equilibrium with the overall
reaction rate dependent on the rate of bond rupture in the successor complex. The
reduction of Cis-[Ru(NH3)4Cl2]2+, Cis-[Ru(NH3)4H2OCl]2– and [Ru(NH3)3Cl]2– by Cr2+
proceeds by Scheme III. (Basolo and Pearson, 1967, Linck 1972).
– 45 –
Energy
Scheme I Scheme II Scheme III
Fig 1:3: – The three pathways for achieving and destroying the activated
complex in inner-sphere electron transfer process (Linck, 1981)
– 46 –
1.4.2 Outer–Sphere Mechanism:
Outer-sphere electron transfer constitutes one large, well-recognized and
extensively studied group of reactions. It is one in which the reactants do not form
an intermediate with bridging functional group to provide a pathway for electron
transfer.
In principle, outer-sphere mechanism involves electron transfer from
reductant to oxidant with the coordination sphere of each staying intact
throughout. Both reactants are inert with respect to substitution or one is relatively
inert and does not present site for the labile reactants.
This mechanism also takes place when the energy barrier involved is low or
is zero. The mode of activation of outer-sphere mechanism can be illustrated as
shown below for the reaction between
[FeII(CN)6]4- + [Ir(IV)Cl6]2- [FeIII(CN)6]3- + [Ir III Cl6]3- …(1.42)
This mechanism occurs in four steps as shown below.
Step I:- Formation of precursor complex
[FeII(CN)6]4- + [IrIVCl6]2- [(NC5) FeII(CN)(Cl- IrIVCl5)]
6- …(1.43)
Step 2:- Activation of precursor complex
[(NC)5FeII(CN)Cl IrIVCl5 ]
6- [ (NC)5FeII(CN)Cl IrCl5 ]
6- …(1.44)
Step 3:- Electron transfer and formation of the successor complex
– 47 –
[(NC)5 FeII (CN)Cl IrIV–Cl5] 6- [ (NC)5FeIII (CN)[Cl- Ir-Cl5]
6- …(1.45)
Step 4:-Decomposition of the successor complex to give final products
[ (NC)5 FeIII (CN) Cl IrIII-Cl5 ]
6- [Fe (CN)6 ]3- + [ IrCl6 ]
3- …(1.46)
In the above, any of the steps can be rate determining step but it is important to
mention the following points:-
The first step is known as the formation of precursor complex, wherein the
distance between the reactant centers is approximately that required for electron
transfer, but their relative orientations and internal structures do not yet permit
electron transfer. (Purcell and Kotz, 1977)
The second and third steps involve structural changes in the precursor to
accommodate the electron transfer. Within the precursor, there must be both a
reorganization of the oxidant and reductant complexes and within those
complexes, structural changes that define the chemical activation process for
electron transfer.
As the transition is passed, there follows the completion of the electron
transfer and final relaxation of the oxidant and reductant structure.
The fourth and final step is the separation of the products, ions or molecules
A typical energy profile for a redox reaction which occurs by the outer sphere
mechanism is illustrated in Fig 1:4
– 48 –
B
C
A
Reaction coordinate
Fig:1.4 Energy profile diagram for a redox reaction which occurs by outer sphere
mechanism
where A. Formation of precursor complex
B. Electron transfer and formation of successor complex
C. Separation of the successor complex to give final product
Theoretically any of the above steps can be the rate-determining step, but
the most usual rate-determining step is B that is the electron transfer and formation
of the successor complex increases the energy of the reaction.
However, several chemists have developed theoretical treatments of
outersphere redox reaction. Probably the best known theory is the version
proposed and developed by Marcus (1956)
Marcus theory determines the rate of outer-sphere redox reaction as the rate
of electron transfer between the oxidized and reduced forms of each of the couples
that make up the redox reaction.
Energy
– 49 –
1. 4.2.1 Proton Coupled Electron Transfer (PCET) Mechanism
This is the mechanism that involves the simultaneous transfer of proton and
electron under the outer-sphere mechanism. Here, it is also possible for threespecies
to come together in the form of an ion pair leading to the final products.
The situation is such that an open coordination will act as an electron-proton
acceptor (Meyer and Taube, 1987). This mechanism has been observed for a
number of reactions (Ghosh et al, 1994, Arabel et al, 1997, Iyun and Lohdip, 1999,
2001).
