The Haber Process
The Haber process is the industrial manufacture of Ammonia from Nitrogen and Hydrogen. The raw materials are:
* air - for nitrogen,
* methane and water - for hydrogen.
The reaction is an exothermic reaction that is reversible and the equation is as follows:
nitrogen + hydrogen ammonia ( + heat)
N2(g) + 3H2(g) 2NH3(g) ( + heat)
We obtain our raw material as follows:
* Nitrogen is obtained from the 79% of air that is nitrogen by fractional distillation.
* Hydrogen is obtained from methane (natural gas) or naphtha. The hydrocarbon is reacted with steam.
Methane + steam carbon dioxide + hydrogen.
CH4(g) + 2H2O(g) CO2(g) + 4H2(g)
The industrial conditions that are preferred and you suggested are:
) A temperature between 450 °C and 500 °C. Heat is treated in this reaction as a product (an exothermic reaction). If you remove heat from an exothermic reaction (cool it down), the equilibrium will shift to produce more product. This will not only produce more heat, but also produce more of the chemical product that you want in the equilibrium mixture.
2) A Pressure of 200 atm (200 atmospheres). For a reversible reaction involving gases, increasing the pressure will shift the equilibrium towards the side of the reaction, which has the smaller volume and produce more to reach equilibrium.
If we refer to the chart and graph we can see that only about 15% of the reactants are converted into products under these conditions. What happens is that Ammonia is cooled and liquefied at the reaction pressure, and then removed as liquid ammonia.
The remaining mix of nitrogen and hydrogen gases (85%) are recycled and fed ...
This is a preview of the whole essay
2) A Pressure of 200 atm (200 atmospheres). For a reversible reaction involving gases, increasing the pressure will shift the equilibrium towards the side of the reaction, which has the smaller volume and produce more to reach equilibrium.
If we refer to the chart and graph we can see that only about 15% of the reactants are converted into products under these conditions. What happens is that Ammonia is cooled and liquefied at the reaction pressure, and then removed as liquid ammonia.
The remaining mix of nitrogen and hydrogen gases (85%) are recycled and fed in at the reactant stage and this process operates continuously. The chart shows, as indicated by the highlighted section, the factors at which we obtain the best yield however it is not possible to work under these conditions due to the amount of product that would be made in relation to cost.
Temperature (°C)
Pressure (Atmospheres)
Yield (mole per cent)
350
0
0
350
00
37
350
200
52
350
300
60
350
400
65
450
0
0
450
00
7
450
200
27
450
300
35
450
400
42
550
0
0
550
00
7
550
200
3
550
300
8
550
400
22
When a reaction is reversible such as this reaction is, it means that it can go either forwards or backwards. The forward reaction is the one we want, as reactants are converted into products. Reactions in both directions occur at the same time and when dynamic equilibrium has been reached, it does not mean that the reactions have stopped, it simply means that the forward reaction is making products in the same quantity that the backward reaction is making reactants. The forward reaction (to form ammonia) is exothermic and since we want ammonia from the Haber Process we need to look at rates of reaction to find the right temperature.
All reactions go faster if the temperature is raised and in a reversible reaction like the Haber Process, raising the temperature will make the equilibrium mixture richer in nitrogen and hydrogen because forming these from ammonia takes heat in. If we cool the reaction down, the proportion of ammonia in the equilibrium mixture will increase, but the rate at which ammonia is formed will decrease as the temperature is lower and it is no good having 90% ammonia if it takes all day to make one bucket full. However it is far better to have 10% ammonia being made very quickly, and at the end of the day you can have thousands of litres. The actual temperature of between 450 °C and 500 °C, is a compromise between the proportion of ammonia in the equilibrium mixture (only 15% because of the high temperature) and the rate at which ammonia is formed (fast because of the high temperature).
If we look at the reaction, the reactants and products are gases. One mole of any gas occupies a volume of 24,000 cm3. On the left side of the equation, there is one mole of nitrogen, and three moles of hydrogen. The total is four moles of reactant. On the right side of the equation (the product), there are two moles of ammonia. So, four moles of reactant give two moles of product. Since one mole of any gas takes up the same volume, the volume of product is only half the volume of reactants. The reason for the difference in moles is that it is in proportion of reactant to product. Increasing the pressure makes the equilibrium mixture richer in ammonia and this is what we want! As we can see, from my graph increased pressure also increases the reaction rate and again, this is what we want!
You may ask why not increase the pressure to 500 atm, and get lots of ammonia really quickly? However in the real world, it all comes down to money because building a high-pressure chemical plant is expensive and running the reaction at about 200 atm gives the highest return on investment capital. The important thing to remember is that any industrial process works in a way which gives the maximum product for the minimum cost. The cost is not entirely dependent on the chemistry of the process, but also includes health and safety, energy, transport and the environment.
