Vapor (or vapour in British English and Canadian English) is a physical state of matter, defined as an aeriform state at a temperature below its critical temperature. The most common example of vapor is steam (water vaporized during boiling or evaporation). In common language, steam is used as a synonym for water vapor.
From the physical point of view gases and vapor are distinguished because the gas can not in any way be condensed (ie reduced to the liquid state) unless after being brought to lower than the critical temperature. For example, the air can be compressed up to thousands of atmospheres of pressure while remaining gas; to make it liquid its temperature must be lower than about -150 °C.
Water vapor (water vapour [UK] or aqueous vapor) is the gaseous phase of water. It is one state of water within the hydrosphere. It is invisible, odorless, and colorless. When water vapor collects in large quantities and mixes with dust, various gases, pollen, combustion residues, it then becomes less transparent, giving rise to the phenomenon of mist or haze.
Water vapor is a relatively common atmospheric constituent, present even in the solar atmosphere as well as every planet in the Solar System and many astronomical objects including natural satellites, comets, and even large asteroids. Likewise, the detection of extrasolar water vapor would indicate a similar distribution in other planetary systems. Water vapor is significant in that it can be indirect evidence supporting the presence of extraterrestrial liquid water in the case of some planetary mass objects.
Water vapor technology was developed since the 17th century and received an effective application in the second half of the 18th century, mainly by British and French scientists and engineers, including Denis Papin and James Watt. The water vapor is obtained by evaporation by boiling the water in special equipment called boilers or more precisely steam generators.
Thermodynamic conditions of water vapor
For technological applications, the distinction between wet saturated, dry saturated and superheated makes a big difference, because the use of water vapor in thermal machines exploits the temperature difference, i.e. the heat it carries (the superheated steam, considered the highest temperature, carries the greater quantity), and therefore with the same quantity of vapor used, those water droplets greatly reduce the available energy and, in the case of fast machines such as steam turbines, those droplets beat violently on the metal, ruining the machines (cavitation).
Fundamentally, vapor pressure is the partial pressure of liquid-vapor (e.g. water vapor). Saturated vapor pressure is the vapor pressure which is in equilibrium with an open liquid surface. Therefore, it is also the pressure at which a liquid will vaporize for a given temperature. If the pressure to which the liquid is exposed is equal to the saturated vapor pressure for a given temperature, then the water will boil. The saturated vapor is a mixture of vapor and liquid droplets and exists when vapor coexists above boiling liquid.
Saturated vapor pressure and boiling point
A liquid boils when its saturated vapor pressure becomes equal to the external pressure on the liquid. When that happens, it enables bubbles of vapor to form throughout the liquid – those are the bubbles you see when a liquid boils. If the external pressure is higher than the saturated vapor pressure, these bubbles are prevented from forming, and you just get evaporation at the surface of the liquid.
If the liquid is in an open container and exposed to normal atmospheric pressure, the liquid boils when its saturated vapor pressure becomes equal to 1 atmosphere (or 101325 Pa or 101.325 kPa or 760 mmHg). This happens with water when the temperature reaches 100°C. But at different pressures, water will boil at different temperatures. For example, at the top of Mount Everest, the pressure is so low that water will boil at about 70°C. Depressions from the Atlantic can easily lower the atmospheric pressure in the UK enough so that water will boil at 99°C – even lower with very deep depressions. Whenever we just talk about “the boiling point” of a liquid, we always assume that it is being measured at exactly 1-atmosphere pressure. In practice, of course, that is rarely exactly true.
Saturated vapor pressure and solids (sublimation)
Solids can also lose particles from their surface to form a vapor, except that in this case, we call the effect sublimation rather than evaporation. Sublimation is the direct change from solid to vapor (or vice versa) without going through the liquid stage.
In most cases, at ordinary temperatures, the saturated vapor pressures of solids range from low to very, very, very low. The forces of attraction in many solids are too high to allow much loss of particles from the surface. However, there are some which do easily form vapors. For example, naphthalene (used in old-fashioned “mothballs” to deter clothes moths) has quite a strong smell. Molecules must be breaking away from the surface as a vapor because otherwise, you wouldn’t be able to smell it. Another fairly common example is solid carbon dioxide (dry ice). This never forms a liquid at atmospheric pressure and always converts directly from solid to vapor. That’s why it is known as dry ice.
Wet saturated vapor
A wet saturated vapor carries liquid globules in suspension; in other words, it is the saturated vapor that contains the maximum quantity of liquid, which is found in the form of minute droplets; examples of this state are the pot steam, fog, and clouds. A wet saturated vapor is a substance in the gaseous state which does not follow the general gas law.
Dry saturated vapor
The dry saturated vapor is free from liquid particles; in other words, it is the saturated vapor with the least quantity of liquid, that is, that which does not contain any water droplet; in these conditions, the vapor is not visible; for example, the sudden disappearance of the fog is due to the transition from wet saturated vapor to dry saturated vapor; it happens in fact that the humidity of the air passes from the saturated humid state (or dew point) to the dry state because the sun’s rays evaporate those droplets and consequently the air becomes transparent. All particles are vaporized, any decrease in vapor temperature or increase in vapor pressure condensates liquid particles in the vapor. The dry saturated vapor is a substance in the gaseous state which does not follow the general gas law.
In superheated vapor, the term “superheated” indicates that the temperature is higher than the boiling point temperature corresponding to the pressure (it is a vapor that is in non-equilibrium conditions). Any additional heat supply, as there is no liquid to vaporize, further raises the temperature of the vapor. The superheated vapor can not exist in contact with the fluid, nor contain fluid particles. An increase in the pressure or decrease in the temperature will not – within limits – condensate out liquid particles in the vapor. Its unique capability is to heat materials to above the normal boiling point of the solvent used. Highly superheated vapors are gases that approximately follow the general gas law.