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I need to calculate the enthalpy of combustion for butane using bond formation energies. Thus I first balance the equation:

$$\ce{C4H10 + 6.5 O2 -> 4 CO2 + 5 H2O}$$

Then, I calculate the total bond energy difference:

$$-3 \cdot 346.0 - 10 \cdot 414.2 - 6.5 \cdot 479.9 + 4 \cdot 2 \cdot 803.3 + 5 \cdot 2 \cdot 464.4 = 2771.05~[\pu{kJ mol-1}]$$

This is different than the correct value of $\pu{2877.5 kJ mol-1}$. Why?

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    $\begingroup$ To me the equation is conceptually correct and I guess the deviation in numbers lies in different sources of information for average bond energies. Note that C-H BEs even in butane only deviate a lot in different sources (from 400.4 to 421.3 kJ/mol), averaging these values bring more error to the table. I'd say you did it correctly and call it a day; if someone will argue, point them to Luo's Comprehensive handbook of chemical bond energies. $\endgroup$ – andselisk Nov 6 '17 at 1:29
  • $\begingroup$ Thank you. That's what I'll do. I see you corrected the markup too, could you point me to a good guide for LaTeX for chemistry? $\endgroup$ – user61143 Nov 6 '17 at 1:33
  • $\begingroup$ Sure, here is the mini-guide for chemistry and a comprehensive tutorial for math. Welcome aboard:) $\endgroup$ – andselisk Nov 6 '17 at 1:39
  • $\begingroup$ I assume you did not mix energies and enthalpies? You found tabulated values for bond enthalpies? $\endgroup$ – TAR86 Nov 6 '17 at 5:53
  • $\begingroup$ The thing is, the values you used are possibly the mean bond enthalpy of those bonds and not the specific bond enthalpy of the bonds in butane. What is the difference? The chemical environment of a bond will affect its enthalpy. As the mean bond enthalpy uses enthalpy of a bond in various compounds it is never the same as the actual bond enthalpy in a given compound. $\endgroup$ – looneysnoop Nov 6 '17 at 11:00

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