The temperature at which the equilibrium vapor pressure between a liquid and its vapor is equal to the external pressure on the liquid. Physically, boiling (or ebullition) cannot begin in a liquid until the temperature is raised to such a point that incipient bubbles forming within the liquid can grow rather than collapse. But for a bubble to grow, its internal vapor pressure must exceed the hydrostatic pressure exerted on the bubble interface. For liquids heated in containers that are open and fairly shallow, this hydrostatic pressure is essentially the same as the external atmospheric pressure, so ebullition begins when the equilibrium vapor pressure equals atmospheric pressure. In liquids perfectly free from foreign particles and contained in a vessel with perfectly smooth walls, boiling will not begin even at the above-described temperature, for boiling resembles condensation in that “nuclei” must exist to initiate the process. When a very pure liquid sample has been heated above its nominal boiling point it is said to have been superheated, a state that is very similar to the state of supersaturation in which a vapor may exist in a nucleus-free environment. Because of the normal decrease of barometric pressure with height, the nominal boiling point of water decreases 3. 0°– 3. 5°C for each kilometer increase of altitude (see hypsometer). The boiling point is a colligative property of a solution; with an increase in dissolved matter, there occurs a raising of the boiling point. The boiling point of pure water at standard pressure is equal to 100°C (212°F) and is a fiducial point for thermometer calibration. Compare ice point.