ABSTRACT
The kinetics of the redox reactions of crystal violet (CV+) with some oxy-anions (S2O82-,
S2O42-, MnO4-) and hydrazine dihydrochloride (N2H4.2HCl) have been studied spectrophotometrically in aqueous hydrochloric acid medium for CV+–S2O82- and CV+– S2O42- systems and aqueous tetraoxosulphate(VI) acid medium for CV+– MnO4- system, while CV+– N2H4.2HCl was not studied in acidic medium. The reactions were carried out at [H+] =1.0 × 10-3mol dm-3 (HCl) and I = 0.50 mol dm-3 (NaCl), for S2O82- and S2O42-, [H+] = 5 × 10-2mol dm-3 (H2SO4) and I = 0.50mol dm-3 (Na2SO4) for MnO4-,T = 36 ± 1ºC for S2O82-, T = 30 ± 1oC for S2O42-, T= 29 ± 1ºC for MnO4-, T = 29±2ºC for N2H4.2HCl, I = 0.50 moldm-3 (NaCl) for N2H4.2HCl and λmax = 585nm The stoichiometric studies showed that one mole of CV+ consumed one mole of peroxydisulphate ion, dithionite ion, permanganate ion and hydrazine dihydrochlorideto yield the products. The reactions were found to be first order with respect to the concentrations of crystal violet, oxy-anions and hydrazine dihydrochloride. Hydrogen ion concentration had no effect on the CV+– S2O82- system, while increase in hydrogen ion concentration of the reaction medium decreased the rate of the reaction for CV+– S2O42- system and increased the rate of the reaction for CV+– MnO4- system. First orders were obtained with respect to [H+] for both CV+– S2O42- and CV+ – MnO4- systems having the rate:
-d[CV+]/dt = k (a [H+]-1) [CV+] [S2O42-] and
-d[CV+]/dt = k (b + c [H+]) [CV+] [MnO4-] respectively.
Increase in ionic strength of the reaction medium decreased the rate of the reaction for all the systems except for CV+–N2H4.2HCl system, where increase in ionic strength of the reaction medium did not affect the reaction rate.
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Added cationscatalysed the rate of the reaction of CV+ – S2O42- and CV+ – MnO4-systems, inhibited the rate of the reaction of CV+– S2O82-system, while added anions inhibited the rate of the reaction of CV+– S2O82-, CV+ – S2O42- and CV+ – MnO4- systems.However, added ions had no effect on the rate of reaction for CV+– N2H4.2HCl system. Polymerisation test suggested the presence of free radicals for the CV+ –S2O82- and CV+ – S2O42- systems and free radicals ions found to be absent inCV+ – MnO4- and CV+ – N2H4.2HCl systems. Michaelis – Menten plots and spectroscopic studiesgave no evidence of intermediate complex formation for all the systems. The results obtained for all the systems investigated are in favour of the outer –sphere mechanism.
