boiling point
Heat a pot of water and at first only the surface evaporates quietly. Then, at a certain temperature, something dramatic happens: bubbles of vapor form deep inside the liquid and rush up. That temperature is the boiling point — the moment vaporization breaks out throughout the whole liquid, not just at the top.
The reason is a simple tug-of-war. A vapor bubble can only survive inside the liquid if its internal pressure can push back against the surrounding atmosphere squeezing in. The boiling point is the temperature at which a liquid's vapor pressure finally rises to equal the outside pressure, so bubbles stop collapsing and grow instead. Because it depends on the outside pressure, boiling point is not one fixed number: water boils at 100 °C at sea level but only about 70 °C atop a high mountain.
To compare substances fairly, chemists quote the normal boiling point — the temperature at which a liquid boils under exactly one standard atmosphere. This is why pressure cookers cook faster (higher pressure raises the boiling point, so the water gets hotter) and why high-altitude recipes need adjusting. A high boiling point signals molecules that cling tightly to one another.
On top of Mount Everest, where the air presses with only a third of its sea-level force, water boils at about 70 °C — too cool to brew a proper cup of tea.
Lower air pressure means a lower boiling point.
Boiling differs from evaporation: evaporation happens only at the surface and at any temperature, while boiling happens throughout the bulk and only once vapor pressure matches the surrounding pressure. Dissolving salt raises the boiling point slightly — a colligative effect.