Calorimetry

The enthalpy change for a reaction (ΔH) can be calculated by measuring the amount of heat energy given out or taken in during a chemical reaction. Recall that exothermic reactions release heat energy to the surroundings and endothermic reactions remove heat energy from the surroundings, or we can say that in an exothermic reaction the system (reactants and products) loses energy and in an endothermic reaction the system gains energy. The amount of heat energy given out/taken in by a chemical reaction can be determined by measuring a change in temperature (ΔT) in a substance, usually water. The measurement of heat flow is called calorimetry and the instrument used to measure it is called a calorimeter.

Specific heat capacity

When different substances are heated the particles that make them up gain kinetic energy; this results in a temperature rise. However when different substances are heated or supplied with the same amount of heat energy they change temperature by different amounts. The temperature change that a substance undergoes when it is heated is determined by its specific heat capacity (symbol C). The specific heat capacity of a substance is defined as the amount of heat energy needed to raise the temperature of 1 gram of the particular substance by 1 degree Kelvin. Its units are Joules per gram per Kelvin (Jg^-1K^-1 or kJkg^-1K^-1). The specific heat capacity of water is 4.18 Jg^-1K^-1, this means it takes 4.18 Joules of heat energy to raise the temperature of 1g of water by 1 degree Kelvin or 1 degree Celsius. We could of course have units of kJ K^-1kg^-1 which would mean that the mass of the substance would have to be expressed in kilograms (kg) and the energy in kilojoules (kJ).

Enthalpy of combustion

A simple calorimeter used to measure the rise in temperature of water when a fuel such as alcohol is burned.

Burning fuels is an exothermic reaction. We can use a simple calorimeter like the one shown opposite to calculate the enthalpy of combustion of a fuel such as ethanol or methanol. We can use the equation below to calculate the amount of heat energy that is transferred to the water from the burning fuel:

q = m x c x ΔT

Where:

q = amount of heat energy taken in by the water (kJ).

m = mass of water in kilograms (kg).

c= specific heat capacity of water in kJkg^-1K^-1

ΔT= change in temperature in degrees Kelvin, but since 1k = 1⁰C then it can be in either without any error in your calculations.

Example 1

1g of ethanol (C₂H₅OH) was burned in a spirit burner. The burner was used to heat 100g of water in a beaker. The initial temperature of the water was 25⁰C and the final temperature was 70⁰C. Calculate the enthalpy change for this combustion reaction.
To calculate the enthalpy change we simply substitute the values into the above equation. Care must be taken with units as this will be the most likely cause of an error in your calculations. We will assume that all the heat energy from the burning ethanol is transferred to the water. So we have:

q = m x c x ΔT
= 0.1kg of water x 4.2 kJkg^-1K^-1 x 45K (the change in temperature)
= 0.1 x 4.18 x 45
= -18.8 kJ (negative sign indicates that the reaction is exothermic)

If all the heat energy lost by the burning ethanol is transferred to the water then the ethanol loses 18.8 kJ of energy and the water gains 18.8kJ of energy. Recall the first law of thermodynamics which states that energy cannot be created or destroyed. This requires then that all the energy lost by the burning fuel is transferred to the water. In reality this is very unlikely as a lot of heat energy will be lost to the surroundings.

We can extend this calculation to work out the enthalpy change per mole of ethanol by simply using the formula:

The M_r of ethanol (C₂H₅OH) is 46, so the mass of 1 mole of ethanol is 46g. In the example above 1g of ethanol was burned and this gave an enthalpy change of -18.9kJ.
So by using the formula above we can easily calculate the number of moles of ethanol present in 1g:

number of moles of ethanol = 1/46 = 0.021 mol.

To calculate the molar enthalpy change (heat energy released by burning 1 mole of ethanol) we use the formula:

formula for molar enthalpy change
A calorimter where steps have been taken to reduce heat loss to the surroundings. The actual enthalpy of combustion of ethanol is -1367 kJ mol^-1

The fact that this is considerably higher than the enthalpy change calculated above should not be a total surprise since I am sure you can spot many ways in which heat from the burning alcohol is lost and not transferred into the water. We have assumed that all the heat energy from the burning alcohol flame has been transferred into the water. This is clearly not the case. The main area of heat loss is likely to be:

Heat loss to the surroundings

However other areas of potential heat loss are:

Evaporation of water from the beaker.
Heat from the burning alcohol flame is used to heat the beaker and the calorimeter not the water.
Incomplete combustion of the alcohol. Spirit burners are not an efficient way to completely burn the alcohol fuel.
The value given for the molar enthalpy of combustion of ethanol (-1367 kJ mol^-1) is under standard conditions. It is unlikely that this reaction was carried out under standard conditions. The temperature is likely to be the main culprit here.

However there are a number of changes that can be made to this simple calorimeter to help reduce heat loss and improve the accuracy of the calculated enthalpy change. These changes are summarised in the diagram opposite:

Key Points

A simple calorimeter can be used to measure the enthalpy of combustion of a fuel
We use the formula:
q = m x c x ΔT to calculate the enthalpy change.
The calculated enthalpy changes are likely to be much less than the actual enthalpy changes due to heat loss, mainly to the surroundings