ABSTRACT
The kinetics and mechanisms of the electron transfer reactions of diaquotetrakis (2, 2’-
bipyridine)-μ-oxodiruthenium(III) ions (hereafter denoted as Ru2O4+or [(H2O)2Ru2O]4+ )
and thiourea (TU),N–methylthiourea (MTU), N–allylthiourea (ATU), N,N’-
dimethylthiourea(DMTU) (collectively denoted by TSH), thiosulphate ions (S2O3
2–),
dithionite ions (S2O4
2–), hypophosphorous acid (H3PO2), methanol (CH3OH), ethanol
(C2H5OH) and propanol (C3H7OH) (collectively denoted as ROH) were studied in
aqueous perchloric acid (HClO4) medium (except for S2O4
2– and the alcohols) at I =
0.5 mol dm–3 and T = (31.0 ± 1) °C. The stoichiometry was found to be 1:2 (Ru2O4+/
reductant) in the TSH and S2O3
2–systems but 1:1 in the S2O4
2–, H3PO2 and ROH
systems. The rate of reaction is first order in oxidants and reductants for all the systems.
Addition of acid has inverse effect on the rates of reaction for the TSH system, but
direct dependence for the S2O3
2– system. The overall rate equation for the TSH reaction
can be given as :
– [Ru2O4+] = (a + b[H+]-1 )[Ru2O4+][Reductant] .
while that for S2O3
2– reaction can be given as:
– [Ru2O4+] = (a + b[H+])[Ru2O4+][S2O3
2-]
Varying the ionic strength (I) and dielectric constant (D)of the reaction medium had no
effect on rates of reaction for the TSH, S2O4
2-, H3PO2 and ROH reactions, while for the
S2O3
2– reaction, increase in I led to a decrease in rate while decrease in D led to increase
in rate of reaction. Added ions had no effect on the TSH and H3PO2 reactions but led to
vii
catalysis and/ or inhibition in S2O32–, S2O42- and ROH systems. Free radicals were identified in the TSH, H3PO2 and ROH reactions only. Spectroscopic information and the results of the Michaelis – Menten plots suggest the lack of formation of intermediate complex prior to electron transfer for all the systems. For all the reactions, [(H2O)2(bipy)2Ru]2+ was found to be the product of Ru2O4+ reduction. For the TSH reaction, disulphides was found to be the oxidation product of TSH while, for H3PO2 reaction, test for phosphorous acid (H3PO3) presence was positive and for the ROH system, test for aldehydes was also positive. The order of reactivity for the TSH system is ATU > DMTU > MTU > TU, while the order of reactivity for the sulphur oxyanions’ reactions is S2O42-> S2O32– and that for the ROH reactions is CH3OH > C2H5OH > C3H7OH. Based on the Michaelis–Menten plots and the interactions with added ions, all the reactions are proposed to have proceeded through the outer–sphere electron transfer mechanism. Plausible mechanisms for all the reactions have been proposed.
viii
TABLE OF CONTENTS
Title Page ii
Declaration iii
Certification iv
Acknowledgements v
Abstract vi
Table of Contents viii
List of Tables xiv
List of Figures xv
List of Abbreviations xxi
CHAPTER ONE
INTRODUCTION 1
1.1 Electron TransferReactions 1
1.2 Types of Electron Transfer Reactions 2
1.2.1 Homonuclear or self-exchange electron transfer reactions 2
1.2.2 Heteronuclear or cross-electron transfer reactions 2
1.3 Mechanisms of Electron Transfer 3
1.3.1 The innersphere mechanism 4
1.3.2 The outersphere mechanism 5
1.3.3 Proton couple electron transfer process 6
1.3.4 Solvated electron theory 6
1.4 Diagnosis of Redox Reaction Mechanism 7
1.4.1 Comparison of the rate of electron transfer and
the rate of substitution (kredox versus ksubs) 7
1.4.2 Identification of binuclear intermediate 8
1.4.3Product analysis 8
ix
1.4.4 Use of ambidentate ligands 9
1.4.5 Effect of added ions 10
1.4.6 Reactivity pattern 11
1.4.7 Activation Parameters 11
1.5Effect of Pressure on Electron Transfer 12
1.6Rate Monitoring Techniques 12
1.6.1 Conventional methods 13
1.6.2 Fast reaction techniques 13
1.7 Justification of the Study 15
1.8 Aims and Objectives of the Study 18
CHAPTER TWO
LITERATURE REVIEW 20
2.1 Ruthenium Chemistry 20
2.1.1 Properties of ruthenium polypyridyl complexes 20
2.1.2 Electron transfer in ruthenium complexes 22
2.1.3 Photosynthesis and water oxidation catalysis 24
2.1.4 Homogenous water oxidation by ruthenium complexes 25
