Haber process

Essay by cornel_stanciu March 2004

download word file, 4 pages 4.3


"¯In the late 19th century, more food was required to feed Europe and North America's growing population.

"¯The addition of nitrogen based fertilizer (NaNO3(s), NH3NO3(s)) increased crop yield.

"¯Soon, there was a high demand for the fertilizer itself.

"¯In 1909, the German company BASF investigated the possibility of producing ammonia from atmospheric nitrogen, N2(g). An year later, Haber, a professor did it.

"¯The observations were: - N2(g) and H2(g) form an equilibrium mixture with ammonia

- Optimum conditions: closed container, suitable catalyst (iron oxide), a 600 ¢XC temperature, a 30 MPa pressure

"¯This method is called the Haber Process

"¯BASF brought the rights to the process and built a plant producing 10 000 t ammonia per year


"¯At low temperatures, the reaction of N2(g) and H2(g) is not economical

"¯Adding heat increases the rate of reaction

"¯The higher the temperature, the lower the ammonia yield

"¯Haber had to balance the rate of a reaction (increased by increasing temperature) against the equilibrium of the reaction (pushed to the right by decreasing temperatures)

"¯Using a catalyst eliminates the need for high temperatures allowing the equilibrium to move to the right at low temperatures

"¯Today, we are using the Haber process to produce huge quantities of ammonia

"¯Ammonia is used to make explosives and fertilizers - dissolves in moisture present in soil (if the soil is acidic it's converted into ammonia ion) and enters the nitrogen cycle where it is converted to nitrate ions by soil bacteria.

Nitrate ions are absorbed by plants through roots and used in proteins, chlorophyll and nucleic acids

"¯Without nitrogen, the plants produce yellow leafs which cause the plant to dies prematurely


1. 1. Add N2(g) - this will increase the reactant concentration. The system will therefore try to balance the equilibrium...