For the reaction H₂(g)+ S(s)$\rarr$ H₂S(g), $\Delta$H° = -20.2 kJ/mol and $\Delta$S° = +43.1 J/K·mol.Which of these statements is true? A)The reaction is only spontaneous at low temperatures. B)The reaction is spontaneous at all temperatures. C)($\Delta$G° becomes less favorable as temperature increases.) D)The reaction is spontaneous only at high temperatures. E)The reaction is at equilibrium at 25°C under standard conditions.

Question

Quizplus · Accepted Answer

The negative value of ΔH° indicates that the reaction is exothermic and the positive value of ΔS° indicates an increase in entropy. When ΔH° is negative and ΔS° is positive, the reaction is spontaneous at all temperatures according to the Gibbs free energy equation ΔG° = ΔH° - TΔS°. Therefore, option B is correct.

For the Reaction H2(g)+ S(s) →\rarr→ H2S(g) Δ\DeltaΔ H° = -202 KJ/mol And H2(g)

For the Reaction H₂(g)+ S(s) $\rarr$ H₂S(g) $\Delta$ H° = -202 KJ/mol And H₂(g)