TABLE OF CONTENTS
Title page iii
Declaration iv
Certification v
Dedication vi
Acknowledgment vii
Abstract viii
Table of Contents x
List of Figures xiv
List ofTables xvii
List of Abbreviations xviii
CHAPTER ONE
1.0 INTRODUCTION 1
1.1 Electron Transfer Reactions 1
1.1.1 Homonuclear or isotopic electron exchange reactions 2
1.1.2 Heteronuclear or cross electron exchange reactions 2
1.2 Theories of Electron Transfer 3
1.2.1 Franck-Condon principle 3
1.2.2 The Marcus correction and correlation 3
1.3 Mechanism of Redox Reaction 4
1.3.1 The outer-sphere mechanism 4
1.3.2 The inner sphere mechanism 5
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1.4 Statement of Problem 5
1.4.1 Justification for the Study 5
1.4.2 Aim and objectives 6
CHAPTER TWO
2.0 LITERATURE REVIEW 7
2.1 Redox Reactions of Crystal Violet 7
2.2 Redox Reactions of Peroxydisulphate Ions 8
2.3 Redox Reactions of Dithionite Ions 9
2.4 Redox Reactions of Permanganate Ions 10
2.5 Redox Reactions of Hydrazine Dihydrochloride 12
CHAPTER THREE
3.0 MATERIALS AND METHODS 15
3.1 Materials 15
3.1.1 Preparation of crystal violet stock solution 15
3.1.2 Preparation of sodium peroxydisulphate stock solution 15
3.1.3 Preparation of sodium dithionite stock solution 16
3.1.4 Preparation of standard potassium permanganate solution 16
3.1.5 Preparation of 0.10 mol dm-3 hydrazine dihydrochloride solution 16
3.1.6 Preparation of salt solutions 16
3.1.7 Preparation of standard sulphuric acid solution 16
3.1.8 Preparation of standard sodium carbonate 16 3.1.9 Preparation of sodium carbonate stock solution 17
3.2 Methods 17
3.2.1 Stoichiometry studies 17
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3.2.2 Product analysis 18
3.2.3 Kinetic studies 18
3.2.3.1 Effect of hydrogen ion concentration on the reaction rate 19
3.2.3.2 Effect of ionic strength and dielectric constant of the reaction medium
on the reaction rate 19
3.2.3.3 Effect of added ions on the reaction rate 20
3.2.4 Test for free radicals (Polymerisation test) 20
3.2.5 Test for intermediate complex formation 20
3.2.5.1 Michaelis – Mentens approach 20
3.2.5.2 Spectrophotometric test 20
CHAPTER FOUR
4.0 RESULTS 21
4.1 Stoichiometry 21
4.2 Product Analysis 26
4.3 Order of Reaction 26
4.4 Effect of Hydrogen Ion Concentration on the Rate of Reaction 39
4.5 Effect of Changes in Ionic Strength of the Reaction Medium on the
Reaction Rates 44
4.6 Effect of Change in Dielectric Constant of the Reaction Medium on the
Reaction Rates 44
4.7 Effect of Added Anions on the Reaction Rate 44
4.8 Effect of Added Cations on the Reaction Rate 63
4.9 Test for Free Radicals (Polymerisation test) 63
4.10 Test for Intermediate Complex Formation 63
4.10.1 Michaelis – Mentens plots 63
4.10.2 Spectrophotometric test 63
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CHAPTER FIVE
5.0 DISCUSSIONS 83
5.1 Crystal Violet – S2O82- System 83
5.2 Crystal Violet – S2O42- System 86
5.3 Crystal Violet – MnO4- System 89
5.4 Crystal Violet – N2H4.2HCl System 92
CHAPTER SIX
6.0 SUMMARY, CONCLUSION AND RECOMMENDATION 95
6.1 Summary 95
6.2 Conclusion 96
6.3 Recommendation 96
REFERENCES 97
CHAPTER ONE
1.0 INTRODUCTION
Chemical kinetics is the study of the rate at which chemical reactions occur and it also sheds light on the reaction mechanism (Theodore et al., 1995).Redox reactions are among the most common types of reactions in chemical and biological processes. They are usually spontaneous and often accompanied by changes in oxidation state of at least two of the reactants (Purcell and Kotz, 1977). These reactions are basically of two types:
i. Reactions involving electron transfer which play important roles in chemical, biological and technological processes.
ii. Reactions involving atom transfer with or without electron transfer.
Various reactions in organic 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 effect control of a wide area of science and technology (Iyun, 1982).
1.1 Electron Transfer Reactions
Electron transfer (ET) occurs when an electron moves from an atom, molecule or ion to another atom, molecule or ion. ET is a mechanistic description of a redox reaction, wherein the oxidation state of reactant and product changes. Numerous biological processes involve electron transfer reactions. ET reaction mechanisms are of importance because of the insight they give into the actual process of transfer (Babatunde, 2005).
ET is very important in polymerisation reaction, photography, electrochemistry, photosynthesis, metabolism,respiration, plant decay, detoxificationand other applications. Based on thermodynamic parameters, ET reactions can be classified into two broad groups. These are homonuclear (isotopic) electron exchange reactions and heteronuclear (cross) electron exchange reactions (Greenwood and Earnshaw, 1997).
