Table of Contents
Latent Heat and Changes of State
1. Introduction
Matter exists in different physical states, principally solid, liquid and gas. Under suitable conditions, a substance can change from one state to another.
These changes are called changes of state or phase changes.
During a change of state, a substance may absorb or release energy without experiencing a temperature change, provided the process occurs at a constant pressure under equilibrium conditions.
The energy involved in such a process is called latent heat.
Latent heat is important in refrigeration, air conditioning, weather systems, cooking, industrial processing and many natural phenomena.
2. States of Matter
The three common states of matter are solid, liquid and gas.
2.1 Solid
A solid has:
- A definite shape.
- A definite volume.
- Particles closely packed together.
- Particles that generally vibrate around fixed equilibrium positions.
Examples include ice, iron, wood and salt.
2.2 Liquid
A liquid has:
- A definite volume.
- No fixed shape.
- The ability to flow.
- Particles that remain relatively close but can move past one another.
Examples include water, oil and mercury.
2.3 Gas
A gas has:
- No definite shape.
- No definite volume.
- Particles that move freely.
- The ability to expand and fill its container.
Examples include oxygen, nitrogen and carbon dioxide.
2.4 Plasma
Plasma is an ionised state of matter containing free electrons and ions.
Examples include:
- The Sun and other stars.
- Lightning.
- Certain types of fluorescent and plasma discharge.
- Some laboratory fusion experiments.
3. Changes of State
A change of state occurs when a substance changes from one physical state to another due to changes in temperature, pressure or both.
The principal changes are:
| Change | Process |
|---|---|
| Solid to liquid | Melting or fusion |
| Liquid to solid | Freezing or solidification |
| Liquid to gas | Vaporisation |
| Gas to liquid | Condensation |
| Solid to gas | Sublimation |
| Gas to solid | Deposition |
4. Melting or Fusion
Melting is the process in which a solid changes into a liquid after absorbing sufficient energy.
For a pure crystalline substance at a specified pressure, melting occurs at a characteristic melting point.
For example, pure ice melts at approximately 0°C under standard atmospheric pressure.
During melting:
- Energy is absorbed.
- The solid structure breaks down.
- Particles gain greater freedom of movement.
- The temperature remains constant during the equilibrium phase transition.
4.1 Melting Point
The melting point is the temperature at which a solid changes into a liquid at a specified pressure.
For pure ice at standard atmospheric pressure:
[
T_m=0^\circ\text{C}
]
5. Freezing or Solidification
Freezing is the process in which a liquid changes into a solid.
During freezing:
- Energy is released to the surroundings.
- Particle motion becomes more restricted.
- The liquid develops a solid structure.
- The temperature remains constant during the phase change under equilibrium conditions.
For pure water at standard atmospheric pressure, freezing occurs at approximately 0°C.
Melting and freezing are reverse processes.
6. Vaporisation
Vaporisation is the process in which a liquid changes into a gas.
It occurs through two principal mechanisms:
- Evaporation.
- Boiling.
6.1 Evaporation
Evaporation is the conversion of a liquid into vapour from its surface.
It can occur at temperatures below the boiling point.
Some surface molecules possess enough kinetic energy to escape from the liquid into the surrounding gas.
Important characteristics include:
- It occurs at the liquid surface.
- It can occur at many temperatures.
- It generally causes cooling.
- Its rate increases with temperature.
- Its rate is influenced by humidity, wind speed and surface area.
6.2 Boiling
Boiling is the rapid formation of vapour throughout a liquid when its vapour pressure becomes equal to the external pressure.
At the boiling point, vapour bubbles can form within the liquid and rise to the surface.
For pure water at standard atmospheric pressure:
[
T_b=100^\circ\text{C}
]
The boiling point changes with external pressure.
7. Difference Between Evaporation and Boiling
| Feature | Evaporation | Boiling |
|---|---|---|
| Location | Surface only | Throughout the liquid |
| Temperature | Can occur at many temperatures | Occurs at the boiling point for the specified pressure |
| Bubble formation | No bulk bubble formation required | Vapour bubbles form within the liquid |
| Rate | Usually gradual | Generally rapid |
| Cooling effect | Commonly produces cooling | Heat must be supplied to sustain boiling |
| Pressure relationship | Occurs below boiling conditions too | Occurs when vapour pressure equals external pressure |
7.1 Why Does Evaporation Cause Cooling?
During evaporation, molecules with relatively high kinetic energy escape from the liquid surface.
