Final answer:
The rate law for the reaction is rate = k[NO]². The units of the rate constant are M⁻¹s⁻¹. The average value of the rate constant can be calculated by averaging the rate constant values obtained from the three data sets. To find the rate of disappearance of NO and O2, the given concentrations can be substituted into the rate law equation and the stoichiometry of the reaction can be used.
Step-by-step explanation:
(a) To determine the rate law for the reaction, we need to examine the effect of the initial concentrations of reactants on the rate. Comparing runs 1 and 2, we see that doubling the initial concentration of NO (from 0.0126 M to 0.0252 M) results in a 4-fold increase in the initial rate. Therefore, the rate is directly proportional to the square of the initial concentration of NO, indicating that the rate law is rate = k[NO]².
(b) To find the units of the rate constant, we analyze the units of the rate equation. Since the rate is expressed in M/s and the concentration of NO is squared, the rate constant, k, must have units of M⁻¹s⁻¹.
(c) The average value of the rate constant can be calculated by taking the average of the rate constant values obtained from the three data sets.
(d) To find the rate of disappearance of NO when [NO] = 0.0750 M and [O2] = 0.0100 M, we substitute these values into the rate law and solve for the rate.
(e) To find the rate of disappearance of O2 at the given concentrations, we can use the stoichiometry of the reaction. Since the coefficient of O2 is 1, the rate of disappearance of O2 is equal to the rate of formation of NO2.