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17 sulfurchlorineargon


Name, symbol, number chlorine, Cl, 17
Chemical serieshalogens
Group, period, block 173, p
Appearanceyellowish green
Standard atomic weight 35.453(2) g·mol−1
Electron configuration [Ne] 3s2 3p5
Electrons per shell 2, 8, 7
Physical properties
Density(0 °C, 101.325 kPa)
3.2 g/L
Melting point171.6 K
(-101.5 °C, -150.7 °F)
Boiling point239.11 K
(-34.4 °C, -29.27 °F)
Critical point416.9 K, 7.991 MPa
Heat of fusion(Cl2) 6.406 kJ·mol−1
Heat of vaporization(Cl2) 20.41 kJ·mol−1
Heat capacity(25 °C) (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
Crystal structureorthorhombic
Oxidation states±1, 3, 5, 7
(strongly acidic oxide)
Electronegativity3.16 (Pauling scale)
Ionization energies
1st: 1251.2 kJ·mol−1
2nd: 2298 kJ·mol−1
3rd: 3822 kJ·mol−1
Atomic radius100 pm
Atomic radius (calc.)79 pm
Covalent radius99 pm
Van der Waals radius175 pm
Magnetic orderingnonmagnetic
Electrical resistivity(20 °C) > 10 Ω·m
Thermal conductivity(300 K) 8.9x10-3  W·m−1·K−1
Speed of sound(gas, 0 °C) 206 m/s
CAS registry number7782-50-5
Selected isotopes
Main article: Isotopes of chlorine
iso NA half-life DM DE (MeV) DP
35Cl 75.77% Cl is stable with 18 neutrons
36Cl syn 3.01×105 y β- 0.709 36Ar
ε - 36S
37Cl 24.23% Cl is stable with 20 neutrons
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Chlorine (IPA: /ˈklɔəriːn/, Greek: χλωρóς chloros, 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 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 common elemental form (Cl2 or "dichlorine") under standard conditions, it is a pale green gas about 2.5 times as dense as air. It has a disagreeable, suffocating odor that is detectable in concentrations as low as 3.5 ppm[1] and is poisonous. Chlorine is a powerful oxidant and is used in bleaching and disinfectants. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine based molecules have been implicated in the destruction of the ozone layer.



Chlorine was discovered in 1774 by Swedish chemist Carl Wilhelm Scheele, who called it dephlogisticated muriatic acid (see phlogiston theory) and mistakenly thought it contained oxygen. Scheele isolated chlorine by reacting MnO2 with HCl. Chlorine was given its current name in 1810 by Sir Humphry Davy, who insisted that it was in fact an element.

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. 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.[2]

Iraq War

Chlorine gas has also been used by insurgents in the Iraq War as a chemical weapon to terrorize the local population and coalition forces. On March 17, 2007, for example, three chlorine filled trucks were detonated in the Anbar province killing 2 and sickening over 350.[3] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.[4] 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.


Main article: Isotopes of chlorine

Chlorine has isotopes with mass numbers ranging from 32 to 40. There are two principal stable isotopes, 35Cl (75.77%) and 37Cl (24.23%), giving chlorine atoms in bulk an apparent atomic weight of 35.5 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.

Notable characteristics

  Chlorine gas is diatomic, with the formula Cl2. It combines readily with all elements except O2 and N2[5] and the noble gases. Compounds with oxygen, nitrogen, and xenon are known but do not form by direct reaction of the elements.[6] Chlorine 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.[7] At 10°C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, and at 30°C, 1 L of water dissolves only 1.77 liters of chlorine.[5]

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


See also Halide minerals.

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.[8]

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


Chlorine gas extraction

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 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.[9][10] The "rocking" cells used have been improved over the years.[11] 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.[12]

Diaphragm cell electrolysis

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates cathode and anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.[13] 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).[14][15] 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 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[15], 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.[16] 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 electrolyzer 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 acids 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 methods

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

4 HCl + O2 → 2 Cl2 + 2 H2O

This reaction is accomplished with the use of CuCl2 as a catalyst and is performed at high temperarature (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.

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

2 NaCl + 2H2SO4 + 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.[17]

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 easily collected by placing the tube inside a flask where it will displace the air. Once full, the collecting flask can be stoppered.

In the laboratory, small amounts of chlorine gas can also be created by adding concentrated hydrochloric acid (typically about 5M) to sodium hypochlorite or sodium chlorate solution.

Industrial production

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 produces simultaneously sodium hydroxide (referred to in the industry as 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 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 chlorate and sulfate and either have treatment systems in place or purging of the brine loop to maintain safe levels, since chlorate can diffuse through the membranes and contaminate the caustic, while sulfate can damage the anode surface coating.

Cell room

The building that houses the 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 electrical current is supplied via rectifiers. Plant load is controlled by varying the current to the cells. As the current is increased flow rates for brine, 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 80º C 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. Cooling also improves the efficiency of the compression and liquefaction stage that follows. Chlorine exiting is ideally between 18º C and 25º C. After cooling the gas stream passes through a series of towers with counter flowing sulfuric acid. These towers progressively remove any remaining moisture from the chlorine gas. After exiting the drying towers the chlorine is filtered to remove any sulfuric acid droplets.

