Question textConsider the decomposition reaction of hydrogen peroxide: 2 H2O2 (aq) → 2 H2O (l) + O2 (g) The reaction cannot occur in one step corresponding to the overall balanced equation. If it did, the rate law would demonstrate that reaction to be second order with respect to H2O2 since the coefficient of the H2O2 in the balanced equation is 2. A reaction mechanism can be constructed which accounts for the rate law and for the experimental detection of the IO− (hypoiodite) ion. This mechanism consists of two bimolecular elementary steps: Step 1  H2O2 (aq) + I− (aq) → H2O (l) + IO− (aq) k1 Step 2  H2O2 (aq) + IO− (aq) → H2O (l) + I− (aq) + O2 (g) k2 Each of the elementary steps describes what actually happens at the molecular level. How do we explain the observed rate dependence? Concentration of the iodide ion [I−] appears in the rate law. Therefore, the elementary step with I− (step 1) Answer 1 Question 5[select: , is, is not] a rate-determining step and Answer 2 Question 5[select: , k1 ≪ k2, k1 ≫ k2, k1 = k2] Therefore, the overall rate of decomposition is completely controlled by the Answer 3 Question 5[select: , first, second] step and Answer 4 Question 5[select: , k = k1, k = k2] And that explains why the overall rate depends on [H2O2] and [I−].Check Question 5Multiple fill-in-the-blank