The sign of ΔG⦵, the standard Gibbs free energy change, has a direct influence on the value of the equilibrium constant, K, for a chemical reaction. When ΔG⦵ is negative, indicating a spontaneous reaction, K is greater than 1, meaning that the products are favoured at equilibrium. If ΔG⦵ is positive, indicating a non-spontaneous reaction, K is less than 1, meaning that the reactants are favoured. When ΔG⦵ is zero, the system is at equilibrium and K is equal to 1. The relationship between ΔG⦵ and K reflects the tendency of a reaction to proceed towards equilibrium, with a lower Gibbs free energy.

A change in pressure can affect the standard Gibbs free energy change, ΔG⦵, for a reaction involving gases due to the impact on enthalpy (ΔH⦵) and entropy (ΔS⦵) changes. For reactions that result in a change in the number of moles of gas, an increase in pressure will favour the formation of fewer moles of gas, potentially altering the values of ΔH⦵ and ΔS⦵, and consequently ΔG⦵. However, under standard conditions, the pressure is fixed at 1 bar, and ΔG⦵ is determined at this constant pressure. Therefore, ΔG⦵ values typically provided in tables are not affected by changes in pressure.

Yes, it is possible for a reaction to have a positive entropy change, ΔS⦵, and still be non-spontaneous. The spontaneity of a reaction depends on the Gibbs free energy change, ΔG⦵, not just on ΔS⦵. Even if ΔS⦵ is positive, contributing to a decrease in ΔG⦵, if the enthalpy change, ΔH⦵, is sufficiently positive or the temperature, T, is low enough, the TΔS⦵ term might not be large enough to offset ΔH⦵, resulting in a positive ΔG⦵. In such cases, the reaction would be non-spontaneous under the given conditions.

A reaction can have a negative ΔG⦵ at all temperatures if its enthalpy change, ΔH⦵, is negative (exothermic) and its entropy change, ΔS⦵, is positive, resulting in a spontaneous reaction under all conditions. In this case, both terms in the Gibbs free energy equation ΔG⦵ = ΔH⦵ - TΔS⦵ contribute to a decrease in ΔG⦵, making it negative. The negative ΔH⦵ lowers ΔG⦵, and the positive TΔS⦵ further lowers ΔG⦵. As temperature increases, the TΔS⦵ term becomes more significant, maintaining the spontaneity of the reaction.

Yes, a reaction can be spontaneous even if both ΔH⦵ (enthalpy change) and ΔS⦵ (entropy change) are negative. Spontaneity is determined by the Gibbs free energy change, ΔG⦵, and not just by enthalpy or entropy changes alone. According to the equation ΔG⦵ = ΔH⦵ - TΔS⦵, a negative ΔH⦵ contributes to a decrease in ΔG⦵, making the reaction more likely to be spontaneous. However, a negative ΔS⦵ tends to increase ΔG⦵, working against spontaneity. Therefore, whether the reaction is spontaneous depends on the magnitude of ΔH⦵ and ΔS⦵ and the temperature of the system. If the magnitude of TΔS⦵ is less than ΔH⦵ and ΔH⦵ is sufficiently negative, the reaction can still be spontaneous.