Chemical Equilibrium
Dynamic equilibrium and Le Chatelier’s principle
In a closed system — where reactants or products can’t escape — reversible reactions can reach dynamic equilibrium.
Dynamic equilibrium describes how reactants are being formed just as fast as the products are being formed, so the concentrations of each substance stays the same (be careful: this doesn’t mean that the concentrations of reactants and products are equal).
At equilibrium:
The forward and reverse reactions are still happening
The rate of the forward and reverse reactions is the same
The concentration of reactants and products remains constant
When a reaction has reached dynamic equilibrium and a change occurs — such as a change in concentration, pressure or temperature — the position of equilibrium will move to counteract the change. This concept is known as Le Chatelier’s principle.
Effect of concentration, pressure and temperature on equilibrium position
Concentration
If the concentration of a reactant is increased, the position of equilibrium shifts to the right to favour the formation of products.
Take a sealed container containing decomposing ammonium chloride. The reaction has reached dynamic equilibrium but then I add more ammonium chloride. The position of equilibrium will move to the right, to make more products and balance things out.
The same is true for the other way round. If I added more product to the container the position of equilibrium would shift to the left to favour the formation of reactants.
Pressure
If the pressure is increased, the position of equilibrium shifts to the side with the fewest moles of gas.
For example, the equation below shows how ammonia is synthesised from nitrogen and hydrogen. There are four moles of gas on the left hand side and only two on the right. If we increase the pressure, the position of equilibrium will shift to the right hand side (towards ammonia) as this is the side with the fewest moles of gas.
Alternatively, if we decrease the pressure, the position of equilibrium would shift to the left hand side where there are more moles.
Temperature
If we increase the temperature of a reaction at equilibrium, the position of equilibrium will shift in the endothermic direction to lower the temperature.
Hydrogen can be manufactured by reacting carbon with steam, as shown in the equation below. The forward reaction is endothermic and the reverse reaction is endothermic. So increases in temperature will shift the position of equilibrium in the endothermic forward direction (towards the products). And if we decrease the temperature, the position of equilibrium will shift in the exothermic reverse direction (towards the reactants).
Catalyst
A catalyst does NOT change the position of equilibrium. This is because it speeds up the forward and reverse reactions by the same amount.
Compromises in reaction conditions
Many industrial processes involve reversible reactions — so manufacturers will think about the temperature, pressure and concentration that maximises yield. But a compromise needs to be reached with reaction rate, and other considerations (such as cost and safety).
For example, low temperatures would maximise the yield of products for exothermic reactions but the products would be formed too slowly. A moderate temperature would therefore be used to get a satisfactory yield of product at a reasonable rate.
Moderate conditions may be used for industrial processes because:
Low pressures result in a slow rate of reaction.
High pressures are expensive to generate, require specialist equipment and may result in unwanted side reactions.
Low temperatures result in a slow rate of reaction.
High temperatures are expensive and involve burning lots of fuel, which releases carbon dioxide and therefore not very environmentally-friendly).
Equilibrium constant, Kc
The position of equilibrium can be given a numerical value – this is the equilibrium constant or Kc. The larger the value of Kc, the more the position of equilibrium lies to the right and the more the forward reaction is being favoured.
Likewise, the smaller the value of Kc, the more the backwards reaction is happening and the position of equilibrium lies further to the left.
For a homogenous reaction — where the reactants and products are in the same state — in dynamic equilibrium, we can calculate the value of Kc using the following equation:
Worked example – calculating Kc
In the Haber process, nitrogen gas is reacted with hydrogen gas to form ammonia to be used for fertiliser. The equation is:
The reaction was left to reach dynamic equilibrium. Calculate the value of Kc for when the following concentrations are present: 0.11 mol dm-3 of N2, 0.33 mol dm-3 of H2 and 0.22 mol dm-3 of NH3.
We first need to write out our Kc expression. Then put the products on the top and divide by the reactant concentrations. Everything should be to the power of their coefficients.
Now we can just pop our concentrations into the equation.
The units can change depending on the reaction we’re dealing with because the number of reactants / products in our Kc expression will vary.
To work out the units, we can put them into our equation and cancel out anything that is the same on the top and the bottom.
Cancel out and bring anything on the bottom of the expression over to the top.
