Chlorine

Chlorine
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17 sulfur ← chlorine → argon
F
↑
Cl
↓
Br
Periodic Table - Extended Periodic Table


General
Name, Symbol, Number chlorine, Cl, 17
Chemical series halogens
Group, Period, Block 17, 3, p
Appearance yellowish green

Standard atomic weight 35.453(2) g·mol−1
Electron configuration [Ne] 3s2 3p5
Electrons per shell 2, 8, 7
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
3.2 g/L
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
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 structure orthorhombic
Oxidation states ±1, 3, 5, 7
(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
Atomic radius 100 pm
Atomic radius (calc.) 79 pm
Covalent radius 99 pm
Van der Waals radius 175 pm
Miscellaneous
Magnetic ordering nonmagnetic
Electrical resistivity (20 °C) > 10Ω·m
Thermal conductivity (300 K) 8.9 m W·m−1·K−1
Speed of sound (gas, 0 °C) 206 m/s
CAS registry number 7782-50-5
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 301×103 y β- 0.709 36Ar
ε - 36S
37Cl 24.23% Cl is stable with 20 neutrons

References
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Chlorine (IPA: [ˈklɔːriːn], 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, it is used in swimming pools to keep them clean. In the upper atmosphere, chlorine atoms have been implicated in destruction of the ozone layer.

Contents [hide]
1 Notable characteristics
2 History
3 Occurrence
4 Isotopes
4.1 36Cl
5 Chlorine gas extraction
5.1 Mercury cell electrolysis
5.2 Diaphragm cell electrolysis
5.3 Membrane cell electrolysis
5.4 Other methods
6 Industrial production
6.1 Brine
6.2 Cell room
6.3 Cooling and drying
6.4 Compression and liquefaction
6.5 Storage and loading
6.6 Caustic handling, evaporation, storage and loading
6.7 Hydrogen handling
7 Applications and uses
7.1 Purification and disinfection
7.2 Chemistry
7.3 World War I
7.4 Iraq War
7.5 Other uses
8 Compounds
8.1 Oxidation states
9 Safety
10 See also
11 References
12 External links



[edit] Notable characteristics
Chlorine gas is diatomic, with the formula Cl2. It combines readily with nearly all other elements, although it is not as extremely reactive as fluorine. At 10 °C and atmospheric pressure one litre of water dissolves 3.10 litres of gaseous chlorine and at 30 °C only 1.77 litres.[2]

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.


[edit] History
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. Chlorine was given its current name in 1810 by Sir Humphry Davy, who insisted that it was in fact an element.

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. It is alleged that his role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide. After its first use, it was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly gases phosgene and mustard gas.[3]


[edit] Occurrence
In nature, chlorine is found mainly 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 chemical equation

2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH
See also Halide minerals.


[edit] Isotopes
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.


[edit] 36Cl
Trace amounts of radioactive [[Chlorine-3636Cl]] 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.


[edit] 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. There are three industrial methods for the extraction of chlorine by electrolysis.


[edit] Mercury cell electrolysis
Mercury cell electrolysis, also known as the Castner-Kellner process, was the first method used to produce chlorine on an industrial scale. Titanium or graphite anodes are located above a liquid mercury cathode. Slate baffles divide the cell into two chambers, in which the anode is in contact with just one. The baffles do not go all the way to the bottom of the cell, but allow the mercury cathode (but not the electrolyte) to flow beneath them. Sodium chloride solution is placed in the anode chamber and water in the other chamber. When an electrical current is applied, chlorine is released at the anodes and sodium dissolves into the mercury cathode forming an amalgam. By rocking the entire cell, the mercury amalgam is exposed to the water chamber, where it reacts to form sodium hydroxide and hydrogen gas as byproducts.[5][6]

This method consumes vast amounts of energy and there are also concerns about mercury emissions.


