Solids, liquids and gases Revision Notes

Solids have a fixed, definite shape and volume. This means a solid object maintains its shape and size regardless of the container it is placed in or external forces applied, provided the forces do not deform or break it.

For example, a wooden block remains the same shape and volume whether placed on a table or inside a box.

• Definite shape: solids do not change shape to fit their container.
• Definite volume: solids occupy a fixed amount of space.

Liquids have a definite volume but no fixed shape. They take the shape of the container they occupy, filling the bottom part of the container due to gravity.

For instance, water poured into a glass takes the shape of the glass but maintains the same volume regardless of the container shape.

• No definite shape: liquids adapt to the shape of their container.
• Definite volume: liquids occupy a fixed amount of space and are not easily compressed.

Gases have neither a definite shape nor a definite volume. They expand to fill the entire volume of their container, regardless of its size or shape.

For example, air in a balloon expands to fill the balloon completely, changing shape and volume as the balloon inflates or deflates.

• No definite shape: gases fill the shape of their container completely.
• No definite volume: gases expand or compress to fill the container volume.

Solids do not flow; they retain their shape and resist deformation under normal conditions.

Liquids flow easily, allowing them to be poured and to take the shape of their container.

Gases also flow and diffuse rapidly, spreading out to fill any available space.

• Solids: rigid, no flow.
• Liquids: flow readily, take container shape.
• Gases: flow and diffuse rapidly, fill container volume.

Solids and liquids are practically incompressible under normal conditions, meaning their volume does not change significantly when pressure is applied.

Gases are highly compressible; their volume decreases significantly when pressure increases, which is observable when pumping air into a bicycle tire or compressing air in a syringe.

• Solids: incompressible.
• Liquids: nearly incompressible.
• Gases: highly compressible.

In solids, particles (atoms, ions or molecules) are tightly packed in a regular, fixed arrangement. This close packing means particles are held in fixed positions by strong forces of attraction, often forming a crystalline lattice structure. The particles cannot move freely but vibrate about fixed points.

This fixed arrangement gives solids a definite shape and volume, as the particles cannot move to fill a container or change shape easily.

• Particles are closely packed in a fixed, regular pattern.
• Strong forces of attraction hold particles in place.
• Particles vibrate about fixed positions but do not move past each other.
• Solids have a definite shape and volume.

In liquids, particles are still close together but not in a fixed arrangement. The forces of attraction between particles are weaker than in solids, allowing particles to move past each other while remaining in contact.

This arrangement means liquids have a definite volume but no fixed shape. They take the shape of their container because particles can flow and rearrange.

The particles remain close enough to maintain a constant volume, unlike gases.

• Particles are close but arranged randomly and not fixed.
• Forces of attraction are weaker than in solids, allowing movement.
• Particles can slide past each other, enabling flow.
• Liquids have a fixed volume but take the shape of their container.

Gas particles are far apart with large spaces between them compared to solids and liquids. The forces of attraction between particles are negligible, allowing particles to move independently and rapidly in all directions.

Because particles are widely separated and move freely, gases have neither a fixed shape nor a fixed volume. They expand to fill the entire volume of their container.

The large separation also makes gases easily compressible, unlike solids and liquids.

• Particles are widely spaced with large gaps between them.
• Forces of attraction are negligible.
• Particles move rapidly and randomly in all directions.
• Gases have no fixed shape or volume and are compressible.

The movement of particles differs significantly between solids, liquids and gases, influencing their physical properties.

In solids, particles vibrate about fixed positions but do not move from place to place, resulting in rigidity and fixed shape.

In liquids, particles move randomly and slide past each other, allowing liquids to flow and change shape while maintaining volume due to close contact between particles. This random motion is continuous and non-ordered but constrained by intermolecular forces keeping particles close together.

• Solid particles vibrate in fixed positions; no translational movement.
• Liquid particles move randomly and slide past each other.
• Gas particles move rapidly and freely in all directions.
• Particle movement explains rigidity in solids, fluidity in liquids, and expansibility in gases.

The arrangement, separation and movement of particles directly determine the observable properties of solids, liquids and gases.

