The Iodide-Persulphate Reaction: the Effect of Concentration on Reaction Rate
Summary
This experiment will measure the rate of oxidation of iodide ions by persulphate ions to derive the rate law for the reaction. Starch will be added to the reaction to facilitate the measure of time during the reaction. The reactant solutions will contain (NH4)2SO4 and KI, represented as:
(NH4)2S2O8 + 2KI -> I2 + (NH4)2SO4 + K2SO4
This can be simplified to:
S2O82- + 2I- -> I2 + 2SO42-
These equations can only be carried out and be visible after the iodine has completely reacted with thiosulphate added – two moles of thiosulphate for every mole of iodine. Once all the thiosulphate has been used up in the reaction, the colour will start to appear.
Results and Conclusions
1. Contained in the following chart:
• Concentrations are calculated by multiplying the concentrations by the amount of liters of that solution divided by the total volume.
• -∆S2O8-2 is calculated by dividing the concentration of the S2O3-2 by the number of moles required, as seen in the chemical equation (2 moles)
• ∆t was determined during the lab
• Rate is calculated by dividing ∆ S2O8-2 by ∆t
• Rate constant ‘k’ is calculated by using the equation: rate = k[S2O8-2]m[I-]n
o m and n are the orders of the reaction with respect to S2O8-2 and I-respectively. When solved based on ratios of each reaction to the change in rate, they both equate to 1. The graphs included later will confirm this.
o The rate equation for k is therefore: k = rate / [S2O8-2][I-]
• Ionic strength is calculated using the equation: u = 0.5åCiZi2
o Ci is the molar concentration of the ith species
o Zi is the ionic charge
Table 2
Run #[S2O8-2]
[M][I-]
[M][S2O3-2]
[M]-∆S2O8-2
[M]∆t
[s]RateRate Constant kIonic Strength
This experiment will measure the rate of oxidation of iodide ions by persulphate ions to derive the rate law for the reaction. Starch will be added to the reaction to facilitate the measure of time during the reaction. The reactant solutions will contain (NH4)2SO4 and KI, represented as:
(NH4)2S2O8 + 2KI -> I2 + (NH4)2SO4 + K2SO4
This can be simplified to:
S2O82- + 2I- -> I2 + 2SO42-
These equations can only be carried out and be visible after the iodine has completely reacted with thiosulphate added – two moles of thiosulphate for every mole of iodine. Once all the thiosulphate has been used up in the reaction, the colour will start to appear.
Results and Conclusions
1. Contained in the following chart:
• Concentrations are calculated by multiplying the concentrations by the amount of liters of that solution divided by the total volume.
• -∆S2O8-2 is calculated by dividing the concentration of the S2O3-2 by the number of moles required, as seen in the chemical equation (2 moles)
• ∆t was determined during the lab
• Rate is calculated by dividing ∆ S2O8-2 by ∆t
• Rate constant ‘k’ is calculated by using the equation: rate = k[S2O8-2]m[I-]n
o m and n are the orders of the reaction with respect to S2O8-2 and I-respectively. When solved based on ratios of each reaction to the change in rate, they both equate to 1. The graphs included later will confirm this.
o The rate equation for k is therefore: k = rate / [S2O8-2][I-]
• Ionic strength is calculated using the equation: u = 0.5åCiZi2
o Ci is the molar concentration of the ith species
o Zi is the ionic charge
Table 2
Run #[S2O8-2]
[M][I-]
[M][S2O3-2]
[M]-∆S2O8-2
[M]∆t
[s]RateRate Constant kIonic Strength