The boiling point is a fundamental physical property of a substance, defined as the temperature at which a liquid changes its state to a gas. This transformation is not mere evaporation, which only occurs at the liquid’s surface, but a bulk phenomenon happening throughout the entire volume of the liquid. This property is constant for a given substance under standard conditions and is one of the most reliable ways to identify pure liquids.
What Causes a Liquid to Boil
The physical mechanism that drives a liquid to boil centers on the concept of vapor pressure, which is the pressure exerted by gas molecules above a liquid in a closed system. Within any liquid, molecules are constantly moving, and some near the surface gain enough kinetic energy to break free and enter the gas phase. As the liquid is heated, the molecules move faster and more of them escape, causing the vapor pressure to increase progressively.
Boiling begins when this internal vapor pressure equals the external pressure pushing down on the liquid, typically the surrounding atmospheric pressure. Once this balance is reached, the liquid can no longer suppress the formation of vapor bubbles within its main body. These bubbles signal the start of the boiling process. For water at sea level, the vapor pressure reaches 101.3 kilopascals (kPa), matching the standard atmospheric pressure, at 100°C.
Adding more heat once the boiling point is reached does not increase the temperature of the liquid. Instead, all the added energy, known as the latent heat of vaporization, is consumed in the process of converting the liquid molecules into gas. This energy provides the force necessary to push the liquid molecules completely apart. The temperature remains constant until the liquid has been converted entirely into the gaseous state.
Why Boiling Points Differ Among Substances
The variation in boiling points among different liquids, such as water at 100°C and ethanol at about 78°C, is due to the unique attractive forces holding the molecules together. These are known as Intermolecular Forces (IMFs) and must be overcome for a molecule to escape into the gas phase. The stronger these forces are, the more energy, and therefore the higher temperature, is needed to separate the molecules and achieve boiling.
The weakest IMFs are the London dispersion forces, temporary attractions caused by fleeting imbalances in electron distribution. Slightly stronger are dipole-dipole forces, which occur between molecules that have a permanent separation of positive and negative charge (a dipole). Both forces contribute to the attraction between molecules in most liquids.
Water’s notably high boiling point is largely a result of the strongest type of IMF, called hydrogen bonding. This is a special, powerful form of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine. These strong hydrogen bonds create a network of attraction that effectively locks the water molecules together. Breaking this network requires significantly more thermal energy compared to breaking the weaker forces found in a substance like ethanol.
How External Pressure Affects Boiling Point
The boiling point of a liquid is dependent on the external pressure applied to its surface. The temperature at which a liquid boils under standard atmospheric pressure (1 atmosphere or 101.3 kPa) is referred to as its standard boiling point. Any change in the surrounding pressure will directly alter the temperature required for the liquid’s vapor pressure to match it.
A decrease in external pressure lowers the boiling point, which is evident at high altitudes. For instance, on the summit of Mount Everest, where the atmospheric pressure is significantly lower than at sea level, water boils at approximately 71°C. Since the water boils at a lower temperature, food takes much longer to cook because the maximum cooking temperature is substantially reduced.
Conversely, increasing the external pressure raises the boiling point of the liquid. This principle is utilized in the pressure cooker, which traps steam and increases the pressure inside the sealed vessel. This higher pressure forces the water to remain a liquid at temperatures well above 100°C, typically reaching around 120°C. Cooking at these elevated temperatures greatly reduces the time needed for food preparation.