Four fuels undergo complete combustion in excess oxygen, O₂, and the energy released is used to heat 1000 g of water - VCE - SSCE Chemistry - Question 22 - 2018 - Paper 1

Using the combustion heats:

• Hydrogen (H₂): Standard enthalpy change is approximately -286 kJ/mol.
• Propane (C₃H₈): Approximately -2040 kJ/mol.
• Methane (CH₄): Approximately -890 kJ/mol.
• Methanol (CH₃OH): Approximately -726 kJ/mol.

Then convert these to energy released per gram or mole as necessary.

Calculate the energy released for each option based on the amounts given and compare:

• For hydrogen:

  • $0.889 g \approx \frac{0.889 g}{2 g/mol} \times -286 kJ/mol \approx -126.4 kJ$

• For propane:

  • $3.95 g \approx \frac{3.95 g}{44 g/mol} \times -2040 kJ/mol \approx -183.1 kJ$

• For methane:

  • $0.282 mol \approx 0.282 mol \times -890 kJ/mol \approx -250.3 kJ$

• For methanol:

  • $0.301 mol \approx 0.301 mol \times -726 kJ/mol \approx -218.4 kJ$

Compare the energy released with 250.8 kJ. Only methanol (D) releases sufficient energy.