Given the following experimental values, for the reaction A + B --> C, figure out the rate constant (k).
Initial values: [A] = 0.0100M; [B] = 0.0100M; Rate = 5.00 * 10^{-6} M/s
Experiment 1: [A] = .0200M; [B] = 0.100M; Rate = 2.50 * 10^{-5} M/s
Experiment 2: [A]= 0.0100M; [B]=0.0200M; Rate = 5.00 * 10^{-6} M/s
Explanation
Answer: B. Solution is as follows:
When [A] doubles, the rate is 4 times as large (2.50 * 10^{-5}/5.00 *10^{-6} = 4). Thus, we have [2A] = 4r. Therefore, we have a second order reaction with respect to A.
When B doubles, the reaction rate doesn’t change, meaning we have a zero order with respect to B.
Thus r=k[A]^{2}
Inputting the initial values:
5.0 *10^{-6} M/s = k[0.0100M]^{2}
solving for k yields 5.00 *10^{-2} with units of 1/(M*s)
Kinetic control vs. thermodynamic control