MCAT General Chemistry Review - Alexander Stone Macnow, MD 2019-2020

Chemical Kinetics
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.

o Intermediates are molecules that exist within the course of a reaction but are neither reactants nor products overall.

o 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.

o For a collision to be effective, molecules must be in the proper orientation and have sufficient kinetic energy to exceed the activation energy.

o 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.

o From the transition state, the reaction can proceed toward products or revert back to reactants.

o The transition state is the highest point on a free energy reaction diagram.

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

o Increasing the concentration of reactant will increase reaction rate (except for zero-order reactions) because there are more effective collisions per time.

o Increasing the temperature will increase reaction rate because the particles’ kinetic energy is increased.

o Changing the medium can increase or decrease reaction rate, depending on how the reactants interact with the medium.

o 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.

o The rate orders usually do not match the stoichiometric coefficients.

o 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.

o The rate of a zero-order reaction can only be affected by changing the temperature or adding a catalyst.

o 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.

o A concentration vs. time curve of a first-order reaction is nonlinear.

o 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.

o A concentration vs. time curve of a second-order reaction is nonlinear.

o 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.