Chlorine

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Chlorine,  ₁₇Cl

Emission line spectra; 400–700 nm
General properties
Name, symbol	chlorine, Cl
Pronunciation	/ˈklɔəriːn/ or /ˈklɔərɨn/ KLOHR-een or KLOHR-ən
Appearance	pale yellow-green gas
Chlorine in the periodic table
F ↑ Cl ↓ Br sulfur ← chlorine → argon
Atomic number	17
Standard atomic weight	35.45[1] (35.446–35.457)[2]
Element category	diatomic nonmetal
Group, block	group 17 (halogens), p-block
Period	period 3
Electron configuration	[Ne] 3s2 3p5
per shell	2, 8, 7
Physical properties
Phase	gas
Melting point	171.6 K ​(−101.5 °C, ​−150.7 °F)
Boiling point	239.11 K ​(−34.04 °C, ​−29.27 °F)
Density at stp (0 °C and 101.325 kPa)	3.2 g·L−1
when liquid, at b.p.	1.5625 g·cm−3[3]
Critical point	416.9 K, 7.991 MPa
Heat of fusion	(Cl2) 6.406 kJ·mol−1
Heat of vaporization	(Cl2) 20.41 kJ·mol−1
Molar heat capacity	(Cl2) 33.949 J·mol−1·K−1
vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 128 139 153 170 197 239
Atomic properties
Oxidation states	7, 6, 5, 4, 3, 2, 1, −1 ​(a strongly acidic oxide)
Electronegativity	Pauling scale: 3.16
Ionization energies	1st: 1251.2 kJ·mol−1 2nd: 2298 kJ·mol−1 3rd: 3822 kJ·mol−1 (more)
Covalent radius	102±4 pm
Van der Waals radius	175 pm
Miscellanea
Crystal structure	​orthorhombic
Speed of sound	206 m·s−1 (gas, at 0 °C)
Thermal conductivity	8.9×10−3 W·m−1·K−1
Electrical resistivity	>10 Ω·m (at 20 °C)
Magnetic ordering	diamagnetic[4]
CAS Registry Number	7782-50-5
History
Discovery and first isolation	Carl Wilhelm Scheele (1774)
Recognized as an element by	Humphry Davy (1808)
Most stable isotopes
Main article: Isotopes of chlorine
iso NA half-life DM DE (MeV) DP 35Cl 75.77% 35Cl is stable with 18 neutrons 36Cl trace 3.01×105 y β− 0.709 36Ar ε – 36S 37Cl 24.23% 37Cl is stable with 20 neutrons

Chlorine is a chemical element with symbol Cl and atomic number 17. Chlorine is in the halogen group (17) and is the second lightest halogen following fluorine. The element is a yellow-green gas under standard conditions, where it forms diatomic molecules. Chlorine has the highest electron affinity and the third highest electronegativity of all the reactive elements. For this reason, chlorine is a strong oxidizing agent. Free chlorine is rare on Earth, and is usually a result of direct or indirect oxidation by oxygen.

The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630 chlorine gas was first synthesized in a chemical reaction, but not recognized as a fundamentally important substance. Characterization of chlorine gas was made in 1774 by Carl Wilhelm Scheele, who supposed it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρóς (khlôros) "pale green".

Nearly all chlorine in the Earth's crust occurs as chloride in various ionic compounds, including table salt. It is the second most abundant halogen and 21st most abundant chemical element in Earth's crust. Elemental chlorine is commercially produced from brine by electrolysis. The high oxidizing potential of elemental chlorine led commercially to free chlorine's bleaching and disinfectant uses, as well as its many uses of an essential reagent in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, as well as many intermediates for production of plastics and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary.
In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils, as part of the immune response against bacteria. Elemental chlorine at high concentrations is extremely dangerous and poisonous for all living organisms, and was used in World War I as the first gaseous chemical warfare agent.

Characteristics

Physical characteristics of chlorine and its compounds

Chlorine, liquefied under a pressure of 7.4 bar at room temperature, displayed in a quartz ampule embedded in acrylic glass.

At standard temperature and pressure, two chlorine atoms form the diatomic molecule Cl₂.^[5] This is a yellow-green gas that has a distinctive strong odor, familiar to most from common household bleach.^[6] The bonding between the two atoms is relatively weak (only 242.580 ± 0.004 kJ/mol), which makes the Cl₂ molecule highly reactive. The boiling point at standard pressure is around −34 ˚C, but it can be liquefied at room temperature with pressures above 740 kPa (107 psi).^[7]

Although elemental chlorine is yellow-green, the chloride ion, in common with other halide ions, has no color in either minerals or solutions (example, table salt). Similarly, (again as with other halogens) chlorine atoms impart no color to organic chlorides when they replace hydrogen atoms in colorless organic compounds, such as tetrachloromethane. The melting point and density of these compounds is increased by substitution of hydrogen in place of chlorine. Compounds of chlorine with other halogens, however, as well as many chlorine oxides, are visibly colored.

Chemical characteristics

Along with fluorine, bromine, iodine, and astatine, chlorine is a member of the halogen series that forms the group 17 (formerly VII, VIIA, or VIIB) of the periodic table. Chlorine forms compounds with almost all of the elements to give compounds that are usually called chlorides. Chlorine gas reacts with most organic compounds, and will even sluggishly support the combustion of hydrocarbons.^[8]

Hydrolysis of free chlorine or disproportionation in water

At 25 °C and atmospheric pressure, one liter of water dissolves 3.26 g or 1.125 L of gaseous chlorine.^[9] Solutions of chlorine in water contain chlorine (Cl₂), hydrochloric acid, and hypochlorous acid:

Cl₂ + H₂O HCl + HClO

This conversion to the right is called disproportionation, because the ingredient chlorine both increases and decreases in formal oxidation state. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide, and in this way, chlorine bleach is produced.^[10]

Cl₂ + 2 OH⁻ → ClO⁻ + Cl⁻ + H₂O

Chlorine gas only exists in a neutral or acidic solution.

