Investigation into the Kinetics of the Reaction Between Peroxodisulphate(vi) Ions and Iodide Ions

PLAN

Introduction:

After having built up knowledge about the kinetics of reactions I decided to do an investigation in this area. I was initially introduced to this particular reaction1 in EP6.4 and then in AA2.1. I was interested in using this reaction as a means of potentially supporting and quantifying some of the theories that I have studied along with also perhaps extending on them.

Aim:

Using a clock reaction I shall:

• Investigate the effect of concentration for each reactant and use the results to find the rate equation for this particular reaction.

• Investigate the effect of temperature on the rate and use the results to find the activation enthalpy for this particular reaction.

Background detail

The Reaction:2

The reaction I am studying is often referred to as an ‘iodine clock reaction.’ A clock reaction is where the time taken to form a definite, small amount of a product at the beginning of a reaction is recorded to work out the rate.

This reaction involves the oxidation of iodide ions to iodine molecules which are soluble in water and are visible as a pale brown clear solution. The formation of the iodine can easily be detected because all other species in the reaction mixture are colourless. The addition of starch to the reaction mixture further enhances the colour change by forming a dark blue-black complex with the iodine. The overall ionic equation is: (the spectator ions K+ have been left out to see the electron transfer clearly)

S2O82- (aq) + 2I- (aq)  2SO42- (aq) + I2 (aq)

The initial rate of the reaction can be measured by measuring the time it takes to produce a fixed small amount of iodine in the reaction as mentioned above. This can be done by adding thiosulphate ions into the reaction system which instantaneously revert the iodine molecules to iodide ions. When the amount of thiosulphate ions run out, iodine is produced and there is a sudden colour change. A sudden colour change makes the time required for the iodine to be produced very obvious. This reaction is shown in the equation:

2S2O32- (aq) + I2 (aq)  S4O62- (aq) + 2I- (aq)

The total amount of iodine produced in the reaction mixture can be calculated by the equivalent amount of thiosulphate added to the reaction mixture. This way the rate can be measured in concentration of iodine produced per unit time rather than just as a reciprocal of time. This is important because it enables me to work out the rate constant, k, in the rate equation which I will discuss later.

The extent to which the reaction is studied can also be controlled using the clock method. It is generally accepted that clock reactions work best within the initial 10-15% of a reaction. The graph below explains this further.

This is called the progress curve for a reaction. The rate is calculated by the gradient of the line. As you can see, the rate decreases as the reaction progresses because the amount of reactants in the mixture starts to decrease which results in a corresponding decrease in the amount of iodine formed. When I am doing my investigation it would be inappropriate to follow the reaction to completion. This would not only take a very long amount of time but also will not allow me to compare the rates of reaction