Equilibria
6.1 The Idea of Equilibrium 6.1 the idea of equilibrium
- A dynamic equilibrium is in which both the 6.2 changing the conditions of an equilibrium reaction
forwards and backwards reactions occur at
6.3 equilibrium reactions in industry
the same time.
- They occur in a closed system – no products, 6.4 the equilibrium constant
reactants or energy can get in or out. 6.5 calculations using equilibrium constant expressions
- The amount of products and reactants does
not change. 6.6 the effect of changing conditions on equilibria
- An equilibrium is a state of balanced change.
- Equilibrium can be approached from either direction, and the final equilibrium position will always be the
same.
- When equilibrium is reached density, concentration, colour and pressure do not change with time.
6.2 Changing the Conditions of an Equilibrium
- It is possible to change the proportion of reactants to products in an equilibrium mixture by changing
the position of the equilibrium.
o If the proportion of products is increased, the equilibrium moves to the right.
o If the proportion of reactants is increased, the equilibrium moves to the left.
- Le Chatelier’s Principle:
o If a system at equilibrium is disturbed, the equilibrium moves in the direction that tends to
reduce the disturbance.
- If any factor affecting the equilibrium changes the position of equilibrium will shift to minimise the
change.
- The factors that affect equilibria also affect the rate of a reaction.
Concentration:
- Increasing a concentration will push the equilibrium away from the increased concentration. E.g.,
increasing the concentration of reactants will push towards the products side.
- In terms of rates, increasing the concentration of a reactant increases the rate of the forward
reaction. This makes the forward reaction faster than the backward one. This means the products will
increase until the rates are balanced again.
Pressure:
- This change only affects gases.
- Increasing pressure pushes the equilibrium to the side with the fewest moles of gas (based on the
balanced equation).
o E.g., 2NO2(g) ⇌ N2O4(g)
o Increasing pressure will push the equilibrium towards N2O4(g).
- Increasing the pressure speeds up the rates of both reactions but it speeds up the reaction with the
most moles of gas more.
Temperature:
- Increasing temperature follows the endothermic reaction.
- You will need to be told which direction is endothermic or exothermic. (a negative number means it is an
exothermic reaction)
- If the forward reaction is endothermic then a higher temperature pushes the equilibrium to the
right/products side.
- If the forward reaction is exothermic then a higher temperature pushes the equilibrium to the
left/reactants side.
Catalyst:
- A catalyst does not change the position of an equilibrium.
- It makes both the forwards and backwards reactions faster.
, - By using a catalyst, it means a lower temperature can be used for a reaction, not only does this save
money but if the forward reaction is exothermic then it will increase the yield of the products.
6.3 Equilibrium Reactions in Industry
- Many industrial processes involve reversible reactions.
- Le Chatelier’s principle can be used to choose the best conditions. The conditions that give the highest
yield may not be used. A compromise is often needed balancing out yield, rate of reaction, as well as the
costs of high temperatures and pressures. Safety is also a consideration.
Ammonia:
- 80% of ammonia is used to make fertilisers such as ammonium nitrate, ammonium sulfate and urea.
- 20% is used to make dyes, explosives, plastics, and synthetic fibres like nylon.
The Haber Process:
- Almost all ammonia is made by the Haber Process.
- N2(g) + 3H2(g) ⇌ 2NH3(g) ∆H𝜃 = -92kJmol-1
- A ratio of 1:3 N2 to H2 is used.
- The N2 comes from the air, and H2 comes from natural gas ~ CH4(g) + H2O(g) à CO(g) + 3H2(g)
- The conditions are: 200 atmosphere pressure (20,000kPa), 400oC, Fe catalyst
- Ammonia is cooled and separated out.
- To get the maximum yield of NH3, a low temperature and high pressure are needed.
- High pressure will push the equilibrium to the side with the fewest moles of gas. Low temperatures will
favour the exothermic reaction (forward). At low temperatures the rate is very slow.
- It is better to get a lower yield faster so the temperature of 400oC is used as a compromise. Higher
pressure would be too costly and have safety issues so 200 atm is used as a compromise.
- A catalysts is used to speed up the reaction. This means a lower temperature can be used which will
give a better yield.
- Due to the continuous flow of N2/H2 over the catalyst the reaction doesn’t actually reach equilibrium.
- The ammonia is removed by cooling, this helps to push the equilibrium to the NH3 side.
- Unreacted N2/H2 are then fed back into the reactor.
Ethanol:
- Ethanol is made by the hydration of ethene. The reaction is reversible.
- It is sped up by a catalyst of phosphoric acid absorbed on silica.
- H2C=CH2(g) + H2O(g) ⇌ CH2CH2OH(g) ∆H𝜃 = -46kJmol-1
- High pressure tends to cause the ethene to polymerise, it also increases the cost of building the plant
and the energy to run it. Low temperatures will reduce the rate of reaction and therefore how quickly
equilibrium is reached, although this is compensated for by the catalyst. Too much steam dilutes the
catalyst.
- 570K and 6500kPa are used, the unreacted ethene is separated and recycled over the catalyst again
and again.
Methanol:
- Methanol is used as a chemical feedstock (starting material to make other chemicals). It is also used to
make plastics or as motor fuel.
- CO(g) + 2H2(g) ⇌ CH3OH(g) ∆H𝜃 = -91kJmol-1
- It uses a copper catalyst
- The highest yield will be at low temperature and high pressure, but compromise is used and so 500K and
10,000kPa are actually used in practice.
6.4 The Equilibrium Constant
- For any reaction that reaches equilibrium we can write the equation in the form:
