Chlorine 1

Chlorine

Chlorine

17Cl

sulfur ← chlorine → argonF↑Cl↓Br

Chlorine in the periodic table

Appearance

pale yellow-green gas

General properties

Name, symbol, number chlorine, Cl, 17

Pronunciation /ˈklɔəriːn/ KLOHR-eenor /ˈklɔərɨn/ KLOHR-ən

Element category diatomic nonmetal

Group, period, block 17 (halogens), 3, p

Standard atomic weight 35.45(1)

Chlorine 2

Electron configuration [Ne] 3s2 3p5

2, 8, 7

History

Discovery Carl Wilhelm Scheele (1774)

First isolation Carl Wilhelm Scheele (1774)

Recognized as an element by Humphry Davy (1808)

Physical properties

Phase gas

Density (0 °C, 101.325 kPa)3.2 g/L

Liquid density at b.p. 1.5625[1] g·cm−3

Melting point 171.6 K, -101.5 °C, -150.7 °F

Boiling point 239.11 K, -34.04 °C, -29.27 °F

Critical point 416.9 K, 7.991 MPa

Heat of fusion (Cl2) 6.406 kJ·mol−1

Heat of vaporization (Cl2) 20.41 kJ·mol−1

Molar heat capacity (Cl2)33.949 J·mol−1·K−1

Vapor pressure

P (Pa) 1 10 100 1 k 10 k 100 k

at T (K) 128 139 153 170 197 239

Atomic properties

Oxidation states 7, 6, 5, 4, 3, 2, 1, -1(strongly acidic oxide)

Electronegativity 3.16 (Pauling scale)

Ionization energies(more)

1st: 1251.2 kJ·mol−1

2nd: 2298 kJ·mol−1

3rd: 3822 kJ·mol−1

Covalent radius 102±4 pm

Chlorine 3

Van der Waals radius 175 pm

Miscellanea

Crystal structure orthorhombic

Magnetic ordering diamagnetic[2]

Electrical resistivity (20 °C) > 10 Ω·m

Thermal conductivity 8.9×10−3  W·m−1·K−1

Speed of sound (gas, 0 °C) 206 m·s−1

CAS registry number 7782-50-5

Most stable isotopes

Main article: Isotopes of chlorine

iso NA half-life DM DE (MeV) DP

35Cl 75.77% 35Cl is stable with 18 neutrons

36Cl trace 3.01×105 y β− 0.709 36Ar

ε - 36S

37Cl 24.23% 37Cl is stable with 20 neutrons

Chlorine is a chemical element with symbol Cl and atomic number 17. Chlorine is in the halogen group (17) and isthe second lightest halogen after fluorine. The element is a yellow-green gas under standard conditions, where itforms diatomic molecules. It has the highest electron affinity and the third highest electronegativity of all theelements; for this reason, chlorine is a strong oxidizing agent. Free chlorine is rare on Earth, and is usually a result ofdirect or indirect oxidation by oxygen.The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times.Around 1630 chlorine gas was first synthesized in a chemical reaction, but not recognized as a fundamentallyimportant substance. Characterization of chlorine gas was made in 1774 by Carl Wilhelm Scheele, who supposed itan oxide of a new element. In 1809 chemists suggested that the gas might be a pure element, and this was confirmedby Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρóς khlôros "pale green".Nearly all chlorine in the Earth's crust occurs as chloride in various ionic compounds, including table salt. It is thesecond most abundant halogen and 21st most abundant chemical element in Earth's crust. Elemental chlorine iscommercially produced from brine by electrolysis. The high oxidizing potential of elemental chlorine ledcommercially to free chlorine's bleaching and disinfectant uses, as well as its many uses of an essential reagent in thechemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds ofthem organic chemicals such as polyvinyl chloride, as well as many intermediates for production of plastics andother end products which do not contain the element. As a common disinfectant, elemental chlorine andchlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary.In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone

Chlorine 4

depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite inneutrophils, as part of the immune response against bacteria. Elemental chlorine at high concentrations is extremelydangerous and poisonous for all living organisms, and was historically used in World War I as the first gaseouschemical warfare agent.

Characteristics

Physical characteristics of chlorine and its compounds

Chlorine, liquefied under a pressure of 7.4 bar atroom temperature, displayed in a quartz ampule

embedded in acrylic glass.