2.1.5 Heterogeneous water oxidation catalysis by
ruthenium complexes 29
2.2 Thiourea and Its Derivatives 30
2.2.1 Electron transfer reactions of thiourea and its derivatives 31
2.3 Electron Transfer Reaction of S2O32– 36
2.4 Electron Transfer Reactions of S2O42– 37
2.5 Electron Transfer Reactions of H3PO2 39
2.6 Electron Transfer Reactions of ROH 42
x
CHAPTER THREE
MATERIALS AND METHODS 44
3.1 Materials and Reagents 44
3.1.1 Diaquotetrakis (2, 2’- bipyridine)-μ-oxo-diruthenium(III) perchlorate 44
3.1.2 Thiourea solution 45
3.1.3 N-methylthiourea solution 45
3.1.4 N,N’- dimethylthiourea solution 45
3.1.5N- allylthiourea solution 45
3.1.6 Sodium thiosulphate solution 45
3.1.7 Sodium dithionite solution 46
3.1.8 Hypophosphorous acid solution 46
3.1.9 Sodium perchlorate solution 46
3.1.10 Magnesium chloride solution 46
3.1.11 Ammonium chloride solution 47
3.1.12 Sodium acetate solution 47
3.1.13 Sodium formate solution 47
3.1.14 Sodium nitrate solution 47
3.1.15 Perchloric acid solution 47
3.1.16 Methanol 48
3.1.17 Ethanol 48
3.1.18 Propan–1–ol 48
3.2 Methods 49
3.2.1 Stoichiometric studies 49
3.2.2 Kinetic measurements 50
3.2.3 Effect of changes in hydrogen ion concentration 51
xi
3.2.4 Effect of changes in ionic strength 51
3.2.5 Effect of changes in dielectric constant of reactionmedium 52
3.2.6 Effect of added ions 52
3.2.7 Test for free radicals 52
3.2.8 Test for Intermediate complex formation 53
3.2.9 Product analysis 53
CHAPTER FOUR
RESULTS 55
4.1 Stability of Ru2O4+ 55
4.2 Stoichiometric Studies 55
4.3 Determination of Pseudo – First Order and Second
Order Rate Constants and Order of Reaction 61
4.3.1 Ru2O4+ reaction with thiourea and its derivatives
(TU, MTU, ATU and DMTU) 61
4.3.2 Reaction of Ru2O4+ with S2O32– and S2O42– 72
4.3.3 Reaction of Ru2O4+ with H3PO2 83
4.3.4 Reaction of Ru2O4+ with ROH 90
4.4 Effect of Changes in Hydrogen Ion Concentration, [H+] 98.
4.4.1 Reaction of Ru2O4+with TU, MTU, ATU and DMTU 98
4.4.2 Reaction of Ru2O4+ with S2O32– 106
4.4.3 Reaction of Ru2O4+ with H3PO2 109
4.5 Effect of Changes in Ionic Strength of Reaction Medium 109
4.5.1 Reaction of Ru2O4+ with thiourea and its derivatives 109
4.5.2 Reaction of Ru2O4+ with S2O32– 109
4.5.3 Reaction of Ru2O4+ with S2O42– 110
xii
4.5.4 Reaction of Ru2O4+ with H3PO2 and ROH 110
4.6 Effect of Changes in Dielectric Constant of Reaction Medium 110
4.6.1 Reaction of Ru2O4+ with TU, MTU, ATU and DMTU 110
4.6.2 Reaction of Ru2O4+ with S2O32– 112
4.6.3 Reaction of Ru2O4+ with S2O42–, H3PO2 and ROH 119
4.7 Effect of Added Ions 119
4.7.1 Reaction of Ru2O4+ with TU, MTU, ATU and DMTU 119
4.7.2 Reaction of Ru2O4+ with S2O32– 119
4.7.3 Reaction of Ru2O4+ with S2O42– 134
4.7.4 Reaction of Ru2O4+ with H3PO2 145
4.7.5 Reaction of Ru2O4+ with the ROH 145
4.8 Tests for the Formation of Intermediate Complex 157
4.8.1 Spectroscopic test 157
4.8.2 Michaelis – Menten plots 157
4.9 Free Radicals Test 168
4.10Product Analysis 168
4.9.1 Reaction of Ru2O4+ with TU, MTU, ATU and DMTU 168
4.9.2 Reaction of Ru2O4+ with H3PO2 168
4.9.3 Reaction of Ru2O4+ with ROH 169
CHAPTER FIVE
DISCUSSION 170
5.1 Ru 2O4+ Reaction with TU, MTU, ATU and DMTU 170
5.2 Ru2O4+ Reaction with S2O32– 180
5.3 Ru2O4+ Reaction with S2O42– 186
xiii
5.4 Ru2O4+ Reaction with H3PO2 189
5.5 Ru2O4+ Reaction with ROH 198
5.6 Comparison of the Redox Reactions of Ru2O4+ and the
Various Reductants under Study 198
CHAPTER SIX
SUMMARY, CONCLUSION AND RECOMMENDATIONS 200
6.1 Summary 200
6.2 Conclusion 201
6.3 Recommendation 201
REFERENCES 202
APPENDIX 221
CHAPTER ONE
INTRODUCTION
1.1 Electron Transfer Reactions
A fundamental understanding of electron transfer (ET) reactions is vital to the inorganic chemist in the context of energy transduction, corrosion processes, metallurgy, redox processes in environmental chemistry, metalloenzymes and metalloproteins involved in ET (Adman, 1979; Bennet, 1972). ET reactions of transition metal complexes are accompanied by a change in the oxidation state of the metal atom and the overall charge on the complex ion.
