The temperature where ice gives in
Warm a block of ice slowly and, at one sharp temperature, it begins to melt. For pure water at ordinary pressure that temperature is 0 °C — the melting point. It is wonderfully reproducible: a clean ice cube anywhere on Earth melts at the same point, which is why it was once used to define a temperature scale.
Melting is a phase transition from solid to liquid, and like all transitions it has an energy price. The energy needed to melt a fixed amount of a substance, at its melting point, is its latent heat of fusion — a particular latent heat. Pour that energy in and the temperature does not budge until the last crystal is gone; only then does the liquid start warming again.
A liquid is always quietly evaporating
Leave a glass of water on a table and it slowly empties, long before it ever boils. Inside the liquid, molecules move at a range of speeds; the fastest few near the surface have enough energy to break free and escape into the air. That is evaporation, and it happens at every temperature.
Now seal the glass with a lid. Escaping molecules build up above the liquid, and some bump back in. Soon escape and return balance out, and the trapped vapor settles to a steady pressure called the vapor pressure. Think of it as the strength of a liquid's outward "breath" — how hard its molecules push to get into the gas phase.
Vapor pressure rises steeply with temperature: hotter molecules escape more eagerly, so the breath grows stronger. A liquid with high vapor pressure (like alcohol or petrol) feels "volatile" — it evaporates fast and smells from across the room. Water's is more modest, which is why a puddle lingers.
Boiling, finally explained
Boiling is more dramatic than quiet evaporation: bubbles of pure vapor form deep inside the liquid and rise. But a bubble can only survive if its vapor pushes out as hard as the surrounding air and water push in. So boiling begins exactly when the liquid's vapor pressure climbs to equal the outside pressure.
The temperature where that happens is the boiling point. At sea level, where the air presses with one atmosphere, water reaches that match at 100 °C. The crucial twist: the boiling point depends on the surrounding pressure, not just on the liquid. Change the pressure and you change where it boils.
How much energy a boil really costs
Turning liquid into gas is energetically expensive, because every molecule must fully tear away from its neighbours. The heat needed to vaporise one mole of a liquid at its boiling point is its enthalpy of vaporization — the boiling cousin of the latent heat of fusion, and usually several times larger.
For water this number is famously large — it takes far more energy to boil a pot dry than merely to bring it to a boil. That single fact powers a lot of nature: it lets oceans store and move enormous amounts of heat, drives the weather, and makes sweating such an effective way to cool a body.
Some solids skip the liquid entirely. Dry ice (frozen carbon dioxide) turns straight to gas in a process called sublimation, with its own latent heat. It happens whenever a solid's vapor pressure reaches the surrounding pressure before the solid ever melts — a clue we will make precise when we draw the phase map.
Pulling the ideas together
Three quantities now describe how a pure substance changes phase. The melting point and boiling point tell you *where* on the temperature scale transitions occur; the latent heats tell you *how much* energy each one demands; and vapor pressure is the hidden link that ties boiling to the pressure of the surroundings.
Hold onto one big realisation: because boiling depends on pressure, the whole behaviour of a substance is really a story told across two dials at once — temperature and pressure. The next guide puts both on a single picture, the phase diagram, where these scattered facts snap into one elegant map.