Vapor

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An ampule of nitrogen oxide vapor: brown nitrogen dioxide and colorless dinitrogen tetroxide, in equilibrium Nitrogen dioxide gas.jpg
An ampule of nitrogen oxide vapor: brown nitrogen dioxide and colorless dinitrogen tetroxide, in equilibrium

In physics, a vapor (American English) or vapour (Commonwealth English; see spelling differences) is a substance in the gas phase at a temperature lower than its critical temperature, [1] which means that the vapor can be condensed to a liquid by increasing the pressure on it without reducing the temperature of the vapor. A vapor is different from an aerosol. [2] An aerosol is a suspension of tiny particles of liquid, solid, or both within a gas. [2]

Contents

For example, water has a critical temperature of 647 K (374 °C; 705 °F), which is the highest temperature at which liquid water can exist at any pressure. In the atmosphere at ordinary temperatures gaseous water (known as water vapor) will condense into a liquid if its partial pressure is increased sufficiently.

A vapor may co-exist with a liquid (or a solid). When this is true, the two phases will be in equilibrium, and the gas-partial pressure will be equal to the equilibrium vapor pressure of the liquid (or solid). [1]

Properties

The vapor-liquid critical point in a pressure-temperature phase diagram is at the high-temperature extreme of the liquid-gas phase boundary (the dotted green line gives the anomalous behaviour of water). Phase-diag2.svg
The vapor-liquid critical point in a pressure-temperature phase diagram is at the high-temperature extreme of the liquid–gas phase boundary (the dotted green line gives the anomalous behaviour of water).

Vapor refers to a gas phase at a temperature where the same substance can also exist in the liquid or solid state, below the critical temperature of the substance. (For example, water has a critical temperature of 374 °C (647 K), which is the highest temperature at which liquid water can exist.) If the vapor is in contact with a liquid or solid phase, the two phases will be in a state of equilibrium. The term gas refers to a compressible fluid phase. Fixed gases are gases for which no liquid or solid can form at the temperature of the gas, such as air at typical ambient temperatures. A liquid or solid does not have to boil to release a vapor.

Vapor is responsible for the familiar processes of cloud formation and condensation. It is commonly employed to carry out the physical processes of distillation and headspace extraction from a liquid sample prior to gas chromatography.

The constituent molecules of a vapor possess vibrational, rotational, and translational motion. These motions are considered in the kinetic theory of gases.

Vapor pressure

Liquid-vapor equilibrium Binary Boiling Point Diagram new.svg
Liquid–vapor equilibrium
If the vapor pressure exceeds the equilibrium value, it becomes supersaturated and condenses on any available nucleation sites e. g. particles of dust. This principle is used in cloud chambers, where particles of radiation are visualized because they nucleate formation of water droplets. Vapor being used in a cloud chamber.jpg
If the vapor pressure exceeds the equilibrium value, it becomes supersaturated and condenses on any available nucleation sites e. g. particles of dust. This principle is used in cloud chambers, where particles of radiation are visualized because they nucleate formation of water droplets.

The vapor pressure is the equilibrium pressure from a liquid or a solid at a specific temperature. The equilibrium vapor pressure of a liquid or solid is not affected by the amount of contact with the liquid or solid interface.

The normal boiling point of a liquid is the temperature at which the vapor pressure is equal to normal atmospheric pressure. [1]

For two-phase systems (e.g., two liquid phases), the vapor pressure of the individual phases are equal. In the absence of stronger inter-species attractions between like-like or like-unlike molecules, the vapor pressure follows Raoult's law, which states that the partial pressure of each component is the product of the vapor pressure of the pure component and its mole fraction in the mixture. The total vapor pressure is the sum of the component partial pressures. [3]

Examples

Invisible water vapor condenses to form visible water droplets called mist Crepuscular Rays Beam through the Mist Blown from Takkakaw Falls.jpg
Invisible water vapor condenses to form visible water droplets called mist

E-cigarettes produce aerosols, not vapors. [2]

Measuring vapor

Since it is in the gas phase, the amount of vapor present is quantified by the partial pressure of the gas. Also, vapors obey the barometric formula in a gravitational field, just as conventional atmospheric gases do.

See also

Related Research Articles

<span class="mw-page-title-main">Boiling point</span> Temperature at which a substance changes from liquid into vapor

The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor.

<span class="mw-page-title-main">Distillation</span> Method of separating mixtures

Distillation, also classical distillation, is the process of separating the component substances of a liquid mixture of two or more chemically discrete substances; the separation process is realized by way of the selective boiling of the mixture and the condensation of the vapors in a still.

