Chlorine

Not to be confused withchloride.

This article is about the chemical element. For other uses, seeChlorine (disambiguation).

"Cl" and "Cl2" redirect here. For other uses, seeCL (disambiguation)andCL2 (disambiguation).

Chlorineis achemical element;it hassymbolClandatomic number17. The second-lightest of thehalogens,it appears betweenfluorineandbrominein the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strongoxidising agent:among the elements, it has the highestelectron affinityand the third-highestelectronegativityon the revisedPauling scale,behind onlyoxygenand fluorine.

Chlorine,₁₇Cl

Chlorine
Pronunciation	/ˈklɔːriːn,-aɪn/​(KLOR-een, -⁠eyen)
Appearance	pale yellow-green gas
Standard atomic weightAr°(Cl)
	[35.446,35.457] 35.45±0.01(abridged)
Chlorine in theperiodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson F ↑ Cl ↓ Br sulfur←chlorine→argon
Atomic number(Z)	17
Group	group 17 (halogens)
Period	period 3
Block	p-block
Electron configuration	[Ne] 3s23p5
Electrons per shell	2, 8, 7
Physical properties
PhaseatSTP	gas
Melting point	(Cl2) 171.6K​(−101.5 °C, ​−150.7 °F)
Boiling point	(Cl2) 239.11 K ​(−34.04 °C, ​−29.27 °F)
Density(at STP)	3.2 g/L
when liquid (atb.p.)	1.5625 g/cm3
Triple point	172.22 K, ​1.392 kPa
Critical point	416.9 K, 7.991 MPa
Heat of fusion	(Cl2) 6.406kJ/mol
Heat of vaporisation	(Cl2) 20.41 kJ/mol
Molar heat capacity	(Cl2) 33.949 J/(mol·K)
Vapour pressure P(Pa) 1 10 100 1 k 10 k 100 k atT(K) 128 139 153 170 197 239
Atomic properties
Oxidation states	common:−1, +1, +3, +5, +7 +2,+4,+6
Electronegativity	Pauling scale: 3.16
Ionisation energies	1st: 1251.2 kJ/mol 2nd: 2298 kJ/mol 3rd: 3822 kJ/mol (more)
Covalent radius	102±4pm
Van der Waals radius	175 pm
Spectral linesof chlorine
Other properties
Natural occurrence	primordial
Crystal structure	​orthorhombic(oS8)
Lattice constants	a= 630.80 pm b= 455.83 pm c= 815.49 pm (at triple point)
Thermal conductivity	8.9×10−3W/(m⋅K)
Electrical resistivity	>10 Ω⋅m (at 20 °C)
Magnetic ordering	diamagnetic
Molar magnetic susceptibility	−40.5×10−6cm3/mol
Speed of sound	206m/s(gas, at 0 °C)
CAS Number	Cl2:7782-50-5
History
Discoveryand first isolation	Carl Wilhelm Scheele(1774)
Recognized as anelementby	Humphry Davy(1808)
Isotopes of chlorine
Main isotopes Decay abun­dance half-life(t1/2) mode pro­duct 35Cl 76% stable 36Cl trace 3.01×105y β− 36Ar ε 36S 37Cl 24% stable

Chlorine played an important role in the experiments conducted by medievalalchemists,which commonly involved the heating of chloridesaltslikeammonium chloride(sal ammoniac) andsodium chloride(common salt), producing various chemical substances containing chlorine such ashydrogen chloride,mercury(II) chloride(corrosive sublimate), andaqua regia.However, the nature of free chlorine gas as a separate substance was only recognised around 1630 byJan Baptist van Helmont.Carl Wilhelm Scheelewrote a description of chlorine gas in 1774, supposing it to be anoxideof a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed bySir Humphry Davyin 1810, who named it after theAncient Greekχλωρός(khlōrós,"pale green" ) because of its colour.

Because of its great reactivity, all chlorine in the Earth's crust is in the form ofionicchloridecompounds, which includes table salt. It is thesecond-most abundanthalogen (after fluorine) and 20th most abundant element in Earth's crust. These crystal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater.

Elemental chlorine is commercially produced frombrinebyelectrolysis,predominantly in thechloralkali process.The high oxidising potential of elemental chlorine led to the development of commercialbleachesanddisinfectants,and areagentfor many processes 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 aspolyvinyl chloride(PVC), many intermediates for the production ofplastics,and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly inswimming poolsto keep themsanitary.Elemental chlorine at highconcentrationis extremely dangerous, andpoisonousto most living organisms. As achemical warfareagent, chlorine was first used inWorld War Ias apoison gasweapon.

In the form of chlorideions,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 theupper atmosphere,chlorine-containing organic molecules such aschlorofluorocarbonshave been implicated inozone depletion.Small quantities of elemental chlorine are generated by oxidation of chloride ions inneutrophilsas part of animmune systemresponse against bacteria.

History

The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence thatrock saltwas used as early as 3000 BC andbrineas early as 6000 BC.

Early discoveries

Around 900, the authors of the Arabic writings attributed toJabir ibn Hayyan(Latin: Geber) and the Persian physician and alchemistAbu Bakr al-Razi(c.865–925, Latin: Rhazes) were experimenting withsal ammoniac(ammonium chloride), which when it was distilled together withvitriol(hydratedsulfatesof various metals) producedhydrogen chloride.However, it appears that in these early experiments with chloridesalts,the gaseous products were discarded, and hydrogen chloride may have been produced many times before it was discovered that it can be put to chemical use.One of the first such uses was the synthesis ofmercury(II) chloride(corrosive sublimate), whose production from the heating ofmercuryeither withalumand ammonium chloride or with vitriol and sodium chloride was first described in theDe aluminibus et salibus( "On Alums and Salts", an eleventh- or twelfth century Arabic text falsely attributed to Abu Bakr al-Razi andtranslated into Latin in the second half of the twelfth centurybyGerard of Cremona,1144–1187).Another important development was the discovery bypseudo-Geber(in theDe inventione veritatis,"On the Discovery of Truth", after c. 1300) that by adding ammonium chloride tonitric acid,a strong solvent capable of dissolving gold (i.e.,aqua regia) could be produced.Althoughaqua regiais an unstable mixture that continually gives off fumes containing free chlorine gas, this chlorine gas appears to have been ignored until c. 1630, when its nature as a separate gaseous substance was recognised by theBrabantianchemist and physicianJan Baptist van Helmont.