Chemistry and compounds

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero and four in chlorine dioxide (see table below, and also structures in chlorite).^[11] Chlorine typically has a −1 oxidation state in compounds, except for compounds containing fluorine, oxygen and nitrogen, all of which are even more electronegative than chlorine. Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elemental chlorine gas.^[12]

 

Oxidation state	Name	Formula	Characteristic compounds
−1	chlorides	Cl−	ionic chlorides, organic chlorides, hydrochloric acid
0	chlorine	Cl2	elemental chlorine
+1	hypochlorites	ClO−	sodium hypochlorite, calcium hypochlorite
+3	chlorites	ClO− 2	sodium chlorite
+4	chlorine(IV)	ClO 2	chlorine dioxide
+5	chloryl, chlorates	ClO− 3 ClO+ 2	potassium chlorate, chloric acid, dichloryl trisulfate [ClO2]2[S3O10].
+6	chlorine(VI)	Cl 2O 6	dichlorine hexoxide (gas). In liquid or solid disproportionates to mix of +5 and +7 oxidation states, as ionic chloryl perchlorate [ClO 2]+[ClO 4]−
+7	perchlorates	ClO− 4	perchloric acid, perchlorate salts such as magnesium perchlorate, dichlorine heptoxide

Chlorine combines with almost all elements to give chlorides. Compounds with oxygen, nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.^[13] Chloride is one of the most common anions in nature. Hydrogen chloride and its aqueous solution, hydrochloric acid, are produced on megaton scale annually both as valued intermediates but sometimes as undesirable pollutants.

Chlorine oxides

Chlorine forms a variety of oxides, as seen above: chlorine dioxide (ClO₂), dichlorine monoxide (Cl₂O), dichlorine hexoxide (Cl₂O₆), dichlorine heptoxide (Cl₂O₇). The anionic derivatives of these same oxides are also well known including chlorate (ClO−
3), chlorite (ClO−
2), hypochlorite (ClO⁻), and perchlorate (ClO−
4). The acid derivatives of these anions are hypochlorous acid (HOCl), chloric acid (HClO₃) and perchloric acid (HClO₄). The chloroxy cation chloryl (ClO₂⁺) is known and has the same structure as chlorite but with a positive charge and chlorine in the +5 oxidation state.^[14] The compound "chlorine trioxide" does not occur, but rather in gas form is found as the dimeric dichlorine hexoxide (Cl₂O₆) with a +6 oxidation state. This compound in liquid or solid form disproportionates to a mixture of +5 and +7 oxidation states, occurring as the ionic compound chloryl perchlorate, [ClO
2]+[ClO
4]−.^[15]

In hot concentrated alkali solution hypochlorite disproportionates:

2 ClO⁻ → Cl⁻ + ClO−
2

ClO⁻ + ClO−
2 → Cl⁻ + ClO−
3

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated to a high temperature, they undergo a further, final disproportionation:

4 ClO−
3 → Cl⁻ + 3 ClO−
4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:^[16]

Reaction	Electrode potential
Cl− + 2 OH− → ClO− + H2O + 2 e−	+0.89 volts
ClO− + 2 OH− → ClO− 2 + H2O + 2 e−	+0.67 volts
ClO− 2 + 2 OH− → ClO− 3 + H2O + 2 e−	+0.33 volts
ClO− 3 + 2 OH− → ClO− 4 + H2O + 2 e−	+0.35 volts

Each step is accompanied at the cathode by

2 H₂O + 2 e⁻ → 2 OH⁻ + H₂ (−0.83 volts)

Interhalogen compounds

Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively. However, it cannot oxidize fluoride salts to fluorine. It makes a variety of interhalogen compounds, such as the chlorine fluorides, chlorine monofluoride (ClF), chlorine trifluoride (ClF
3), chlorine pentafluoride (ClF
5). Chlorides of bromine and iodine are also known.^[17]

Organochlorine compounds

Chlorine is used extensively in organic chemistry in substitution and addition reactions. Chlorine often imparts many desired properties to an organic compound, in part owing to its electronegativity.Like the other halides, chlorine undergoes electrophilic addition reactions, the most notable one being the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide.^{[citation needed]}

Occurrence

Essentially no chlorine was created in the Big Bang. Chlorine in the universe is created and distributed through the interstellar medium from creation in supernovae, via the r-process.^[18] This chlorine provides the supply found in the Solar System.
In meteorites and on Earth, chlorine is found primarily as the chloride ion which occurs in minerals. In the Earth's crust, chlorine is present at average concentrations of about 126 parts per million,^[19] predominantly in such minerals as halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate).

Chloride is a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground.
Over 2000 naturally occurring organic chlorine compounds are known.^[20]

Isotopes

Chlorine has a wide range of isotopes. The two stable isotopes are ³⁵Cl (75.77%) and ³⁷Cl (24.23%).^[21] Together they give chlorine an atomic weight of 35.4527 g/mol. The half-integer value for chlorine's weight caused some confusion in the early days of chemistry, when it had been postulated that atoms were composed of even units of hydrogen (see Proust's law), and the existence of chemical isotopes was unsuspected.^[22]
Trace amounts of radioactive ³⁶Cl exist in the environment, in a ratio of about 7x10⁻¹³ to 1 with stable isotopes. ³⁶Cl is produced in the atmosphere by spallation of ³⁶Ar by interactions with cosmic ray protons. In the subsurface environment, ³⁶Cl is generated primarily as a result of neutron capture by ³⁵Cl or muon capture by ⁴⁰Ca. ³⁶Cl decays to ³⁶S and to ³⁶Ar, with a combined half-life of 308,000 years. The half-life of this isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of ³⁶Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of ³⁶Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and groundwater, ³⁶Cl is also useful for dating waters less than 50 years before the present. ³⁶Cl has seen use in other areas of the geological sciences, including dating ice and sediments.^[21]

