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## Homework Statement

"Estimate the standard reaction Gibbs energy of the following reaction:

N

_{2}+ 3 H

_{2}‒‒> 2 NH

_{3}

at 100K and at 1000K."

## Homework Equations

ΔS(T2) = S°(T1) + ∫ n C

_{p}dT/T

ΔG = ΔH ‒ TΔS

Given data: http://imgur.com/MBakUEB (may need to right-click and select "Open in new tab")

## The Attempt at a Solution

So, I realize that we likely have to use the formula ΔS(T2) = S°(T1) + ∫ n C

_{p}dT/T in our calculations to find the ΔS at the different temperatures. What I did was I calculated the ΔS at 100K for NH3, H2, and N2 separately, then did Δ

_{r}G = Δ

_{prod}G ‒ Δ

_{reac}G. I had to obtain values for C

_{p}, Δ

_{f}H

_{m}°, and ΔS

_{m}° from a table in the back of my textbook.

Here's my math:

Δ

_{r}H = 2 mol × ‒46.11 kJ/mol, because it is 0 for hydrogen and nitrogen.

http://imgur.com/TV8cfmB

Forgive the non-matching formatting. I found it was easier to go into Word and use its equation editor than to try and learn LaTEX, as time is somewhat of the essence here.

So is my work for the 100K case correct? I just realized that I may have needed to also convert my H value from the 273K to 100K, because the book gave standard enthalpies of reaction, and not that at 100K... Would I use the Kirchoff equation for that, then?

Thanks for the help!