The average kinetic energy of the remaining liquid molecules decreases, resulting in a reduction in temperature if sufficient heat is not supplied from elsewhere.
This is why sweating cools the human body.
8. Condensation
Condensation is the conversion of a gas or vapour into a liquid.
During condensation:
- Energy is released.
- Particles move closer together.
- Intermolecular attractions become more influential.
- The substance changes from the gaseous state to the liquid state.
Examples include:
- Water droplets forming on a cold glass.
- Dew forming on grass.
- Clouds developing through condensation of atmospheric water vapour.
- Steam forming water droplets on a cool surface.
Condensation is the reverse of vaporisation.
9. Sublimation
Sublimation is the direct conversion of a solid into a gas without passing through the liquid state.
Examples include:
- Dry ice, which is solid carbon dioxide.
- Naphthalene under suitable conditions.
- Certain forms of iodine under laboratory conditions.
Sublimation is useful in specialised drying and purification processes.
9.1 Dry Ice
Dry ice is solid carbon dioxide.
At ordinary atmospheric pressure, it changes directly from solid to gas at approximately −78.5°C.
It is widely used for cooling and transporting temperature-sensitive materials.
10. Deposition
Deposition is the direct conversion of a gas into a solid without passing through the liquid state.
It is the reverse of sublimation.
Examples include:
- Formation of frost from water vapour.
- Formation of ice crystals under suitable atmospheric conditions.
11. Latent Heat
Latent heat is the energy absorbed or released by a substance during a change of state without a change in temperature, under constant-pressure equilibrium conditions.
The word latent means hidden.
The energy involved changes the arrangement and intermolecular potential energy of particles rather than increasing their average kinetic energy in the usual way.
12. Specific Latent Heat
Specific latent heat is the amount of energy required to change the state of one kilogram of a substance without changing its temperature.
It is represented by (L).
The formula is:
[
L=\frac{Q}{m}
]
Therefore:
[
Q=mL
]
Where:
- (Q) = heat energy transferred in joules
- (m) = mass in kilograms
- (L) = specific latent heat
The SI unit is:
[
\text{J kg}^{-1}
]
13. Types of Specific Latent Heat
Two important types are:
- Specific latent heat of fusion.
- Specific latent heat of vaporisation.
13.1 Specific Latent Heat of Fusion
Specific latent heat of fusion is the energy required to convert one kilogram of a solid into liquid at its melting point without a temperature change.
It is represented by (L_f).
For water:
[
L_f\approx3.34\times10^5\ \text{J kg}^{-1}
]
The same amount of energy per kilogram is released when water freezes at the same pressure and temperature.
13.2 Specific Latent Heat of Vaporisation
Specific latent heat of vaporisation is the energy required to convert one kilogram of a liquid into vapour at its boiling point without a temperature change.
It is represented by (L_v).
For water at approximately 100°C and standard atmospheric pressure:
[
L_v\approx2.26\times10^6\ \text{J kg}^{-1}
]
The same amount of energy per kilogram is released during condensation under corresponding conditions.
13.3 Comparison
| Property | Latent heat of fusion | Latent heat of vaporisation |
|---|---|---|
| Symbol | (L_f) | (L_v) |
| Process | Solid to liquid | Liquid to gas |
| Reverse process | Freezing | Condensation |
| Approximate value for water | (3.34\times10^5) J kg(^{-1}) | (2.26\times10^6) J kg(^{-1}) |
| Energy requirement | Lower for water | Higher for water |
These values are approximate and depend on pressure and other physical conditions.
14. Why Does Temperature Remain Constant During a Change of State?
When a substance is heated, its particles generally gain kinetic energy and its temperature increases.
However, during a phase change, energy is used primarily to alter the arrangement of particles and overcome or modify intermolecular forces.