At standard temperature and pressure, two chlorine atoms form thediatomic molecule Cl2.[3] This is a yellow-green gas that has adistinctive strong odor, familiar to most from common householdbleach.[4] The bonding between the two atoms is relatively weak (only242.580 ± 0.004 kJ/mol), which makes the Cl2 molecule highlyreactive. The boiling point at regular atmosphere is around −34 ˚C, butit can be liquefied at room temperature with pressures above 740kPa.[5]

Although elemental chlorine is yellow-green, chloride ion, in commonwith other halide ions, has no color in either minerals or solutions(example, table salt). Similarly, (again as with other halogens) chlorineatoms impart no color to organic chlorides when they replace hydrogenatoms in colorless organic compounds, such as tetrachloromethane.The melting point and density of these compounds is increased by

substitution of hydrogen in place of chlorine. Compounds of chlorine with other halogens, however, as well as manychlorine oxides, are visibly colored.

Chemical characteristicsAlong with fluorine, bromine, iodine, and astatine, chlorine is a member of the halogen series that forms the group17 (formerly VII, VIIA, or VIIB) of the periodic table. Chlorine forms compounds with almost all of the elements togive compounds that are usually called chlorides. Chlorine gas reacts with most organic compounds, and will evensluggishly support the combustion of hydrocarbons.[]

Hydrolysis of free chlorine or disproportionation in water

At 25 °C and atmospheric pressure, one liter of water dissolves 3.26 g or 1.125 L of gaseous chlorine.[6] Solutions ofchlorine in water contain chlorine (Cl2), hydrochloric acid, and hypochlorous acid:

Cl2 + H2O HCl + HClOThis conversion to the right is called disproportionation, because the ingredient chlorine both increases and decreasesin formal oxidation state. The solubility of chlorine in water is increased if the water contains dissolved alkalihydroxide, and in this way, chlorine bleach is produced.[7]

Cl2 + 2 OH– → ClO– + Cl– + H2OChlorine gas only exists in a neutral or acidic solution.

Chlorine 5

Chemistry and compounds

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero and four inchlorine dioxide (see table below, and also structures in chlorite).[8] Progressing through the states, hydrochloric acidcan be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elementalchlorine gas.[9]

Oxidationstate

Name Formula Characteristic compounds

−1 chlorides Cl− ionic chlorides, organic chlorides, hydrochloric acid

0 chlorine Cl2 elemental chlorine

+1 hypochlorites ClO− sodium hypochlorite, calcium hypochlorite

+3 chlorites ClO−2

sodium chlorite

+4 chlorine(IV) ClO2

chlorine dioxide

+5 chloryl,chlorates

ClO−3 ClO+

2

potassium chlorate, chloric acid, dichloryl trisulfate [ClO2]2[S3O10].

+6 chlorine(VI) Cl2O6

dichlorine hexoxide (gas). In liquid or solid disproportionates to mix of +5 and +7 oxidation states, asionic chloryl perchlorate [ClO2]+[ClO4]−

+7 perchlorates ClO−4

perchloric acid, perchlorate salts such as magnesium perchlorate, dichlorine heptoxide

Chlorides

Chlorine combines with almost all elements to give chlorides. Compounds with oxygen, nitrogen, xenon, andkrypton are known, but do not form by direct reaction of the elements.[10] Chloride is one of the most commonanions in nature. Hydrogen chloride and its aqueous solution, hydrochloric acid, are produced on megaton scaleannually both as valued intermediates but sometimes as undesirable pollutants.

Chlorine oxides

Chlorine forms a variety of oxides, as seen above: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine hexoxide (Cl2O6), dichlorine heptoxide (Cl2O7). The anionic derivatives of these same oxides are also well known including chlorate (ClO− 3), chlorite (ClO− 2), hypochlorite (ClO−), and perchlorate (ClO− 4). The acid derivatives of these anions are hypochlorous acid (HOCl), chloric acid (HClO3) and perchloric acid (HClO4). The chloroxy cation chloryl (ClO2

+) is known and has the same structure as chlorite but with a positive charge and chlorine in the +5 oxidation state.[11] The compound "chlorine trioxide" does not occur, but rather in gas form is found as the dimeric dichlorine hexoxide (Cl2O6) with a +6 oxidation state. This compound in liquid or solid form disproportionates to a mixture of +5 and +7 oxidation states, occurring as the ionic compound chloryl perchlorate, [ClO 2]+ [ClO 4]−

Chlorine 6

.[12]

