Question #44840

A kinetic study of the oxidation of a large excess of an organic substrate by Cr(VI) is performed in a heavily buffered aqueous solution at a constant pH of 5.0 at 298 K. The total analytical concentration of Cr(VI), denoted as $C = [\ce{Cr(VI)}]_{T}$, is monitored over time. The active oxidant is strictly the dichromate ion ($\ce{Cr2O7^{2-}}$), and the intrinsic redox reaction follows pseudo-first-order kinetics: Rate of redox consumption $= k_{redox}[\ce{Cr2O7^{2-}}]$. The solution maintains the rapid equilibrium: $2\ce{CrO4^{2-}} + 2\ce{H+} \rightleftharpoons \ce{Cr2O7^{2-}} + \ce{H2O}$. Assuming $\ce{CrO4^{2-}}$ and $\ce{Cr2O7^{2-}}$ are the only significant Cr(VI) species present, which of the following statements correctly describes the apparent kinetic order with respect to $C$ and the corresponding graphical behavior as the reaction progresses from $t=0$ to near completion?