Problem: Consider two reactions for the production of ethanol:C2H4 (g) + H2O (g) → CH3CH2OH (l)C2H6 (g) + H2O (g) → CH3CH2OH (l) + H2 (g)Which would be the more thermodynamically feasible at standard conditions? Why?

FREE Expert Solution

We are asked which would be the more thermodynamically feasible at standard conditions among the two reactions for the production of ethanol. 


C2H4 (g) + H2O (g) → CH3CH2OH (l)
C2H6 (g) + H2O (g) → CH3CH2OH (l) + H2 (g)


Calculate the ΔG: 

C2H4 (g) + H2O (g) → CH3CH2OH (l)


Given: 

ΔG f ° for C2H4 (g) =  +68.4 kJ/mol

ΔG f ° for H2O (g) = -228.5 kJ/mol

ΔG f ° for CH3CH2OH (l) = -172.8 kJ/mol


 ΔG˚rxn =  ΔG˚f, products - ΔG˚f, reactants ΔG˚rxn =  [ΔG˚f, CH3CH2OH ] - [ΔG˚f, C2H4 + ΔG˚f, H2O]ΔG˚rxn = [ (1 mol)(-172.8 kJ/mol)]             -  [(1 mol)(+68.4 kJ/mol) + (1 mol)(-228.5 kJ/mol) ] 

ΔG˚rxn = -12.7 kJ


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Problem Details

Consider two reactions for the production of ethanol:

C2H4 (g) + H2O (g) → CH3CH2OH (l)
C2H6 (g) + H2O (g) → CH3CH2OH (l) + H2 (g)

Which would be the more thermodynamically feasible at standard conditions? Why?

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