In hot concentrated alkali solution hypochlorite disproportionates:2 ClO− → Cl− + ClO−

2

ClO− + ClO−

2 → Cl− + ClO−

3

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

4 ClO−

3 → Cl− + 3 ClO−

4

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

Reaction Electrodepotential

Cl− + 2 OH− → ClO− + H2O + 2 e− +0.89 volts

ClO− + 2 OH− → ClO−2 + H2O + 2 e−

+0.67 volts

ClO−2 + 2 OH− → ClO−

3 + H2O + 2 e−

+0.33 volts

ClO−3 + 2 OH− → ClO−

4 + H2O + 2 e−

+0.35 volts

Each step is accompanied at the cathode by2 H2O + 2 e− → 2 OH− + H2 (−0.83 volts)

Interhalogen compounds

Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively. However, it cannot oxidize fluoridesalts to fluorine. It makes a variety of interhalogen compounds, such as the chlorine fluorides, chlorine monofluoride(ClF), chlorine trifluoride (ClF3), chlorine pentafluoride (ClF5). Chlorides of bromine and iodine are also known.[13]

Organochlorine compounds

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

Chlorine 7

OccurrenceIn the interstellar medium, chlorine is produced in supernovae via the r-process.[14]

In meteorites and on Earth, chlorine is found primarily as the chloride ion which occurs in minerals. In the Earth'scrust, chlorine is present at average concentrations of about 126 parts per million,[15] predominantly in such mineralsas halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloridehexahydrate).Chloride is a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the massof seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in undergroundbrine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found inabundance in dry climates or deep underground.Over 2000 naturally occurring organic chlorine compounds are known.[]

IsotopesChlorine has a wide range of isotopes. The two stable isotopes are 35Cl (75.77%) and 37Cl (24.23%).[] Together theygive chlorine an atomic weight of 35.4527 g/mol. The half-integer value for chlorine's weight caused some confusionin the early days of chemistry, when it had been postulated that atoms were composed of even units of hydrogen (seeProust's law), and the existence of chemical isotopes was unsuspected.[]

Trace amounts of radioactive 36Cl exist in the environment, in a ratio of about 7x10−13 to 1 with stable isotopes. 36Clis produced in the atmosphere by spallation of 36Ar by interactions with cosmic ray protons. In the subsurfaceenvironment, 36Cl is generated primarily as a result of neutron capture by 35Cl or muon capture by 40Ca. 36Cl decaysto 36S and to 36Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotopemakes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of 36Clwere produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and1958. The residence time of 36Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soiland ground water, 36Cl is also useful for dating waters less than 50 years before the present. 36Cl has seen use inother areas of the geological sciences, including dating ice and sediments.[]

HistoryThe most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists havefound evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.[16] Around 1630, chlorinewas recognized as a gas by the Belgian chemist and physician Jan Baptist van Helmont.[17]

Carl Wilhelm Scheele

Elemental chlorine was first prepared and studied in 1774 by Swedish chemistCarl Wilhelm Scheele, and, therefore, he is credited for its discovery.[] Hecalled it "dephlogisticated muriatic acid air" since it is a gas (then called "airs")and it came from hydrochloric acid (then known as "muriatic acid").[] However,he failed to establish chlorine as an element, mistakenly thinking that it was theoxide obtained from the hydrochloric acid (see phlogiston theory).[] He namedthe new element within this oxide as muriaticum.[] Regardless of what hethought, Scheele did isolate chlorine by reacting MnO2 (as the mineralpyrolusite) with HCl:[17]

4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2Scheele observed several of the properties of chlorine: the bleaching effect onlitmus, the deadly effect on insects, the yellow green color, and the smell similarto aqua regia.[18]

Chlorine 8

At the time, common chemical theory was: any acid is a compound that contains oxygen (still sounding in theGerman and Dutch names of oxygen: sauerstoff or zuurstof, both translating into English as acid stuff), so a numberof chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be acombination of oxygen and the yet undiscovered element, muriaticum.[19][20][]

In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid airby reacting it with charcoal to release the free element muriaticum (and carbon dioxide).[] They did not succeed andpublished a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, butwere not convinced.[21]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it is an element, and not acompound.[] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaninggreen-yellow.[22] The name halogen, meaning "salt producer," was originally used for chlorine in 1811 by JohannSalomo Christoph Schweigger. However, this term was later used as a generic term to describe all the elements inthe chlorine family (fluorine, bromine, iodine), after a suggestion by Jöns Jakob Berzelius in 1842.[23][24] In 1823,Michael Faraday liquefied chlorine for the first time,[25][26] and demonstrated that what was then known as "solidchlorine" had a structure of chlorine hydrate (Cl2·H2O).[17]

