Chlorine (, from the Greek word 'χλωρóς' (khlôros) meaning '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 (Merck Index of Chemicals and Drugs, 9th ed., monograph 2065) 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.

Characteristics

 Chlorine gas is diatomic, with the formula Cl2. It combines readily with all elements except O2 and N2 (

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.

Isotopes

• ³⁶Cl

Occurrence

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]Create? + 2 H₂O → Cl₂ + H₂ + 2 [NaOH]Create?

Production

Gas extraction

:Cathode: 2 H⁺ (aq) + 2 e⁻ → H₂ (g) :Anode: 2 Cl⁻ (aq) → Cl₂ (g) + 2 e⁻ Overall process: 2 [NaCl]Create? (or KCl) + 2 H₂O → Cl₂ + H₂ + 2 [NaOH]Create? (or KOH)

• Mercury cell electrolysis

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. ( (#6) )

• Diaphragm cell electrolysis

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 , but large amounts of steam are required if the caustic has to be evaporated to the commercial concentration of 50%.

• Membrane cell electrolysis

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

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

:4 HCl + O₂ → 2 Cl₂ + 2 H₂O

This reaction is accomplished with the use of copper(II) chloride ([CuCl]Create?₂) 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). (J. Catal. 255, 29 (2008))

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

:2 [NaCl]Create? + 2H₂SO₄ + MnO₂ → Na₂SO₄ + [MnSO]Create?₄ + 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. ( (#11) )

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.

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 production

• Brine

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 bothchlorate 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

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

• Compression and liquefaction

• Storage and loading

• Caustic handling, evaporation, storage and loading

• Hydrogen handling

• Energy consumption

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.

Compounds

Other chlorine-containing compounds include:

• 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)

Oxidation states

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: (Cotton, F. Albert and Wilkinson, Geoffrey, Advanced Inorganic Chemistry 2nd ed. John Wiley & sons, p568)

:{| class="wikitable" align="left" ! Reaction !! Electrode
potential

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

Applications and uses

Production of industrial and consumer products

Purification and disinfection

Chemistry

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 weapon

• World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 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 which makes the deaths by chlorine gas much lower then 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. ( (#20) )

• 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. ( (#21) ) Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions. ( (#22) ) 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.

Other uses

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

History

:4 HCl + MnO₂ → [MnCl]Create?₂ + 2 H₂O + Cl₂

Scheele observed several of the properties of chlorine. The bleaching effect on litmus and the deadly effect on insects additional to the yellow green colour and the smell similar to aqua regina. Chlorine was given its current name in 1810 by Sir Humphry Davy, who insisted that it was in fact an element.

Safety

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.

The number of people allergic to chlorine is very small. People who are allergic to chlorine cannot drink tap water, bathe in tap water or swim in pools. Dechlorinating bath salts are used to neutralize the chlorine in bath water. Otherwise, fresh water is boiled and cooled.

Chlorine cracking

See also

• Chloride
• Polymer degradation

References

External links

• 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
• National Institute for Occupational Safety and Health - Chlorine Page

Category:Chemical elements Category:Halogens

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