Rates of reaction - kinetics (A2)
Read our general notes on Risk Assessment
I am doing an investigation of the rate of reaction between manganate(VII) and oxalate ions. However, today when we tried it we came upon a whole load of problems. Firstly, we found increasing the concentration of the reactants increased the time for the solution to turn colourless. Surely this goes against everything in the collision theory. Also, before we left school, we found that even when using 0.1 M solutions of ethanedioic acid, sulphuric acid, and iodine solution, with a 0.02 M solution of potassium manganate(VII) produced similar titres when we remove samples of the mixture at three 30-second intervals after the reaction started. What are we doing wrong? We can't seem to obtain any titres that are useful and the reaction occurs more slowly at higher concentrations.
Igloo writes ...
It isn't clear from your question whether or not you are increasing the concentrations of these two ions at once. I suspect that you may be.
Rather than using a titration approach I strongly suggest that in all your experiments you try adding a fixed volume of manganate(VII) ions to a mixture of a fixed volume of oxalate and hydrogen ions (both of which are in excess), so that after the addition has taken place you start the clock and then time how long it takes for the colour of the MnO4- ions to disappear.
For example, you could add 10 cm3 of manganate(VII) ions of concentration 0.02 mol dm-3 to 10 cm3 of oxalate ions of concentration 0.50 mol dm-3 mixed with 10 cm3 of hydrogen ions of concentration 0.50 mol dm-3. Since the oxalate and hydrogen ions are 25 times more concentrated than the manganate(VII) you can see that they're in excess. You do need to carry out a trial beforehand to make sure that, at room temperature, the colour disappears in a reasonable time. If not, you will have to alter the concentrations accordingly. I'll leave you to overcome this potential problem.
Right, that's the introduction. Now, you could start off by keeping [C2O42-] and [H+] constant, and vary [MnO4-]. If you double the concentration of manganate(VII) ions, you have twice as many of these coloured ions to destroy, so, if all other concentrations remain the same, and if the reaction rate is independent of [MnO4-], i.e. if the reaction is zero order wrt [MnO4-] the time will double. If the time remains the same, then, since twice as many of these ions were destroyed, it must be going twice as fast, so you can say that the reaction is first order wrt to [MnO4-].
In another set of experiments you can keep [MnO4-] constant and vary [C2O42-]. On these occasions, the same number of coloured ions needs to be destroyed, so you should be able to work out how the times should respond if the reaction is (1) zero order, or (2) first order wrt [C2O42-].
You should now be able to apply similar reasoning to the effect of altering [H+].
I do not know for certain the kinetics of this particular reaction, but I am fairly sure that it not be more than first order wrt to each of the reactants.
I assume that you know that this reaction is autocatalytic, and I hope that you are intending to study this effect. This phenomenon is best followed using the titration approach, but, hopefully, buoyed with success from your first set of experiments (!) you should now be able to work things out for yourself.
Always carry out a risk assessment and check with your teacher before starting any practical work.
Risk assessment
Before attempting any practical work based on the advice and suggestions on this website, you must do the following. Identify any hazards, assess the risks from these hazards, and then decide appropriate control measures to reduce the risks. You must have these approved by those in authority in your school or college laboratory. Do not rely on what is said on this website.
For further guidance see our tutorial on Risk Assessment.
back to Rates of reaction - kinetics (A2)
Rate this page or react
Share your views on this page, 0 ratings so far
updated: 23 February 2007
