At GCSE we were very careful to say that most chemical reactions were not reversible – they could not go back to the reactants once the products are formed. In the case of the vast majority of chemical reactions this is true, the reaction of methane and oxygen for example:
CH4(g) + O2(g) ® CO2(g) + 2H2O(l)
It is almost impossible to return the carbon dioxide and water to the original methane and oxygen.
Some chemical reactions however will go backwards and forwards depending on the conditions.
CoCl2·6H2O(s) ® CoCl2(s) + 6H2O(l)
If paper soaked in pink hydrated cobalt (II) chloride is heated in an oven, it gives up its water of crystallisation and turns blue. If this paper is exposed to moisture once it has turned blue, it turns back to pink as it regains water of crystallisation.
Another example is the reaction between ammonia and hydrogen chloride gas. As room temperature a white cloud of ammonium chloride is formed.
NH3(g) + HCl(g) ® NH4Cl(s)
If ammonium chloride is heated, then the compound decomposes to give the 2 gases back.
NH4Cl(s) ® NH3(g) + HCl(g)
Important: The chemicals on the left hand side of a chemical reaction are called the reactants and those on the right the products. The left to right reaction the forward reaction and the right to left the backward reaction.
In most reversible reactions balance points exist between the forward and backward reactions.
A simple experiment can be carried out.
I2(in hexane) + I-(aq) « I3-(aq)
It doesn’t matter whether the KI solution or the hexane is added first, the end equilibrium will be the same.
This shows two important factors of dynamic equilibrium
This definition to be learnt like a parrot!
Changing the conditions of an established equilibrium can disturb a system. Any outside influence can shift the balance of the forward and backward reactions in one direction or the other.
Your syllabus is very specific on this, you have to know how changing the following conditions affects the balance of the equilibrium.
Important: A catalyst does not affect the position of equilibrium; it just alters the speed (rate) at which the equilibrium is reached.
A French gentleman by the name of Henri le Chatelier (1850 – 1936) designed a system of rules that allows us to predict qualitatively how a system will change when an outside influence acts on it.
Definition to be learnt like a parrot:
Using a simplified system:
A + B « C + D
So by adding one of the ingredients that go to make up the reaction, the equilibrium will act to shift the equilibrium such that the amount of added material is reduced.
Exactly the same principles apply, just in a slightly more complicated way.
N2(g) + 3H2(g) « 2NH3(g) DH = -92.4 kJ mol-1
Using this example of the Haber Process we can see what will happen if the position of equilibrium changes.
The first step in determining which way the equilibrium will go if pressure is changed is to see how many gaseous moles there are on each side of the equation.
· Products = 2 moles
· Reactants = 4 moles
So if the pressure is increased the equilibrium will shift to reduce the pressure – move to the side with the least number of gaseous moles, to the right to produce more ammonia
If the pressure is decreased the reverse must be true, the equilibrium will move in the direction that increases the number of gaseous moles, towards to reactants (4 moles rather than 2).
This is a little more complex, but not if you follow these rules.
· If ammonia is produced, the reaction will increase in temperature, forward reaction exothermic.
· If nitrogen and hydrogen are produced the reaction will decrease in temperature, the backward reaction is endothermic.
· So before even thinking about what would happen write the following:
o Forward increase in temperature – exothermic
o Backward decrease in temperature – endothermic.
So if asked what will happed if the temperature is increased, we know that the reaction will move to oppose the change.
(a) Increase in temperature, the equilibrium will shift to decrease the temperature, in the endothermic (backwards) direction towards the reactants.
(b) Decrease in temperature, the equilibrium will shift to increase the temperature, in the exothermic (forward) direction towards the products.
So for this particular reaction, High temperature = lots of reactant, Low temperature = lots of product.
If asked what would be the ideal conditions to produce the most amount of product quickly and cost effectively, le Chatelier’s principle states, that we would need:
· Low Temperature
· High Pressure
· Catalyst (all industrial processes have their rate increased using a catalyst, iron is used in this case).
This would give the biggest yield (most amount of product produced), but would it give the quickest reaction?
In reality a compromise is reached between the chemical equilibrium and the rate or kinetics of the process.
At low temperature the rate is very slow (although gives a big yield), so the process typically uses the following conditions:
· 70 to 200 atmospheres
· 400°C to 600°C
This allows for a respectable rate and decent yield. The process is helped as the product is removed as it is produced.
Ammonia is produced on a vast scale, and has four major uses:
1. Fertiliser production (80%)
2. Nitric acid production (5%)
3. Nylon (polyamide) manufacture (7%)
4. Explosives (8%)
Now try the questions