Solids have fixed shape and volume due to fixed particle positions and strong attractions.

Liquids have fixed volume but variable shape because particles are close but can move past each other, allowing flow and shape change to fit containers without volume change.

• Fixed particle positions in solids cause fixed shape and volume.
• Close but mobile particles in liquids cause fixed volume but variable shape.
• Widely separated, freely moving particles in gases cause variable shape and volume.
• Compressibility is high in gases due to large particle separation, negligible in solids and liquids.

Melting is the process where a solid changes into a liquid when heat is supplied. This occurs at a specific temperature called the melting point, which is characteristic of the substance.

During melting, the solid absorbs heat energy, causing it to change state without a change in temperature until the entire solid has melted.

A common example is ice melting to form liquid water at 0°C under standard atmospheric pressure (1 atm).

• Melting point is fixed for pure substances (e.g., ice melts at 0°C).
• Heat is absorbed (endothermic process).
• Solid changes to liquid with no temperature change during the process.
• Observation: solid becomes liquid; shape changes from fixed to fluid.

Boiling is the rapid change of a liquid into a gas that occurs throughout the liquid at a specific temperature called the boiling point.

The boiling point depends on the pressure; at 1 atm, water boils at 100°C.

Boiling involves bubbles of vapour forming inside the liquid and rising to the surface, producing vigorous bubbling and steam.

• Boiling point is fixed for a pure liquid at a given pressure.
• Heat is absorbed (endothermic process).
• Bubbles form throughout the liquid, not just at the surface.
• Observation: vigorous bubbling and steam production at boiling point.

Evaporation is the gradual change of a liquid into a gas at temperatures below the boiling point.

It occurs only at the surface of the liquid, producing vapour without bubbles.

Evaporation happens at all temperatures but is faster at higher temperatures, with increased surface area or lower humidity.

• Occurs at any temperature below boiling point.
• Evaporation occurs at the liquid surface and does not form bubbles.
• Heat is absorbed from surroundings (endothermic).
• Observation: liquid level decreases slowly without bubbling.

Freezing is the process where a liquid changes into a solid when heat is removed.

This occurs at the freezing point, which is the same temperature as the melting point for a pure substance.

During freezing, the liquid releases heat to the surroundings (exothermic process).

• Freezing point equals melting point for pure substances.
• Heat is released (exothermic process).
• Liquid changes to solid with no temperature change during the process.
• Observation: liquid solidifies, forming a fixed shape.

Condensation is the change of a gas into a liquid when it loses heat.

It occurs when a gas is cooled to its dew point or comes into contact with a cooler surface.

Condensation releases heat to the surroundings (exothermic process).

• Gas changes to liquid upon cooling.
• Heat is released (exothermic process).
• Observation: droplets form on cool surfaces (e.g., water droplets on a cold glass).

When the pressure on a gas is kept constant, its volume changes directly with temperature. This relationship is known as Charles's Law.

As the temperature of a gas increases, the kinetic energy of its particles increases, causing them to move faster and push outwards, expanding the gas volume.

Conversely, lowering the temperature reduces particle movement and decreases the gas volume, provided the pressure remains constant.

• Charles's Law states: V ∝ T (at constant pressure), where V is volume and T is temperature in kelvin.
• Mathematically: \( \frac{V_1}{T_1} = \frac{V_2}{T_2} \)
• Temperature must be measured in kelvin (K) for calculations.
• Example: A gas occupies 2.0 dm³ at 300 K. If heated to 450 K at constant pressure, the new volume \( V_2 \) is calculated as:
• \( V_2 = V_1 \times \frac{T_2}{T_1} = 2.0 \times \frac{450}{300} = 3.0 \, \text{dm}^3 \)

At constant temperature, the volume of a gas is inversely proportional to its pressure. This is described by Boyle's Law.

Increasing the pressure compresses the gas particles into a smaller volume, while decreasing the pressure allows the gas to expand.

This occurs because the gas particles collide more frequently with container walls under higher pressure, reducing the space they occupy.