[edit] Diaphragm cell electrolysis
In diaphragm cell electrolysis, an asbestos (or other porous material) diaphragm separates cathode and anode, preventing the chlorine forming at the anode and the sodium hydroxide forming at the cathode from re-mixing. 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).[7][8] The cells vary in construction and placement of the diaphragm, with some having the diaphragm in direct contact with the cathode. Each of these uses the same principle of allowing sodium ions to diffuse through the porous diaphragm from the anode side containing the chloride to the cathode side containing the hydroxide. The concentration of sodium ions in the cathode side is kept low by constantly removing some hydroxide solution and replacing it with water. The sodium ion concentration on the anode side is kept high by adding sodium chloride to keep the solution saturated. Sodium ions are driven by the electric current to flow toward the cathode, whereas the chloride ions are driven in the opposite direction. Despite this some diffusion of chloride and hypochlorite ions through the diaphragm is unavoidable. As a result diaphragm methods produce alkali of less purity than do mercury cell methods. But diaphragm cells are not burdened with the cost of mercury and 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.[8]


[edit] Membrane cell electrolysis
The electrolysis cell is divided into two by a cation permeable membrane acting as an ion exchanger. Saturated sodium chloride solution is passed through the anode compartment leaving a lower concentration. Sodium hydroxide solution is circulated through the cathode compartment exiting at a higher concentration. A portion of this concentrated sodium hydroxide solution is diverted as product while the remainder is diluted with deionized water and passed through the electrolyzer again.

This method is nearly as efficient as the diaphragm cell and produces very pure sodium hydroxide but requires very pure sodium chloride solution.

Cathode: 2 H+(aq) + 2 e− → H2 (g)
Anode: 2 Cl− → Cl2 (g) + 2 e−
Overall equation: 2 NaCl + 2H2O → Cl2 + H2 + 2 NaOH


[edit] 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 in 400°C. The amount of extracted chlorine is approximately at 80%. Due to the extremely corrosive reaction mixture, industrial use of this method is difficult.

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.[9]

The Downs process for producing sodium metal electrolytically from fused sodium chloride produces chlorine as a byproduct.

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

Small amounts of chlorine gas can also 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.


[edit] 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 will also produce sodium hydroxide (referred to in the industry as caustic), hydrogen gas and sodium hypochlorite. A typical plant consists of brine production, cell operations, chlorine cooling & drying, chlorine compression & liquefaction, liquid chlorine storage & loading, caustic handling, evaporation, storage & loading and hydrogen handling.


[edit] Brine
Key to the production of chlorine is the operation of the brine saturation 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 treated with sodium hydroxide, sodium carbonate and a polymer flocculant to reduce calcium, magnesium and other impurities. The brine proceeds to a large clarifier where the impurities are allowed to settle out. After overflowing the clarifier the brine is mechanically filtered via a sand bed filter and a leaf pressure filter before entering ion exchangers to further remove impurities. At several points in this process, the brine is tested for hardness and strength.

The brine is then considered pure, and is placed in large 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 for a given electrical load. Brine exiting the cell room must be treated to remove residual chlorine, control pH, and be returned to the pile. 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 these too can damage cells.


[edit] Cell room
A cell room or cell house is a large building that houses the many electrolytic cells. It contains support structures for the cells and connections for supplying electrical power to the cells. It also contains piping for supplying feed brine, feed caustic and the removal of spent brine, exit caustic as well as wet chlorine gas and wet hydrogen gas. Monitoring and control of the temperatures of the feed caustic and brine is done to control exit temperatures. The exit temperatures are usually monitored for each cell and controlled by adjusting the feed temperatures and flow rates. Also monitored are the voltages of each cell which vary with the electrical load on the cell room which 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 a rectifier with the cells connected in series. Plant load is controlled by varying the current to the cell room based on the ratio of amperage to square area of membrane surface in the cell room. As the current is increased flow rates for brine, caustic and deionized water are increased while lowering the feed temperatures.


[edit] 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 to condense and fall out of the gas stream where it can be piped off and placed in the exit brine to recover some of the chlorine and conserve deionized water. Cooling also improves the efficiency of the compression and liquefaction stage that follows. This usually accomplished via a multistage tube and shell set of coolers fed with cooling water on the first stage and chilled water on the second. Chlorine exiting the second stage is ideally between 18º C and 25º C. After cooling the gas stream passes through a series of packed towers with counter flowing sulfuric acid. These towers progressively remove any remaining moisture from the chlorine gas. After exiting the drying towers a filter system to remove any sulfuric acid droplets from the flow is usually employed.


