Chemical Equilibrium notes: Chapter 15
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 (k f [A]) and the reverse reaction (kr[B])
become equal.
AB
This equilibrium occurs whether we start with only A in the closed container, or with only B. We can even
start with a mixture of the two. When equilibrium has been established, the ratio of concentrations will
equal a constant value.
15.2 The equilibrium constant
The constant ratio of the concentration of products and reactants is described by the law
of mass action. 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:
where Kc is the equilibrium constant. The subscript c stands for concentration (molarity).
The equilibrium expression is essentially products over reactants, each raised to its
coefficient in the balanced equation. Thus, the equilibrium expression for the Haber
process is:
The equilibrium expression of a chemical reaction depends only upon the stoichiometry
of the equation, and not on the mechanism by which the reaction takes place.
The value of the equilibrium constant depends only on the particular reaction and on
temperature. It does not depend on starting concentrations of reactants or products.
When all of the species in the equilibrium expression are gases, it is possible to write the
expression in terms of partial pressures, rather than molar concentrations. When we do
this, we use a subscript p for pressure. For the Haber process the equilibrium expression
in terms of gas pressures is
Both Kc and Kp expressions can be written for equilibria involving only gases. In
general, Kc and Kp are not equal to one another, although they can be equal under certain
circumstances.
We can convert from pressure to molarity using the ideal-gas equation
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. The constant
concentrations of solids and liquids are embodied in the value of the equilibrium
constant.
15.4 Calculating equilibrium
To solve equilibrium problems for an unknown concentration, we use the following
procedure:
1. Tabulate the known initial and equilibrium concentrations of all species involved in
the equilibrium.
2. For those species for which both the initial and equilibrium concentrations are known,
calculate the change in concentration that occurs as the system reaches equilibrium.
3. Use the stoichiometry of the reaction to calculate the changes in concentration for all
the other species in the equilibrium.
4. From the initial concentrations and the changes in concentration, calculate the
equilibrium concentrations. These are used to evaluate the equilibrium constant.
15.5 Applications of equilibrium constants
Using a technique similar to the one in the previous section, we can use a known
equilibrium constant to predict the equilibrium concentrations of reactants and products.
The Kc for the decomposition of HI gas at 450°C is 0.0195. If we were to place 0.560
mol of HI in a 5.00-L vessel and allow it to equilibrate at this temperature, what would
the equilibrium concentrations of HI, H2, and I2 be?
We plug these expressions of equilibrium concentration into the equilibrium-constant
expression.
To solve for x, we take the square root of both sides of the equation.
Solving for x gives 0.0122.
15.6 Le Châtlier’s Principle
Le Châtelier's Principle - When a system at equilibrium is stressed, the equilibrium will
shift in such a way as to minimize the effect of the stress.
Equilibria can be stressed in a number of different ways. When more reactant is added the
reaction proceeds to the right, consuming some of the additional reactant (along with a
commensurate amount of the other reactant) and producing more product. When more
product is added, the equilibrium shifts to the left, consuming some of the additional
product and producing more reactants. When a reactant or product is removed from the
reaction mixture, the equilibrium shifts in such a way as to replace some of the removed
species.
A change in pressure can stress an equilibrium. Consider the equilibrium between N2O4
and NO2. When the volume of a vessel containing an equilibrium mixture of the two
gases is suddenly halved (and the pressure of both gases suddenly doubled), the
concentration of NO2 is momentarily increased by a factor of 2. The equilibrium responds
to this change in pressure by shifting to the side with the smallest number of moles of
gas; in this case to the left (toward N2O4). At constant temperature a volume decrease will
shift an equilibrium toward the side that has the smallest number of moles of gas.
Equilibria that have the same number of moles of gas on both sides are not shifted by a
change in volume.