• Boyle's Law states: V ∝ \( \frac{1}{P} \) (at constant temperature), where P is pressure.
• Mathematically: \( P_1 V_1 = P_2 V_2 \)
• Pressure must be in consistent units (e.g., kPa or atm) for calculations.
• Example: A gas has a volume of 4.0 dm³ at 100 kPa. If the pressure increases to 200 kPa at constant temperature, the new volume \( V_2 \) is:
• \( V_2 = \frac{P_1 V_1}{P_2} = \frac{100 \times 4.0}{200} = 2.0 \, \text{dm}^3 \)

In laboratory experiments, gas volume changes can be observed using a gas syringe or a sealed container with a movable piston.

When heating a gas at constant pressure, the gas expands and pushes the piston outward, increasing volume.

When compressing a gas at constant temperature, the piston moves inward as volume decreases and pressure increases.

• Ensure temperature is measured accurately with a thermometer and pressure with a manometer or pressure sensor.
• Repeat measurements to improve accuracy and calculate averages.
• Safety: Avoid heating gases in sealed containers without pressure relief to prevent explosions.
• Control variables carefully: keep pressure constant when studying temperature effects, and keep temperature constant when studying pressure effects.

A sample of gas occupies 5.0 dm³ at 300 K and 100 kPa. Calculate the volume when the temperature is increased to 360 K and the pressure is decreased to 80 kPa.

Step 1: Use combined gas law since both temperature and pressure change:

\( \frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2} \) Rearranged to find \( V_2 \): \( V_2 = \frac{P_1 V_1 T_2}{P_2 T_1} \).

• Substitute values: \( V_2 = \frac{100 \times 5.0 \times 360}{80 \times 300} \)
• Calculate numerator: 100 × 5.0 × 360 = 180,000
• Calculate denominator: 80 × 300 = 24,000
• Divide: 180,000 ÷ 24,000 = 7.5 dm³
• Answer: The new volume is 7.5 dm³.
• Common error: forgetting to convert temperature to kelvin or mixing units for pressure.

Kinetic particle theory states that all matter consists of tiny particles (atoms, molecules or ions) that are in constant motion. The state of matter (solid, liquid or gas) depends on the arrangement, movement and energy of these particles.

In solids, particles are closely packed in a fixed, regular arrangement and vibrate about fixed positions. In liquids, particles are close but can move past each other randomly, allowing flow. In gases, particles are far apart and move rapidly in all directions.

The energy of particles increases with temperature, causing changes in their motion and arrangement, which leads to changes of state.

• Particles in solids vibrate about fixed positions.
• Particles in liquids move randomly but remain close.
• Particles in gases move rapidly and are widely separated.
• Temperature increase raises particle kinetic energy.

When a solid is heated, its particles gain kinetic energy and vibrate more vigorously until they overcome the forces holding them in fixed positions. At the melting point, the solid changes to a liquid as particles move freely but remain close.

Further heating increases particle energy until the liquid boils. At the boiling point, particles have enough energy to overcome intermolecular forces and escape as gas.

Cooling reverses these processes: gas particles lose energy and condense to liquid; liquid particles lose energy and freeze to solid.

• Melting: solid to liquid; particles gain energy and move freely.
• Boiling: liquid to gas; particles gain enough energy to separate completely.
• Condensation: gas to liquid; particles lose energy and come closer.
• Freezing: liquid to solid; particles lose energy and become fixed.

During heating, temperature rises as particle kinetic energy increases until a change of state occurs. At this point, temperature remains constant despite continued heating because energy is used to break intermolecular forces, not increase kinetic energy.

Similarly, during cooling, temperature falls until a phase change occurs, then remains constant as energy is released when bonds form.

These energy changes are called latent heat: latent heat of fusion for melting/freezing and latent heat of vaporisation for boiling/condensation.

• Temperature constant during phase change despite heat input or loss.
• Energy used to overcome or form intermolecular forces.
• Latent heat of fusion: energy for melting/freezing.
• Latent heat of vaporisation: energy for boiling/condensation.

A heating curve plots temperature against heat energy added to a substance. It shows distinct regions: temperature rises in solid, liquid and gas phases, and flat plateaus where phase changes occur.