History

The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.^[23] Around 1630, chlorine was recognized as a gas by the Flemish chemist and physician Jan Baptist van Helmont.^[24]

Carl Wilhelm Scheele

Elemental chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and, therefore, he is credited for its discovery.^[25] He called it "dephlogisticated muriatic acid air" since it is a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid").^[25] However, he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory).^[25] He named the new element within this oxide as muriaticum.^[25] Regardless of what he thought, Scheele did isolate chlorine by reacting MnO₂ (as the mineral pyrolusite) with HCl:^[24]

4 HCl + MnO₂ → MnCl₂ + 2 H₂O + Cl₂

Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow green color, and the smell similar to aqua regia.^[26]
At the time, common chemical theory was: any acid is a compound that contains oxygen (still sounding in the German and Dutch names of oxygen: sauerstoff or zuurstof, both translating into English as acid substance), so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum.^[27][28][29]

In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide).^[25] They did not succeed and published a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, but were not convinced.^[30]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it is an element, and not a compound.^[25] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow.^[31] The name halogen, meaning "salt producer," was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger. However, this term was later used as a generic term to describe all the elements in the chlorine family (fluorine, bromine, iodine), after a suggestion by Jöns Jakob Berzelius in 1842.^[32][33] In 1823, Michael Faraday liquefied chlorine for the first time,^[34][35] and demonstrated that what was then known as "solid chlorine" had a structure of chlorine hydrate (Cl₂·H₂O).^[24]

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.^[36][37] Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel (now part of Paris, France), by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of sodium hypochlorite. However, this process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of calcium hypochlorite ("chlorinated lime"), then solid calcium hypochlorite (bleaching powder).^[36] These compounds produced low levels of elemental chlorine, and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidizer. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite.^[38] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of essentially all modern elemental chlorine and sodium hydroxide production (a related low-temperature electrolysis reaction, the Hooker process, is now responsible for bleach and sodium hypochlorite production).

Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti-putrification agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This work is mainly due to Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations for the purpose (for a more complete history, see below). Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) as well as public sanitation (especially of swimming and drinking water).

In 1826, silver chloride was used to produce photographic images for the first time.^[39] Chloroform was first used as an anesthetic in 1847.^[39]

Polyvinyl chloride (PVC) was invented in 1912, initially without a purpose.^[39]

Chlorine gas was first introduced as a weapon on April 22, 1915, at Ypres by the German Army,^[40][41] and the results of this weapon were disastrous because gas masks had not been mass distributed and were tricky to get on quickly.

Production

Liquid chlorine analysis

In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. This method, the chloralkali process industrialized in 1892, now provides essentially all industrial chlorine gas.^[42] Along with chlorine, the method yields hydrogen gas and sodium hydroxide (with sodium hydroxide actually being the most crucial of the three industrial products produced by the process). The process proceeds according to the following chemical equation:^[12]

2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

The electrolysis of chloride solutions all proceed according to the following equations:

Cathode: 2 H⁺_(aq) + 2 e⁻ → H_2(g)

Anode: 2 Cl⁻_(aq) → Cl_2(g) + 2 e⁻

Overall process: 2 NaCl (or KCl) + 2 H₂O → Cl₂ + H₂ + 2 NaOH (or KOH)

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.^[43] The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali but they are not burdened with the problem of preventing mercury discharge into the environment and they are more energy efficient. Membrane cell electrolysis employ permeable membrane as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.^[44] This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine.

Membrane cell process for chloralkali production

Laboratory methods

Small amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid and manganese dioxide. Alternatively a strong acid such as sulfuric acid or hydrochloric acid reacts with sodium hypochlorite solution to release chlorine gas but reacts with sodium chlorate to produce chlorine gas and chlorine dioxide gas as well. In the home, accidents occur when hypochlorite bleach solutions are combined with certain acidic drain-cleaners.

Applications

Production of industrial and consumer products

Principal applications of chlorine are in the production of a wide range of industrial and consumer products.^[45][46] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.

Many important industrial products are produced via organochlorine intermediates. Examples include polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose, and propylene oxide. Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organic compounds. When applied to organic substrates, reaction is often—but not invariably—non-regioselective, and, hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions are common. If the different reaction products are easily separated, e.g., by distillation, substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride, chloroform, and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene, and tetrachloroethylene from 1,2-dichloroethane.

Quantitatively, about 63% and 18% of all elemental chlorine produced is used in the manufacture of organic and inorganic chlorine compounds, respectively.^[42] About 15,000 chlorine compounds are being used commercially.^[26] The remaining 19% is used for bleaches and disinfection products.^[42] The most significant of organic compounds in terms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes, and trichlorobenzenes. The major inorganic compounds include HCl, Cl₂O, HOCl, NaClO₃, chlorinated isocyanurates, AlCl₃, SiCl₄, SnCl₄, PCl₃, PCl₅, POCl₃, AsCl₃, SbCl₃, SbCl₅, BiCl₃, S₂Cl₂, SCl₂, SOCI₂, ClF₃, ICl, ICl₃, TiCl₃, TiCl₄, MoCl₅, FeCl₃, ZnCl₂, etc.^[42][47]

Pulp bleaching was done often with elemental chlorine in the past. However, this tends to produce organochlorine pollution, and today environmental laws make it prohibitive. Chlorine is used either in chlorine dioxide and sodium hypochlorite stages in elemental chlorine free (ECF) bleaching, or not at all (total chlorine free or TCF bleaching).

