Solids, Liquids and Gases
1.1 Solids, Liquids and Gases
1. Distinguishing Properties of Solids, Liquids and Gases
Matter exists in three common states under everyday conditions: solid, liquid, and gas. Each state has characteristic physical properties which arise from the arrangement, separation, and motion of particles.
1.1.1 Shape
- Solid:
- Has a fixed shape which does not change unless an external force deforms or breaks it.
- Example: A metal cube stays cubic whether placed on a table or in a box.
- Liquid:
- Has no fixed shape; takes the shape of the part of the container it occupies.
- Example: Milk poured into a bowl spreads out to fit the shape of the bowl.
- Gas:
- Has no fixed shape; fills the whole container evenly, regardless of its size or shape.
- Example: Carbon dioxide from a soda spreads throughout a room if released.
1.1.2 Volume
- Solid:
- Fixed volume; cannot be squeezed into a smaller volume under normal conditions.
- Liquid:
- Fixed volume; resists compression, but shape changes easily.
- Gas:
- No fixed volume; expands or contracts to fill the container.
1.1.3 Density
- Solid: Usually high density (particles closely packed).
- Example: Gold ~19.3 g/cm³.
- Liquid: Moderate to high density; usually slightly lower than solid form of the same substance.
- Example: Water ~1.00 g/cm³ at 4°C; ice floats because it is less dense (~0.92 g/cm³).
- Gas: Very low density compared to solids and liquids.
- Example: Air ~0.0012 g/cm³ at room temperature.
1.1.4 Compressibility
- Solid: Almost incompressible; particles are already very close together.
- Liquid: Slightly compressible, but volume changes are minimal even under high pressure.
- Gas: Highly compressible; volume decreases dramatically when pressure is applied.
1.1.5 Ability to Flow
- Solid: Does not flow; particles are fixed in position.
- Liquid: Flows easily; particles can slide past one another.
- Gas: Flows very easily; particles move rapidly in all directions.
Summary Table:
| Property | Solid | Liquid | Gas |
|---|---|---|---|
| Shape | Fixed | Not fixed | Not fixed |
| Volume | Fixed | Fixed | Not fixed |
| Density | High | Moderate/High | Low |
| Compressibility | Almost none | Slight | High |
| Flow | None | Flows | Flows easily |
2. Structure of Solids, Liquids and Gases (Particle Separation, Arrangement, Motion)
2.1 Solids
- Separation: Particles extremely close (distance about 0.1–0.3 nm).
- Arrangement: Regular and ordered (crystalline solids) or irregular (amorphous solids like glass).
- Motion: Particles vibrate in fixed positions; no translational movement.
- Forces: Very strong forces (ionic, covalent, metallic, or strong van der Waals).
- Energy: Lowest kinetic energy of all three states.
Text Diagram (Solid Structure):
[O][O][O][O]
[O][O][O][O]
[O][O][O][O]
Particles fixed in position in a lattice.
2.2 Liquids
- Separation: Slightly greater than in solids.
- Arrangement: Random; particles not fixed.
- Motion: Particles slide/roll over each other; vibrational and rotational motion.
- Forces: Moderate; strong enough to maintain fixed volume but weak enough to allow flow.
- Energy: Higher kinetic energy than solids but less than gases.
Text Diagram (Liquid Structure):
O O O O
O O O
O O O
Random arrangement, close spacing.
2.3 Gases
- Separation: Very large compared to particle size.
- Arrangement: Random, far apart.
- Motion: Rapid and free in all directions; constant collisions with container walls.
- Forces: Negligible except during collisions.
- Energy: Highest kinetic energy of all three states.
Text Diagram (Gas Structure):
O O
O O
O O
Particles widely spaced and moving rapidly.
3. Changes of State in Terms of Kinetic Particle Theory
Kinetic Particle Theory explains changes of state through changes in particle energy and movement.
3.1 Melting (Solid → Liquid)
- Process:
- Heating increases kinetic energy → particles vibrate more strongly.