Carl Wilhelm Scheele,discoverer of chlorine

Isolation

The element was first studied in detail in 1774 by Swedish chemistCarl Wilhelm Scheele,and he is credited with the discovery.Scheele produced chlorine by reactingMnO₂(as the mineralpyrolusite) with HCl:

4 HCl + MnO₂→ MnCl₂+ 2 H₂O + Cl₂

Scheele observed several of the properties of chlorine: the bleaching effect onlitmus,the deadly effect on insects, the yellow-green colour, and the smell similar toaqua regia.He called it "dephlogisticated muriatic acid air"since it is a gas (then called" airs ") and it came fromhydrochloric acid(then known as "muriatic acid" ).He failed to establish chlorine as an element.

Common chemical theory at that time held that an acid is a compound that contains oxygen (remnants of this survive in the German and Dutch names ofoxygen:sauerstofforzuurstof,both translating into English asacid substance), so a number of chemists, includingClaude Berthollet,suggested that Scheele'sdephlogisticated muriatic acid airmust be a combination of oxygen and the yet undiscovered element,muriaticum.

In 1809,Joseph Louis Gay-LussacandLouis-Jacques Thénardtried to decomposedephlogisticated muriatic acid airby reacting it with charcoal to release the free elementmuriaticum(and carbon dioxide).They did not succeed and published a report in which they considered the possibility thatdephlogisticated muriatic acid airis an element, but were not convinced.

In 1810,Sir Humphry Davytried the same experiment again, and concluded that the substance was an element, and not a compound.He announced his results to the Royal Society on 15 November that year.At that time, he named this new element "chlorine", from the Greek word χλωρος (chlōros,"green-yellow" ), in reference to its colour.The name "halogen",meaning" salt producer ", was originally used for chlorine in 1811 byJohann Salomo Christoph Schweigger.This term was later used as a generic term to describe all the elements in the chlorine family (fluorine, bromine, iodine), after a suggestion byJöns Jakob Berzeliusin 1826.In 1823,Michael Faradayliquefied chlorine for the first time,and demonstrated that what was then known as "solid chlorine" had a structure ofchlorine hydrate(Cl₂·H₂O).

Later uses

Chlorine gas was first used by French chemistClaude Bertholletto bleach textiles in 1785.Modern bleaches resulted from further work by Berthollet, who first producedsodium hypochloritein 1789 in his laboratory in the town ofJavel(now part ofParis,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 ofsodium hypochlorite.This process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialistCharles Tennantfirst produced a solution ofcalcium hypochlorite( "chlorinated lime" ), then solid calcium hypochlorite (bleaching powder).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 ofbrineto producesodium hydroxideand chlorine gas, which then mixed to form sodium hypochlorite.This is known as thechloralkali process,first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide.In 1884 Chemischen Fabrik Griesheim of Germany developed anotherchloralkali processwhich entered commercial production in 1888.

Elemental chlorine solutions dissolved inchemically basicwater (sodium andcalcium hypochlorite) were first used as anti-putrefactionagents anddisinfectantsin the 1820s, in France, long before the establishment of thegerm theory of disease.This practice was pioneered byAntoine-Germain Labarraque,who adapted Berthollet's "Javel water" bleach and other chlorine preparations.Elemental chlorine has since served a continuous function in topicalantisepsis(wound irrigation solutions and the like) and public sanitation, particularly in swimming and drinking water.

Chlorine gas was first used as a weapon on April 22, 1915, at theSecond Battle of Ypresby theGerman Army.The effect on the allies was devastating because the existinggas maskswere difficult to deploy and had not been broadly distributed.

Properties

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

Gaseous chlorine stored inside a 500-mL water bottle. It is not recommended to store chlorine in this manner.

Solid chlorine at −150 °C

Chlorine is the secondhalogen,being anonmetalin group 17 of the periodic table. Its properties are thus similar tofluorine,bromine,andiodine,and are largely intermediate between those of the first two. Chlorine has the electron configuration [Ne]3s²3p⁵,with the seven electrons in the third and outermost shell acting as itsvalence electrons.Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell.Corresponding toperiodic trends,it is intermediate inelectronegativitybetween fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine. It is also a weaker oxidising agent than fluorine, but a stronger one than bromine. Conversely, thechlorideion is a weaker reducing agent than bromide, but a stronger one than fluoride.It is intermediate inatomic radiusbetween fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as firstionisation energy,electron affinity,enthalpy of dissociation of the X₂molecule (X = Cl, Br, I), ionic radius, and X–X bond length. (Fluorine is anomalous due to its small size.)

All four stable halogens experience intermolecularvan der Waals forcesof attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: chlorine melts at −101.0 °C and boils at −34.0 °C. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure.The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow-green. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group.Specifically, the colour of a halogen, such as chlorine, results from theelectron transitionbetween thehighest occupiedantibondingπ_gmolecular orbital and the lowest vacant antibondingσ_umolecular orbital.The colour fades at low temperatures, so that solid chlorine at −195 °C is almost colourless.

Like solid bromine and iodine, solid chlorine crystallises in theorthorhombic crystal system,in a layered lattice of Cl₂molecules. The Cl–Cl distance is 198 pm (close to the gaseous Cl–Cl distance of 199 pm) and the Cl···Cl distance between molecules is 332 pm within a layer and 382 pm between layers (compare the van der Waals radius of chlorine, 180 pm). This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable.

Isotopes

Main article:Isotopes of chlorine

Chlorine has two stable isotopes,³⁵Cl and³⁷Cl. These are its only two natural isotopes occurring in quantity, with³⁵Cl making up 76% of natural chlorine and³⁷Cl making up the remaining 24%. Both are synthesised in stars in theoxygen-burningandsilicon-burning processes.Both have nuclear spin 3/2+ and thus may be used fornuclear magnetic resonance,although the spin magnitude being greater than 1/2 results in non-spherical nuclear charge distribution and thus resonance broadening as a result of a nonzeronuclear quadrupole momentand resultant quadrupolar relaxation. The other chlorine isotopes are all radioactive, withhalf-livestoo short to occur in natureprimordially.Of these, the most commonly used in the laboratory are³⁶Cl (t_1/2= 3.0×10⁵y) and³⁸Cl (t_1/2= 37.2 min), which may be produced from theneutron activationof natural chlorine.

The most stable chlorine radioisotope is³⁶Cl. The primary decay mode of isotopes lighter than³⁵Cl iselectron captureto isotopes ofsulfur;that of isotopes heavier than³⁷Cl isbeta decayto isotopes ofargon;and³⁶Cl may decay by either mode to stable³⁶S or³⁶Ar.³⁶Cloccurs in trace quantities in nature as acosmogenic nuclidein a ratio of about (7–10) × 10⁻¹³to 1 with stable chlorine isotopes: it is produced in the atmosphere byspallationof³⁶Arby interactions withcosmic rayprotons.In the top meter of thelithosphere,³⁶Cl is generated primarily bythermal neutronactivation of³⁵Cl and spallation of³⁹Kand⁴⁰Ca.In the subsurface environment,muon captureby⁴⁰Cabecomes more important as a way to generate³⁶Cl.

