How can you calculate the total bond energy of CO2? I've searched for the bond energy of C=O and found it was 745, but sometimes 799. Which one is the right one to use? And do those numbers mean the bond energy of CO2?

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    $\begingroup$ there are many C=O bonds with different chemical environments, bearing slightly different energies. Usually the reported value is an average of many instances of that bond. $\endgroup$ – khaverim Feb 7 '16 at 21:47

You're asking a pretty complex question. The short answer is that to calculate such an energy you have to utilize quantum mechanical calculations, or otherwise perform macroscopic experiments on bulk $\ce{CO2}$.

A quick ab-initio calculation of a O=C=O molecule with bond-length 1.164 Angstroems (quite close to experimental and theoretically deduced length) yields the following Hartree-Fock SCF energy (using the 6-31G* basis set):

-187.634176090515 Hartrees, which is -492633.5668524824 kJ/mol

This is a huge negative energy, and rightly so -- it is how much energy is theoretically required to break the $\ce{CO2}$ molecule entirely into subatomic particles. In other words, it is the energy holding $\ce{CO2}$ together -- the energy that defines the molecule.

Better approximations can be made with higher levels of theory but don't deviate from these values by much.

The simplistic route is to add up collective bonds from a data table provided in your textbook, e.g. $2 \times$ C=O = total molecular bond energy. This is not the same thing as the total molecular energy though: the latter includes energetic contribution of the nuclear-electron interactions, not just atomic bonds.

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I wouldn't worry too much about the slight differences in the bond enthalpy for the C$=$O. I will discuss why after the next paragraph.

The bond enthalpy is the amount of energy (in kJ) that is required to break 1 mol of that bond. Therefore, it is also a measurement of how much energy is released as 1 mol of that same bond is formed. The bond enthalpy is dependent upon the strength of the bond. The higher the strength, the larger the bond enthalpy.

The bond enthalpy is more of an average value than an accurate figure. For example, the energy required to break the first double bond between the central carbon atom and the oxygen is different from the energy required to break the second double bond. This is because the strength of the bond is dependent upon its environment. When you break off that first oxygen atom, it changes the environment.

So you can use any one of the those bond enthalpy values.

Source(s): Any good textbook in general Chemistry should discuss how enthalpy bonds are calculated. Personally, I used the Pearson Baccalaureate Textbook for IB Chemistry HL

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  • $\begingroup$ Thanks! This helped. But what is the total bond energy of CO2? $\endgroup$ – user3658408 Feb 7 '16 at 21:52
  • $\begingroup$ In $\ce{CO2}$, the potential energy in the carbon-oxygen double bonds is equal to the number of such bonds in the molecule multiplied with the average bond enthalpy. If you are in doubt, it is always useful to draw the Lewis Structure of the molecule to help define the bonds. $\endgroup$ – Mattias Feb 7 '16 at 22:02

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