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If a reaction has an enthalpy of -54.32 kJ/mol and an entropy of -354.2 J/(K*mol), what is the Gibbs free Energy at 54.3(degrees c)?
DeltaG = DeltaH - TDeltaS
dG = -54.32 kJ/mol - (54'32+273)K(-354.2J/molK)
NB Thevtemperature is quoted in Kelvin(K) and the Entropy must be converted to kJ by dividing by '1000'/
Hence
dG = - 54.32kJ/mol - (327.32K)(-0.3542 kJ/molK)
NB The 'K' cancels out. Then maker the multiplication
dG = -54/32 kJ/mol - - 115.94 kJ/mol Note the double minus; it becomes plus(+).
Hence
dG = -54.32kj/mol + 115.94 kJ/mol
dG = (+)61.61 kJ/mol
Since dG is positive, the reaction is NOT thermodynamically feasible.
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