ABSTRACT
The kinetics and mechanisms of the redox reactions of 3,7-bis
(dimethylamine)phenothionium chloride (MB+) with the oxyanions, BrO3
-, S2O8
2- and
SO3
2-, have been studied. The rate of the reaction is first order in each oxidant and
reductant concentrations for the three reactions. MB+ was oxidized by BrO3
– and S2O8
2-
and the order of reactivity is k2 (BrO3
-) k2 (S2O8
2-). The MB+- S2O8
2- reaction was Ag+
catalysed. SO3
2- reduced MB+. The rates of the redox reactions of SO3
2- an S2O8
2- with
MB+ show no acid dependence, while that with BrO3
– ion, showed acid independent and
a second order acid dependence. All the reactions are affected by change in ionic
strength. The stoichiometry is 2:3 for MB+: BrO3
– and 1:6 for MB+: S2O8
2-. The
stoichiometry for the MB+- SO3
2- could not be determined. The overall rate equation for
the three reactions can be represented as shown below:
For the MB+: BrO3
– reaction:
-1d[MB+] = k2[MB+][BrO3
-]
2 dt
where k2 = 10.16 0.29 dm3 mol-1 s-1, at [H+] = 0.20 mol dm-3 μ = 0.50 mol dm-3
(NaCI) and T = 29.0 1.00C.
For the Ag+ catalysed MB+: S2O8
2- reaction
-d[MB+] = k2[MB+][S2O8
2-][Ag+]
dt
where k2 = 0.37 0.03 dm3 mol s-1 at [H+] = 1.0 x 10-4 mol dm-3 μ = 0.10 mol dm-3
(NaNO3), [Ag+] = 5.0 x 10-3 mol dm-3 and T= 26.0 1.00C
7
For the MB+: SO3
2- reaction
-d[MB+] = k2[MB+][SO3
2-]
dt
where k2 = (3.02 0.22) x 10-3 dm3 mol-1 s-1 at [H+] = 1.0 x 10-3 mol dm3 μ = 0.10 mol
dm-3 (NaCI) and T = 25.0 1.00C
Free radicals were detected in the MB+ reactions with S2O8
2- and SO3
2-.
Spectroscopic investigation and kinetic evidence suggested the absence of short-lived
intermediate complex in all the systems studied.
The experimental results obtained for all the systems are in favour of the outersphere
mechanism.
TABLE OF CONTENTS
Content
Page
Title page – – – – – i
Declaration – – – – – – ii
Certification – – – – – – iii
Dedication- – – – – iv
Acknowledgment – – – – – – v
Abstract- – – – – – – vi
Tale of contents – – – – – viii
List of Tables- – – – – – xiii
List of Figures – – – – – xiv
Abbreviations – – – – – xvi
CHAPTER ONE
1.0 Introduction- – – – 1
1.1 Rates monitoring techniques – – 1
1.1.1 Conventional methods- – – – 2
1.1.2 Monitoring rates of fast reactions – – 2
1.2 Theoretical consideration in electron transfer process- – 3
1.2.1 Frank-Condon principle – – – 3
1.2.2 Electron tunneling theory – – – 4
1.2.3 Solvated electron theory – – – — 5
1.3 Oxidation-reduction reaction- – — 5
1.3.1 Homonuclear (Isotopic) exchange reaction – – 5
1.3.2 Heteronuclear (cross) reactions – — 6
9
1.4 Mechanisms of redox reactions – – – 7
1.4.1 The outer-sphere mechanism – – – 7
1.4.2 The inner-sphere mechanism – – – 8
1.5 Diagnosis of redox reaction mechanism – – 10
1.5.1 kredox versus ksubstitution – – – – – 11
1.5.2 Identification of binuclear intermediate – 11
1.5.3 Product analysis – – – – 12
1.5.4 Activation parameters – – – – 12
1.5.5 Reactivity pattern – – – – – 13
1.5.6 Marcus theory of electron transfer rates – – 14
1.6 Research objective – – – – 15
CHAPTER TWO
2.0 Literature Review- – – – 16
2.1 Biological stains – – – 16
2.2 Redox reactions of methlyene blue – – 18
2.3 Electron transfer reactions of oxyanions – – – 19
2.3.1 Electron transfer reactions of bromate ion – – 20
2.3.2 Electron transfer reactions of peroxydisulphate ion – 21
2.2.3 Electron transfer reactions of sulphite ion – – 22
CHAPTER THREE
3.0 Experimental – – – – 23
3.1 Materials and reagents – – – – 23
3.1.1 Preparation of 9.36x 10-6 mol dm-3 methylene blue solution 23
3.1.2 Preparation and standardization of 0.01 mol dm-3 sodium
