Potassium permanganate is used to react with antimony (III) oxide, Sb2O3, in a kinetics experiment.
MnO4- (aq) + 1.25 Sb2O3 (aq) + 3 H+ (aq) → Mn2+ (aq) + 1.25 Sb2O5 (aq) + 1.5 H2O (l)
The reaction is very fast, so the temperature was lowered to a constant value of 5°C for the following experiments. The time needed for the purple color of MnO4- to disappear was recorded.
Experiment # mLs 0.130 M KMnO4 mLs 0.450 M Sb2O3 mLs total volume Time (s) Rate
1 1.00 5.00 12.0 74
2 1.00 10.00 12.0 18
3 2.00 5.00 12.0 71
6. Calculate [MnO4-] under the initial conditions of experiment 1.
7. Calculate [Sb2O3] under the same conditions.
8. Calculate the average reaction rate, in M/s, for the disappearance of MnO4- during Experiment 1.
9. Determine the rate law, with proper exponent orders, using the results of the 3 reactions. Note that in Experiment 3, the change in [MnO4-] is twice as large as in the other experiments. You may leave the rate constant as the generic variable k for now.
10. Using any one of the experiments, calculate the value of the rate constant k.
11. Suppose that for a different group of students, in experiment 2, the temperature was allowed to rise to 8°C unknowingly, while kept at 5°C for the other experiments. What effect will this have on their time of reaction for Experiment 2 compared to the "true" time they should have measured? Explain your answer.
