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