Chemical Equilibria

Reversible reactions

Reaching an equilibrium state

Dynamic Equilibrium

Factors affecting equilibria

How can we predict how an equilibrium is going to change?

Changing the Concentration

Changing the temperature and pressure.

The industrial importance of ammonia


Reversible reactions

 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)

     pink                              blue

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.

Reaching an equilibrium state.

In most reversible reactions balance points exist between the forward and backward reactions.

  • Reactants and products appear together
  • The reaction appears  to have stopped
  • Neither forward or backward reaction is complete
  • This is chemical equilibrium

A simple experiment can be carried out.

  • Crystal of iodine placed in a test tube.
  • This is then shaken with a mixture of hexane and potassium iodide solution (they are immiscible).
  • The iodine become distributed throughout the mixture and reaches an equilibrium.
  • Iodine is purple in hexane, and a yellowy brown in potassium iodide solution
  • The iodine in the hexane reacts with the iodide in the solution to give tri-iodide ions.

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

  • At equilibrium the concentration of reactants and products does not change
  • The same equilibrium can be reached from either the ‘reactant’ side or the ‘product’ side.

Dynamic Equilibrium

This definition to be learnt like a parrot!

In dynamic equilibrium the forward and backwards reactions continue at equal rates so the overall effect does not change. On a molecular scale there is continuous change. On the macroscopic scale nothing appears to be happening. The system needs to be closed – isolated from the outside world.

Factors affecting equilibria

How can we predict how an equilibrium is going to change?

Changing the Concentration

Changing the temperature and pressure.


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.

How can we predict how an equilibrium is going to change?

Your syllabus is very specific on this, you have to know how changing the following conditions affects the balance of the equilibrium.

  • Concentration
  • Temperature
  • Pressure

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:

The principle states that when the conditions of a system at equilibrium change, the position of equilibrium shifts in the direction that tends to  counteract the change.

Changing the concentration

Using a simplified system:

A + B «  C + D


How does the equilibrium mixture respond?

The result

Concentration of A increases

It moves to the right. Some A is used by reaction with B

More C and D is formed

Concentration of D increases

It moves to the left. Some of the added D is used up by reaction with C

More A and B are formed

Concentration of D decreases

It moves to the right to make up for the loss of D

There is more C and less A and B in the new equilibrium

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.

Changing the temperature and pressure.

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.

1.                  Pressure

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).

2.                  Temperature

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.

The industrial importance of ammonia

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