<span class="mw-page-title-main">Evaporation</span> Type of vaporization of a liquid that occurs from its surface; surface phenomenon

Evaporation is a type of vaporization that occurs on the surface of a liquid as it changes into the gas phase. A high concentration of the evaporating substance in the surrounding gas significantly slows down evaporation, such as when humidity affects rate of evaporation of water. When the molecules of the liquid collide, they transfer energy to each other based on how they collide. When a molecule near the surface absorbs enough energy to overcome the vapor pressure, it will escape and enter the surrounding air as a gas. When evaporation occurs, the energy removed from the vaporized liquid will reduce the temperature of the liquid, resulting in evaporative cooling.

<span class="mw-page-title-main">Phase (matter)</span> Region of uniform physical properties

In the physical sciences, a phase is a region of material that is chemically uniform, physically distinct, and (often) mechanically separable. In a system consisting of ice and water in a glass jar, the ice cubes are one phase, the water is a second phase, and the humid air is a third phase over the ice and water. The glass of the jar is another separate phase.

<span class="mw-page-title-main">Solution (chemistry)</span> Homogeneous mixture of a solute and a solvent

In chemistry, a solution is a special type of homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent. If the attractive forces between the solvent and solute particles are greater than the attractive forces holding the solute particles together, the solvent particles pull the solute particles apart and surround them. These surrounded solute particles then move away from the solid solute and out into the solution. The mixing process of a solution happens at a scale where the effects of chemical polarity are involved, resulting in interactions that are specific to solvation. The solution usually has the state of the solvent when the solvent is the larger fraction of the mixture, as is commonly the case. One important parameter of a solution is the concentration, which is a measure of the amount of solute in a given amount of solution or solvent. The term "aqueous solution" is used when one of the solvents is water.

<span class="mw-page-title-main">Triple point</span> Thermodynamic point where three matter phases exist

In thermodynamics, the triple point of a substance is the temperature and pressure at which the three phases of that substance coexist in thermodynamic equilibrium. It is that temperature and pressure at which the sublimation, fusion, and vaporisation curves meet. For example, the triple point of mercury occurs at a temperature of −38.8 °C (−37.8 °F) and a pressure of 0.165 mPa.

<span class="mw-page-title-main">Enthalpy of vaporization</span> Energy to convert a liquid substance to a gas at a given pressure

In thermodynamics, the enthalpy of vaporization, also known as the (latent) heat of vaporization or heat of evaporation, is the amount of energy (enthalpy) that must be added to a liquid substance to transform a quantity of that substance into a gas. The enthalpy of vaporization is a function of the pressure and temperature at which the transformation takes place.

<span class="mw-page-title-main">Vapor pressure</span> Pressure exerted by a vapor in thermodynamic equilibrium

Vapor pressure or equilibrium vapor pressure is the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases at a given temperature in a closed system. The equilibrium vapor pressure is an indication of a liquid's thermodynamic tendency to evaporate. It relates to the balance of particles escaping from the liquid in equilibrium with those in a coexisting vapor phase. A substance with a high vapor pressure at normal temperatures is often referred to as volatile. The pressure exhibited by vapor present above a liquid surface is known as vapor pressure. As the temperature of a liquid increases, the attractive interactions between liquid molecules become less significant in comparison to the entropy of those molecules in the gas phase, increasing the vapor pressure. Thus, liquids with strong intermolecular interactions are likely to have smaller vapor pressures, with the reverse true for weaker interactions.

<span class="mw-page-title-main">Partial pressure</span> Pressure of a component gas in a mixture

In a mixture of gases, each constituent gas has a partial pressure which is the notional pressure of that constituent gas as if it alone occupied the entire volume of the original mixture at the same temperature. The total pressure of an ideal gas mixture is the sum of the partial pressures of the gases in the mixture.

<span class="mw-page-title-main">Condensation</span> Change of state of matter from a gas phase into a liquid phase

Condensation is the change of the state of matter from the gas phase into the liquid phase, and is the reverse of vaporization. The word most often refers to the water cycle. It can also be defined as the change in the state of water vapor to liquid water when in contact with a liquid or solid surface or cloud condensation nuclei within the atmosphere. When the transition happens from the gaseous phase into the solid phase directly, the change is called deposition.

<span class="mw-page-title-main">Phase diagram</span> Chart used to show conditions at which physical phases of a substance occur

A phase diagram in physical chemistry, engineering, mineralogy, and materials science is a type of chart used to show conditions at which thermodynamically distinct phases occur and coexist at equilibrium.

