At what temperature does a reaction become spontaneous at standard conditions if only the Gibbs free energy change (ΔG) is known?

• A. When ΔG is greater than the enthalpy change (ΔH)
• B. When ΔG is zero
• C. When ΔG is negative
• D. When ΔG is positive

Answer: (B)

Verified step by step guidance

1. Understand the concept of Gibbs free energy (ΔG). It is a thermodynamic quantity that can predict the spontaneity of a reaction. A negative ΔG indicates a spontaneous reaction, while a positive ΔG indicates a non-spontaneous reaction.
2. Recall the Gibbs free energy equation: $\Delta G=\Delta H-T\Delta S$, where ΔH is the change in enthalpy, T is the temperature in Kelvin, and ΔS is the change in entropy.
3. To find the temperature at which a reaction becomes spontaneous, set ΔG to zero, as this is the point where the reaction transitions from non-spontaneous to spontaneous.
4. Rearrange the Gibbs free energy equation to solve for temperature: $T=\frac{\Delta H}{\Delta S}$. This equation shows that the temperature at which ΔG becomes zero is dependent on the ratio of ΔH to ΔS.
5. Consider the signs of ΔH and ΔS. If ΔH is positive and ΔS is positive, the reaction becomes spontaneous at high temperatures. If ΔH is negative and ΔS is negative, the reaction becomes spontaneous at low temperatures. If ΔH is negative and ΔS is positive, the reaction is spontaneous at all temperatures.