Boiling Point
Definition and meaning of Boiling Point in chemistry.
Boiling point is the temperature at which a liquid's vapor pressure equals the pressure of the surrounding environment. At this point, the liquid vaporizes throughout its entire volume, not just at the surface. Boiling point depends on external pressure, so it is not a fixed property of a substance alone.
In more detail
As a liquid heats up, its molecules gain kinetic energy and move faster. This raises the liquid's vapor pressure, which is the pressure exerted by molecules escaping into vapor above the liquid's surface. Boiling begins the instant this vapor pressure matches the external pressure pushing down on the liquid.
At that moment, bubbles of vapor can form anywhere within the liquid, not just at the exposed surface. This is why boiling looks so much more dramatic than gentle evaporation. Because atmospheric pressure changes with altitude and weather, a liquid's boiling point is not truly a fixed number.
Lower external pressure lets vapor pressure catch up sooner, so the liquid boils at a lower temperature. This is why chemists specify the normal boiling point. It is defined as the boiling point measured specifically at 1 atmosphere of pressure, 101.325 kilopascals, to allow fair comparisons between substances.
Trends in boiling point across different substances reveal a lot about intermolecular forces, the attractions between separate molecules. Substances with strong hydrogen bonding, like water, or strong dipole-dipole attraction tend to boil at higher temperatures. More energy is needed to separate their molecules into gas.
Heavier molecules also tend to have higher boiling points, because they generally have stronger dispersion forces. A common misconception is that boiling requires reaching a specific temperature no matter the conditions. In reality, that specific temperature always assumes standard atmospheric pressure and shifts noticeably at high altitude or under vacuum.
Key facts
| Field | Physical Chemistry |
|---|---|
| Normal boiling point of water | 100 C (212 F) at 1 atm |
| Formula (water example) | H2O |
| Pressure dependence | Boiling point decreases as external pressure decreases |
| Governing condition | Vapor pressure equals external pressure |
| Related to | Strength of intermolecular forces |
Water has a normal boiling point of 100 degrees Celsius, or 212 degrees Fahrenheit, at 1 atmosphere of pressure. Atop Mount Everest, where atmospheric pressure is much lower, water boils at only about 70 degrees Celsius. This is why food takes longer to cook there.
Frequently asked questions
Why does water boil at a lower temperature at high altitude?
Atmospheric pressure is lower at high altitude. Water's vapor pressure reaches that lower value at a lower temperature, so the water boils below 100 degrees Celsius.
Is boiling point the same as evaporation?
No. Evaporation happens only at a liquid's surface and can occur at any temperature. Boiling happens throughout the liquid, and only at the specific temperature where vapor pressure equals external pressure.
Why do stronger intermolecular forces raise boiling point?
Stronger attractions between molecules take more energy to overcome. This means the liquid must be heated to a higher temperature before enough molecules can escape into the gas phase to match external pressure.