<span class="mw-page-title-main">Water vapor</span> Gaseous phase of water

Water vapor, water vapour or aqueous vapor is the gaseous phase of water. It is one state of water within the hydrosphere. Water vapor can be produced from the evaporation or boiling of liquid water or from the sublimation of ice. Water vapor is transparent, like most constituents of the atmosphere. Under typical atmospheric conditions, water vapor is continuously generated by evaporation and removed by condensation. It is less dense than most of the other constituents of air and triggers convection currents that can lead to clouds and fog.

<span class="mw-page-title-main">Latent heat</span> Thermodynamic phase transition energy

Latent heat is energy released or absorbed, by a body or a thermodynamic system, during a constant-temperature process—usually a first-order phase transition, like melting or condensation.

Vaporization 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 a surface phenomenon, where as boiling is a bulk phenomenon.

<span class="mw-page-title-main">Flash evaporation</span> Partial vapor due to reduction in pressure

Flash evaporation is the partial vapor that occurs when a saturated liquid stream undergoes a reduction in pressure by passing through a throttling valve or other throttling device. This process is one of the simplest unit operations. If the throttling valve or device is located at the entry into a pressure vessel so that the flash evaporation occurs within the vessel, then the vessel is often referred to as a flash drum.

In thermodynamics, the phase rule is a general principle governing "pVT" systems, whose thermodynamic states are completely described by the variables pressure, volume and temperature, in thermodynamic equilibrium. If F is the number of degrees of freedom, C is the number of components and P is the number of phases, then

<span class="mw-page-title-main">Sublimation (phase transition)</span> Transition from solid to gas

Sublimation is the transition of a substance directly from the solid to the gas state, without passing through the liquid state. The verb form of sublimation is sublime, or less preferably, sublimate. Sublimate also refers to the product obtained by sublimation. The point at which sublimation occurs rapidly is called critical sublimation point, or simply sublimation point. Notable examples include sublimation of dry ice at room temperature and atmospheric pressure, and that of solid iodine with heating.

<span class="mw-page-title-main">Critical point (thermodynamics)</span> Temperature and pressure point where phase boundaries disappear

In thermodynamics, a critical point is the end point of a phase equilibrium curve. One example is the liquid–vapor critical point, the end point of the pressure–temperature curve that designates conditions under which a liquid and its vapor can coexist. At higher temperatures, the gas cannot be liquefied by pressure alone; at lower pressures, it cannot be liquefied by temperature alone. At the critical point, defined by a critical temperatureTc and a critical pressurepc, phase boundaries vanish. Other examples include the liquid–liquid critical points in mixtures, and the ferromagnet–paramagnet transition in the absence of an external magnetic field.

<span class="mw-page-title-main">Volatility (chemistry)</span> Tendency of a substance to vaporize

In chemistry, volatility is a material quality which describes how readily a substance vaporizes. At a given temperature and pressure, a substance with high volatility is more likely to exist as a vapour, while a substance with low volatility is more likely to be a liquid or solid. Volatility can also describe the tendency of a vapor to condense into a liquid or solid; less volatile substances will more readily condense from a vapor than highly volatile ones. Differences in volatility can be observed by comparing how fast substances within a group evaporate when exposed to the atmosphere. A highly volatile substance such as rubbing alcohol will quickly evaporate, while a substance with low volatility such as vegetable oil will remain condensed. In general, solids are much less volatile than liquids, but there are some exceptions. Solids that sublimate such as dry ice or iodine can vaporize at a similar rate as some liquids under standard conditions.

The Glossary of fuel cell terms lists the definitions of many terms used within the fuel cell industry. The terms in this fuel cell glossary may be used by fuel cell industry associations, in education material and fuel cell codes and standards to name but a few.

References

  1. 1 2 3 R. H. Petrucci, W. S. Harwood, and F. G. Herring, General Chemistry, Prentice-Hall, 8th ed. 2002, p. 483–86.
  2. 1 2 3 Cheng, T. (2014). "Chemical evaluation of electronic cigarettes". Tobacco Control. 23 (Supplement 2): ii11–ii17. doi:10.1136/tobaccocontrol-2013-051482. ISSN   0964-4563. PMC   3995255 . PMID   24732157.
  3. Thomas Engel and Philip Reid, Physical Chemistry, Pearson Benjamin-Cummings, 2006, p.194
  4. Ferguson, Lon H.; Janicak, Christopher A. (2005-09-01). Fundamentals of Fire Protection for the Safety Professional. Government Institutes. ISBN   9781591919605.