Two alternative liquid fuels, X and Y, are being evaluated for use in a remote power generator. Complete combustion in excess O₂ produces CO₂(g) and H₂O(l).
Fuel X molecular formula: C₂H₄O (molar mass = 60.1 g mol⁻¹)
Fuel Y molecular formula: C₂H₆ (molar mass = 30.1 g mol⁻¹)
A data dashboard presents stacked-bar representations of average bond energies (kJ mol⁻¹). The numerical heights for each distinct bond are listed below:
Bonds in fuels
• C=C 614
• C–C 348
• C–H 413
• C–O 358
• O–H 463 (present only after combustion)
Bonds in products
• C=O (in CO₂) 799
• O–H (in H₂O) 463
• O=O 498
Fuels contain the bonds indicated by their formulas: X has one C=C, one C–O, and four C–H bonds; Y has one C–C and six C–H bonds. Oxygen molecules contribute one O=O bond each.
Combustion reactions (unbalanced):
C₂H₄O(l) + O₂(g) → CO₂(g) + H₂O(l)
C₂H₆(l) + O₂(g) → CO₂(g) + H₂O(l)
(a) Write the balanced chemical equation for complete combustion of one mole of Fuel X, including physical states.
(b) Using the bond-energy data, determine the enthalpy of combustion per mole of Fuel X. Show the totals for bonds broken and bonds formed.
(c) Repeat the calculations in part (b) for Fuel Y.
(d) On the basis of your results, identify which fuel releases more energy per gram when burned completely. Justify your answer with numerical evidence.