Given the equilibrium concentrations of A and B as 0.0354 M and 0.121 M, respectively, for the reaction A(g) ⇌ B(g), what is the standard Gibbs free energy change (ΔG°) at 25°C?

Gibbs free energy is a measure of the spontaneity of a chemical reaction. It is the chemical potential for a reaction and is minimized at equilibrium. The change in Gibbs free energy can be calculated by ΔG°_rxn = ΔH°_rxn - TΔS°_rxn where ΔG°_rxn is:

Given ΔH°rxn = -95 kJ, ΔS°rxn = 157 J/K, and T = 855 K, calculate the change in Gibbs free energy (ΔG) and predict whether this reaction is spontaneous at the given temperature. Choose from the following options:

Calculate the Gibbs free energy change associated with the formation of 2.0 g of CO2 from the reaction: 2 CO(g) + O2(g) -> 2 CO2(g); given ΔG° = -514.4 kJ for the formation of 2 moles of CO2.

The dissolution of ammonium nitrate is given by the reaction: NH4NO3(s) -> NH4+(aq) + NO3-(aq). Assuming that the values of ΔH° and ΔS° do not change appreciably with temperature, which of the following is the correct expression to calculate the ΔG° value for the reaction?

Given that the equilibrium constant, Keq, for a reaction at 25°C (298 Kelvin) is 740, what is the standard Gibbs free energy change, ΔG°', for the reaction? Provide your answer to two decimal places.

Calculate the value of ΔG° at 281.0 °C for the formation of calcium chloride from its constituent elements: Ca(s) + Cl2(g) → CaCl2(s), given that at 25.0 °C, ΔH° is -795.8 kJ/mol and ΔG° is -748.1 kJ/mol.

Calculate the value of ΔG° at 241.0°C for the formation of calcium chloride from its constituent elements: Ca(s) + Cl2(g) → CaCl2(s). Given that at 25.0°C, ΔH° is -795.8 kJ/mol, ΔG° is -748.1 kJ/mol, and ΔS° is -159.8 J/K.

Given the reaction IF7(g) + I2(g) → IF5(g) + 2 IF(g) with ΔH°rxn = -89 kJ, ΔH°f for IF7(g) = -941 kJ/mol, and ΔH°f for IF5(g) = -840 kJ/mol, calculate the standard enthalpy of formation (ΔH°f) for IF(g).