Chemical equilibria, Le Chatelier’s principle and Kc

Equilibrium

Equilibrium

When rate of forwards reaction is equal to rate of backwards reaction in a closed system

Concentration of reactants and products remain constant

Does a catalyst affect the position of equilibrium

No

Rate

Products / reactants

From point at which reactants are mixed, what must happen to the concentration of reactants

• concentration of reactants must decrease until equilibrium is reached

• Forwards reaction is at a higher rate than reverse reaction

• Reactants are turning into products at a higher rate than products being made into reactants

• At equilibrium concentration stays constant

Hydrogen can be prepared on an industrial scale using the reversible reaction between methane and steam.

CH4(g) + H2O(g) ⇌ CO(g) + 3 H2(g)

ΔH = +206 kJ mol−1

The reaction is done at a temperature of 800 °C and a low pressure of 300 kPa in the presence of a nickel catalyst. Explain, in terms of equilibrium yield and cost, why these conditions are used

Temperature:

• The reaction is endothermic (so equilibrium shifts to RHS to reduce the temperature)

• So, higher temperature increases the yield

• High temperatures are costly (so compromise temperature used)

Pressure:

• More moles of gas on the right hand side, (so equilibrium shifts to RHS to increase the yield)

• So, lower pressure increases the yield AQA

• A low pressure means a low cost

Catalyst :

• Catalyst has no effect on yield.

• Adding a catalyst allows a lower temperature to be used

• So, this lowers the cost

Kp brackets

round

Kc brackets

square

ICE TABLE

initial change equilibrium

Kp units

Pa (could be Pa-1 or kPa etc)

Kc units

mol dm-3