Measuring enthalpy changes practice questions.

A student dissolves solid ammonium nitrate in water and records that the temperature of the solution falls. State what this observation indicates about the dissolving process.

• It is exothermic, ΔH is positive
• It is exothermic, ΔH is negative
• It is endothermic, ΔH is positive
• It is endothermic, ΔH is negative

Answer: It is endothermic, ΔH is positive

A temperature fall means the system absorbs heat from the surroundings, so the process is endothermic and ΔH is positive (products have higher enthalpy than reactants).

In a coffee-cup calorimeter, 100.0 g of solution (assume c = 4.18 J g⁻¹ K⁻¹) rises in temperature by 6.5 K when 0.0500 mol of acid is neutralised. Calculate the enthalpy of neutralisation, in kJ mol⁻¹.

• -2717 kJ mol⁻¹
• -54.3 kJ mol⁻¹
• -27.2 kJ mol⁻¹
• +54.3 kJ mol⁻¹

Answer: -54.3 kJ mol⁻¹

q = 100.0 × 4.18 × 6.5 = 2717 J = 2.717 kJ; ΔH = -2.717 / 0.0500 = -54.3 kJ mol⁻¹ (exothermic, so negative).

Which statement about an endothermic reaction is correct?

• The surroundings become warmer and ΔH is negative
• The system releases heat to the surroundings
• The products have greater enthalpy than the reactants
• The reaction cannot occur at room temperature

Answer: The products have greater enthalpy than the reactants

In an endothermic reaction the system absorbs heat, so products lie at higher enthalpy than reactants and ΔH is positive; the surroundings cool.

A calorimetry experiment measuring the enthalpy of combustion of a fuel usually gives a value less exothermic (smaller magnitude) than the data-book value. Suggest the main reason.

• Heat is lost to the surroundings and the container
• The fuel undergoes only complete combustion
• The specific heat capacity of water is too high
• Too many moles of fuel are burned

Answer: Heat is lost to the surroundings and the container

Heat loss to the surroundings and calorimeter (plus incomplete combustion and evaporation) means not all released energy heats the water, so the measured value is smaller in magnitude.

50.0 cm³ of 1.00 mol dm⁻³ HCl is mixed with 50.0 cm³ of 1.00 mol dm⁻³ NaOH; the temperature rises by 6.8 K. Taking the total mass as 100.0 g and c = 4.18 J g⁻¹ K⁻¹, calculate the enthalpy of neutralisation, in kJ mol⁻¹.

• -2842 kJ mol⁻¹
• -28.4 kJ mol⁻¹
• -56.8 kJ mol⁻¹
• -142 kJ mol⁻¹

Answer: -56.8 kJ mol⁻¹

q = 100.0 × 4.18 × 6.8 = 2842 J = 2.842 kJ; n = 0.0500 mol; ΔH = -2.842 / 0.0500 = -56.8 kJ mol⁻¹.

Which change would be expected to have a positive enthalpy change (endothermic)?

• Condensation of steam to liquid water
• Freezing of liquid water to ice
• Combustion of methane in oxygen
• Melting of ice to liquid water

Answer: Melting of ice to liquid water

Melting absorbs energy to overcome intermolecular forces, so ΔH > 0. Condensation, freezing and combustion all release energy (ΔH < 0).