RESEARCH ARTICLE The Kinetics of Oxidation of Iodide ion by Dichromate Ion in an Acidic Medium

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1 International Journal of Advanced Chemical Research Vol. 4, No. 10, PP , October RESEARCH ARTICLE The Kinetics of Oxidation of Iodide ion by Dichromate Ion in an Acidic Medium 1,*Isa baba koki and 2 Kolo A. Madu 1 Department of Chemistry, Northwest University Kano PMB 3220 Kano, Nigeria 2 Department of Pure and Industrial Chemistry, Abubakar Tafawa Balewa University Bauchi, PMB 0248 Bauchi, Nigeria Abstract The kinetics of dichromate-iodide reaction in an acidic medium was studied by iodometric technique. The rate has been studied by appropriate choice of concentration conditions at constant temperature and ionic strength. The kinetic studies showed half order reaction with respect to dichromate ion and first order both for iodide ion and hydrogen ion concentrations. The observed rate law is given: d[l 2 ] = K[Cr 2 O 7 ] ½ [I - ][H + ] dt The value of the rate constant was 8.67mol ½ S 2 L -½ The dependence of reaction rate on temperature was monitored and Ea, A, ΔS, ΔH and ΔG were evaluated to be KJmol -1, 2.0x10 12, KJmol -1, KJmol -1 and KJmol -1 respectively, which implies that the reaction is spontaneous with negative value of ΔG, and Exothermic. Key words: Kinetics, Oxidation, Hydrogen, Iodide, Dichromate, Concentration. INTRODUCTION The subject of chemical kinetics is concerned with the quantitative study of the rates of chemical reactions and of the factors upon which they depend such as concentration, temperature etc or Chemical kinetics is the study and discussion of chemical reactions with respect to reaction rates, effect of various variables, re-arrangement of atoms, formation of intermediates etc. Of more fundamental interest are those kinetic studies of chemical reactions in which the objective is to arrive at the reaction mechanism. Of more basic significance are studies whose object is to throw light on the general principles of reactivity (Laidler, 1973). The rates of most chemical reactions are sensitive to the temperature, so in conventional experiments the temperature of the reaction mixture must be held constant throughout the course of the reaction (Atkins and De Paula, 2006). Potassium dichromate (Orange) is used as a primary standard in volumetric analysis for estimating reducing agents: Cr 2 O 7 (aq) +14H + (aq)+6e - For example I - I 2,Fe 2+ 2Cr 3+ (aq)+7h 2 O (l) Fe 3+ etc It has the advantage over potassium permanganate that it is more stable in solution, and since it oxidizes chloride ions in dilute solution much more slowly, it can be used in the presence of dilute hydrochloric acid. However, dichromate titrations require an indicator (Rogers,1982). *Corresponding author: Isa baba koki, Department of Chemistry, Northwest University Kano PMB 3220 Kano, Nigeria Therefore in experimental studies of kinetics, as much information as possible should be obtained about reaction intermediate, a product, including how their concentration changes with extent of the overall reaction (Rogers, 1982). The order of reaction is strictly an experimental quantity, and merely provides information about the way in which the rate depends on concentration. The experimental dependence of rate of reaction on concentration of the reactant is the order of that reaction (Laidler, 1973). The oxidation of chromium (vi) and chromium (v) is almost similar in which a single oxidation product is produced, both valency of chromium must react in a similar way, even though detail of the mechanism may be different (Espenson and King, 1963). In some few cases induced oxidation occurs when the selectivity s of chromium (vi) and chromium (v) are different. In oxidation reaction, chromium usually undergoes a valence change from Cr (VI) to Cr (III), which is a three electron change only making the oxidation a non-complementary one. A two electron transfer process appears most reasonable similar to that proposed by Haight for the Cr (VI)-S (IV) system (Haight, 1965). Far less likely one of the four electron transfer producing Cr (II) and the one electron transfer corresponding to a rate limiting Cr (VI) Cr (V) There is a considerable evidence to show that it is reduction to the Cr (IV) state, involving a change from four fold to six fold coordination which is rate limiting (Espenson and King,1965; Espenson, 1964).

