Thanks for showing your work; it makes it easy to spot the error.
#1. I don't believe you converted T from C to K correctly. 273.15 + 195.1 = 468.25.
I didn't check further.
For a particular reaction at 195.1 °C, ΔG = -1488.09 kJ/mol, and ΔS = 288.67 J/(mol·K).?
Calculate ΔG for this reaction at -20.0 °C.
delta G = delta H - (T*delta S)
-1488.09 kJ/mol = (delta H) - ((461.8K)*(+288.67 J/(mol x K))*(1 kJ/1000 J))
-1488.09 kJ/mol = (delta H) - (133.3 kJ/mol)
-1354.7 kJ/mol = delta H
delta G = ((-1354.7 kJ/mol) + (253 K)*(+288.67 J/(mol x K))*(1 kJ/1000 J))
delta G = -1281.7 kJ/mol
Thats what I get for an answer but it tells me its wrong, can you tell me what I'm doing wrong?
4 answers
i got -1,279.89 and still got it wrong..doing the same calculations just fixing the 468.25 where it belongs
I don't get that.
I get -1352.9 for dH.
I didn't throw away the 0.15 from 273.15. Since dS is measured to so many places I used (253.15K) for the second T.
-1352.92 kJ/mol
Then dG = -1352.92 -(253.15)(0.28867)
Is this what you have? I don't think that is -1279.89.
I get -1352.9 for dH.
I didn't throw away the 0.15 from 273.15. Since dS is measured to so many places I used (253.15K) for the second T.
-1352.92 kJ/mol
Then dG = -1352.92 -(253.15)(0.28867)
Is this what you have? I don't think that is -1279.89.
Oh I see what I did wrong! Thanks!