Since the late 1940s, ET reactions’ has grown greatly in biochemical processes. This development, as well as its relation to the study of other kinds of chemical reactions, represents a very interesting history in which many unanswered puzzles have been pieced together.
ET reactions play very important roles in many biological processes including collagen synthesis, steroid metabolism, the immune response, drug activation, neurotransmitter metabolism, nitrogen fixation, respiration and photosynthesis (Prince and George, 1990). The latter two processes are fundamentally significant as they provide most of the energy that is required for the maintenance of life. The extraction of energy from organic compounds, carried out by several catabolic pathways (e.g. the citric acid cycle) involves the oxidation of these compounds to CO2 and H2O with the concomitant production of water-soluble reductants (Simondsen and Tollin, 1980). These reductants
2
donate electrons to components of the mitochondrial electron-transfer chain, resulting in
the reduction of oxygen to water.
1.2 Types of Electron Transfer Reactions
1.2.1 Homonuclear or self–exchange electron transfer reactions
Transfer of electrons can occur between two identical metal ion centres existing in
different oxidation states, in what is usually referred to as homonuclear or self–
exchange electron transfer reactions. For such reactions, the reactant and product ion
concentrations are the same and hence the overall free energy of the reaction is
approximately zero with the possibility of very small differences in the free energies of
the reactant and products existing as a result of mixing. Isotopic labelling techniques
have been employed to monitor the rate of such electron exchange reactions when they
are slow. For fast exchange reactions, however, the nuclear magnetic resonance
spectroscopic technique is employed (Meyer and Taube, 1987; Wilkins, 1974).
Examples of self–exchange reactions are depicted in Equations 1.1 – 1.3.
V2+ + *V3+ V3+ + *V2+ …(1.1)
FeEDTA2- + *FeEDTA- FeEDTA- + *FeEDTA2- …(1.2)
PtCl4
2- + *PtCl6
2- PtCl6
2- + *PtCl4
2- …(1.3)
1.2.2 Heteronuclear or cross electron transfer reactions
Here, different metal ion centres are involved in electron transfer reaction. The
difference in the concentrations of the reactants and the products observed in such
reactions implies that there must be a net change in the free energy of the reaction, i.e.
ΔGө ≠ 0. It has been found that for most reactions, ΔGө< 0. However, in some reactions
3
ΔGө> 0 and the driving force for such reactions is provided by the necessity to maintain
the equilibrium concentration of certain intermediate species in one or more of the
reaction steps (Ramaraj et al., 1986a, 1986b). Examples of such electron transfer
reactions in which ΔGө< 0 are represented by Equations 1.4 – 1.6,
[Co(NH3)5X]2+ + Cr2+ + H+
CrX2+ + Co2+ + 5NH4
+ ..(1.4)
FeIIcyt – c + Ru(NH3)6
3+ FeIIIcyt – c + Ru(NH3)6 2+ …(1.5)
Co(en)3
3+ + Ru(NH3)6
2+ Co(en)3
2+ + Ru(NH3)6
3+ ..(1.6)
Heteronuclear electron transfer reactions can be either complementary, in which the
oxidant and the reductant undergo equal changes in oxidation state. Equations (1.4) to
(1.6) above are examples of complementary reactions. On the other hand, noncomplementary
reactions are those in which there are unequal changes in oxidation
states by the oxidant and the reactant. Examples of non–complementary reactions are
given in Equations (1.7 – 1.9).
[Fe(LL)3]3+ + 2I-
[Fe(LL)3]2+ + I2 (1.7)
2V2+ + Br2 2V3+ + 2Br- ..(1.8)
2Fe2+ + Tl3+ 2Fe3+ + Tl+ . . (1.9)
1.3 Mechanisms of Electron Transfer
In order to state the mechanism of electron transfer, after knowing that a particular
electron transfer is a self-exchange reaction (homonuclear) or a cross-reaction
(heteronuclear), complementary or non–complementary, certain very important
considerations have to be made:
4
(1) Comparison of the rate of electron transfer with the rate of substitution into the inner coordination sphere of the reactant ions.