The requirement for the occurrence of the PCET pathway is that the
molecule or substrate must contain
(i). a protonable moiety,
(ii). be capable of accepting an electron or contain acidic protron(s)
or
(iii). have easily removable electron(s)
(iv). have easily oxidisable group(s) or easily removable
electron(s).(Chaudhuri et al, 1995, Arabel et al, 1997, Iyun et al,
1997)
It has been found that mixed valence manganese complexes fulfill
two of the above conditions as shown in the equations (1.47)–(1.51).
[L2MnIIIO2MnIVL2]3+ +H+ + e– [L2MnIIO(OH)MnIII L2]3+ …(1.47)
L = bipyrindine (bpy)
[L2MnIIIO2MnIVL2 +H+]3++ e– [L2MnIIO(OH)MnIII L2]3+ …(1.48)
[ L2MnIIO(OH)MnIIIL2]3+ +H+ + e– [ L2MnIII(OH)MnIII L2]
3+ (1.49)
L=Phenanthroline (phen)
– 50 –
[L2MnIIIO2MnIVL2]3++ e– [ L2MnIIIO2MnIII L2] 2+ …(1.50)
[L2MnIIIO2MnIIIL2 ]2+ + H+ [ L2MnIIO(OH) MnIII L2] 2+ …(1.51)
L = N, N-bis (2- methyl pyridyl) ethane-1,2-diamine.
The redox reactions of these systems have been shown to occur by PCET
pathway (Chaudhuri et al 1995, Arabel et al, 1997 Iyun et al, 1997). Other
substrates which have been shown to promote the PCET mechanism include thiols
and hydroxy-acids because both groups of compounds contain acidic protons and
easily oxidisable groups.
1.4.3 Solvated Electron Theory
In addition to the two mechanisms for electron transfer reactions,
another possibility is the solvated electron concept (Latimer 1952). In this
theory, the reducing agent is assumed to eject an electron into the solvents
which solvate it and holds it until the oxidizing agent picks it up. The redox
process involved in this model is thought to occur by the following steps
equations (1.52)–(1.54).
A + S A+ + S- …(1.52)
B + S- [B + S-] …(1.53)
[B + S-] B + S …(1.54)
– 51 –
Where A = reductant, B= oxidant and S = Solvent.
This process has been found to occur readily in non-aqueous solvents.
(e.g. liquid ammonia and liquid halides), but it is unlikely to occur in
aqueous medium. There is no available evidence to show that in aqueous
solutions, the electron from a reducing agent is released and becomes
solvated before reacting with an oxidizing agent (Cooke, 1979). This is
because the aquated electron (Eo = – 2.7V), will rapidly reduce water to
hydrogen (Latimer, 1952) as in equation (1.55).
e–
(aq)+H2O ½H2(g) + OH- …(1.55)
1.5 FACTORS AFFECTING OUTER-SPHERE REACTIVITY
In reactions occurring by the outersphere mechanism, it is possible to assess
the factors (at the molecular level) that determine rate of electron transfer. The
outer-sphere reaction is far easier to treat than reactions occurring by inner-sphere
mechanism because of its relative simplicity and the absence of bond-making or
bond-breaking steps. (Wilkinson, 1987)
There are three main factors that play important roles in determining the
rate of electron transfer.
1.5.1 Intramolecular Vibrational Trapping
A commonly used example for illustrating intra molecular vibrational
trapping induced by structural changes is [Fe(H2O)6]2+ and [Fe(H2O)6]3+ selfexchange
reactions. X-ray crystallographic studies in solution show that the most
significant structural difference between the aquated FeII and FeIII ions is a
– 52 –
symmetrical decrease in the Fe-OH2 bond distance upon oxidation. Electron
transfer is accompanied by net structural changes at each redox site since both
change their electron content during the reaction e.g the site labelled B in equation
(1:56) will have the equilibrium structure of FeII before electron transfer and the
equilibrium structure of FeIII after electron transfer.
[Fe(H2O)6]3+,[Fe(H2O)6]2+ [Fe(H2O)6]2+.[Fe(H2O)6]3+…(1.56)
A B A B
The changes in structure that must occur create barriers to electron transfer.
In order to understand the origin of the barrier and to treat it quantitatively, it is
necessary to recall that the structural changes at each reactant can be resolved into
a linear combination of its normal vibrational modes.
