Rates of reaction - kinetics (A2)
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Hi, I am investigating the reaction between KI and H2O2 in the presence of a strong acid. I am looking to find the order of reaction, the rate constant, and the enthalpy of activation. I am having problems finding a value for the rate constant.
Ulex replies
Here's how the student tackled the problem:
1. To find the order of reaction with respect to each reactant I used a colorimetric method, measuring the time it took to reach a set value of absorbance (of iodine) whilst varying the concentrations of each reactant in turn. Plotting graphs of 1/time vs concentration, gave me me three graphs all indicating first order.
2. To find the rate constant, k, for the reaction, using the equation: rate = k[H2O2]1[I-]1[H+]1, I planned to take my rate value from one of my graphs of "rate" against concentration. Using this value, I would obtain the rate constant for a set of known reactant concentrations. However, I obtained my values for the ‘rate’ from absorbance/time, which is only proportional to rate, not equal to it, with a unit of s-1. For the rate equation, the rate must be in mol dm-3s-1.
To enable me to calculate the change in concentration (of iodine) in the measured time, I would need to use the Beer-Lambert law, which relates absorbance, A, to concentration (c) as follows: A = e x l x C, where l is the pathlength of the sample and e is the molar absorbtivity of iodine. This would allow me to calculate the concentration of iodine at the level of absorbance reached in the set time. Knowing the time for this change in concentration (from zero), I would be able to calculate the rate of reaction.
My problem is that I don't know the value for e, so I cannot calculate the change in concentration of iodine. I believe that there are experimental methods of calculating the value for e. If you know the value or where to find it, or can help with any other suggestions I would be most grateful.
Ulex: I do not think there is any way of calculating a value for e, it is not a universal constant but only one which applies to this particular solution, wavelength(s) of the filter and the instrumentation.
Fortunately, the minimum requirement for calculating a value for e is to know the absorbance at any one known iodine concentration. Your best bet is to beg a further half-hour's time with the colorimeterand a standard solution of iodine in KI. If no such solution of accurately known concentration is available you will need to standardise it by titration against standard sodium thiosulphate, which will, of course, take a little longer. Starch solution as indicator would be desirable but by no means essential.
There are two points which are worth bringing out in your subsequent account:
1. It would obviously be better to use several measurements with different known concentrations of iodine solution, preferably in the region of absorbance which you used as your target in your experiments.
2. (Which is actually linked to 1). Beer's Law, postulating a linear relationship between absorbance and concentration, is only approximately true in real cases. Fortunately, your results are based on only one absorbance value so they are not in error even if Beer's Law is not obeyed. The method I have suggested for finding e does, however, depend on the linear relationship.
Good luck with this. Your method is ingenious and seems to have worked well. If you are not able to find a value for e, I have every confidence that you will be able to put up a convincing case to explain away your inability to do so!
I have just read through your account again and have noticed that you wish to find the activation energy. You have, presumably, realised that you don't need to have k in the correct units in order to do this. Using values which are proportional to k will enable you to do an Arrhenius plot and get exactly the same value of Ea as you would have got with 'proper' values of k. In one sense, therefore, obtaining the 'proper' values is an irrelevant side-issue. How's that for a convincing argument?!
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updated: 09 March 2007
