Melting, boiling and evaporation Revision Notes

Matter exists in three common states: solid, liquid, and gas. Changing from one state to another is called a phase change. Two important phase changes are melting (solid to liquid) and boiling (liquid to gas). Both require energy input to occur.

During these phase changes, the temperature of the substance remains constant even though energy is being added. This is a key concept in thermal physics and is explained by the energy being used to change the internal structure of the substance rather than increasing kinetic energy of particles.

• Phase changes involve changes in the arrangement and energy of particles.
• Melting and boiling require energy input to overcome forces between particles.
• Temperature remains constant during these phase changes.

Melting occurs when a solid absorbs enough energy to change into a liquid. The energy supplied is called the latent heat of fusion.

In a solid, particles are closely packed in a fixed, orderly arrangement and vibrate about fixed positions. To melt, energy must be supplied to overcome the strong forces holding particles in place.

During melting, the temperature remains constant at the melting point because the energy input is used to break these intermolecular bonds rather than increase particle kinetic energy. Once all the solid has melted, further energy input will raise the temperature of the liquid.

• Latent heat of fusion is the energy required to melt 1 kg of a solid without temperature change (unit: J/kg).
• Particles gain potential energy to overcome forces but kinetic energy (temperature) stays constant.
• Melting point is the temperature at which solid and liquid coexist in equilibrium.

Boiling is the process where a liquid changes into a gas when it absorbs enough energy called the latent heat of vaporisation.

In a liquid, particles are close but can move past each other. To boil, energy is needed to overcome the intermolecular forces completely so particles can move freely as a gas.

During boiling, the temperature remains constant at the boiling point because the energy supplied breaks the intermolecular forces rather than increasing particle speed. After boiling, further energy increases the temperature of the gas.

• Latent heat of vaporisation is the energy required to boil 1 kg of a liquid without temperature change (unit: J/kg).
• Energy increases potential energy to separate particles fully, not kinetic energy.
• Boiling point is the temperature at which liquid and gas coexist in equilibrium under given pressure.

Temperature measures the average kinetic energy of particles. During melting and boiling, energy input does not increase kinetic energy but increases potential energy by breaking intermolecular bonds.

This explains why temperature remains constant during these phase changes despite continuous energy input.

The energy required per kilogram for melting or boiling is characteristic of the substance and is called latent heat. The total energy Q needed is calculated by: Q = mL, where m is mass (kg) and L is latent heat (J/kg).

• Temperature remains constant during melting and boiling because kinetic energy is unchanged.
• Energy input increases potential energy to change state.
• Q = mL calculates energy for phase change, where L is latent heat of fusion or vaporisation.

In solids, particles vibrate in fixed positions with strong forces between them.

In liquids, particles move randomly but remain close, with weaker forces than solids.

In gases, particles move freely and are far apart with negligible forces between them during boiling.

• Melting: particles gain energy to move from fixed positions to flow past each other.
• Boiling: particles gain enough energy to separate completely and move independently.
• Diagrams should show particles tightly packed in solids, close but mobile in liquids, and widely spaced in gases.

Matter can exist in different states: solid, liquid, and gas. When a substance changes from one state to another, it undergoes a phase change.

The melting point of a substance is the temperature at which it changes from solid to liquid at a given pressure. At this temperature, the solid and liquid phases coexist in equilibrium.

The boiling point of a substance is the temperature at which it changes from liquid to gas (boils) at a given pressure. At this temperature, the liquid and gas phases coexist in equilibrium.

• Melting point: temperature where solid turns to liquid.
• Boiling point: temperature where liquid turns to gas.

Water melts at 0 degrees Celsius (°C) under standard atmospheric pressure, which is defined as 1 atmosphere (atm) or 101.3 kilopascals (kPa).

At this temperature, ice (solid water) and liquid water coexist. Adding heat at 0°C causes ice to melt without increasing temperature until all ice has turned to liquid.