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1.1.1 Homonuclear or isotopic electron exchange reactions
In homonuclear electron exchange reactions, transfer of electrons occurs between two identical metal ion centers existing in different oxidation states. Examples of these reactions are given in equation 1.1 – 1.3.
*Co(NH3)62+ + Co(NH3)63+ → *Co(NH3)63+ + Co(NH3)62+ 1.1
*Cr2+ + CrCl22+ → *CrCl2 2+ + Cr2+ 1.2
Ti+ + *Ti3+ → Ti3+ + *Ti+ 1.3
In the above reactions, there is no net chemical change and the rate constant for the forward and backward reactions are equal. The reactants and products have same nuclei and concentrations, hence, the equilibrium constant = 1. The only change in free energy is due to mixing and therefore the overall free energy is approximately zero. Slow exchange reactions of this type are studied by isotopic labeling techniques and nuclear magnetic resonance. (Burgess, 1978).
1.1.2 Heteronuclear or cross electron exchange reactions
Heteronuclear or cross electron exchange reactions, involve transfer of electrons between different metal ion centers and the products are chemically distinct from the reactants. The net change in free energy in most cases is less than zero (ΔG < O). Examples of these reactions are given in equation1.4 and 1.5.
[Fe(CN)6]4-+ [IrCI6]2- → [Fe(CN)6]3- + [IrCI6]3- 1.4
2Fe2+ + Ti3+ → 2Fe3+ + Ti+ 1.5
Reactions of the type in which the oxidant and reductant change oxidation state by the same number of electrons are termed complementary and the stoichiometry of such reactions is always 1:1 as represented in equations 1.6 and 1.7.
Sn2+ + Ti3+ →Sn4+ + Ti+ 1.6
[Co (en)3]3+ + [Ru (NH3)6]2+→ [Co (en)3]2+ + [Ru (NH3)6]3+ 1.7
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When the oxidant and reductant undergo unequal changes in oxidation state and the stoichiometry is not 1:1, the reactions are termed non-complementary reactions. Examples of such reactions are given in equation 1.8 and 1.9.
2V2+ + Br2 → 2V3+ + 2Br- 1.8
Sn2+ + 2Fe3+→ Sn4+ + 2Fe2+ 1.9
1.2 Theories of Electron Transfer
1.2.1 Franck–Condon principle
In inorganic electron transfer reactions, it is established that reactants must undergo reorganisation of the hydration (solvation) shell surrounding them before electron transfer in such a way that their energies in the transition state become identical, thus minimizing the energy change on electron transfer (Reynolds and Lumry, 1966). It was Libby (1952) who accounted for this phenomenon based on the Franck–Condon principles suggesting that, before the fast electron transfer can occur, the slower nuclear rearrangement of water molecules in the hydration shells must take place, the proton being 1.836 times as massive as the electron (Ji, 2012).
The principle states that electron transfer takes place in a much shorter time (10-15s-1) than the time (10-13s-1) required for the nuclei to move hence electron transfer occurs without appreciable movement of the nuclei (Platzman and Franck, 1954; Sutin, 1966). The principle shows that the position of the nuclei remains intact during the process of electron transfer.
1.2.2 Marcus theory
The idea of Libby turns out to be incorrect, because for reactions occurring in the dark the energy is not conserved: the ions would be formed in the wrong high–energy environment, but the only way such a non–conserving event could happen would be by
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absorption of light (vertical transition) and not in the dark. To clarify this, the energy transfer reaction of equation 1.10 is considered.
