Physical Chemistry , 1st ed.

Because for dilute aqueous concentrations the molality is approximately

equal to the molarity, it is not uncommon to write equilibrium concentrations

in units of molarity. (In fact, this is how it is usually done in introductory

courses.) However, this adds an additional approximation in our expression of

reaction quotients and equilibrium constants.

Example 5.8

What is the proper expression for the equilibrium constant, in terms of pres-

sures, for the following chemical equilibrium? Assume that conditions are

near standard pressures.

Fe 2 (SO 4 ) 3 (s) Fe 2 O 3 (s) 3SO 3 (g)

Solution

The correct expression for the equilibrium constant is

K (aSO 3 )^3 

p

p

SO

°

^3



3

The other species in the equilibrium are condensed phases and, if we are close

to standard pressures, do not affect the numerical value ofK.

Example 5.9

What is the proper expression for the equilibrium constant for the following

chemical equilibrium in terms of concentration and partial pressures? This

equilibrium is partly responsible for the atmospheric production of acid rain.

2H 2 O () 4NO (g) 3O 2 (g) 4H(aq) 4NO 3 (aq)

Solution

The proper equilibrium expression is

K

4



4

K



4



3

As a condensed phase, H 2 O () does not appear in the expression.

5.5 Changes in Equilibrium Constants

Despite their names, the numerical values of equilibrium constants can vary

depending on conditions, usually with varying temperatures. The effects of

temperature on equilibria are easy to model. In the last chapter, we derived the

Gibbs-Helmholtz equation as







T



T

G



p



T

H

 2

When applied to a chemical reaction under conditions of standard pressure, it

can be rewritten it as







T

^ rx

T

nG°

p

(^) r
T
xn
2
H°
pO 2

p°
pNO

p°
NO 3 mNO 3 

m°
HmH

m°

JQPJ

(aSO 3 )^3 aFe 2 O 3



aFe 2 (SO 4 ) 3

JQPJ

132 CHAPTER 5 Introduction to Chemical Equilibrium