The vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor are in dynamic equilibrium (see below). If we were to place a substance in closed container, some of it would vaporize. The pressure in the space above the liquid would increase from zero and eventually stabilize at a constant value, the vapor pressure. If a liquid is not in a closed container it still has a vapor pressure, however, the liquid would eventually all evaporate away (turn into gas).
Even though the pressure in our closed container is constant, molecules of the vapor are still condensing into the liquid phase and molecules of the liquid are still evaporating into the vapor phase. However, the rate of these two processes is equal, so there is no net change in the amount of vapor or liquid. This process is called dynamic equilibrium. For this reason, the term equilibrium vapor pressure is sometimes used.
Vapor pressure and boiling point have an intimate relationship. The boiling point is the temperature at which the vapor pressure of the liquid equals the external pressure. For example, because the air pressure is lower in a city far above sea level such as Denver, the boiling point of water is lower than in a sea level city such as New York.
The most common unit for vapor pressure is the torr. 1 torr = 1 mm Hg (one millimeter of mercury).
Most materials have very low vapor pressures. For example, water has a vapor pressure of approximately 15 torr at room temperature. But remember that vapor pressures increase with temperature; water will have a vapor pressure of 760 torr = 1 atm at its boiling point of 100 oC (212 oF).
Solids are bound together much tighter and stronger than liquids. As a general rule, the vapor pressure of solids is much lower than the vapor pressure of a liquid.
In general, the higher the vapor pressure of a material at a given temperature, the lower the boiling point. In other words, compounds with high vapor pressures are volatile, forming a high concentration of vapor above the liquid; this can sometimes pose a fire hazard.
Specific heat is the ratio of the quantity of heat required to raise the temperature of a body one degree to that required to raise the temperature of an equal mass of water one degree. The term is also used in a narrower sense to mean the amount of heat, in calories, required to raise the temperature of one gram of a substance by one Celsius degree. The Scottish scientist Joseph Black, in the 18th century, noticed that equal masses of different substances needed different amounts of heat to raise them through the same temperature interval, and, from this observation, he founded the concept of specific heat.
The ability of water to stabilize temperature depends on its relatively high specific heat. The specific heat of a substance is defined at the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1º C. The specific heat of water is 1.00 cal/g ºC. Compared with most other substances, water has an unusually high specific heat. For example, ethyl alcohol, the type in alcoholic beverages, has a specific heat of 0.6 cal/g ºC.
Because of the high specific heat of water relative to other materials, water will change its temperature less when it absorbs or loses a given amount of heat. The reason you can burn your finger by touching the metal handle of a pot on the stove when the water in the pot is still lukewarm is that the specific heat of water is ten times greater than that of iron. In other words, it will take only 0.1 cal to raise the temperature of 1 g of iron 1ºC. Specific heat can be thought of as a measure of how well a substance resists changing its temperature when it absorbs energy.