- At melting point, vibrations overcome forces holding particles fixed.
- Structure breaks down into a liquid.
- Temperature: Constant during melting; energy supplied is used to overcome forces (latent heat of fusion).
- Example: Ice melts at 0°C under 1 atm pressure.
3.2 Boiling (Liquid → Gas)
- Process:
- Heating increases kinetic energy of particles.
- At boiling point, particles have enough energy to overcome all intermolecular forces.
- Gas bubbles form throughout the liquid and rise to the surface.
- Pressure Effect: Boiling point decreases at lower pressure.
- Example: Water boils at 100°C at sea level; ~70°C at Mount Everest summit.
3.3 Evaporation (Liquid → Gas, below boiling point)
- Process:
- High-energy particles at the surface escape into the gas phase.
- Leaves behind particles with lower average kinetic energy → cooling effect.
- Factors increasing rate:
- Higher temperature
- Larger surface area
- Lower humidity
- Faster airflow
- Example: Sweat evaporating from skin cools the body.
3.4 Freezing (Liquid → Solid)
- Process:
- Cooling reduces kinetic energy.
- Intermolecular forces fix particles into ordered solid arrangement.
- Temperature: Constant at freezing point (same as melting point for a pure substance).
3.5 Condensation (Gas → Liquid)
- Process:
- Cooling reduces particle speed.
- Forces bring particles together into a liquid.
- Example: Water vapour condenses on a cold glass of water.
4. Heating and Cooling Curves in Terms of Kinetic Particle Theory
4.1 Heating Curve
- Solid heating:
- Temperature rises; particles vibrate faster.
- Melting plateau:
- Temperature constant; energy used to break bonds (latent heat of fusion).
- Liquid heating:
- Temperature rises; particles move faster.
- Boiling plateau:
- Temperature constant; energy used to separate particles completely (latent heat of vaporisation).
- Gas heating:
- Temperature rises; particles move even faster.
4.2 Cooling Curve
- Gas cooling:
- Temperature falls; particles slow.
- Condensation plateau:
- Temperature constant; energy released as gas becomes liquid.
- Liquid cooling:
- Temperature falls further.
- Freezing plateau:
- Temperature constant; energy released as liquid becomes solid.
- Solid cooling:
- Temperature falls; particles vibrate less.
Example:
Heating ice from −20°C to steam at 120°C involves:
- Heating solid ice
- Melting at 0°C
- Heating liquid water
- Boiling at 100°C
- Heating steam
5. Effects of Temperature and Pressure on the Volume of a Gas
5.1 Temperature Effect (Charles’s Law)
- Law: V ∝ T (in kelvin) at constant pressure.
- Higher temperature → faster particle movement → greater volume.
- Formula: V₁/T₁ = V₂/T₂
(T in kelvin) - Example Calculation:
- A gas at 300 K has a volume of 2.0 dm³.
- What is its volume at 450 K?
V₂ = V₁ × (T₂/T₁) = 2.0 × (450/300) = 3.0 dm³.
5.2 Pressure Effect (Boyle’s Law)
- Law: P ∝ 1/V at constant temperature.
- Increasing pressure compresses gas → decreases volume.
- Formula: P₁V₁ = P₂V₂
- Example Calculation:
- A gas at 100 kPa has a volume of 5.0 dm³.
- What volume will it have at 250 kPa?
V₂ = (P₁V₁) / P₂ = (100 × 5.0) / 250 = 2.0 dm³.
5.3 Combined Gas Law
- Formula: (P₁V₁)/T₁ = (P₂V₂)/T₂
- Allows simultaneous changes in pressure, volume, and temperature.
5.4 Real-life Applications
- Aerosol cans: Dangerous when heated; pressure inside rises due to increased particle speed.
- Hot air balloons: Heated air expands, lowering density, making the balloon rise.
- Syringes: Pulling plunger increases volume → pressure drops → liquid enters.
- Refrigerators: Use evaporation and condensation cycles to remove heat