Template:Infobox chlorine Chlorine (pronounced /ˈklɔəriːn/ Template:Respell, from the Greek word 'χλωρóς' (khlôros, meaning 'pale green'), is the chemical element with atomic number 17 and symbol Cl. It is a halogen, found in the periodic table in group 17 (formerly VII, VIIa, or VIIb). As the chloride ion, which is part of common salt and other compounds, it is abundant in nature and necessary to most forms of life, including humans. In its elemental form (Cl2 or "dichlorine") under standard conditions, chlorine is a powerful oxidant and is used in bleaching and disinfectants, as well as an essential reagent in the chemical industry. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine-containing molecules such as chlorofluorocarbons have been implicated in the destruction of the ozone layer.


At standard temperature and pressure, two chlorine atoms form the diatomic molecule Cl2. This is a pale yellow-green gas that has its distinctive strong smell, the smell of bleach. The bonding between the two atoms is relatively weak (only of 242.580 ±0.004 kJ/mol) which makes the Cl2 molecule highly reactive.

Along with fluorine, bromine, iodine and astatine, chlorine is a member of the halogen series that forms the group 17 of the periodic table—the most reactive group of elements. It combines readily with nearly all elements.

Compounds with oxygen, nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.[1] Chlorine, though very reactive, is not as extremely reactive as fluorine. Pure chlorine gas does, however, support combustion of organic compounds such as hydrocarbons, although the carbon component tends to burn incompletely, with much of it remaining as soot.[2] At 10 °C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, and at 30 °C (Template:Convert/LoffT), 1 L of water dissolves only 1.77 liters of chlorine.[3]

Chlorine is a member of the salt-forming halogen series and is extracted from chlorides through oxidation often by electrolysis. With metals, it forms salts called chlorides. As the chloride ion, Cl, it is also the most abundant dissolved ion in ocean water.


Main article: Isotopes of chlorine

Chlorine has a wide range of isotopes, the two principal stable isotopes being 35Cl (75.77%) and 37Cl (24.23%); they give chlorine atoms an apparent atomic weight of 35.4527 g/mol.

Trace amounts of radioactive 36Cl exist in the environment, in a ratio of about 7x10−13 to 1 with stable isotopes. 36Cl is produced in the atmosphere by spallation of 36Ar by interactions with cosmic ray protons. In the subsurface environment, 36Cl is generated primarily as a result of neutron capture by 35Cl or muon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of 36Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of 36Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and ground water, 36Cl is also useful for dating waters less than 50 years before the present. 36Cl has seen use in other areas of the geological sciences, including dating ice and sediments.


In nature, chlorine is found primarily as the chloride ion, a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate). Over 2000 naturally-occurring organic chlorine compounds are known.[4]

Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the following chemical equation:

2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH


File:Chlor 1a.jpg

The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.[5] Hydrochloric acid was probably known to alchemist Abu Musa Jabir ibn Hayyan (Geber) around 800 AD.[6] Before 1400 AD, aqua regia (a mixture of nitric acid and hydrochloric acid) began to be used to dissolve gold,[7] and today this is still one of the few reagents that will dissolve gold. Upon dissolving gold in aqua regia, chlorine gas is released along with other nauseating and irritating gases, but this wasn't known until much more recently.

Chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and therefore he is credited for its discovery.[8] He called it "dephlogisticated muriatic acid air" since it was a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid").[8] However, he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory).[8] He named the new element within this oxide as muriaticum.[8] Regardless of what he thought, Scheele did isolate chlorine by reacting MnO2 (as the mineral pyrolusite) with HCl:

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.

