Chlorine

Around 900, the authors of the Arabic writings attributed to Jabir ibn Hayyan (Latin: Geber) and the Persian physician and alchemist Abu Bakr al-Razi (c. 865–925, Latin: Rhazes) were experimenting with sal ammoniac (ammonium chloride), which when it was distilled together with vitriol (hydrated sulfates of various metals) produced hydrogen chloride.[7] However, it appears that in these early experiments with chloride salts, 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.[8] One of the first such uses was the synthesis of mercury(II) chloride (corrosive sublimate), whose production from the heating of mercury either with alum and ammonium chloride or with vitriol and sodium chloride was first described in the De aluminibus et salibus ("On Alums and Salts", an eleventh- or twelfth century Arabic text falsely attributed to Abu Bakr al-Razi and translated into Latin in the second half of the twelfth century by Gerard of Cremona, 1144–1187).[9] Another important development was the discovery by pseudo-Geber (in the De inventione veritatis, "On the Discovery of Truth", after c. 1300) that by adding ammonium chloride to nitric acid, a strong solvent capable of dissolving gold (i.e., aqua regia) could be produced.[10] Although aqua regia is 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 the Brabantian chemist and physician Jan Baptist van Helmont.[11][en 1]

Carl Wilhelm Scheele, discoverer of chlorine

The element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, and he is credited with the discovery.[12][13] Scheele produced chlorine by reacting MnO₂ (as the mineral pyrolusite) with HCl:[11]

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.[14] 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").[13] He failed to establish chlorine as an element.[13]

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 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.[15][16]

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).[13] 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.[17]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that the substance was an element, and not a compound.[13] He announced his results to the Royal Society on 15 November that year.[11] At that time, he named this new element "chlorine", from the Greek word χλωρος (chlōros, "green-yellow"), in reference to its color.[18] The name "halogen", meaning "salt producer", was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger.[19] 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 1826.[20][21] In 1823, Michael Faraday liquefied chlorine for the first time,[22][23][24] and demonstrated that what was then known as "solid chlorine" had a structure of chlorine hydrate (Cl₂·H₂O).[11]

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.[25][26] 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. 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).[25] 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.[27] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide.[28] In 1884 Chemischen Fabrik Griesheim of Germany developed another chloralkali process which entered commercial production in 1888.[29]

Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti-putrefaction agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This practice was pioneered by Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations.[30] Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) and public sanitation, particularly in swimming and drinking water.[14]

Chlorine gas was first used as a weapon on April 22, 1915 at the Second Battle of Ypres by the German Army.[31][32] The effect on the allies was devastating because the existing gas masks were difficult to deploy and had not been broadly distributed.[33][34]

Chlorine has two stable isotopes, ³⁵Cl and ³⁷Cl. These are its only two natural isotopes occurring in quantity, with ³⁵Cl making up 76% of natural chlorine and ³⁷Cl making up the remaining 24%. Both are synthesised in stars in the oxygen-burning and silicon-burning processes.[37] Both have nuclear spin 3/2+ and thus may be used for nuclear 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 nonzero nuclear quadrupole moment and resultant quadrupolar relaxation. The other chlorine isotopes are all radioactive, with half-lives too short to occur in nature primordially. Of these, the most commonly used in the laboratory are ³⁶Cl (t_1/2 = 3.0×10⁵ y) and ³⁸Cl (t_1/2 = 37.2 min), which may be produced from the neutron activation of natural chlorine.[35]

The most stable chlorine radioisotope is ³⁶Cl. The primary decay mode of isotopes lighter than ³⁵Cl is electron capture to isotopes of sulfur; that of isotopes heavier than ³⁷Cl is beta decay to isotopes of argon; and ³⁶Cl may decay by either mode to stable ³⁶S or ³⁶Ar.[38] ³⁶Cl occurs in trace quantities in nature as a cosmogenic nuclide in a ratio of about (7–10) × 10⁻¹³ to 1 with stable chlorine isotopes: it is produced in the atmosphere by spallation of ³⁶Ar by interactions with cosmic ray protons. In the top meter of the lithosphere, ³⁶Cl is generated primarily by thermal neutron activation of ³⁵Cl and spallation of ³⁹K and ⁴⁰Ca. In the subsurface environment, muon capture by ⁴⁰Ca becomes more important as a way to generate ³⁶Cl.[39][40]

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

The simplest chlorine compound is hydrogen chloride, HCl, a major chemical in industry as well as in the laboratory, both as a gas and dissolved in water as hydrochloric acid. It is often produced by burning hydrogen gas in chlorine gas, or as a byproduct of chlorinating hydrocarbons. Another approach is to treat sodium chloride with concentrated sulfuric acid to produce hydrochloric acid, also known as the "salt-cake" process:[42]