Chemistry and compounds

Halogen bond energies (kJ/mol)

X	XX	HX	BX3	AlX3	CX4
F	159	574	645	582	456
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

Chlorine is intermediate in reactivity between fluorine and bromine, and is one of the most reactive elements. Chlorine is a weaker oxidising agent than fluorine but a stronger one than bromine or iodine. This can be seen from thestandard electrode potentialsof the X₂/X⁻couples (F, +2.866  V; Cl, +1.395 V; Br, +1.087  V; I, +0.615 V;At,approximately +0.3  V). However, this trend is not shown in the bond energies because fluorine is singular due to its small size, low polarisability, and inability to showhypervalence.As another difference, chlorine has a significant chemistry in positive oxidation states while fluorine does not. Chlorination often leads to higher oxidation states than bromination or iodination but lower oxidation states than fluorination. Chlorine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Cl bonds.

Given that E°(⁠1/2⁠O₂/H₂O) = +1.229 V, which is less than +1.395 V, it would be expected that chlorine should be able to oxidise water to oxygen and hydrochloric acid. However, the kinetics of this reaction are unfavorable, and there is also a bubbleoverpotentialeffect to consider, so that electrolysis of aqueous chloride solutions evolves chlorine gas and not oxygen gas, a fact that is very useful for the industrial production of chlorine.

Hydrogen chloride

Structure of solid deuterium chloride, with D···Cl hydrogen bonds

The simplest chlorine compound ishydrogen chloride,HCl, a major chemical in industry as well as in the laboratory, both as a gas and dissolved in water ashydrochloric acid.It is often produced by burning hydrogen gas in chlorine gas, or as a byproduct of chlorinatinghydrocarbons.Another approach is to treatsodium chloridewith concentratedsulfuric acidto produce hydrochloric acid, also known as the "salt-cake" process:

NaCl + H₂SO₄150 °C⟶ NaHSO₄+ HCl

NaCl + NaHSO₄540–600 °C⟶ Na₂SO₄+ HCl

In the laboratory, hydrogen chloride gas may be made by drying the acid with concentrated sulfuric acid. Deuterium chloride, DCl, may be produced by reactingbenzoyl chloridewithheavy water(D₂O).

At room temperature, hydrogen chloride is a colourless gas, like all the hydrogen halides apart fromhydrogen fluoride,since hydrogen cannot form stronghydrogen bondsto the larger electronegative chlorine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen chloride at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised.Hydrochloric acid is a strong acid (pK_a= −7) because the hydrogen bonds to chlorine are too weak to inhibit dissociation. The HCl/H₂O system has many hydrates HCl·nH₂O forn= 1, 2, 3, 4, and 6. Beyond a 1:1 mixture of HCl and H₂O, the system separates completely into two separate liquid phases. Hydrochloric acid forms anazeotropewith boiling point 108.58 °C at 20.22 g HCl per 100 g solution; thus hydrochloric acid cannot be concentrated beyond this point by distillation.

Unlike hydrogen fluoride, anhydrous liquid hydrogen chloride is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, itsdielectric constantis low and it does not dissociate appreciably into H₂Cl⁺andHCl⁻
₂ions – the latter, in any case, are much less stable than thebifluorideions (HF⁻
₂) due to the very weak hydrogen bonding between hydrogen and chlorine, though its salts with very large and weakly polarising cations such asCs⁺andNR⁺
₄(R =Me,Et,Buⁿ) may still be isolated. Anhydrous hydrogen chloride is a poor solvent, only able to dissolve small molecular compounds such asnitrosyl chlorideandphenol,or salts with very lowlattice energiessuch as tetraalkylammonium halides. It readily protonateselectrophilescontaining lone-pairs or π bonds.Solvolysis,ligandreplacement reactions, and oxidations are well-characterised in hydrogen chloride solution:

Ph₃SnCl + HCl ⟶ Ph₂SnCl₂+ PhH (solvolysis)

Ph₃COH + 3 HCl ⟶Ph
₃C⁺
HCl⁻
₂+ H₃O⁺Cl⁻(solvolysis)

Me
₄N⁺
HCl⁻
₂+ BCl₃⟶Me
₄N⁺
BCl⁻
₄+ HCl (ligand replacement)

PCl₃+ Cl₂+ HCl ⟶PCl⁺
₄HCl⁻
₂(oxidation)

Other binary chlorides

Hydratednickel(II) chloride,NiCl₂(H₂O)₆

Nearly all elements in the periodic table form binary chlorides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (thenoble gases,with the exception ofxenonin the highly unstableXeCl₂and XeCl₄); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyondbismuth); and having an electronegativity higher than chlorine's (oxygenandfluorine) so that the resultant binary compounds are formally not chlorides but rather oxides or fluorides of chlorine.Even thoughnitrogenin NCl₃is bearing a negative charge, the compound is usually callednitrogen trichloride.

Chlorination of metals with Cl₂usually leads to a higher oxidation state than bromination with Br₂when multiple oxidation states are available, such as inMoCl₅andMoBr₃.Chlorides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrochloric acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen chloride gas. These methods work best when the chloride product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative chlorination of the element with chlorine or hydrogen chloride, high-temperature chlorination of a metal oxide or other halide by chlorine, a volatile metal chloride,carbon tetrachloride,or an organic chloride. For instance,zirconium dioxidereacts with chlorine at standard conditions to producezirconium tetrachloride,anduranium trioxidereacts withhexachloropropenewhen heated underrefluxto giveuranium tetrachloride.The second example also involves a reduction inoxidation state,which can also be achieved by reducing a higher chloride using hydrogen or a metal as a reducing agent. This may also be achieved by thermal decomposition or disproportionation as follows:

EuCl₃+⁠1/2⁠H₂⟶ EuCl₂+ HCl

ReCl₅at "bp"⟶ ReCl₃+ Cl₂

AuCl₃160 °C⟶ AuCl + Cl₂

Most metal chlorides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular chlorides, as do metals in high oxidation states from +3 and above. Both ionic and covalent chlorides are known for metals in oxidation state +3 (e.g.scandium chlorideis mostly ionic, butaluminium chlorideis not).Silver chlorideis very insoluble in water and is thus often used as a qualitative test for chlorine.

