9.4.B - The Haber Process

The Haber process, also called the Haber-Bosch process, is the industrial process used to produce ammonia from nitrogen gas and hydrogen gas, over an enriched iron or ruthenium catalyst.

Despite the fact that 78.1% of the air we breathe is nitrogen, the gas is relatively unavailable because it is so unreactive: nitrogen molecules are held together by strong triple bonds. It was not until the early 20th century that the Haber process was developed to harness the atmospheric abundance of nitrogen to create ammonia, which can then be oxidized to make the nitrates and nitrites essential for the production of nitrate fertilizer and explosives.

Prior to the discovery of the Haber process, ammonia had been difficult to produce on an industrial scale. The Haber process is important today because the fertilizer generated from ammonia is responsible for sustaining one-third of the Earth's population. It is estimated that half of the protein within human beings is made of nitrogen that was originally fixed by this process, the remainder was produced by nitrogen fixing bacteria and archaea.

In this topic students:

  • Evaluate the significance of this discovery at that time in world history
  • Identify some of the uses of ammonia
  • Use Le Chatelier's Principle to predict the effect of changing reaction conditions on an equilibrium system
  • Describe and explain the reaction conditions that maximise both rate and yield in the production of ammonia in the Haber Process
  • Describe the process of producing ammonia on an industrial scale


Ammonia or azane is a weak base and compound of nitrogen and hydrogen with the formula NH3. It is a colourless gas with a characteristic pungent odour. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to fertilisers. Ammonia is also a building block for the synthesis of explosives and many pharmaceuticals and is also used in many commercial cleaning products. In 2006, worldwide production was estimated at 146.5 million tonnes

Although in wide use, ammonia is both caustic and hazardous. Ammonia is often called anhydrous ammonia because of the absence of water in the material. Because ammonia boils at -33.34°C at atmospheric pressure, the liquid must be stored under high pressure or at low temperature. Household ammonia or ammonium hydroxide is a solution of NH3 in water.


Why Make Ammonia?

Approximately 83% of industrially produced ammonia is used to make fertilisers. When applied to soil, it helps provide increased yields of crops such as corn and wheat. All plants need nitrogen to make amino acids, proteins and DNA, but the nitrogen in the atmosphere is not in a form that they can use. Nitrogen fixation is the name given to any process by which nitrogen (N2) in the atmosphere is converted into ammonia. Atmospheric nitrogen is relatively inert and it does not easily react with other chemicals to form new compounds. Fixation processes free up the nitrogen atoms from their diatomic form to be used in other ways.

Nitrogen fixation is essential for all forms of life because nitrogen is required to biosynthesize basic building blocks of plants, animals and other life forms such as nucleotides for DNA and RNA and amino acids for proteins. Nitrogen fixation through the Haber process is essential for agriculture and the manufacture of fertiliser as the ammonia acts as the precursor to various ammonium salts such as ammonium nitrate and sulfate that provide usable nitrogen to plants and then to animals.

Ammonia also acts directly or indirectly as the precursor to most nitrogen-containing compounds. Virtually all synthetic nitrogen compounds are derived from ammonia. An important derivative is nitric acid. This key material is generated via the Ostwald process by oxidation of ammonia with air over a platinum catalyst at 700–850°C

NH3 + 2O2 ⇒ HNO3 + H2O

Both nitric acid and ammonia are used in the production of explosives such as TNT and ammonium nitrate.

Household ammonia is a solution of ammonia in water (aqueous ammonium hydroxide) and is used as a general purpose cleaner for many surfaces. Because ammonia results in a relatively streak-free shine, one of its most common uses is to clean glass, porcelain and stainless steel. It is also frequently used for cleaning ovens and soaking items to loosen baked-on grime.


Significance in World History

In 1912, Europe was on the brink of World War I. Since ammonia was used in the manufacture of explosives and nitrogen-based fertilisers, demand was far greater than actual supply. At the time most of the world's supply of nitrates came from a giant salt petre on the coast of Chile. At the time of WWI the allied forces had blocked the supply route and Germany needed a new source. Haber developed the process for the synthesis of ammonia in response to the need for more nitrogen compounds. The manufacture of fertilisers for continued food production and the continued manufacture of explosives were high priorities for Germany on the brink of war. Haber's contribution overcame supply problems for Germany's war effort and is held by many historians to be responsible for prolonging the war.