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.[][] Modern bleachesresulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in thetown of Javel (now part of Paris, France), by passing chlorine gas through a solution of sodium carbonate. Theresulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of sodium hypochlorite. However,this process was not very efficient, and alternative production methods were sought. Scottish chemist andindustrialist Charles Tennant first produced a solution of calcium hypochlorite ("chlorinated lime"), then solidcalcium hypochlorite (bleaching powder).[] These compounds produced low levels of elemental chlorine, and couldbe more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purifiedto 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 sodiumhydroxide and chlorine gas, which then mixed to form sodium hypochlorite.[27] This is known as the chloralkaliprocess, first introduced on an industrial scale in 1892, and now the source of essentially all modern elementalchlorine and sodium hydroxide production (a related low-temperature electrolysis reaction, the Hooker process, isnow responsible for bleach and sodium hypochlorite production).Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first usedas anti-putrification agents and disinfectants in the 1820s, in France, long before the establishment of the germtheory of disease. This work is mainly due to Antoine-Germain Labarraque, who adapted Berthollet's "Javel water"bleach and other chlorine preparations for the purpose (for a more complete history, see below). Elemental chlorinehas since served a continuous function in topical antisepsis (wound irrigation solutions and the like) as well as publicsanitation (especially of swimming and drinking water).In 1826, silver chloride was used to produce photographic images for the first time.[] Chloroform was first used as ananesthetic in 1847.[]

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

Chlorine gas was first introduced as a weapon on April 22, 1915, at Ypres by the German Army,[28][29] and theresults of this weapon were disastrous because gas masks had not been mass distributed and were tricky to get onquickly.

Chlorine 9

Production

Liquid chlorine analysis

In industry, elemental chlorine is usually produced by the electrolysis of sodiumchloride dissolved in water. This method, the chloralkali process industrialized in1892, now provides essentially all industrial chlorine gas.[30] Along withchlorine, the method yields hydrogen gas and sodium hydroxide (with sodiumhydroxide actually being the most crucial of the three industrial productsproduced by the process). The process proceeds according to the followingchemical equation:[9]

2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOHThe electrolysis of chloride solutions all proceed according to the followingequations:

Cathode: 2 H+ (aq) + 2 e− → H2 (g)Anode: 2 Cl− (aq) → Cl2 (g) + 2 e−

Overall process: 2 NaCl (or KCl) + 2 H2O → Cl2 + H2 + 2 NaOH (or KOH)In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragmseparates a cathode and an anode, preventing the chlorine forming at the anodefrom re-mixing with the sodium hydroxide and the hydrogen formed at thecathode.[] The salt solution (brine) is continuously fed to the anode compartmentand flows through the diaphragm to the cathode compartment, where the causticalkali is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkalibut they are not burdened with the problem of preventing mercury discharge into the environment and they are moreenergy efficient. Membrane cell electrolysis employ permeable membrane as an ion exchanger. Saturated sodium (orpotassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.[] Thismethod is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about32% concentration, but requires very pure brine.

Membrane cell process for chloralkali production

Chlorine 10

Laboratory methodsSmall amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid and manganesedioxide. Alternatively a strong acid such as sulfuric acid or hydrochloric acid reacts with sodium hypochloritesolution to release chlorine gas but reacts with sodium chlorate to produce chlorine gas and chlorine dioxide gas aswell. In the home, accidents occur when hypochlorite bleach solutions are combined with certain acidicdrain-cleaners.

Applications

Production of industrial and consumer productsPrincipal applications of chlorine are in the production of a wide range of industrial and consumer products.[][] Forexample, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals andpharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.Many important industrial products are produced via organochlorine intermediates. Examples includepolycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose, and propylene oxide.Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organiccompounds. When applied to organic substrates, reaction is often—but not invariably—non-regioselective, and,hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well,so multiple substitutions are common. If the different reaction products are easily separated, e.g., by distillation,substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) maybe a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride,chloroform, and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene, andtetrachloroethylene from 1,2-dichloroethane.Quantitatively, about 63% and 18% of all elemental chlorine produced is used in the manufacture of organic andinorganic chlorine compounds, respectively,[30]. About 15,000 chlorine compounds are being used commercially.[18]