thiosulphate solution – – – 23
10
3.1.3 Preparation of 0.1 mol dm-3 potassium bromate solution – 24
3.1.4 Preparation of 0.1 mol dm-3 sodium peroxydisulphate solution 24
3.1.5 Preparation of 0.1 mol dm-3 sodium sulphite solution – 24
3.1.6 Preparation 0.1 mol dm-3 iodine solution – – 25
3.1.7 Preparation of 1.0 mol dm-3 sodium chloride solution- – 25
3.1.8 Preparation of 1.0 mol dm-3 sodium nitrate solution- – 25
3.1.9 Preparation of 0.5 mol dm-3 sodium carbonate solution- – 25
3.1.10 Preparation of 1.0 mol dm-3 hydrochloric acid solution – 25
3.1.11 Preparation of 1.0 mol dm-3 nitric acid solution – – 26
3.1.12 Preparation of 0.1 mol dm-3 sodium formate solution- – 26
3.1.13 Preparation of 0.1 mol dm-3 sodium sulphate solution- – 26
3.1.14 Preparation of 0.50 mol dm-3 silver nitrate solution – – 26
3.2 Stoichiometric studies- – – – – 26
3.3 Kinetic measurements – – – – 27
3.4 Test for the presence of intermediate complex – – 30
3.4.1 Spectrophotometic investigation – – – 30
3.4.2 Free radial test – – – – 30
CHAPTER FOUR
4.0 Results – – – – – 32
4.1 Stoichiometry- – – – – 32
4.1.1 Methylene blue-bromate system- – – – 32
4.1.2 Silver(I) catalysed methylene blue-proxydisulphate system – 32
4.1.3 Methylene blue-sulphite system – – – 32
4.2 Determination of order with respect to reactants – – 35
4.2.1 Methylene blue-bromate system – – 35
11
4.2.2 Silver(1) catalysed methylene blue-peroxydisulphate system– 35
4.2.3 Methlyene blue-sulphite system – – — 40
4.3 Effect of hydrogen ion concentration on the rate of the reactions 47
4.3.1 Methylene blue-bromate system – – – 47
4.3.2 Silver(I) catalysed methylene blue- peroxydisulphate system 53
4.3.3 Methylene blue-sulphite system – – – 53
4.4 Effect of added anions – – – – 53
4.4.1 Methylene blue-bromate system – – – 53
4.4.2 Silver(I) catalysed methylene blue- peroxydisulphate system- 53
4.4.3 Methylene blue-sulphite system – – – 55
4.5 Test for intermediate complex – – – 55
4.5.1 Spectrophotometric investigation – – – 55
4.5.2 The Michaelis-Menten plot — – – 55
4.6 Free radical test – — – 56
4.7 The effect of ionic strength — – – 56
4.7.1 Methylene blue-bromate system – – – 56
4.7.2 Silver(I) catalysed methylene blue- peroxydisulphate system- 61
4.7.3 Methylene blue-sulphite system – – — 61
4.8 Effect of change in dielectric constant of the reaction medium– 61
4.9 Product analysis- – – – – 61
CHAPTER FIVE
5.0 Discussion – – – – — 66
5.1 Methylene blue-bromate ion reaction – — 66
5.2 Silver(I) catalysed methylene blue-peroxydisulphate ion reaction 68
5.3 Methylene blue-sulphite ion reaction – – 70
12
5.4 Comparison of the MB+- BrO3
– and the MB+ – S2O8
2- systems 72
5.5 Summary and conclusion – – – 73
CHAPTER ONE
1.0 Introduction
A reaction mechanism is the detailed stepwise process involving molecules,
atoms, radicals or ions that occurs simultaneously or consecutively and culminates in
the observed overall reaction (Cooke, 1979). The experimentally determined rate
equation, the exact nature of both reactants and products, the presence of any
equilibrium, and the stoichiometry of the reaction are all important information required
for proper elucidation of mechanisms (Basolo and Johnson, 1964). Electron transfer is
only one of many pathways by which redox reactions can occur (Basolo and Pearson,
1967; Burgess, 1978). Various reactions in inorganic and biological systems involve the
transfer of electron at one stage or the other and proper understanding of these electron
transfer processes helps in the understanding, development, and eventual effective
