Summary
Reaction Rates and Chemical Equilibrium

For two molecules to react, they must collide and the electrons rearrange so that old bonds are broken and new bonds formed. For a reaction to occur spontaneously, the enthalpy of the reaction, the temperature at which it is to occur, and the entropy change must be such that the free energy of the reaction as calculated by the equation G = H - TS is less than zero. Often activation energy must be added to initiate the reaction. The course of a chemical reaction can be plotted versus the energy of the reactants and products. This plot shows the activation energy necessary for an effective collision between reactants, as well as the exothermic or endothermic nature of the reaction. The number of collisions between reactants can be increased by increasing their concentrations. Increasing the temperature of a reaction will increase the number of molecules that are sufficiently energetic to collide effectively. The addition of a catalyst provides a reaction pathway that requires a lower activation energy. Catalysts in biological systems (enzymes) are particularly important as they allow complex reactions to occur at body temperature.

Many chemical reactions are reversible. Allowed to proceed spontaneously, they reach an equilibrium state in which the rate of the forward reaction equals the rate of the reverse reaction. At equilibrium, a relationship exists between the concentrations of the components of the reaction known as the equilibrium constant. For the reaction

aA + bB cC + dD

the equilibrium constant expression is

Le Chatelier's Principle predicts how an equilibrium will adjust to stresses imposed by changing concentrations, pressure, or temperature. A catalyst does not shift an equilbrium.

Equilibrium constants
Name of Constant	Symbol	Typical Equilibrium	Expression of constant
equilibrium constant	Keq	A2 + B2 2 AB	[AB]2    [A2[B2]
acid dissociation constant	Ka	HA H+ + A-	[H+][A-]        [HA}
ionization constant of water	Kw	H2O(l) H+ + OH-	[H+][OH-]
solubility product constant	Ksp	MaNb(s) aMb+ + bNa-	[Mb+][Na-]b

Ionic equilibria may involve sparingly soluble salts. The equilibrium constant for the equilibrium between a sparingly soluble salt and its ions in solution is known as the solubility constant, K_sp. The concentration of the undissolved salt does not appear in this constant. Ionic equilibria also involve weak electrolytes. These equilbibria were discussed in Chapter 12 and repreated here in a discussion of buffers.

A buffer solution resists change. an acid-base buffer resists changes in pH. Buffers, prepared from weak acids or bases and their salts, are particularly effective within one pH unit of the pK_a of the acid. The table lists thetypes of equilibrium constants discussed.