The remaining 19% is used for bleaches and disinfection products.[30] The most significant of organic compounds interms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Otherparticularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride,trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes, andtrichlorobenzenes. The major inorganic compounds include HCl, Cl2O, HOCl, NaClO3, chlorinated isocyanurates,AlCl3, SiCl4, SnCl4, PCl3, PCl5, POCl3, AsCl3, SbCl3, SbCl5, BiCl3, S2Cl2, SCl2, SOCI2, CIF3, ICl, ICl3, TiCl3,TiCl4, MoCl5, FeCl3, ZnCl2, etc.[30][31]

Chlorine 11

Public sanitation, disinfection, and antisepsis

Combating putrefaction

Antoine-Germain Labarraque

In France (as elsewhere) there was a need to process animal guts inorder to make musical instrument strings, Goldbeater's skin and otherproducts. This was carried out in "gut factories" (boyauderies) as anodiferous and unhealthy business. In or about 1820, the Sociétéd'encouragement pour l'industrie nationale offered a prize for thediscovery of a method, chemical or mechanical, that could be used toseparate the peritoneal membrane of animal intestines without causingputrefaction.[][] It was won by Antoine-Germain Labarraque, a 44year-old French chemist and pharmacist who had discovered thatBerthollet's chlorinated bleaching solutions ("Eau de Javel") not onlydestroyed the smell of putrefaction of animal tissue decomposition, butalso retarded the decomposition process itself.[][]

Labarraque's research resulted in chlorides and hypochlorites of lime(calcium hypochlorite) and of sodium (sodium hypochlorite) beingemployed not only in the boyauderies but also for the routine disinfection and deodorisation of latrines, sewers,markets, abattoirs, anatomical theatres and morgues.[] They were also used, with success, in hospitals, lazarets,prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and for exhumations,[32]

embalming, during outbreaks of epidemic illness, fever, blackleg in cattle, etc.[]

Against infection and contagion

Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called"contagious infection", and presumed to be transmitted by "miasmas") and also to treat putrefaction of existingwounds, including septic wounds.[33] In this 1828 work, Labarraque recommended for the doctor to breathe chlorine,wash his hands with chlorinated lime, and even sprinkle chlorinated lime about the patient's bed, in cases of"contagious infection." In 1828, it was well known that some infections were contagious, even though the agency ofthe microbe was not to be realized or discovered for more than half a century.During the Paris cholera outbreak of 1832, large quantities of so-called chloride of lime were used to disinfect thecapital. This was not simply modern calcium chloride, but contained chlorine gas dissolved in lime-water (dilutecalcium hydroxide) to form calcium hypochlorite (chlorinated lime). Labarraque's discovery helped to remove theterrible stench of decay from hospitals and dissecting rooms, and, by doing so, effectively deodorised the LatinQuarter of Paris.[34] These "putrid miasmas" were thought by many to be responsible for the spread of "contagion"and "infection" – both words used before the germ theory of infection. The use of chloride of lime was based ondestruction of odors and "putrid matter." One source has claimed that chloride of lime was used by Dr. John Snow todisinfect water from the cholera-contaminated well feeding the Broad Street pump in 1854 London.[35] Threereputable sources that described the famous Broad Street pump cholera epidemic do not mention Snow performingany disinfection of water from that well.[36][37][38] Instead, one reference makes it clear that chloride of lime wasused to disinfect the offal and filth in the streets surrounding the Broad Street pump—a common practice inmid-nineteenth century England.[36]:296

Chlorine 12

Semmelweis and experiments with antisepsis

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine andchemical base solutions was in 1847, when Ignaz Semmelweis used(first) chlorine-water (simply chlorine dissolved in pure water), thencheaper chlorinated lime solutions, to deodorize the hands of Austriandoctors, which Semmelweis noticed still carried the stench ofdecomposition from the dissection rooms to the patient examinationrooms. Semmelweis, still long before the germ theory of disease, hadtheorized that "cadaveric particles" were somehow transmitting decayfrom fresh medical cadavers to living patients, and he used thewell-known "Labarraque's solutions" as the only known method toremove the smell of decay and tissue decomposition (which he foundthat soap did not). The solutions proved to be far more effectivegermicide antiseptics than soap (Semmelweis was also aware of theirgreater efficacy, but not the reason), and this resulted in Semmelweis's(later) celebrated success in stopping the transmission of childbed fever("puerperal fever") in the maternity wards of Vienna General Hospitalin Austria in 1847.[39]

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 (whogave full credit to Labarraque's prior work in this area). Called Dakin's solution, the method of wound irrigation withchlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modernantibiotic era. A modified version of this solution continues to be employed in wound irrigation in the modern era,where it remains effective against multiply antibiotic resistant bacteria (see Century Pharmaceuticals).

