The Foundations of Chemistry

THE HABER PROCESS: A PRACTICAL APPLICATION

OF EQUILIBRIUM

Nitrogen, N 2 , is very unreactive. The Haber process is the economically important indus-

trial process by which atmospheric N 2 is converted to ammonia, NH 3 , a soluble, reactive

compound. Innumerable dyes, plastics, explosives, fertilizers, and synthetic fibers are made

from ammonia. The Haber process provides insight into kinetic and thermodynamic

factors that influence reaction rates and the positions of equilibria. In this process the

reaction between N 2 and H 2 to produce NH 3 is never allowed to reach equilibrium, but

moves toward it.

N 2 (g)3H 2 (g) 34 2NH 3 (g) 
H^0 92 kJ/mol

Kc3.6 108 (at 25°C)

The process is diagrammed in Figure 17-3. The reaction is carried out at about 450°C

under pressures ranging from 200 to 1000 atmospheres. Hydrogen is obtained from coal

gas or petroleum refining and nitrogen from liquefied air.

[NH 3 ]^2



[N 2 ][H 2 ]^3

17-7

Fritz Haber (1868–1934) developed
the process to provide a cheaper and
more reliable source of explosives as
Germany prepared for World War I.
(Britain controlled the seas and thus
the access to the natural nitrates in
India and Chile that were needed to
prepare explosives.) The current use of
the process is more humanitarian: most
NH 3 is used to produce fertilizers. In
the United States, approximately 135
pounds of NH 3 is required per person
per year.

728 CHAPTER 17: Chemical Equilibrium

Heat

exchanger

Catalyst

Heating

coil

N 2 and H 2

Cooling coil

Liquid ammonia

Recirculating

pump

Uncombined

N 2 and H 2

H 2 , N 2 , and ammonia

Figure 17-3 A simplified
representation of the Haber process
for synthesizing ammonia.