The melting point is a fixed physical property of pure water at this pressure.

• Melting point of water = 0°C (32°F) at 1 atm pressure.
• Temperature remains constant during melting until phase change completes.

Water boils at 100 degrees Celsius (°C) under standard atmospheric pressure (1 atm).

At this temperature, liquid water and water vapour (gas) coexist. Adding heat at 100°C causes water to boil without increasing temperature until all liquid has vaporised.

The boiling point depends on pressure; at lower pressures, boiling occurs at lower temperatures.

• Boiling point of water = 100°C (212°F) at 1 atm pressure.
• Temperature remains constant during boiling until phase change completes.

Standard atmospheric pressure is the average pressure at sea level, equal to 1 atm or 101.3 kPa.

Melting and boiling points are defined at this pressure because pressure affects these temperatures.

For example, at higher altitudes where pressure is lower, water boils at temperatures below 100°C.

• Standard atmospheric pressure = 1 atm = 101.3 kPa.
• Melting and boiling points vary with pressure.
• Standard conditions allow consistent reference values.

To measure melting or boiling points, pure water is heated or cooled under controlled pressure (1 atm).

Temperature is recorded using a calibrated thermometer as the substance changes state.

During melting or boiling, temperature remains constant despite heat input, indicating energy is used to change phase (latent heat).

• Use pure water and standard pressure for accurate measurements.
• Temperature remains constant during phase change.
• Thermometer must be accurate and calibrated.

Matter exists in different states: solid, liquid, and gas. Changes between these states are called phase changes and involve energy transfer and changes in particle behavior.

Condensation and solidification are two such phase changes where a substance loses energy and changes state: condensation is gas to liquid, and solidification is liquid to solid.

• Phase changes involve energy transfer without changing the chemical identity of the substance.
• Particle behavior—movement and arrangement—changes during phase changes.

Condensation occurs when a gas cools and changes into a liquid. As the gas loses thermal energy, the particles slow down and move closer together.

The decrease in kinetic energy allows intermolecular forces to pull particles closer, forming a liquid with a definite volume but no fixed shape.

Particles in the liquid phase are still able to move past each other but are much closer than in the gas phase.

• Gas particles have high kinetic energy and move freely and far apart.
• Cooling reduces particle kinetic energy, causing particles to come closer.
• Intermolecular forces become significant, leading to liquid formation.

Solidification, also called freezing, happens when a liquid loses enough energy to become a solid.

As the liquid cools, particles lose kinetic energy and move more slowly, allowing them to arrange into a fixed, regular pattern called a lattice.

In the solid state, particles vibrate about fixed positions and have a definite shape and volume.

• Liquid particles move randomly but are close together.
• Energy loss reduces particle motion, enabling fixed arrangement.
• Strong intermolecular forces hold particles in a rigid structure.

Both condensation and solidification involve the removal of thermal energy from the substance.

During condensation, energy is released as gas particles slow down and form liquid bonds.

During solidification, further energy is lost as particles settle into a solid lattice structure, releasing latent heat of fusion.

• Energy loss causes particles to slow and come closer.
• Latent heat is the energy released during phase changes without temperature change.
• Energy transfer explains why temperature remains constant during the phase change itself.

Condensation: gas particles lose energy, slow down, and come closer to form a liquid with weaker intermolecular bonds than a solid.

Solidification: liquid particles lose more energy, slow further, and arrange into a fixed, regular structure forming a solid.

Both processes illustrate how energy loss leads to decreased particle motion and increased particle attraction, changing the state of matter.

• Condensation: gas → liquid, energy lost, particles closer, more ordered.
• Solidification: liquid → solid, more energy lost, fixed particle arrangement.
• Understanding particle behavior explains macroscopic properties of states of matter.

Evaporation is a phase change where particles in a liquid gain enough energy to overcome intermolecular forces and escape into the air as gas particles.

Unlike boiling, evaporation occurs only at the surface of the liquid and can happen at any temperature, not just at the boiling point.

