fertiliser production

Chemistry only

Making fertilisers in the lab and industrially

Ammonium sulfate and ammonium nitrate are two common compounds used as fertilisers. They are both easy to make in the school lab. The diagram below outlines how to make ammonium sulfate by carrying out a simple neutralisation reaction: making ammonium sulfate in the lab

You can alter the method above to make ammonium nitrate by simply using nitric acid instead of sulfuric in step 1.

Making fertilisers industrially

The Ostwald process for making nitric acid

Ammonium nitrate is perhaps the most common compound found in fertilisers. It is made by reacting ammonium hydroxide with nitric acid. The equation for the reaction is given below:
ammonium hydroxide(aq) + nitric acid(aq) → ammonium nitrate(aq) + water(l)
NH4OH(aq) + HNO3(aq) → NH4NO3(aq) + H2O(l)
Ammonium hydroxide is simply made by dissolving ammonia in water:
ammonia(g) + water(l) ammonium hydroxide(aq)
NH3(s) + H2O(g) NH4OH(aq)
Obtaining large amounts of ammonium hydroxide will be straight forward. Ammonia is readily obtained from the Haber process. The other reactants, nitric acid is also obtained thanks to the Haber process. To make an acid a non-metal oxide is dissolved in water e.g.
Non-metal oxide + wateracid
Carbon dioxide + watercarbonic acid
Sulfur dioxide + watersulfurous acid
sulfur trioxide + watersulfuric acid
nitrogen dioxide + waternitric acid

The problem with making nitric acid is actually getting the nitrogen dioxide gas. As nitrogen gas is a very unreactive gas so simply burning nitrogen gas in oxygen (often called burning air!) is really not feasible since large amounts of energy are needed.

nitrogen(g) + oxygen(g) nitrogen dioxide (g)

So what is needed is another, easier way of preparing nitrogen dioxide gas. What about burning ammonia? Ammonia burns in oxygen with a yellowish coloured flame. ammonia burning in oxygen gas

However there is a problem, ammonia burns to produce nitrogen gas and water. No nitrogen dioxide is produced, as might have been expected:

ammonia(g) + oxygen(g) → nitrogen(aq) + water(l)
However by altering the conditions above we can obtain nitrogen dioxide gas, the gas needed to make nitric acid. All that is needed is the introduction of a platinum catalys and some heat. The apparatus is shown below:

oxidation of ammonia in presence of a Pt catalyst

In the presence of a platinum catalyst the ammonia is oxidised to give:

ammonia(g) + oxygen(g) → nitrogen monoxide(g) + water(l)
4NH3(s) + 5O2(g) → 4NO(g) + 6H2O(l)
Nitrogen monoxide gas, often called nitric oxide, is a colourless gas that forms inside the combustion tube. However on exposure to air nitrogen monoxide is immediately oxidised to form brown nitrogen dioxide gas.
nitrogen monoxide(g) + oxygen(g) nitrogen dioxide(g)
2NO(g) + O2(g) 2NO2(g)
Nitrogen dioxide is a reddish-brown toxic gas with a bleachy smell. It dissolves in water to form nitric acid:
nitrogen dioxide(g) + water(l) + oxygen(g) nitric acid(aq)
4NO2(g) + 2H2O(l) + O2(g) 4HNO3(aq)
The industrial process for making nitric acid is called the Ostwald process, after Wilhelm Ostwald, a German Nobel prize winning scientist. He developed a process based on the reactions above to manufacture nitric acid. An outline of the Ostwald process is shown below:

Ostwald process oxidation of ammonia

Starting from the left hand-side of the image:

  1. Oxygen from the air is compressed to between 4-10 atmospheres pressure and then pre-heated before it enters the reactor.
  2. Liquid ammonia from the Haber process enters the vapouriser where it is turned into a gas. Next the oxygen and gaseous ammonia enter the reactor. The ammonia is oxidised to nitrogen monoxide gas in a high temperature catalysed reaction. A platinum/rhodium catalyst is used and temperatures are in the range 800-950OC. This reaction is highly exothermic and releases a large amount of heat energy. This heat can be used to generate electricity or used as a heat source to pre-heat gases elsewhere in the reaction.
    ammonia(g) + oxygen(g) → nitrogen monoxide(g) + water(l)
    4NH3(s) + 5O2(g) → 4NO(g) + 6H2O(l)
  3. The nitrogen monoxide leaves the reactor and is cooled in the cooler. Here water is turned into steam as the hot nitrogen monoxide loses heat energy. This cool nitrogen monoxide gas now joins with oxygen to form nitrogen dioxide gas:
    nitrogen monoxide(g) + oxygen(g) nitrogen dioxide(g)
    2NO(g) + O2(g) 2NO2(g)
  4. The final vessel in the image is called the absorption tower. Here a shower of water falls from the top of the tower and meets the nitrogen dioxide gas as it rises up the tower. The nitrogen dioxide gas dissolves in the shower of water to form nitric acid. The nitric acid leaves at the base of the absorption tower and is collected in a large tank. The nitric acid produced can then be reacted with ammonium hydroxide solution, made by dissolving ammonia in water. This neutralisation reaction will form ammonium nitrate:
    ammonium hydroxide(aq) + nitric acid(aq) → ammonium nitrate(aq) + water(l)
    NH4OH (aq) + HNO3(aq) → NH4NO3(aq) + H2O(l)

Fertilisers from phosphate rock

Phosphorus is one of the three main essential elements needed by plants. Phosphate rock is a sedimentary rock that contains high amounts of minerals rich in phosphorus. Limestones and mudstones are two common examples of phosphate rocks. However phosphate rock is insoluble and so cannot be taken in by the plant roots. To be useful as a source of phosphorus for the plant the insoluble phosphate rock needs to be converted into soluble compounds containing phosphate ions (PO43-) which will supply the phosphorus needed by the plant. To produce soluble compounds containing the phosphate ion (PO43-) the phosphate rock is reacted with acids e.g.

Phosphate rock + sulfuric acid → calcium sulfate + calcium phosphate
The mixture of calcium sulfate and calcium phosphate is often called superphosphate or single superphosphate.
Phosphate rock + nitric acid → calcium nitrate + phosphoric acid
Phosphate rock + phosphoric acid→ calcium phosphate (often called triple superphosphate).
The actual reactions taking place here are fairly complex and are unlikely to be asked for in a gcse chemistry paper. As an example consider the industrial preparation of superphosphate (calcium sulfate/calcium phosphate mixture).

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