Chemical Equilibrium AP Chapter 15
Chemical Equilibrium • Chemical Equilibrium occurs when opposing reactions are proceeding at equal rates. • It results in the formation of an equilibrium mixture of the reactants and products of the reaction. • The composition mixture does not change with time.
The Haber Process • N2(g) + 3H2(g) ↔ 2NH3(g) • This reaction involves the presence of a catalyst, a pressure of several hundred atmospheres and a temperature of several hundred degrees Celsius. • This equilibrium mixture can be reached regardless of where one starts.
Law of Mass Action • The relationship between the concentrations of the reactants and the products of a system in equilibrium is given by the law of mass action. • aA + bB ↔ dD + eE • Kc = Kc = equilibrium constant [D]d[E]e ← products [A]a[B]b ← reactants
Example • N2O4(g) ↔ 2NO2(g) • Kc = = = 0.211 [NO2]2 [N2O4] [0.0172]2 [0.00140]
Equilibrium Constants - Pressure • When the equilibrium system consists of gases, it is convenient to express the concentrations of reactants and products in terms of gas pressures: • Kp = • Kc and Kp are related by: Kp = Kc(RT)Δn (PD)d(PE)e (PA)a(PB)b
Working with Equilibrium Constants • CO(g) + Cl2(g) ↔ COCl2(g) • Kc = = 4.56 x 109 • For the equilibrium constant to be so large, the numerator must be much larger than the denominator. Therefore, the equilibrium concentration of COCl2 must be greater than that of CO or Cl2. [COCl2] [CO][Cl2]
Relating Chemical Equations and Equilibrium Constants • A large value for the equilibrium constant indicates that the mixture contains more products than reactants and therefore lies towards the product side of the equation. • A small value for the equilibrium constant means the mixture contains less products than reactants and therefore lies toward the reactant side.
Homogeneous vs. Heterogeneous • Equilibria for which all substances are in the same phase are called homogeneous equilibria. • Equilibria in which 2 or more phases are present are called heterogeneous equilibria. • The concentrations of pure solids and liquids are left out of the equilibrium constant expression for a heterogeneous equation.
The equilibrium pressure of CO2 (g) is the same in both bell jars, at the same temperature. The equilibrium constant expression is Kp = PCO2.
Calculating Equilibrium Constants • If the concentrations of all species in an equilibrium are known, the equilibrium-constant expression can be used to calculate the value of the equilibrium constant. • The changes in the concentrations of reactants and products in the process of achieving equilibrium are governed by the stoichiometry of the equation.
Applying Equilibrium Constants • The reaction quotient, Q, is found by substituting reactant and product concentrations or partial pressures at any point during a reaction into the equilibrium-constant expression. • If the system is at equilibrium, Q = K. • Q ≠ K, the system is not at equilibrium. • Q < K, the reaction will move toward equilibrium by forming more products (it moves from left to right.) • Q > K, the reaction will proceed from right to left.
Reaction Quotients, continued • Knowing the value of K makes it possible to calculate the equilibrium amounts of the reactants and products, often by solving an equation where the unknown is the change in a partial pressure or concentration.
LeChatlier’s Principle • LeChatlier’s principle states: • If a system at equilibrium is disturbed by a change in temperature, pressure or concentration of one of the components, the system will shift its equilibrium position to counteract the effect of the disturbance.
Change in reactant or product • If a chemical system is at equilibrium and the concentration of a substance is increased (either product or reactant), the system responds by consuming some of the substance. • If some of the concentration is decreased, the system will respond by producing some of the substance. • N2 + 3H2↔ 2NH3
Change in Volume and Pressure • At constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas. • The system will always favor the side with fewer moles of a gas in order to re-equilibrate.
Temperature Changes • When the temperature of a system in equilibrium is increased, the system reacts as though a reactant were added to an endothermic reaction or a product was added to an exothermic reaction. • It will shift in the direction to consume the excess reactant or product, which is heat. • Endothermic: reactants + heat↔ products • Exothermic: reactants ↔ products + heat
Catalysts • Adding a catalyst increases the rate at which equilibrium is achieved, but it does not change the composition of the equilibrium mixture.