Consider the following mechanism for the reaction between nitric oxide and ozone: 03(g) + NO(g) O(g) + NO3(g) O(g) + O2(g) → 20₂(g) NO3(g) + NO(g) - - 2 NO₂(g) Write the chemical equation of the overall reaction: Are there any intermediates in this mechanism? If there are intermediates, write down their chemical formulas. Put a comma between each chemical formula, if there's more than one. 0 yes no (1) (2) (3)

13.102 Chlorine oxide (CIO), which plays an important role in the depletion of ozone (see Problem 13.101), decays rapidly at room temperature according to the equation 2C10(g) Cl₂(g) + O₂(g) From the following data, determine the reaction order and calculate the rate constant of the reaction. Time (s) 0.12 x 10-³ 0.96 x 10-³ 2.24 x 10-³ 3.20 x 10-³ 4.00 x 10-³ [CIO] (M) 8.49 x 10-6 7.10 x 10-6 5.79 x 10-6 5.20 x 10-6 4.77 x 10-6

A. Consider the following proposed mechanism for the production of chlorine dixoide from chlorine and ozone. Draw a reaction coordinate diagram for this reaction. Label reactants, products, transition states, intermediates, and activation energy on your sketch. Assume the reaction is endothermic. NOTE: CI is chlorine, not Carbon lodide Overall: Cl₂ (g) + 203 (g) → 2 CIO₂ (g) + O₂ (g) Step 1: Cl₂ (g) → 2 Cl (g) fast Step 2: Cl (g) + Os (g) CIO₂ +O fast Step 3: O (g) + O3 (g) → 202 (g) fast Step 4: Cl (g) + O2 (g) → CIO2 (g) slow

Given the rate of change of the concentration of C, determine the rate of change of the concentration of D.3 A (g) + 7 B (g) → 6 C (g) + 5 D (g)         Δ[C]/Δt = 5.458x10 -6 M/s

Consider the mechanism. Step 1: Step 2: А +В — С slow А +С — D fast Overall: 2A + В — D Determine the rate law for the overall reaction, where the overall rate constant is represented as k. rate =

Given the rate of change of the concentration of C, determine the rate of change of the concentration of A. 3 A (g) + 5 B (g) → 2 C (g) + 6 D (g)         Δ[C]/Δt = 7.984x10 -5 M/s

Suppose the reaction between nitric oxide and oxygen proceeds by the following mechanism: step elementary reaction rate constant 1 NO(g) + O,(g) –→ NO2(g) + O(g) k1 NO(g) + O(g) → NO2(g) Suppose also k,«k,. That is, the first step is much slower than the second. Write the balanced chemical equation for the overall chemical reaction: Write the experimentally- observable rate law for the overall chemical reaction. rate = k || Note: your answer should not contain the concentrations of any intermediates. Express the rate constant k for the overall chemical reaction in terms of k1, k2, and (if necessary) the rate constants k.1 and k-2 for k = | %3D the reverse of the two elementary reactions in the mechanism.

The mechanism for the reaction described by NO₂(g) + CO(g) → CO₂(g) + NO(g) is suggested to be k₁ (1) 2 NO₂(g) NO₂(g) + NO(g) k₂ (2) NO3(g) + CO(g) → NO₂(g) + CO₂(g) Assuming that [NO] is governed by steady-state conditions, derive the rate law for the production of CO₂ (g). rate of CO₂ (g) production =

Consider the following mechanism for the formation of nitrogen dioxide: 2 NO(g) N2O2(g) (1) N,O2(9) + O2(9) → 2 NO2(g) (2) Write the chemical equation of the overall reaction: yes 0,0,.. Are there any intermediates in this mechanism? no ? If there are intermediates, write down their chemical formulas. Put a comma between each chemical formula, if there's more than one.