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?