(2) The possibility of identifying or inferring the closeness of approach of reactant metal centres in the activated complex prior to electron transfer.
(3) Is the reaction thermodynamically feasible? If not, what drives the reaction?
(4) How many electrons are transferred at a time? (if more than one electron is transferred).
(5) What dictates the interpretation of the acid and base catalysis in some of these reactions? Do the structures of the reactants, transition state products form a basis for rationalisation?
(6) Are the large variations of rates which usually accompany the variation of reactant ions indicative of the effect of the variation of electronic structure of reactant ions on the rate and mode of electron transfer?
The fundamental classification scheme for the stoichiometric mechanism of redox reactions was established in the 1950s by the Nobel Prize winner in Chemistry of 1983, Henry Taube (Taube, 1959). Taube delineated the innersphere mechanism against the outersphere mechanism for electron transfer reactions.
1.3.1 The innersphere mechanism
Taube in the 1950s recognised that many oxidation–reduction reactions occur by a ligand–bridging mechanism in which substitution of the coordination shell of one of the metal ions occurs. An intermediate or a bridged binuclear complex is therefore formed in which a ligand forms a bridge between one metal ion and the other metal prior to electron transfer. Electron transfer processes occurring through this bridging ligand are
5
classified as innersphere electron transfer reactions (Lagford, 1979). The bridging
ligand is usually, but not always, transferred from one reactant to the other.
1.3.2 The outersphere mechanism
Since electrons are capable of passing through barriers such as ligands via quantum
mechanical process of tunnelling, some electron transfers could proceed without any
change in the coordination spheres of either reactant. Consider the reactions in
Equations 1.10 and 1.11.
[Fe(bipy)3]3+ + [Co(tpy)2]2+-
Fe(bipy)3]2+ + [Co(tpy)2]3+ ..(1.10)
W(CN)8
4- + FeIII(CN)6
3-
WV(CN)8
3- + FeII(CN)6
4- ..(1.11)
Reactions, such as in Equations (1.10) and (1.11), during which the ligands remain
coordinated in the primary spheres of the respective central atoms (i.e. no ligand
exchange or substitution occurs), are classified as outersphere electron transfer reactions
(Taube et al,, 1953). The coordination shells of the two reactants remain essentially
intact before and after the electron transfer (Larsen and Wahl, 1965). When either or
both the oxidant and the reductant are inert to substitution during the time sufficient for
electron transfer to take place, outersphere electron transfer is most likely to take place.
The activated complex that results during the electron transfer process is designated ‘the
outersphere activated complex’.
Outersphere reactions between complexes of different metals are usually very fast: for
these reactions are accompanied by a decrease in standard free energy.
6
1.3.3 Proton couple electron transfer process
In this model, a proton and electron are transferred simultaneously during redox reaction. The study of this model has received a great deal of attention and consideration in the past two decades by researchers (Ghosh et al., 1994; Lohdip and Iyun, 1998 and 2003; Ukoha and Iyun, 2002; Babatunde and Iyun, 2004; Lohdip and Ogara, 2004; Lohdip and Shamle, 2004). The importance of this coupling between movement of proton, which may be a proton transfer (PT) reaction, and electron transfer (ET) is a well elaborated theme in the study of a number of biological assemblies (Babcock et al., 1989; Wikstrom, 1989; Morgan and Wikstrom, 1991). The coupling of PT and ET is referred to as proton coupled electron transfer (PCET) reaction (Turro et al., 1992). Coupling between proton motion and electron transfer is a fundamental mechanism for energy conversion in biological and chemical systems. The translocation of protons across biological membranes in the proteins involved in photosynthesis (Babcock et al., 1989; Okamura and Feher, 1992) and respiration (Wikstrom, 1989; Babcock and Wikstrom, 1992; Malmstrom, 1993) is based on PCET.
1.3.4 Solvated electron theory
A solvated electron is a free electron in solution. Solvated electrons are widely occurring and represent a more diffuse charge distribution than do electrons in molecules. The solvated electron theory presents a mechanism of electron transfer that occurs in non-aqueous media like NH3 and halides (Ayoko, 1981). The basic assumption of this theory is that there is ejection of an electron by the reducing agent into the solvent which solvates the electron and holds it until an oxidising agent picks it up.
7
1.4 Diagnosis of Redox Reaction Mechanisms
Assigning outersphere or innersphere to redox reactions can be adequately done by
considering some criteria.
1.4.1 Comparison of the rate of electron transfer and the rate of substitution (kredox
versus ksubs)
Comparison of the rate of electron transfer versus the rate of substitution into the inner
coordination shell of the more labile of the reactants can be used to assign mechanism
of the redox reaction since in the innersphere mechanism, substitution into the inner
coordination shell of one of the reactant ion precedes electron transfer.
