Doc Brown's Chemistry KS4 science GCSE/IGCSE/O level Revision Notes


The industrial manufacture of ammonia is described in terms of the reaction conditions employed in the Haber Synthesis of ammonia from hydrogen and nitrogen e.g. temperature, pressure, iron catalyst, recycling unreacted gases, condensing out of ammonia. Equilibrium theory and rates of reaction factors are explained and discussed to arrive at the optimum (most efficient) reaction conditions to make ammonia gas. These revision notes on the Haber synthesis and artificial synthetic nitrogen fixation from air, should prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE chemistry (Gateway & 21st Century) GCSE (91), (9-5) & (5-1) science courses.

Index: A Reversible Reactions  *  B Reversible reactions & equilibrium  *  C Haber Synthesis of ammonia (this page)

D(a) The Uses of ammonia-nitric acid-fertilisers  *  D(b) Fertilisers-environmental problems  *  E The nitrogen cycle 

Advanced A Level Notes - Chemical Equilibrium * Advanced A Level Notes - Nitrogen & Ammonia

PART C 3a. The Synthesis of ammonia - The Haber Process of Nitrogen Fixation


The Haber process is used by the chemical industry to manufacture ammonia which is used to make nitric acid and fertilisers. The raw materials for the Haber process are nitrogen from air and hydrogen is usually obtained from natural gas. The purified gases are passed over a catalyst of iron at a high temperature (typically 450C) and a very high pressure (typically 200 atmospheres). Some of the hydrogen reacts with nitrogen reacts to form ammonia (nitrogen + hydrogen ==> ammonia). Unfortunately the reaction is reversible i.e. some of the ammonia produced decomposes back into nitrogen and hydrogen. On cooling the reaction mixture, the ammonia liquefies and is removed. To increase efficiency the remaining unreacted hydrogen and nitrogen are recycled. The rules on equilibria can be applied to predict the best conditions to get the highest yield of ammonia but unfortunately there has to be a trade-off between rate of production and position of equilibrium.

  • Ammonia gas is synthesised in the chemical industry by reacting nitrogen gas with hydrogen gas in what is known as the Haber-Bosch Process, named after two highly inventive and subsequently famous chemists.
    • The Haber synthesis of ammonia is important for agriculture because nitrogen is an important element for plant growth.
    • But, it is a very stable molecule and only a few plants like legumes (peas, beans etc.) can directly 'fix' nitrogen the from air and incorporate it into protein molecules.
    • The Haber synthesis allows the efficient mass production of 'artificial fertilisers'.
  • If the chemical feedstocks for the Haber Process are nitrogen and hydrogen, where do we get these materials from?
  • The nitrogen was once obtained from the fractional distillation of liquified air (80% N2).
    • the air is filtered to remove dust and then compressed under high pressure.
    • The filtered air is cooled and water condensed out, and then carbon dioxide freezes out at -78oC.
    • The air further cooled to form liquid air (liquefaction) at around -200oC.
    • The liquid air is fractionally distilled at low temperature to separate oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia), Noble Gases e.g. argon for light bulbs, helium for balloons).
    • Oxygen and argon are very close in boiling point and initially come out in the same fraction so a further fractional distillation is needed to separate them.
    • However, nitrogen is also produced by 'deoxygenating' air by combustion with methane.
  • The hydrogen is made by (i) reacting methane (natural gas) and water or (ii) from cracking hydrocarbons (both reactions are done at high temperature with a catalyst).
    • (i) methane + water (steam) ==> hydrogen + carbon monoxide
      • CH4 + H2O ====> 3H2 + CO
      • The 'deadly' carbon monoxide can be reacted with water in a 2nd stage to make more hydrogen.
        • CO + H2O ====> CO2 + H2
          • and the carbon dioxide is removed to give the desired hydrogen gas.
          • These are called 'reforming' reactions.
    • or (ii) from cracking an alkane hydrocarbon from crude oil e.g.
      • C8H18 ====> C8H16 + H2
    • AND, but just in passing and convenient ....
    • Other uses of hydrogen
      • Hydrogen-oxygen fuel cells to make electricity on small-scale.
      • Hydrogenation vegetable oils to make margarine.
      • Reducing metal oxides to free the metal.
      • An atomic hydrogen-oxygen welding torch.
      • Inflating weather balloons.
  • HABER PROCESS: The Haber process is based on the reversible reaction between nitrogen and hydrogen to form ammonia and using an iron catalyst.
  • The balanced equation for the Haber Synthesis reversible reaction is ...
    • N2(g) + 3H2(g) (c) doc b 2NH3(g)    (plus 92 kJ of heat energy given out, exothermic reaction)

