A kinetic study of the reaction: A → products provides the data: t = 0 s, [A] = 2.00 M; 500 s, 1.00 M; 1500 s, 0.50 M; 3500 s, 0.25 M. In the simplest possible way determine, whether this reaction is of

second order

Given data:
\(t = 0\) s, \([A] = 2.00\) M;
\(t = 500\) s, \([A] = 1.00\) M;
\(t = 1500\) s, \([A] = 0.50\) M;
\(t = 3500\) s, \([A] = 0.25\) M.

To determine the order of the reaction, we need to examine the change in concentration with respect to time. Let's calculate the reaction rate for each data point:

Rate at \(t = 0\) s: \(\frac{{[A]_t - [A]_0}}{{t - t_0}} = \frac{{1.00 - 2.00}}{{500 - 0}} = -0.002\) M/s
Rate at \(t = 500\) s: \(\frac{{[A]_t - [A]_0}}{{t - t_0}} = \frac{{0.50 - 1.00}}{{1500 - 500}} = -0.002\) M/s
Rate at \(t = 1500\) s: \(\frac{{[A]_t - [A]_0}}{{t - t_0}} = \frac{{0.25 - 0.50}}{{3500 - 1500}} = -0.002\) M/s

From the data, we can observe that the reaction rate remains constant \((-0.002\) M/s\) throughout the entire process. This suggests that the reaction follows second-order kinetics.

Therefore, the correct answer is (3) second order.