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How does the activation enthalpy and the rate of the iodine-clock reaction vary with the concentration of reactants, catalysts, and the presence of different catalysts?

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Salter's Chemistry A-Level Individual Investigation Report How does the activation enthalpy and the rate of the iodine-clock reaction vary with the concentration of reactants, catalysts, and the presence of different catalysts? Investigation How does the activation enthalpy and the rate of reaction vary with the concentration of reactants, catalysts, and the presence of different catalysts? Aim In this investigation I have set out to run a few standard tests and attempt to explain the mechanism by which the transition metal catalysts work in this particular instance. Theory The iodine-clock reaction. This reaction has been known for quite a long time, and it involves the oxidation of iodide ions to iodine molecules, which are soluble in water and will show its presence by a pale brown colour. The addition of starch to the reaction mixture makes the colour change far more visible by forming a dark-blue complex with iodine. The reaction can be represented by the following half-equation: 2I-(aq) ---> I2 (aq) + 2e- E standard /V = +0.54 2 I-(aq) > I2 (aq) + 2 e- E O /V = +0.54 ........................ Equation 1 There are a range of oxidation agents available to carry out this reaction. In fact, almost all transition metal ions could be used to oxidise iodide ions. A quick scan through a table of electrode potentials will yield the necessary list of possible oxidation agents. However, in order to investigate the catalytic properties of the transition metal ions, another type of oxidation agent had to be chosen. The standard iodine-clock experimentA1 uses the reduction of the peroxy-disulphate ions (S2O82-), for various reasons which will be described later. The half-equation of this reaction is: 2e- + S2O82-(aq) ---> 2SO42-(aq) E standard /V = +2.01 2 e- + S2O82-(aq) > 2 SO42-(aq) E O /V = +2.01 ........................ Equation 2 Therefore, the overall equation of the reaction is: 2I-(aq) + S2O82-(aq) ...read more.


For example, solid silver metal has a much lower entropy than its ionic counterparts, and therefore silver ions may not act as a catalyst in this reaction as it would require silver in the solid state be produced in an intermittent stage of the reaction. As systems naturally go towards the side with higher entropy the silver should not allow itself be reduced by iodide ions. The initial rate of reaction. As in most reactions the rate of reaction is dependent on the concentration of the chemicals it is the initial rate of reaction that we are interested in, because that's when the rate of reaction is a true reflection of the rate given by that particular concentration. This is partly why the iodine-clock reaction and the peroxydisulphate ion is chosen as a sample, as we are able to delay the expiry of the clock by adding Sodium Thiosulphate to the mixture to mop-up any iodine produced and reduce it back to iodide. The equation for this side-reaction is as follows: I2 (aq) + 2Na2S2O3 (aq) ---> 2Na+I-(aq) + Na2S4O6 (aq) I2(aq) + 2 Na2S2O3(aq) > 2 Na+I- (aq) + Na2S4O6(aq) ........................ Equation 8 Note that iodide ions (I-) are produced again on the right hand side as Sodium Iodide, replenishing the supply of iodide ions in the reaction mixture. This enables us to maintain an approximate concentration of iodide ions and enables us to measure the "initial rate" for a period, before all the Thiosulphate ions are used up. This technique will only work if the reaction of Thiosulphate with iodine is much faster than the oxidation of iodide by the persulphate ions. This appears to be the case, and may go part of the way to explain why erratic readings are sometimes obtained in reactions with Cupric ions as catalyst, because the catalyst is extremely efficient in this case, and this side-reaction may not be able to keep up with the rate at which iodine is being produced. ...read more.


Reaction Time. The human reaction time has a minimum value of 0.2 seconds (this can be measured using a simple computer programme) but often depends on whether one is anticipating the event. In order to prevent the reaction time from distorting the result during the practicals I must arrange for the timer to be unseen from where I am so that I can't anticipate when it would go blue. This error can then be quantified as a systematic error of approximately 0.5 seconds. The order in which the reactants are mixed. By using Cupric ions as catalyst we introduced another problem which I was not aware of until the practicals started: The iodide ions in the solution reacted with the Cupric ions in the solution in significant amounts - enough to alter the results quite dramatically, almost halving the reaction time needed. This is due to the fact that Cupric ions react with any iodide ions present to form a light blue precipitate, and therefore decreasing the concentration of iodide ions. The problem could be solved by changing the order in which the reactants are mixed: if I added the Cupric ions before the addition of Potassium Iodide, but after Sodium Thiosulphate has been added, the Cupric ions will be complexed by the Thiosulphate ions present, and will therefore be prevented from reacting with the iodide as the iodide enters. Again, as long as within a set of data the order of mixing stays one way, the data still can be compared across data sets because the activation enthalpy is not affected by the concentration of either iodide ions or Cupric ions. It seemed that the relative local concentration of the Cupric ions had an effect, too. When investigating the effect of concentration of the catalyst I had made up some 0.0039M Cupric Sulphate solution, and this did not seem to give the precipitate even though it is injected into the reaction mixture after the addition of iodide ions. ...read more.

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