Particles in a liquid have a range of kinetic energies due to thermal motion; only those with higher-than-average energy can escape the surface.

• Evaporation happens at the liquid surface.
• Only particles with sufficient kinetic energy escape.
• It can occur at temperatures below boiling point.

In a liquid, particles move randomly with different speeds, described by a distribution of kinetic energies.

Some particles near the surface have enough energy to overcome the attractive forces binding them to other particles and escape into the air as vapour.

This selective escape of higher-energy particles reduces the average kinetic energy of the remaining liquid, causing cooling.

• Particles have a range of kinetic energies.
• Higher-energy particles escape during evaporation.
• Remaining liquid cools as energetic particles leave.

The rate of evaporation depends on several factors: temperature, surface area, and air movement.

Higher temperature increases the average kinetic energy of particles, so more particles have enough energy to evaporate.

A larger surface area exposes more particles to the surface, increasing evaporation rate as more particles can escape simultaneously.

• Increasing temperature raises evaporation rate.
• Larger surface area increases evaporation.
• Air movement removes vapour, enhancing evaporation.

Because the most energetic particles leave the liquid, the average kinetic energy of the remaining particles decreases, lowering the liquid's temperature.

This is why evaporation causes cooling, for example, when sweat evaporates from skin.

This cooling effect is important in natural and technological processes.

• Evaporation removes energetic particles.
• Average kinetic energy and temperature of liquid decrease.
• Evaporation causes cooling effects.

A simple diagram to draw shows a container with liquid and particles inside.

At the surface, some particles are shown with arrows pointing away, representing escape into the air.

Particles inside the liquid have varying speeds, indicated by different arrow lengths, showing energy distribution.

• Draw liquid surface with particles inside.
• Show some particles escaping from surface.
• Use arrows of different lengths to indicate particle speeds.

Evaporation is the process where molecules at the surface of a liquid gain enough kinetic energy to overcome the intermolecular forces holding them in the liquid and escape into the air as gas molecules.

Unlike boiling, evaporation can occur at any temperature and only involves molecules at the surface of the liquid.

The molecules that escape are those with the highest kinetic energy, leaving behind molecules with lower average kinetic energy.

• Evaporation occurs at the liquid surface.
• Only the most energetic molecules escape.
• Evaporation can happen below boiling point.

When molecules evaporate, they take energy with them in the form of latent heat of vaporisation.

This energy is absorbed from the liquid itself and its surroundings, causing the average kinetic energy of the remaining liquid molecules to decrease.

Since temperature is a measure of average kinetic energy, the temperature of the liquid falls as evaporation proceeds, resulting in cooling.

• Evaporation is an endothermic process.
• Energy is taken from the liquid and surroundings.
• Temperature of the liquid decreases.

The cooling effect is due to the loss of high-energy molecules, which reduces the average kinetic energy and thus the temperature of the liquid.

This effect can be observed when sweat evaporates from the skin, removing heat and cooling the body.

Similarly, evaporation is used in cooling systems such as evaporative coolers and in natural processes like transpiration in plants.

• Loss of energetic molecules lowers temperature.
• Sweating cools the body by evaporation.
• Evaporation is used in cooling technologies.

The rate of evaporation and thus the cooling effect depends on factors such as temperature, surface area, humidity, and air movement.

Higher temperature increases molecular energy, increasing evaporation rate.

Greater surface area allows more molecules to escape simultaneously, increasing cooling rate.

• Higher temperature increases evaporation rate.
• Larger surface area increases evaporation.
• Dry air and air movement enhance evaporation.

Evaporation is the process by which molecules at the surface of a liquid gain sufficient kinetic energy to overcome the intermolecular forces binding them and escape into the gas phase. This process occurs only at the surface of the liquid and can happen at any temperature, not necessarily at the boiling point.

