Gibbs free-energy change

• If ΔG is negative, the reaction is spontaneous and can occur without the input of energy.
• The formula for Gibbs free energy change is:

  • ΔG = ΔH − TΔS

  where ΔG is Gibbs free energy, ΔH is enthalpy change, T is temperature in Kelvin, and ΔS is entropy change.

  • ΔG and ΔH are measured in kJ mol⁻¹.
  • ΔS is measured in J K⁻¹ mol⁻¹ (so it must be converted to kJ by dividing by 1000 when used in the Gibbs free energy equation).
• If ΔG is positive, the reaction is non-spontaneous and requires energy input to occur.
• If ΔG = 0, the system is at equilibrium, meaning no net change occurs, and the forward and reverse reactions happen at the same rate.
• Temperature (T) directly affects ΔG. As temperature increases, the term TΔS becomes larger, which can make a reaction more likely to be spontaneous if entropy (ΔS) is positive.
• Entropy (ΔS) is the measure of disorder in a system. If ΔS is positive, the system becomes more disordered, which can lower ΔG and make the reaction more spontaneous, especially at higher temperatures.
• How does enthalpy (ΔH) influence Gibbs free energy?

  • Enthalpy (ΔH) reflects the heat absorbed or released during a reaction.
  • If ΔH is negative (exothermic), it tends to make ΔG more negative, increasing the likelihood of the reaction being spontaneous.
• What conditions favour a reaction being spontaneous at all temperatures?

  A reaction is spontaneous at all temperatures if ΔH is negative (exothermic) and ΔS is positive (increase in disorder), making ΔG negative at any temperature.
• Under what conditions will a reaction be non-spontaneous at all temperatures?

  A reaction will be non-spontaneous at all temperatures if ΔH is positive (endothermic) and ΔS is negative (decrease in disorder), making ΔG positive at any temperature.
• How can a reaction be spontaneous at high temperatures but not at low temperatures?

  If ΔH is positive (endothermic) and ΔS is positive (increase in disorder), the reaction may only be spontaneous at high temperatures, where the TΔS term can outweigh the positive ΔH.
• How can a reaction be spontaneous at low temperatures but not at high temperatures?

  If ΔH is negative (exothermic) and ΔS is negative (decrease in disorder), the reaction may only be spontaneous at low temperatures, where the TΔS term is small enough that ΔH dominates and makes ΔG negative.
• How do you calculate the temperature at which a reaction becomes feasible (ΔG = 0)?

  Set ΔG = 0 in the Gibbs equation:

  • ΔG = ΔH − TΔS

  Rearrange to solve for temperature:

  • T = ΔH/ΔST
• The standard Gibbs free energy change is the Gibbs free energy change when reactants and products are in their standard states (1 atm pressure, 298 K, 1 M concentration for solutions).
• What is the relationship between ΔG and reaction spontaneity?

  • If ΔG < 0, the reaction is spontaneous.
  • If ΔG > 0, the reaction is non-spontaneous.
  • If ΔG = 0, the system is at equilibrium.
• ΔS = sum of ΔS of products - the sum of ΔS of reactants