control of a wide area of science and technology (Iyun, 1982). Also development in
redox chemistry has shown that the basic principles which are operative for simpler
systems involving metal ions in solution extends with some modifications to
biologically important processes like the reactions of metalloproteins, photo-induced
electron transfer processes, intervalence transfer and the reactions of the hydrated
electron (Lippards, 1973).
1.1 Rate monitoring techniques
The method used to monitor the rate of a reaction depends on the species
involved and the rapidity with which their concentration change with time (Atkins,
1997).
18
1.1.1 Conventional methods
These methods involve the measurement of the concentration or any physical
property of one or more of the reactants or products that depends on their concentration
as a function of time. The most important property of a complex ion which has been
used for kinetic measurement is its absorption characteristics (Wilkins, 1962). The
absorbance of any of the reactants or products is measured and related directly to its
concentration (Wilkins, 1974).
1.1.2 Monitoring rates of fast reactions
Some reactions are so fast that special techniques have to be employed. Such
techniques are of two main types. The first type employs the same principles as are used
for slow reactions, the methods being modified to make them suitable for more rapid
reactions. The second type is of a different character and involves special principles like
temperature jump (Laidler and Meiser, 1982).
Flow technique: In the flow method the reactants are mixed as they flow together in a
chamber. Different flow techniques can be employed, depending on the treatment given
to the reaction after mixing (Caldin, 1964). In the quenching method, the reaction is
stopped after it has been allowed to proceed for a certain time, and the composition is
analyzed, by any conventional method.
In the continuous flow method the reaction is allowed to continue as the
thoroughly mixed solution flow through the outlet tube, and the composition at different
positions along the tube is observed. Spectrophotometric methods are commonly
employed for the observation (Candlin and Halpern, 1965; Campion et. al., 1964;
McAuley and Hill, 1969).
19
Relaxation methods: These methods are used for systems at equilibrium (Eigen and
Maeyer, 1963). In these cases, the equilibrium is displaced by some external parameters
and the readjustment to a new equilibrium position is directly observed. Two main types
of physical perturbation have been used: The stepwise or transient methods and the
stationary methods.
Resonance technique: Both nuclear magnetic resonance and electron spin resonance
techniques have been used to study complex formation (Pearson and Buch, 1962) and
the aquation of metal ions (Connick and Poulson, 1959). The width of the resonance
absorption line is related in n.m.r to the lifetime of the nucleus in a given spin state and
in e.s.r to the lifetime of the paramagnetic species in a given energy state, and any
reduction of the lifetime of these states by chemical interaction results in line
broadening.
1.2 Theoretical consideration in election transfer processes
1.2.1 Franck-Condon principle
The principle states that nuclear motion is slow as compared to electronic
motion and that election transfer occurs without an appreciable movement of the nuclei
(Sutin, 1966). This implies that electron transitions are rapid compared with nuclear
motions and electron transfer occurs without significant movement of the atoms.
