In a reversible reaction, reactants react together to produce products (forward reaction) and those products also react to produce the original reactants (reverse reaction).
At equilibrium, the rates of both the forward and reverse reactions are the same – the concentrations of all reactants and products do not change.
The position of equilibrium describes how much one reaction is favoured compared to the other.
Changing the conditions of an equilibrium system causes the position of equilibrium to move to oppose this change (Le Chatelier’s Principle).
Increasing the temperature of a system forces the position of equilibrium to move to favor the endothermic reaction.
Increasing the concentration of a reactant forces the position of equilibrium to move to favor the reaction that uses up the reactant.
Increasing the pressure of a system forces the position of equilibrium to move to favor the reaction that produces the fewest moles of gas.
Some reactions are reversible. In a reversible reaction, when reactants are reacting to produce new products, these products are also reacting to produce the reactants of the reaction.
Imagine it as being a bit like a motorway – cars are driving to London from Birmingham. At the same time however, cars are also driving from London to Birmingham.
When these two reactions are happening at the same rate, we say the reactions are at equilibrium. The number of cars heading into London from Birmingham is the same as the number of cars leaving London going into Birmingham.
Because we are using-up reactants at the same rate as we are producing reactants, the concentrations of the reactants and the products do not change. It is very important to remember that the reactions are still happening! They just cancel each other out.
Le Chatelier’s principle:
A French chemist, Le Chatelier, discovered that we can change how much of each reaction is occurring.
He stated that the position of equilibrium will always move to oppose any change to an equilibrium system. This just means that the equilibrium will always want to do the opposite of whatever is being done to the system.
If, for example, the temperature is increased, the position of equilibrium will move to decrease the temperature.
If the pressure in a gaseous system is increased, the position of equilibrium will move to decrease the pressure, and so on.
This principle can be very useful for predicting what will happen in certain conditions and how the positions of equilibrium can be tweaked to produce as much desired product as possible.
Effect of temperature change on equilibrium:
Reactions have an enthalpy change, and they must either be exothermic (release energy in the form of heat) or endothermic (take in energy in the form of heat). With a reversible reaction, there are two reactions that are basically the opposites of themselves. Therefore, if in one direction the reaction is exothermic, in the other it must be endothermic.
Exothermic reactions release energy to the surroundings, so they increase the temperature of the equilibrium system.
Endothermic reactions take in energy from the surroundings, so they decrease the temperature of the equilibrium system.
If the temperature of an equilibrium system is increased, Le Chatelier’s principle tells us that the position of equilibrium will move to oppose that increase in temperature. This means the position of equilibrium will move to favour the endothermic reaction – effectively ‘removing’ the increased temperature, thus lowering the temperature of the equilibrium.
If the temperature of an equilibrium system is decreased, the opposite is true. The position of equilibrium will move to favour the exothermic reaction.
Effect of pressure change on equilibrium:
In a homogenous (all substances are in the same state) gaseous equilibrium, if the pressure is changed, the position of equilibrium will move to oppose this change.
At A-level chemistry, we assume that all gases occupy the same volume at a given temperature and pressure – irrespective of their individual size. This means that 1 mole of methane (CH4) would occupy the same amount of space as 1 mole of butane (C4H10). Even though butane is actually four times the size of methane, the difference in physical size of a molecule is very small compared the volume gases occupy.
Because of this assumption, if the pressure of an equilibrium system is increased, the position of equilibrium will move to produce the least number of gaseous molecules and lower the pressure of the system. If the pressure is increased, the reaction that has the least number of gaseous products will be favoured.
If the pressure of the system is decreased, the opposite is true; the reaction that produces the greatest number of moles of gas will be favoured.
Effect of concentration change on equilibrium:
If the concentration of a reactant (or product) is changed, then how much of that reactant present also changes. Le Chatelier’s principle tells us that the position of equilibrium will move to oppose a change. So, if the concentration of a reactant is increased, the position of equilibrium will move to favour the reaction that uses up that reactant and lower its concentration. If the concentration of a reactant is decreased, the reaction that produces this reactant will be favoured, therefore increasing the concentration and opposing the change.
Effect of a catalyst on equilibrium:
A catalyst does not have any effect on the position of equilibrium. A catalyst just changes the rates of each reaction and both directions are effected in equal amounts.