For a particular reaction, ΔH = -144.6 kJ and ΔS = -301.2 J/K. Calculate ΔG for this reaction at 298 K.
Answers
The Gibbs free energy change for this reaction at 298 K is approximately -54,484.8 J.
Determining the Gibbs free energy change
Calculating the Gibbs free energy change (ΔG) for a reaction at a given temperature using the Gibbs-Helmholtz equation,
ΔG = ΔH - TΔS
ΔG = (-144,600 J) - (298 K)(-301.2 J/K)
ΔG = -144,600 J + 90,115.2 J
ΔG = -54,484.8 J
ΔG = -54,484.8 J
Converting to kilojoules:
ΔG = -54,484.8 J / 1,000
= -54.48 kJ
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Either way, use the Gibb's Formula:
ΔG= ΔH-TΔS
If the final answer is supposed to be in J/K then plug-in your values to the formula and simply solve.
However, if your final answer is supposed to be in kJ, your first step would be to convert ΔS from J/K into kJ.
This is simply done by dividing the given ΔS by 1000 ----> (-301.2/1000).
(we divide by 1000 because by definition that's what a "kilo" is )
New ΔS value= -0.3012 kJ
Now all you have to do is plug-in all the values into the formula.
ΔG= ΔH-TΔS
ΔG= (-144.6 kJ) - 298K(-0.3012kJ)
= -54.84 kJ
Since the ΔG is negative, this means that the reaction is spontaneous.
Hope this helps! :)
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I hope this helps.
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