Public sanitation, disinfection, and antisepsis

Combating putrefaction

Antoine-Germain Labarraque

In France (as elsewhere) there was a need to process animal guts in order to make musical instrument strings, Goldbeater's skin and other products. This was carried out in "gut factories" (boyauderies) as an odiferous and unhealthy business. In or about 1820, the Société d'encouragement pour l'industrie nationale offered a prize for the discovery of a method, chemical or mechanical, that could be used to separate the peritoneal membrane of animal intestines without causing putrefaction.^[48][49] It was won by Antoine-Germain Labarraque, a 44 year-old French chemist and pharmacist who had discovered that Berthollet's chlorinated bleaching solutions ("Eau de Javel") not only destroyed the smell of putrefaction of animal tissue decomposition, but also retarded the decomposition process itself.^[49][50]

Labarraque's research resulted in chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) being employed not only in the boyauderies but also for the routine disinfection and deodorisation of latrines, sewers, markets, abattoirs, anatomical theatres and morgues.^[51] They were also used, with success, in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and for exhumations,^[52] embalming, during outbreaks of epidemic illness, fever, blackleg in cattle, etc.^[48]

Against infection and contagion

Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", and presumed to be transmitted by "miasmas") and also to treat putrefaction of existing wounds, including septic wounds.^[53] In this 1828 work, Labarraque recommended for the doctor to breathe chlorine, wash his hands with chlorinated lime, and even sprinkle chlorinated lime about the patient's bed, in cases of "contagious infection." In 1828, it was well known that some infections were contagious, even though the agency of the microbe was not to be realized or discovered for more than half a century.

During the Paris cholera outbreak of 1832, large quantities of so-called chloride of lime were used to disinfect the capital. This was not simply modern calcium chloride, but contained chlorine gas dissolved in lime-water (dilute calcium hydroxide) to form calcium hypochlorite (chlorinated lime). Labarraque's discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and, by doing so, effectively deodorised the Latin Quarter of Paris.^[54] These "putrid miasmas" were thought by many to be responsible for the spread of "contagion" and "infection" – both words used before the germ theory of infection. The use of chloride of lime was based on destruction of odors and "putrid matter." One source has claimed that chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well feeding the Broad Street pump in 1854 London.^[55] Three reputable sources that described the famous Broad Street pump cholera epidemic do not mention Snow performing any disinfection of water from that well.^[56][57][58]
Instead, one reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump—a common practice in mid-nineteenth century England.^[56]:296

Semmelweis and experiments with antisepsis

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine and chemical base solutions was in 1847, when Ignaz Semmelweis used (first) chlorine-water (simply chlorine dissolved in pure water), then cheaper chlorinated lime solutions, to deodorize the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Semmelweis, still long before the germ theory of disease, had theorized that "cadaveric particles" were somehow transmitting decay from fresh medical cadavers to living patients, and he used the well-known "Labarraque's solutions" as the only known method to remove the smell of decay and tissue decomposition (which he found that soap did not). The solutions proved to be far more effective germicide antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis's (later) celebrated success in stopping the transmission of childbed fever ("puerperal fever") in the maternity wards of Vienna General Hospital in Austria in 1847.^[59]

Much later, during World War I in 1916, a standardized and diluted modification of Labarraque's solution, containing hypochlorite (0.5%) and boric acid as an acidic stabilizer, was developed by Henry Drysdale Dakin (who gave full credit to Labarraque's prior work in this area). Called Dakin's solution, the method of wound irrigation with chlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modern antibiotic era. A modified version of this solution continues to be employed in wound irrigation in the modern era, where it remains effective against multiply antibiotic resistant bacteria (see Century Pharmaceuticals).

Public sanitation

By 1918, the US Department of Treasury called for all drinking water to be disinfected with chlorine. Chlorine is presently an important chemical for water purification (such as in water treatment plants), in disinfectants, and in bleach. Chlorine in water is more than three times as effective as a disinfectant against Escherichia coli than an equivalent concentration of bromine, and is more than six times more effective than an equivalent concentration of iodine.^[60]

Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather sodium hypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. The drawback of using chlorine in swimming pools is that the chlorine reacts with the proteins in human hair and skin (see Hypochlorous acid). Once the chlorine reacts with the hair and skin, it becomes chemically bonded. Even small water supplies are now routinely chlorinated.^[8]

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include hypochlorite solutions, which gradually release chlorine into the water, and compounds like sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred to as "dichlor", and trichloro-s-triazinetrione, sometimes referred to as "trichlor". These compounds are stable while solid and may be used in powdered, granular, or tablet form. When added in small amounts to pool water or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forming hypochlorous acid (HOCl), which acts as a general biocide, killing germs, micro-organisms, algae, and so on.^[61][62]

Use as a weapon

World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres.^[63] As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form hydrochloric acid, an irritant that can be lethal. The damage done by chlorine gas can be prevented by the activated charcoal commonly found in gas masks, or other filtration methods, which makes the overall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide.^[64] After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly phosgene and mustard gas.^[65]

Iraq War

Chlorine gas has also been used by insurgents against the local population and coalition forces in the Iraq War in the form of chlorine bombs. On March 17, 2007, for example, three chlorine-filled trucks were detonated in the Anbar province killing two and sickening over 350.^[66] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.^[67] Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.

Syrian Civil War

There have been allegations of chlorine gas attacks during the Syrian Civil War such as the 2014 Kafr Zita chemical attack.

Islamic State of Iraq and the Levant (ISIL/ISIS)

On October 24, 2014 it was reported that the Islamic State of Iraq and the Levant had used chlorine gas in the town of Duluiyah, Iraq.^[68]

Health effects of the free element and hazards

NFPA 704 "fire diamond"
0 3 0 OX

Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.^[69]

Chlorine is detectable with measuring devices in concentrations of as low as 0.2 parts per million (ppm), and by smell at 3 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can be fatal after a few deep breaths of the gas.^[26] Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.^[70] The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations above 30 ppm, it begins to react with water and cells, which change it into hydrochloric acid (HCl) and hypochlorous acid (HClO).

When used at specified levels for water disinfection, the reaction of chlorine with water is not a major concern for human health. Other materials present in the water may generate disinfection by-products that are associated with negative effects on human health.^[71][72]

Chlorine induced cracking in structural materials

Chlorine "attack" on an acetal resin plumbing joint.

