$$ \ce{X(g) + 2Y(g) -> XY2(g)} $$

In order to determine the order of the reaction represented above, the initial rate of formation of \(\ce{XY2}\) is measured using different initial values of \([\ce{X}]\) and \([\ce{Y}]\). The results of the experiment are shown in the table below.

$$ \begin{array}{|c|c|c|c|} \hline \text{Trial} & [\ce{X}] & [\ce{Y}] & \text{Initial Rate of Formation of }\ce{XY2} \\ & & & (M \pu{ s^-1}) \\ \hline 1 & 0.50 & 0.50 & 8.0\times 10^{-3} \\ \hline 2 & 1.00 & 0.50 & 3.2 \times 10^{-2} \\ \hline 3 & 1.00 & 1.00 & 6.4\times 10^{-2} \\ \hline \end{array} $$

In trial 2 which of the reactants would be consumed more rapidly, and why?

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Using the method of initial rates, we could confirm that the reaction is second order with respect to X and first order with respect to Y.

Using that method is timely and unncessary. We can simply look at the stoichiometric ratios of the balanced reaction. 2 molecules of Y are consumed for every molecule of X. Therefore, Y will be consumed more rapidly.