Along with fluorine, bromine, iodine and astatine, chlorine is a member of the halogen series that forms the group 17 of the periodic table - 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. 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. 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.
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.
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.
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:
Chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and therefore he is credited for its discovery. 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"). However, he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory). He named the new element within this oxide as muriaticum. Regardless of what he thought, Scheele did isolate chlorine by reacting MnO2 with HCl:
Claude Berthollet suggested that Scheele's dephlogisticated muratic acid air must be a combination of oxygen and an undiscovered element, muriaticum.
In 1809 Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muratic acid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide). They did not succeed and published a report in which they considered the possibility that dephlogisticated muratic acid air is an element, but were not convinced.
In 1810, Sir Humphrey Davy tried the same experiment again, and concluded that it was an element, and not a compound. He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow. 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. In 1822, Michael Faraday liquefied chlorine for the first time.
Chlorine was first used to bleach textiles in 1785. In 1826, silver chloride was used to produce photographic images for the first time. Chloroform was first used as an anesthetic in 1847. Chlorine 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, 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. Polyvinylchloride (PVC) was invented in 1912, initially without a purpose. Chlorine gas was first introduced as a weapon on April 22, 1915 at Ypres by the German Army. and the results of this weapon were disastrous because gas masks had not yet been invented.
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.
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%.
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).
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. This demand was met by capturing hydrochloric acid driven off as a gas during the production of alkali by the Leblanc process, oxidising 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.
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.
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.
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.
Other chlorine-containing compounds include:
|−1||chlorides||Cl−||ionic chlorides, organic chlorides, hydrochloric acid|
|+1||hypochlorites||ClO−||sodium hypochlorite, calcium hypochlorite|
|+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:
In hot concentrated alkali solution disproportionation continues:
This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:
|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|
Chlorine is usually 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. (See also chlorination)
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.
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 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.
Chlorine gas has also been used 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. Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions. 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.
The element is widely used for purifying water owing to its powerful oxidising 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. Some polymers, however, are 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 90's. 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 moulding 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.
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.
Chlorine is detectable in concentrations of as low as 1 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. Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.