The value of Kcis 3.6 108 at 25°C. This very large value of Kcindicates that at equi-
libriumvirtually all of the N 2 and H 2 (mixed in a 13 mole ratio) would be converted into
NH 3. At 25°C, the reaction occurs so slowly, however, that no measurable amount of NH 3
is produced within a reasonable time. Thus, the large equilibrium constant (a thermody-
namic factor) indicates that the reaction proceeds toward the right almost completely. It
tells us nothing,however, about how fast the reaction occurs (a kinetic factor).
There are four moles of gas on the left side of the equation and only two moles of gas
on the right, so increasing the pressure favors the production of NH 3. The Haber process
is therefore carried out at very high pressures, as high as the equipment will safely stand.
The reaction is exothermic (H^0 rxnis negative), so increasing the temperature favors
the decompositionof NH 3 (the reverse reaction). But, the rates of both forward and reverse
reactions increase as the temperature increases.
The addition of a catalyst of finely divided iron and small amounts of selected oxides
also speeds up both the forward and reverse reactions. This allows NH 3 to be produced
not only faster but at a lower temperature, which increases the yield of NH 3 and extends
the life of the equipment.
Table 17-1 shows the effects of increases in temperature and pressure on the equilib-
rium yield of NH 3 , starting with 13 mole ratios of N 2 H 2. Kcdecreases by more than
ten orders of magnitude, from 3.6 108 at 25°C to only 1.4 10 ^2 at 758°C. This tells
us that the reaction proceeds very far to the leftat high temperatures. Casual examination
of the data might suggest that the reaction should be run at lower temperatures, because
a higher percentage of the N 2 and H 2 is converted into NH 3. The reaction occurs so
slowly, however, even in the presence of a catalyst, that it cannot be run economically at
temperatures below about 450°C.
The emerging reaction mixture is cooled down, and NH 3 (bp33.43°C) is removed
as a liquid. This prevents the reaction from reaching equilibrium and favors the forward
reaction. The unreacted N 2 and H 2 are recycled. Excess N 2 is used to favor the reaction
to the right.
APPLICATION OF STRESS TO A SYSTEM
AT EQUILIBRIUMWe can use equilibrium constants to determine new equilibrium concentrations that result
from adding one or more species to, or removing one or more species from, a system at
equilibrium.
17-8
In practice, the mixed reactants are
compressed by special pumps and
injected into the heated reaction vessel.Ten orders of magnitude is 10^10 , that
is, 10 billion.
1 1010 10,000,000,00017-8 Application of Stress to a System at Equilibrium 729TABLE 17-1 Effect of T and P on Yield of AmmoniaMole % NH 3 in Equilibrium Mixture°C Kc 10 atm 100 atm 1000 atm209 650 51 82 98
467 0.5 4 25 80
758 0.014 0.5 5 13A nighttime photo of a large plant
for the commercial production of
ammonia, NH 3. Such an installation
can produce up to 7000 metric tons
of ammonia per day. There are
nearly 100 such plants in the world.