---
title: "A student investigates the dissolution equilibrium of lithium sulfate in water:  \\[ \\text{Li}_2\\text{SO}_4(s) \\rightleftharpoons 2\\,\\text{Li}^+(aq) + \\text{SO}_4^{2-}(aq) \\]  Thermodynamic data at \\(298\\text{ K}\\) are provided in the table below.  | Substance | \\(\\Delta H_f^\\circ\\text{ (kJ/mol)}\\) | \\(S^\\circ\\text{ (J/(mol}\\cdot\\text{K))}\\) | |—|—|—| | \\(\\text{Li}_2\\text{SO}_4(s)\\) | \\(-1436\\) | \\(114\\) | | \\(\\text{Li}^+(aq)\\) | \\(-278\\) | \\(13\\) | | \\(\\text{SO}_4^{2-}(aq)\\) | \\(-909\\) | \\(20\\) |  Which of the following correctly predicts and explains the effect of increasing temperature on the solubility of \\(\\text{Li}_2\\text{SO}_4(s)\\) in water?"
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date_modified: "2026-08-23T04:23:11+00:00"
---

# A student investigates the dissolution equilibrium of lithium sulfate in water:

\[ \text{Li}_2\text{SO}_4(s) \rightleftharpoons 2\,\text{Li}^+(aq) + \text{SO}_4^{2-}(aq) \]

Thermodynamic data at \(298\text{ K}\) are provided in the table below.

| Substance | \(\Delta H_f^\circ\text{ (kJ/mol)}\) | \(S^\circ\text{ (J/(mol}\cdot\text{K))}\) |
|—|—|—|
| \(\text{Li}_2\text{SO}_4(s)\) | \(-1436\) | \(114\) |
| \(\text{Li}^+(aq)\) | \(-278\) | \(13\) |
| \(\text{SO}_4^{2-}(aq)\) | \(-909\) | \(20\) |

Which of the following correctly predicts and explains the effect of increasing temperature on the solubility of \(\text{Li}_2\text{SO}_4(s)\) in water?

A student investigates the dissolution equilibrium of lithium sulfate in water:

\[ \text{Li}_2\text{SO}_4(s) \rightleftharpoons 2\,\text{Li}^+(aq) + \text{SO}_4^{2-}(aq) \]

Thermodynamic data at \(298\text{ K}\) are provided in the table below.

| Substance | \(\Delta H_f^\circ\text{ (kJ/mol)}\) | \(S^\circ\text{ (J/(mol}\cdot\text{K))}\) |
|---|---|---|
| \(\text{Li}_2\text{SO}_4(s)\) | \(-1436\) | \(114\) |
| \(\text{Li}^+(aq)\) | \(-278\) | \(13\) |
| \(\text{SO}_4^{2-}(aq)\) | \(-909\) | \(20\) |

Which of the following correctly predicts and explains the effect of increasing temperature on the solubility of \(\text{Li}_2\text{SO}_4(s)\) in water?

- **A.** The solubility will increase because \(\Delta S^\circ_{\text{dissolution}} > 0\), making \(\Delta G^\circ_{\text{dissolution}}\) more negative at higher temperatures.
- **B.** The solubility will decrease because \(\Delta S^\circ_{\text{dissolution}} < 0\), making \(\Delta G^\circ_{\text{dissolution}}\) more positive at higher temperatures.
- **C.** The solubility will increase because \(\Delta H^\circ_{\text{dissolution}} < 0\), providing additional thermal energy to overcome ion-ion interactions at higher temperatures.
- **D.** The solubility will decrease because \(\Delta H^\circ_{\text{dissolution}} > 0\), requiring the absorption of heat that is unfavorable at higher temperatures.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/120333/*