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:

Which of the following conditions indicates that a chemical reaction is spontaneous at constant temperature and pressure?

Which of the following conditions indicates that a chemical reaction is spontaneous at constant temperature and pressure?

Which of the following conditions will result in a spontaneous reaction at standard conditions according to the Gibbs free energy equation ΔG°rxn = ΔH°rxn - TΔS°rxn?

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:

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 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?

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.