A number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and an undiscovered element, muriaticum.[9]

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).[8] 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.[10]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it was an element, and not a compound.[8] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow.[11] The name halogen, meaning salt producer, was originally defined for chlorine (in 1811 by Johann Salomo Christoph Schweigger), and it was later applied to the rest of the elements in this family.[12] In 1822, Michael Faraday liquefied chlorine for the first time.[13]

Chlorine was first used to bleach textiles in 1785.[14] In 1826, silver chloride was used to produce photographic images for the first time.[15] Chloroform was first used as an anesthetic in 1847.[15] A chlorine solution in lime-water (hypochlorite) was first used as a germicide to prevent the spread of puerperal fever in the maternity wards of Vienna General Hospital in Austria in 1847,[16] and in 1850 by John Snow to disinfect the water supply in London after an outbreak of cholera. The US Department of Treasury called for all water to be disinfected with chlorine by 1918.[15] Polyvinylchloride (PVC) was invented in 1912, initially without a purpose.[15] Chlorine gas was first introduced as a weapon on April 22, 1915 at Ypres by the German Army,[17][18] and the results of this weapon were disastrous because gas masks had not yet been invented.


Main article: Chlorine production



Chlorine can be manufactured by electrolysis of a sodium chloride solution (brine). The production of chlorine results in the co-products caustic soda (sodium hydroxide, NaOH) and hydrogen gas (H2). These two products, as well as chlorine itself, are highly reactive. Chlorine can also be produced by the electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash (potassium hydroxide). There are three industrial methods for the extraction of chlorine by electrolysis of chloride solutions, all proceeding according to the following equations:

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)

Mercury cell electrolysis

Mercury cell electrolysis, also known as the Castner-Kellner process, was the first method used at the end of the nineteenth century to produce chlorine on an industrial scale.[19][20] The "rocking" cells used have been improved over the years.[21] Today, in the "primary cell", titanium anodes (formerly graphite ones) are placed in a sodium (or potassium) chloride solution flowing over a liquid mercury cathode. When a potential difference is applied and current flows, chlorine is released at the titanium anode and sodium (or potassium) dissolves in the mercury cathode forming an amalgam. This flows continuously into a separate reactor ("denuder" or "secondary cell"), where it is usually converted back to mercury by reaction with water, producing hydrogen and sodium (or potassium) hydroxide at a commercially useful concentration (50% by weight). The mercury is then recycled to the primary cell.

The mercury process is the least energy-efficient of the three main technologies (mercury, diaphragm and membrane) and there are also concerns about mercury emissions.

It is estimated that there are still around 100 mercury-cell plants operating worldwide. In Japan, mercury-based chloralkali production was virtually phased out by 1987 (except for the last two potassium chloride units shut down in 2003). In the United States, there will be only five mercury plants remaining in operation by the end of 2008. In Europe, mercury cells accounted for 43% of capacity in 2006 and Western European producers have committed to closing or converting all remaining chloralkali mercury plants by 2020.[22]

Diaphragm cell electrolysis

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.[23] This technology was also developed at the end of the nineteenth century. There are several variants of this process: the Le Sueur cell (1893), the Hargreaves-Bird cell (1901), the Gibbs cell (1908), and the Townsend cell (1904).[24][25] The cells vary in construction and placement of the diaphragm, with some having the diaphragm in direct contact with the cathode.

The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted.

As a result, diaphragm methods produce alkali that is quite dilute (about 12%) and of lower purity than do mercury cell methods. But diaphragm cells are not burdened with the problem of preventing mercury discharge into the environment. They also operate at a lower voltage, resulting in an energy savings over the mercury cell method,[25] but large amounts of steam are required if the caustic has to be evaporated to the commercial concentration of 50%.

Membrane cell electrolysis

Development of this technology began in the 1970s. The electrolysis cell is divided into two "rooms" by a cation permeable membrane acting as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.[26] Sodium (or potassium) hydroxide solution is circulated through the cathode compartment, exiting at a higher concentration. A portion of the concentrated sodium hydroxide solution leaving the cell is diverted as product, while the remainder is diluted with deionized water and passed through the electrolysis apparatus again.

This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine.

