Production of Ammonia 2
Ammonia is found in fertilisers as ammonium nitrate. Ammonia is reacted with nitric acid to produce ammonium nitrate.

Production of Ammonia 2

The production of ammonia is one of the important industrial processes that you need to know about for your GCSE Chemistry exam. Ammonia is used to make many other materials such as nitric acid, explosives, nitrogen containing organic chemicals and fertilisers. Without the fertilisers made using ammonia, the world would probably not provide enough food to support the human population. It is used directly in many other products and processes including industrial fermentation and household cleaning products.

Fritz Haber, the man who invented the Haber process for production of ammonia, was a talented chemist. If you study A-level chemistry, you will come across him in several places. He was married to a very accomplished scientist, Clara. Unfortunately at the time, women's achievements were rarely acknowledged and she did not get any credit for the help she gave her husband in his work. Her husband was involved with the development of chemical weapons and is sometimes referred to as the 'father of chemical warfare'. Clara was staunchly opposed to this and committed suicide shortly after the first use of poisonous gas in World War I.

The industrial production of ammonia from nitrogen and hydrogen requires high temperatures, high pressures, an iron catalyst and is a reversible reaction. The nitrogen comes from the air whilst the hydrogen is usually obtained by reacting steam with natural gas. The two gases are continuously cycled and recycled through the reaction vessel. During the recycling, they are cooled to remove the ammonia and fresh nitrogen and hydrogen are added. The process is therefore continuous and you are expected to know the factors that affect the manufacturing costs. Since the equipment runs all day and every day, automating the process is a good way for manufacturers to save money. The recycling of the gases is also a simple way of keeping costs down.

The pressures and temperatures used are always a compromise in order to make the process economically worthwhile and are the key factors that increase the costs. Using higher pressures would increase the amount of ammonia produced, however, the extra cost of building the equipment that could generate and withstand such pressures would not be balanced by the sale of the extra ammonia. Energy costs could be reduced by lowering the temperature but the reaction would slow dow too much and it would not be possible to satisfy the demands of customers. The iron catalyst is useful here as it speeds up the reaction, meaning that temperatures can be kept relatively low.

The conditions that are chosen for the Haber process, 450oC and 200atm, are the best conditions for what?
The highest yield of ammonia
The fastest yield of ammonia
The slowest reverse reaction
The most ammonia as quickly as possible
Other conditions may give a higher yield, but won't be as fast, or will be faster, but give a smaller yield, so a compromise needs to be reached
If a high pressure increases the amount of ammonia produced, why is the process not performed at a much higher pressure than 200 atmospheres?
It is too dangerous
The yield is too high
It is too expensive
Too much hydrogen and nitrogen are used
Very high pressures need reaction vessels and pipes strong enough to cope with the high pressure and these are very expensive. It would also take a lot more energy to produce and maintain the higher pressures which would increase costs even further
If the Haber process was carried out at low temperature, it would increase the amount of ammonia produced. Why is the process NOT carried out at low temperature?
The forward reaction would be too slow
The forward reaction would be too fast
The reverse reaction would give out too much energy
The forward reaction would give out too much energy
Even though the yield of ammonia would be greater, overall, less ammonia would be produced as the reaction would be a lot slower
Ammonia is found in fertilisers as ammonium nitrate. Name the acid that the ammonia is reacted with to produce ammonium nitrate.
Hydrochloric acid
Sulfuric acid
Ethanoic acid
Nitric acid
Ammonia + water ⇌ ammonium hydroxide
Ammonium hydroxide + nitric acid → ammonium nitrate + water
The Haber process produces a smaller volume of gas than is reacted together. What effect will increasing the pressure have on this process?
The amount of ammonia is smaller
The amount of product is increased
The amount of reactant produced is larger
The reverse reaction occurs more quickly
In reversible reactions, increasing the pressure favours the reaction direction that leads to a smaller volume
Here is the equation for the Haber process: N2 + 3H2 ⇌ 2NH3   How many moles of gases are there on each side of the equation?
LHS: 3
RHS: 4
LHS: 4
RHS: 2
LHS: 2
RHS: 4
LHS: 2
RHS: 2
There is 1 mole of nitrogen and 3 moles of hydrogen on the LHS, and 2 moles of ammonia on the RHS
Which of the following statements is true about the role of the iron catalyst in the Haber process?
It speeds up only the forward reaction
It affects the amount of ammonia in the equilibrium mixture
It is used up during the reaction
It increases the rate of reaction in both directions
During reversible reactions, a catalyst has the same effect on both the forward and the reverse reaction. They are used to make sure that the equilibrium point is reached faster than would be possible without the catalyst
What happens to the hydrogen and nitrogen that are not used in the reaction?
They are recycled
They are put into the air
They are used to make other chemicals
They are dissolved in water
These gases are returned to the reaction vessel so that they can be reacted to make more ammonia, this helps to keep manufacturing costs down
Ammonia can also be reacted with sulfuric acid to produce ammonium sulfate. Pick the correct equation for this reaction.
2NH3 + H2SO4 → (NH4SO4)2 + H2O
NH3 + H2SO4 → (NH4SO4)2 + H2O
NH3 + H2SO4 → NH4SO4 + H2O
2NH3 + H2SO4 → (NH4)2SO4
Don't try to remember all of the chemical reactions that you have met during your studies, it is much easier to learn the rules for working them out - they are completely predictable
The forward reaction of the Haber process is exothermic. What does this mean?
It is faster than the reverse reaction
It gives out energy
It takes in energy
It is slower than the reverse reaction
In any reversible reaction, the exothermic direction is favoured by a lower temperature
Author:  Kate Gardiner

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