I am trying to determine two of the theoretical values for the enthalpy of the combustion of methanol.

I have calculated (using Hess' law) that the $ \Delta H_{\mathrm {reaction}}\; = -1453.418~\mathrm{kJ~mol^{-1}} $, which is correct according to several sources. However, when I tried to calculate the $ \Delta H_{\mathrm{reaction}} $ using the bond dissociation and formation method: $ \Delta H_{\mathrm{reaction}}\; =\; \sum{\Delta H_{\mathrm{bonds\; broken}}}\;-\; \sum{\Delta H_{\mathrm{bonds\; formed}}} $

I got a value of $\mathrm{-464.9~kJ~mol^{-1}}$. I am just wondering if this is correct, or if I have made a mistake somewhere?

Ps: My balanced equation looks like this: $$\ce{2CH3OH + 3O2 -> 4H2O + 2CO2}$$

I used this table to get the known values for the various bond energies.


  • $\begingroup$ Could you show us your actual calculations? Otherwise it is difficult to see where the mistake is. $\endgroup$ – bon Sep 6 '15 at 13:45
  • $\begingroup$ Could the discrepancy be partly due to the fact that the bond energies in your table are given in kcal per mole? $\endgroup$ – Ivan Neretin Sep 6 '15 at 14:00

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