Calorimetry experiments

Calorimetry is a way of finding out how much energy a reaction releases or absorbs without measuring the energy directly. Instead, you measure something easy — a temperature change.

The idea is simple:

• An exothermic reaction releases heat → the heat warms up the water (or solution) → the temperature rises.
• An endothermic reaction absorbs heat → it takes heat from the water → the temperature falls.

So by measuring the temperature change (ΔT) of a known amount of water, you can tell whether a reaction gives out or takes in energy, and compare different reactions.

Two standard set-ups appear in Double Award Chemistry:

Reaction type	Calorimeter	Why
Combustion of a fuel	metal can of water above a burner	a metal can conducts the flame's heat efficiently into the water
Neutralisation / displacement / dissolving	polystyrene cup	polystyrene insulates — it reduces heat loss to the surroundings

The thing you actually measure is always the temperature of the water/solution — never the temperature of the flame or the reactants themselves.

To compare fuels (or find how much energy a fuel gives out), you burn a measured amount of it under a known mass of water and measure how much the water temperature rises.

Apparatus: a spirit burner containing the fuel, a metal (e.g. copper) can holding the water, a thermometer, a balance, and ideally a draught shield.

Method (combustion):

1. Measure a known volume/mass of water into the can (e.g. 100 g).
2. Record the starting temperature of the water.
3. Weigh the spirit burner + fuel.
4. Light the burner and let it heat the water.
5. Stir, then record the highest temperature reached → work out ΔT (rise).
6. Re-weigh the burner to find the mass of fuel burned.

What this tells you: a bigger temperature rise (for the same mass of water and fuel) means the fuel released more energy. This is how you fairly compare the energy given out by different fuels.

For reactions that happen in solution — neutralisation, displacement and dissolving — you do not need a flame. You simply mix the reactants in an insulated cup and watch the temperature change.

Apparatus: a polystyrene (expanded plastic) cup, usually sitting in a beaker for support, with a lid and a thermometer.

Method (solution reactions):

1. Measure the first reactant into the cup (e.g. 50 cm³ of acid) and record its starting temperature.
2. Add the second reactant (e.g. 50 cm³ of alkali, or a measured mass of solid for dissolving).
3. Stir with the thermometer.
4. Record the highest temperature if it warms up (exothermic) or the lowest temperature if it cools down (endothermic) → find ΔT.

Typical examples:

• Neutralisation: acid + alkali → temperature rises (exothermic).
• Displacement: a more reactive metal pushes out a less reactive one, e.g. zinc added to copper sulfate solution → temperature rises (exothermic).
• Dissolving: some solids warm the water; others, such as ammonium nitrate, make it colder (endothermic) — this is how instant cold packs work.

The polystyrene cup matters: because it insulates, very little of the reaction's heat leaks away, so your temperature change is more accurate.

When you compare your experimental result with the data-book (true) value, your value is almost always smaller — the temperature rise you measured is less than it should have been. This is the single most important evaluation point in calorimetry.

Why? Heat is lost (and not all goes into the water):

• Heat lost to the surroundings — to the air and the apparatus (the can, cup, thermometer, stand). This is the biggest cause, especially in combustion where lots of heat escapes around the can.
• Incomplete combustion — if the fuel does not burn completely, it releases less energy than it should (and may leave soot).
• Evaporation of water from an open cup carries energy away.

Because energy escapes, less heat reaches the water, the temperature rise is smaller, and so the calculated energy comes out lower than the data-book value.

How to reduce heat loss (improvements):

• Put a lid on the cup/can (cuts heat loss and evaporation).
• Insulate the calorimeter (e.g. lag it with cotton wool; use a polystyrene cup, not glass — glass conducts heat away).
• Use a draught shield around the flame so heat is not blown away.
• For combustion, position the can close to the flame and shield it so more of the heat reaches the water.

Exam phrasing that scores: "Heat is lost to the surroundings, so less heat warms the water, the temperature rise is smaller, and the calculated energy is lower than the data-book value." Vague answers like "the experiment wasn't accurate" earn nothing.