Three hypothetical gas-phase reactions

CHEMISTRY · Multi-part · A Level · Chemistry · Thermodynamics · Thermodynamics

Problem

Three hypothetical gas-phase reactions, all measured at 298 K298\,\text{K} and 1 bar1\,\text{bar}, have the following standard Gibbs free-energy changes: Reaction I: ΔG∘=−37.8 kJ⋅mol−1\Delta G^\circ = -37.8\,\text{kJ}\cdot\text{mol}^{-1} Reaction II: ΔG∘=+12.4 kJ⋅mol−1\Delta G^\circ = +12.4\,\text{kJ}\cdot\text{mol}^{-1} Reaction III: ΔG∘=−4.60 kJ⋅mol−1\Delta G^\circ = -4.60\,\text{kJ}\cdot\text{mol}^{-1} The universal gas constant is R=8.314 J⋅mol−1⋅K−1R = 8.314\,\text{J}\cdot\text{mol}^{-1}\cdot\text{K}^{-1}. (a) Rank the three reactions from largest to smallest equilibrium constant (KK) at 298 K298\,\text{K}. Justify your ordering qualitatively without performing a numerical calculation. (b) Identify which reaction mixture would contain the greatest proportion of products at equilibrium. Explain your reasoning in terms of thermodynamic favorability. (c) For Reaction II, describe whether the forward or reverse process is spontaneous under standard conditions and relate this to the sign of ΔG∘\Delta G^\circ. (d) Suppose Reaction I is exothermic (ΔH∘<0\Delta H^\circ < 0). Predict qualitatively how its equilibrium constant will change if the temperature is increased to 350 K350\,\text{K}. Support your prediction with thermodynamic reasoning.

Related theory

Browse the problem library