When the rate of electron transfer (rate of redox process) is much greater than the rate of
substitution i.e. when kredox >>> ksubs, the outer–sphere mechanism of electron transfer is
implicated (Rossenheim et al., 1974), while if kredox<<< ksubs, the innersphere
mechanism is said to be operating. However, if kredox = ksubs, both the innersphere and
the outersphere mechanisms are simultaneously in operation (Endicott and Taube,
1964). In the electron transfer reaction between Fe(phen)3
2+ and *Fe(phen)3
3+ (Larsen
and Wahl, 1965) given in Equation (1.12), the rate constants for the substitution of phen
Fe(phen)3
2+ + *Fe(phen)3
3+- Fe(phen)3
3+ + *Fe(phen)3]2+ ..(1.12)
for *Fe(phen)3
2+ and Fe(phen)3
3+ are 7.5 x 10-5 s-1 and 5.01 x 10-5 s-1, respectively,
while the rate constant for electron exchange is 105 dm3 mol-1 s-1. Since in this reaction
kred » ksub, it has been explained in favour of outersphere. This was also observed for
8
the electron exchange reaction between Fe(CN)64- and Fe(CN)63- (Campion et al., 1967; Deck and Wahl, 1954).
1.4.2 Identification of binuclear intermediate
Innersphere mechanism is said to be in operation if there is an evidence of the presence of a binuclear complex either as a transient or stable intermediate along a pathway between reactants and products (Haim, 1983). In the reduction process involving Fe(OH2)62+ (Haim and Sutin, 1966) and V(OH2)62+ (Espenson, 1967; Price and Taube, 1968), detection of transient intermediates from ligand transfer processes was possible by the use of rapid flow techniques. In the course of the reaction between VO2+ and Cr2+, an intermediate, VOCr4+ was characterized (Espenson, 1965). The possibility of identifying binuclear intermediates in the inner–sphere mechanism occurs when reduced oxidant and the oxidised reductant are relatively inert to substitution. Initial investigations suggest that most of the intermediates formed are successor complexes. For example, halide-bridged successor complexes have been found in the reactions of IrCl62- with Co(CN)52- (Grossman and Haim, 1970) and IrBr62- and Cr(OH2)62+ or Co(CN)52- (Melvin and Haim, 1977) and the rate of decomposition of the intermediate is slower than the rate of electron transfer. In situations where the binuclear intermediate is transient in nature, their presence can only be confirmed indirectly from kinetics data or from empirical rate law.
1.4.3 Product analysis
In the innersphere mechanism where there is the formation of a bridging ligand, transfer of the latter from one metal ion to another provides an empirical proof that the
9
innersphere mechanism is in operation. Taube and co-workers (Taube et al., 1953) have
shown that reaction (Equation 1.13) occurs.
[Co(NH3)5X]2+ + Cr2+
(aq) + 5H+ [Cr(H2O)5X]2+ + Co2+ + 5NH4
+ ..(1.13)
where X = Halide, SO4
2-, NCS-, N3
-, PO4
3-, P2O7
2-, CH3COO-, CH3C7CO2
-, succinate,
maleate and oxalate.
Also, the same researchers reported that in the redox reaction involving Co(NH3)5Cl2+
and Cr2+, the CrCl2+ product identified has been associated with the binuclear
intermediate (NH3)5Co–Cl–Cr4+ proffering evidence in favour of innersphere
mechanism (Taube et al., 1953).
1.4.4 Use of ambidentate ligands
The difference in the rate of electron transfer observed with symmetrical and
unsymmetrical bridging ligands such as azide, thiocyanate, and isothiocyanate provide a
basis for making a distinction between the innersphere and the outersphere mechanisms.
If there is no difference between the ratio of the rate constant for the reduction of the
azide complex to that of the reduction of the thiocyanate complex, (kN/ kNCS ) or for
isothiocyanate and thiocyanate complex (kSCN/ kNCS ), then the outer sphere
mechanism is suspected to be operating. This scenario is observed in the outer-sphere
reduction of [Co(NH3)5NCS]2+ and [Co(NH3)5N3]2+ by [Cr(bpy)3]2+ , [Ru(NH3)6]2+,
[Ru(en)3]2+ and [Ru(NH3)5OH2]2+ , in which there is no discrimination between NCS-
and N3
– (Haim, 1983). However, if the ratios of the rate constants are dissimilar, then
10
the innersphere mechanism is suspected (Caldin et al., 1964; Espenson, 1965; Sutin, 1966).
