The Haber process

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  1. Equilibrium in real life chemistry.
  2. The Haber process: introduction.
  3. Changing conditions in the Haber process.
  4. How do we maximise production?
Topic Notes

In this lesson, we will learn:

  • To recall the Haber process and the chemicals involved.
  • How to explain how changing reaction conditions affects the Haber process.
  • How to explain the optimum conditions of the Haber process using Le Chatelier's principle.


  • Some very important chemical processes happen in an equilibrium. One example is the Haber process, which converts nitrogen and hydrogen gas into ammonia, a common fertilizer now used to increase crop yield.

    N2(g)+3H2(g) 2NH3(g)\mathrm{N_{2(g)} + 3H_{2(g)} \ \rightleftharpoons 2NH_{3(g)}}

    The Haber process was invented by a German chemist called Fritz Haber in the early 1900's as a way to get to nitrates which were needed to make explosives at the time. Today, over a hundred million tons of ammonia are made every year, and the figure has steadily increased since the 1940s.
    The Haber process is an exothermic reaction that happens in equilibrium, which creates the following problem: increasing the temperature will make the reaction faster, but will shift the equilibrium to make less products!
  • As we saw in C12.1.9: Entropy and spontaneous reactions, when enthalpy and entropy favor opposite sides of a reaction, an equilibrium is usually the result. Chemical companies want and need to run the Haber process to make as much ammonia as possible using the least resources as possible. This means:
    • Make it cheap to run, using less heat, pressure, energy and chemicals.
    • Make it low-maintenance, so things don't need replacing, processing or support to work properly.
    • Do it quickly, with the fastest reaction rate possible.
  • Using our knowledge of equilibrium, what can we do to the reaction conditions to make as much ammonia as possible?
    • Increasing the operating temperature…
      • …will increase the rate of reaction, which will speed up both the forward and the reverse reaction.
      • …will shift the equilibrium to the left, which favors the reverse reaction and makes more reactants. The reverse reaction is endothermic and will absorb the added heat from the increased temperature.
    • Decreasing the operating temperature…
      • …will lower the rate of reaction, slowing both forward and reverse reactions and the time taken for the process to reach equilibrium.
      • …will shift the equilibrium to the right, making more products by favoring the forward reaction. Favoring the exothermic forward reaction releases heat, countering the decreasing temperature.
    • Increasing the pressure…
      • …will shift the equilibrium to the right to favor the products. There are less moles of gas in the products, so this counters the increased pressure applied to the system.
      • …will be extremely expensive and quite dangerous for a reaction running on a large scale!
    • Decreasing the pressure…
      • Will shift the equilibrium to the left to favor the reactants to counter the decrease in pressure.
    What else could be added to the equilibrium mixture to make the forward reaction (or any reaction!) run faster?
  • Resolving these issues, chemical industry uses compromise conditions to extract the most ammonia from this process:
    • Temperature is set at around 450°C, which is high enough for a high rate of reaction without excessively favoring the reactants in the equilibrium.
    • Pressure is around 200 atmospheres (200 times higher than normal air pressure). This is high enough to favor the products in the reaction, but not too high that it is extremely expensive or dangerous to run machinery.
    • An iron catalyst is added. This speeds up the rate of the reaction both ways and decreases the time taken to get to equilibrium.