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Case Study: The Haber Process

The Haber process is the production of ammonia from a reaction between nitrogen and hydrogen, because of an iron substitute. This process is known for the commercial synthesis of ammonia. There is great abundance of nitrogen in the air when it is combined with hydrogen under extreme pressure and high temperature. This process is a great example of chemical equilibrium.


The Haber process, also known as the Haber-Bosch Process, was founded by Fritz Haber and Carl Bosch, both who were German Chemists. Haber discovered the conditions for the formation of ammonia, and Bosch discovered the work of high-pressure on chemical reactions (developed into industrial process). Both were awarded the Nobel Prize. During the 1920’s, there was a shortage of the world's supply for fixed nitrogen. Nitrogen was mainly used for fertilizer. Fertilizer was used in order to produce food, so that in WWI people could continue to fight. It only requires 1 percent of the world's energy to make 500 million tons of artificial fertilizer per year, which, in turn, helps feed 40 percent of the world's population.


The Haber process takes nitrogen gas from air and combines it with molecular hydrogen gas to form ammonia gas. This is an exothermic reaction, meaning it releases energy so that the sum of the enthalpies of N2 and H2 (the reactants) is greater than the enthalpy of NH3 (the products).

N2(g) + 3H2(g) → 2NH3(g)     ΔH=-92.4 kJ

which is a reversible reaction:

2NH3(g) → N2(g) + 3H2(g)     ΔH=+ 92.4 kJ mol-1

Here's a visual to help convey the process:


From the flow chart above, we can see that methane and steam combine to form hydrogen and carbon monoxide, which in turn releases hydrogen. The hydrogen then combines with oxygen from the air to produce water. Finally, nitrogen gas is released which combines with hydrogen gas to form ammonia. This takes place under high pressure and temperature and with an iron catalyst*** (mentioned later on).

Le Châtelier's Principle

The Haber process incorporates Le Chatlier's Principle, which is a good example of equilibrium principles. Uses of Le Chatlier's Principle are reversible reactions and reversible reactions involving gases. Chemical equilibrium is when a reaction has no tendency to change the quantity of the products and reactants, so the reaction can go both ways.

  • Increasing the pressure and decreasing the temperature results in the higher yield of ammonia by causing a move of the reaction to the right.
  • Because there are more molecules on the left side than the right side, when the pressure is increased, the system adapts to the change by moving the molecules left to right to decrease the overall pressure.
  • For temperature, it moves from right to left when the temperature drops is because of the process being exothermic, where heat is released.
  • The system adjusts to lessen the change, so it would make more heat to compensate for the energy lost, since that is the product of this. If more energy is made, then that would mean more ammonia is made, too. Even though decreasing the temperature is a slow reaction, if the temperature was increased to speed up the reaction, it would produce a smaller amount of ammonia yield.

Examples of Le Châtelier's Principle

N2 (g) + 3H2 (g) ↔ 2NH3 (g)

  • If the volume is decreased here, it has the same result as when the pressure is decreased.
  • If the pressure is increased, in this equation, it will move right because there are fewer gas molecules are produced going to the right then the backwards one.

When you increase the pressure so that the least amount of molecules will be formed, there won't be an increase in collisions. However, if more gas molecules are formed, there will be an increase in collisions, thus moving the way that will produce the least amount of molecules.

PCl3(g) + Cl2(g) ↔ PCl5(g) + energy

  • If the temperature increased here, it would shift to the left because it would use the extra energy that is left over in the equation.
  • In the equation it can go both ways, left and right, so one way is endothermic and the other is exothermic.
  • If the temperature is increased, it would benefit the endothermic reactions, so there would be more energy for the reaction to take in. * If the temperature is decreased, it would be in the favor of an exothermic reaction, because then the reaction can release heat.

H2+ I2 ↔ 2HI

  • Removing H2 from the system will cause it to move towards the left, so more H2 can be made.
  • By lowering the concentration of one substance, the equation will shift in that direction so that it can produce more of that which was lowered.
  • If one concentration increases, it will move in the direction that would help lower the concentration.


A catalyst is used to speed up a reaction by lowering the activation energy. So, in this reaction, the iron catalyst is used to lower the activation energy so that the N2 and H2 can be easier to break down.


From the above diagram the blue curve represents what the activation energy would look like with an iron catalyst. The purple curve shows what the activation energy would look like if a catalyst was not involved. Note that without the catalyst, the activation energy is much bigger. A catalyst is needed to break down the nitrogen and hydrogen gases.

Economic Effects

Nitric Acid

By mixing one part ammonia to nine parts air with the use of a catalyst, the ammonia will get oxidized to nitric acid.

4 NH3 + 5 O2 → 4 NO + 6 H2O

2 NO + O2 → 2 NO2

2 NO2 + 2 H2O → 2 HNO3 + H2


Ammonia is mixed with irrigation water to form a solution for fertilizer ingredients. Fertilizer is mainly used in fields that grow crops such as corn, barley, sorghum, rapeseed, soybean, and sunflower.


  1. Petrucci, et al. General Chemistry: Principles & Modern Applications: AIE (Hardcover). Upper Saddle River: Pearson/Prentice Hall, 2007


  • Nathalie Interiano (UCD)

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Last Modified
12:37, 16 Dec 2013

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