Ammonia Synthesis

The synthesis of ammonia is a reversible reaction with the forward reaction being exothermic.

N2 (g) + 3H2 (g) ⇔ 2NH3 (g) ΔH = -92 kJ mol-1

Starting materials for the synthesis of

Despite the fact that 78.1% of the air we breathe is nitrogen, the gas is relatively unavailable because it is so unreactive: nitrogen molecules are held together by strong triple bonds. It was not until the early 20th century that the Haber process was developed to harness the atmospheric abundance of nitrogen to create ammonia, which can then be oxidized to make the nitrates and nitrites essential for the production of nitrate fertiliser and explosives.

Ammonia is synthesised from gaseous nitrogen and hydrogen. The nitrogen is extracted from air by distillation and the hydrogen can by produced by a number of methods including the reaction of super heated steam with methane forming carbon monoxide and hydrogen.

Ammonia is a gas at room temperature so to synthesise and collect it the system must be closed. Closed systems bring reversible reactions to equilibrium so it is important to consider the yield of ammonia when the system comes to equilibrium. Under standard conditions of temperature and pressure, the reaction is slow and has a low yield (an equilibrium position favouring the reactants).

It was Haber who used the principles of chemical kinetics and equilibrium to ensure the reaction conditions were set so as to maximise both the rate and the yield.


The Haber Process - A Delicate Balancing Act

An increase in temperature always increase the rate of a chemical reaction. Increasing the temperature for the synthesis of ammonia will increase the rate of the forward and reverse reactions and ensure the system gets to equilibrium quickly. However, because the forward reaction is exothermic, an increase in temperature will cause a shift toward the reactants and decrease the yield. According to Le Chatelier's principle, when a change is made to a system in equilibrium, the system will shift to partially counteract the change. This reaction shifts in reverse because the reverse reaction is endothermic and will therefore counteract the increased temperature by absorbing some of the heat. Therefore, increasing temperature speeds up the reaction that produces ammonia but it reduces the yield of ammonia. If the reaction were carried out at a low temperature to maximise the yield, the molecules would not have sufficient energy to react.

Haber was able to overcome the temperature issue by developing a suitable catalyst for the reaction. The catalyst lowered the activation energy for the reaction and helped increase the rate that had been lowered by the need for a lower temperature to increase the yield. The catalyst developed was a mixture of Fe2O3 and Fe3O4.

The reaction is usually run at about 400°C, this temperature being a compromise between a higher temperature which would reduce the yield of ammonia and a lower temperature which would reduce the rate of the reaction.

Pressure is another factor that will affect the yield of ammonia in the Haber Process. An increase in pressure will increase the rate of the reaction because there will be more successful collisions between reactant molecules. It will also cause a shift in the equilibrium toward the production of ammonia. This is because when pressure is raised, the equilibrium will shift to partially counteract the high pressure. The forward shift will use four moles of molecules and form two moles of molecules, reducing pressure. In the process more ammonia is formed and its yield is increased. The favoured pressure is around 250 atm.


Industrial Production

In 1909, the Haber process was purchased by the German chemical company BASF, which assigned Carl Bosch the task of scaling up Haber's tabletop machine to industrial-level production.

The Haber process on an industrial scale showing each of the reaction conditions being maintained to ensure the
best rate and yeild of ammonia.

The diagram above shows how ammonia is produced via a continuous industrial process designed to ensure that the system never gets to equilibrium. The reaction vessel must be continually monitored to ensure that optimal conditions are met for the production of ammonia and to ensure a safe environment for workers with minimal waste. The factors that need to be monitored include:

  • Temperature - This needs to be maintained at around 400°C. If it is too high the yield of ammonia drops. If it is too low the reaction rate drops.

  • Pressure - This needs to be maintained at around 250 atm. If it is too high the vessel may not be strong enough to hold the gases and an explosion may occur. If it is too low the yield of ammonia will drop.

  • Reactants - The ratio of nitrogen to hydrogen needs to be maintained at 1:3 otherwise one reactant will be in excess and wasted. These gases are also pressurised to concentrate them before the enter the system.

  • Product Removal - The ammonia gas is liquefied and removed which reduces its concentration and drives the equilibrium to the right increasing the yield.