Exchange of electron can occur when the two particles have nearly identical structural
and electronic configurations and the total change in energy involved in the process can
be represented by equation (1.2) (Marcus, 1956)
20
G = Gt
+ Gi
+ Go
————————————————- (1.2)
Gt
= Association free energy
Gi
= lnner–sphere reorganisation energy
Go
= Outer–sphere reorganisation energy
The principle also assumes that no angular momentum can be transferred to or
from the transition state during electron transfer. It is also important that the two ions
have identical spin states.
1.2.2 Electron tunneling theory
This theory states that solvent and ligands produce an electronic
energy barrier that the transferring electron must penetrate (Weiss, 1954; Marcus,
1956). The implication is that the electron will be able to travel distances much greater
than would correspond to the actual collision of reactants (Basolo and Pearson, 1967).
Reorganization of the solvent molecules had to occur prior to electron transfer (Marcus,
1957). In terms of the transition state theory, their results may be written in the form
k = KTk1exp (- Gr
≠ – Ge
≠) —————————————- (1.3)
h RT RT
k1 = electron transmission coefficient
k = rate constant
K = Boltzman constant
T = absolute temperature
Gr
= activation free energy for rearrangement of the hydration and
coordination shells
Ge
= activation free energy for overcoming electric repulsion between the
ions.
21
The transmission coefficient, which always remains less then one, increases, as
the exchanging partners come closer together.
1.2.3 Solvated electron theory
This theory attributes the mechanism of electron transfer in solution to the
ejection of an electron by the reductant into the solvent (equation 1.4). The solvent then
solvates the electron and holds it until the oxidant picks it up (equation 1.5). The
possibility that this mechanism operates in aqueous solution is remote, due to the fact
that solvated electron with estimated potential of –2.70V (Latimer, 1952; Burgess,
1978) cannot exist in water without it reducing water (equation 1.6)
A + S = A+ + S- —————————————————— (1.4)
S – + B = [S- + B]≠ S + B ——————————– (1.5)
where A = Reductant
B = Oxidant
S = Solvent
e-
(aq) + H2O ½ H2(g)
+ OH- ———————————– ( 1.6)
1.3 Oxidation –reduction reactions
Redox reactions are those in which the oxidation states of some atoms change.
Generally, they can be classified into two broad groups: homonuclear (isotopic)
exchange reacion and chemical or cross reactions (Edwards, 1964).
1.3.1 Homonuclear (Isotopic) exchange reactions.
This class of reactions involves transfer of electrons between two identical metal
ion centers existing in different oxidation states. The reactants and products are the
same and identical hence have the same concentrations. The free energy change for such
22
a reaction is mainly due to mixing and therefore is approximately zero (Sharpe, 1982).
Examples of these reactions are given in equations (1.7) and (1.8).
[FeEDTA]2-+ *[FeEDTA]- = [FeEDTA]- + *[FeEDTA]2- ——————-(1.7)
TlI + *TIIII = T1111 +*T11 ——————————————————(1.8)
(where * is an isotopically labeled specie)
In these types of electron transfers, there is no net chemical change and the
equilibrium constant is one, since the rate constants for the forward and backward
reactions are equal (Edwards, 1964). This class of reactions is studied by isotopic
labeling (Burgess, 1978).
1.3.2 Heteronuclear (Cross) reactions
This class of reactions involves transfer of electrons between different metal ion
centers and the products are chemically distinct from the reactants (Cooke, 1979). The
net change in free energy in most cases is less than zero (G < 0). Examples of these
reactions are given in equations (1.9) and (1.10)
2Fe2+ + Tl3+ =2Fe3+ + Tl+ ———————————————— (1.9)
2[Fe(phen)3]3+ + 2l- = 2[Fe(phen)3]2+ + I2 ———————————- (1.10)
Reactions of the type in which the oxidant and reductant change oxidation state by the
same number of electrons are termed complimentary, as represented in equations (1.11)
and (l.12) (Halpern, 1961).
[Co(en)3]3+ + [Ru(NH3)6]2+ = [Co(en)3]2+ + [Ru(NH3)6]3+ —————(1.11)
FeIIcyt–c + [Ru(NH3)6]3+=FeIIIcyt-c + [Ru(NH3)6]2+ ——————— (1.12)
When the oxidant and reductant differ in their change in oxidation state (equations 1.13
and 1.14) the reactions are termed non-complementary reactions. These generally
23
proceed by complicated mechanisms (Burgess, 1978).
