Chapter 14, Problem 14.68SP

The oxidation of iodide ion by hydrogen peroxide in an acidic solution is described by the balanced equation $H_2{O}_2\left(aq\right)+3I^{-1}\left(aq\right)+2H^{+}\left(aq\right)\to I_3^{-}\left(aq\right)+2H_{2}O\left(l\right)$
The rate of formation of the red triiodide ion, $\Delta \left[I_3^{-}\right]/\Delta t,$ can be determined by measuring the rate of appearance of the color.

A sequence of photographs showing the progress of thereaction of hydrogen peroxide (H₂O₂) and iodide ion ( ${\text{I}}^{-}$ ). As time passes, the red color due to the triiodide ion ( $I_3^{-}$ ) increases in intensity.

Initial rate data at 25 °C and a constant ${\text{[H}}^{\text{+}}\text{]}$ are as follows:

(a) What is the rate law?
(b) What is the value of the rate constant?
(c) What is the initial rate when the initial concentrations are ${\text{[H}}_2{\text{O}}_2\text{] = 0}{\text{.300 M and [I}}^{-}\text{] = 0}\text{.400 M?}$