Molecules in a liquid have a range of kinetic energies. Those with higher energy near the surface can break free and become vapour. Evaporation causes cooling of the remaining liquid because the molecules with the highest energy leave, lowering the average kinetic energy and thus temperature.

• Occurs only at the liquid surface.
• Can happen at any temperature.
• Molecules with sufficient energy escape into the air.
• Results in cooling of the liquid.

Boiling is a rapid phase change from liquid to gas that occurs throughout the entire volume of the liquid when it reaches a specific temperature called the boiling point. At this temperature, the vapour pressure of the liquid equals the atmospheric pressure, allowing bubbles of vapour to form inside the liquid and rise to the surface.

During boiling, bubbles of vapour form at nucleation sites within the liquid and grow as the liquid absorbs heat. The temperature remains constant during boiling until all the liquid has changed to gas.

• Occurs throughout the liquid, not just at the surface.
• Happens only at the boiling point temperature.
• Vapour bubbles form inside the liquid and rise to the surface.
• Temperature remains constant during boiling.

Evaporation can occur at any temperature because some molecules always have enough energy to escape the surface. The rate of evaporation increases with temperature, surface area, and air movement.

Boiling occurs only at a fixed temperature for a given pressure, known as the boiling point. For water at standard atmospheric pressure, this is 100°C. If pressure changes, the boiling point changes accordingly.

The boiling point is the temperature at which the vapour pressure of the liquid equals the external pressure.

• Evaporation occurs at all temperatures; boiling only at boiling point.
• Boiling point depends on external pressure.
• Evaporation rate increases with temperature, surface area, and airflow.
• Boiling point for water at 1 atm is 100°C.

Both evaporation and boiling require energy input to overcome intermolecular forces. This energy is called latent heat of vaporisation.

In evaporation, energy is absorbed by surface molecules individually, allowing them to escape. In boiling, energy is absorbed throughout the liquid, enabling bubble formation inside the liquid.

During boiling, temperature remains constant because energy is used to change the state, not to increase temperature.

• Energy input is latent heat of vaporisation.
• Evaporation absorbs energy from surroundings, causing cooling.
• Boiling uses energy to form vapour bubbles inside the liquid.
• Temperature remains constant during boiling despite energy input.

Evaporation is a slow, invisible process with no bubble formation. It is often observed as gradual drying or cooling.

Boiling is a rapid process with visible bubbles rising and breaking at the surface, producing vapour or steam.

In practical terms, boiling is used to sterilise water or cook food, while evaporation is important in processes like drying and cooling by sweat.

• Evaporation: no bubbles, slow, surface only.
• Boiling: bubbles form, rapid, throughout liquid.
• Boiling produces steam; evaporation produces vapour.
• Boiling is used in cooking; evaporation in cooling and drying.

Evaporation is the process by which molecules at the surface of a liquid gain enough kinetic energy to overcome intermolecular forces and escape into the gas phase.

Only molecules with higher-than-average kinetic energy can evaporate, leaving behind molecules with lower average kinetic energy, which causes cooling of the remaining liquid.

Evaporation occurs at all temperatures but is faster at higher temperatures due to increased molecular energy.

• Evaporation involves surface molecules escaping into the air.
• Molecules must have sufficient kinetic energy to evaporate.
• Evaporation causes cooling of the liquid.

Temperature is a measure of the average kinetic energy of molecules in a substance.

As temperature increases, more molecules have enough energy to escape from the liquid surface, increasing the evaporation rate.

This relationship means evaporation rate increases exponentially with temperature, as described by the Arrhenius equation in advanced contexts, but qualitatively it is sufficient to know higher temperature means faster evaporation.

• Higher temperature increases average molecular kinetic energy.
• More molecules can overcome intermolecular forces at higher temperatures.
• Evaporation rate increases as temperature rises.

Surface area is the area of the liquid exposed to air.

A larger surface area means more molecules are at the surface and able to escape into the air at any moment.

Therefore, increasing the surface area increases the number of molecules evaporating per second, raising the evaporation rate.