2V2+ + Br2 = 2V3+ + 2Br – ————————————————–(1.13)
2Eu2+ + C12 = 2Eu3+ + 2Cl- —————————————–(1.14)
1.4 Mechanism of redox reactions
The outer-sphere and the inner–sphere mechanism have been identified as the
basic mechanisms that are operative when electrons are transferred between two metal
ions in solution (Taube et.al., 1953; Taube and Meyers, 1954; Taube, 1959).
1.4.1 The outer–sphere mechanism
This type of mechanism occurs when both reactants are inert with respect to
substitution or when one of the reactants is relatively inert and does not present site for
the labile reactant (Larsen and Wahl, 1965). The outer-sphere mechanism is
characterized by low or zero G or E (Sharpe, 1982). The mode of activation of the
outer-sphere mechanism can be illustrated as shown below for the reaction between
Fe(CN)6
4- and IrCl6
2-:
[FeII(CN)6]4- + [IrIVCI6]2- [FeIII(CN)6]3- + [IrIIICI6]3- ———-(1.15)
The above reaction takes place via the outer-sphere mechanism, and just like every
other outer-sphere reaction, is thought to occur in four steps as shown below;
Step 1: Formation of precursor complex
[FeII(CN)6]4- + [IrIVCl6]2- [(NC)5FeII(CN)(ClIrIVCI5)]6- ———–(1.16)
Step 2: Activation of precursor complex
[(NC)5FeII(CN)(ClIrIVCI5)]6- [(NC)5FeII(CN)(ClIrIV-Cl5)]6-≠ —(1.17)
Step 3: Electron transfer and formation of successor complex
[(NC)5FeII(CN)(ClIrIV-Cl5)]6- [(NC)5FeIII(CN)(ClIrIII-Cl5)]6- —–(1.18)
24
Step 4: Decomposition of successor complex to give final products
[(NC)5FeIII(CN)(ClIrIII-Cl5)]6- [FeIII(CN)6]3- + [IrIIICI6]3- ——-(1.19)
Reaction Cordinates
Energy
Fig. 1.1:- Energy profile diagram for a redox reaction which occur by the outer–
sphere mechanism.
Although any of the steps can be the rate determining step, step 3- the electron transfer,
is usually the rate determining step.
1.4.2. Inner-sphere mechanism
In this mechanism, electron transfer is preceeded by substitution into the
coordination shell of one of the reactants, with the formation of a bridged intermediate
in which two metal ions are linked by a common ligand (Halpern, 1961). Typical
reactions that occur via this mechanism include reactions 1.20 and 1.21 below
≠1
≠2
≠3
≠4
25
[CoIII(NH3)5Cl]2+ +[CrII(H2O)6]2+ [CoII(H2O)6]2+ +5NH4
+ +[CrIII(H2O)5Cl]2+ —- (1.20)
[(Cr(OH2)5Cl]2+ +[Cr*(OH2)6]2+ [(Cr(OH2)6]2+ + [Cr*(OH2)5Cl]2+——- (1.21)
In general, the inner–sphere mechanism can be represented by the following reaction
scheme.
Step 1: Formation of encounter or collision complex
[L5MIIIX]m+ + [NII(H2O)6]n+ [L5MIIIX//NII (H2O)6](m + n) + ————–(1.22)
Step 2: Formation of a bridged precursor complex
[L5MIIIX//NII(H2O)6](m+n)+ H2O+[L5MIII-X-NII(H2O)5](m+n)+———(1.23)
Step 3: Activation of the precursor complex.