NaCl + H₂SO₄ 150 °C⟶  NaHSO₄ + HCl

NaCl + NaHSO₄ 540–600 °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 reacting benzoyl chloride with heavy water (D₂O).[42]

At room temperature, hydrogen chloride is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to 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.[42] 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 for n = 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 an azeotrope with 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.[43]

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, its dielectric constant is low and it does not dissociate appreciably into H₂Cl⁺ and HCl⁻
₂ ions – the latter, in any case, are much less stable than the bifluoride ions (HF⁻
₂) due to the very weak hydrogen bonding between hydrogen and chlorine, though its salts with very large and weakly polarising cations such as Cs⁺ and NR⁺
₄ (R = Me, Et, Buⁿ) may still be isolated. Anhydrous hydrogen chloride is a poor solvent, only able to dissolve small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides. It readily protonates electrophiles containing lone-pairs or π bonds. Solvolysis, ligand replacement reactions, and oxidations are well-characterised in hydrogen chloride solution:[44]

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)

Hydrated nickel(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 (the noble gases, with the exception of xenon in the highly unstable XeCl₂ and XeCl₄); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than chlorine's (oxygen and fluorine) so that the resultant binary compounds are formally not chlorides but rather oxides or fluorides of chlorine.[45] Even though nitrogen in NCl₃ is bearing a negative charge, the compound is usually called nitrogen trichloride.

Chlorination of metals with Cl₂ usually leads to a higher oxidation state than bromination with Br₂ when multiple oxidation states are available, such as in MoCl₅ and MoBr₃. 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 dioxide reacts with chlorine at standard conditions to produce zirconium tetrachloride, and uranium trioxide reacts with hexachloropropene when heated under reflux to give uranium tetrachloride. The second example also involves a reduction in oxidation 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:[45]

EuCl₃ + 1/2 H₂ ⟶ EuCl₂ + HCl

ReCl₅ at "bp"⟶  ReCl₃ + Cl₂

AuCl₃ 160 °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 chloride is mostly ionic, but aluminium chloride is not). Silver chloride is very insoluble in water and is thus often used as a qualitative test for chlorine.[45]

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:[46]

Cl₂ + ClF + AsF₅ −78 °C⟶ [Cl₃]⁺[AsF₆]⁻

This reaction is conducted in the oxidising solvent arsenic pentafluoride. The trichloride anion, [Cl₃]⁻, has also been characterised; it is analogous to triiodide.[47]

The three fluorides of chlorine form a subset of the interhalogen compounds, all of which are diamagnetic.[47] Some cationic and anionic derivatives are known, such as ClF⁻
₂, ClF⁻
₄, ClF⁺
₂, and Cl₂F⁺.[48] Some pseudohalides of chlorine are also known, such as cyanogen chloride (ClCN, linear), chlorine cyanate (ClNCO), chlorine thiocyanate (ClSCN, unlike its oxygen counterpart), and chlorine azide (ClN₃).[47]

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 from chlorine trifluoride and 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 attack carbon monoxide to form carbonyl chlorofluoride, COFCl. It will react analogously with hexafluoroacetone, (CF₃)₂CO, with a potassium fluoride catalyst to produce heptafluoroisopropyl hypochlorite, (CF₃)₂CFOCl; with nitriles RCN to produce RCF₂NCl₂; and with the sulfur oxides SO₂ and SO₃ to produce ClSO₂F and ClOSO₂F respectively. It will also react exothermically with compounds containing –OH and –NH groups, such as water:[47]

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, containing hydrogen, potassium, phosphorus, arsenic, antimony, sulfur, selenium, tellurium, bromine, iodine, and powdered molybdenum, tungsten, rhodium, iridium, and iron. It will also ignite water, along with many substances which in ordinary circumstances would be considered chemically inert such as asbestos, concrete, glass, and sand. When heated, it will even corrode noble metals as palladium, platinum, and gold, and even the noble gases xenon and radon do not escape fluorination. An impermeable fluoride layer is formed by sodium, magnesium, aluminium, zinc, tin, and silver, 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 with hydrazine to form hydrogen fluoride, nitrogen, and chlorine gases was used in experimental rocket engine, but has problems largely stemming from its extreme hypergolicity resulting in ignition without any measurable delay. Today, it is mostly used in nuclear fuel processing, to oxidise uranium to uranium hexafluoride for its enriching and to separate it from plutonium, as well as in the semiconductor industry, where it is used to clean chemical vapor deposition chambers.[49] It can act as a fluoride ion donor or acceptor (Lewis base or acid), although it does not dissociate appreciably into ClF⁺
₂ and ClF⁻
₄ ions.[50]

Chlorine pentafluoride (ClF₅) is made on a large scale by direct fluorination of chlorine with excess fluorine gas 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 pentafluoride and antimony pentafluoride form ionic adducts of the form [ClF₄]⁺[MF₆]⁻ (M = As, Sb) and water reacts vigorously as follows:[51]

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 unreactive perchloryl 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.[52]

Yellow chlorine dioxide (ClO₂) gas above a solution containing chlorine dioxide.