1.4.5 Effect of added ions
The rates of redox reactions are reported to be markedly affected by added ions for a variety of reactions known to occur by the outersphere mechanism. The effects of added ions on the rates of reactions occurring through the innersphere pathway are not so glaring. Thus, observation of the dependence of rate of redox reaction on added cation or anion has been used to distinguish between the two pathways (Lohdip et al., 1998; Ukoha and Iyun, 2001 and 2002; Ukoha and Ibrahim, 2004). However, addition of Cl– can lead to serious complication in the interpretation of observed rates and kinetics (Thakuria and Gupta, 1975; Adegite et al., 1977). It should be noted that outersphere electron transfer reactions are, theoretically, easier to determine than innersphere reactions because in the latter bond breaking and bond forming steps are pronounced. In the outersphere reaction, the redox nuclei must be sufficiently close to enhance electronic interaction which results in the delocalisation of exchanging electrons. Consequently, reactions proceeding by outersphere mechanism can be catalysed by added ions which shorten the distance to which electron can be transferred (Meyer and Taube, 1987). Added cations affect the electron transfer reaction involving two positive nuclei by the cation repelling away the oxidant, thus reducing the degree of repulsion between the redox centres. This enhances the electron transfer reaction. Also, added anions enhance the electron transfer reaction by the anion coordinating between the redox centres by electrostatic attraction. However, when two oppositely charged ions are involved, added ions could retard the rate of electron transfer because coordination to any of the reactants could reduce the degree of attraction between the redox centres.
11
This, expectedly, would increase the distance between the redox partners thereby slowing down the electron transfer process (Ukoha, 1999).
1.4.6 Reactivity pattern
Reactivity patterns can successfully be used to assign mechanisms to redox reactions, some of which include:
(1) Reactivity pattern with a wide range of reactants
Comparing the rates of reactions when similar complexes like Co(NH3)5X2+ (where X = Cl–, F–, Br–or NO3–) are reacted with a particular reductant gives an idea into which of the mechanisms is operating. Since the outersphere mechanism does not depend on the identity of the bridging ligands, it means that the outersphere mechanism is implicated when the rates of the reactions are similar. Conversely, when the rates of reactions are different, the innersphere is suspected since the rates of the reactions in the inner–sphere mechanism depend on the nature of the bridging group (Sykes, 1967; Shea and Haim, 1973)
(2) Relative rates of reactions of hydroxyl and aquo complexes
Since the hydroxyl group (OH–) is a better bridging ligand than water (H2O), it
follows that hydroxo complexes are expected to react faster via the innersphere
mechanism. Therefore, where kOH<< kH2O, the outer – sphere mechanism is
said to be operating while the converse applies when kOH>> kH2O (Endicott
and Taube, 1964).
1.4.7 Activation parameters
There seems to be no strong correlation between activation parameters ΔH#, ΔG# and ΔS# and the type of mechanism operating in a particular redox process. However, the
12
signs or the magnitudes of the activation parameters could give a clue as to which mechanism is inherent in a reaction. Formation of precursor complex as in innersphere mechanism is indicated by a negative ΔH# (Wilkins, 1974). However, despite the difference in mechanisms for the reaction of Cr2+ and V2+ with Ru3+ complexes, the ΔS# are almost the same. Measurements of the volume of activation (ΔV#) for the reduction of various complexes has been applied as diagnostic tool in reaction kinetics (Hubbard et al., 1991).
1.5 Effect of Pressure on Electron Transfer
The value of mechanistic information that emerges from kinetic measurements over a series of elevated pressures for solution reactions in inorganic chemistry has been realised for some time (Stranks, 1974). However, many inorganic reactions are too fast to follow using conventional instrumentation. The momentum regarding investigation at high pressures vis-à-vis organic reactions was delayed somewhat until adaptation of rapid reaction techniques for operation at high pressures had been achieved, mostly in the period from 1975 to 1985 (Kotowski and van Eldik, 1989; van Eldik et al.,1989; van Eldik and Merbach, 1992). Studies on the effect of pressure on electron transfer reactions can assist the elucidation of the intimate reaction mechanisms of such processes.
1.6 Rate Monitoring Techniques
A variety of experimental techniques have been developed and used to investigate redox reactions (Sykes, 1966, 1967; Rosseinsky, 1972; Wilkins, 1974). The range of techniques used depends on how fast the reactions go. Many authors have extensively
13
reviewed the techniques (Stranks, 1960; Caldin, 1974; Wilkins, 1974), which are summarized below:
1.6.1 Conventional methods
The conventional methods involve the measurement, as a function of time, the concentration of one or more of the reactants or products or any physical property like absorbance which is directly related to the concentration. Conventional methods, which are used to monitor slow redox reactions with half – lives approximately 30 seconds, include photometric, spectrometric, polarometric and radiometric techniques. Other physical properties that are directly related to concentration are conductivity, pressure changes, changes in volume, refractive index, optical activity etc.
1.6.2 Fast reaction techniques
These techniques are used to monitor fast reactions with half – lives of about a millisecond. These techniques are in turn classified into (a) flow, (b) relaxation and (c) resonance methods.