Other electrolytic processes

Although a much lower production scale is involved, electrolytic diaphragm and membrane technologies are also used industrially to recover chlorine from hydrochloric acid solutions, producing hydrogen (but no caustic alkali) as a co-product.

Furthermore, electrolysis of fused chloride salts (Downs process) also enables chlorine to be produced, in this case as a by-product of the manufacture of metallic sodium or magnesium.

Other methodsEdit

Before electrolytic methods were used for chlorine production, the direct oxidation of hydrogen chloride with oxygen or air was used in the Deacon process:

4 HCl + O2 → 2 Cl2 + 2 H2O

This reaction is accomplished with the use of copper(II) chloride (CuCl2) as a catalyst and is performed at high temperature (about 400 °C). The amount of extracted chlorine is approximately 80%. Due to the extremely corrosive reaction mixture, industrial use of this method is difficult and several pilot trials failed in the past. Nevertheless, recent developments are promising. Recently Sumitomo patented a catalyst for the Deacon process using ruthenium(IV) oxide (RuO2).[27]

Another earlier process to produce chlorine was to heat brine with acid and manganese dioxide.

2 NaCl + 2 H2SO4 + MnO2 → Na2SO4 + MnSO4 + 2 H2O + Cl2

Using this process, chemist Carl Wilhelm Scheele was the first to isolate chlorine in a laboratory. The manganese can be recovered by the Weldon process.[28]

In the latter half of the 19th century, prior to the adoption of electrolytic methods of chlorine production, there was substantial production of chlorine by these reactions to meet demand for bleach and bleaching powder for use by textile industries; by the 1880s the UK, as well as supporting its own (then not inconsiderable) domestic textile production was exporting 70,000 tons per year of bleaching powder.[29] This demand was met by capturing hydrochloric acid driven off as a gas during the production of alkali by the Leblanc process, oxidizing this to chlorine (originally by reaction with manganese dioxide), later by direct oxidation by air using the Deacon process (in which case impurities capable of poisoning the catalyst had first to be removed), and subsequently absorbing the chlorine onto lime.

Small amounts of chlorine gas can be made in the laboratory by putting concentrated hydrochloric acid in a flask with a side arm and rubber tubing attached. Manganese dioxide is then added and the flask stoppered. The reaction is not greatly exothermic. As chlorine is denser than air, it can be collected by placing the tube inside a flask where it will displace the air. Once full, the collecting flask can be stoppered.

Another method for producing small amounts of chlorine gas in a lab is by adding concentrated hydrochloric acid (typically about 5M) to sodium hypochlorite or sodium chlorate solution.

Industrial productionEdit

File:Liquid chlorine in flask.jpg

Large-scale production of chlorine involves several steps and many pieces of equipment. The description below is typical of a membrane plant. The plant also simultaneously produces sodium hydroxide (caustic soda) and hydrogen gas. A typical plant consists of brine production/treatment, cell operations, chlorine cooling & drying, chlorine compression & liquefaction, liquid chlorine storage & loading, caustic handling, evaporation, storage & loading and hydrogen handling.


Key to the production of chlorine is the operation of the brine saturation/treatment system. Maintaining a properly saturated solution with the correct purity is vital, especially for membrane cells. Many plants have a salt pile which is sprayed with recycled brine. Others have slurry tanks that are fed raw salt.

The raw brine is partially or totally treated with sodium hydroxide, sodium carbonate and a flocculant to reduce calcium, magnesium and other impurities. The brine proceeds to a large clarifier or a filter where the impurities are removed. The total brine is additionally filtered before entering ion exchangers to further remove impurities. At several points in this process, the brine is tested for hardness and strength.

After the ion exchangers, the brine is considered pure, and is transferred to storage tanks to be pumped into the cell room. Brine, fed to the cell line, is heated to the correct temperature to control exit brine temperatures according to the electrical load. Brine exiting the cell room must be treated to remove residual chlorine and control pH levels before being returned to the saturation stage. This can be accomplished via dechlorination towers with acid and sodium bisulfite addition. Failure to remove chlorine can result in damage to the cells. Brine should be monitored for accumulation of both chlorate anions and sulfate anions, and either have a treatment system in place, or purging of the brine loop to maintain safe levels, since chlorate anions can diffuse through the membranes and contaminate the caustic, while sulfate anions can damage the anode surface coating.

