Chapter 12.3, Problem 12.4P

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

${\text{H}}_2{\text{O}}_2\text{(}aq\text{)}+3{\text{I}}^{-}\text{(}aq\text{)}+2{\text{H}}^{+}(aq)\stackrel{}{\to }{\text{I}}_{\text{3}}{}^{-}\text{(}aq\text{)}+2{\text{H}}_2\text{O(}l\text{)}$

The rate of formation of the red triiodide ion, Δ[I₃₊]/Δt, can be determined by measuring the rate of appearance of the color (Figure 12.5).

Figure 12.5

A sequence of photographs showing the progress of the reaction of hydrogen peroxide (H₂O₂) and iodide ion (I⁻). As time passes (left to right), the red color due to the triiodide ion (I₃⁻) increases in intensity.

Initial rate data at 25 °C and a constant H⁺ ion concentration are as follows:

1. (a) What is the rate law for the formation of I₃⁻?
2. (b) What is the value of the rate constant?
3. (c) What is the initial rate of formation of I₃⁻ when the initial concentrations are [H₂O₂] = 0.300 M and [I⁻] = 0.400 M?