  • .. which means an equilibrium will form, so there is no chance of 100% yield even if you use, as you actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3 !
  • In forming ammonia 92kJ of heat energy is given out (i.e. exothermic, 46kJ of heat released per mole of ammonia formed).
  • Also, four moles of 'reactant' gas form two moles of 'product' gas, so there is a net decrease in gas molecules on forming ammonia.
  • Actual values of % ammonia in an equilibrium mixture with nitrogen and hydrogen are shown in the graph above.
  • Each line of the graph gives you the % of ammonia in the equilibrium mixture as the pressure is gradually increased at a constant temperature (in this case 100oC to 800oC at 100 degree intervals)..
  • These graphs should match your prediction  from applying the equilibrium rules explained in section B.
  • So applying the equilibrium rules, the formation of ammonia should be favoured by  ...
    • (a) Using high pressure because you are going from 4 to 2 gas molecules, so high pressure favours the forward reaction to give fewer gas molecules.
      • The high pressure also speeds up the reaction because it effectively increases the concentration of the gas molecules,
      • but, the higher pressure means more dangerous and more costly engineering, so a compromise needed.
      • Does this prediction match the graph?
      • Note: Increasing the pressure also increases the rate of the reactions, because increasing pressure is effectively increasing the concentration of the reactants.
      • So increasing pressure has two positive effects in terms of yield and getting to the maximum possible % equilibrium of ammonia.
    • (b) Carrying out the reaction at a low temperature, favouring the forward reaction,
      • because it is an exothermic reaction favoured by lowering the temperature,
      • Does this prediction match the graph?
      • However, this may produce too slow a rate of reaction, so a compromise is needed.
    • Therefore, the idea is to use a set of optimum conditions to get the most efficient yield of ammonia and this involves getting a low % yield (e.g. 8% - 15% conversion) but fast.
    • Described below are the conditions to give the most economic production of ammonia.
    • These arguments make the point that the yield of an equilibrium reaction depends on the conditions used.
      • The word 'yield' means how much product you get compared to the theoretical maximum possible if the forward reaction goes 100% that way.
      • For more on chemical economics see Extra Industrial Chemistry page.
  • In industry pressures of 200 - 300 times normal atmospheric pressure are used in line with the theory (200-300 atm).
  • Theoretically a low temperature would give a high yield of ammonia BUT ...
    • Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low temperatures.
    • To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400-450oC).
    • The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!
    • Note: a catalyst does NOT affect the yield of a reaction, i.e. the equilibrium position BUT you do get to the equilibrium position a lot  faster!
  • With reference to the HABER SYNTHESIS chemical plant DIAGRAM
    • Hydrogen and nitrogen gases are mixed in the ratio 3:1 (to fit in with the molecular equation mole ratio) and the gaseous mixture fed into the top of the reaction chamber.
    • The gases are pumped down through the reaction chamber filled with lots of beds ('shelves') coated in the iron catalyst.
    • The hydrogen and nitrogen gases react on the surface of the iron catalyst to form ammonia.
      • N2(g) + 3H2(g) (c) doc b 2NH3(g)

      • The atom economy is 100%.

      • The initial yield is 6%-8%, but unreacted gases are recycled to raise this to nearer 100% eventually ... read on ...

