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4D.23 7th Ed.

Posted: Tue Jan 29, 2019 7:55 pm
by JulieAljamal1E
The question asks to calculate the standard enthalpy formation of dinitrogen pentoxide from the info:
2NO+O2–>2NO2 delta H= -114.1kJ
4NO2+O2->2N2O5 delta H= -110.2 kJ
I’m confused on how to do this. The solutions show to add the first reaction with half of the second, but why are you allowed to multiply the second equation by 1/2. Then after it adds the two up it says the enthalpy of the reaction equals the enthalpy of formation of N2O5 minus twice the enthalpy formation of NO, how is that so? Can you step by step explain this problem, thank you!

Re: 4D.23 7th Ed.

Posted: Wed Jan 30, 2019 10:17 am
by Sarah Kiamanesh 1D
For this problem, we want to cancel out NO2 using Hess's Law because that's the only component for which we don't know the standard enthalpy of formation for. To cancel this, we multiply the second equation by 1/2 and add the two equations. Remember to adjust the standard enthalpies of the reactions as you manipulate the equations. Adding the equations results in
2NO + 3/2O2 --> N2O5
Delta H = -114.1 + (-55.1) = -169.2KJ
Now, we use Delta Hrxn = (nHf(products))-(nHf(reactants))
-169.2KJ = (1)Delta Hf(N2O5) - (3/2)Delta Hf(O2) - (2)Delta Hf(NO)
-169.2KJ = Delta Hf(N2O5) - 0 - 2(90.25)
Delta Hf(N2O5) = 11.3KJ