Definition
According to kinetic theory, the basic difference between solid, liquid, and gas comes down to the strength of inter-particle attraction and the amount of kinetic energy the particles have — solids have the strongest attractions and least particle energy (particles only vibrate in fixed positions), liquids have intermediate attraction and energy (particles can move past each other but stay close), and gases have the weakest attractions and the most energy (particles move freely, far apart).
When a solid is heated, its vibrating particles gain more kinetic energy and vibrate more vigorously, until eventually they gain enough energy to break free of their fixed positions — this is melting.
Remember
Solids and liquids can diffuse too, not just gases, but far more slowly, since their particles are much more closely packed and have far less freedom to move past each other compared to a gas's widely spaced, fast-moving particles.
Definition
Melting: solid liquid. Freezing: liquid solid. Vaporization (evaporation/boiling): liquid gas. Condensation: gas liquid.
Sublimation: solid gas directly, skipping the liquid phase. Deposition: gas solid directly (cooling vapours of some substances turns them straight into solid) — this is what causes frost, and is also why naphthalene balls slowly disappear (sublime) into the air over time without ever visibly melting.
Example 1 — identifying phase changes
In the sequence gas liquid solid, identify phase changes and .
- Gas turning to liquid is condensation.
- Liquid turning to solid is freezing.
Answer: = condensation, = freezing
Example 2 — frost disappearing in the morning
Frost forms on the ground at 6:30 am and disappears by 8:30 am. Which phase change turned the frost into water vapour?
- Frost is solid ice; as the morning warms, it turns directly into water vapour (gas), skipping the liquid stage.
Answer: Sublimation
Remember
Dew forms in the early morning because the air cools overnight, and cooler air can hold less water vapour — the excess vapour condenses onto cool surfaces (grass, car windows) as tiny liquid droplets once the air's temperature drops below its dew point.
Definition
Evaporation happens only at the liquid's surface, at any temperature, and depends strongly on the force of attraction between molecules — weaker inter-molecular attraction lets molecules escape the surface more easily, so liquids with weak attractions evaporate faster.
Boiling happens throughout the liquid (not just the surface), only at a specific temperature (the boiling point) where the liquid's vapour pressure equals the external (atmospheric) pressure.
Boiling point depends on external pressure: lower external pressure (e.g. at high altitude) means the vapour pressure needed to match it is reached at a lower temperature, so liquids boil at a lower temperature at high altitude — and conversely, a sealed pressure cooker raises internal pressure, raising the boiling point and cooking food faster.
Internal energy of a substance includes both the kinetic energy of its particles (from their motion/vibration) and the potential energy stored in the attractive forces between them (from their arrangement/separation).
Remember
Evaporation doesn't need externally added energy because the fastest-moving, most energetic molecules already at the surface can escape using only the kinetic energy they already have — energy that came from the liquid's own thermal energy (which is why the remaining liquid cools down as evaporation proceeds).
Evaporation is faster when heated because a higher temperature means more molecules have enough energy to escape the surface at any given moment.
Wet clothes take longer to dry in humid/cold weather (like winter or a rainy season) because the air is already closer to holding as much water vapour as it can, and cooler air holds less vapour overall, both of which slow the net rate of evaporation from the clothes.
Definition
During a phase change, all the energy supplied (or removed) goes into breaking (or forming) the attractive bonds between particles — changing their potential energy, not their kinetic energy — so the temperature (which reflects average kinetic energy) stays constant throughout the change, even though heat keeps flowing in or out.
This is exactly why heating and cooling curves for a substance follow the same flat plateaus at the melting and boiling points in both directions: the same amount of energy is absorbed to break the bonds on heating as is released re-forming them on cooling.
Remember
Water vapour at C causes more severe burns than liquid water at the same C, because the vapour carries extra hidden (latent) heat energy on top of its temperature — when it condenses on skin, it releases this large additional quantity of energy (the latent heat of vaporization) before even starting to cool down as a liquid, delivering far more total heat to the skin.
Remember
Ice melts under pressure (e.g. under an ice skate's blade) because, unusually for most substances, water's solid form (ice) is less dense than its liquid form — squeezing ice favours the denser liquid state, so applying enough pressure can melt it even without adding heat.
Rule
Boyle's Law: at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure — squeezing a gas into half the volume doubles its pressure, and vice versa.
Charles's Law: at constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute (Kelvin) temperature — heating a gas at constant pressure makes it expand.
Avogadro's Law: equal volumes of different gases, at the same temperature and pressure, contain equal numbers of molecules — regardless of what the gas actually is.
Boyle's Law:
Charles's Law:
Example 3 — Boyle's law: doubling the pressure
In Boyle's experiment, trapped air has volume at pressure . What happens to the volume if the pressure is doubled (temperature constant)?
- .
Answer: The volume is halved
Example 4 — Boyle's law calculation
A gas occupies L at a pressure of atm. Find its new volume if the pressure is increased to atm at constant temperature.
- .
- .
- .
Answer: L
Example 5 — Charles's law calculation
A gas has volume L at K. Find its volume at K, at constant pressure.
- .
- .
- .
Answer: L
Definition
Diffusion (gases mixing/spreading due to random particle motion) happens fastest for the lightest gas, since lighter molecules move faster (at the same temperature) than heavier ones, letting them spread through a space more quickly. This is also why lighter drug molecules can diffuse through the body and take effect faster than heavier ones, all else being equal.
Definition
Sublimation is used in printing (some inks/dyes sublime directly onto fabric or paper under heat, in sublimation printing) and in air fresheners (solid fragrance blocks slowly sublime, continuously releasing scent molecules into the air without ever needing to melt).
Phase change materials (PCMs) absorb or release large amounts of latent heat while changing state at a nearly constant temperature, making them useful for thermal regulation — e.g. PCM-infused fabrics or building materials that absorb excess heat by melting during the day and release it back by solidifying at night, helping keep temperature steady.
Remember
Cotton clothes feel more comfortable in summer because cotton fibres absorb sweat and let it evaporate easily from a larger exposed surface area, and evaporation itself cools the skin by removing latent heat — synthetic fabrics that trap moisture close to the skin don't allow this cooling as effectively.
Quick Recap
- ●Kinetic theory: solid (strong attraction, low energy) liquid gas (weak attraction, high energy)
- ●Phase changes: melting/freezing, vaporization/condensation, sublimation/deposition
- ●Evaporation: surface only, any temperature, depends on inter-molecular attraction, doesn't need added energy
- ●Boiling: throughout the liquid, at boiling point where vapour pressure external pressure
- ●Temperature constant during phase change: energy changes potential energy (bonds), not kinetic energy
- ●Boyle's Law: . Charles's Law: . Avogadro's Law: equal volumes equal molecules
