Initial Rate Of The Reaction example essay topic

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 reaction 1 in EP 6.4 and then in AA 2.1. 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: o Investigate the effect of concentration for each reactant and use the results to find the rate equation for this particular reaction. o 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) S 2 O 82- (aq) + 2 I- (aq) 2 SO 42- (aq) + I 2 (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: 2 S 2 O 32- (aq) + I 2 (aq) S 4 O 62- (aq) + 2 I- (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 at different concentrations of a reactant because the rate would gradually become very similar for each concentration. As mentioned above, an ideal extent of the reaction to be studied is 12.5%. This is where the initial rate occurs. It is during this time that the rate is constant and the fastest which allows me to compare data appropriately for different concentrations while also using time efficiently.

The percentage of the reaction studied can be varied by altering the concentrations and volumes of the solutions used. In my particular reaction, the amount of iodine that would be produced if the reaction went to completion can be calculated; this would be the 100% mark on the graph above. This would depend on the concentration of the reactant not in excess in the reaction mixture, i.e. either iodide or peroxodisulphate ions. This is because once the amount of the reactant not in excess runs out then no further products can be formed and the reaction ends.

Then the amount of iodine actually produced in each reaction mixture can be calculated by the corresponding amount on thiosulphate used (as mentioned above). To work out the percentage extent of the reaction: Amount of iodine produced in the reaction mixture x 100 The amount of iodine produced if the reaction went to completion shall do this in my method to ensure a suitable percentage of reaction is being studied. Percentages higher than 15% give values for rate that are lower than they should be because although the rate starts to decrease (the line in the above graph starts to curve), it is assumed that it is constant. Percentages lower than 10% would theoretically provide more accurate results because the rate would be even more constant and the line in the above graph shows that an even 'straighter' part of the reaction would be studied. However, progressively lower percentages means using smaller concentrations which lead to larger percentage uncertainties. It would be interesting to see whether these lower percentages are more or less reliable than the standard 12.5%.

I shall decide whether I will do this or not after the trial. Before progressing onto the actual method for the investigation, I shall discuss some of the theories that explain why concentration and temperature have an effect on the reaction rate and whether there are other factors that also affect it that need to be controlled.