AP Chemistry Notes
Unit 8 Textbook Notes
12.1 The N2O4-NO2 Equilibrium System
When you put a sample of N2O4, a colorless gas, in a closed container at 100C a reddish-brown color starts to show. This is due to NO2 formed by the decomp. of part of the original substance. The forward and reverse reactions are taking place at the same rate. The concentrations of species present remain constant with time. These concentrations are independent of the direction from which equilibrium is approached. The equilibrium constant K is where the partial pressures in atmospheres, is a constant, independent of the original composition, the volume of the container, or the total pressure.
12.2 The Equilibrium Constant Expression
An equilibrium constant expression can be written for every gaseous chemical system and it states that the conditions that must be attained at equilibrium. Partial pressures must be expressed in atmospheres. The equilibrium partial pressures of products appear in the numerator. The equilibrium atrial pressures of reactants appear in the denominator. Each partial pressure is raised to a power equal to its coefficient in the balanced equation. This constant is usually symbolized by Kp to show that it involves partial pressures. Kc is used to symbolize that concentration is used. Kp=Kc(RT)change in n. The expression for K depends on the form of the chemical equation written to describe the equilibrium system. The coefficient rule states that if the coefficients in a balanced equation are multiplied by a factor n, the equilibrium constant is raised to the nth power (K’=Kn). The reciprocal rule states that the equilibrium constants for forward and reverse reactions are the reciprocals of each other (Kn=1/K). The rule of multiple equilibria states that if a reaction can be expressed as the sum of two or more reactions, K for the overall reaction is the product of the equilibrium constants of the individual reactions {K(reaction 3)=K(reaction 1)x K(reaction 2)}. The equilibrium
12.1 The N2O4-NO2 Equilibrium System
When you put a sample of N2O4, a colorless gas, in a closed container at 100C a reddish-brown color starts to show. This is due to NO2 formed by the decomp. of part of the original substance. The forward and reverse reactions are taking place at the same rate. The concentrations of species present remain constant with time. These concentrations are independent of the direction from which equilibrium is approached. The equilibrium constant K is where the partial pressures in atmospheres, is a constant, independent of the original composition, the volume of the container, or the total pressure.
12.2 The Equilibrium Constant Expression
An equilibrium constant expression can be written for every gaseous chemical system and it states that the conditions that must be attained at equilibrium. Partial pressures must be expressed in atmospheres. The equilibrium partial pressures of products appear in the numerator. The equilibrium atrial pressures of reactants appear in the denominator. Each partial pressure is raised to a power equal to its coefficient in the balanced equation. This constant is usually symbolized by Kp to show that it involves partial pressures. Kc is used to symbolize that concentration is used. Kp=Kc(RT)change in n. The expression for K depends on the form of the chemical equation written to describe the equilibrium system. The coefficient rule states that if the coefficients in a balanced equation are multiplied by a factor n, the equilibrium constant is raised to the nth power (K’=Kn). The reciprocal rule states that the equilibrium constants for forward and reverse reactions are the reciprocals of each other (Kn=1/K). The rule of multiple equilibria states that if a reaction can be expressed as the sum of two or more reactions, K for the overall reaction is the product of the equilibrium constants of the individual reactions {K(reaction 3)=K(reaction 1)x K(reaction 2)}. The equilibrium