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

The chlorine oxides are well-studied in spite of their instability (all of them are endothermic compounds). They are important because they are produced when chlorofluorocarbons undergo photolysis in the upper atmosphere and cause the destruction of the ozone layer. None of them can be made from directly reacting the elements.[53]

Dichlorine monoxide (Cl₂O) is a brownish-yellow gas (red-brown when solid or liquid) which may be obtained by reacting chlorine gas with yellow mercury(II) oxide. It is very soluble in water, in which it is in equilibrium with hypochlorous acid (HOCl), of which it is the anhydride. It is thus an effective bleach and is mostly used to make hypochlorites. It explodes on heating or sparking or in the presence of ammonia gas.[53]

Chlorine dioxide (ClO₂) was the first chlorine oxide to be discovered in 1811 by Humphry 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 a chlorate as follows:[53]

ClO⁻
₃ + Cl⁻ + 2 H⁺ ⟶ ClO₂ + 1/2 Cl₂ + H₂O

Its production is thus intimately linked to the redox reactions of the chlorine oxoacids. It is a strong oxidising agent, reacting with sulfur, phosphorus, phosphorus halides, and potassium 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 ClO₂ molecules result in the radicals ClO and ClOO, while at room temperature mostly chlorine, oxygen, and some ClO₃ and Cl₂O₆ are produced. Cl₂O₃ is 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:[53]

Cl• + O₃ ⟶ ClO• + O₂

ClO• + O• ⟶ Cl• + O₂

Chlorine perchlorate (ClOClO₃) is a pale yellow liquid that is less stable than ClO₂ and decomposes at room temperature to form chlorine, oxygen, and dichlorine hexoxide (Cl₂O₆).[53] Chlorine perchlorate may also be considered a chlorine derivative of perchloric acid (HOClO₃), similar to the thermally unstable chlorine derivatives of other oxoacids: examples include chlorine nitrate (ClONO₂, vigorously reactive and explosive), and chlorine fluorosulfate (ClOSO₂F, more stable but still moisture-sensitive and highly reactive).[54] 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 anhydrous hydrogen fluoride does not proceed to completion.[53]

Dichlorine heptoxide (Cl₂O₇) is the anhydride of perchloric acid (HClO₄) and can readily be obtained from it by dehydrating it with phosphoric acid at −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 ClO₃ and ClO₄ which immediately decompose to the elements through intermediate oxides.[53]

Chlorine forms four oxoacids: hypochlorous acid (HOCl), chlorous acid (HOClO), chloric acid (HOClO₂), and perchloric 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:[41]

Cl2 + H2O	⇌ HOCl + H+ + Cl−	Kac = 4.2 × 10−4 mol2 l−2
Cl2 + 2 OH−	⇌ OCl− + H2O + Cl−	Kalk = 7.5 × 1015 mol−1 l

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 (4 ClO⁻
₃ ⇌ Cl⁻ + 3 ClO⁻
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.[41]

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 chlorite is 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 is sodium 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:[55]

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 anhydrous sodium perchlorate or barium perchlorate with 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, sets hydrogen iodide and thionyl chloride on fire and even oxidises silver and gold. Although it is a weak ligand, weaker than water, a few compounds involving coordinated ClO⁻
₄ are known.[55]

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

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

Like the other carbon–halogen bonds, the C–Cl bond is a common functional group that forms part of core organic 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 thus electrophilic. Chlorination modifies the physical properties of hydrocarbons in several ways: chlorocarbons are typically denser than water due to the higher atomic weight of chlorine versus hydrogen, and aliphatic organochlorides are alkylating agents because chloride is a leaving group.[56]

Alkanes and aryl alkanes may be chlorinated under free-radical conditions, with UV light. However, the extent of chlorination is difficult to control: the reaction is not regioselective and 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 the Friedel-Crafts halogenation, using chlorine and a Lewis acid catalyst.[56] The haloform reaction, using chlorine and sodium hydroxide, is also able to generate alkyl halides from methyl ketones, and related compounds. Chlorine adds to the multiple bonds on alkenes and alkynes as 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 as phosphorus pentachloride (PCl₅) or thionyl chloride (SOCl₂). The last is very convenient in the laboratory because all side products are gaseous and do not have to be distilled out.[56]

Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.[57][58] Chlorinated organic compounds are found in nearly every class of biomolecules including alkaloids, terpenes, amino acids, flavonoids, steroids, and fatty acids.[57][59] Organochlorides, including dioxins, 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.[60] In addition, a variety of simple chlorinated hydrocarbons including dichloromethane, chloroform, and carbon tetrachloride have been isolated from marine algae.[61] A majority of the chloromethane in the environment is produced naturally by biological decomposition, forest fires, and volcanoes.[62]

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 as DDT, are persistent organic pollutants which 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.[63] Due to the ready homolytic fission of the C–Cl bond to create chlorine radicals in the upper atmosphere, chlorofluorocarbons have been phased out due to the harm they do to the ozone layer.[53]

In France (as elsewhere), animal intestines were processed to make musical instrument strings, Goldbeater's skin and other products. This was done in "gut factories" (boyauderies), and it was an odiferous and unhealthy process. In or about 1820, the Société d'encouragement pour l'industrie nationale offered a prize for the discovery of a method, chemical or mechanical, for separating the peritoneal membrane of animal intestines without putrefaction.[72][73] The prize 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 actually retarded the decomposition.[73][30]

Labarraque's research resulted in the use of chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) in the boyauderies. The same chemicals were found to be useful in the routine disinfection and deodorization of latrines, sewers, markets, abattoirs, anatomical theatres, and morgues.[74] They were successful in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and they were beneficial during exhumations,[75] embalming, outbreaks of epidemic disease, fever, and blackleg in cattle.[72]

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 treat putrefaction of existing wounds, including septic wounds.[76] 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 the microbe was not discovered until more than half a century later.

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 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.[77] 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,[78] though three other reputable sources that describe that famous cholera epidemic do not mention the incident.[79][80][81] 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.[79]^: 296

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine and chemical base solutions was in 1847, when Ignaz Semmelweis used 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 of childbed fever ("puerperal fever") in the maternity wards of Vienna General Hospital in Austria in 1847.[82]

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 modern times, where it remains effective against bacteria that are resistant to multiple antibiotics (see Century Pharmaceuticals).[83]

Liquid Pool Chlorine

The first continuous application of chlorination to drinking U.S. water was installed in Jersey City, New Jersey, in 1908.[84] 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. Even small water supplies are now routinely chlorinated.[85]

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 amino acids in 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 of chloramine, a chemical compound produced by the reaction of free dissolved chlorine with amines in organic substances including those in urine and sweat.[86] As a disinfectant in water, chlorine is more than three times as effective against Escherichia coli as bromine, and more than six times as effective as iodine.[87] Increasingly, monochloramine itself is being directly added to drinking water for purposes of disinfection, a process known as chloramination.[88]

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, microorganisms, algae, and so on.[89][90]

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.[91][92] 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 the mucosa of the lungs to form hydrochloric acid, destructive to living tissue and potentially lethal. Human respiratory systems can be protected from chlorine gas by gas masks with activated charcoal or 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 Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, which developed methods for discharging chlorine gas against an entrenched enemy.[93] After its first use, both sides in the conflict used chlorine as a chemical weapon, but it was soon replaced by the more deadly phosgene and mustard gas.[94]

Chlorine gas was also used during the Iraq War in Anbar Province in 2007, with insurgents packing truck bombs with mortar shells 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.[95][96]

On 23 October 2014, it was reported that the Islamic State of Iraq and the Levant had used chlorine gas in the town of Duluiyah, Iraq.[97] Laboratory analysis of clothing and soil samples confirmed the use of chlorine gas against Kurdish Peshmerga Forces in a vehicle-borne improvised explosive device attack on 23 January 2015 at the Highway 47 Kiske Junction near Mosul.[98]

Another country in the middle east, Syria, has used chlorine as a chemical weapon[99] delivered from barrel bombs and rockets.[100][101] In 2016, the OPCW-UN Joint Investigative Mechanism concluded that the Syrian government used chlorine as a chemical weapon in three separate attacks.[102] Later investigations from the OPCW's Investigation and Identification Team concluded that the Syrian Air Force was responsible for chlorine attacks in 2017 and 2018.[103]

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-induced stress corrosion cracking of stainless steel suspension rods.[119] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic plumbing, and stress corrosion cracking caused widespread failures in the US in the 1980s and 1990s.[120]

Chlorine "attack" on an acetal resin plumbing joint resulting from a fractured acetal joint in a water supply system which started at an injection molding defect 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

The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.[121][122] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework that carries chlorine gas is made of steel.[121][122]