Page 244 - Theory and Problems of BEGINNING CHEMISTRY

CHAP. 16] RATES AND EQUILIBRIUM 233

EXAMPLE 16.3. Suppose that, under a certain set of conditions, a mixture of nitrogen, hydrogen, and ammonia is at
equilibrium. The concentration of hydrogen is 0.0350 mol/L; that of nitrogen is 0.0100 mol/L, and that of ammonia is
0.0250 mol/L. Now 0.0003 mol of hydrogen is added to 1.00 L of the mixture. What is the widest possible range of hydrogen
concentration in the new equilibrium?

Ans. Before addition of the extra hydrogen, its concentration was 0.0350 mol/L. After addition of the extra hydrogen,
but before any equilibrium shift could take place, there is 0.0353 mol/L. The shift of the equilibrium uses up some
but not all of the added hydrogen, and so the final hydrogen concentration must be above 0.0350 mol/L and below
0.0353 mol/L. Some nitrogen has been used up and its final concentration must be less than its original concentration,
0.0100 mol/L. Some additional ammonia has been formed, and so its final concentration has been increased over
0.0250 mol/L. Notice that Le Chˆatelier’s principle does not tell us how much of a shift there will be, but only
qualitatively in which direction a shift will occur.

Temperature Change
To consider the effect of temperature change, let us rewrite one of the preceeding equations including the
heat involved:
−→
3H 2 + N 2 ←− 2NH 3 + heat
How do we get the temperature of the system to rise? By adding heat. When we add heat, this equilibrium system
reacts to reduce that stress, that is, to use up some of the added heat. It can use up heat in the reverse reaction, the
decomposition of ammonia to hydrogen and nitrogen. When the substances written as products of the reaction
(on the right side of the equation) react to produce more reactants (on the left side of the equation), we say that
the reaction has shifted to the left. When the opposite process occurs, we say that the equilibrium has shifted to
the right. Thus, raising the temperature on this system already at equilibrium causes a shift to the left; some of
the ammonia decomposes without being replaced.

EXAMPLE 16.4. What happens to the above system at equilibrium if the temperature is lowered?
Ans. We lower the temperature of a system by removing heat. The equilibrium tries to minimize that change by restoring
some heat; it can do that by shifting to the right. Thus, more nitrogen and hydrogen are converted to ammonia.

The Effect of Pressure
Pressure affects gases in a system much more than it affects liquid or solids. We will investigate the same
ammonia, hydrogen, nitrogen system discussed before. If the system is at equilibrium, what will an increase in
pressure by the chemist do to the equilibrium? The system will shift to try to reduce the stress, as required by
Le Chˆatelier’s principle. How can this system reduce its own pressure? By reducing the total number of moles
present. It can shift to the right to produce 2 mol of gas for every 4 mol used up:

−→
3H 2 + N 2 ←− 2NH 3
Of course, it need not shift much. For example, if 0.00030 mol H 2 reacts with 0.00010 mol N 2 to produce
0.00020 mol NH 3 , the total number of moles will have been reduced (by 0.00020 mol), and the pressure will
therefore have been reduced.

EXAMPLE 16.5. What will be the effect of increased pressure caused by decreasing the volume on the following system
at equilibrium?
−→
2NH 3 ←− 3H 2 + N 2
Ans. Again some nitrogen and hydrogen will be converted to ammonia. This time, since the equation is written in the
reverse of the way it was written above, the equilibrium will shift to the left. Of course, the same physical effect
is produced. More ammonia is formed. But the answer in terms of the direction of the shift is different since the
equation is written “backward.”