The normal modes constitute a complete, orthonormal set of molecular
motions into which any changes in intramolecular structure can be resolved. From
the essentially symmetrical decrease in all six Fe-O bonds between FeII and FeIII, it
can be concluded that the only significant normal mode contributing to the
structural change between FeII and FeIII is the totally symmetric (Fe-O) iron –
oxygen breathing mode.
1.5.2 Solvent Trapping
The solvent plays a role in trapping the exchanged electron on one site
similar to that of intramolecular structural changes. An ion can be thought of as
creating a polarization field in the surrounding solvent. At the molecular level,
there are three ways that the solvent molecules respond to the electrostastic field
of the ion.
a. orientation of permanent solvent dipole
b. electrostatically induced structural distortions within the molecules.
c. induced electrostatic polarization of electronic clouds including
bounding and non-bounding electron density.
– 53 –
The strength of the ion-solvent interaction depends on the charge on the
ion, when electron transfer occurs, changes must also occur in the orientation of
the surrounding solvent dipoles since in the electron transfer act, the polarization
field at the reactants are interchanged. The necessary re-orientation of solvents are
closely related to rotations of molecules in the gas phase but are necessarily
collective in nature because of molecule-molecule interaction in the condensed
phase of the solution.
The orientation component arising from solvent dipoles must adopt a
non-equilibrium distribution before electron transfer can occur. The
orientation of solvent dipoles contributes to the energy of activation through
the time scale for dipole re-orientation. It can also contribute to the preexponential
or frequency factor for electron transfer.
1.5.3 Electron coupling:-
For electron transfer to occur between reactants, an electronic
Interaction must exist which tends to delocalise the exchanging electron
between sites. Neglecting the overlap, the magnitude of interaction is given
by equation (1.57)
v = (QD/v/QA) …(1.57)
where
= electrostastic operator that describes the electronic pertubation
between the electron donor and acceptor and causes electron transfer
to occur.
QA = electronic wave function for the acceptor
QD = electronic wave function for the donor.
The electronic interaction leads to two new electronic states, an upper
state and a lower state that combine to give bonding and anti bonding
– 54 –
molecular orbital. If the overlap between QD and QA is negligible, the
energies of the two new states including the electronic interaction but
neglecting the vibrational energies are given by equation (1.58)
Ue + V and Ue – V …(1.58)
where Ue = Ue
p =Ue
R for a self-exchange reaction.
The existence of vibrational trapping creates an unusual situation
compared with other molecular orbital-type problems. If electronic coupling
between the two sites is weak, the effect of vibrational trapping will be to
localize the exchanging electron either on the electron donor site or on the
acceptor site. The delocalization of the exchanging electron is in fact the
electron transfer process itself.
If the electronic coupling is weak, electron transfer becomes an
occasional event that can only occur at an appreciable rate from electron
donors and acceptors having the non-equilibrium vibrational distributions
and solvent dipole orientations appropriate for maximizing the electron
transfer rate.
1.6 CRITERIA FOR ASSIGNING THE MECHANISM OF
ELECTRON TRANSFER REACTIONS
Kinetic studies of electron transfer reactions try to assign the system
under study to one of the two established reaction mechanisms. i.e. the
outer-sphere and inner-sphere mechanism. The following criteria discussed
below have been used for assessing and distinguishing the types of
mechanisms for electron transfer reactions.
– 55 –
1.6.1. Rate of redox process versus rate of substitution (kred vs ksub)
For the inner-sphere mechanism to take place, there must be
substitution of a ligand into the co-ordination shell of one of the metal ions
before the electron transfer process. This is to form ligand bridged precursor
complex.
If the rate of redox process is greater than rate of substitution, then the
reaction cannot proceed through the inner-sphere mechanism. The outersphere
mechanism must be presumed to be operating.
An example of a reaction with outer sphere mechanism is as given in
equation (1.59).
[*Fe(CN)6]4– + [Fe(CN)6]3- [*Fe(CN)6]3- + [Fe(CN)6]4– …(1.59)
k red = 103 dm 3 mol-1 s-1.
[Fe(CN)6]3– +*CN– [Fe*(CN)(CN)5]4–+ CN– …(1.60)
k sub = 7.8 x 10 –6 dm 3 mol-1 s-1.
