Kinetics is the study of the rate of reaction and the factors that affect it.
In order for a reaction to occur, the molecules must collide in the correct orientation with the minimum energy needed for the transition from the reactants to the products (the activation energy). Only a very few collisions meet these requirements and result in a reaction.
Reaction Coordinate Diagrams
Reaction coordinate diagrams show the energy of the reactants, the activation energy up to the activated complex, or transition state (the in-between state between the reactants and the products), and the energy of the products. The overall energy change of the reaction is also shown.
Factors affecting rate
Increasing the temperature increases the number of collisions, and also the number of collisions with the needed energy. Therefore, increasing temperature increases the rate of reaction. Increasing the concentration or the surface area also increases the number of collisions, therefore increasing the chance that a successful collision will occur—which increases rate. Adding a catalyst, a species that increases the rate of reaction without being used up in the reaction, also increases the rate.
Dynamic equilibrium is when the rate of the forward and reverse reactions are equal. The reaction will appear to have stopped, as the concentrations of each species won’t change, but the reaction continues to proceed in both directions.
Equilibrium constant expressions are the ratio of the concentrations of products to reactants using their coefficients from the balanced equation as exponents. Pure solids and liquids are not included in the equilibrium constant expression. Equilibrium concentrations are plugged into the expression to solve for the equilibrium constant. A large equilibrium constant means that the reaction “lies to the right” with more products than reactants at equilibrium. A small constant means that the reaction “lies to the left” with more reactants the products present at equilibrium.
The reaction quotient is the value when concentrations at any time are plugged into the equilibrium constant expression. If the reaction quotient equals the equilibrium constant, then the system is at equilibrium. If the reaction quotient is too large, the reaction will proceed to the left to produce more reactants (and reduce products). The opposite will happen if the reaction quotient is too small.
Le Chatelier’s Principle
Le Chatelier’s principle says that if a system at equilibrium is stressed or changed, the system will shift to reach equilibrium again. Whatever you do to the system, it will un-do it. For example, if you add more reactants, the reaction will proceed to the right to get rid of the extra reactants and make more products.