Given the standard enthalpies of formation for the following substances, determine the change in enthalpy for the combustion of 1.0 mol of butane. $\begin{array}{lc}
\text { Substance } & \Delta \boldsymbol{H} \
& \circ \
& \mathrm{f} \
& \text { in } \mathbf{k J} / \mathbf{m o l} \
\mathrm{C}_{4} \mathrm{H}_{10}(g) & -124.7 \
\mathrm{CO}_{2}(g) & -393.5 \
\mathrm{H}_{2} \mathrm{O}(g)& -241.8
\end{array}$

Question

Given the standard enthalpies of formation for the following substances, determine the change in enthalpy for the combustion of 1.0 mol of butane. $\begin{array}{lc}
\text { Substance } & \Delta \boldsymbol{H} \
& \circ \
& \mathrm{f} \
& \text { in } \mathbf{k J} / \mathbf{m o l} \
\mathrm{C}_{4} \mathrm{H}_{10}(g)  & -124.7 \
\mathrm{CO}_{2}(g) & -393.5 \
\mathrm{H}_{2} \mathrm{O}(g)& -241.8
\end{array}$

Quizplus · Accepted Answer

The balanced combustion reaction for butane is: $2 \mathrm{C}_{4} \mathrm{H}_{10}(g) + 13 \mathrm{O}_{2}(g) \rightarrow 8 \mathrm{CO}_{2}(g) + 10 \mathrm{H}_{2} \mathrm{O}(g)$. The change in enthalpy is calculated using the formula: $\Delta H = \sum \Delta H_f (\text{products}) - \sum \Delta H_f (\text{reactants})$. For 1 mol of butane, $\Delta H = [8(-393.5) + 10(-241.8)] - [2(-124.7)] = -2,658 \text{ kJ/mol}$.

Given the Standard Enthalpies of Formation for the Following Substances