A student investigates the decomposition of a gaseous reactant, \(\text{X(g)}\), catalyzed by a solid platinum surface, \(\text{Pt(s)}\), in a rigid container of fixed volume at constant temperature:
\[ \text{X(g)} \xrightarrow{\text{Pt(s)}} \text{Y(g)} + \text{Z(g)} \]
The accepted mechanism for the reaction consists of two elementary steps:
\[ \text{Step 1: } \text{X(g)} + \text{Pt(surface)} \rightleftharpoons \text{X}\cdot\text{Pt(surface)} \quad (\text{fast equilibrium}) \]
\[ \text{Step 2: } \text{X}\cdot\text{Pt(surface)} \rightarrow \text{Y(g)} + \text{Z(g)} + \text{Pt(surface)} \quad (\text{slow}) \]
At low initial partial pressures of \(\text{X(g)}\), the rate of decomposition is experimentally observed to be first order with respect to \(\text{X}\). Which of the following best predicts and explains the effect on the apparent order of the reaction with respect to \(\text{X}\) if the initial partial pressure of \(\text{X(g)}\) is increased to extremely high values while the total surface area of \(\text{Pt(s)}\) remains constant?