[Fe(H2O)6]2+ + [Fe*(H2O)6]3+→ [Fe(H2O)6]3+ + [Fe*(H2O)6]2+ 1.10
Here the Fe–O equilibrium distance in the Fe2+ is 2.21Å while it is 2.05Å for the Fe3+. If the electron transfer takes place at their equilibrium positions, then Fe2+ would have been compressed while Fe3+on the other hand would be stretched. These are in vibrational excited states and would release energy, if the electron transfer occurs in dark, then the energy is not conserved. Thus the reactants must match their energies before electron transfer can occur. In the limit of weak electron interactions between the reactants both the Franck–Condon principle (vertical transition) and energy conservation must be satisfied. Fluctuations had to occur in the various nuclear coordination as well as in the orientation coordinates of the solvent molecules and in any other coordinates whose most probable distribution for the products differs from that of the reactants. With such fluctuations, values of the coordinates (i.e. reaction coordinates) could be reached which satisfy both the Franck–Condon and energy conservation conditions in order to permit ET to occur in the dark (Marcus, 1956).
1.3 Mechanism of Redox Reaction
The outer–sphere and the inner–sphere mechanism have been established by Taube and his co–workers as the basic mechanisms that are operative when electrons are transferred between two metal ions in solution (Taube et at., 1953).
1.3.1 The outer–sphere mechanism
This type of mechanism involves electron transfer from reductant to oxidant with the coordination sphere of each staying intact before, during and after electron transfer. Both reactants are inert with respect to substitution or one is relatively inert and does not present
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site for the labile reactants. This mechanism is one which the reactants do not form an intermediate with bridging functional group to provide a path way for electron transfer.
1.3.2 The inner–sphere mechanism
An inner–sphere mechanism is one in which the reductant and oxidant are linked together by at least one bridging ligand common to their coordination shells. In this case, ligand displacement is faster than electron transfer process (Jagannatham, 2012). The main feature of this mechanism is that, substitution takes place at one of the metal centers to give binuclear ligand – bridged species (intermediate) before electron transfer.
1.4 Statement of the Research Problem
The kinetic data on the redox reaction of crystal violet are scanty, besides there is no kinetic data on the redox reactions of crystal violet with the oxy-anions S2O82-, S2O42- and MnO4-, and hydrazine dihydrochloride. Moreover, there is no established mechanism of reactions between the dye with the oxy-anions (S2O82-, S2O42-, MnO4-) and hydrazine dihydrochloride. This has constituted a great impediment with respect to proper understanding of some important kinetic information about the dye
1.4.1 Justification for the Study
Crystal violet (CV+),also called gentian violet, is widely used as a biological stain in human and veterinary medicine as an acid–base indicator, and as an agent against infection by bacteria, fungi, pinworms and other parasites. It is also used in several industrial processes for different applications mostly as a textile dye. It is believed that CV+ is harmful to humans and animals (Hall and Hamilton, 1982).
Considerable studies have been carried out on the kinetics and mechanisms of redox reaction of CV+. These studies include CV+– BrO3- (Adetoroet al.,2014), CV+ – Cr2O72- (Mohammed and Komolafe, 2010), CV+– S2O52-, CV+– IO4-, CV+ – C1O-, CV+– C1O2- (Abdulsalam, 2015).
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In spite of these studies, the kinetic and mechanism of the redox reaction of crystal violet and the reactions involving the dye and S2O82-, S2O42-, MnO4-and hydrazine dihydrochloride have not been reported to the best of our knowledge. In view of the outlined uses and wide applications of the dye, there is a need to carry out further work on its redox reaction. The desire to gain further insight in the kinetics and mechanisms of the redox reaction of the dye motivate us to embark on this study. The kinetic data generated and the subsequent mechanistic studies will assist in the better understanding and more efficient utilisation of the dye especially as microscopic stain and in the dye industry.
1.4.2 Aim and objectives
The aim of the study is to propose mechanistic pathways for the redox reaction of crystal violet with peroxydisulphate, dithionite, permanganate ions and hydrazine dihydrochloride.
This aim will be achieved through the following objectives:
(i) Determination of the stoichiometry of the reactions,
(ii) Determination of the order of the reactions,
(iii)Effect of changes in acid concentration on the reaction rates,
(iv) Effect of added cations and anions on the reaction rates,
(v) Effect of changes in ionic strength on the reaction rates,
(vi) Investigation of intermediate complex and free radical formations in the reactions,
(vii) Product analyses.
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