Cell room

The building that houses many electrolytic cells is usually called a cell room or cell house, although some plants are built outdoors. This building contains support structures for the cells, connections for supplying electrical power to the cells and piping for the fluids. Monitoring and control of the temperatures of the feed caustic and brine is done to control exit temperatures. Also monitored are the voltages of each cell which vary with the electrical load on the cell room that is used to control the rate of production. Monitoring and control of the pressures in the chlorine and hydrogen headers is also done via pressure control valves.

Direct current is supplied via a rectified power source. Plant load is controlled by varying the current to the cells. As the current is increased, flow rates for brine and caustic and deionized water are increased, while lowering the feed temperatures.

Cooling and drying

Chlorine gas exiting the cell line must be cooled and dried since the exit gas can be over 80 °C (Template:Convert/LoffT) and contains moisture that allows chlorine gas to be corrosive to iron piping. Cooling the gas allows for a large amount of moisture from the brine to condense out of the gas stream. This reduces both the cooling requirements and feed flow of sulfuric acid required in the drying towers. Cooling also improves the efficiency of both the compression and the liquefaction stage that follows. Chlorine exiting is ideally between 18 °C (Template:Convert/LoffT) and 25 °C (Template:Convert/LoffT). After cooling the gas stream passes through a series of towers with counter flowing sulfuric acid. The sulfuric acid is fed into the final tower at 98% and the first tower typically has a strength between 66% and 76% depending on materials of construction. These towers progressively remove any remaining moisture from the chlorine gas. After exiting the drying towers the chlorine is filtered to remove any remaining sulfuric acid.

Compression and liquefaction

Several methods of compression may be used: liquid ring, reciprocating, or centrifugal. The chlorine gas is compressed at this stage and may be further cooled by inter- and after-coolers. After compression it flows to the liquefiers, where it is cooled enough to liquefy. Non condensable gases and remaining chlorine gas are vented off as part of the pressure control of the liquefaction systems. These gases are routed to a gas scrubber, producing sodium hypochlorite, or used in the production of hydrochloric acid (by combustion with hydrogen) or ethylene dichloride (by reaction with ethylene).

Storage and loading

Liquid chlorine is typically gravity-fed to storage tanks. It can be loaded into rail or road tankers via pumps or padded with compressed dry gas.

Caustic handling, evaporation, storage and loading

Caustic, fed to the cell room flows in a loop that is simultaneously bled off to storage with a part diluted with deionized water and returned to the cell line for strengthening within the cells. The caustic exiting the cell line must be monitored for strength, to maintain safe concentrations. Too strong or too weak a solution may damage the membranes. Membrane cells typically produce caustic in the range of 30% to 33% by weight. The feed caustic flow is heated at low electrical loads to control its exit temperature. Higher loads require the caustic to be cooled, to maintain correct exit temperatures. The caustic exiting to storage is pulled from a storage tank and may be diluted for sale to customers who require weak caustic or for use on site. Another stream may be pumped into a multiple effect evaporator set to produce commercial 50% caustic. Rail cars and tanker trucks are loaded at loading stations via pumps.

Hydrogen handling

Hydrogen produced may be vented unprocessed directly to the atmosphere or cooled, compressed and dried for use in other processes on site or sold to a customer via pipeline, cylinders or trucks. Some possible uses include the manufacture of hydrochloric acid or hydrogen peroxide, as well as desulfurization of petroleum oils, or use as a fuel in boilers or fuel cells. In Porsgrunn the byproduct is used for the hydrogen fueling station at Hynor.

