Energy from Fuels · Paper 1 — Free Chemistry SL Practice Questions

A student burns ethanol ($\text{C}_2\text{H}_5\text{OH}$) in a simple calorimeter and obtains an experimental enthalpy change of combustion of $-1360 \ \text{kJ mol}^{-1}$. The literature value is $-1367 \ \text{kJ mol}^{-1}$. Which explanation best accounts for the experimental value being less exothermic than the literature value?

Equal masses of $1.00\ \text{g}$ of methane ($\text{CH}_4$), propane ($\text{C}_3\text{H}_8$), and octane ($\text{C}_8\text{H}_{18}$) are each burned completely in separate calorimeters containing identical volumes of water. Using the standard enthalpies of combustion below, which option correctly ranks the resulting temperature increases from smallest to largest? Fuel — $\Delta H_c^\circ\ /\ \text{kJ mol}^{-1}$ $\text{CH}_4$ — $-890$ $\text{C}_3\text{H}_8$ — $-2220$ $\text{C}_8\text{H}_{18}$ — $-5470$

A student burns $0.50\ \text{g}$ of methane, $\text{CH}_4$, completely in excess oxygen. The standard enthalpy of combustion of methane is $-890\ \text{kJ mol}^{-1}$. The energy released heats $200\ \text{g}$ of water with a specific heat capacity of $4.18\ \text{J g}^{-1}\text{K}^{-1}$. Assuming no heat loss to the surroundings, what is the expected temperature increase of the water? (Molar mass of $\text{CH}_4 = 16.0\ \text{g mol}^{-1}$)

When 1.00 mol of methane undergoes complete combustion under standard conditions, the enthalpy change is $-890 \text{ kJ mol}^{-1}$. The heat released by burning 0.0500 mol of methane is transferred entirely to 500 g of water with a specific heat capacity of $4.18 \text{ J g}^{-1} \text{ K}^{-1}$. What is the expected temperature rise of the water?