Question

The oxidation of iodide ion by hydrogen peroxide in an acidic solution is described by the balanced equation

H2O2(aq)+3I−(aq)+2H+(aq)⟶I3−(aq)+2H2O(l)H2O2(aq)+3I−(aq)+2H+(aq)⟶I3−(aq)+2H2O(l)

The rate of formation of the red triiodide ion, Δ[I3−]/ΔtΔ[I3−]/Δt, can be determined by measuring the rate of appearance of the color.

A sequence of photographs showing the progress of the reaction of hydrogen peroxide (H2O2)(H2O2) and iodide ion (I−)(I−).

As time passes (left to right), the red color due to the triiodide ion (I3−)(I3−) increases in intensity.

Initial rate data at 25 ∘C∘C and a constant [H+][H+] are as follows:

Experiment	Initial [H2O2][H2O2]	Initial [I−][I−]	Initial Reaction Rate (M/s)(M/s)
1	0.100	0.100	1.15×10−41.15×10−4
2	0.100	0.200	2.30×10−42.30×10−4
3	0.200	0.100	2.30×10−42.30×10−4
4	0.200	0.200	4.60×10−44.60×10−4

What is the value of the rate constant?

Answer 1

EXP.	[H2O2]	[I-]	Rate of Reaction (MS-1)
1	0.100	0.100	1.15×10−4
2	0.100	0.200	2.30×10−4
3	0.200	0.100	2.30×10−4
4	0.200	0.200	4.60×10−4

Rate = k [H2O2]a [I-]b

From 1 & 2 ; 1.15×10−4 = k [0.1]a [0.1]b ; 2.30×10−4 = k [0.1]a [0.2]b ; b = 1

From 1 & 3 ; 1.15×10−4 = k [0.1]a [0.1]b ; 2.30×10−4 = k [0.2]a [0.1]b ; a = 1

Rate = k [H2O2]1 [I-]1

From 1:  

1.15×10−4 = k [0.1]1 [0.1]1

k = 1.15×10−2