Public sanitation

By 1918, the US Department of Treasury called for all drinking water to be disinfected with chlorine. Chlorine ispresently an important chemical for water purification (such as in water treatment plants), in disinfectants, and inbleach. Chlorine in water is more than three times as effective as a disinfectant against Escherichia coli than anequivalent concentration of bromine, and is more than six times more effective than an equivalent concentration ofiodine.[40]

Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking watersupplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather sodiumhypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. Thedrawback of using chlorine in swimming pools is that the chlorine reacts with a human's hair and skin because hairand skin are made from protein (see Hypochlorous acid#Reaction with protein amino groups). Once the chlorinereacts with the hair and skin, it becomes chemically bonded. Even small water supplies are now routinelychlorinated.[]

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of addingchlorine are used. These include hypochlorite solutions, which gradually release chlorine into the water, andcompounds like sodium 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 beused 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 ageneral biocide, killing germs, micro-organisms, algae, and so on.[41][42]

Chlorine 13

Use as a weapon

World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 inthe Second Battle of Ypres.[43] As described by the soldiers it had a distinctive smell of a mixture between pepperand pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in themucosa of the lungs to form hydrochloric acid, an irritant that can be lethal. The damage done by chlorine gas can beprevented by a gas mask, or other filtration method, which makes the overall chance of death by chlorine gas muchlower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate,Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IGFarben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber'srole in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide.[44] After its first use,chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly phosgene andmustard gas.[]

Iraq War

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

Chlorine induced cracking in structural materials

Chlorine "attack" on an acetal resin plumbingjoint.

The element is widely used for purifying water owing to its powerfuloxidizing properties, especially potable water supplies and water usedin swimming pools. Several catastrophic collapses of swimming poolceilings have occurred owing to chlorine induced stress corrosioncracking of stainless steel rods used to suspend them.[45] Somepolymers are also sensitive to attack, including acetal resin andpolybutene. Both materials were used in hot and cold water domesticsupplies, and stress corrosion cracking caused widespread failures inthe USA in the 1980s and 1990s. The picture on the right shows anacetal joint in a water supply system, which, when it fractured, causedsubstantial physical damage to computers in the labs below the supply.The cracks started at injection molding defects in the joint and slowly grew until finally triggered. The fracturesurface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.[46]

Health effects of the free element and hazards

NFPA 704

Chlorine 14

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

Chlorine is detectable with measuring devices in concentrations of as low as 0.2 parts per million (ppm), and bysmell at 3 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can befatal after a few deep breaths of the gas.[18] Breathing lower concentrations can aggravate the respiratory system, andexposure to the gas can irritate the eyes.[] The toxicity of chlorine comes from its oxidizing power. When chlorine isinhaled at concentrations above 30 ppm, it begins to react with water and cells, which change it into hydrochloricacid (HCl) and hypochlorous acid (HClO).When used at specified levels for water disinfection, the reaction of chlorine with water is not a major concern forhuman health. Other materials present in the water may generate disinfection by-products that are associated withnegative effects on human health,[47][48] however, the health risk is far lower than drinking undisinfected water.

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

References[1] Chlorine (http:/ / encyclopedia. airliquide. com/ Encyclopedia. asp?LanguageID=11& CountryID=19& Formula=& GasID=13&

UNNumber=& EquivGasID=3& PressionBox=& VolLiquideBox=& MasseLiquideBox=& VolGasBox=& MasseGasBox=& RD20=29&RD9=8& RD6=64& RD4=2& RD3=22& RD8=27& RD2=20& RD18=41& RD7=18& RD13=71& RD16=35& RD12=31& RD19=34&RD24=62& RD25=77& RD26=78& RD28=81& RD29=82), Gas Encyclopaedia, Air Liquide

[2] Magnetic susceptibility of the elements and inorganic compounds (http:/ / www-d0. fnal. gov/ hardware/ cal/ lvps_info/ engineering/elementmagn. pdf), in