The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to chlorine induced stress corrosion cracking of stainless steel rods used to suspend them.^[73] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic supplies, and stress corrosion cracking caused widespread failures in the USA in the 1980s and 1990s. The picture on the right shows an acetal joint in a water supply system, which, when it fractured, caused substantial physical damage to computers in the labs below the supply. The cracks started at injection molding defects in the joint and slowly grew until finally triggered. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.^[74]

Chlorine-iron fire

The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.^[75][76] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework used to carry chlorine gas is made of steel.^[75][76]

Organochlorine compounds as pollutants

Some organochlorine compounds are serious pollutants. These are produced either as by-products or end products of industrial processes which are persistent in the environment, such as certain chlorinated pesticides and chlorofluorocarbons. Chlorine is added both to pesticides and pharmaceuticals to make the molecules more resistant to enzymatic degradation by bacteria, insects, and mammals, but this property also has the effect of prolonging the residence time of these compounds when they enter the environment. In this respect chlorinated organics have some resemblance to fluorinated organics.

Water vapor

From Wikipedia, the free encyclopedia

Water vapor (H2O)
Invisible water vapor condenses to form visible clouds of liquid water droplets
Systematic name	Water vapor
Liquid state	water
Solid state	ice
Properties[1]
Molecular formula	H2O
Molar mass	18.01528(33) g/mol
Melting point	0.00 °C (273.15 K)[2]
Boiling point	99.98 °C (373.13 K)[2]
specific gas constant	461.5 J/(kg·K)
Heat of vaporization	2.27 MJ/kg
Heat capacity at 300 K	1.864 kJ/(kg·K)[3]

Water vapor, or 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. Unlike other forms of water, water vapor is invisible.^[4] Under typical atmospheric conditions, water vapor is continuously generated by evaporation and removed by condensation. It is lighter than air and triggers convection currents that can lead to clouds.

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.

Being a component of Earth's hydrosphere and hydrologic cycle, it is particularly abundant in Earth's atmosphere where it is also a potent greenhouse gas along with other gases such as carbon dioxide and methane. Use of water vapor, as steam, has been important to humans for cooking and as a major component in energy production and transport systems since the industrial revolution.

Occurrence

Water vapor is a significant component of the Earth's atmosphere and a greenhouse gas. It is also common in the Solar System and by extension, other planetary systems. Its signature has been detected in the atmospheres of the Sun, occurring in sunspots. The presence of water vapor has been detected in the atmospheres of Mercury, Venus, Earth (and Moon^[5]), Mars, Jupiter, Saturn, Uranus and Neptune, the planets of the Solar System, although typically in only trace amounts.

Plume of water aerosol created by condensing vapor on Europa (artist concept) (December 12, 2013).^[6]

Artist's illustration of the signatures of water in exoplanet atmospheres detectable by instruments such as the Hubble Space Telescope.^[7]

Water vapor ejected into the atmosphere of some planetary mass objects (such as water ejected from geological formations such as geysers) may indicate the presence of substantial quantities of subsurface water. Plumes of water vapor have been detected on Europa^[6] (a moon of Jupiter) and are similar to plumes of water vapor detected on Enceladus^[6] (a moon of Saturn). Traces of water vapor have also been detected in the stratosphere of Titan.^[8] Water vapor has been found to be a major constituent of the atmosphere of dwarf planet, Ceres, largest object in the asteroid belt^[9] The detection was made by using the far-infrared abilities of the Herschel Space Observatory.^[10] The finding is unexpected because comets, not asteroids, are typically considered to "sprout jets and plumes." According to one of the scientists, "The lines are becoming more and more blurred between comets and asteroids."^[10] Scientists studying Mars hypothesize that if water moves about the planet, it does so as vapor.^[11]

The brilliance of comet tails comes largely from water vapor. On approach to the sun, the ice many comets carry sublimates to vapor, which reflects light from the sun. Knowing a comet's distance from the sun, astronomers may deduce a comet's water content from its brilliance.^[12]

Water vapor has also been confirmed outside the Solar System. Spectroscopic analysis of HD 209458 b, an extrasolar planet in the constellation Pegasus, provides the first evidence of atmospheric water vapor beyond the Solar System. A star called CW Leonis was found to have a ring of vast quantities of water vapor circling the aging, massive star. A NASA satellite designed to study chemicals in interstellar gas clouds, made the discovery with an onboard spectrometer. Most likely, "the water vapor was vaporized from the surfaces of orbiting comets."^[13] HAT-P-11b a relatively small exoplanet has also been found to possess water vapour.^[14]

General properties of water vaporization

Evaporation and sublimation

Whenever a water molecule leaves a surface and diffuses into a surrounding gas, it is said to have evaporated. Each individual water molecule which transitions between a more associated (liquid) and a less associated (vapor/gas) state does so through the absorption or release of kinetic energy. The aggregate measurement of this kinetic energy transfer is defined as thermal energy and occurs only when there is differential in the temperature of the water molecules. Liquid water that becomes water vapor takes a parcel of heat with it, in a process called evaporative cooling.^[15] The amount of water vapor in the air determines how fast each molecule will return to the surface. When a net evaporation occurs, the body of water will undergo a net cooling directly related to the loss of water.

In the US, the National Weather Service measures the actual rate of evaporation from a standardized "pan" open water surface outdoors, at various locations nationwide. Others do likewise around the world. The US data is collected and compiled into an annual evaporation map.^[16] The measurements range from under 30 to over 120 inches per year. Formulas can be used for calculating the rate of evaporation from a water surface such as a swimming pool.^[17][18] In some countries, the evaporation rate far exceeds the precipitation rate.

