Calorimetry & Enthalpy Change – Lecture Review

These flashcards review key equations, procedures, assumptions, error sources, and calculation steps for calorimetry and enthalpy change experiments in the laboratory.

What is the general laboratory method for measuring an enthalpy change (ΔH)?

Measure the temperature change of a known quantity of substance reacting or dissolving, calculate heat energy (q) with q = mcΔT, then convert to ΔH using ΔH = –q ⁄ n.

Which equation is used to calculate heat energy transferred in a calorimetry experiment?

q = mcΔT, where q = heat energy (J), m = mass of solution (g), c = specific heat capacity (4.18 J g⁻¹ K⁻¹ for water), and ΔT = temperature change (°C or K).

How is enthalpy change per mole (ΔH) calculated from experimental data?

ΔH = –q ⁄ n, where q is the heat energy in joules and n is the moles of the limiting reagent; the negative sign indicates heat released in exothermic reactions.

What are the usual units for enthalpy change (ΔH)?

kJ mol⁻¹ (convert q from J to kJ by dividing by 1000 before applying ΔH = –q ⁄ n).

List four key assumptions made when using q = mcΔT in solution calorimetry.

1) All heat transfer occurs within the solution, 2) The solution’s specific heat capacity equals that of water (4.18 J g⁻¹ K⁻¹), 3) No heat is lost to surroundings, 4) The reaction goes to completion.

Outline the experimental procedure to measure the enthalpy change of neutralisation.

Use an insulated polystyrene cup, measure equal known volumes of acid and alkali, record their initial temperatures, mix and stir, record maximum temperature, then apply q = mcΔT (using total mass ≈ volume) and ΔH = –q ⁄ moles of water formed.

Why is a polystyrene cup preferred in solution-based enthalpy experiments?

Polystyrene is a good thermal insulator, reducing heat exchange with surroundings and improving accuracy of measured temperature changes.

Why do chemists often assume the density of dilute aqueous solutions is 1 g cm⁻³?

It lets them equate volume in cm³ directly to mass in grams, simplifying calorimetric calculations and introducing minimal error for dilute solutions.

Describe how to measure the enthalpy of combustion using a simple calorimeter.

Fill a metal calorimeter with a known mass of water, weigh a spirit burner before and after burning, burn the fuel beneath the water, record the temperature rise, calculate q = mcΔT for the water, determine moles of fuel burned, then compute ΔH = –q ⁄ n.

Give four sources of error when measuring an enthalpy of combustion with a simple calorimeter.

Heat loss to surroundings, incomplete combustion of fuel, evaporation of fuel or water, and inadequate insulation.

State five methods to improve the accuracy of calorimetry experiments.

Use a draught shield, place a lid on the calorimeter, insulate with foil or polystyrene, employ a digital thermometer/temperature probe, or use a bomb calorimeter for combustion.

Why might an experimentally determined ΔH be less exothermic than the accepted value?

Heat loss to surroundings, evaporation of liquids, measurement inaccuracies, or the reaction being incomplete or slow.

In the example where 25 cm³ of 1.00 mol dm⁻³ HCl reacts with 25 cm³ of 1.00 mol dm⁻³ NaOH and ΔT = 6.5 °C, what is the calculated ΔH?

–54.3 kJ mol⁻¹ (q = 1.358 kJ; moles of water formed = 0.025 mol; ΔH = –1.358 ⁄ 0.025).

What specific heat capacity value is usually assumed for aqueous solutions in calorimetry?

4.18 J g⁻¹ K⁻¹, the specific heat capacity of water.

In ΔH = –q ⁄ n, which quantity of substance is used for n?

The number of moles of the limiting reagent (or, for neutralisation, moles of water produced; for combustion, moles of fuel burned).

What does the negative sign in ΔH = –q ⁄ n signify in an exothermic reaction?

That heat is released by the system to the surroundings, making ΔH negative for exothermic processes.