A proposed mechanism for the gas-phase reaction \(2\text{NO}(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\) consists of an initial fast equilibrium followed by a slow elementary step:

\[ \begin{aligned} \text{Step 1: } & 2\text{NO}(g) \rightleftharpoons \text{N}_2\text{O}_2(g) & & (\text{fast equilibrium, } \Delta H_1^\circ < 0) \\ \text{Step 2: } & \text{N}_2\text{O}_2(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g) & & (\text{slow, rate constant } k_2) \end{aligned} \]

Under constant initial concentrations of \(\text{NO}(g)\) and \(\text{O}_2(g)\), the initial rate of formation of \(\text{NO}_2(g)\) is observed to decrease as the temperature is increased. Which of the following best accounts for this observed outcome in terms of the elementary rate constant \(k_2\) and the equilibrium intermediate concentration \([\text{N}_2\text{O}_2]\)?