For example, when ice melts:
- Energy is supplied.
- The solid structure gradually breaks down.
- The energy increases the system’s microscopic potential energy.
- The temperature remains approximately constant during the melting process.
After all the ice has melted, additional energy supplied to the water increases its temperature.
The reverse occurs during freezing and condensation, when energy is released.
15. Heating Curve of Water
A heating curve shows how the temperature of a substance changes as heat is supplied.
For ice initially below 0°C and heated at standard atmospheric pressure, the process can be divided into five stages.
Stage 1: Heating Ice
The temperature of ice increases until it reaches its melting point.
Stage 2: Melting
The ice changes into liquid water.
The temperature remains approximately 0°C while melting continues.
Stage 3: Heating Water
After all the ice has melted, the water temperature increases from 0°C to 100°C.
Stage 4: Boiling
Water changes into steam at approximately 100°C.
The temperature remains approximately constant while boiling continues at standard atmospheric pressure.
Stage 5: Heating Steam
After all the liquid water has vaporised, further heating increases the temperature of the steam.
15.1 Interpretation of the Heating Curve
The sloping sections of the curve represent temperature increases.
The horizontal sections represent changes of state during which energy is absorbed without a temperature increase.
The same principles apply in reverse during cooling.
16. Effect of Pressure on Boiling Point
The boiling point of a liquid depends on external pressure.
A liquid boils when its vapour pressure equals the surrounding pressure.
16.1 High Altitudes
Atmospheric pressure decreases with increasing altitude.
As a result, water boils at a temperature lower than 100°C in many high-altitude locations.
Food may therefore take longer to cook in boiling water because the water temperature is lower.
16.2 Pressure Cooker
A pressure cooker increases the pressure above the cooking liquid.
This raises the boiling point of water, allowing food to cook at a temperature above 100°C.
16.3 Vacuum Distillation
Reducing pressure lowers the boiling point of liquids.
This principle is used in vacuum distillation to separate or purify substances that may decompose at high temperatures.
17. Factors Affecting Evaporation
17.1 Temperature
Higher temperature increases the average kinetic energy of molecules and generally increases the evaporation rate.
17.2 Surface Area
A larger exposed surface allows more molecules to escape at a given time.
17.3 Humidity
High humidity reduces the rate of evaporation because the surrounding air already contains a large amount of water vapour.
17.4 Wind Speed
Moving air removes vapour from the surface and generally increases evaporation.
17.5 Nature of the Liquid
Different liquids have different intermolecular forces and vapour pressures. More volatile liquids generally evaporate more readily.
18. Applications of Latent Heat
18.1 Refrigeration
Refrigerators use the evaporation of a refrigerant to absorb heat from the refrigerated space.
The refrigerant later releases heat during condensation in the external part of the system.
18.2 Air Conditioning
Air-conditioning systems use phase changes of refrigerants to transfer heat from indoor spaces to the outside environment.
18.3 Sweating
Evaporation of sweat removes energy from the skin and helps regulate body temperature.
18.4 Steam Heating
Steam releases substantial energy when it condenses into liquid water.
This property makes steam useful in heating systems and industrial processes.
18.5 Ice for Cooling
Ice absorbs energy from its surroundings when it melts.
This makes ice useful for cooling drinks, food and other materials.
18.6 Freeze Drying
Freeze drying involves freezing a material and removing ice through sublimation under reduced pressure.
It is used in food preservation, pharmaceuticals and biological materials.
18.7 Distillation
Changes between liquid and vapour states are used to separate substances according to their volatility and boiling characteristics.
19. Worked Examples
Example 1: Energy Required to Melt Ice
Calculate the energy required to melt 0.5 kg of ice at 0°C into water at 0°C.
Take:
[
L_f=3.34\times10^5\ \text{J kg}^{-1}
]
Given:
[
m=0.5\ \text{kg}
]
[
L_f=3.34\times10^5\ \text{J kg}^{-1}
]
Using:
[
Q=mL_f
]
[
Q=0.5\times3.34\times10^5
]
[
\boxed{Q=167,000\ \text{J}}
]
The energy required is 167 kJ.
