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Uses and Production of Ammonia by the Haber Process

Key Concepts

●In 1909 Fritz Haber established the conditions under which nitrogen, N₂(g), and hydrogen, H₂(g), would combine to produce ammonia, NH₃(g) using: (i) medium temperature (≈500^oC) (ii) very high pressure (≈250 atmospheres, ≈25,500kPa) (iii) a catalyst (a porous iron catalyst prepared by reducing magnetite, Fe₃O₄). Osmium is a much better catalyst for the reaction but is very expensive. ●This process produces an ammonia, NH₃(g), yield of approximately 10-20%. ●The Haber synthesis was developed into an industrial process by Carl Bosch. Please do not block ads on this website. No ads = no money for us = no free stuff for you!

Equilibrium Considerations

The reaction between nitrogen gas and hydrogen gas to produce ammonia gas is an exothermic equilibrium reaction, releasing 92.4 kJ mol^-1 of energy at 298 K (25^oC).

N₂(g)
nitrogen

3H₂(g)
hydrogen

heat, pressure, catalyst

2NH₃(g)
ammonia

ΔH = -92.4 kJ mol^-1

OR, including the enthalpy change as a product of the reaction:

N₂(g)
nitrogen

3H₂(g)
hydrogen

heat, pressure, catalyst

2NH₃(g)
ammonia

+ 92.4 kJ mol^-1

ByLe Chetalier's Principle, increasing the pressure on the reaction mixture favours the formation of ammonia gas: Increasing the pressure causes the equilibrium position to move to the right resulting in a higher yeild of ammonia since there are more gas molecules on the left hand side of the equation (4 in total) than there are on the right hand side of the equation (2). Increasing the pressure means the system adjusts to reduce the effect of the change, that is, to reduce the pressure by having fewer gas molecules. ByLe Chetalier's Principle, decreasing the temperature of the reaction mixture favours the formation of ammonia gas: Decreasing the temperature causes the equilibrium position to move to the right resulting in a higher yield of ammonia since the reaction is exothermic (releases heat). Reducing the temperature means the system will adjust to minimise the effect of the change, that is, it will produce more heat since energy is a product of the reaction, and will therefore produce more ammonia gas as well. However, the rate of the reaction at lower temperatures is extremely slow, so a higher temperature must be used to speed up the reaction which results in a lower yield of ammonia. The equilibrium expression for this reaction is:

K_c =

[NH_3(g)]²
[N_2(g)][H_2(g)]³

As the temperature increases, the value of the equilibrium constant decreases as the yield of ammonia decreases.

Temperature (^oC)	K_c	trend
25	6.4 x 10²	higher K_c
200	4.4 x 10^-1	↓
300	4.3 x 10^-3	↓
400	1.6 x 10^-4	↓
500	1.5 x 10^-5	lower K_c

In summary, higher pressure favours formation of ammonia gas, lower temperature favours formation of ammonia gas BUT the reaction rate is slow at lower temperatures. Do you know this? Join AUS-e-TUTE! Play the game now!

Reaction Rate (Kinetic) Considerations

●A catalyst such as an iron catalyst is used to speed up the reaction by lowering the activation energy so that the N₂ bonds and H₂ bonds can be more readily broken. ●Increased temperature means more reactant molecules have sufficient energy to overcome the energy barrier to reacting (activation energy) so the reaction is faster at higher temperatures (but the yield of ammonia is lower as discussed above). A temperature range of 400-500^oC is a compromise designed to achieve an acceptable yield of ammonia (10-20%) within an acceptable time period.

Refer to the graph on the right which records the yield of ammonia as a percentage at different pressures:

At 200^oC and pressures above 750 atm there is an almost 100% conversion of reactants to the ammonia product.

Since there are difficulties associated with containing larger amounts of materials at this high pressure, lower pressures of around 200 atm are used industrially.

By using a pressure of around 200 atm and a temperature of about 500^oC, the yield of ammonia is 10-20%, while costs and safety concerns in the building and during operation of the plant are minimised

During industrial production of ammonia, the reaction never reaches equilibrium as the gas mixture leaving the reactor is cooled to liquefy and remove the ammonia. The remaining mixture of reactant gases are recycled through the reactor. The heat released by the reaction is removed and used to heat the incoming gas mixture. Do you understand this? Join AUS-e-TUTE! Take the test now!