Polychlorine compounds

Although dichlorine is a strong oxidising agent with a high first ionisation energy, it may be oxidised under extreme conditions to form the[Cl₂]⁺cation. This is very unstable and has only been characterised by its electronic band spectrum when produced in a low-pressure discharge tube. The yellow[Cl₃]⁺cation is more stable and may be produced as follows:

Cl₂+ ClF + AsF₅−78 °C⟶[Cl₃]⁺[AsF₆]⁻

This reaction is conducted in the oxidising solventarsenic pentafluoride.The trichloride anion,[Cl₃]⁻,has also been characterised; it is analogous totriiodide.

Chlorine fluorides

The three fluorides of chlorine form a subset of theinterhalogencompounds, all of which arediamagnetic.Some cationic and anionic derivatives are known, such asClF⁻
₂,ClF⁻
₄,ClF⁺
₂,and Cl₂F⁺.Somepseudohalidesof chlorine are also known, such ascyanogen chloride(ClCN, linear), chlorinecyanate(ClNCO), chlorinethiocyanate(ClSCN, unlike its oxygen counterpart), and chlorineazide(ClN₃).

Chlorine monofluoride(ClF) is extremely thermally stable, and is sold commercially in 500-gram steel lecture bottles. It is a colourless gas that melts at −155.6 °C and boils at −100.1 °C. It may be produced by the reaction of its elements at 225 °C, though it must then be separated and purified fromchlorine trifluorideand its reactants. Its properties are mostly intermediate between those of chlorine and fluorine. It will react with many metals and nonmetals from room temperature and above, fluorinating them and liberating chlorine. It will also act as a chlorofluorinating agent, adding chlorine and fluorine across a multiple bond or by oxidation: for example, it will attackcarbon monoxideto form carbonyl chlorofluoride, COFCl. It will react analogously withhexafluoroacetone,(CF₃)₂CO, with apotassium fluoridecatalyst to produce heptafluoroisopropyl hypochlorite, (CF₃)₂CFOCl; withnitrilesRCN to produce RCF₂NCl₂;and with the sulfur oxides SO₂and SO₃to produce ClSO₂F and ClOSO₂F respectively. It will also react exothermically with compounds containing –OH and –NH groups, such as water:

H₂O + 2 ClF ⟶ 2 HF + Cl₂O

Chlorine trifluoride(ClF₃) is a volatile colourless molecular liquid which melts at −76.3 °C and boils at 11.8  °C. It may be formed by directly fluorinating gaseous chlorine or chlorine monofluoride at 200–300 °C. One of the most reactive chemical compounds known, the list of elements it sets on fire is diverse, containinghydrogen,potassium,phosphorus,arsenic,antimony,sulfur,selenium,tellurium,bromine,iodine,and powderedmolybdenum,tungsten,rhodium,iridium,andiron.It will also ignite water, along with many substances which in ordinary circumstances would be considered chemically inert such asasbestos,concrete, glass, and sand. When heated, it will even corrodenoble metalsaspalladium,platinum,andgold,and even thenoble gasesxenonandradondo not escape fluorination. An impermeable fluoride layer is formed bysodium,magnesium,aluminium,zinc,tin,andsilver,which may be removed by heating.Nickel,copper, and steel containers are usually used due to their great resistance to attack by chlorine trifluoride, stemming from the formation of an unreactive layer of metal fluoride. Its reaction withhydrazineto form hydrogen fluoride, nitrogen, and chlorine gases was used in experimental rocket engine, but has problems largely stemming from its extremehypergolicityresulting in ignition without any measurable delay. Today, it is mostly used in nuclear fuel processing, to oxidiseuraniumtouranium hexafluoridefor its enriching and to separate it fromplutonium,as well as in the semiconductor industry, where it is used to cleanchemical vapor depositionchambers.It can act as a fluoride ion donor or acceptor (Lewis base or acid), although it does not dissociate appreciably intoClF⁺
₂andClF⁻
₄ions.

Chlorine pentafluoride(ClF₅) is made on a large scale by direct fluorination of chlorine with excessfluorinegas at 350 °C and 250 atm, and on a small scale by reacting metal chlorides with fluorine gas at 100–300 °C. It melts at −103 °C and boils at −13.1 °C. It is a very strong fluorinating agent, although it is still not as effective as chlorine trifluoride. Only a few specific stoichiometric reactions have been characterised.Arsenic pentafluorideandantimony pentafluorideform ionic adducts of the form [ClF₄]⁺[MF₆]⁻(M = As, Sb) and water reacts vigorously as follows:

2 H₂O + ClF₅⟶ 4 HF + FClO₂

The product,chloryl fluoride,is one of the five known chlorine oxide fluorides. These range from the thermally unstable FClO to the chemically unreactiveperchloryl fluoride(FClO₃), the other three being FClO₂,F₃ClO, and F₃ClO₂.All five behave similarly to the chlorine fluorides, both structurally and chemically, and may act as Lewis acids or bases by gaining or losing fluoride ions respectively or as very strong oxidising and fluorinating agents.

Chlorine oxides

Yellowchlorine dioxide(ClO₂) gas above a solution of hydrochloric acid and sodium chlorite in water, also containing dissolved chlorine dioxide

Structure ofdichlorine heptoxide,Cl₂O₇,the most stable of the chlorine oxides

Thechlorine oxidesare well-studied in spite of their instability (all of them are endothermic compounds). They are important because they are produced whenchlorofluorocarbonsundergo photolysis in the upper atmosphere and cause the destruction of the ozone layer. None of them can be made from directly reacting the elements.

Dichlorine monoxide(Cl₂O) is a brownish-yellow gas (red-brown when solid or liquid) which may be obtained by reacting chlorine gas with yellowmercury(II) oxide.It is very soluble in water, in which it is in equilibrium withhypochlorous acid(HOCl), of which it is the anhydride. It is thus an effective bleach and is mostly used to makehypochlorites.It explodes on heating or sparking or in the presence of ammonia gas.

