Heat Capacity and Calorimetry

Full Notes

Heat Transfer and Specific Heat

When a substance is heated or cooled, its temperature changes in proportion to the heat added or removed, the amount of substance, and the material's ability to absorb energy.

The equation for heat transfer is:
q = mcΔT

• q is the heat absorbed (+) or released (−) in joules (J)
• m is the mass of the substance in grams (g)
• c is the specific heat capacity in J/g·°C
• ΔT is the temperature change (final − initial)

A larger specific heat capacity means the substance resists temperature change – it needs more energy to heat up or cool down.

ExampleWater has a much higher specific heat capacity than copper. This means that, for equal masses, copper will heat up more quickly than water when the same amount of energy is applied.

Calorimetry

Calorimetry is the experimental technique used to measure heat transfer.

During an experiment, the temperature change of a reaction’s surroundings (e.g., a solution) is measured.

• If temperature increases, the process is exothermic (releasing heat).
• If temperature decreases, the process is endothermic (absorbing heat).

Matt’s exam tip

For calculations, assume no heat is lost to the surroundings unless told otherwise. All heat lost by one part of the system is gained by another.

First Law of Thermodynamics

The first law of thermodynamics states that energy is conserved. This means energy cannot be created or destroyed — it can only be transferred or transformed.

Heat lost by one substance = heat gained by another (in a closed calorimeter system).

Specific vs. Molar Heat Capacity

• Specific heat capacity: energy needed to raise 1 g of a substance by 1°C.
• Molar heat capacity: energy needed to raise 1 mol of a substance by 1°C.

Units:

• Specific heat (c): J/g·°C
• Molar heat capacity: J/mol·°C

Be careful to match your units to the type of heat capacity you're using.

Dissolution and Heat Flow

In calorimetry experiments involving dissolution:

• If temperature of the solution increases, the process is exothermic.
• If temperature decreases, the process is endothermic.

This tells us whether the solute releases or absorbs heat during dissolving. See examples below

Examples of Calorimetry

Combustion of Ethanol

Method:

1. Measure a known volume of water in a calorimeter (beaker or copper can).
2. Record the starting temperature of the water.
3. Weigh the spirit burner containing the fuel.
4. Light the burner and allow it to heat the water.
5. Stir and measure the final temperature of the water.
6. Reweigh the burner to determine mass of fuel burned.
7. Calculate q using q = mcΔT

Sources of Error:

• Heat loss to surroundings (e.g., air, beaker).
• Incomplete combustion (producing CO instead of CO₂).
• Evaporation of fuel from the wick.

Neutralisation

Method:

1. Use a polystyrene cup (to reduce heat loss).
2. Add a known volume of acid and record the starting temperature.
3. Add a known volume of alkali, stir, and record the maximum temperature.
4. Use q = mcΔT to calculate heat energy change.

Sources of Error:

• Heat loss to surroundings.
• Assumption that the solution has the same specific heat capacity as water.

Measuring Energy Change of Dissolution

Method:

1. Add a known mass of solute to a known volume of water in a polystyrene cup.
2. Stir and record the temperature change.
3. Use q = mcΔT to calculate the heat energy change.

Sources of Error:

• Heat loss to surroundings.
• Incomplete dissolution of solute.

Summary

Heat transfer during heating or cooling can be calculated using the equation q = mcΔT, where q is the heat exchanged, m is mass, c is specific heat, and ΔT is the temperature change.

Calorimetry allows us to measure this heat exchange in experiments. The specific heat capacity determines how much energy is needed for a substance to change temperature.

Energy is conserved in all processes, as described by the first law of thermodynamics, and the direction of heat flow can reveal whether a process is endothermic or exothermic.