[6][6] Wiberg 2001, p. 409.[7] Greenwood 1997, pp. 857–858.[8][8] Greenwood 1997, p. 806.[9][9] Wiberg 2001, p. 408.[11] Greenwood 1997, pp. 844–850.[12][12] Greenwood 1997, p. 849.[15][15] Greenwood 1997, p. 795.[17][17] Greenwood 1997, p. 790.[18][18] Greenwood 1997, p. 793.[19][19] Greenwood 1997, p. 792.[23][23] Greenwood 1997, p. 789.[30][30] Greenwood 1997, p. 798.[31][31] Wiberg 2001, p. 412.[33] Scott, James (trans.). On the disinfecting properties of Labarraque's preparations of chlorine (http:/ / books. google. co. uk/

books?id=pD0XAQAAMAAJ& printsec=frontcover#v=onepage& q& f=false) (S. Highley, 1828) Accessed Nov 1, 2011.[34] Corbin, Alain. The Foul and the Fragrant: Odor and the French Social Imagination (http:/ / books. google. com/

books?id=LI1M4sLcvPAC& printsec=frontcover#v=onepage& q& f=false) (Harvard University Press, 1988) pp. 121–2.[35] Black and Veatch Corp. (2010). White's Handbook of Chlorination and Alternative Disinfectants. Hoboken, NJ:Wiley. first page, Chapter 9.[36] Vinten-Johansen, Peter, Howard Brody, Nigel Paneth, Stephen Rachman and Michael Rip. (2003). Cholera, Chloroform, and the Science of

Medicine. New York:Oxford University.[37] Hemphill, Sandra. (2007). The Strange Case of the Broad Street Pump: John Snow and the Mystery of Cholera. Los Angeles:University of

California[38] Johnson, Steven. (2006). The Ghost Map: The Story of London’s Most Terrifying Epidemic and How It Changed Science, Cities and the

Modern World. New York :Riverhead Books

Chlorine 15

[41][41] Greenwood 1997, p. 860.[42][42] Wiberg 2001, p. 411.[43] "Battle of Ypres" The Canadian Encyclopedia

Bibliography• Greenwood, Norman N; Earnshaw, Alan (1997). Chemistry of the Elements (2 ed.). Oxford:

Butterworth-Heinemann. ISBN 0-08-037941-9.• Wiberg, Egon; Wiberg, Nils and Holleman, Arnold Frederick (2001). Inorganic Chemistry. Academic Press.

ISBN 0-12-352651-5.

External links• Chlorine (http:/ / www. periodicvideos. com/ videos/ 017. htm) at The Periodic Table of Videos (University of

Nottingham)• Agency for Toxic Substances and Disease Registry: Chlorine (http:/ / www. atsdr. cdc. gov/ substances/

toxsubstance. asp?toxid=36)• Electrolytic production (http:/ / electrochem. cwru. edu/ encycl/ art-b01-brine. htm)• Production and liquefaction of chlorine (http:/ / www. amazingrust. com/ Experiments/ how_to/ Liquid_Cl2.

html)• Chlorine Production Using Mercury, Environmental Considerations and Alternatives (http:/ / www. oceana. org/

chlorine)• National Pollutant Inventory – Chlorine (http:/ / www. npi. gov. au/ database/ substance-info/ profiles/ 20. html)• National Institute for Occupational Safety and Health – Chlorine Page (http:/ / www. cdc. gov/ niosh/ topics/

chlorine/ )• Chlorine Institute (http:/ / www. chlorineinstitute. org/ ) – Trade association representing the chlorine industry• Chlorine Online (http:/ / www. eurochlor. org/ ) – the web portal of Eurochlor – the business association of the