[edit] Compression and liquefaction
Several methods of compression my be used: for small production levels a liquid ring compressor, for medium levels a reciprocating compressor, for large levels typically a centrifugal compressor. 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, which are large freon chillers, where it is cooled enough to liquefy. Non condensible gases and trace chlorine gas are vented off as part of the pressure control of the liquefaction systems. This pressure control helps maintain the correct pressure for liquefaction to occur. These gases are routed to a gas scrubber, producing sodium hypochlorite, or used in the production of hydrochloric acid via an hydrochloric acid burner or oven.


[edit] Storage and loading
Liquid chlorine is typically gravity-fed to storage tanks or directly to chlorine rail cars. Liquid chlorine can be loaded into rail cars via pumps or padded into cars with compressed air.


[edit] 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 with in the cells. The caustic exiting the cell line must be monitored for strength, to maintain safe concentrations. Too strong or too weak of a solution may damage the cells. Membrane cells typically produce caustic in the range of 30% to 33% by weight. The caustic is heated at low electrical loads to control the exit temperature of the caustic. Higher loads require the caustic to be cooled, to maintain correct exit temperatures. A tube and shell heat exchanger is typically used in a dual mode configuration. The caustic exiting to storage is pulled from a storage tank and may be diluted to 25% solution for sale for use on site. Another stream may be pumped into a multiple effect evaporator set to produce 50% caustic, improving storage and allowing for economical transportation. The 50% and 25% caustic are stored in separate tanks and rail cars and tanker trucks are loaded at loading stations via pumps.


[edit] 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. Some possible uses are hydrochloric acid or hydrogen peroxide production and use as a fuel in boilers or fuel cells.


[edit] Applications and uses

[edit] Purification and disinfection
Chlorine is an important chemical for some processes of water purification, in disinfectants, and in bleach.

Chlorine is also used widely in the manufacture of many every-day items, or to purify water in various forms.

Used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and swimming pools. However, in most non-commercial swimming pools chlorine itself is not used, but rather sodium hypochlorite (household bleach), a compound of chlorine with sodium and oxygen. Calcium hypochlorite is also used as a cheaper alternative. Even small water supplies are now routinely chlorinated.[10] (See also chlorination)
Used widely in paper product production, antiseptic, dyestuffs, food, insecticides, paints, petroleum products, plastics, medicines, textiles, solvents, and many other consumer products.

[edit] Chemistry
Chlorine gas 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 organic compounds. This reaction usually results in a mixture of products based on a statistical distribution. It is often difficult to control the degree of substitution as well — multiple substitutions are common. Unless the different reaction products are easily separated, e.g. by distillation, free-radical chlorination is not a preferred synthetic route.

Like the other halides, chlorine undergoes electrophilic additions reactions, most notably, the halogenation of alkenes, and the halogenation of aromatic compounds with a Lewis acid catalyst. Organic chloride compounds tend to be less reactive than the corresponding bromides or iodides, 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. Excess chlorine is removed from water with sulfur dioxide.


[edit] World War I
Main article: Use of poison gas in World War I
Chlorine became the first killing agent to be employed during World War I. German chemical conglomerate IG Farben had been producing chlorine as a by-product of their dye manufacturing. In cooperation with Fritz Haber of the Kaiser Wilhelm Institute for Chemistry in Berlin, they developed methods of discharging chlorine gas against an entrenched enemy.


[edit] Iraq War
Main article: 2007 chlorine bombings in Iraq
Chlorine gas has been used by insurgents in the Iraq War as a chemical weapon to increase the capability to threaten the local population and coalition forces in their attacks. On March 17, 2007, for example, three chlorine filled trucks were detonated in the Anbar province killing 2 and sickening over 350.[11] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.[12] While fatalities are common, chlorine bombings do not inflict a massive loss of life, and are primarily intended to create widespread panic.


[edit] Other uses
It is also used in the production of chlorates, chloroform, carbon tetrachloride, and in bromine extraction.


[edit] 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).

See also:

Fluorides: chlorine monofluoride (ClF), chlorine trifluoride (ClF3), chlorine pentafluoride (ClF5)
Oxides: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine heptoxide (Cl2O7)
Acids: hydrochloric acid (HCl), chloric acid (HClO3), and perchloric acid (HClO4)
See also Chlorine compounds.


[edit] Oxidation states
Oxidation
state 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, organic perchlorates, ammonium perchlorate, magnesium 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−
Potassium chlorate can be crystalized from solutions formed by the above reactions. If its crystals are heated, it undergoes the final disproportionation step.