(a) Flow Method
This technique involves the mixing of separate solutions of two reactants inside a mixing device. Different flow techniques exist, depending on the treatment the reaction mixture is subjected to (Caldin et al.,1964). The commonest flow methods used are:
(i) Continuous Flow Method
Here, the reaction mixture flows continuously along an observation tube
while conventional monitoring is made either at different points along the
observation tube with the flow maintained at constant rate or at a fixed
14
point on the tube with varied flow rate. In either case, a series of values
for the extent of reaction is obtained at different times. These values
constitute the kinetic data required.
(ii) Stopped – Flow Method
In this method, the solutions are mixed together in a mixing device and as the mixed solutions flow along the tube, it is abruptly stopped so that the solution comes to rest within a few milliseconds. The rate of the flow along the observation tube is such that when the solution is stopped, a segment within 1 cm of mixing device has been mixed for only 1–2 milliseconds. Any reaction taking place in this segment of solution is then monitored and the signal relayed to an oscilloscope.
(iii) Quenched – Flow Method
In this method, the reaction solution is quenched after a pre –
determined time and the quenched solution is analysed by any
conventional method.
(b) Relaxation Method
Very fast reversible reactions with half – lives (t1/2) as short as 10–9 seconds are
measured using the relaxation methods (Eigen and De Maeyer, 1963). Any equilibrium attained during the reaction is perturbed by sudden variation of a physical parameter such as temperature, pressure, or electric field intensity and the time taken to readjust to a new equilibrium is monitored, which is then related to the rate constants of the forward and reverse reactions. The dependence of a particular equilibrium on a chosen external physical parameter largely determines the method of perturbation used for such equilibrium and hence these
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methods are referred to as ‘temperature–jump’, ‘pressure–jump’ or ‘electric field–jump’.
(c) Resonance Methods
Burlamacchi et al. (1967) reported the use of both nuclear magnetic resonance (N.M.R) and electron spin resonance (E.S.R) in the study of rates of reactions. In the N.M.R. technique, the resonance absorption line is related to the life-time of the nucleus in a given state, while in E.S.R it is related to the life-time of paramagnetic species in a given energy state. Line broadening is always the result of any reduction of life-time of these states by a chemical interaction. By adding an increasing amount of a reagent, measurement of the corresponding increase in width of the line due to the second reagent can be made. Hnmr line broadening, for example, has been used to measure the rate of exchange of several mono- and bidentate nitrogen and oxygen donor ligands coordinated to MnII, FeII, CoII, NiII and CuII (Caldin et al., 1964).
1.7 Justification of the Study
Transition metal systems in which metal ions are linked by a bridging ligand can differ significantly with regard to the nature and/ or extent of metal – metal interactions. For nonorganometallic first row transition metal ions’ systems like Fe–O–Fe, relatively weak metal – metal interactions are commonly found. The nature and extent of the interactions have been studied extensively using magnetic techniques (Weaver et al., 1975). Although interactions do exist, the component ions usually have chemical and electronic properties similar to the properties expected for isolated monomeric complexes. The same strong modification in properties can also occur in ligand-bridged complexes if the metal-metal interaction across the bridging ligand is sufficiently
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strong. Griffiths (1970) reported that oxo – bridged complexes of the first row transition element series, like Fe–O–Fe, exhibit only weak metal-metal interactions, but the interactions seem to be much stronger for the second and third rows, consecutively. Compared to the Ru and Os system in which the extent of metal-metal interaction between Ru(III) ions and Os(III) ions is sufficient to change appreciably the chemical and electronic properties of the dimeric ions, the metal – metal interactions between the Fe(III) ions is apparently insufficient to change the properties of the dimeric ions to an appreciable extent. This is probably because of a greater d-orbital radial extension for Ru(III) and Os(III) when compared to Fe(III) (Weaver et al., 1975). Consequently, the Ru(III) and Os(III) have been grouped in Class II of Robin and Day (1967) classification, while the Fe–O–Fe system has been classified in Class III of the Robin and Day Scheme. An example of a Ru–O–Ru system is the blue dimer, [(H2O(bipy)2Ru–O–Ru(bipy)2H2O]4+ or diaquotetrakis (2,2’- bipyridine)-μ-oxodiruthenium(III) ions. This dimer has generated a lot of interest due to its versatility, including its ability to mimic photosynthesis. The catalytic oxidation of water and chloride with the blue dimer was reported by the Meyers research team (Gersten et al., 1982; Gilbert et al., 1985). Consequent upon Meyers and his co–researchers report on the catalytic activity of the dimer, many studies have been carried out on the mechanism of this catalysis with CeIV and CoIII as oxidants (Nagoshi et al., 1999; Binstead et al., 2000; Yamada et al., 2001; Meyer and Huynh, 2003). Theversatility of the blue dimer cuts across its potential use in diverse areas such as photosensitisers for photochemical conversion of solar energy (Meyer, 1990; Balzani et al., 1996; Kalyanasundaram and Gratzel, 1998; Juris et al., 1998; Hammerstrom et al., 2000 and Islam et al., 2003), molecular electronic devices (Barigelleti and Flamigni, 2000; El–Ghayoury et al., 2000; Mishra et al., 2003 and Newkome et al., 2004).