European chlor-alkali industry

Article Sources and Contributors 16

Article Sources and ContributorsChlorine  Source: http://en.wikipedia.org/w/index.php?oldid=569686846  Contributors: 000n, 0612, 12dstring, 2001:558:6008:4:1DC0:E03D:91CA:6783, 213.106.152.xxx, 28bytes,2L84MEANDU, A5b, A8UDI, ABF, APendleton, ATSDR, Aa35te, Aboutmovies, Acdx, Ace of Spades IV, Acroterion, ActivExpression, Adashiel, AdjustShift, Adrian.benko, Ageofthewolf,Ahoerstemeier, Ajrocke, Alan Liefting, Alan Peakall, Alansohn, Alchemist-hp, Ale jrb, Alexlewin, Alfie66, Aljasm, Allstarecho, AltiusBimm, Alureiter, Alxndr, Anand Karia, Anastrophe,Anaxial, Andres, Andrewc141, Andrewtindal, Anon0096, Antandrus, Anturiaethwr, Any Moose, Archimerged, Aristotle28, Audiosmurf, Audirs8, Aussie Alchemist, Avicennasis, AxelBoldt,Az1568, B'Rat Ud, Bcorr, Beetstra, Bejnar, Benbest, Bender235, Benjah-bmm27, Bennybp, Bhadani, Big Bird, Bighalonut, Billy431, Biodragon, BlueOrb, BluefieldWV, Bobblewik, Bomac,Bonadea, Bovineboy2008, Bowlhover, Bradydavis, Brandonrush, Brian0918, Brvman, Bryan 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Insomesia, Iridescent, Isaac, Itfc+canes=me,Ithunn, Itub, IvoShandor, J Di, J-stan, J.delanoy, J8079s, JAn Dudík, JForget, JSarek, Jackollie, James Lum K.H., James086, Jaraalbe, Jauhienij, Javierito92, Jayron32, Jeff G., Jenniferrrxb,Jim1138, Jimjamjak, Jirt, Jjron, Jmundo, Jncraton, Joanjoc, Jodupouy, John, John Millikin, John Quincy Adding Machine, John Vandenberg, JohnSRoberts99, Jonathan tcn, JonoBurger, Jordanp,Jordi.1991, Jose77, Jovianeye, Jrobbinz123, Jrugordon, Julesd, Justin00220, Kaelia, Kakofonous, Karlhahn, Karuna8, Kazikameuk, Keenan Pepper, Kelly Martin, Kenz0198, KevinCable,Khazar2, Kingcobra333, Kingpin13, Kirk Hilliard, Klemen Kocjancic, Kntrabssi, KostasG, Kozuch, Krukowski, KuRiZu, Kukini, Kurykh, Kwamikagami, L Kensington, LA2, La Parka YourCar, LachlanA, Lamro, Lanthanum-138, Laughcosts, Lavateraguy, LeadSongDog, LeaveSleaves, Leszek Jańczuk, LiDaobing, Lightmouse, LittleOldMe, Llort, Lockesdonkey, Logical2u,Longhair, LorenzoB, Lucius Winslow, LuigiManiac, Luk, 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Image Sources, Licenses and ContributorsFile:Transparent.gif  Source: http://en.wikipedia.org/w/index.php?title=File:Transparent.gif  License: Public Domain  Contributors: Edokterfile:Chlorine ampoule.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Chlorine_ampoule.jpg  License: Creative Commons Attribution-Sharealike 3.0  Contributors:Chemicalinterest, Soerfm, WoelenFile:Electron shell 017 Chlorine (diatomic nonmetal) - no label.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Electron_shell_017_Chlorine_(diatomic_nonmetal)_-_no_label.svg License: Creative Commons Attribution-Sharealike 3.0  Contributors: User:DePiepFile:Orthorhombic.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Orthorhombic.svg  License: GNU Free Documentation License  Contributors: Original PNGs by Daniel Mayer,traced in Inkscape by User:StanneredFile:Chlorine liquid in an ampoule.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Chlorine_liquid_in_an_ampoule.jpg  License: Free Art License  Contributors: Alchemist-hp(talk) ( www.pse-mendelejew.de)Image:Equilibrium.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Equilibrium.svg  License: Public Domain  Contributors: L'AquatiqueFile:PSM V31 D740 Carl Wilhelm Scheele.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:PSM_V31_D740_Carl_Wilhelm_Scheele.jpg  License: Public Domain  Contributors:Ineuw, Kilom691File:Liquid chlorine in flask.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Liquid_chlorine_in_flask.jpg  License: Creative Commons Attribution 3.0  Contributors:Workingclass91File:Chloralkali membrane.svg  Source: http://en.wikipedia.org/w/index.php?title=File:Chloralkali_membrane.svg  License: Creative Commons Attribution-Sharealike 3.0  Contributors:JkwchuiFile:A G Barraque.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:A_G_Barraque.jpg  License: Public Domain  Contributors: UnknownFile:Ignaz Semmelweis 1860.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Ignaz_Semmelweis_1860.jpg  License: Public Domain  Contributors: Jenő DobyFile:Chlorine attack1.jpg  Source: http://en.wikipedia.org/w/index.php?title=File:Chlorine_attack1.jpg  License: Public Domain  Contributors: Peterlewis

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