Uses of Ammonia

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A Brief History of Ammonia Production

At the beginning of the 20^th century there was a shortage of naturally occurring, nitrogen-rich fertilisers, such as Chile saltpetre, which prompted the German Chemist Fritz Haber, and others, to look for ways of combining the nitrogen in the air with hydrogen to form ammonia, which is a convenient starting point in the manufacture of fertilisers. This process was also of interest to the German chemical industry as Germany was preparing for World War I and nitrogen compounds were needed for explosives. The hydrogen for the ammonia synthesis was made by the water-gas process (a Carl Bosch invention) which involves blowing steam through a bed of red hot coke resulting in the separation of hydrogen from oxygen. The nitrogen was obtained by distillation of liquid air, then by cooling and compressing air. These days, the hydrogen is produced by reforming light petroleum fractions or natural gas (methane, CH₄) by adding steam:

Industry

Use

Fertilser

production of:

ammonium sulfate, (NH₄)₂SO₄

ammonium phosphate, (NH₄)₃PO₄

ammonium nitrate, NH₄NO₃

urea, (NH₂)₂CO, used in the production of barbiturates (sedatives), is made by the reaction of ammonia with carbon dioxide as shown below:

CO₂
carbon dioxide

2NH₃
ammonia

H₂NCOONH₄
ammonium carbonate

heat,
pressure

(NH₂)₂CO
urea

Chemicals

synthesis of: ●nitric acid, HNO₃, which is used in making explosives such as TNT (2,4,6-trinitrotoluene), nitroglycerine which is also used as a vasodilator (a substance that dilates blood vessels) and PETN (pentaerythritol nitrate). ●sodium hydrogen carbonate (sodium bicarbonate), NaHCO₃ ●sodium carbonate, Na₂CO₃ (Solvay Process) ●hydrogen cyanide (hydrocyanic acid), HCN ●hydrazine, N₂H₄ (used in rocket propulsion systems)

Explosives

ammonium nitrate, NH₄NO₃

Fibres and Plastics

nylon, -[(CH₂)₄-CO-NH-(CH₂)₆-NH-CO]-,and other polyamides

Refrigeration

used for making ice, large scale refrigeration plants, air-conditioning units in buildings and plants

Pharmaceuticals

used in the manufacture of drugs such as sulfonamide which inhibit the growth and multiplication of bacteria that require p-aminobenzoic acid (PABA) for the biosynthesis of folic acids, anti-malarials and vitamins such as the B vitamins nicotinamide (niacinamide) and thiamine.

Pulp and Paper

ammonium hydrogen sulfite, NH₄HSO₃, enables some hardwoods to be used

Mining and Metallurgy

used in nitriding (bright annealing) steel, used in zinc and nickel extraction

Cleaning

ammonia in solution is used as a cleaning agent such as in 'cloudy ammonia'

CH₄(g) + H₂O(g)

Ni catalyst
→
700^oC

CO(g) + 3H₂(g)

Enough steam is used to react with about 45% of the methane (CH₄), the rest of the methane is reacted with air:

2CH₄(g) +

O₂(g) + 4N₂(g)
(air)

Ni catalyst
→

2CO(g) + 4H₂(g) + 4N₂(g)

All the carbon monoxide (CO) in the mixture is oxidised to CO₂ using steam and an iron oxide catalyst:

CO(g) + H₂O(g)

iron oxide catalyst
→

H₂(g) + CO₂(g)

The carbon dioxide (CO₂) is removed using a suitable base so that only the nitrogen gas (N₂) and hydrogen gas (H₂) remain and are used in the production of ammonia (NH₃). In ammonia production the hydrogen and nitrogen are mixed together in a ratio of 3:1 by volume and compressed to around 200 times atmospheric pressure. Do you know this? Join AUS-e-TUTE! Play the game now! Warning! Some content on this page could not be displayed. Please enable javascript and pop-ups to view all page content.