one reactant varies while the concentration of the other reactant remains constant. The concentration of A molecules in vessel (2) is twice that in vessel (1) while the concentration of B remains constant. Because the reaction rate in vessel (2) is twice that in vessel (1), the rate is proportional to [A], and therefore the reaction is first order in A. When the concentration of B is doubled while the concentration of A remains constant [compare vessels (1) and (3)], the rate doubles, so the reaction is first order in B. When the concentrations of both A and B are doubled, the rate increases by a factor of 4 [compare vessels (1) and (4)], in accord with a reaction that is first order in A and first order in B. The overall reaction order is the sum of the orders in A and B, or 1 + 1 = 2. Since the reaction is first order in A and first order in B, the rate law is rate = k[A][B]. Note that the exponents in the rate law differ from the coefficients in the balanced chemical equation, A + 2 B rightarrow products. The oxidation of iodide ion by hydrogen peroxide in an acidic solution is described by the balanced equation H_2O_2(aq) + 3 I^-(aq) + 2 H^+ (aq) rightarrow I_3^-(aq) + 2 H_2O(l) The rate of formation of the red triiodide ion, A [I_3^-]/Delta t, can be determined by measuring the rate of appearance of the color (Figure 12.5). A sequence of photographs showing the progress of the reaction of hydrogen peroxide (H_2O_2) and iodide (I^-).