[L5MIII-X-NII(H2O)5](m+n)+ [L5MIII-X-NII(H2O)5](m+n)+ —————- (1.24)
Step 4: Electron transfer and formation of successor complex
[L5MIII-X-NII(H2O)5](m+n)+ [L5MII-X-NIII(H2O)5](m+n)+ ————– (1.25)
Step 5: Deactivation of successor complex
[L5MII-X-NIII(H2O)5](m+n)+ [L5MII-X-NIII(H2O)5](m+n)+ ————(1.26)
Step 6: Decomposition of activated complex
[L5MII-X-NIII(H2O)5](m+n)+ [L5MII(H2O)]m++ [XNIII(H2O)5]n+ —(1.27)
Investigations have revealed that the electron transfer step is usually the rate
determining step (Iyun, 1982). Typical energy profile for a redox reaction which occurs
by the inner–sphere mechanism can be illustrated by Figure (1.2)
26
Reaction Cordinates
Energy
Fig. 1.2:-Typical energy profile for a redox reaction which occurs by the
inner–sphere can be illustrated by figure
≠1 = Collision complex
≠2 = Bridged precursor complex formation
≠3 = Successor complex formation and the act of electron-transfer
≠4 = Decomposition of successor complex giving the products.
1.5 Diagnosis of redox reaction mechanism
Electron-transfer reactions may be occurring by either the outer- or the innersphere
mechanisms. Some of the criteria that are used in assigning a mechanism for
reactions are discussed below:
≠1
≠2
≠3
≠4
27
1.5.1 kredox versus ksubstitution
If the rate of electron-transfer is greater than the rate of substitution into the
inner-coordination sphere of either of the complexes, the outer-sphere mechanism is
indicated (equation 1.28) (Rosenheim et.al., 1974).
V2+ + Fe3+ V3+ + Fe2+ ————-(1.28)
kredox= 1.8 x 104 dm3 mol-1 s-1
ksubstitution= 100 s-1
The outer-sphere mechanism is established when rapid electron transfer occurs between
two substantially inert complexes (Burgess, 1978).
However for reactions in which the rate of substitution into the innercoordination
sphere of either of the complexes is faster than the rate of electron transfer
and where a bridging ligand is present, the inner–sphere mechanism is a possibility. For
the reaction (equation 1.20), the reductant CrII is labile (ksubstitution= 1.0 x 108 s-1) and the
reaction was assumed to involve rapid and reversible formation of the bridged
intermediate followed by a slow electron transfer (kredox= 6.0 x 105 dm3 mol-1 s-1)
(Cooke, 1979).
1.5.1 Identification of binuclear intermediate
The detection of an intermediate with a suitable bridging ligand is an indication
that the redox reaction has occurred by means of a bridged activated complex and this
provides a convincing evidence for the operation of the inner-sphere mechanism
(Burgess, 1980).
The presence of an intermediate is also indicated if the rate law contains inverse
orders, non-integral orders or orders larger than three (Edwards, 1964).
28
1.5.3 Product analysis.
Hints about the possible pathway for the redox reactions could be obtained from
the isolation and characterization of the products of the reaction. It is also possible to
establish the details of how electron transfer reactions occur based on isotopic labelling.
The use of a particular isotope enables the observation of bond-breaking processes and
the identification of atom or group of atoms transferred from one metal center to
another. An example is equation (1.20) based on the reduction of substitutionally inert
amine complexes of CoIII by substitutionally labile Cr(H2O)6
2+ and the introduction of
labelled Cl- into the reaction mixture. The appearance of unlabelled Cl- in the inert CrIII
product indicated that both metal centers must have been bonded simultaneously to Clas
a bridging ligand when electron transfer occurred. (Taube, et. al., 1953).
However, there are inner–sphere reactions which are not accompanied by atom
transfer, equation (1.29) (Cooke, 1979). In this reaction it was demonstrated that bridge
formation occurred but the bridging ligand was not transferred.
[Co(EDTA)]2- + [Fe(CN)6]3- [(EDTA)CoNCFe(CN)5]5-
[Co(EDTA)]- + [Fe(CN)6]4- ————————-(1.29)
1.5.4 Activation parameters.
The use of activation parameters, ΔH≠, ΔG≠ and ΔS≠ to determine the type of
mechanism operating in a particular redox process is of little use, because there seems
to be no direct correlation between activation parameters and the type of mechanism
(Cotton and Wilkinson, 1980). However change in molar volume of activation (ΔV≠)
has been used in mechanistic diagnosis (Hubbard et. al., 1991).
