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=                              Chlorine                              =
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                            Introduction
======================================================================
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 (sal ammoniac) and sodium chloride
(common salt), producing various chemical substances containing
chlorine such as hydrogen chloride, mercury(II) chloride (corrosive
sublimate), and . 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  (,
"pale green") 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 (after fluorine) 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 (PVC), 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 (Latin: Geber) and the Persian physician and alchemist Abu Bakr
al-Razi ( 865-925, Latin: Rhazes) were experimenting with sal ammoniac
(ammonium chloride), which when it was distilled together with vitriol
(hydrated sulfates of various metals) 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(II) chloride (corrosive sublimate), 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' ("On Alums and Salts", an eleventh- or twelfth century
Arabic text falsely attributed to Abu Bakr al-Razi and translated into
Latin in the second half of the twelfth century by Gerard of Cremona,
1144-1187). Another important development was the discovery by
pseudo-Geber (in the 'De inventione veritatis', "On the Discovery of
Truth", after c. 1300) that by adding ammonium chloride to nitric
acid, a strong solvent capable of dissolving gold (i.e., 'aqua regia')
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. From
'"Complexionum atque mistionum elementalium figmentum."' (Formation of
combinations and of mixtures of elements), §37,
[https://books.google.com/books?id=Qy5AAAAAcAAJ&pg=PA105 p. 105:]
'"Accipe salis petrae, vitrioli, & alumnis partes aequas:
exsiccato singula, & connexis simul, distilla aquam. Quae nil
aliud est, quam merum sal volatile. Hujus accipe uncias quatuor, salis
armeniaci unciam junge, in forti vitro, alembico, per caementum (ex
cera, colophonia, & vitri pulverre) calidissime affusum, firmato;
mox, etiam in frigore, Gas excitatur, & vas, utut forte, dissilit
cum fragore."' (Take equal parts of saltpeter [i.e., sodium nitrate],
vitriol [i.e., concentrated sulfuric acid], and alum: dry each and
combine simultaneously; distill off the water [i.e., liquid]. That
[distillate] is nothing else than pure volatile salt [i.e., spirit of
nitre, nitric acid]. Take four ounces of this [viz, nitric acid], add
one ounce of Armenian salt [i.e., ammonium chloride], [place it] in a
strong glass alembic sealed by cement ([made] from wax, rosin, and
powdered glass) [that has been] poured very hot; soon, even in the
cold, gas is stimulated, and the vessel, however strong, bursts into
fragments.) From '"De Flatibus"' (On gases),
[https://books.google.com/books?id=Qy5AAAAAcAAJ&pg=PA408 p. 408] :
'"Sal armeniacus enim, & aqua chrysulca, quae singula per se
distillari, possunt, & pati calorem: sin autem jungantur, &
intepescant, non possunt non, quin statim in Gas sylvestre, sive
incoercibilem flatum transmutentur."' (Truly Armenian salt [i.e.,
ammonium chloride] and nitric acid, each of which can be distilled by
itself, and submitted to heat; but if, on the other hand, they be
combined and become warm, they cannot but be changed immediately into
carbon dioxide [note: van Helmont's identification of the gas is
mistaken] or an incondensable gas.)
See also:
*
[https://www.encyclopedia.com/people/science-and-technology/chemistry-biographies/johannes-joan-baptista-van-helmont
Helmont, Johannes (Joan) Baptista Van, Encyclopedia.Com] : "Others
were chlorine gas from the reaction of nitric acid and sal ammoniac;
... "
* Wisniak, Jaime (2009) "Carl Wilhelm Scheele," 'Revista CENIC
Ciencias Químicas', 40 (3): 165-73; see p. 168: "Early in the
seventeenth century Johannes Baptiste van Helmont (1579-1644)
mentioned that when sal marin (sodium chloride) or sal ammoniacus and
aqua chrysulca (nitric acid) were mixed together, a flatus incoercible
(non-condensable gas) was evolved."


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 (as the mineral pyrolusite) with
HCl:

:4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2

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 (then called
"airs") and it came from hydrochloric acid (then known as "muriatic
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 (remnants of this survive in the German and Dutch
names of oxygen:  or ', both translating into English as 'acid
substance'), 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' (and carbon
dioxide). 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. See: §
'De la nature et des propriétés de l'acide muriatique et de l'acide
muriatique oxigéné' (On the nature and properties of muriatic acid and
of oxidized muriatic acid), pp. 339-58. From pp. 357-58: '"Le gaz
muriatique oxigéné n'est pas, en effect, décomposé ... comme un corps
composé."' ("In fact, oxygenated muriatic acid is not decomposed by
charcoal, and it might be supposed, from this fact and those that are
communicated in this Memoir, that this gas is a simple body. The
phenomena that it presents can be explained well enough on this
hypothesis; we shall not seek to defend it, however, as it appears to
us that they are still better explained by regarding oxygenated
muriatic acid as a compound body.") For a full English translation of
this section, see: [http://web.lemoyne.edu/~giunta/thenard.html Joseph
Louis Gay-Lussac and Louis Jacques Thénard, "On the nature and the
properties of muriatic acid and of oxygenated muriatic acid" (Lemoyne
College, Syracuse, New York)]

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 ('chlōros', "green-yellow"), 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 (fluorine, bromine, iodine), 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 (Cl2·H2O).