Evaporative cooling is restricted by atmospheric conditions. Humidity is the amount of water vapor in the air. The vapor content of air is measured with devices known as hygrometers. The measurements are usually expressed as specific humidity or percent relative humidity. The temperatures of the atmosphere and the water surface determine the equilibrium vapor pressure; 100% relative humidity occurs when the partial pressure of water vapor is equal to the equilibrium vapor pressure. This condition is often referred to as complete saturation. Humidity ranges from 0 gram per cubic metre in dry air to 30 grams per cubic metre (0.03 ounce per cubic foot) when the vapor is saturated at 30 °C.^[19] (See also Absolute Humidity table)

Meteorite Recovery Antarctica

Electron micrograph of tight junctions in blood–brain barrier, prepared by sublimation in freeze-etching process.

Another form of evaporation is sublimation, by which water molecules become gaseous directly, leaving the surface of ice without first becoming liquid water. Sublimation accounts for the slow mid-winter disappearance of ice and snow at temperatures too low to cause melting. Antarctica shows this effect to a unique degree because it is by far the continent with the lowest rate of precipitation on Earth. As a result there are large areas where millennial layers of snow have sublimed, leaving behind whatever non-volatile materials they had contained. This is extremely valuable to certain scientific disciplines, a dramatic example being the collection of meteorites that are left exposed in unparalleled numbers and excellent states of preservation.

Sublimation is of importance in the preparation of certain classes of biological specimens for scanning electron microscopy. Typically the specimens are prepared by cryofixation and freeze-fracture, after which the broken surface is freeze-etched, being eroded by exposure to vacuum till it shows the required level of detail. This technique can display protein molecules, organelle structures and lipid bilayers with very low degrees of distortion.

Condensation

Clouds, formed by condensed water vapor.

Water vapor will only condense onto another surface when that surface is cooler than the dew point temperature, or when the water vapor equilibrium in air has been exceeded. When water vapor condenses onto a surface, a net warming occurs on that surface. The water molecule brings heat energy with it. In turn, the temperature of the atmosphere drops slightly.^[20] In the atmosphere, condensation produces clouds, fog and precipitation (usually only when facilitated by cloud condensation nuclei). The dew point of an air parcel is the temperature to which it must cool before water vapor in the air begins to condense.

Also, a net condensation of water vapor occurs on surfaces when the temperature of the surface is at or below the dew point temperature of the atmosphere. Deposition, the direct formation of ice from water vapor, is a type of condensation. Frost and snow are examples of deposition.

Chemical reactions

A number of chemical reactions have water as a product. If the reactions take place at temperatures higher than the dew point of the surrounding air the water will be formed as vapor and increase the local humidity, if below the dew point local condensation will occur. Typical reactions that result in water formation are the burning of hydrogen or many other hydrocarbons in air itself or in combination with oxygen or other oxidisers.

In a similar fashion other chemical or physical reactions can take place in the presence of water vapor resulting in new chemicals forming such as rust on iron or steel, polymerisation occurring (certain polyurethane foams and cyanoacrylate glues cure with exposure to atmospheric humidity) or forms changing such as where anhydrous chemicals may absorb enough vapor to form a crystalline structure or alter an existing one, sometimes resulting in characteristic color changes that can be used for measurement.

Measurement

Measuring the quantity of water vapor in a medium can be done directly or remotely with varying degrees of accuracy. Remote methods such electromagnetic absorption are possible from satellites above planetary atmospheres. Direct methods may use electronic transducers, moistened thermometers or hygroscopic materials measuring changes in physical properties or dimensions.

	medium	temperature range (degC)	measurement uncertainty	typical measurement frequency	system cost	notes
sling psychrometer	air	−10 to 50	low to moderate	hourly	low
satellite-based spectroscopy	air	−80 to 60	low		very high
capacitive sensor	air/gases	−40 to 50	moderate	2 to 0.05 Hz	medium	prone to becoming saturated/contaminated over time
warmed capacitive sensor	air/gases	−15 to 50	moderate to low	2 to 0.05 Hz (temp dependant)	medium to high	prone to becoming saturated/contaminated over time
resistive sensor	air/gases	−10 to 50	moderate	60 seconds	medium	prone to contamination
lithium chloride dewcell	air	−30 to 50	moderate	continuous	medium	see dewcell
Cobalt(II) chloride	air/gases	0 to 50	high	5 minutes	very low	often used in Humidity indicator card
Absorption spectroscopy	air/gases		moderate		high
Aluminum oxide	air/gases		moderate		medium	see Moisture analysis
silicon oxide	air/gases		moderate		medium	see Moisture analysis
Piezoelectric sorption	air/gases		moderate		medium	see Moisture analysis
Electrolytic	air/gases		moderate		medium	see Moisture analysis
hair tension	air	0 to 40	high	continuous	low to medium	Affected by temperature. Adversely affected by prolonged high concentrations
Nephelometer	air/other gases		low		very high
Goldbeater's skin (cow Peritoneum)	air	−20 to 30	moderate (with corrections)	slow, slower at lower temperatures	low	ref:WMO Guide to Meteorological Instruments and Methods of Observation No. 8 2006, (pages 1.12–1)
Lyman-alpha				high frequency	high	http://amsglossary.allenpress.com/glossary/search?id=lyman-alpha-hygrometer1 Requires frequent calibration
Gravimetric Hygrometer			very low		very high	often called primary source, national independent standards developed in US,UK,EU & Japan
	medium	temperature range (degC)	measurement uncertainty	typical measurement frequency	system cost	notes

Water vapor density

Water vapor is lighter or less dense than dry air.^[21][22] At equivalent temperatures it is buoyant with respect to dry air, whereby the density of dry air at standard temperature and pressure is 1.27 g/l and water vapor at standard temperature and pressure has the much lower density of .804 g/l.

Water vapor and dry air density calculations at 0 °C

The molar mass of water is 18.02 g/mol, as calculated from the sum of the atomic masses of its constituent atoms.

The average molecular mass of air (approx. 79% nitrogen, N₂; 21% oxygen, O₂) is 28.57 g/mol at standard temperature and pressure (STP).

