Kinetics
Chapter 12
Kinetics helps us to answer questions such as: how much time will it take for this reaction to happen? After 20 minutes, how much product will form? When will half of the initial reactant be used up? If we double the amount of one reactant, how will this impact the rate of the reaction? If we decrease the temperature by a few degrees, how much will this impact the rate of the reaction? What are the individual collisions that occur to cause products to form in this reaction? How does adding a catalyst change this reaction?
Important Equations
Above is a generic differential rate law for a reaction involving two reactants: A and B. k is the rate constant for the reaction, [A] is the initial molarity of reactant A, x is the partial order of reactant A, [B] is the initial molarity of reactant B, and y is the partial order of reactant B.
Above is a zero order integrated rate law. This equation should only be used if the partial order of reactant A is zero. k is the rate constant for the reaction, t is time, [A] is the molarity of reactant A remaining at time t, and [A]0 is the initial molarity of reactant A.
Above is a first order integrated rate law. This equation should only be used if the partial order of reactant A is one. k is the rate constant for the reaction, t is time, [A] is the molarity of reactant A remaining at time t, and [A]0 is the initial molarity of reactant A.
Above is a second order integrated rate law. This equation should only be used if the partial order of reactant A is two. k is the rate constant for the reaction, t is time, [A] is the molarity of reactant A remaining at time t, and [A]0 is the initial molarity of reactant A.
Above is the Arrhenius equation which can be used to compare two rate constants for the same reaction occur at two temperatures. Ea is the activation energy of the reaction, R is the gas constant, k1 is the rate constant when the reaction occurs at temperature T1, and k2 is the rate constant when the reaction occurs at temperature T2.