7. Measuring Rate of Reaction – Initial Rate Method
Iodine Clock
• Hydrogen peroxide reacts with iodide ions to form iodine: H2O2 + 2H+ + 2I- → I2 + 2H2O
• The thiosulfate ion immediately reacts with the iodine: 2S2O32- + I2 → 2I- + S4O62-
• Once all of the thiosulfate ions have been reacted, excess I2 will remain in the solution.
• The presence of this I2 is indicated by a reaction with starch, forming a blue-black colour.
Method
1) Measure out the required volumes of KI, Na2S2O3, starch, H2SO4 & H2O and add them into a
conical flask.
By varying the concentrations of a reactant (i.e. I -), and keeping the other reactants at a constant
concentration, you can determine the order of reaction with respect to that reactant.
2) Measure out the required volume of H2O2 into a test tube.
H2O2 is the last reagent to be added, as the reaction will begin immediately after.
3) Pour the H2O2 into the flask and immediately begin timing.
4) Stir the mixture.
To ensure that there is an equal concentration of reactants throughout the solution.
5) Record the time taken for the blue-black colour of the starch-iodine complex to appear.
Analysis
1) Calculate 1/time (s-1).
This is a measure of the rate of reaction.
2) Plot a graph of 1/time vs. [I-].
The gradient of the graph will show the order of the reaction with respect to [I -]
Control Variables
• Temperature (the value of the rate constant changes as temperature changes)
• How well the solution is mixed (affects the rate of reaction)
Safety Hazards
• Hydrogen peroxide is corrosive.
• Sulfuric acid is an irritant.
• Iodine is harmful.
Initial Rates
• The initial rate can be calculated from the gradient of a concentration vs. time graph at time = zero.
• Initial rates are preferable since the exact concentrations of all reactants are known.
• If the concentration of one reactant is kept in a large excess in a reaction with several reactants, that
reactant will appear not to affect the rate, since its concentration stays virtually constant.
Iodine Clock
• Hydrogen peroxide reacts with iodide ions to form iodine: H2O2 + 2H+ + 2I- → I2 + 2H2O
• The thiosulfate ion immediately reacts with the iodine: 2S2O32- + I2 → 2I- + S4O62-
• Once all of the thiosulfate ions have been reacted, excess I2 will remain in the solution.
• The presence of this I2 is indicated by a reaction with starch, forming a blue-black colour.
Method
1) Measure out the required volumes of KI, Na2S2O3, starch, H2SO4 & H2O and add them into a
conical flask.
By varying the concentrations of a reactant (i.e. I -), and keeping the other reactants at a constant
concentration, you can determine the order of reaction with respect to that reactant.
2) Measure out the required volume of H2O2 into a test tube.
H2O2 is the last reagent to be added, as the reaction will begin immediately after.
3) Pour the H2O2 into the flask and immediately begin timing.
4) Stir the mixture.
To ensure that there is an equal concentration of reactants throughout the solution.
5) Record the time taken for the blue-black colour of the starch-iodine complex to appear.
Analysis
1) Calculate 1/time (s-1).
This is a measure of the rate of reaction.
2) Plot a graph of 1/time vs. [I-].
The gradient of the graph will show the order of the reaction with respect to [I -]
Control Variables
• Temperature (the value of the rate constant changes as temperature changes)
• How well the solution is mixed (affects the rate of reaction)
Safety Hazards
• Hydrogen peroxide is corrosive.
• Sulfuric acid is an irritant.
• Iodine is harmful.
Initial Rates
• The initial rate can be calculated from the gradient of a concentration vs. time graph at time = zero.
• Initial rates are preferable since the exact concentrations of all reactants are known.
• If the concentration of one reactant is kept in a large excess in a reaction with several reactants, that
reactant will appear not to affect the rate, since its concentration stays virtually constant.
