Part 2: Calorimetry & practical methods

1. q = mcΔT

A temperature change is numerically the same in K and °C: a rise from 20.0 °C to 26.6 °C is 6.6 K. Do not add 273 to ΔT.

For a solution, the mass is usually found from its volume using the density you are given (often 1.00 g cm⁻³). When two solutions are mixed, use the total volume of both: 25.0 cm³ + 25.0 cm³ → 50.0 g.

2. From q to ΔH

In an ideal calorimeter, the heat released by the reaction all goes into the solution: q_reaction = −q_surroundings. A temperature rise means ΔH is negative.

n must match the mole basis you are asked for:

• Per mole of a named substance: use n of that substance that actually reacted. Find the limiting reactant with the equation first.
• Per mole of the reaction as written: use the extent of reaction, n_i ÷ (coefficient of i), not simply the raw moles of the limiting reactant.

3. Worked examples: reactions in solution

4. Worked examples: burning fuels

5. Required Practical 2: measuring an enthalpy change

Reaction in solution

1. Measure a known volume of one solution (pipette or burette) into a polystyrene cup inside a beaker, for stability. Use a lid to reduce heat loss.
2. Record the temperature every minute for about 3 minutes before mixing, to set a baseline.
3. At a recorded time (say 4 minutes), add a known volume of the second solution, or a weighed mass of solid (weigh by difference).
4. Stir, and keep the thermometer in the liquid, not touching the cup's wall.
5. Keep recording every minute for several minutes after the maximum (or minimum).

Both solutions should start at the same temperature. If they do not, you would need a proper weighted heat balance; a plain average of two temperatures is only valid when the masses (and heat capacities) are equal.

Burning a fuel

• Measure the volume (and so the mass) of water in a copper can; record its starting temperature.
• Weigh the spirit burner with its cap on, before and after.
• Use a draught shield, leaving a gap so oxygen still reaches the flame. Keep the flame a fixed short distance below the can.
• Stir the water. Put the cap back on as soon as you put the flame out, to reduce evaporation of fuel before reweighing.

6. Temperature–time graphs and extrapolation

For an exothermic reaction, the cup loses heat while the reaction is still going, so the highest reading is lower than the true value. (The endothermic case, where the mixture warms back up, is explained below the graph.) To correct for this:

1. Plot temperature (y) against time (x).
2. Draw a best-fit line through the steady points before mixing; extend it to the mixing time.
3. Draw a best-fit line through the cooling points after the maximum; extrapolate it back to the mixing time.
4. ΔT is the vertical gap between the two lines at the mixing time, at the same moment.

This estimates the temperature the mixture would have reached if the reaction had been instant and no heat was lost.

7. Errors and their effects

Match each improvement to its error: a lid or insulation reduces heat loss; a draught shield reduces heat loss from the flame; extrapolation corrects for cooling; capping the burner reduces evaporation. Repeating the experiment improves reliability, but it does not remove a systematic error such as heat loss.

8. Uncertainty in ΔT

ΔT comes from two readings, so if each reading is ±0.1 °C, ΔT is ±0.2 °C.

For ΔT = 6.6 °C: (0.2 ÷ 6.6) × 100 = 3.0%. A bigger ΔT gives a smaller percentage uncertainty. If a question asks for the uncertainty in a product or quotient such as q = mcΔT, add the percentage uncertainties of the measured quantities.

ΔT is bigger when the same amount of reactant reacts completely in less water, or when more reactant reacts in the same volume (e.g. more concentrated solutions). Simply halving both solution volumes at the same concentration halves the heat and the mass together, so ΔT does not increase.

Quick checks

These are Finesse practice questions. The step-by-step answers are indicative worked solutions, not official AQA mark allocations.

Sources