periodic table heading

Chemistry only

ammonia molecule Ammonia is a colourless gas with a choking and distinctive smell which makes it immediately recognizable. It is a simple molecule made up of only nitrogen and hydrogen atoms, its molecular formula is NH3. Ammonia is an excellent base, it is very very soluble in water where it dissolves to form the weak alkali ammonium hydroxide

ammonia(g) + water(l) ammonium hydroxide(aq)
NH3(g) + H2O(l) NH4OH(aq)

Ammonia is easy to liquefy, this is done by compressing gaseous ammonia or by cooling it down to -330C (its boiling point). Ammonia is manufactured on a large scale in the UK by a method called the Haber process. One of the main uses of ammonia is in the manufacture of fertiliser and explosives.

Making ammonia

In industry ammonia is made on a large scale using the process devised by the German chemist Fritz Haber in 1909. Being a small and simple molecule it may appear at first glance that making ammonia (NH3) would be straightforward, simply react nitrogen gas (N2) and hydrogen gas (H2) together to make ammonia (NH3). However there are a number of problems that need to be overcome before large amounts of ammonia can be made.

1. Large amounts of hydrogen gas will be needed. The hydrogen is obtained by reacting methane (CH4) or natural gas with steam using a nickel catalyst. This reaction is often referred to as steam-reforming:

methane + steam carbon monoxide + hydrogen
CH4(g) + H2O(g)CO(g) + 3H2(g)
The carbon monoxide produce can be reacted further with steam to yield more hydrogen.
Carbon monoxide + steamcarbon dioxide + hydrogen
CO(g) + H2O(g) CO2(g) + H2(g)
2. The nitrogen gas needed to make ammonia is obtained from the air. Air is approximately 78% nitrogen. All that is required is to remove the other gases, mainly oxygen, carbon dioxide and water vapour which would be considered impurities and this will leave the nitrogen gas.

3. The equation for the Haber process used to make ammonia is shown below.
Nitrogen(g) + hydrogen(g) ammonia (g)
N2(g) + 3H2(g) 2NH3(g)

nitrogen molecule Perhaps the first point to note from this equation is that the reaction is reversible. Under normal lab conditions nitrogen and hydrogen will react to produce an equilibrium mixture which contains very very little ammonia. The main reason for this is the fact that nitrogen is a particularly unreactive gas. Each nitrogen molecule is held together by a triple covalent bond which requires a large amount of energy to be broken. This triple covalent bond results in a very high activation energy for the reaction and so under lab conditions there is not enough energy to break many of these triple bonds so th reaction is very very slow.

The Haber process uses a catalyst to help lower the activation energy and to break this triple covalent bond and get the reaction going at a reasonable rate. However in order to have any ammonia to sell and make a profit there are a number of other factors you need to think about before setting up a chemical plant to manufacture ammonia.

How temperature affects the yield of ammonia4. As a manager of a chemical plant producing ammonia there are a number of factors you need to think about in order to produce a reasonable amount of ammonia to sell for a profit at the end of the day. Consider the reversible reaction for the formation of ammonia given below:

N2(g) + 3H2(g) 2NH3(g) ΔH= -92KJ mol-1
The forward reaction is exothermic, so according to Le Chatelier's principle lowering the temperature will cause the position of equilibrium to shift to the right and more ammonia will be produced (see graph opposite). The problem is that lowering the temperature will also cause the rate of reaction to slow down. So while lowering the temperature will push the position of equilibrium to the right it may take weeks or even years to happen. Also to work efficiently the catalys used in the Haber process needs a temperature of between 400-5000C.

So a low temperature is not really an option if you want to produce ammonia quickly. However a high temperature is not an option either. According to Le Chatelier's principle, an increase in temperature will force the position of equilibrium to the left, that is reduce the amount of product, so less ammonia.

The answer is to use a compromise temperature. That is a temperature that will produce a reasonable yield of ammonia in a reasonable time. A temperature of around 4500C is usually used in the Haber process as a compromise temperature. The graph opposite shows that as the temperature increases the amount of ammonia present in the equilibrium mixture decreases. However even with the use of a compromise temperature most of the nitrogen and hydrogen are still unreacted and little ammonia is produced. So what else can be done to try and increase the amount of ammonia in the equilibrium mixture Well if you look at the equation for the Haber Process you will see that there are 4 moles of gas on the reactants side and only 2 moles of gas on the product side of the equation:

N2(g) + 3H2(g) 2NH3(g) ΔH= -92KJ mol-1
How pressure affects the yield of ammonia So again using Le Chatelier's principle we can alter the amount of reactants and products in the equilibrium mixture by changing the pressure. If we increase the pressure hen the system (the reacting chemicals) will respond by attempting to lower the pressur, this can be done by forcing the position of equilibrium to the side of the equation with the least number of moles of gas present. That is the product side. So we can say that the higher the pressure the more ammonia will be present in the equilibrium mixture.

However high pressure brings its own problems. It will be very expensive to build pipes, valves and reactors that can withstand and maintain high pressures not to mention the cost of the compressors needed to maintain these high pressure. High pressure also brings an increased risk of explosion.
So as with temperature a balance needs to be struck. There is not point operating a chemical plant at very high pressures due to the expense and danger resulting from these pressures, so a compromise pressur of around 200-250 atmospheres is normally used in the Haber process. This is a balance between economics, safety and need to make a reasonable amount of ammonia.

5. The catalyst.
A catalyst is necessary in the Haber process to speed up the reaction. A catalyst will have no effect on the position of equilibrium or alter the amounts of reactant or product at equilibrium. The catalyst will however speed up the reaction rate considerably. The catalyst used in the Haber process is an activated iron catalyst. It works most efficiently around 4000C.
A simplified outline of the Haber process is shown in the image below: The Haber process

The nitrogen and hydrogen are mixed in the ratio of 3:1, just as in the equation for the Haber process. These gases are then compressed to as high a pressure as is ecomonically viable for the plant, normally around 200 atmospheres. The mixture of gases are then pre-heated before they enter the reactor at around 4500C. On the surface of the iron catalyst the gases react and around 15% yield of ammonia is produced, however about 85% of the nitrogen and hydrogen gases are unreacted.
Once the gases leave the reactor they enter a condenser (cooler). At this point there is a mixture of ammonia and unreacted nitrogen and hydrogen. The ammonia present will easily liquefy under pressure after leaving the cooler. It will then be stored in a separate refrigeration vessel. The unreacted nitrogen and hydrogen gases are then recycled back through the reactor again. With continual recycling of the nitrogen and hydrogen up to 98% of it can be turned into ammonia.

Key points

Practice questions

Check your understanding - Questions on The Haber process