13.1 - The Equilibrium Condition

• Equilibrium is a dynamic situation

• The concept of chemical equilibrium is analogous to the flow of cars across a bridge connecting two island cities

• Although the equilibrium position lies far to the right, the concentrations of reactants never go to zero

• A double arrow is used to show that a reaction can occur in either direction

• The equilibrium position of a reaction is determined by many factors: the initial concentrations, the relative energies of the reactants and products, and the relative degree of “organization” of the reactants and products

• Energy and organization come into play because nature tries to achieve minimum energy and maximum disorder

• The United States produces about 20 million tons of ammonia annually

• A pure liquid or solid is never included in the equilibrium expression

• There are two possible reasons why the concentrations of the reactants and products of a given chemical reaction remain unchanged when mixed

  • The system is at chemical equilibrium

  • The forward and reverse reactions are so slow that the system moves toward equilibrium at a rate that cannot be detected.

• Dynamic state: Reactants and products are interconverted continually

13.2 - The Equilibrium Constant

• The law of mass action is based on experimental observation

• The law of mass action applies to solution and gaseous equilibria

• The square brackets indicate concentration in units of mol/L

• When the balanced equation for a reaction is multiplied by a factor n, the equilibrium expression for the new reaction is the original expression raised to the ninth power although the special ratio of products to reactants defined by the equilibrium expression is constant for a given reaction system at a given temperature, the equilibrium concentrations will not always be the same

• For a reaction at a given temperature, there are many equilibrium positions but only one value for K

• Each set of equilibrium concentrations is called an equilibrium position

13.3 - Equilibrium Expressions Involving Pressures

• K involves concentrations; Kp involves pressures. In some books, the symbol Kc is used instead of K

• The symbol Kp represents an equilibrium constant in terms of partial pressures

• In the equilibrium expression for the ammonia synthesis reaction, the sum of the powers in the numerator is different from that in the denominator, and K does not equal Kp

13.4 - Heterogeneous Equilibria

• Many equilibria involve more than one phase and are called heterogeneous equilibria

• Lime is among the top five chemicals manufactured in the United States in terms of the amount produced.

• The concentrations of pure liquids and solids are constant

• Experimental results show that the position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present

  • The fundamental reason for this behavior is that the concentrations of pure solids and liquids cannot change

• If pure solids or pure liquids are involved in a chemical reaction, their concentrations are not included in the equilibrium expression for the reaction

  • This simplification occurs only with pure solids or liquids, not with solutions or gases

13.5 - Applications of the Equilibrium Constant

• The inherent tendency for a reaction to occur is indicated by the magnitude of the equilibrium constant

• A value of K much larger than 1 means that at equilibrium the reaction system will consist of mostly products—the equilibrium lies to the right

• A very small value of K means that the system at equilibrium will consist of mostly reactants—the equilibrium position is far to the left

• It is important to understand that the size of K and the time required to reach equilibrium are not directly related

  • The time required to achieve equilibrium depends on the reaction rate, which is determined by the size of the activation energy

  • The size of K is determined by thermodynamic factors such as the difference in energy between products and reactants

• The reaction quotient is obtained by applying the law of mass action using initial concentrations instead of equilibrium concentrations

• Sometimes we are not given any of the equilibrium concentrations (or pressures), only the initial values

  • Then we must use the stoichiometry of the reaction to express concentrations at equilibrium in terms of the initial values

• Since the coefficients in the balanced equation are all 1, the magnitude of the change is the same for all species

13.6 - Solving Equilibrium Problems

• Procedure for Solving Equilibrium Problems: Write the balanced equation for the reaction

  • Write the equilibrium expression using the law of mass action

  • List the initial concentrations

  • Calculate Q, and determine the direction of the shift to equilibrium

  • Define the change needed to reach equilibrium, and define the equilibrium concentrations by applying the change to the initial concentrations

  • Substitute the equilibrium concentrations into the equilibrium expression, and solve for the unknown

• Check your calculated equilibrium concentrations by making sure they give the correct value of K

• It is equally valid to use pressures for a gas phase system at constant temperature and volume

• Approximations can simplify complicated math, but their validity should be checked carefully

13.7 - Le Chatelier's Principle

• We can qualitatively predict the effects of changes in concentration, pressure, and temperature on a system at equilibrium by using Le Châtelier’s principle,

  • It states that if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change

• Since the change imposed is the addition of nitrogen, Le Châtelier’s principle predicts that the system will shift in a direction that consumes nitrogen

• This reduces the effect of the addition

• The system shifts in the direction that compensates for the imposed change

• Another way of stating Le Châtelier’s principle is to say that if a component is added to a reaction system at equilibrium, the equilibrium position will shift in the direction that lowers the concentration of that component

  • When an inert gas is added, there is no effect on the equilibrium position.

• The addition of an inert gas increases the total pressure but has no effect on the concentrations or partial pressures of the reactants or products

• The central idea is that when the volume of the container holding a gaseous system is reduced, the system responds by reducing its own volume

  • This is done by decreasing the total number of gaseous molecules in the system

• If energy is added to this system at equilibrium by heating it, Le Châtelier’s principle predicts that the shift will be in the direction that consumes energy, that is, to the left

• Although Le Châtelier’s principle cannot predict the size of the change in K, it does correctly predict the direction of the change.