Cliffs AP Chemistry, 3rd Edition

Part III: AP Chemistry Laboratory Experiments

A Sample Lab Report*

I. Title: Chemical Kinetics: An Iodine Clock Reaction

II. Date:2/3/98

III. Purpose:Determine the initial rate of a chemical reaction using an internal indicator.

Examine the dependence of the initial reaction rate upon the initial concentrations of the reactants.

Find the reaction order, the partial orders, and the experimental rate constant of the reaction.

IV. Theory:The rate of the reaction:

32 3IHO HO I HO++ 232 " 3 2

-+-+ (1)

may be determined by measuring the time required for a fixed amount of S 2 O 32 - to

react with the I 3 – formed in (1)

2 SO 232 --++I 346 "SO^2 - 3I- (2)

The moment the S 2 O 22 −is all used up, excess I 3 – is free to react with starch to form

a dark-blue complex. The coefficients in the rate equation.

Rate = k[H 2 O 2 ]x[H 3 O+]z[I–]y (3)

May be determined by the initial rate method, where the

4

rate t sec

I

t

SO M

∆

∆

2 ∆

==^3315102 =#

77 --AA^2 -

V. Procedure:Chemical Education Resources, Inc., “Chemical Kinetics: An Iodine Clock
Reaction,” 1988.**

VI. Results:

Part I Standardization of Hydrogen Peroxide

Average concentration of H 2 O 2 0.10 M

Part II. Reaction Order with Respect to Hydrogen Peroxide

Beaker A Beaker B

0.05 M KI [I–] 0.05 M Na 2 S 2 O 3 Buffer1 0.1 M H 2 O 2 [H 2 O 2 ] Time Initial rate

(mL) (M) (mL) (mL) (mL) (M) (s) (M/s)

10.0 0.010 1.0 19.0 20.0 0.040 47 1.0638 × 10 –5

10.0 0.010 1.0 24.0 15.0 0.030 51 9.8039 × 10 –6

10.0 0.010 1.0 28.0 11.0 0.022 60 8.3333 × 10 –6

10.0 0.010 1.0 30.00 9.00 0.018 70 7.1429 × 10 –6

10.0 0.010 1.0 32.0 7.00 0.014 90 5.5556 × 10 –6