Example 2: Energy Required to Vaporise Water
Calculate the energy required to convert 0.1 kg of water at 100°C into steam at 100°C.
Take:
[
L_v=2.26\times10^6\ \text{J kg}^{-1}
]
Using:
[
Q=mL_v
]
[
Q=0.1\times2.26\times10^6
]
[
\boxed{Q=226,000\ \text{J}}
]
The energy required is 226 kJ.
Example 3: Heat Required to Melt Ice and Then Warm the Water
Calculate the total heat required to convert 0.2 kg of ice at 0°C into water at 20°C.
Take:
[
L_f=3.34\times10^5\ \text{J kg}^{-1}
]
[
c_{\text{water}}=4200\ \text{J kg}^{-1}\text{K}^{-1}
]
Step 1: Heat required to melt the ice
[
Q_1=mL_f
]
[
Q_1=0.2\times334000
]
[
Q_1=66800\ \text{J}
]
Step 2: Heat required to warm the water
[
Q_2=mc\Delta T
]
[
Q_2=0.2\times4200\times20
]
[
Q_2=16800\ \text{J}
]
Step 3: Total heat
[
Q_{\text{total}}=Q_1+Q_2
]
[
Q_{\text{total}}=66800+16800
]
[
\boxed{Q_{\text{total}}=83600\ \text{J}}
]
The total energy required is 83.6 kJ.
20. Important Examination Distinctions
| Concept | Key distinction |
|---|---|
| Melting | Solid changes into liquid |
| Freezing | Liquid changes into solid |
| Vaporisation | Liquid changes into gas |
| Condensation | Gas changes into liquid |
| Evaporation | Surface conversion of liquid into vapour |
| Boiling | Rapid vaporisation throughout a liquid |
| Sublimation | Solid changes directly into gas |
| Deposition | Gas changes directly into solid |
| Latent heat | Energy transferred during a phase change without temperature change |
| Specific latent heat | Latent heat per unit mass |
36 Key Takeaways for Exams
- Matter commonly exists as solid, liquid and gas.
- Plasma is an ionised state of matter.
- A change of state is also called a phase change.
- Melting is the conversion of a solid into a liquid.
- Freezing is the conversion of a liquid into a solid.
- Vaporisation is the conversion of a liquid into a gas.
- Condensation is the conversion of a gas into a liquid.
- Sublimation is the direct conversion of a solid into a gas.
- Deposition is the direct conversion of a gas into a solid.
- Melting and freezing are reverse processes.
- Vaporisation and condensation are reverse processes.
- Sublimation and deposition are reverse processes.
- Evaporation occurs at the surface of a liquid.
- Evaporation can occur below the boiling point.
- Boiling occurs throughout a liquid.
- A liquid boils when its vapour pressure equals the external pressure.
- Evaporation generally causes cooling.
- Higher temperature generally increases the rate of evaporation.
- Greater surface area generally increases evaporation.
- High humidity generally reduces evaporation.
- Increased wind speed generally increases evaporation.
- Latent heat is energy absorbed or released during a change of state without a temperature change under equilibrium conditions.
- Specific latent heat is the energy required to change the state of one kilogram of a substance.
- The formula for latent heat is (Q=mL).
- The SI unit of specific latent heat is J kg(^{-1}).
- Specific latent heat of fusion is represented by (L_f).
- Specific latent heat of vaporisation is represented by (L_v).
- The approximate latent heat of fusion of ice is (3.34\times10^5) J kg(^{-1}).
- The approximate latent heat of vaporisation of water at 100°C is (2.26\times10^6) J kg(^{-1}).
- Water’s latent heat of vaporisation is greater than its latent heat of fusion.
- Temperature remains approximately constant during a phase change of a pure substance at constant pressure under equilibrium conditions.
- During melting, supplied energy changes the microscopic arrangement and potential energy of the substance.
- The boiling point of water decreases at high altitudes because atmospheric pressure is lower.
- A pressure cooker raises the boiling point of water by increasing pressure.
- Refrigeration systems use evaporation and condensation of refrigerants to transfer heat.
- Sweat cools the body through evaporation, while steam can cause severe burns because of the energy released during condensation.