Manufacturing Ammonia

Haber Process

Ammonia ($\text{NH}_{3}$) is manufactured from nitrogen and hydrogen by the Haber process (which is a reversible reaction).

Starting Ingredients – Nitrogen & Hydrogen

Nitrogen is obtained from the fractional distillation of liquid air.

Hydrogen is obtained from

  • steam reforming of natural gas: $\text{CH}_{4}(\text{g}) + 2\text{H}_{2}\text{O}(\text{g}) \rightarrow \text{CO}_{2}(\text{g})+4\text{H}_{2}(\text{g})$
  • cracking of petroleum fractions: $\text{C}_{2}\text{H}_{6}(\text{g}) \rightarrow \text{C}_{2}\text{H}_{4}(\text{g}) + \text{H}_{2}(\text{g})$

Equation of Reaction In Haber Process

A mixture of 1 volume of nitrogen and 3 volumes of hydrogen are reacted according to the equation below:

$$\text{N}_{2}(\text{g}) + 3 \text{H}_{2} (\text{g}) \rightleftharpoons 2 \text{NH}_{3} (\text{g}) \tag{1}$$

Reaction Conditions in Haber Process

Temperature: $500^{\circ}\text{C}$

Pressure: 250 atm

Catalyst: Finely divided iron

The ammonia formed is removed by cooling the reaction mixture. Ammonia will liquefy first and the liquid ammonia is run off.

Unreacted nitrogen and hydrogen are pumped back into the reaction chamber for further reaction.

Haber Process is Exothermic

The equation of reaction (Equation 1) is exothermic with a negative change in enthalpy.

$$\Delta \text{H} = \, – 92 \, \text{kJ mol}^{-1}$$

The above equation means that when 1 mole of nitrogen gas reacts with 3 moles of hydrogren to form 2 moles of ammonia, 92 kJ of heat energy is liberated to the surroundings.

Factors Affecting Speed of Reaction In Haber Process

$$\text{N}_{2}(\text{g}) + 3 \text{H}_{2} (\text{g}) \rightleftharpoons 2 \text{NH}_{3} (\text{g}) \tag{1}$$

As Equation 1 is a reversible equation, by Le Chatelier’s Principle, if the conditions of a system in dynamic equilibrium are altered, the system will move so as to oppose the change.

High Pressure

Increasing pressure favours the forward reaction as 4 moles of reactant gases are changed into 2 moles of product gas, decreasing the gas volume/pressure.

Hence, the yield of ammonia increases with increasing pressure.

Lower Temperature

Since the reaction is exothermic, a lower temperature would favour a shift of equilibrium to the right to increase the yield of ammonia.

However, if the temperature is too low, the rate of reaction becomes too slow. The conditions stated above represent optimum conditions to produce a reasonable amount of ammonia in a reasonable amount of time.

Presence of Catalyst

The addition of a catalyst would speed up the reaction carried out at a low temperature.

Laboratory Preparation of Ammonia

Ammonia can be prepared by heating any ammonium salt with a base. This works via displacement of ammonia from its salt.


$$\text{NH}_{4}\text{Cl} (\text{s}) + \text{NaOH}(\text{s}) \rightarrow \text{NaCl}(\text{s}) + \text{NH}_{3}(\text{g}) + \text{H}_{2}\text{O} (\text{l})$$

$$2\text{NH}_{4}\text{Cl}(\text{s}) + \text{Ca}(\text{OH})_{2}(\text{s}) \rightarrow \text{CaCl}_{2}(\text{aq}) + 2\text{H}_{2}\text{O}(\text{l}) + 2\text{NH}_{3}(\text{g})$$


  • Grind a mixture of ammonium chloride (any ammonium salt) and calcium hydroxide (any base) and place it in a round-bottomed flask
  • Mount the flask on a stand
  • Heat the mixture in the flask
  • Pass the resulting gas (ammonia gas) through calcium oxide (quicklime) to dry it
  • Collect the gas by upward delivery or downward displacement of gas

Other Notes

  • Ammonium fertilisers such as ammonium phosphates will lose their nitrogen if the farmer adds calcium hydroxide to the soil (through the process highlighted above)
  • Concentrated sulphuric acid will neutralize basic ammonia forming ammonium suphate. Hence, it cannot be used to dry ammonia gas. The equation of reaction will be $2 \text{NH}_{3}(\text{g})+\text{H}_{2}\text{SO}_{4}(\text{aq}) \rightarrow (\text{NH}_{4})_{2}\text{SO}_{4}(\text{aq})$
  • Anhydrous calcium chloride cannot be used as the drying agent for ammonia gas as it forms complex compound $\text{CaCl}_{2}.4\text{NH}_{3}$ with ammonia.

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