Chlorine dioxide(ClO₂) was the first chlorine oxide to be discovered in 1811 byHumphry Davy.It is a yellow paramagnetic gas (deep-red as a solid or liquid), as expected from its having an odd number of electrons: it is stable towards dimerisation due to the delocalisation of the unpaired electron. It explodes above −40 °C as a liquid and under pressure as a gas and therefore must be made at low concentrations for wood-pulp bleaching and water treatment. It is usually prepared by reducing achlorateas follows:

ClO⁻
₃+ Cl⁻+ 2 H⁺⟶ ClO₂+⁠1/2⁠Cl₂+ H₂O

Its production is thus intimately linked to the redox reactions of the chlorine oxoacids. It is a strong oxidising agent, reacting withsulfur,phosphorus,phosphorus halides, andpotassium borohydride.It dissolves exothermically in water to form dark-green solutions that very slowly decompose in the dark. Crystalline clathrate hydrates ClO₂·nH₂O (n≈ 6–10) separate out at low temperatures. However, in the presence of light, these solutions rapidly photodecompose to form a mixture of chloric and hydrochloric acids. Photolysis of individual ClO₂molecules result in the radicals ClO and ClOO, while at room temperature mostly chlorine, oxygen, and some ClO₃and Cl₂O₆are produced. Cl₂O₃is also produced when photolysing the solid at −78 °C: it is a dark brown solid that explodes below 0 °C. The ClO radical leads to the depletion of atmospheric ozone and is thus environmentally important as follows:

Cl• + O₃⟶ ClO• + O₂

ClO• + O• ⟶ Cl• + O₂

Chlorine perchlorate(ClOClO₃) is a pale yellow liquid that is less stable than ClO₂and decomposes at room temperature to form chlorine, oxygen, anddichlorine hexoxide(Cl₂O₆).Chlorine perchlorate may also be considered a chlorine derivative ofperchloric acid(HOClO₃), similar to the thermally unstable chlorine derivatives of other oxoacids: examples includechlorine nitrate(ClONO₂,vigorously reactive and explosive), and chlorine fluorosulfate (ClOSO₂F, more stable but still moisture-sensitive and highly reactive).Dichlorine hexoxide is a dark-red liquid that freezes to form a solid which turns yellow at −180 °C: it is usually made by reaction of chlorine dioxide with oxygen. Despite attempts to rationalise it as the dimer of ClO₃,it reacts more as though it were chloryl perchlorate, [ClO₂]⁺[ClO₄]⁻,which has been confirmed to be the correct structure of the solid. It hydrolyses in water to give a mixture of chloric and perchloric acids: the analogous reaction with anhydroushydrogen fluoridedoes not proceed to completion.

Dichlorine heptoxide(Cl₂O₇) is the anhydride ofperchloric acid(HClO₄) and can readily be obtained from it by dehydrating it withphosphoric acidat −10 °C and then distilling the product at −35 °C and 1 mmHg. It is a shock-sensitive, colourless oily liquid. It is the least reactive of the chlorine oxides, being the only one to not set organic materials on fire at room temperature. It may be dissolved in water to regenerate perchloric acid or in aqueous alkalis to regenerate perchlorates. However, it thermally decomposes explosively by breaking one of the central Cl–O bonds, producing the radicals ClO₃and ClO₄which immediately decompose to the elements through intermediate oxides.

Chlorine oxoacids and oxyanions

Standard reduction potentials for aqueous Cl species

E°(couple)	a(H+) = 1 (acid)	E°(couple)	a(OH−) = 1 (base)
Cl2/Cl−	+1.358	Cl2/Cl−	+1.358
HOCl/Cl−	+1.484	ClO−/Cl−	+0.890
ClO− 3/Cl−	+1.459
HOCl/Cl2	+1.630	ClO−/Cl2	+0.421
HClO2/Cl2	+1.659
ClO− 3/Cl2	+1.468
ClO− 4/Cl2	+1.277
HClO2/HOCl	+1.701	ClO− 2/ClO−	+0.681
		ClO− 3/ClO−	+0.488
ClO− 3/HClO2	+1.181	ClO− 3/ClO− 2	+0.295
ClO− 4/ClO− 3	+1.201	ClO− 4/ClO− 3	+0.374

Chlorine forms four oxoacids:hypochlorous acid(HOCl),chlorous acid(HOClO),chloric acid(HOClO₂), andperchloric acid(HOClO₃). As can be seen from the redox potentials given in the adjacent table, chlorine is much more stable towards disproportionation in acidic solutions than in alkaline solutions:

Cl2+ H2O	⇌ HOCl + H++ Cl−	Kac= 4.2 × 10−4mol2l−2
Cl2+ 2 OH−	⇌ OCl−+ H2O + Cl−	Kalk= 7.5 × 1015mol−1l

The hypochlorite ions also disproportionate further to produce chloride and chlorate (3 ClO⁻⇌ 2 Cl⁻+ClO⁻
₃) but this reaction is quite slow at temperatures below 70 °C in spite of the very favourable equilibrium constant of 10²⁷.The chlorate ions may themselves disproportionate to form chloride and perchlorate (4ClO⁻
₃⇌ Cl⁻+ 3ClO⁻
₄) but this is still very slow even at 100 °C despite the very favourable equilibrium constant of 10²⁰.The rates of reaction for the chlorine oxyanions increases as the oxidation state of chlorine decreases. The strengths of the chlorine oxyacids increase very quickly as the oxidation state of chlorine increases due to the increasing delocalisation of charge over more and more oxygen atoms in their conjugate bases.

Most of the chlorine oxoacids may be produced by exploiting these disproportionation reactions. Hypochlorous acid (HOCl) is highly reactive and quite unstable; its salts are mostly used for their bleaching and sterilising abilities. They are very strong oxidising agents, transferring an oxygen atom to most inorganic species. Chlorous acid (HOClO) is even more unstable and cannot be isolated or concentrated without decomposition: it is known from the decomposition of aqueous chlorine dioxide. However,sodium chloriteis a stable salt and is useful for bleaching and stripping textiles, as an oxidising agent, and as a source of chlorine dioxide. Chloric acid (HOClO₂) is a strong acid that is quite stable in cold water up to 30% concentration, but on warming gives chlorine and chlorine dioxide. Evaporation under reduced pressure allows it to be concentrated further to about 40%, but then it decomposes to perchloric acid, chlorine, oxygen, water, and chlorine dioxide. Its most important salt issodium chlorate,mostly used to make chlorine dioxide to bleach paper pulp. The decomposition of chlorate to chloride and oxygen is a common way to produce oxygen in the laboratory on a small scale. Chloride and chlorate may comproportionate to form chlorine as follows:

