Iodine Clock Reaction

Topics: Iodine, Chemistry, Sodium Pages: 5 (1364 words) Published: March 26, 2011

This Project Page first appeared in the November 1996 issue of Chemistry Review, Volume 6, Number 2, Pages 14 and 15. Chemistry Review is published four times during the academic year by Philip Allan Updates and is a journal for post-16 students. It contains a variety of interesting and colourful articles aimed at 16-19 year-olds taking mainly AS and A2 courses in chemistry.

NOTE: Project Page is designed to help you think about your investigation. It is not intended to be a set of instructions for practical work and does not include a list of safety precautions. CHEMISTRY REVIEW accepts no responsibility if Project Page is used in any way as a set of instructions.

Clock reactions

If you choose a project that explores the kinetics of a chemical reaction you will need a way of measuring the rate of the reaction. Clock reactions provide an interesting way of doing this for some systems.

In a typical reaction the first part of a graph showing the concentration of product against time is approximately a straight line (see Figure 1). If you choose any value of concentration that lies on this straight line (say c1) the initial rate of reaction can be found by dividing this concentration by the time taken to reach it (t1).


Figure 1

If you measure the time taken for the same concentration to be reached in a series of reactions, you will be finding the time for the same amount of product to be formed for each reaction. The shorter the time, the faster the reaction is occurring. You can therefore take 1/t as a relative measure of the initial rate of reaction.

The trick, of course, is knowing when the fixed amount of product has been formed. The following examples illustrate how this can be done.

Appearing blue
There are a number of so called 'iodine clock' reactions in which molecular iodine is one of the products. Probably the most famous of these is the reaction involving hydrogen peroxide and iodide ions in acid solution:

H2O2 + 2I– + 2H+ → I2 + 2H2O

The kinetics of this reaction was first investigated by Vernon Harcourt and William Essen, and the reaction is still referred to as the Harcourt-Essen reaction. In your project you will need to add the same, fixed amount of sodium thiosulphate solution together with a little starch solution to your reactants in each experiment. The molecular iodine produced by the main reaction between hydrogen peroxide, iodide and acid immediately reacts with the thiosulphate ions:

I2 + 2S2O32–- → S4O62– + 2I–

When all of the thiosulphate has been used up, the iodine accumulates in the solution and reacts with the starch to give a distinctive blue-black colour. The time from mixing the reactants to the appearance of the blue colour is therefore the time for a fixed amount of iodine to be formed. The appearance of the blue colour is like a motor race chequered flag; it tells when a particular amount of product has been formed however long it takes for this to happen. You can simply use 1/t as a measure of the rate of reaction. Now that you have a method of monitoring the rate of reaction you can look at how different factors affect it. You could separately change the concentrations of each reactant, you could try the experiment at different temperatures or investigate the effects of a catalyst such as ammonium molybdate(VI). Another iodine clock reaction is that between peroxydisulphate(VI) and iodide ions:

S2O82– + 2I– → 2SO42– + I2

Again, the same, small, fixed amount of thiosulphate ions and some starch solution are added to the reaction mixture in each experiment. If you measure the time from mixing to appearance of the blue-black starch-iodine complex you can again use 1/t as your initial reaction rate. As well as looking at the effects on the rate of changing concentrations and ternperature, you might like to explore which d block ions catalyse the reaction. When you first...
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