Vaporization (or vaporisation in British English) of an element or compound is a phase transition from the liquid phase to vapor. There are two types of vaporization: evaporation and boiling.
Evaporation is the transition from liquid to vapor that occurs slowly and smoothly, affecting only the surface layers of the liquid. The surface particles of the liquid, which are less bound than the internal particles, can leave the liquid and become vapor. Evaporation occurs at all temperatures, but with different intensities: an increase in temperature increases the kinetic energy of the liquid molecules and consequently increases evaporation. The amount of particles that evaporate also depends on the free surface area of the liquid: the greater the surface area, the greater the probability that evaporation will occur. When evaporation occurs in a closed environment, a state of equilibrium is reached between the liquid and the vapor such that the number of liquid particles that become vapor is maintained over time equal to the number of vapor particles that become liquid. Under these conditions, the vapor is said to be saturated, and the pressure at which equilibrium occurs is called the vapor pressure (or saturated vapor pressure).
Vapor pressure is different for each liquid and increases with temperature, but it is completely independent of the mass of the liquid. The vapor pressure of a substance measures its volatility, or ability to evaporate, at a given temperature: alcohol, for example, is more volatile than water at room temperature.The latent heat of evaporation is defined as the amount of heat required to evaporate a unit mass of liquid. When a liquid evaporates, it removes from the environment an amount of heat equal to the latent heat of evaporation (this heat is returned in the reverse process). This explains, for example, the cold sensation we feel when sweat evaporates from our skin.
Boiling is the transition from liquid to aeriform (gas or vapor), which occurs rapidly and turbulently and affects the entire mass of the liquid. All liquids contain gas bubbles in which molecules of the liquid are trapped in the gas or vapor state. As the temperature of the liquid rises, the bubbles expand, and when their vapor pressure becomes equal to the external pressure, the phenomenon of boiling occurs, in which the bubbles rise to the surface and release the vapor or gas they contain. The boiling of a liquid at a given external pressure occurs at a certain temperature, called the boiling temperature (or boiling point), which remains constant throughout the boiling process. The liquid is said to boil at that particular temperature.
The boiling temperature is the temperature at which the vapor pressure of the liquid is equal to the pressure at the surface of the liquid. The boiling temperature varies with pressure: it increases as the external pressure increases and decreases as the external pressure decreases. As the external pressure decreases, the pressure at which boiling can occur decreases, and consequently boiling can occur at a lower temperature, and vice versa as the external pressure increases. For example, the boiling temperature of water at normal atmospheric pressure (101.32 kPa) is 100 °C; at pressures half of normal atmospheric pressure (about 50 kPa), a condition that occurs at about 5500 m altitude, for example, water boils at 86 °C. By increasing the external pressure, water boils at temperatures above 100 °C, as in a pressure cooker. The amount of heat required to completely boil a unit mass of liquid is called the latent heat of boiling.