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?

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.

Oh I see what I did wrong! Thanks!

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.

i got -1,279.89 and still got it wrong..doing the same calculations just fixing the 468.25 where it belongs

Based on the equation you provided, ΔG = ΔH - (T * ΔS), it seems like you made a mistake in your calculations.

To calculate ΔG at -20.0 °C, you need to convert the temperature from Celsius to Kelvin. Remember that the Kelvin scale does not have negative values, so you need to add 273.15 to the temperature in Celsius.

Also, in the equation, make sure that the units of ΔS match the units of temperature. In this case, ΔS is given in J/(mol·K), so the temperature needs to be in Kelvin.

Let's go through the calculations again:

Temperature at -20.0 °C = (-20.0 + 273.15) K = 253.15 K

Now, let's calculate ΔH:

ΔH = ΔG + (T * ΔS)
ΔH = -1488.09 kJ/mol + (253.15 K * 0.28867 kJ/(mol x K))
ΔH = -1488.09 kJ/mol + 73.1054 kJ/mol
ΔH = -1414.9846 kJ/mol

Finally, let's substitute the values back into the equation to calculate ΔG:

ΔG = ΔH - (T * ΔS)
ΔG = -1414.9846 kJ/mol - (253.15 K * 0.28867 kJ/(mol x K))
ΔG = -1414.9846 kJ/mol - 73.1054 kJ/mol
ΔG = -1488.09 kJ/mol

So, the correct value of ΔG for this reaction at -20.0 °C is -1488.09 kJ/mol, which is the same value given in the question.