Step 1

$\mathbf{Q}\mathbf{}\mathbf{=}\mathbf{}{\mathbf{C}}_{\mathbf{cal}}\mathbf{\u2206}\mathbf{T}\phantom{\rule{0ex}{0ex}}\mathbf{Q}\mathbf{}\mathbf{=}\mathbf{}\mathbf{(}\mathbf{6}\mathbf{.}\mathbf{21}\mathbf{}\frac{\mathbf{kJ}}{\overline{)\mathbf{\xb0}\mathbf{C}}}\mathbf{)}\mathbf{(}\mathbf{27}\mathbf{.}\mathbf{1}\overline{)\mathbf{\xb0}\mathbf{C}}\mathbf{)}$

**Q = 168.291 kJ**

Step 2

$\mathbf{3}\mathbf{.}\mathbf{80}\mathbf{}\overline{)\mathbf{g}\mathbf{}{\mathbf{C}}_{\mathbf{8}}{\mathbf{H}}_{\mathbf{8}}}\mathbf{}\mathbf{\times}\frac{\mathbf{1}\mathbf{}\mathbf{mol}\mathbf{}{\mathbf{C}}_{\mathbf{8}}{\mathbf{H}}_{\mathbf{8}}}{\mathbf{104}\mathbf{.}\mathbf{144}\mathbf{}\overline{)\mathbf{g}\mathbf{}{\mathbf{C}}_{\mathbf{8}}{\mathbf{H}}_{\mathbf{8}}}}$** = 0.0365 mol C _{8}H_{8}**

When a 3.80-g sample of liquid octane (C_{8}H_{18}) is burned in a bomb calorimeter, the temperature of the calorimeter rises by 27.1 °C. The heat capacity of the calorimeter, measured in a separate experiment, is 6.21 kJ/°C. Determine ΔE for octane combustion in units of kJ/mol octane.

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