1. **State the problem:** Calculate the heat of reaction ($\Delta H_{rxn}$) for the combustion of methanol using bond enthalpies. 2. **Formula:** $\Delta H_{rxn} = \text{Energy to break bonds} - \text{Energy to make bonds}$ 3. **Given data:** - Energy to break bonds (endothermic): $2803$ kJ - Energy to make bonds (exothermic): $-3451$ kJ 4. **Calculate $\Delta H_{rxn}$:** $$\Delta H_{rxn} = 2803 - (-3451) = 2803 + 3451 = 6254\text{ kJ}$$ 5. **Check for cancellation:** Since the formula is subtraction of a negative, no factors to cancel here. 6. **Interpretation:** The positive value indicates the reaction absorbs energy, but this contradicts typical combustion reactions which are exothermic. The problem's tables show bond breaking as positive and bond making as negative, so the correct formula is: $$\Delta H_{rxn} = \text{Energy to break bonds} + \text{Energy to make bonds} = 2803 + (-3451) = -648\text{ kJ}$$ 7. **Conclusion:** The heat of reaction is $-648$ kJ/mol, meaning the combustion of methanol is **exothermic**. **Final answer:** $$\Delta H_{rxn} = -648\text{ kJ/mol}$$ **Combustion of Methanol is exothermic.**