ClO⁻
₃+ 5 Cl⁻+ 6 H⁺⟶ 3 Cl₂+ 3 H₂O

Perchlorates and perchloric acid (HOClO₃) are the most stable oxo-compounds of chlorine, in keeping with the fact that chlorine compounds are most stable when the chlorine atom is in its lowest (−1) or highest (+7) possible oxidation states. Perchloric acid and aqueous perchlorates are vigorous and sometimes violent oxidising agents when heated, in stark contrast to their mostly inactive nature at room temperature due to the high activation energies for these reactions for kinetic reasons. Perchlorates are made by electrolytically oxidising sodium chlorate, and perchloric acid is made by reacting anhydroussodium perchlorateorbarium perchloratewith concentrated hydrochloric acid, filtering away the chloride precipitated and distilling the filtrate to concentrate it. Anhydrous perchloric acid is a colourless mobile liquid that is sensitive to shock that explodes on contact with most organic compounds, setshydrogen iodideandthionyl chlorideon fire and even oxidises silver and gold. Although it is a weak ligand, weaker than water, a few compounds involving coordinatedClO⁻
₄are known.The Table below presents typical oxidation states for chlorine element as given in the secondary schools or colleges. There are more complex chemical compounds, the structure of which can only be explained using modern quantum chemical methods, for example, cluster technetium chloride [(CH₃)₄N]₃[Tc₆Cl₁₄], in which 6 of the 14 chlorine atoms are formally divalent, and oxidation states are fractional.In addition, all the above chemical regularities are valid for "normal" or close to normal conditions, while at ultra-high pressures (for example, in the cores of large planets), chlorine can exhibit an oxidation state of -3, forming a Na3Cl compound with sodium, which does not fit into traditional concepts of chemistry.

Chlorine oxidation state	−1	+1	+3	+5	+7
Name	chloride	hypochlorite	chlorite	chlorate	perchlorate
Formula	Cl−	ClO−	ClO− 2	ClO− 3	ClO− 4
Structure

Organochlorine compounds

Main article:Organochlorine compound

Suggested mechanism for the chlorination of a carboxylic acid by phosphorus pentachloride to form anacyl chloride

Like the other carbon–halogen bonds, the C–Cl bond is a common functional group that forms part of coreorganic chemistry.Formally, compounds with this functional group may be considered organic derivatives of the chloride anion. Due to the difference of electronegativity between chlorine (3.16) and carbon (2.55), the carbon in a C–Cl bond is electron-deficient and thuselectrophilic.Chlorinationmodifies the physical properties of hydrocarbons in several ways: chlorocarbons are typically denser thanwaterdue to the higher atomic weight of chlorine versus hydrogen, and aliphatic organochlorides arealkylating agentsbecause chloride is aleaving group.

Alkanesandarylalkanes may be chlorinated underfree-radicalconditions, with UV light. However, the extent of chlorination is difficult to control: the reaction is notregioselectiveand often results in a mixture of various isomers with different degrees of chlorination, though this may be permissible if the products are easily separated. Aryl chlorides may be prepared by theFriedel-Crafts halogenation,using chlorine and aLewis acidcatalyst.Thehaloform reaction,using chlorine andsodium hydroxide,is also able to generate alkyl halides from methyl ketones, and related compounds. Chlorine adds to the multiple bonds onalkenesandalkynesas well, giving di- or tetrachloro compounds. However, due to the expense and reactivity of chlorine, organochlorine compounds are more commonly produced by using hydrogen chloride, or with chlorinating agents such asphosphorus pentachloride(PCl₅) orthionyl chloride(SOCl₂). The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out.

Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.Chlorinated organic compounds are found in nearly every class of biomolecules includingalkaloids,terpenes,amino acids,flavonoids,steroids,andfatty acids.Organochlorides, includingdioxins,are produced in the high temperature environment of forest fires, and dioxins have been found in the preserved ashes of lightning-ignited fires that predate synthetic dioxins.In addition, a variety of simple chlorinated hydrocarbons including dichloromethane, chloroform, andcarbon tetrachloridehave been isolated from marine algae.A majority of thechloromethanein the environment is produced naturally by biological decomposition, forest fires, and volcanoes.

Some types of organochlorides, though not all, have significant toxicity to plants or animals, including humans. Dioxins, produced when organic matter is burned in the presence of chlorine, and some insecticides, such asDDT,arepersistent organic pollutantswhich pose dangers when they are released into the environment. For example, DDT, which was widely used to control insects in the mid 20th century, also accumulates in food chains, and causes reproductive problems (e.g., eggshell thinning) in certain bird species.Due to the ready homolytic fission of the C–Cl bond to create chlorine radicals in the upper atmosphere,chlorofluorocarbonshave been phased out due to the harm they do to the ozone layer.

Occurrence

Liquid chlorine analysis

Chlorine is too reactive to occur as the free element in nature but is very abundant in the form of its chloride salts. It is the 20th most abundant elementin Earth's crust and makes up 126parts per millionof it, through the large deposits of chloride minerals, especiallysodium chloride,that have been evaporated from water bodies. All of these pale in comparison to the reserves of chloride ions in seawater: smaller amounts at higher concentrations occur in some inland seas and undergroundbrinewells, such as theGreat Salt Lakein Utah and theDead Seain Israel.

Small batches of chlorine gas are prepared in the laboratory by combining hydrochloric acid andmanganese dioxide,but the need rarely arises due to its ready availability. In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. This method, thechloralkali processindustrialized in 1892, now provides most industrial chlorine gas.Along with chlorine, the method yieldshydrogengas andsodium hydroxide,which is the most valuable product. The process proceeds according to the followingchemical equation:

2 NaCl + 2 H₂O → Cl₂+ H₂+ 2 NaOH

Production

Main article:Chlorine production

Chlorine is primarily produced by thechloralkali process,although non-chloralkali processes exist. Global 2006 production was estimated to be 65 million short tons.The most visible use of chlorine is inwater disinfection.35-40 % of chlorine produced is used to makepoly(vinyl chloride)throughethylene dichlorideandvinyl chloride.The chlorine produced is available in cylinders from sizes ranging from 450 g to 70 kg, as well as drums (865 kg), tank wagons (15 tonnes on roads; 27–90 tonnes by rail), and barges (600–1200 tonnes). Due to the difficulty in transporting elemental chlorine, production is typically located near where it is consumed. As examples, vinyl chloride producers such asWestlake ChemicalandFormosa Plasticshave integrated chloralkali assets.

Chloralkali processes

Main article:Chloralkali process

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

Cathode: 2 H₂O + 2 e⁻→ H₂+ 2 OH⁻

Anode: 2 Cl⁻→ Cl₂+ 2 e⁻

In the conventional case where sodium chloride is electrolyzed,sodium hydroxideand chlorine are coproducts.

Industrially, there are three chloralkali processes:

• TheCastner–Kellner processthat utilizes a mercury electrode
• The diaphragm cell process that utilizes an asbestos diaphragm that separates the cathode and anode
• The membrane cell process that uses anion exchangemembrane in place of the diaphragm

The Castner–Kellner process was the first method used at the end of the nineteenth century to produce chlorine on an industrial scale.Mercury (that is toxic) was used as an electrode toamalgamatethe sodium product, preventing undesirable side reactions.