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 condensible 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 are hydrochloric acid or hydrogen peroxide production, desulfurization of petroleum oils and use as a fuel in boilers or fuel cells.

Energy consumption

Production of chlorine is extremely energy intensive.[18] Energy consumption per unit weight of product is not far below that for iron and steel manufacture[19] and greater than for the production of glass[20] or cement.[21]

Since electricity is an indispensable raw material for the production of chlorine, the energy consumption corresponding to the electrochemical reaction cannot be reduced. Energy savings arise primarily through applying more efficient technologies and reducing ancillary energy use.


See also Chlorine compounds

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).[22]

Other chlorine-containing compounds include:

Oxidation states

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
+5 chlorates ClO3 sodium chlorate, potassium chlorate, chloric acid
+7 perchlorates ClO4 potassium perchlorate, perchloric acid,magnesium perchlorate
organic perchlorates, ammonium perchlorate

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 + 2OH → Cl + ClO + H2O

In hot concentrated alkali solution disproportionation continues:

2ClO → 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.

4ClO3 → Cl + 3ClO4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[23]

Reaction Electrode
Cl + 2OH → ClO + H2O + 2e +0.89 volts
ClO + 2OH → ClO2 + H2O + 2e +0.67 volts
ClO2 + 2OH → ClO3 + H2O + 2e +0.33 volts
ClO3 + 2OH → ClO4 + H2O + 2e +0.35 volts

Each step is accompanied at the cathode by

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

Applications and uses

Production of industrial and consumer products

Chlorine's principal applications are in the production of a wide range of industrial and consumer products.[24] [25] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, etc.

Purification and disinfection

Chlorine is an important chemical for water purification, in disinfectants, and in bleach. It is used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. However, in most private swimming pools chlorine itself is not used, but rather sodium hypochlorite (household bleach), formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. Even small water supplies are now routinely chlorinated.[26] (See also chlorination)


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.

Other uses

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.


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.[27]

Never use ABC Dry Chemical to fight a chlorine fire, the resulting chemical reaction with the ammonium phosphate will release toxic gases and/or result in an explosion. Water fogs or CAFS should be used to extinguish the material.[27]

See also


  1. ^ Merck Index of Chemicals and Drugs, 9th ed., monograph 2065
  2. ^ Weapons of War: Poison Gas. First World Retrieved on 2007-08-12.
  3. ^ Mahdi, Basim. "Iraq gas attack makes hundreds ill", CNN, 2007-03-17. Retrieved on 2007-03-17. 
  4. ^ "'Chlorine bomb' hits Iraq village", BBC News, 2007-05-17. Retrieved on 2007-05-17. 
  5. ^ a b – Chlorine. Mark Winter [The University of Sheffield and WebElements Ltd, UK]. Retrieved on 2007-03-17.
  6. ^ Merck Index of Chemicals and Drugs, 9th ed.
  7. ^ Lange's Handbook of Chemistry, 10th ed
  8. ^ Risk assessment and the cycling of natural organochlorines. Euro Chlor. Retrieved on 2007-08-12.
  9. ^ Pauling, Linus, General Chemistry, 1970 ed., Dover publications
  10. ^ Electrolytic Processes for Chlorine and Caustic Soda. Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. Retrieved on 2007-03-17.
  11. ^ Mercury cell. Euro Chlor. Retrieved on 2007-08-15.
  12. ^ Regional Awareness-raising Workshop on Mercury Pollution. UNEP. Retrieved on 2007-10-28.
  13. ^ Diaphragm cell. Euro Chlor. Retrieved on 2007-08-15.
  14. ^ The Electrolysis of Brine. Salt Manufacturers' Association. Retrieved on 2007-03-17.
  15. ^ a b Kiefer, David M.. When the Industry Charged Ahead. Chemistry Chronicles. Retrieved on 2007-03-17.
  16. ^ Membrane cell. Euro Chlor. Retrieved on 2007-08-15.
  17. ^ The Chlorine Industry. Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. Retrieved on 2007-03-17.
  18. ^ Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Chlor-Alkali Manufacturing Industry. European Commission. Retrieved on 2007-09-02.
  19. ^ Integrated Pollution Prevention and Control (IPPC) - Best Available Techniques Reference Document on the Production of Iron and Steel. European Commission. Retrieved on 2007-09-02.
  20. ^ Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Glass Manufacturing Industry. European Commission. Retrieved on 2007-09-02.
  21. ^ Integrated Pollution Prevention and Control (IPPC) - Reference Document on Best Available Techniques in the Cement and Lime Manufacturing Industries. European Commission. Retrieved on 2007-09-02.
  22. ^ Chlorine compounds of the month. Euro Chlor. Retrieved on 2007-08-29.
  23. ^ Cotton, F. Albert and Wilkinson, Geoffrey, Advanced Inorganic Chemistry 2nd ed. John Wiley & sons, p568
  24. ^ Uses. Euro Chlor. Retrieved on 2007-08-20.
  25. ^ Chlorine Tree. Chlorine Tree. Retrieved on 2007-08-20.
  26. ^ Chlorine. Los Alamos National Laboratory. Retrieved on 2007-03-17.
  27. ^ a b "Chlorine." MSDS. Issued on October 23, 1997; Revised on November 1, 1999; Retrieved on September 8, 2007.

This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Chlorine". A list of authors is available in Wikipedia.
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