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The oxo-bridged ruthenium dimer has been synthesised, characterised and a lot of kinetic data have been documented as regards its redox behaviour (Weaver et al., 1975; Iyun et al.1992a, 1992b, 1992c, 1992d, 1995a, 1995b, 1996). It has been reported that in oxo–bridged systems having the Ru–O–Ru moiety, chemically significant interactions exist between the Ru atoms through the bridging ligand (Meyer, and Huynh, 2003). Such compounds manifest unusual physical and chemical properties. This is possibly as a result of the strong Ru–O–Ru interaction which makes the separate Ru3+ ions lose their identity since the valence levels which largely determine these properties are molecular orbitals delocalised over the Ru–O–Ru linkage.
The use of thiourea and its derivatives, thiosulphates ions, dithionite ions (all oxyanions of sulphur), hypophosphorous acid (phosphorous oxyacid) and the alcohols as reductants of choice in this work is based on the role they play both in chemical and biochemical systems. Thiourea and its derivatives have been used as effective scavenger of reactive oxygen intermediates (Fox, 1984). Due to their reducing properties, they have been used in the textile industry (Arifoglu et al., 1992), as corrosion inhibitors (Ayres, 1970) and in industrial equipment such as boilers which develop scales due to corrosion. Besides these, several thiourea derivatives have various agricultural uses as fungicides, herbicides and rodenticides and industrial uses which include applications in rubber industries as accelerators, and in photography as fixing agents and to remove stains from negatives. Gaining an insight into the redox pattern of this important class of reductants would be of immense value to knowledge, thus including them as reductants in this study. Thiosulphatess and dithionites are oxyanions of sulphur which are good reducing agents. The reducing property of the thiosulphate ions has made it gain wide application in photographic processing as a fixer and also used in gold extraction. By
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this study, we hope to promote the knowledge of the redox behaviour of these sulphur oxyanions. The peculiar reduction property of hypophosphorous acid as reported by Caroll and Thomas (1966), involving the active H3PO2 species obtained by acid induced activation of the inactive H3PO2 through a reversible tautomeric shift has made hypophosphorous acid attractive to us for this study. Various possibilities exist in the oxidation of alcohols. These possibilities include mechanisms that involve hydride ion transfer to the oxidant prior to the rate determining decomposition of the intermediate to products (Lee and Spitzer, 1975; Sengupta et al., 1986; Lee and Congson, 1990; Nimbalkar and Chavan, 1998; Saraswat et al., 2003; Dharmaraja et al, 2008; Kothari and Banerji, 2011; Dhage et al., 2013 and Bijudas, 2014), attack of ClO2– on an α – H in ethanol in the oxidation of ethanol by ClO4– (Gaswick and Krueger, 1969) and initial equilibrium formation of chromate ester followed by a rate determining decomposition to products (Yusuf et al., 2004). The above highlights have stirred up our interest in investigating the kinetics and mechanisms of oxidation of methanol, ethanol and propanol by the oxo–bridged ruthenium dimer. Understanding the mechanisms of these reactions may assist in the improvement of breadth analysis for detecting the level of alcohol in the system. It is our hope that redox kinetics and mechanistic studies involving alcohols like earlier ones (Iyun and Ukoha, 1999; Yusuf et al., 2004) will assist in further understanding of these complex but seemingly simple reactions.
1.8 Aim and Objectives of the Study
The aim of the research is to generate kinetic data which would give an insight into the mechanisms of the electron transfer reactions of diaquotetrakis (2,2’- bipyridine) – μ – oxodiruthenium (III) ions and some reductants. The findings would contribute to knowledge of the redox behaviour of this versatile μ-oxo–bridged ruthenium complex
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and μ-oxo–bridged systems, in general. The objectives to achieve the aim of the study include the following:
i. To synthesise and characterize the diaquotetrakis(2,2’-bipyridine)-μ-oxodiruthenium(III) complex ions.
ii. To determine the stoichiometries of the reactions.
iii. To determine the pseudo–first order and second order rate constans of the reactions.
iv. To determine the order of reactions’ with respect to each of the reductants.
v. To determine the effects of changing the hydrogen ion concentration on the rates of the reactions that took place in acid medium.
vi. To determine the effects of changes of ionic strength and dielectric constant of reaction medium on the rates of the reactions.
vii. To determine the effect of added ions to the reaction medium on the rates of reactions.
viii. To determine the participation of free radicals in the various reactions.
ix. To determine the formation of intermediate complexes in the course of the reactions.
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