• Surface area is the exposed liquid surface.
• Larger surface area allows more molecules to evaporate simultaneously.
• Evaporation rate is directly proportional to surface area.

Air movement above the liquid surface removes water vapour molecules, reducing the concentration of vapour near the surface.

This reduction in vapour concentration decreases the rate of condensation back into the liquid, allowing more molecules to evaporate.

Stronger air currents increase evaporation rate by continuously replacing saturated air with drier air.

• Air movement removes vapour molecules from above the liquid.
• Lower vapour concentration near the surface increases net evaporation.
• Increased air movement leads to faster evaporation.

The rate of evaporation depends on temperature, surface area, and air movement.

In practical terms, evaporation can be increased by heating the liquid, spreading it over a larger area, or increasing airflow (e.g., using a fan).

Understanding these factors is important in processes such as drying clothes, cooling by evaporation, and industrial drying techniques.

• Evaporation rate ∝ temperature, surface area, and air movement.
• Practical methods to increase evaporation include heating, spreading, and fanning.
• Evaporation causes cooling due to loss of energetic molecules.

Evaporation is the process where molecules at the surface of a liquid gain enough kinetic energy to overcome intermolecular forces and escape into the gas phase.

This phase change from liquid to gas requires energy called the latent heat of vaporisation, which is absorbed from the surroundings.

When a liquid is in contact with an object, the liquid absorbs thermal energy from the object to supply this latent heat, causing the object to lose heat and cool down.

• Latent heat of vaporisation (L_v) is the energy needed to change 1 kg of liquid into gas without temperature change, measured in joules per kilogram (J/kg).
• Energy absorbed during evaporation = mass evaporated × latent heat of vaporisation (Q = m × L_v).
• The object’s temperature decreases because it loses thermal energy to the evaporating liquid.

When the liquid evaporates, it takes energy from the object’s thermal store, reducing the object’s internal energy and thus its temperature.

This cooling effect is commonly experienced when sweat evaporates from skin, removing heat and cooling the body.

The rate of cooling depends on how quickly evaporation occurs, which in turn depends on several factors.

• Cooling occurs because evaporation is an endothermic process requiring energy input from the object.
• The temperature drop in the object is proportional to the amount of liquid evaporated and the latent heat of vaporisation.
• Continuous evaporation maintains cooling as long as liquid and suitable conditions persist.

Several factors influence how fast evaporation and thus cooling occurs: temperature, surface area, and air movement.

Higher temperature increases the average kinetic energy of liquid molecules, increasing evaporation rate.

A larger surface area exposes more molecules to the air, increasing evaporation rate and cooling effect.

• Increased airflow removes vapour molecules from the surface, maintaining a concentration gradient that favours evaporation.
• Lower humidity in the surrounding air increases evaporation rate, enhancing cooling.
• If the liquid is saturated with vapour, evaporation slows and cooling reduces.

Molecules in a liquid have a range of kinetic energies; those with the highest energies escape first during evaporation.

Removing these higher-energy molecules lowers the average kinetic energy of the remaining molecules, which corresponds to a lower temperature.

This explains why evaporation cools both the liquid and any object in thermal contact with it.

• Temperature is a measure of average kinetic energy of particles.
• Evaporation selectively removes higher-energy molecules, reducing average kinetic energy.
• The object loses heat to supply energy for evaporation, causing its temperature to drop.

In practical situations, evaporation can be used for cooling, such as in sweat evaporation or cooling towers.

To observe cooling by evaporation, one can place a wet cloth on a warm object and feel the temperature drop as the water evaporates.

Safety considerations include avoiding flammable liquids and ensuring good ventilation when evaporation involves volatile substances.

• Use water or safe liquids for demonstration of evaporative cooling.
• Measure temperature changes with a thermometer before and after evaporation.
• Repeat measurements to improve accuracy and consider environmental factors like airflow and humidity.