In diaphragm cell electrolysis, anasbestos(or polymer-fiber) diaphragm separates a cathode and ananode,preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where thecausticalkaliis produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali, but they are not burdened with the problem ofmercurydisposal and they are more energy efficient.

Membrane cell electrolysis employspermeable membraneas anion exchanger.Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lowerconcentration.This method also produces very pure sodium (or potassium) hydroxide but has the disadvantage of requiring very pure brine at high concentrations.

Membrane cell process for chloralkali production

Non-chloralkali processes

In theDeacon process,hydrogen chloride recovered from the production oforganochlorine compoundsis recovered as chlorine. The process relies on oxidation using oxygen:

4 HCl + O₂→ 2 Cl₂+ 2 H₂O

The reaction requires a catalyst. As introduced by Deacon, early catalysts were based on copper. Commercial processes, such as the Mitsui MT-Chlorine Process, have switched to chromium and ruthenium-based catalysts.

Applications

A railwaytank carcarrying chlorine, displayinghazardous materialsinformation including a diamond-shapedU.S. DOT placardshowing aUN number

Sodium chloride is the most common chlorine compound, and is the main source of chlorine for the demand by the chemical industry. About 15000 chlorine-containing compounds are commercially traded, including such diverse compounds as chlorinatedmethane,ethanes,vinyl chloride,polyvinyl chloride(PVC),aluminium trichlorideforcatalysis,the chlorides ofmagnesium,titanium,zirconium,andhafniumwhich are the precursors for producing the pure form of those elements.

Quantitatively, of all elemental chlorine produced, about 63% is used in the manufacture of organic compounds, and 18% in the manufacture of inorganic chlorine compounds.About 15,000 chlorine compounds are used commercially.The remaining 19% of chlorine produced is used for bleaches and disinfection products.The most significant of organic compounds in terms of production volume are1,2-dichloroethaneandvinyl chloride,intermediates in the production ofPVC.Other particularly important organochlorines aremethyl chloride,methylene chloride,chloroform,vinylidene chloride,trichloroethylene,perchloroethylene,allyl chloride,epichlorohydrin,chlorobenzene,dichlorobenzenes,andtrichlorobenzenes.The major inorganic compounds include HCl, Cl₂O, HOCl, NaClO₃,AlCl₃,SiCl₄,SnCl₄,PCl₃,PCl₅,POCl₃,AsCl₃,SbCl₃,SbCl₅,BiCl₃,andZnCl₂.

Sanitation, disinfection, and antisepsis

Main articles:Water chlorinationandBleach

Combating putrefaction

In France (as elsewhere),animal intestineswere processed to make musical instrument strings,Goldbeater's skinand other products. This was done in "gut factories" (boyauderies), and it was an odiferous and unhealthy process. In or about 1820, theSociété d'encouragement pour l'industrie nationaleoffered a prize for the discovery of a method, chemical or mechanical, for separating theperitonealmembrane of animal intestines withoutputrefaction.The prize was won byAntoine-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 actually retarded the decomposition.

Labarraque's research resulted in the use of chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) in theboyauderies.The same chemicals were found to be useful in the routinedisinfectionand deodorization oflatrines,sewers,markets,abattoirs,anatomical theatres,and morgues.They were successful inhospitals,lazarets,prisons,infirmaries(both on land and at sea),magnaneries,stables,cattle-sheds, etc.; and they were beneficial duringexhumations,embalming,outbreaks of epidemic disease, fever, andblacklegin cattle.

Disinfection

Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", presumed to be transmitted by "miasmas"), and to treatputrefactionof existing wounds, including septic wounds.In his 1828 work, Labarraque recommended that doctors breathe chlorine, wash their hands in chlorinated lime, and even sprinkle chlorinated lime about the patients' beds in cases of "contagious infection". In 1828, the contagion of infections was well known, even though the agency of themicrobewas not discovered until more than half a century later.

During theParis cholera outbreakof 1832, large quantities of so-calledchloride of limewere used to disinfect the capital. This was not simply moderncalcium chloride,but chlorine gas dissolved in lime-water (dilutecalcium hydroxide) to formcalcium 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 theLatin Quarterof Paris.These "putrid miasmas" were thought by many to cause the spread of "contagion" and "infection" – both words used before the germ theory of infection. Chloride of lime was used for destroying odors and "putrid matter". One source claims chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well that was feeding the Broad Street pump in 1854 London,though three other reputable sources that describe that famous cholera epidemic do not mention the incident.One reference makes it clear that chloride of lime was used to disinfect theoffaland filth in the streets surrounding the Broad Street pump – a common practice in mid-nineteenth century England.

Semmelweis and experiments with antisepsis

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine andchemical basesolutions was in 1847, whenIgnaz Semmelweisused chlorine-water (chlorine dissolved in pure water, which was cheaper than chlorinated lime solutions) to disinfect the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Long before the germ theory of disease, Semmelweis theorized that "cadaveric particles" were 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 antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis's celebrated success in stopping the transmission ofchildbed fever( "puerperal fever" ) in the maternity wards ofVienna General HospitalinAustriain 1847.

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 byHenry Drysdale Dakin(who gave full credit to Labarraque's prior work in this area). CalledDakin'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 modern times, where it remains effective against bacteria that are resistant to multiple antibiotics (seeCentury Pharmaceuticals).

Public sanitation

Chlorinated water is used inswimming poolsto disinfect water from microbial contaminants

Liquid pool chlorine

The first continuous application of chlorination to drinking U.S. water was installed inJersey City,New Jersey, in 1908.By 1918, theUS Department of Treasurycalled for all drinking water to be disinfected with chlorine. Chlorine is presently an important chemical forwater purification(such as in water treatment plants), indisinfectants,and inbleach.Even small water supplies are now routinely chlorinated.

Chlorine is usually used (in the form ofhypochlorous acid) to killbacteriaand other microbes indrinking watersupplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rathersodium hypochlorite,formed from chlorine andsodium hydroxide,or solid tablets of chlorinated isocyanurates. The drawback of using chlorine in swimming pools is that the chlorine reacts with theamino acidsin proteins in human hair and skin. Contrary to popular belief, the distinctive "chlorine aroma" associated with swimming pools is not the result of elemental chlorine itself, but ofchloramine,a chemical compound produced by the reaction of free dissolved chlorine with amines in organic substances including those in urine and sweat.As a disinfectant in water, chlorine is more than three times as effective againstEscherichia coliasbromine,and more than six times as effective asiodine.Increasingly,monochloramineitself is being directly added to drinking water for purposes of disinfection, a process known aschloramination.

