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title: "A student investigates the thermodynamics of gas-phase hydrogenation reactions. The student compares the reaction of ethane with hydrogen to the reaction of ethyne with hydrogen, as represented by the following balanced equations:  \\[ \\text{Reaction 1: } \\text{C}_2\\text{H}_6\\text{(g)} + \\text{H}_2\\text{(g)} \\rightarrow 2\\,\\text{CH}_4\\text{(g)} \\quad \\Delta H^\\circ_1 = -42\\text{ kJ/mol}_{\\text{rxn}} \\] \\[ \\text{Reaction 2: } \\text{C}_2\\text{H}_2\\text{(g)} + \\text{H}_2\\text{(g)} \\rightarrow \\text{C}_2\\text{H}_4\\text{(g)} \\quad \\Delta H^\\circ_2 \\]  Both reactions break \\(1\\text{ mol}\\) of \\(\\text{H}-\\text{H}\\) bonds and form \\(2\\text{ mol}\\) of \\(\\text{C}-\\text{H}\\) bonds per mole of reaction. Which of the following best predicts and explains the value of \\(\\Delta H^\\circ_2\\) relative to \\(\\Delta H^\\circ_1\\)?"
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url: "https://nerd-notes.com/ubq/123845/"
date_modified: "2026-09-28T12:30:31+00:00"
---

# A student investigates the thermodynamics of gas-phase hydrogenation reactions. The student compares the reaction of ethane with hydrogen to the reaction of ethyne with hydrogen, as represented by the following balanced equations:

\[ \text{Reaction 1: } \text{C}_2\text{H}_6\text{(g)} + \text{H}_2\text{(g)} \rightarrow 2\,\text{CH}_4\text{(g)} \quad \Delta H^\circ_1 = -42\text{ kJ/mol}_{\text{rxn}} \]
\[ \text{Reaction 2: } \text{C}_2\text{H}_2\text{(g)} + \text{H}_2\text{(g)} \rightarrow \text{C}_2\text{H}_4\text{(g)} \quad \Delta H^\circ_2 \]

Both reactions break \(1\text{ mol}\) of \(\text{H}-\text{H}\) bonds and form \(2\text{ mol}\) of \(\text{C}-\text{H}\) bonds per mole of reaction. Which of the following best predicts and explains the value of \(\Delta H^\circ_2\) relative to \(\Delta H^\circ_1\)?

A student investigates the thermodynamics of gas-phase hydrogenation reactions. The student compares the reaction of ethane with hydrogen to the reaction of ethyne with hydrogen, as represented by the following balanced equations:

\[ \text{Reaction 1: } \text{C}_2\text{H}_6\text{(g)} + \text{H}_2\text{(g)} \rightarrow 2\,\text{CH}_4\text{(g)} \quad \Delta H^\circ_1 = -42\text{ kJ/mol}_{\text{rxn}} \]
\[ \text{Reaction 2: } \text{C}_2\text{H}_2\text{(g)} + \text{H}_2\text{(g)} \rightarrow \text{C}_2\text{H}_4\text{(g)} \quad \Delta H^\circ_2 \]

Both reactions break \(1\text{ mol}\) of \(\text{H}-\text{H}\) bonds and form \(2\text{ mol}\) of \(\text{C}-\text{H}\) bonds per mole of reaction. Which of the following best predicts and explains the value of \(\Delta H^\circ_2\) relative to \(\Delta H^\circ_1\)?

- **A.** \(\Delta H^\circ_2\) is more negative than \(\Delta H^\circ_1\) because the \(\text{C}-\text{H}\) bonds formed in \(\text{C}_2\text{H}_4\text{(g)}\) have a higher bond order and release significantly more energy upon formation than the \(\text{C}-\text{H}\) bonds formed in \(\text{CH}_4\text{(g)}\).
- **B.** \(\Delta H^\circ_2\) is more negative than \(\Delta H^\circ_1\) because breaking a \(\text{C}-\text{C}\) \(\pi\) bond in \(\text{C}_2\text{H}_2\text{(g)}\) requires less energy than breaking the \(\text{C}-\text{C}\) \(\sigma\) bond in \(\text{C}_2\text{H}_6\text{(g)}\).
- **C.** \(\Delta H^\circ_2\) is less negative than \(\Delta H^\circ_1\) because the \(\pi\) bond in \(\text{C}_2\text{H}_2\text{(g)}\) has greater electron density between the carbon nuclei, requiring more energy to break than the \(\sigma\) bond in \(\text{C}_2\text{H}_6\text{(g)}\).
- **D.** \(\Delta H^\circ_2\) is less negative than \(\Delta H^\circ_1\) because the \(\text{C}\equiv\text{C}\) triple bond in \(\text{C}_2\text{H}_2\text{(g)}\) must be completely broken into separate carbon atoms before \(\text{C}_2\text{H}_4\text{(g)}\) can form.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/123845/*