Concept Summary

Chemical Kinetics

• The change in Gibbs free energy (ΔG) determines whether or not a reaction is spontaneous.

• Chemical mechanisms propose a series of steps that make up the overall reaction.

  • Intermediates are molecules that exist within the course of a reaction but are neither reactants nor products overall.
  • The slowest step, also known as the rate-determining step, limits the maximum rate at which the reaction can proceed.

• The collision theory states that a reaction rate is proportional to the number of effective collisions between the reacting molecules.

  • For a collision to be effective, molecules must be in the proper orientation and have sufficient kinetic energy to exceed the activation energy.
  • The Arrhenius equation is a mathematical way of representing collision theory.

• The transition state theory states that molecules form a transition state or activated complex during a reaction in which the old bonds are partially dissociated and the new bonds are partially formed.

  • From the transition state, the reaction can proceed toward products or revert back to reactants.
  • The transition state is the highest point on a free energy reaction diagram.

• Reaction rates can be affected by a number of factors.

  • Increasing the concentration of reactant will increase reaction rate (except for zero-order reactions) because there are more effective collisions per time.
  • Increasing the temperature will increase reaction rate because the particles’ kinetic energy is increased.
  • Changing the medium can increase or decrease reaction rate, depending on how the reactants interact with the medium.
  • Adding a catalyst increases reaction rate because it lowers the activation energy. Homogeneous catalysts are the same phase as the reactants; heterogeneous catalysts are a different phase.

Reaction Rates

• Reaction rates are measured in terms of the rate of disappearance of a reactant or appearance of a product.

• Rate laws take the form of rate = k[A]^x[B]^y.

  • The rate orders usually do not match the stoichiometric coefficients.
  • Rate laws must be determined from experimental data.
• The rate order of a reaction is the sum of all individual rate orders in the rate law.

• Zero-order reactions have a constant rate that does not depend on the concentration of reactant.

  • The rate of a zero-order reaction can only be affected by changing the temperature or adding a catalyst.
  • A concentration vs. time curve of a zero-order reaction is a straight line; the slope of such a line is equal to –k.

• First-order reactions have a nonconstant rate that depends on the concentration of reactant.

  • A concentration vs. time curve of a first-order reaction is nonlinear.
  • The slope of a ln [A] vs. time plot is –k for a first-order reaction.

• Second-order reactions have a nonconstant rate that depends on the concentration of reactant.

  • A concentration vs. time curve of a second-order reaction is nonlinear.
  • The slope of a vs. time plot is k for a second-order reaction.
• Broken-order reactions are those with noninteger orders.
• Mixed-order reactions are those that have a rate order that changes over time.