Heat Capacity (VCE SSCE Chemistry): Revision Notes

Heat Capacity

Introduction

Water has a remarkably high heat capacity, which is one of the key properties that makes it essential for life on Earth. Heat capacity refers to how much heat energy a substance can absorb before its temperature changes significantly. Water can soak up large amounts of heat energy whilst experiencing only small temperature increases.

You've probably noticed this property at the beach on a hot summer's day. The dry sand can be scorching hot under your feet, whilst the shallow water at the sea's edge remains pleasantly cool. Both the sand and water receive similar amounts of heat energy from the Sun, but they respond very differently because they have different heat capacities.

Comparing the heat capacity of different substances

Heat capacity measures a substance's ability to absorb and store heat energy. When you apply the same amount of heat energy to two different substances, they will undergo different temperature changes depending on their heat capacities.

We can demonstrate this in the laboratory by comparing water with ethanediol (a liquid used as coolant in car radiators). Water has a higher heat capacity than ethanediol. When we heat $100 \text{ g}$ of each liquid with identical amounts of energy, the water's temperature increases by a smaller amount than the ethanediol's temperature. This is shown in the experimental setup below.

Specific heat capacity

Definition and units

Specific heat capacity $(c)$ is a more precise way of measuring how substances respond to heat energy. It tells us exactly how much energy (measured in joules) is needed to raise the temperature of $1 \text{ g}$ of a substance by $1°\text{C}$.

The units for specific heat capacity are joules per gram per degree Celsius: $\text{J g}^{-1} °\text{C}^{-1}$

Heat capacity values for common substances

Different materials have very different specific heat capacities, which reflects their internal structure and bonding. Here are some examples:

Metals:

• Iron: $0.45 \text{ J g}^{-1} °\text{C}^{-1}$
• Copper: $0.39 \text{ J g}^{-1} °\text{C}^{-1}$
• Aluminium: $0.90 \text{ J g}^{-1} °\text{C}^{-1}$

Molecular substances:

• Water: $4.18 \text{ J g}^{-1} °\text{C}^{-1}$
• Ethanediol (antifreeze): $2.42 \text{ J g}^{-1} °\text{C}^{-1}$
• Chlorofluorocarbon (CCl₂F₂): $0.60 \text{ J g}^{-1} °\text{C}^{-1}$

Composite materials:

• Wood: $1.3-2.0 \text{ J g}^{-1} °\text{C}^{-1}$
• Concrete: $0.8 \text{ J g}^{-1} °\text{C}^{-1}$
• Glass: $0.84 \text{ J g}^{-1} °\text{C}^{-1}$

Why does water have such a high heat capacity?

The specific heat capacity of a substance depends on the types of bonds holding its molecules, ions, or atoms together. For molecular substances like water, the heat capacity depends on the strength of the intermolecular forces between molecules.

Water's unusually high specific heat capacity is due to the extensive network of hydrogen bonds between water molecules. Hydrogen bonds are stronger than other types of intermolecular forces. When heat energy is added to water, much of that energy is used to stretch and vibrate these hydrogen bonds rather than increasing the kinetic energy of the molecules (which would raise the temperature).

Calculations using specific heat capacity

The heat capacity formula

Heat energy is represented by the symbol $q$. We can calculate the heat energy needed to change the temperature of a substance using this formula:

$\text{Heat energy} = \text{mass} \times \text{specific heat capacity} \times \text{temperature change}$

Using symbols:

$q = m \times c \times \Delta T$

where:

• $q$ = amount of heat energy in joules (J)
• $m$ = mass in grams (g)
• $c$ = specific heat capacity ($\text{J g}^{-1} °\text{C}^{-1}$)
• $\Delta T$ = temperature change in degrees Celsius (°C)

Working with water volumes

Because water is a liquid, we often measure it by volume (in mL or L) rather than by mass. Fortunately, water has a convenient property that makes conversions easy:

$1 \text{ mL of water} = \text{approximately } 1 \text{ g}$

This means:

• $100 \text{ mL of water} \approx 100 \text{ g}$
• $250 \text{ mL of water} \approx 250 \text{ g}$
• $1 \text{ L of water} \approx 1000 \text{ g}$

Worked example

Question: Calculate the heat energy, in kJ, needed to increase the temperature of $200 \text{ g}$ of water by $15.0°\text{C}$.

Practical applications

Water as a coolant in car engines

Water's high heat capacity makes it an excellent coolant for car engines. When fuel combusts in an engine, it releases enormous amounts of heat energy through a highly exothermic reaction. If this heat isn't removed, the engine will overheat and suffer permanent damage.

Water circulates through the engine via a system of pipes, absorbing excess heat energy. Because of its high heat capacity, water can absorb large amounts of this heat without its own temperature rising too much. The heated water then flows to the radiator, where the heat energy transfers from the water to the metal radiator and then to the air passing through the vehicle's grill.

Remember!