When ozone is heated in a closed vessel, it slowly turns back into oxygen: $2\,\text{O}_3(g) \rightarrow 3\,\text{O}_2(g)$. A proposed mechanism has a fast equilibrium followed by a slow step: $\displaystyle \text{O}_3 \rightleftharpoons \text{O}_2 + \text{O} \quad (\text{fast, equilibrium constant } K)$ $\displaystyle \text{O} + \text{O}_3 \rightarrow 2\,\text{O}_2 \quad (\text{slow, rate constant } k_2)$ Oxygen atoms form and react again so fast in the first step that it stays at equilibrium, with $K = \dfrac{[\text{O}_2][\text{O}]}{[\text{O}_3]}$. Only a tiny share of the oxygen is ever present as free atoms. The rate of the overall reaction is the rate of the slow step, and the temperature stays fixed throughout. Which of the following statements are correct?