This is a question from a sample exam that I can't figure out. The answer is 1297 kJ/mol, but I don't know how to solve it. Any help is appreciated.
The following are average bond energies (kJ/mol):
C−H 413 O−H 463 C=C 614
C−C 348 O−O 146 C=O 799
C−O 358 O2 495
What is the amount of heat released in the complete combustion of ethylene, H2C=CH2?
H(reaction) = H(products) - H(reactants)
The Attempt at a Solution
I've written the balanced equation...
C2H4 + 3O2 -> 2H2O + 2CO2
Now, I think I need to break down the reaction and add all of the energies from broken bonds (reactants) together, and then subtract that sum from the energy released from bonds made (products).
Hf of reactants :
C2H4 -> one C=C bond, four C-H bonds = (614 + 4*413) = 2266 kJ/mol * 1mol = 2266 kJ
O2 -> this is an element in its normal state, so the enthalpy is zero
Hf of products :
H2O -> two O-H bonds = (2*463) = 926 kJ/mol * 2mol = 1852 kJ
CO2 -> two C=O bonds = (2*799) = 1598 kJ/mol * 2mol = 3196 kJ
I continued like this, but the answer I got was wrong. Am I completely off base? I have a feeling I'm misunderstanding the question and approaching it in the wrong way.