Chlorine

Chlorine is a chemical element; it has symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the revised Pauling scale, behind only oxygen and fluorine.
Chlorine played an important role in the experiments conducted by medieval alchemists, which commonly involved the heating of chloride salts like ammonium chloride and sodium chloride, producing various chemical substances containing chlorine such as hydrogen chloride, mercury chloride, and aqua regia. However, the nature of free chlorine gas as a separate substance was only recognised around 1630 by Jan Baptist van Helmont. Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it after the Ancient Greek χλωρός because of its colour.
Because of its great reactivity, all chlorine in the Earth's crust is in the form of ionic chloride compounds, which includes table salt. It is the second-most abundant halogen and 20th most abundant element in Earth's crust. These crystal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater.
Elemental chlorine is commercially produced from brine by electrolysis, predominantly in the chloralkali process. The high oxidising potential of elemental chlorine led to the development of commercial bleaches and disinfectants, and a reagent for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, many intermediates for the production of plastics, and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them sanitary. Elemental chlorine at high concentration is extremely dangerous, and poisonous to most living organisms. As a chemical warfare agent, chlorine was first used in World War I as a poison gas weapon.
In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride ions in neutrophils as part of an immune system response against bacteria.

History

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.

Early discoveries

Around 900, the authors of the Arabic writings attributed to Jabir ibn Hayyan and the Persian physician and alchemist Abu Bakr al-Razi were experimenting with sal ammoniac, which when it was distilled together with vitriol produced hydrogen chloride. However, it appears that in these early experiments with chloride salts, the gaseous products were discarded, and hydrogen chloride may have been produced many times before it was discovered that it can be put to chemical use. One of the first such uses was the synthesis of mercury chloride, whose production from the heating of mercury either with alum and ammonium chloride or with vitriol and sodium chloride was first described in the De aluminibus et salibus. Another important development was the discovery by pseudo-Geber that by adding ammonium chloride to nitric acid, a strong solvent capable of dissolving gold could be produced. Although aqua regia is an unstable mixture that continually gives off fumes containing free chlorine gas, this chlorine gas appears to have been ignored until c. 1630, when its nature as a separate gaseous substance was recognised by the Brabantian chemist and physician Jan Baptist van Helmont.

Isolation

The element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, and he is credited with the discovery. Scheele produced chlorine by reacting MnO₂ with HCl:
Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow-green colour, and the smell similar to aqua regia. He called it "dephlogisticated muriatic acid air" since it is a gas and it came from hydrochloric acid. He failed to establish chlorine as an element.
Common chemical theory at that time held that an acid is a compound that contains oxygen, so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum.
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. 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.
In 1810, Sir Humphry Davy tried the same experiment again, and concluded that the substance was an element, and not a compound. He announced his results to the Royal Society on 15 November that year. At that time, he named this new element "chlorine", from the Greek word χλωρος, in reference to its colour. The name "halogen", meaning "salt producer", was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger. This term was later used as a generic term to describe all the elements in the chlorine family, after a suggestion by Jöns Jakob Berzelius in 1826. In 1823, Michael Faraday liquefied chlorine for the first time, and demonstrated that what was then known as "solid chlorine" had a structure of chlorine hydrate.

Later uses

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785. Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel, by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "Eau de Javel", was a weak solution of sodium hypochlorite. This process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of calcium hypochlorite, then solid calcium hypochlorite. These compounds produced low levels of elemental chlorine and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidiser. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite. This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide. In 1884 Chemischen Fabrik Griesheim of Germany developed another chloralkali process which entered commercial production in 1888.
Elemental chlorine solutions dissolved in chemically basic water were first used as anti-putrefaction agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This practice was pioneered by Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations. Elemental chlorine has since served a continuous function in topical antisepsis and public sanitation, particularly in swimming and drinking water.
Chlorine gas was first used as a weapon on 22 April 1915, at the Second Battle of Ypres by the German Army. The effect on the allies was devastating because the existing gas masks were difficult to deploy and had not been broadly distributed.

Properties

Chlorine is the second halogen, being a nonmetal in group 17 of the periodic table. Its properties are thus similar to fluorine, bromine, and iodine, and are largely intermediate between those of the first two. Chlorine has the electron configuration 3s²3p⁵, with the seven electrons in the third and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between fluorine and bromine, and is less reactive than fluorine and more reactive than bromine. It is also a weaker oxidising agent than fluorine, but a stronger one than bromine. Conversely, the chloride ion is a weaker reducing agent than bromide, but a stronger one than fluoride. It is intermediate in atomic radius between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X₂ molecule, ionic radius, and X–X bond length.
All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: chlorine melts at −101.0 °C and boils at −34.0 °C. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low thanks to their diatomic molecular structure. The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow-green. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Specifically, the colour of a halogen, such as chlorine, results from the electron transition between the highest occupied antibonding π_g molecular orbital and the lowest vacant antibonding σ_u molecular orbital. The colour fades at low temperatures, so that solid chlorine at −195 °C is almost colourless.
Like solid bromine and iodine, solid chlorine crystallises in the orthorhombic crystal system, in a layered lattice of Cl₂ molecules. The Cl–Cl distance is 198 pm and the Cl···Cl distance between molecules is 332 pm within a layer and 382 pm between layers. This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable.