[mrmo123]

hey...i just took a test today an really wanna know why doesn't this work...

through hess's law i found that the delta h of N2 + 2O2 => 2NO2 is around 68 (using thermo gives a value of 66.4)

however...when i wrote out the lewis structures, and found delta h using the bond energies, i got a totally different answer...is this because of the lone electron on NO2??

***N triple bond is 945 kJ/mol
O double bond is 498 kj/mol
N-O bond is 201 kj/mol and N double bond O is 607 kj/mol***
so 945 + 2(498) - 2(201) - 2 (607) = 325 kj/mol



- thx for all the help in advance

[Yggdrasil]

NO₂ has a resonance structure, meaning that a conjugated pi bond exists in the molecule.  The conjugation adds to the stability of NO₂, lowering the enthalpy of the molecule.  Since using only bond energies does not take this conjugation into account, it makes sense that your estimate overestimates the delta H.

[mrmo123]

but doesn't the higher bond energy of the double bond relative to a single bond take this into account??...and the estimate value is ~5x higher...isn't that unreasonable?  -thanks for your reply

[Yggdrasil]

The higher bond energy of the double bond does not take conjugation into account.  It's one of the situations where the whole is greater than the sum of its parts.

Also, it's not relevant to say that the energy is 5x higher since the zero of any energy scale is arbitrary.

Anyway, this example is one of many that shows that estimates of reaction enthalpies from bond energies can be fairly inaccurate.