4ClO3− → Cl− + 3ClO4−
This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[13]

Reaction Electrode
potential
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

[edit] Safety
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. For more information see an MSDS.


[edit] See also
[hide] v • d • e Diatomic Elements
Hydrogen
H2
|

Nitrogen
N2
|

Oxygen
O2
|

Fluorine
F2

Chlorine
Cl2
|

Bromine
Br2
|

Iodine
I2
|

Astatine
At2


Chloride

[edit] References
^ Merck Index of Chemicals and Drugs, 9th ed., monograph 2065
^ WebElements.com – Chlorine. Mark Winter [The University of Sheffield and WebElements Ltd, UK]. Retrieved on 2007-03-17.
^ Weapons of War: Poison Gas. First World War.com. Retrieved on 2007-08-12.
^ Risk assessment and the cycling of natural organochlorines. Euro Chlor. Retrieved on 2007-08-12.
^ Pauling, Linus, General Chemistry, 1970 ed., Dover publications
^ 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.
^ The Electrolysis of Brine. Salt Manufacturers' Association. Retrieved on 2007-03-17.
^ a b Kiefer, David M.. When the Industry Charged Ahead. Chemistry Chronicles. Retrieved on 2007-03-17.
^ The Chlorine Industry. Lenntech Water treatment & air purification Holding B.V., Rotterdamseweg 402 M, 2629 HH Delft, The Netherlands. Retrieved on 2007-03-17.
^ Chlorine. Los Alamos National Laboratory. Retrieved on 2007-03-17.
^ Mahdi, Basim. "Iraq gas attack makes hundreds ill", CNN, 2007-03-17. Retrieved on 2007-03-17.
^ "'Chlorine bomb' hits Iraq village", BBC News, 2007-05-17. Retrieved on 2007-05-17.
^ Cotton, F. Albert and Wilkinson, Geoffrey, Advanced Inorganic Chemistry 2nd ed. John Wiley & sons, p568

[edit] External links
Wikimedia Commons has media related to:
ChlorineLook up chlorine in
Wiktionary, the free dictionary.Chlorine Institute - Trade association and lobby group representing the interests of the chlorine industry
Chlorine Online - Chlorine Online is an information resource produced by Eurochlor - the business association of the European chlor-alkali industry
Computational Chemistry Wiki
Chlorine Production Using Mercury, Environmental Considerations and Alternatives
National Pollutant Inventory - Chlorine



This article forms part of the series

Chemical warfare
Blood agents: Cyanogen chloride (CK) – Hydrogen cyanide (AC)
Blister agents: Lewisite (L) – Sulfur mustard gas (HD, H, HT, HL, HQ) – Nitrogen mustard gas (HN1, HN2, HN3)
Nerve agents: G-Agents: Tabun (GA) – Sarin (GB) – Soman (GD) – Cyclosarin (GF) – GV | V-Agents: VE – VG – VM – VX | Novichok agents
Pulmonary agents: Chlorine – Chloropicrin (PS) – Phosgene (CG) – Diphosgene (DP)
Incapacitating agents: Agent 15 (BZ) – KOLOKOL-1
Riot control agents: Pepper spray (OC) – CS gas – CN gas (mace) – CR gas v • d • e

v • d • eE numbers[show]
Colours (E100–199) • Preservatives (E200–299) • Antioxidants & Acidity regulators (E300–399) • Thickeners, stabilisers & emulsifiers (E400–499) • pH regulators & anti-caking agents (E500–599) • Flavour enhancers (E600–699) • Miscellaneous (E900–999) • Additional chemicals (E1100–1599)


--------------------------------------------------------------------------------

Waxes (E900–909) • Synthetic glazes (E910–919) • Improving agents (E920–929) • Packaging gases (E930–949) • Sweeteners (E950–969) • Foaming agents (E990–999)


--------------------------------------------------------------------------------

L-cysteine (E920) • L-cystine (E921) • Potassium persulfate (E922) • Ammonium persulfate (E923) • Potassium bromate (E924) • Chlorine (E925) • Chlorine dioxide (E926) • Azodicarbonamide (E927) • Carbamide (E927b) • Benzoyl peroxide (E928)


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