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These includehypochloritesolutions, which gradually release chlorine into the water, and compounds likesodium dichloro-s-triazinetrione(dihydrate or anhydrous), sometimes referred to as "dichlor", andtrichloro-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 generalbiocide,killing germs, microorganisms, algae, and so on.

Use as a weapon

World War I

Main article:Chemical weapons in World War I

Chlorine gas, also known as bertholite, was firstused as a weaponinWorld War Iby Germany on April 22, 1915, in theSecond Battle of Ypres.As described by the soldiers, it had the distinctive smell of a mixture of pepper and pineapple.It also tasted metallic and stung the back of the throat and chest. Chlorine reacts with water in themucosaof the lungs to formhydrochloric acid,destructive to living tissue and potentially lethal. Human respiratory systems can be protected from chlorine gas bygas maskswithactivated charcoalor other filters, which makes chlorine gas much less lethal than other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate,Fritz Haberof theKaiser Wilhelm Institutein Berlin, in collaboration with the German chemical conglomerateIG Farben,which developed methods for discharging chlorine gas against anentrenchedenemy.After its first use, both sides in the conflict used chlorine as a chemical weapon, but it was soon replaced by the more deadlyphosgeneandmustard gas.

Middle east

Main articles:Chlorine bombings in IraqandUse of chemical weapons in the Syrian Civil War

Chlorine gas was also used during theIraq War in Anbar Provincein 2007, with insurgents packingtruck bombswithmortarshells and chlorine tanks. The attacks killed two people from the explosives and sickened more than 350. 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. In some bombings, over a hundred civilians were hospitalized due to breathing difficulties. The Iraqi authorities tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.

On 23 October 2014, it was reported that theIslamic State of Iraq and the Levanthad used chlorine gas in the town of Duluiyah,Iraq.Laboratory analysis of clothing and soil samples confirmed the use of chlorine gas against KurdishPeshmergaForces in a vehicle-borne improvised explosive device attack on 23 January 2015 at the Highway 47 Kiske Junction near Mosul.

Another country in the middle east,Syria,has used chlorine as achemical weapondelivered frombarrel bombsand rockets.In 2016, theOPCW-UN Joint Investigative Mechanismconcluded that the Syrian government used chlorine as a chemical weapon in three separate attacks.Later investigations from the OPCW's Investigation and Identification Team concluded that theSyrian Air Forcewas responsible for chlorine attacks in 2017 and 2018.

Biological role

Thechlorideanion is anessential nutrientfor metabolism. Chlorine is needed for the production ofhydrochloric acidin the stomach and in cellular pump functions.The main dietary source is table salt, or sodium chloride. Overly low or high concentrations of chloride in the blood are examples ofelectrolyte disturbances.Hypochloremia(having too little chloride) rarely occurs in the absence of other abnormalities. It is sometimes associated withhypoventilation.It can be associated with chronicrespiratory acidosis.Hyperchloremia(having too much chloride) usually does not produce symptoms. When symptoms do occur, they tend to resemble those ofhypernatremia(having too muchsodium). Reduction in blood chloride leads to cerebral dehydration; symptoms are most often caused by rapid rehydration which results incerebral edema.Hyperchloremia can affect oxygen transport.

Hazards

Chlorine

Hazards
GHSlabelling:
Pictograms
Signal word	Danger
Hazard statements	H270,H315,H319,H330,H335,H400
Precautionary statements	P220,P233,P244,P261,P304,P312,P340,P403,P410
NFPA 704(fire diamond)	3 0 0 OX

Chlorine is a toxic gas that attacks the respiratory system, eyes, and skin.Because it is denser 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.

Chlorine is detectable with measuring devices in concentrations 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.TheIDLH(immediately dangerous to life and health) concentration is 10 ppm.Breathing lower concentrations can aggravate the respiratory system and exposure to the gas can irritate the eyes.When chlorine is inhaled at concentrations greater than 30 ppm, it reacts with water within the lungs, producinghydrochloric acid(HCl) andhypochlorous acid(HOCl).

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 generatedisinfection by-productsthat are associated with negative effects on human health.

In the United States, theOccupational Safety and Health Administration(OSHA) has set thepermissible exposure limitfor elemental chlorine at 1 ppm, or 3 mg/m³.TheNational Institute for Occupational Safety and Healthhas designated arecommended exposure limitof 0.5 ppm over 15 minutes.

In the home, accidents occur when hypochlorite bleach solutions come into contact with certain acidic drain-cleaners to produce chlorine gas.Hypochlorite bleach (a popularlaundryadditive) combined withammonia(another popular laundry additive) produceschloramines,another toxic group of chemicals.

Chlorine-induced cracking in structural materials

Chlorine "attack" on an acetal resin plumbing joint resulting from a fractured acetal joint in a water supply system which started at aninjection moldingdefect in the joint and slowly grew until the part failed. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure and are the indirect result of the chlorine attack.

Chlorine is widely used for purifying water, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred from chlorine-inducedstress corrosion crackingofstainless steelsuspension rods.Somepolymersare also sensitive to attack, includingacetal resinandpolybutene.Both materials were used in hot and cold water domestic plumbing, andstress corrosion crackingcaused widespread failures in the US in the 1980s and 1990s.

Chlorine-iron fire

The elementironcan combine with chlorine at high temperatures in a strong exothermic reaction, creating achlorine-iron fire.Chlorine-iron fires are a risk in chemical process plants, where much of the pipework that carries chlorine gas is made of steel.

See also

• Chlorine cycle
• Chlorine gas poisoning
• Industrial gas
• Polymer degradation
• Reductive dechlorination

References

Bibliography

• Greenwood, Norman N.;Earnshaw, Alan (1997).Chemistry of the Elements(2nd ed.).Butterworth-Heinemann.ISBN978-0-08-037941-8.

External links

• ChlorineatThe Periodic Table of Videos(University of Nottingham)
• Agency for Toxic Substances and Disease Registry:Chlorine
• Electrolytic production
• Production and liquefaction of chlorine
• Chlorine Production Using Mercury, Environmental Considerations and Alternatives
• National Pollutant Inventory – Chlorine
• National Institute for Occupational Safety and Health – Chlorine Page
• Chlorine InstituteArchived2012-08-27 at theWayback Machine– Trade association representing the chlorine industry
• Chlorine Online– the web portal of Eurochlor – the business association of the European chlor-alkali industry