29
1.5.5 Reactivity pattern
(i) Relative rates of reaction of azido and thiocyanato complexes
Comparison of the effect of azido and thiocyanato ligands on the rates of
electron transfer reactions involving metal ions and their complexes have been widely
used as a basis for determining whether a particular reaction proceeds by the inner– or
the outer–sphere mechanism (Ball and King, 1958; Espenson, 1965; Douglas and Sutin,
1970). If the metal centers of the oxidizing and reducing agents are hard and if transfer
of the bridging group occurs during the reaction, then an inner-sphere mechanism
should proceed faster when the bridging group is azido than when it is isothiocyanato.
(ii) The effect of added anion
The effect of added anions and cations give an indication of the type of
mechanism operating in a redox process (Pennington and Haim, 1967; Przystas and
Sutin, 1973; Adegite et. al., 1977). The rate of oxidation of acetaldehyde by BrO3
– was
found to be inhibited by the added formate and nitrate ions. The reaction was therefore
suggested to have occurred via the outer-sphere mechanism (Lohdip and Iyun, 1993).
Although there are exceptions (Halpern, 1961) generally anion catalyzed reactions are
presumed to be operating via the outer-sphere mechanism (Dulz et. al., 1964; Przystas
and Sutin, 1973).
(iii) Relative rates of reaction of hydroxo and aquo complexes
The hydroxyl group (OH-) is a better bridging ligand than H2O, so the hydroxo
complexes are expected to react faster via the inner–sphere mechanism, which involves
the formation of bridged intermediate, thus where kOH >> kH O
the inner sphere
mechanism is said to be operating (Cotton and Wilkinson, 1980).
30
(iv) Reactivity pattern with a wide range of reactants
The relative rates of oxidation of a series of reducing agents by two different
oxidants should be independent of the identity of the reducing agent if both set of
reactions are by the outer–sphere mechanism (Marcus, 1963). For the inner-sphere
mechanism the rate is very sensitive to the nature of the reductant, the oxidant and the
bridged atom (Edwards, 1964).
1.5.6 Marcus theory of electron transfer rates
This theory predicts a simple relationship (equation 1.30) between the rate
constant k12 for an electron–transfer reaction, K12, the equilibrium constant for the
reaction, and k11 and k22, the self exchange rate constant for the reductant and oxidant
complexes (Marcus, 1963)
*Ox1 + Red2 *Red1 + Ox2——————– (1.32)
*Ox1 + Red1 *Red1 + Ox1 ——————– (1.33)
*Ox2 + Red2 *Red2 + Ox2 ——————– (1.34)
k12 = (k11k22K12f12)½ ————————————— (1.30)
logf12 = (logK12)2
4log(k11k22/z2) ————————————— (1.31)
Z = collision frequency generally taken as 1011 dm3 mol-1 s– 1
The theory is an adiabatic theory of electron transfer and assumes that within the
activated complex for electron transfer, the probability of electron transfer is one. It also
assumes that work terms for the self exchange and cross reactions are the same and that
k11
k
k22
K12
31
the electron transfer reagents may be treated as spherical, structureless reactants.
Other theoretical treatments and modifications of Marcus theory include those of
Hush (1961) and Sutin (1966). Good agreement between the observed rate and that
calculated by Marcus-Cross relation suggests outer-sphere mechanism (Cotton and
Wilkinson, 1980).
1.6 Research objective
Our main interest in this work is to study electron transfer reactions of
methylene blue. We intend to gain more insight into the kinetics and mechanisms of its
reactions by investigating the dynamics of its reactions with the oxyanions such as
BrO3
-, S2O8
2- and SO3
2-. This will also enable us to compare the reactions of these
oxyanions with other reported electron transfer reactions of oxyanions.
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