---
title: "The decomposition of sulfuryl chloride gas is represented by the following equation:  \\[\\text{SO}_2\\text{Cl}_2(g) \\rightleftharpoons \\text{SO}_2(g) + \\text{Cl}_2(g)\\]  For this reaction, \\(\\Delta H^\\circ = +90\\text{ kJ/mol}_{\\text{rxn}}\\) and \\(\\Delta S^\\circ = +150\\text{ J/(mol}_{\\text{rxn}}\\cdot\\text{K)}\\). Which of the following correctly predicts the sign of \\(\\Delta G^\\circ\\) and the relative magnitude of the equilibrium constant, \\(K\\), as the temperature of the reaction system is changed from \\(500\\text{ K}\\) to \\(700\\text{ K}\\)?"
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url: "https://nerd-notes.com/ubq/121480/"
date_modified: "2026-08-23T05:04:47+00:00"
---

# The decomposition of sulfuryl chloride gas is represented by the following equation:

\[\text{SO}_2\text{Cl}_2(g) \rightleftharpoons \text{SO}_2(g) + \text{Cl}_2(g)\]

For this reaction, \(\Delta H^\circ = +90\text{ kJ/mol}_{\text{rxn}}\) and \(\Delta S^\circ = +150\text{ J/(mol}_{\text{rxn}}\cdot\text{K)}\). Which of the following correctly predicts the sign of \(\Delta G^\circ\) and the relative magnitude of the equilibrium constant, \(K\), as the temperature of the reaction system is changed from \(500\text{ K}\) to \(700\text{ K}\)?

The decomposition of sulfuryl chloride gas is represented by the following equation:

\[\text{SO}_2\text{Cl}_2(g) \rightleftharpoons \text{SO}_2(g) + \text{Cl}_2(g)\]

For this reaction, \(\Delta H^\circ = +90\text{ kJ/mol}_{\text{rxn}}\) and \(\Delta S^\circ = +150\text{ J/(mol}_{\text{rxn}}\cdot\text{K)}\). Which of the following correctly predicts the sign of \(\Delta G^\circ\) and the relative magnitude of the equilibrium constant, \(K\), as the temperature of the reaction system is changed from \(500\text{ K}\) to \(700\text{ K}\)?

- **A.** At \(500\text{ K}\), \(\Delta G^\circ < 0\) and \(K > 1\); at \(700\text{ K}\), \(\Delta G^\circ < 0\) and \(K > 1\), because the entropy term outweighs the enthalpy change at both temperatures.
- **B.** At \(500\text{ K}\), \(\Delta G^\circ > 0\) and \(K < 1\); at \(700\text{ K}\), \(\Delta G^\circ > 0\) and \(K < 1\), because the positive value of \(\Delta H^\circ\) prevents the reaction from being thermodynamically favorable at either temperature.
- **C.** At \(500\text{ K}\), \(\Delta G^\circ > 0\) and \(K < 1\); at \(700\text{ K}\), \(\Delta G^\circ < 0\) and \(K > 1\), because the magnitude of \(-T\Delta S^\circ\) is smaller than \(\Delta H^\circ\) at \(500\text{ K}\) but exceeds \(\Delta H^\circ\) at \(700\text{ K}\).
- **D.** At \(500\text{ K}\), \(\Delta G^\circ < 0\) and \(K > 1\); at \(700\text{ K}\), \(\Delta G^\circ > 0\) and \(K < 1\), because increasing the temperature makes \(-T\Delta S^\circ\) more positive.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/121480/*