Energy consumption

Production of chlorine consumes a large amount of energy.[30] Energy consumption per unit weight of product is not far below that for iron and steel manufacture[31] and greater than for the production of glass[32] or cement.[33]

The amount of electrical energy required to produce a given amount of chlorine is fixed by the nature of the electrochemical reaction. Any energy savings, therefore, can only be made by improving the efficiency of the process and reducing ancillary energy use.


For general references to the chloride ion (Cl), including references to specific chlorides, see chloride. For other chlorine compounds see chlorate (ClO3-), chlorite (ClO2-), hypochlorite (ClO), and perchlorate (ClO4-), and chloramine (NH2Cl).[34]

Other chlorine-containing compounds include:

Oxidation statesEdit

Name Formula Example compounds
−1 chlorides Cl ionic chlorides, organic chlorides, hydrochloric acid
0 chlorine Cl2 elemental chlorine
+1 hypochlorites ClO sodium hypochlorite, calcium hypochlorite
+3 chlorites ClO2- sodium chlorite
+4 chlorine dioxide ClO2
+5 chlorates ClO3- sodium chlorate, potassium chlorate, chloric acid
+7 perchlorates ClO4- potassium perchlorate, perchloric acid, magnesium perchlorate
organic perchlorates, ammonium perchlorate, dichlorine heptoxide

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero. Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elemental chlorine gas. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide. This is due to disproportionation:

Cl2 + 2 OH → Cl + ClO + H2O

In hot concentrated alkali solution disproportionation continues:

2 ClO → Cl + ClO2-
ClO + ClO2- → Cl + ClO3-

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated, they undergo the final disproportionation step.

4 ClO3- → Cl + 3 ClO4-

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[35]

Reaction Electrode
Cl + 2 OH → ClO + H2O + 2 e +0.89 volts
ClO + 2 OHClO2- + H2O + 2 e +0.67 volts
ClO2- + 2 OHClO3- + H2O + 2 e +0.33 volts
ClO3- + 2 OHClO4- + H2O + 2 e +0.35 volts

Each step is accompanied at the cathode by

2 H2O + 2 e → 2 OH + H2 (−0.83 volts)

Applications and usesEdit

Production of industrial and consumer products Edit

Chlorine's principal applications are in the production of a wide range of industrial and consumer products.[36][37] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.

Purification and disinfectionEdit

Chlorine is an important chemical for water purification (such as water treatment plants), in disinfectants, and in bleach. Chlorine in water is more than three times more effective as a disinfectant against Escherichia coli than an equivalent concentration of bromine, and is more than six times more effective than an equivalent concentration of iodine.[38]

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. Even small water supplies are now routinely chlorinated.[2] (See also chlorination)

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


Elemental chlorine is an oxidizer. It undergoes halogen substitution reactions with lower halide salts. For example, chlorine gas bubbled through a solution of bromide or iodide anions oxidizes them to bromine and iodine respectively.

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

Like the other halides, chlorine undergoes electrophilic additions reactions, most notably, the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide.

Chlorine is used extensively in organic and inorganic chemistry as an oxidizing agent and in substitution reactions because chlorine often imparts many desired properties to an organic compound, due to its electronegativity.

Chlorine compounds are used as intermediates in the production of a number of important commercial products that do not contain chlorine. Examples are: polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose and propylene oxide.

Use as a weaponEdit

  • World War I
Main article: Poison gas in World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres. As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form hydrochloric acid, an irritant which can be lethal. The damage done by chlorine gas can be prevented by a gas mask, or other filtration method, which makes the overall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide. After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly gases phosgene and mustard gas.[39]

  • Iraq War
Main article: 2007 chlorine bombings in Iraq

Chlorine gas has also been used by insurgents against the local population and coalition forces in the Iraq War in the form of chlorine bombs. On March 17, 2007, for example, three chlorine filled trucks were detonated in the Anbar province killing two and sickening over 350.[40] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.[41] Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened up security for chlorine, which is essential for providing safe drinking water for the population.