Using Avogadro's Law and the ideal gas law, water vapor and air will have a molar volume of 22.414 L/mol at STP. A molar mass of air and water vapor occupy the same volume of 22.414 litres. The density (mass/volume) of water vapor is 0.804 g/L, which is significantly less than that of dry air at 1.27 g/L at STP. This means water vapor is lighter than air.

STP conditions imply a temperature of 0 °C, at which the ability of water to become vapor is very restricted. Its concentration in air is very low at 0 °C. The red line on the chart to the right is the maximum concentration of water vapor expected for a given temperature. The water vapor concentration increases significantly as the temperature rises, approaching 100% (steam, pure water vapor) at 100 °C. However the difference in densities between air and water vapor would still exist.

Air and water vapor density interactions at equal temperatures

At the same temperature, a column of dry air will be denser or heavier than a column of air containing any water vapor, the molar mass of diatomic nitrogen and diatomic oxygen both being greater than the molar mass of water. Thus, any volume of dry air will sink if placed in a larger volume of moist air. Also, a volume of moist air will rise or be buoyant if placed in a larger region of dry air. As the temperature rises the proportion of water vapor in the air increases, and its buoyancy will increase. The increase in buoyancy can have a significant atmospheric impact, giving rise to powerful, moisture rich, upward air currents when the air temperature and sea temperature reaches 25 °C or above. This phenomenon provides a significant motivating force for cyclonic and anticyclonic weather systems (typhoons and hurricanes).

Water vapor and respiration or breathing

Water vapor is a by-product of respiration in plants and animals. Its contribution to the pressure, increases as its concentration increases. Its partial pressure contribution to air pressure increases, lowering the partial pressure contribution of the other atmospheric gases (Dalton's Law). The total air pressure must remain constant. The presence of water vapor in the air naturally dilutes or displaces the other air components as its concentration increases.

This can have an effect on respiration. In very warm air (35 °C) the proportion of water vapor is large enough to give rise to the stuffiness that can be experienced in humid jungle conditions or in poorly ventilated buildings.

Lifting gas

Water vapor has lower density than that of air and is therefore buoyant in air but has lower vapor pressure than that of air. When water vapor is used as a lifting gas for use by a thermal airship the water vapor is heated to form steam so that its vapor pressure is greater than the surrounding air pressure in order to pressurize and to maintain the shape a theoretical "steam balloon", which yields approximately 60% the lift of helium and twice that of hot air.^[23]

General discussion

The amount of water vapor in an atmosphere is constrained by the restrictions of partial pressures and temperature. Dew point temperature and relative humidity act as guidelines for the process of water vapor in the water cycle. Energy input, such as sunlight, can trigger more evaporation on an ocean surface or more sublimation on a chunk of ice on top of a mountain. The balance between condensation and evaporation gives the quantity called vapor partial pressure.

The maximum partial pressure (saturation pressure) of water vapor in air varies with temperature of the air and water vapor mixture. A variety of empirical formulas exist for this quantity; the most used reference formula is the Goff-Gratch equation for the SVP over liquid water below zero degree Celsius:

$\log _{{10}}\left(p\right)=$	$-7.90298\left({\frac {373.16}{T}}-1\right)+5.02808\log _{{10}}{\frac {373.16}{T}}$
	$-1.3816\times 10^{{-7}}\left(10^{{11.344\left(1-{\frac {T}{373.16}}\right)}}-1\right)$
	$+8.1328\times 10^{{-3}}\left(10^{{-3.49149\left({\frac {373.16}{T}}-1\right)}}-1\right)$
	$+\log _{{10}}\left(1013.246\right)$

Where T, temperature of the moist air, is given in units of kelvins, and p is given in units of millibars (hectopascals).

The formula is valid from about −50 to 102 °C; however there are a very limited number of measurements of the vapor pressure of water over supercooled liquid water. There are a number of other formulae which can be used.^[24]

Under certain conditions, such as when the boiling temperature of water is reached, a net evaporation will always occur during standard atmospheric conditions regardless of the percent of relative humidity. This immediate process will dispel massive amounts of water vapor into a cooler atmosphere.

Exhaled air is almost fully at equilibrium with water vapor at the body temperature. In the cold air the exhaled vapor quickly condenses, thus showing up as a fog or mist of water droplets and as condensation or frost on surfaces. Forcibly condensing these water droplets from exhaled breath is the basis of exhaled breath condensate, an evolving medical diagnostic test.

Controlling water vapor in air is a key concern in the heating, ventilating, and air-conditioning (HVAC) industry. Thermal comfort depends on the moist air conditions. Non-human comfort situations are called refrigeration, and also are affected by water vapor. For example many food stores, like supermarkets, utilize open chiller cabinets, or food cases, which can significantly lower the water vapor pressure (lowering humidity). This practice delivers several benefits as well as problems.

Water vapor in Earth's atmosphere

Evidence for increasing amounts of stratospheric water vapor over time in Boulder, Colorado.

Play media

These maps show the average amount of water vapor in a column of atmosphere in a given month. The units are given in centimeters, which is the equivalent amount of water that could be produced if all the water vapor in the column were to condense. The lowest amounts of water vapor (0 centimeters) appear in yellow, and the highest amounts (6 centimeters) appear in dark blue. Areas of missing data appear in shades of gray. The maps are based on data collected by the Moderate Resolution Imaging Spectroradiometer (MODIS) sensor on NASA’s Aqua satellite. The most noticeable pattern in the time series is the influence of seasonal temperature changes and incoming sunlight on water vapor. In the tropics, a band of extremely humid air wobbles north and south of the equator as the seasons change. This band of humidity is part of the Intertropical Convergence Zone, where the easterly trade winds from each hemisphere converge and produce near-daily thunderstorms and clouds. Farther from the equator, water vapor concentrations are high in the hemisphere experiencing summer and low in the one experiencing winter. Another pattern that shows up in the time series is that water vapor amounts over land areas decrease more in winter months than adjacent ocean areas do. This is largely because air temperatures over land drop more in the winter than temperatures over the ocean. Water vapor condenses more rapidly in colder air.^[25]

