IB Physics HLTopic 2 — Matter, Heat & ElectricityPaper 1 & 2Changes of State~8 min read
Changing State
Ice, water and steam again — but this time we care about the crossings between them. Melting, freezing, boiling, condensing: each is a swap between solid, liquid and gas, driven by energy flowing in or out. And there’s a curious twist that falls straight out of internal energy — while a substance is changing state, its temperature flatly refuses to move.
📚 What you need to know
A change of state (phase change) is when matter changes between solid, liquid and gas.
Energy is transferred during every phase change — absorbed (in) or released (out).
That energy changes the particles’ potential energy (spacing and bonds), not their kinetic energy — so the temperature stays constant throughout.
The four main changes: melting (solid→liquid, in), freezing (liquid→solid, out), vaporisation / boiling (liquid→gas, in), condensation (gas→liquid, out).
Melting & freezing happen at the melting/freezing point; vaporisation & condensation at the boiling point.
Every substance has its own fixed points — water: 0 °C (melt/freeze), 100 °C (boil/condense).
You may also meet sublimation (solid→gas) and deposition (gas→solid).
Four ways to change state
A phase change is just matter swapping between its three states. Climb the ladder from solid to liquid to gas and you have to pull particles apart against the forces holding them — that always takes energy in. Come back down and the particles fall together, giving energy out.
The four main phase changes. Going towards gas absorbs energy (red, top); going towards solid releases it (blue, bottom). Solid and gas can also swap directly by sublimation and deposition.
Here’s a neat symmetry to lean on: whatever energy a change absorbs one way, it releases going back. The energy given out when steam condenses is exactly the energy it took to boil that water in the first place. Up the ladder costs energy; down the ladder pays it back.
Why the temperature holds still
This is the idea examiners love. During a phase change you can pour energy in and watch a thermometer sit dead still. It’s not a paradox — it follows straight from internal energy. The supplied energy goes entirely into potential energy, prising particles apart, and none into kinetic energy. Since temperature measures average kinetic energy, it doesn’t budge until the change is complete.
Melting: the energy breaks the forces between particles (potential energy rises) rather than speeding them up (kinetic energy is unchanged), so the temperature stays put.
Say it in the exam like this: “the energy supplied increases the particles’ potential energy by overcoming the intermolecular forces, not their kinetic energy, so the temperature stays constant.” That one sentence lands the marks nearly every time.
Melting and boiling points
Each substance flips state at its own set temperatures. The melting point and freezing point are the same temperature — just crossed in opposite directions — and likewise the boiling point equals the condensation point. For water these sit at the familiar 0 °C and 100 °C.
Water’s fixed points: solid ice below 0 °C, liquid water between, steam above 100 °C. Melting and freezing share 0 °C; boiling and condensation share 100 °C.
Don’t assume every substance uses water’s numbers. Iron melts near 1500 °C, oxygen boils at around −183 °C. The ideas are identical for every material — only the fixed-point temperatures change.
Worked examples
WE 1
For each change, name the process and state whether thermal energy is absorbed or released: (a) liquid water → steam, (b) water vapour → dew droplets, (c) molten iron → solid iron, (d) ice → water.
(a) liquid → gas: vaporisation (boiling) — energy ABSORBED(b) gas → liquid: condensation — energy RELEASED(c) liquid → solid: freezing — energy RELEASED(d) solid → liquid: melting — energy ABSORBEDtowards gas = absorb · towards solid = releaseThat single rule of thumb gets you the energy direction for any phase change.
WE 2
When a pure solid melts, energy is supplied continuously yet its temperature stays constant. Explain why, using kinetic and potential energy.
Temperature is a measure of the average KINETIC energy of the particles.During melting the supplied energy goes into POTENTIAL energy — overcoming the forces holding the lattice together.The average kinetic energy doesn’t change, so neither does the temperature.energy → PE (not KE) → temperature constantOnly once melting is finished does extra energy start raising the KE again — and the temperature climbs.
WE 3
Steam at 100 °C causes far worse burns than liquid water at 100 °C, even though they’re the same temperature. Explain why.
When steam meets your skin it first CONDENSES into water.Condensation is a change of state, so it RELEASES energy — the same energy that was absorbed to boil the water originally.That large burst is delivered before the water even begins to cool from 100 °C.steam dumps extra energy as it condenses → worse burnThe size of that extra energy is the latent heat of vaporisation — put on a number on a later page.
🔧 Working out a phase change
Which phases? Note the state before and after (solid / liquid / gas).
Which direction? Towards gas → energy IN (absorbed). Towards solid → energy OUT (released).
Name it: melting / freezing (solid⇆liquid) or vaporising / condensing (liquid⇆gas).
Temperature? Constant during the change — the energy is potential, not kinetic.
At what temperature? The melting point (solid⇆liquid) or boiling point (liquid⇆gas).
SOLID
melt → +E ← freeze −E
LIQUID
boil → +E ← condense −E
GAS
Quick recap: A change of state swaps matter between solid, liquid and gas. Moving towards gas absorbs energy; moving towards solid releases it. Throughout, the energy shifts the particles’ potential energy (spacing), not their kinetic energy, so the temperature stays constant. Melting and freezing share one fixed point; boiling and condensation share another — 0 °C and 100 °C for water.
💡 Top tips
Up the ladder = energy in; down = energy out. Melting and boiling absorb; freezing and condensing release.
Temperature is constant during any phase change — the energy is potential, not kinetic.
Melting point = freezing point, and boiling point = condensation point — same temperature, opposite direction.
Fixed points are per-substance. Don’t assume 0 °C / 100 °C for anything but water.
Condensing and freezing release energy — easy to forget, and the reason steam burns are so nasty.
⚠ Common mistakes
Saying the temperature keeps rising while ice melts — it holds at the melting point
Thinking the energy “disappears” during a phase change — it becomes potential energy (or is released from it)
Mixing up directions — melting absorbs, freezing releases
Assuming every substance melts at 0 °C and boils at 100 °C — those are water’s values only
Forgetting that condensation and freezing release energy, not absorb it
You can now name every crossing between solid, liquid and gas, and you know the temperature sits still while it happens. The next questions are the quantitative ones: how much energy does it take to warm a substance by a set amount, and how much to change its state entirely? First up: specific heat capacity and the mcΔT equation for plain temperature changes — then latent heat for the state changes themselves.
Want this to actually click before the exam?
Book a free meeting and let’s work through the tricky bits together.