It should be made clear that the kinetic data do not allow any choice to be made in the relative order of addition of H + and I - nor do they discriminate between attack by iodide at the Cr centre as

2 International Journal of Advanced Chemical Research 045 In oxidation reaction chromium usually undergoes a valency change from (VI) to (III). It is therefore obvious that an intermediate chromium valence state to appear and take part in the oxidation process (Gaswick and Kruger, 1969). It should be made clear that the kinetic data do not allow any choice to be made in the relative order of addition of H + and I - nor do they discriminate between attack by iodide at the Cr centre as opposed to attack at an oxygen centre. Possible steps leading to the activated complex are: ICrO 3 + H + + I - rate determining the species I - CrO 3 seems a likely intermediate by analogy with the known Cl - CrO 3 species. The nucleophile I - can have an appreciable reactivity at Cr (VI) is an evidence by the catalysis of the hydrolysis of Cr 2 O 7 by thiourea (Perlmutter and Wolff, 1965). There is convincing evidence that I 2 does induce the O 2 oxidation of I - in acidic aqueous solution. Further support for the existence of below reaction, Cr IV + I - Cr III + I are derived from a comparison of the results with respect to retardation of the reaction by added iodine. The study of intermediate oxidation state established that oxidation by chromium (VI) is in fact an oxidation by chromium (V) to an extent of 67% (Rocek and Rich, 1966). The presumed steps of Cr (VI) - I - reaction are derived from the behavior of systems in which induced oxidation of iodide occurs (Espenson, 1964). A study was carried out to provide a direct examination of this important reaction. 2HCrO I - +14H + +4H 2 O 2 Cr[H 2 O] I 2 An earlier study of this reaction was made (Rocek and Rich, 1966) and the rate law was reported as, -d[hcro 4 - ] = [HCrO 4 - ] [K 1 [H + ][I - ]+K 2 [H + ] 2 [I - ] 2 ]. dt It was then felt desirable to work at low concentrations of Cr (VI) to minimize some of the complexities occurring in the system (Tong and Johnson, 1966). Rocek and Radkowsky (1968) had presented evidence that the classical sequence is not operative in the chromium (VI) oxidation of cyclobutanol. This work therefore determined the oxidation of iodide ion by dichromate ion in an acidic medium at constant temperature and ionic strength. The concentration studies of the reaction enabled us to determine the respective orders of the reaction. From the orders of the reaction obtained, the observed rate equation for the oxidation reaction was then established. Temperature dependence of the reaction was also monitored to evaluate the Arrhenius parameters. Also the effect of ionic strength on the rate of the reaction was studied to be able to determine the nature of the activated complex of the reaction (the changes of the species involved in the activated complex) and its influence on the rate of the reaction. MATERIALS AND METHODS The order of the reactants must be known first before determining the rate law. If the rate of the reaction doubled, when the concentration is doubled then it is first order. Also if the rate quadruples when the concentration is doubled, the order of the reaction is second in one of the reactant that varied. This means that the rate of formation of iodine is proportional to the first or second power of the concentration of reactant that varied. A small amount of sodium thiosulphate is added, it is known that thiosulphate ion react extremely rapidly with iodine as shown by the equation. 2S 2 O 3 (aq) + I 2(aq) S 4 O 6 (aq) + 2I - (aq) Thiosulphate convert any I 2 formed to I - (Mc Alpine, 1945).The presence of iodine in the solution can be made visible by the addition of 10cm 3 3% starch solution, which with iodine gives a blue colour. As iodine is formed it will quickly be consumed by thiosulphate ion. The moment all the thiosulphate is used up, iodine will increase and give a blue colour. Determination of iodine formation dependence on concentration was carried out by preparing series of solutions containing different initial concentrations of the reactant. And to each solution a constant amount of Sodium thiosulphate solution was added and with a stop clock, the time required to consume this amount of thiosulphate is measured. The appearance of blue colour of the starch indicator indicates the completion of the reaction. We therefore know the time required to produce a specific amount of iodine. Procedure for Determination of Rate law Six burettes were used to dispense the reagents, 20cm 3 each of Sulphuric acid and Potassium dichromate solutions were measured into the first 250ml beaker. 20cm 3 each of Potassium Iodide and Sodium thiosulphate, and 5cm 3 Sodium Sulphate solutions were measured into the second beaker and 10cm 3 3% starch solution was added. The two 250ml beakers were placed side by side on the display table, and a stirring rod was placed in the second beaker, and quickly the content of the first beaker was poured into the second beaker and stirred gently to increase mixing of the two solutions. The time for the reaction was noted from the time of mixing to the first appearance of the blue colour. Procedure for Temperature Studies The effect of temperature change on the rate of redox reaction was studied in this experiment. The same volume and concentration of reagents used in the previous experiment were used. The procedure followed was also the same as the previous experiment except that the temperature was varied. The temperatures used were 15 0 C, 20 0 C and 35 0 C. Individual runs were carried out on the water bath maintained at 15 0 C. The solutions in the first and second beaker were mixed and placed in the water bath and timing was started immediately until the first appearance of the blue colour. The time for the reaction was then recorded. Reactions were carried out at the remaining temperature following the same experimental procedure.