Chlorine crackingEdit

File:Chlorine attack1.jpg

The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to stress corrosion cracking of stainless steel rods used to suspend them.[42] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic supplies, and stress corrosion cracking caused widespread failures in the USA in the 1980s and 90s. One example shows an acetal joint in a water supply system, which when it fractured, caused substantial physical damage to computers in the labs below the supply. The cracks started at injection molding defects in the joint and grew slowly until finally triggered. The fracture surface shows iron and calcium salts which were deposited in the leaking joint from the water supply before failure.

Other usesEdit

Chlorine is used in the manufacture of numerous organic chlorine compounds, the most significant of which in terms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes and trichlorobenzenes.

Chlorine is also used in the production of chlorates and in bromine extraction.

Health effectsEdit

NFPA 704
Template:NFPA 704

Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.[43]

Chlorine is detectable in concentrations of as low as 0.2 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.[3] Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.[44]

Chlorine's toxicity comes from its oxidizing power. When chlorine is inhaled at concentrations above 30ppm it begins to react with water and cells which change it into hydrochloric acid (HCl) and hypochlorous acid (HClO).

When used at specified levels for water disinfection, although chlorine reaction with water itself usually doesn't represent a major concern for human health, other materials present in the water can generate disinfection by-products that can damage human health.[45][46]

See alsoEdit


  1. Windholz, Martha et al., ed. (1976). Merck Index of Chemicals and Drugs, 9th ed.. Rahway, N.J.: Merck & Co.. ISBN 0911910263. 
  2. 2.0 2.1 Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics 81st edition, CRC press. ISBN 0849304814. 
  3. 3.0 3.1 " – Chlorine". Mark Winter [The University of Sheffield and WebElements Ltd, UK]. Retrieved on 2007-03-17.
  4. "Risk assessment and the cycling of natural organochlorines". Euro Chlor. Retrieved on 2007-08-12.
  5. "The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania". Retrieved on 2008-07-10.
  6. Pereira, Jonathan (1854). The elements of materia medica and therapeutics, Volume 1, Longman, Brown, Green, and Longmans. p. 387, 
  7. Hoover, Herbert Clark (2003). Georgius Agricola de Re Metallica, Kessinger Publishing. p. 354. ISBN 0766131971, 
  8. 8.0 8.1 8.2 8.3 8.4 8.5 "17 Chlorine". Retrieved on 2008-09-12.
  9. Ihde, Aaron John (1984). The development of modern chemistry, Courier Dover Publications. p. 158. ISBN 0486642356, 
  10. Gay-Lussac, Joseph Louis; Thénard, Louis-Jacques (1809). "On the nature and the properties of muriatic acid and of oxygenated muriatic acid". Mémoires de Physique et de Chimie de la Société d'Arcueil 2: 339–358, 
  11. Sir Humphry Davy (1811). "On a Combination of Oxymuriatic Gas and Oxygene Gas". Philosophical Transactions of the Royal Society 101: 155–162. doi:10.1098/rstl.1811.0008, 
  12. Snelders, H. A. M. (1971). "J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter". Isis 62 (3): 328, 
  13. "Discovery of Chlorine". Retrieved on 2008-07-10.
  14. "History of Chlorine". Archived from the original on 2001-08-16. Retrieved on 2008-07-10.
  15. 15.0 15.1 15.2 15.3 Jacqueline Brazin. "Chlorine & its Consequences". Retrieved on 2008-07-10.
  16. "Chlorine Story". americanchemistry. Retrieved on 2008-07-10.
  17. "Chlorine - History". Retrieved on 2008-07-10.
  18. "Weaponry: Use of Chlorine Gas Cylinders in World War I". Retrieved on 2008-07-10.
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External linksEdit

Template:Chlorine compounds Template:Diatomicelements Template:U.S. chemical weapons Template:Chemical warfare Template:E number infobox 920-929

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