Gaseous water represents a small but environmentally significant constituent of the atmosphere. The percentage water vapor in surface air varies from .01% at -42 °C (-44 °F)^[26] to 4.24% when the dew point is 30 °C (86 °F).^[27] Approximately 99.13% of it is contained in the troposphere. The condensation of water vapor to the liquid or ice phase is responsible for clouds, rain, snow, and other precipitation, all of which count among the most significant elements of what we experience as weather. Less obviously, the latent heat of vaporization, which is released to the atmosphere whenever condensation occurs, is one of the most important terms in the atmospheric energy budget on both local and global scales. For example, latent heat release in atmospheric convection is directly responsible for powering destructive storms such as tropical cyclones and severe thunderstorms. Water vapor is also the most potent greenhouse gas owing to the presence of the hydroxyl bond which strongly absorbs in the infra-red region of the light spectrum.

Water in Earth's atmosphere is not merely below its boiling point (100 °C), but at altitude it goes below its freezing point (0 °C), due to water's highly polar attraction. When combined with its quantity, water vapor then has a relevant dew point and frost point, unlike e. g., carbon dioxide and methane. Water vapor thus has a scale height a fraction of that of the bulk atmosphere,^{[28][29][30][31]} as the water condenses and exits, primarily in the troposphere, the lowest layer of the atmosphere.^[32] Carbon dioxide and methane, being non-polar, rise above water vapor. The absorption and emission of CO₂ and CH₄ contribute to Earth's emission to space, and thus the planetary greenhouse effect.^[33] Conversely, adding water vapor at high altitudes has a disproportionate impact, which is why methane (rising, then oxidizing to CO₂ and two water molecules) and jet traffic have disproportionately high warming effects.

It is less clear how cloudiness would respond to a warming climate; depending on the nature of the response, clouds could either further amplify or partly mitigate warming from long-lived greenhouse gases.

In the absence of other greenhouse gases, Earth's water vapor would condense to the surface;^{[34][35][36][37][38][39]} this has likely happened, possibly more than once. Scientists thus distinguish between non-condensable (driving) and condensable (driven) greenhouse gases- i. e., the above water vapor feedback.^[40][41][42]

Fog and clouds form through condensation around cloud condensation nuclei. In the absence of nuclei, condensation will only occur at much lower temperatures. Under persistent condensation or deposition, cloud droplets or snowflakes form, which precipitate when they reach a critical mass.

The water content of the atmosphere as a whole is constantly depleted by precipitation. At the same time it is constantly replenished by evaporation, most prominently from seas, lakes, rivers, and moist earth. Other sources of atmospheric water include combustion, respiration, volcanic eruptions, the transpiration of plants, and various other biological and geological processes. The mean global content of water vapor in the atmosphere is roughly sufficient to cover the surface of the planet with a layer of liquid water about 25 mm deep. The mean annual precipitation for the planet is about 1 meter, which implies a rapid turnover of water in the air – on average, the residence time of a water molecule in the troposphere is about 9 to 10 days.

Episodes of surface geothermal activity, such as volcanic eruptions and geysers, release variable amounts of water vapor into the atmosphere. Such eruptions may be large in human terms, and major explosive eruptions may inject exceptionally large masses of water exceptionally high into the atmosphere, but as a percentage of total atmospheric water, the role of such processes is minor. The relative concentrations of the various gases emitted by volcanoes varies considerably according to the site and according to the particular event at any one site. However, water vapor is consistently the commonest volcanic gas; as a rule, it comprises more than 60% of total emissions during a subaerial eruption.^[43]

Atmospheric water vapor content is expressed using various measures. These include vapor pressure, specific humidity, mixing ratio, dew point temperature, and relative humidity.

Radar and satellite imaging

MODIS/Terra global mean atmospheric water vapor

Because water molecules absorb microwaves and other radio wave frequencies, water in the atmosphere attenuates radar signals.^[44] In addition, atmospheric water will reflect and refract signals to an extent that depends on whether it is vapor, liquid or solid.

Generally, radar signals lose strength progressively the farther they travel through the troposphere. Different frequencies attenuate at different rates, such that some components of air are opaque to some frequencies and transparent to others. Radio waves used for broadcasting and other communication experience the same effect.

Water vapor reflects radar to a lesser extent than do water's other two phases. In the form of drops and ice crystals, water acts as a prism, which it does not do as an individual molecule; however, the existence of water vapor in the atmosphere causes the atmosphere to act as a giant prism.^[45]

A comparison of GOES-12 satellite images shows the distribution of atmospheric water vapor relative to the oceans, clouds and continents of the Earth. Vapor surrounds the planet but is unevenly distributed.

Lightning generation

Water vapor plays a key role in lightning production in the atmosphere. From cloud physics, usually, clouds are the real generators of static charge as found in Earth's atmosphere. But the ability, or capability of clouds to hold massive amounts of electrical energy is directly related to the amount of water vapor present in the local system.

The amount of water vapor directly controls the permittivity of the air. During times of low humidity, static discharge is quick and easy. During times of higher humidity, fewer static discharges occur.
Permittivity and capacitance work hand in hand to produce the megawatt outputs of lightning.^[46]

After a cloud, for instance, has started its way to becoming a lightning generator, atmospheric water vapor acts as a substance (or insulator) that decreases the ability of the cloud to discharge its electrical energy. Over a certain amount of time, if the cloud continues to generate and store more static electricity, the barrier that was created by the atmospheric water vapor will ultimately break down from the stored electrical potential energy.^[47][48] This energy will be released to a locally, oppositely charged region in the form of lightning. The strength of each discharge is directly related to the atmospheric permittivity, capacitance, and the source's charge generating ability.^[49]