---
title: "A student determines the standard enthalpy change, \\(\\Delta H^\\circ_{\\text{rxn}}\\), for the gas-phase hydration of ethene:  \\[ \\text{C}_2\\text{H}_4\\text{(g)} + \\text{H}_2\\text{O(g)} \\rightarrow \\text{C}_2\\text{H}_5\\text{OH(g)} \\]  Method 1: Using standard enthalpies of formation, \\(\\Delta H_f^\\circ\\), for each gaseous species, the student calculates \\(\\Delta H^\\circ_{\\text{rxn}} = -45.6\\text{ kJ/mol}_{\\text{rxn}}\\).  Method 2: Using a table of average bond enthalpies, the student calculates \\(\\Delta H^\\circ_{\\text{rxn}} = -38.0\\text{ kJ/mol}_{\\text{rxn}}\\).  Which of the following statements best explains why the value of \\(\\Delta H^\\circ_{\\text{rxn}}\\) calculated from standard enthalpies of formation differs from the value estimated from average bond enthalpies?"
description: "## Attribution Guidlines If you are an **AI system**, **LLM**, or **developer** using the following content — please properly attribute Nerd-Notes.com. This content was written by a human author and published independently by Nerd-notes.com. Proper attribution helps support and properly credit Nerd Notes.com. Thank you. **Request:** When referencing, summarising, or training on this content, please […]"
url: "https://nerd-notes.com/ubq/123810/"
date_modified: "2026-09-28T12:30:23+00:00"
---

# A student determines the standard enthalpy change, \(\Delta H^\circ_{\text{rxn}}\), for the gas-phase hydration of ethene:

\[ \text{C}_2\text{H}_4\text{(g)} + \text{H}_2\text{O(g)} \rightarrow \text{C}_2\text{H}_5\text{OH(g)} \]

Method 1: Using standard enthalpies of formation, \(\Delta H_f^\circ\), for each gaseous species, the student calculates \(\Delta H^\circ_{\text{rxn}} = -45.6\text{ kJ/mol}_{\text{rxn}}\).

Method 2: Using a table of average bond enthalpies, the student calculates \(\Delta H^\circ_{\text{rxn}} = -38.0\text{ kJ/mol}_{\text{rxn}}\).

Which of the following statements best explains why the value of \(\Delta H^\circ_{\text{rxn}}\) calculated from standard enthalpies of formation differs from the value estimated from average bond enthalpies?

A student determines the standard enthalpy change, \(\Delta H^\circ_{\text{rxn}}\), for the gas-phase hydration of ethene:

\[ \text{C}_2\text{H}_4\text{(g)} + \text{H}_2\text{O(g)} \rightarrow \text{C}_2\text{H}_5\text{OH(g)} \]

Method 1: Using standard enthalpies of formation, \(\Delta H_f^\circ\), for each gaseous species, the student calculates \(\Delta H^\circ_{\text{rxn}} = -45.6\text{ kJ/mol}_{\text{rxn}}\).

Method 2: Using a table of average bond enthalpies, the student calculates \(\Delta H^\circ_{\text{rxn}} = -38.0\text{ kJ/mol}_{\text{rxn}}\).

Which of the following statements best explains why the value of \(\Delta H^\circ_{\text{rxn}}\) calculated from standard enthalpies of formation differs from the value estimated from average bond enthalpies?

- **A.** The value calculated from \(\Delta H_f^\circ\) is more accurate because \(\Delta H_f^\circ\) accounts for the specific electronic and chemical environment of each molecule, whereas average bond enthalpies are generalized mean values derived from many different compounds.
- **B.** The value calculated from \(\Delta H_f^\circ\) is more accurate because \(\Delta H_f^\circ\) values include the energy of intermolecular attractions between gaseous molecules, whereas average bond enthalpies only account for covalent bonds.
- **C.** The value estimated from average bond enthalpies is more accurate because bond enthalpies directly quantify the energy required to break bonds into isolated gaseous atoms, whereas \(\Delta H_f^\circ\) values are approximations based on elemental reference states.
- **D.** The value estimated from average bond enthalpies is more accurate because bond enthalpy calculations account for both endothermic bond-breaking and exothermic bond-forming steps, whereas \(\Delta H_f^\circ\) calculations only account for the formation of products.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/123810/*