Making Ammonia - The Haber Process

In this reaction nitrogen-N2 (from the air) and hydrogen-H2 (from methane) are reacted together in the ratio of 1:3 to give ammonia-NH3.

The double arrow means that the reaction is reversible.  This has an impact on the yield of product (amount of product) obtained.  In general where a reaction is reversible you will not obtain a high yield without controlling the conditions of the reaction.

In industrial processes yield and atom economy are important considerations.

% yield	=	Mass of product obtained	x 100
		Mass of product expected

The greater the yield the more product a company has to sell and therefore make money.

Atom economy	=	Molecular mass of useful products	x 100
		Molecular mass of all product

Ideally the greater the atom economy the better, as there are no waste atoms (molecules) which have to be dealt with.  If there are waste molecules it could be the case that they are toxic for example hydrogen chloride or contribute to the green house effect for example carbon dioxide  and will need to be treated carefully to remove them.  This will be at the expense of the company.

Yield V's Atom Economy

	Low Yield	High Yield
Low Atom Economy	Undesirable - too much waste not enough product.	Not ideal as the by-products still have to be dealt with.  Producing by-product with some other industrial use would help in this situation.
High Atom Economy	Not ideal as companies need to produce useful product to make profit.  If the reaction is reversible any unreacted product could be fed back into the start of the process to help increase yield.	Most desirable – maximum product with little waste.

With only 1 product formed in the Haber process it has a 100% atom economy.

Conditions for the Haber Process

This reaction is exothermic (gives off energy (heat)) in the forward direction.

The following graph shows the % yield for ammonia production under different conditions of pressure and temperature:

High pressures (reducing the volume the gases move in) force the reaction in the forward direction (making product).  This is because there are fewer gaseous molecules in the products so they will take up less volume.  So high pressures are favoured, the graph shows that 400 atmospheres (pressure 400 times greater than normal atmospheric pressure) produces the greater yield. However it can be expensive to build machinery to produce and maintain high pressures.

The reaction is exothermic in the forward direction and endothermic in the reverse direction.  Increasing the temperature of the reaction will increase the rate of reaction, but as you are increasing energy it will favour the reverse (endothermic) reaction, so the yield will be lower.  This is a difficult condition, to get a high yield temperature must be low however the reaction will be too slow.

A compromise is used:

Temperature: 450ºC

Pressure: 200 atmospheres

With an iron catalyst to speed up the reaction.

The process occurs in the following system:

To try and increase the yield any unreacted hydrogen and nitrogen are recycled to the beginning of the process.

Calculating Reacting Masses

The slide show below shows how you can work out reacting masses in chemical equations.  The important thing to remember is that you need to have a correctly balanced equation (these will usually be given to you in an exam, you should also be given the atomic masses as part of the question, e.g. [Na = 23, O = 16], if not you will have to use the Periodic table at the back of the exam paper).  The clip is only 15 seconds long so pause it at each step and replay as often as you need.