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 \(\Delta H^\circ_1\)?