unit 5: kinetics

a species which appears in the mechanism of a reaction, but not in the overall balanced equation. intermediates are not in overall related rates equations!! even if they are in the slow step, they are not included!! if they are not part of the first reactants OR final products, do NOT include them!

describes how fast the reaction occurs. measured in change in reactant or change in product over time.

1. concentration of reactants: included in the rate law equation as a variable

   1. more reactants = more collisions = higher reaction rate

2. temperature: factored into k (rate constant)

   1. greater temperature = more collisions = greater reaction rate

   2. temperature is a measure of the kinetic energy in a container/system

3. physical state matters: factored into unique rate constant (k)

   1. solid reactants = less surface area = fewer collisions = slower reaction rate

4. presence of a catalyst: increases rate of reaction by lowering the energy required for a successful reaction and collision

• an equation that relates the reactants to the reaction rate

  • must be determined experimentally

if the amount of [A] reactant doubles, and the rate stays the same. units are M/s

if the amount of [B] reactant doubles, and so does the rate. units are 1/s

if the amount of [C] reactant doubles, and the rate quadruples. units are 1/Ms

[A](t) - [A](0) = -kt

ln[A](t) - ln[A](0) = -kt

1/[A](t) - 1/[A](0) = kt

a reaction with a single reaction step that has a single intermediate/transition state. rate law is determined by the reactant concentrations and coefficients in a balanced chemical equation

a reaction that requires 2+ elementary steps. rate law is not discernible from coefficients in the balanced chemical equation

a sequence of elementary reactions (steps) that, when summed together, result in the overall net reaction

predicts reaction rate based on frequency of effective collisions (collisions resulting in a chemical reaction are effective). depends on:

1. frequency of collisions: more collisions = higher reaction rate

2. kinetic energy of colliding particles: higher temperature/kinetic energy = higher reaction rate (since chemical bonds require energy to break)

3. orientation of colliding particles: correctly orientated particles = higher chance of correct collision = higher reaction rate

the slowest elementary reaction with the highest activation energy (determines rate and rate law). has the greatest activation energy

the second step of the reaction in the slow/rate limiting step

increases the reaction rate by creating an alternative reaction pathway/environment that reduces the activation energy required for the reaction