Chapter 15

Chemical Equilibrium

Overview• 15.1 The Concept of Equilibrium• When a reaction takes place, both the forward process (the reaction as we have written it) and the reverse reaction occur. Equilibrium is the point at

which both processes are occurring at the same rate, with no net change of reactant or product concentration. Equilibrium can be achieved starting with only reactants, only products, or some of each. You will use a simulation in this section to plot concentrations of reactants and products as equilibrium is achieved.

• 15.2 The Equilibrium Constant• There is a constant relationship between equilibrium concentrations of reactants and products at a given temperature. You will learn to write

equilibrium expressions for chemical reactions, and a simulation will illustrate the significance of an equilibrium constant's magnitude.

• 15.3 Heterogeneous Equilibria• Some equilibria involve substances all in the same phase. Many others involve substances in different phases. The concentrations of liquids and solids

do not appear in equilibrium expressions for a reaction, and you'll see some examples of heterogeneous equilibria and their equilibrium expressions.

• 15.4 Calculating Equilibrium Constants• Often, only one of the species present in an equilibrium mixture can be measured. We can generally use the stoichiometry of the reaction to deduce

the concentrations of the other species in the chemical equation. We can also calculate the equilibrium constant for such a reaction.

• 15.5 Applications of Equilibrium Constants• Equilibrium constants can be used to predict the equilibrium concentrations of reactants and products. You will learn how to use K values to calculate

the composition of an equilibrium mixture for several different situations.

• 15.6 Le Châtelier's Principle• An important principle throughout chemistry, Le Châtelier's principle is introduced in this section. Two animations illustrate the effects of certain

types of stress that can be applied to a system at equilibrium.

15.1 The Concept of Equilibrium

• The reaction of pure compound A, with initial concentration [A]0 . After a time the concentrations of A and B do not change. The reason is that the rates of the forward reaction (kf [A]) and the reverse reaction (kr[B]) become equal.

15.2 The Equilibrium Constant

• For the general reaction aA + bB ↔ cC + dD, the concentrations of reactants and products at equilibrium are related by the equilibrium-constant expression or simply the equilibrium expression

15.3 Heterogeneous Equilibria• Equilibria that involve more than one phase are

called heterogeneous equilibria. When writing the equilibrium expression for a heterogeneous equilibrium, we do not include the concentrations of solids or liquids. Regardless of how much solid or liquid is present, its concentration in terms of moles per liter remains constant.

Ex: SnO2 (s) + 2CO (g) ↔ Sn (s) + 2 CO2 (g)

Kc = [CO2 (g)]2/[CO (g)]2

15.4 Calculating Equilibrium Constants

• ICE box!!!

Initial (Initial molarity/pressure of reactants and products)

Change (Change in molarity/pressure, calculated from the stoichiometry of the reaction)

Equilibrium (Molarity/Pressure at Equilibrium = Initial – change)

Example

• 0.100 mol of N2O4 gas is placed in a 2.00-L vessel and allowed to equilibrate at 110°C. At equilibrium the concentration of NO2 is 0.072 M. Calculate the value of Kc for the reaction at this temperature.

N2O4 (g) 2 NO2 (g)Start (0.100 mol/2.00L) =

0.0500 M0 M (If there is none mentioned in the problem, assume there is none present)

Change + 0.072 M

Equilibrium 0.072 M

Example continued

N2O4 (g) 2 NO2 (g)

• Change in N2O4 (g) can be calculated from the stoichiometry of the equation

• Then the rest of the box can be filled in

Start (0.100 mol/2.00L) = 0.0500 M

0 M

Change (-0.072M/2) = - 0.036 M + 0.072 M

Equilibrium 0.014 M 0.072 M

Example finished

Start (0.100 mol/2.00L) = 0.0500 M

0 M

Change (-0.072M/2) = - 0.036 M + 0.072 M

Equilibrium 0.014 M 0.072 M

15.6 Le Châtelier's Principle

• Le Châtelier's Principle. The principle can be stated as follows: When a system at equilibrium is stressed, the equilibrium will shift in such a way as to minimize the effect of the stress.

Figure 15. 12. When H2 is added to an equilibrium mixture of N2, H2, and NH3, a portion of the H2 reacts with N2 to form NH3, thereby establishing a new equilibrium position.

Equilibrium & Temperature

• The heat given off by an exothermic chemical reaction can be thought of as a product of the reaction.

• Likewise, the heat consumed by an endothermic process can be thought of as a reactant.

• The addition of a catalyst changes the rate at which equilibrium is established, but it does not affect the position of the equilibrium.