For example, heating ice at 0°C shows a plateau where ice melts to water; temperature remains constant until all ice melts. Then temperature rises until water boils at 100°C, where another plateau occurs during boiling.

The flat sections correspond to latent heat absorption; the sloped sections correspond to temperature increase and kinetic energy rise.

• Sloped regions: temperature increases; particles gain kinetic energy.
• Plateaus: temperature constant; energy used for phase change.
• Melting plateau at melting point.
• Boiling plateau at boiling point.

Cooling curves plot temperature against heat energy removed. They mirror heating curves but show temperature decreasing.

During cooling, temperature falls in gas and liquid phases until condensation occurs, shown by a plateau where temperature remains constant as gas condenses to liquid.

Further cooling leads to freezing, another plateau where liquid solidifies. After freezing, temperature decreases in the solid phase again as particles lose kinetic energy.

• Temperature decreases in gas, liquid, and solid phases.
• Plateaus indicate condensation and freezing.
• Energy released during phase changes.
• Particle motion slows as temperature falls.

Kinetic particle theory states that gases consist of tiny particles (atoms or molecules) in constant, random motion. These particles are widely spaced compared to their size, allowing them to move freely and rapidly in all directions. The particles collide elastically with each other and with the container walls, exerting pressure.

The average kinetic energy of gas particles is directly proportional to the absolute temperature (measured in kelvin). As temperature increases, particles move faster; as temperature decreases, they move slower. This relationship is fundamental to understanding gas behaviour.

• Gas particles are far apart and move randomly at high speeds.
• Particles collide elastically, transferring energy without loss.
• Average kinetic energy ∝ absolute temperature (K).

When the temperature of a gas increases at constant pressure, the particles gain kinetic energy and move faster. This causes the gas to expand because the particles collide with the container walls more forcefully and more frequently, pushing the walls outward if flexible, or increasing volume if the container allows expansion.

This relationship is described by Charles's Law: at constant pressure, the volume of a gas is directly proportional to its absolute temperature (V ∝ T).

For example, heating a gas from 300 K to 600 K at constant pressure doubles its volume.

• Increasing temperature increases particle speed and kinetic energy.
• Gas volume increases if pressure is constant (Charles's Law).
• Volume and absolute temperature are directly proportional: V₁/T₁ = V₂/T₂.

If the pressure on a gas increases while the temperature remains constant, the gas particles are forced closer together. This reduces the space between particles, decreasing the gas volume.

According to Boyle's Law, at constant temperature, the volume of a gas is inversely proportional to its pressure (V ∝ 1/P).

For example, doubling the pressure on a gas at constant temperature halves its volume.

• Increasing pressure compresses gas particles closer together.
• Gas volume decreases if temperature is constant (Boyle's Law).
• Pressure and volume are inversely proportional: P₁V₁ = P₂V₂.

A gas occupies 2.0 dm³ at 300 K and 1 atm pressure. What volume will it occupy at 450 K if pressure remains constant?

Using Charles's Law: V₁/T₁ = V₂/T₂, rearranged to V₂ = V₁ × (T₂/T₁).

Substitute values: V₂ = 2.0 dm³ × (450 K / 300 K) = 2.0 dm³ × 1.5 = 3.0 dm³.

• Formula: V₂ = V₁ × (T₂/T₁)
• Substitution: V₂ = 2.0 × (450/300)
• Answer: V₂ = 3.0 dm³
• Common error: Using °C instead of K in calculations.

A gas occupies 4.0 dm³ at 1.0 atm and 300 K. What volume will it occupy at 2.0 atm if temperature remains constant?

Using Boyle's Law: P₁V₁ = P₂V₂, rearranged to V₂ = (P₁V₁)/P₂.

Substitute values: V₂ = (1.0 atm × 4.0 dm³) / 2.0 atm = 2.0 dm³.

• Formula: V₂ = (P₁V₁)/P₂
• Substitution: V₂ = (1.0 × 4.0) / 2.0
• Answer: V₂ = 2.0 dm³
• Common error: Mixing units or not keeping temperature constant.