    • At the end of the process, when the gases emerge from the bottom of the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when only the ammonia liquefies and is so can be removed, tapped off from the cooled compression chamber and stored in cylinders for use e.g. making fertilisers.
    • Because the reaction is reversible, not all the nitrogen and hydrogen are converted to ammonia.
    • Any unreacted nitrogen and hydrogen (NOT liquified), is recycled back through the reactor chamber, very little is wasted!
      • Nitrogen (-196oC) and hydrogen (-252oC) have much lower boiling points than ammonia (-33oC) and stay as gases.
      • Boiling points increase with pressure, but these normal atmospheric pressure values offer a fair comparison and the higher the boiling point of the liquid, the higher condensation point of the gas.
      • The temperature in the lower chamber is never low enough to condense out the unreacted hydrogen or nitrogen so only the desired product, ammonia gas condenses out, then the liquid ammonia is drained off at the bottom of the.
      • Since the hydrogen and nitrogen are still gases above the liquid ammonia, they are easily pumped around and mixed with new hydrogen and nitrogen and hence recycled through the reactor.
      • This means non of the original hydrogen and nitrogen reactants is wasted, despite the reaction being an equilibrium.
        • In fact the yield of ammonia can be as little as 6% conversion, but FAST, and the other 94% of reactant gases is recycled FAST.
  • To sum up: A low % yield of ammonia is produced quickly at moderately high temperatures and pressure in the presence of an iron catalyst, and is more economic than getting a higher % equilibrium yield of ammonia at a more costly high pressure and a slower lower temperature reaction.
    • Using an effective iron catalyst can the reduce the cost of manufacturing ammonia by increasing the rate of reaction (more efficient) and lowering the energy requirements if the process can be done at lower temperatures (activation energy reduced).
    • Increasing the rate of reaction saves time and operating at a lower temperature saves energy and therefore saves money.
    • However, catalysts can be very specialised and expensive to produce and they get contaminated ('poisoned') and become less efficient, in this case sulfur compounds contaminate the iron catalyst.
    • So the iron catalyst might have to be extracted and cleaned up, but if a true catalyst (and it is), this' refurbishment' should enable the iron catalyst to be reused.
      • Remember, theoretically catalysts take part in the reaction, but are not consumed in the reaction and can be reused over and over again.
    • See 'chemical economics' for other commercial aspects of chemical production.
  • Detailed notes on "Rates of Reaction" for further reading.
  • AND there are some more general notes on Chemical Economics on the Industrial Chemistry page.


  • Nitrogen fixation is the process of turning atmospheric nitrogen molecules (N2) into useful nitrogenous compounds like ammonia and ammonium salts (fertilisers) and ammonia is used in the chemical industry in the manufacture of explosives, pharmaceutical products and plastics.

    • The Haber process described above is a non-biological way of 'fixing' nitrogen from the air and uses an iron catalyst rather than a biological catalyst like the enzymes involved in the naturally occurring nitrogen cycle.

  • Most industrially produced ammonia is used to make fertilisers, so important to increase crop yields to feed the world's growing population. However the overuse of these synthetic-artificial fertiliser can cause pollution problems.

  • The fixation of nitrogen by e.g. legume plants is quite slow, but using modern chemical technology e.g. high pressure & temperature reactors and an iron catalyst, the whole process is speeded up enormously.

    • The idea of using optimum conditions has been described in detail in section 3a.

  • Ideally it would be great to develop catalyst as powerful as enzymes to 'fix' the nitrogen at room temperature and atmospheric pressure, thereby saving on energy and more costly reactor technology making the overall process cheaper and more efficient.

    • Perhaps genetically modified cells can be developed with more powerful enzymes???

  • Since huge quantities of ammonia are manufactured world-wide, using lots of energy AND oil based hydrocarbons, how sustainable is the Haber process?

    • The hydrogen is made from methane, a non-renewable fossil fuel resource. This is a finite resource, though the nitrogen comes in abundance from the air and would never run out!

    • The atom economy is very good, theoretically it is 100% and in practice the yield is made very high by recycling any unreacted hydrogen and nitrogen gases.

    • There are no waste products because there is only one product possible (no by-products) and even the unreacted gases are recycled.

    • The process is costly in terms of energy needs to maintain a reactor temperature of 200-450oC and 200-400 atmosphere pressure and these process conditions come with a range of health and safety issues to do with high temperatures and pressures - chemical plants are always potentially dangerous unless managed very carefully.

    • The process itself does not damage the environment directly, but misuse/overuse of fertilisers causes problems, but is this not a separate issue from the Haber process itself? unless you get into debate about natural organic fertilisers versus synthetic fertilisers from ammonium salts?

    • Despite, some negative issues mentioned, making ammonia is very profitable because of the huge and increasing world-wide demand for food, especially growing it rapidly with high crop yields.

Keywords & formulae: The Haber Synthesis of ammonia H2 + N2 ==> NH3 N2 + 3H2 ==> 2NH3 reaction conditions temperature pressure iron catalyst recycling unreacted gases

Index: A Reversible Reactions  *  B Reversible reactions and Equilibrium  *  C The Haber Synthesis of ammonia  *  D(a) The Uses of ammonia-nitric acid-fertilisers  *  D(b) Fertilisers-environmental problems  *  E The nitrogen cycle 

(c) doc b Foundation tier (easier) multiple choice QUIZ on ammonia, nitric acid and fertilisers etc.

(c) doc b Higher tier (harder) multiple choice QUIZ on ammonia, nitric acid and fertilisers etc.

Advanced A Level Notes - Equilibrium (use indexes)

Advanced A Level Chemistry Notes p-block nitrogen & ammonia

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