6.3.1 The Haber process as a reversible reaction

Ammonia is formed from nitrogen and hydrogen

1. The Haber process brings nitrogen and hydrogen together to form ammonia.
2. The equation is written with the reversible arrow: $$N_2(g)+3H_2(g)\rightleftharpoons 2{NH}_3(g)$$
3. One molecule of nitrogen combines with three of hydrogen to give two of ammonia.
4. Both reactants are gases, and so is the ammonia formed.
5. Ammonia is the starting point for fertiliser manufacture, which is why the process matters.

The reaction runs in both directions at once

1. Nitrogen and hydrogen combine, and at the same time ammonia breaks back down into them.
2. The forward reaction forms ammonia and the backward reaction destroys it.
3. Once some ammonia has formed, both run simultaneously under the same conditions.
4. As ammonia builds up, the backward reaction speeds up.
5. As nitrogen and hydrogen are used up, the forward reaction slows down.

The mixture reaches a dynamic equilibrium

1. The two rates eventually become equal, and the amounts then stop changing.
2. Both reactions are still running at that point, which is what dynamic means.
3. The mixture at equilibrium holds ammonia together with unreacted nitrogen and hydrogen.
4. The reactor is sealed while the gases are inside it, which is what allows equilibrium to be reached.
5. Reaching equilibrium sooner does not change what the equilibrium mixture holds.

No single pass converts all of the reactants

1. Because the backward reaction is always running, some nitrogen and hydrogen remain.
2. The proportion converted in one pass is therefore well below one hundred percent.
3. Ammonia is separated by cooling the mixture until it condenses to a liquid.
4. The nitrogen and hydrogen stay as gases and are recycled into the reactor.
5. Recycling means the raw materials are eventually used up even though each pass converts only part.

A reversible reaction is still worth running

1. A low conversion per pass would waste raw material if the gases were simply discarded.
2. Recycling removes that objection, so the overall use of raw material is efficient.
3. The nitrogen is drawn from the air, which costs nothing but the energy to separate it.
4. Running the reaction continuously rather than in batches keeps the plant working at full output.
5. The scale of world demand for fertiliser is what makes the process worth the energy it uses.

6.3.2 Conditions and rate in industrial equilibria

Rate of attainment is how quickly equilibrium is reached

1. The rate of attainment of equilibrium is how long the mixture takes to settle.
2. It is a separate question from the position of the equilibrium, which is what the mixture settles at.
3. A change can affect the rate, the position, or both.
4. Anything that makes particles collide more often, or more energetically, makes equilibrium arrive sooner.
5. An industrial process needs an acceptable yield in an acceptable time, so both questions are asked together.

Temperature and pressure act on rate and on position

1. Raising the temperature gives particles more energy, so equilibrium is reached sooner.
2. Raising the temperature also shifts the position towards the endothermic direction.
3. For the Haber process the forward reaction is exothermic, so a higher temperature lowers the yield.
4. Raising the pressure pushes gas particles closer together, so collisions are more frequent and equilibrium comes sooner.
5. Raising the pressure also shifts the position towards the side with fewer gas molecules, which is ammonia.

Concentration and catalysts act on the rate

1. Raising the concentration of a reactant makes collisions more frequent, so equilibrium is reached sooner.
2. A catalyst provides a different route with a lower activation energy, which speeds the reaction up.
3. It speeds the forward and backward reactions by the same factor.
4. Because both rates rise together, the composition at which they become equal is unchanged.
5. A catalyst therefore brings equilibrium sooner without altering the yield at all.

The Haber conditions are a compromise

1. A low temperature would give the best yield, because the forward reaction is exothermic.
2. At a low temperature the reaction would be far too slow to be worth running.
3. $450{}^{\circ }C$ is chosen as a compromise, accepting a lower yield for a workable rate.
4. A high pressure helps both the yield and the rate, so pressure is pushed as far as is practical.
5. $200$ atmospheres is chosen because higher pressures need stronger, costlier plant and are more dangerous.

Cost of raw materials and energy shapes the choice

1. Maintaining a high temperature and a high pressure both consume energy that has to be paid for.
2. Stronger vessels and compressors raise the capital cost of building the plant.
3. Nitrogen from the air is effectively free, while hydrogen from natural gas carries a real cost.
4. Recycling the unreacted gases means a modest conversion per pass is acceptable.
5. The conditions chosen are the ones that make the product at the lowest overall cost, not the ones that give the highest yield.

6.3.3 Fertilisers and the preparation of ammonium sulfate

Fertilisers supply the elements plants need

1. Growing crops take nitrogen, phosphorus and potassium out of the soil.
2. A fertiliser returns those elements as soluble compounds that roots can absorb.
3. Nitrogen compounds support leaf growth, so they are the largest part of most fertilisers.
4. Compounds of phosphorus and potassium are supplied alongside them.
5. A compound only works as a fertiliser if it dissolves, because roots take up solutions.

Ammonia reacts with nitric acid to give a fertiliser salt

1. Ammonia solution is an alkali, so it reacts with an acid to form a salt.
2. Ammonia and nitric acid give ammonium nitrate: $${NH}_3+{HNO}_3\to {NH}_4{NO}_3$$
3. Ammonium nitrate is especially valuable because both of its ions supply nitrogen.
4. Ammonia is unusual among bases in giving a salt and no water.
5. Being soluble, it dissolves in soil water and reaches the roots of the crop.

Preparing ammonium sulfate in the laboratory

1. A measured volume of ammonia solution is placed in a flask with a few drops of indicator.
2. Dilute sulfuric acid is added from a burette until the indicator just changes colour.
3. The reaction taking place is: $$2{NH}_3+H_2{SO}_4\to ({NH}_4{)}_2{SO}_4$$
4. Repeating the run with the same volumes and no indicator keeps the product uncontaminated.
5. The solution is evaporated to the point of crystallisation and left to form crystals, which are then dried.

Producing ammonium sulfate industrially

1. Industry makes the same salt from ammonia and sulfuric acid, but on a far larger scale.
2. The ammonia has first to be manufactured from nitrogen and hydrogen by the Haber process.
3. The sulfuric acid has to be manufactured from its own raw materials in several stages.
4. The neutralisation is therefore only the last step of a much longer sequence.
5. The process runs continuously, with the reactants flowing in and the product flowing out.

Comparing the two scales

1. The chemistry is the same in both cases, and so is the equation.
2. The laboratory route uses bought-in ammonia and acid, while industry manufactures both.
3. A laboratory preparation is a batch, made once and then worked up.
4. Industry runs continuously, because starting and stopping a large plant wastes energy.
5. Purification differs too: crystals are filtered and dried by hand, while industry uses large scale separation.