Understandings:

• Le Châtelier’s principle for changes in concentration can be explained by the equilibrium law.

• The position of equilibrium corresponds to a maximum value of entropy and a minimum in the value of the Gibbs free energy.

• The Gibbs free energy change of a reaction and the equilibrium constant can both be used to measure the position of an equilibrium reaction and are related by the equation, ∆𝐺 = −𝑅T ln𝐾

Applications and skills:

• Solution of homogeneous equilibrium problems using the expression for Kc.

• Relationship between ∆G and the equilibrium constant.

• Calculations using the equation ∆𝐺 = −𝑅T ln𝐾.

Guidance:

• The expression ∆𝐺 = −𝑅T ln𝐾 is given in the data booklet in section 1.

• Students will not be expected to derive the expression ∆𝐺 = −𝑅T ln𝐾.

• The use of quadratic equations will not be assessed.

Where:

R = Gas Constant = 8.314 J^-1 mol^-1 K^-1

T = Temperature in Kelvin (^oC + 273)

∆𝐺 = Standard Free Energy Change NOTE This is not the ∆𝐺 for a given temperature

lnK = Natural log of K - you will also need to reverse this using the exponential function

The Data Booklet has values in KJmol^-1

However Standard Units give this as Jmol^-1 so you will need to convert the data booklet values by x1000

A reaction is at Equilibrium when:

• Kc = 1
• ∆𝐺 = 0

Remember when ∆𝐺 = 0 there is no observable change in Enthalpy or Entropy which are the two key factors of ∆𝐺

• Reactions only proceed when ∆𝐺 < 0
• Larger Kc Values indicate the reactions equilibrium position is far to the right
• Lower Kc values indicate the reactions equilibrium position is far to the left

Therefore - we can summise that reactions are unlikely to be at equilibrium if Kc < 1 and ∆𝐺 > 0