Because kred ksub, the electron transfer in the above reactions
cannot occur via formation of a binuclear intermediate hence the reaction
has been classified as occurring by the outer-sphere mechanism. (Rosenhein
et al, 1974)
According to Larsen and Wahl in 1965, electron exchange reaction between
[Fe(phen)3 ]2– and [*Fe(phen)3 ]3– can also serve as another example.
[Fe(phen)3]2+ + [*Fe(phen)3]3+ [Fe(phen)3]3+ + [*Fe(phen)3]2+…(1.61)
– 56 –
The rate constant for the substitution of phen are 7.5×10–5s–1 for [*Fe(phen)3]2– and 5.01
x 10–5s–1 for [Fe(phen)3]3+ while k for the exchange is 105 dm3 mol–1s–1. This is another
example where the outersphere mechanism is operative because kred>>ksub.
When ksub>> kred, the reaction will follow inner-sphere mechanism for example,
the reduction of [Co(NH3)5Cl]2+ by Cr2+ as reported by Taube et al (1954). When kred =
ksub, such reaction is said to be substitution controlled. An example of such reaction is the
oxidation of Ru2+ by ClO4
–.
2Ru2+ + ClO4
–.+ 2H+ 2Ru3+ + ClO3
–.+ H2O …(1.62)
Rate = k[Ru2+ ][ ClO4
–.] …(1.63)
This reaction is believed to be substitution controlled, the rate constant k and activation
parameters being similar to those of halide / anion reactions.
1.6.2 Reactivity Pattern.
i) Reactivity pattern with a wide range of reactants .
If a wide variety of complexes such as [Co(NH3)5X]2+ (X = Cl–, F–,
Br–, NO3
–) is reacted with another metal ion, the rate of the reactions for
inner-sphere mechanism would be expected to be sensitive to the nature
of X (the bridging ligand), while in the case of outer-sphere mechanism,
the rate will be similar irrespective of the identity of X (Sykes, 1967, Shea
and Haim, 1973, Adegite et al, 1977).
ii) Relative rates of reaction of hydroxyl and aquo complexes.
The hydroxyl group (OH–) is a better bridging ligand than H2O so it
is expected that the hydroxo complexes react faster via the inner-sphere
mechanism. Thus where kOH kH O 2 the outer-sphere mechanism is said
to be operating while the converse is true when kOH kH O 2 .
– 57 –
iii) The effect of added anions and cations gives clue to the type of
mechanism in a given redox reaction (Pennington and Haim, 1967;
Prizstas and Sutin, 1973 and Adegite et al, 1977). Effect of added anions
on inner-sphere redox reactions are much smaller than that of outer-sphere
reactions. Anion catalyzed reactions are presumed to be operating via
outer-sphere mechanism. Thus an examination of the dependence of redox
rate on added anion such as chloride ion has been used for the
identification of the mechanism.
iv) The use of ambidentate bridging ligand.
The distinction between the inner sphere and outer sphere mechanism can be
made on the basis of the difference in the rate of electron transfer observed with
symmetrical and unsymmetrical bridging ligands e.g. azide, thiocyanate and
isothiocyanate complex. If the ratio of the rate constant for the reduction of azide
complex to that for reduction of the thiocyanate complex
NCS
N
k
k 3 or for
isothiocyanate and thiocyanate complex
NCS
SCN
k
k shows no difference or are
similar, then the outer sphere mechanism should be operating while the converse is
true for reactions of the inner sphere mechanism (Sutin, 1968, Wang and Esperson;
1965, Candlin et al; 1964, Fay, 1970).
1.6.3 Identification of binuclear intermediate.
The presence of binuclear intermediate confirms that the reaction proceeds via an
inner sphere mechanism. The identification of the binuclear intermediate in the inner
sphere mechanism is however possible if the reduced oxidant and the oxidized reductant
are relatively inert and the rate of decomposition of the intermediate is slower than the
rate of electron transfer. In some cases, where the binuclear intermediate has transient
– 58 –
nature, their presence can be inferred indirectly from kinetic data or from empirical rate
law. An example of such indirect identified binuclear intermediate includes CrOHV5+
(Espenson 1965) and OHTi and [Co(NH3)6]3+ (Orhanovic and Earley 1975).
At times, the intermediates are kinetically stable enough to be detected during the
reaction studies which provides a convincing evidence for the interpretation of the innersphere
mechanism
IF YOU CAN'T FIND YOUR TOPIC, CLICK HERE TO HIRE A WRITER»