3 International Journal of Advanced Chemical Research 046 RESULTS Table 1 showed the effect of dichromate ion concentration on rate, the rate of the reaction increases by half when the concentration of dichromate ion is doubled. The order of the reaction was determined from a plot of log rate vs log [Cr 2 O 7 ] as shown in figure 1. The linearity of the plot indicates that the reaction is half order with respect to Cr 2 O 7 concentration under the given experimental condition. The plot gave a slope 0.57, showing that the reaction is half order. The least square plot of the result also confirmed a half order with respect to dichromate ion concentration. The rate of the reaction was found to increase with increase in hydrogen ion concentration. The rate doubled when the hydrogen ion concentration was doubled. The order with respect to hydrogen ion concentration was obtained from a plot of Rate vs log [H + ] as shown in figure 3. Table 1. Effect of [Cr 2 O 7 ] on rate [I 2 formation] RUN [Cr 2O 7 ]M RATE (MS -1 ) x x x x x x x x x x10-6 I = M, [Na 2SO 4] = 1.0x10-2 M, [Na 2S 2O 3] = 2.0x10-4 M, [H 2SO 4] = 2.0x10-2 M, [KI] = 5.0x10-3 M Figure 2. Plot of LOG rate versus LOG [I-] Table 3. Effect of [H + ] on rate [I 2 ] Formation] RUN [H + ]M RATE (MS -1 ) 1 5.0x x x x x x x x x x10-5 I = M, [Na 2SO 4] = 1.0x10-2 M, [Na 2S 2O 3] = 2.0x10-4 M [K 2Cr 2O 7] = 5.0x10-4 M, [KI]= 5.0x10-3 M Figure 1. Plot of log rate versus LOG [Cr 2 O 7 ] The dependence of rate of reaction on iodine ion concentration is shown in Table 2. The rate of the reaction doubled when iodide ion concentration is doubled. The order with respect to iodine ion concentration was obtained from a plot of log rate vs. log [I - ] as shown in figure 2. The plot was linear and the slope of the plot was approximately one, which implies that the order with respect to iodine ion is first order. The least square plot confirmed a first order reaction. Table 2. Effect of [I] on rate [I 2 Formation] RUN [I - ]M RATE (MS -1 ) 1 2.5x x x x x x x x x x10-5 I = 0.049M, [Na 2SO 4] = 1.0x10-2 M, [Na 2S 2O 3] = 2.0x10-4 M, [K 2Cr 2O 7] = 5.0x10-4 M, [H 2SO 4] = 2.0x10-2 M LOG [H + ] Figure 3. Plot of LOG rate versus LOG [H + ]

4 International Journal of Advanced Chemical Research 047 Table 4. Effect of Temperature on rate [I 2 ] Formation] RUN TEM (K) RATE (MS -1 ) K exp x x x x x I = M, [Na 2SO 4] = 1.0x10-2 M, [Na 2S 2O 3] = 2.0x10-4 M [H 2SO 4] = 2.0x10-2 M, [K 2Cr 2O 7] = 5.0x10-4 M, [KI] = 5.0x10-3 M Table 5. Effect of Ionic strength on rate [I 2 ] Formation] RUN I(M) RATE (MS -1 ) ( ) K x x x x x Na 2SO 4] = 1.0x10-2 M, [Na 2S 2O 3] = 2.0x10-4 M, [H 2SO 4]=2.0x10-2 M [K 2Cr 2O 7] = 5.0x10-4 M, [KI] = 5.0x10-3 M The slope of the plot is 0.96, which implies that the order of the reaction with respect to hydrogen ion concentration is first. The least square plot confirmed a first order rate law. The rate of the reaction increases with increase in temperature as shown in the above table. A plot of log K exp against the reciprocal of temperature was made. From the slope of the plot, the Arrhenius parameters were evaluated. The effect of ionic strength on rate the reaction was investigated in the range of M M. In all the reactions, the concentrations of other reactants except the Na 2 SO 4 were kept constant. The rates constant were found to decrease with increasing ionic strength. The reaction predicts a negative kinetic salt effect that is decreasing rate constant with increasing salt concentration. The log of rate constant K was plotted against the square root of total ionic strength. The slope of the plot was found to be -2 as in figure 5. DISCUSSION The kinetic studies of the system revealed half order for dichromate ion concentration, and first order both for iodide ion and hydrogen ion concentrations, The observation rate law can be written as: 1/T X 10 4 Figure 4. PLOT OF LOG K exp VERSUS 1/T d[i 2 ] = K [Cr 2 O 7 ]½ [I - ] [H + ] dt The result on iodide and hydrogen ion concentrations are similar to that reported in Chromic acid oxidation of iodide ion over a range of hydrogen ion concentration from 0.3M to 1.0M (Vandegrift and Rocek, 1977). The reaction rate was first order with respect to [HCrO 4 - ], [I] and [H + ]. Similar [H + ] dependence on reaction rate has been established in the redox reaction of Toluidine blue with acidic Bromate and 4-oxy-4- arylbutanoic acid respectively (Jonnalagadda and Musengiwa, 1998) and (Cherkupally and Padma, 2010). The result is also similar to that reported in oxidation of Iodide ion by Manganese oxide over P H range of (Patricia et al., 2009). The rate of the reaction varied with both P H and MnO 2 concentration, with faster oxidation occurring at lower P H and higher MnO 2 concentration. The reaction rate was first order with respect to [H + ], [I - ] and [MnO 2 ]. Under the experiment conditions employed for this study, reaction rate dependence on temperature was studied and Ea, A, ΔS, ΔH and ΔG were evaluated to be KJmol -1, 2.0x10 12, KJmol -1, KJmol -1 and KJmol -1 respectively, which implies that the reaction is spontaneous with negative value of ΔG, and Exothermic. I Figure 5. Plot of LOG K Versus Conclusion The kinetic studies was investigated for dichromate ion and iodide ion, kinetic evidence has been obtained for the reaction with half order for dichromate ion, and first order for both hydrogen ion and iodide ion concentrations. The effect of

5 International Journal of Advanced Chemical Research 048 ionic strength on the rate of redox reaction was studied; the ionic strength was adjusted by varying the concentration of the sulphate to minimum, since increasing the concentration of any of the reactant increases the ionic strength. The effect of temperature on the rate of redox reaction was studied and the activation energy was evaluated to be KJmol -1 and frequency factor was found to be 2.0x The method involved in this research work is of the advantage that it can be employed easily without the use of any sophisticated instruments. REFERENCES Atkins,P. and DePaula, J Physical Chemistry, 8 th edition, Oxford University Press, NewYork, Pp. 79 Cherkupally, S.R. and Padma, S.M Kinetics and Mechanism of Acid Bromate Oxidation of Substituted 4- Oxoacids. Indian Journal of Chemistry. Vol. 49A, PP Espenson, J. H Mechanism of reaction of Vanadium (IV) and Chromium (VI) and of the Espenson, J. H. and King, E. L Kinetics and Mechanism of Uncatalyzed Chromium (VI) induced oxidation of iodide ion. Journal of American Chemical Society 86:5101. Gaswick, D. C. and Kruger, J. H The oxidation of iodide ion by some transition metals. Journal of American Chemical Society 9:2240. Haight, G. P Effect of added Mn(II) ion on the rate of reaction. Journal of AmericanChemical Society87: Jonnalagadda, S.B. and Musengiwa, N Kinetics and Mechanism of Reaction of Toluidine Blue with acidic Bromate.International Journal of Chemical Kinetics.30: Laidler, K. J Chemical Kinetics, 2 nd edition; Pp. 1-5 McGraw Hill Publishing Company, New Delhi. Mc Alpine R. K Journal of Chemical Education 22(8): oxidation of iodide. Journal of American Chemical Society 85:2228. Patricia M. F., James A. D., and George W.L The kinetics of Iodide oxidation by Manganese Oxide. GeochemicaetcosmochimicaActa. Vol 73 Issue 10: Perlmutter, H. B. and Wolf M.A Catalysis of the hydrolysis of Cr 2 O 7 by Thiourea. Canadian Journal of Chemistry 43:2913. Rocek, J. and Radkowsky, A. E Chromium (VI) oxidation of cyclobutanol. Journal ofamerican Chemical Society 90: Rocek, J. and Rich, S.A Tetrahedron letters, Pengamon Press, Great Britain, No. 14 Pp Rogers, M. J. et. al., Chemistry Facts, Patterns and Principles. First edition; Pp , , Addidison- Wesley Publishers Limited, London. Tong, J. Y. and Johnson, R. I Chronic acid oxidation of iodide ion. Journal of Inorganic Chemistry 5:1902. Vandegrift, G. F. and Rocek, J Chromic acid oxidation of iodide ion. Journal of American Chemical Society 99(1):143. Westheimer, F. H Induced Oxidation of Iodide ion by Transition Metals. Journal of Physical Chemistry 45:419. *******

Hydrated nickel(ii) salts are green in colour. Give the electron configuration of a nickel(ii) ion and hence state why the ion is coloured.

1. This question is about nickel compounds. (a) Hydrated nickel(ii) salts are green in colour. Give the electron configuration of a nickel(ii) ion and hence state why the ion is coloured.......... (b)