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 (now part of Paris, France), by
passing chlorine gas through a solution of sodium carbonate. The
resulting liquid, known as "" ("Javel water"), 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 ("chlorinated lime"), then solid calcium hypochlorite
(bleaching powder). 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
oxidizer. 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
(sodium and calcium hypochlorite) 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
(wound irrigation solutions and the like) and public sanitation,
particularly in swimming and drinking water.

Chlorine gas was first used as a weapon on April 22, 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 [Ne]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 (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), 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 (X = Cl, Br, I), ionic
radius, and X-X bond length. (Fluorine is anomalous due to its small
size.)

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 (leading to high volatility) 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 'πg'
molecular orbital and the lowest vacant antibonding 'σu' 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 (close to the gaseous Cl-Cl distance of
199 pm) and the Cl···Cl distance between molecules is 332 pm within a
layer and 382 pm between layers (compare the van der Waals radius of
chlorine, 180 pm). This structure means that chlorine is a very poor
conductor of electricity, and indeed its conductivity is so low as to
be practically unmeasurable.


Isotopes
==========
Chlorine has two stable isotopes, 35Cl and 37Cl. These are its only
two natural isotopes occurring in quantity, with 35Cl making up 76% of
natural chlorine and 37Cl making up the remaining 24%. Both are
synthesised in stars in the oxygen-burning and silicon-burning
processes. Both have nuclear spin 3/2+ and thus may be used for
nuclear magnetic resonance, although the spin magnitude being greater
than 1/2 results in non-spherical nuclear charge distribution and thus
resonance broadening as a result of a nonzero nuclear quadrupole
moment and resultant quadrupolar relaxation. The other chlorine
isotopes are all radioactive, with half-lives too short to occur in
nature primordially. Of these, the most commonly used in the
laboratory are 36Cl ('t'1/2 = 3.0×105 y) and 38Cl ('t'1/2 = 37.2 min),
which may be produced from the neutron activation of natural chlorine.

The most stable chlorine radioisotope is 36Cl. The primary decay mode
of isotopes lighter than 35Cl is electron capture to isotopes of
sulfur; that of isotopes heavier than 37Cl is beta decay to isotopes
of argon; and 36Cl may decay by either mode to stable 36S or 36Ar.
36Cl occurs in trace quantities in nature as a cosmogenic nuclide in a
ratio of about (7-10) × 10−13 to 1 with stable chlorine isotopes: it
is produced in the atmosphere by spallation of 36Ar by interactions
with cosmic ray protons. In the top meter of the lithosphere, 36Cl is
generated primarily by thermal neutron activation of 35Cl and
spallation of 39K and 40Ca. In the subsurface environment, muon
capture by 40Ca becomes more important as a way to generate 36Cl.


                      Chemistry and compounds
======================================================================
Halogen bond energies (kJ/mol)
X       XX      HX      BX3     AlX3    CX4
F       159     574     645     582     456
Cl      |243    |428    |444    |427    |327
Br      |193    |363    |368    |360    |272
I       |151    |294    |272    |285    |239
Chlorine is intermediate in reactivity between fluorine and bromine,
and is one of the most reactive elements. Chlorine is a weaker
oxidising agent than fluorine but a stronger one than bromine or
iodine. This can be seen from the standard electrode potentials of the
X2/X− couples (F, +2.866  V; Cl, +1.395 V; Br, +1.087  V; I, +0.615 V;
At, approximately +0.3  V). However, this trend is not shown in the
bond energies because fluorine is singular due to its small size, low
polarisability, and inability to show hypervalence. As another
difference, chlorine has a significant chemistry in positive oxidation
states while fluorine does not. Chlorination often leads to higher
oxidation states than bromination or iodination but lower oxidation
states than fluorination. Chlorine tends to react with compounds
including M-M, M-H, or M-C bonds to form M-Cl bonds.

Given that E°(O2/H2O) = +1.229 V, which is less than +1.395 V, it
would be expected that chlorine should be able to oxidise water to
oxygen and hydrochloric acid. However, the kinetics of this reaction
are unfavorable, and there is also a bubble overpotential effect to
consider, so that electrolysis of aqueous chloride solutions evolves
chlorine gas and not oxygen gas, a fact that is very useful for the
industrial production of chlorine.


Hydrogen chloride
===================
The simplest chlorine compound is hydrogen chloride, HCl, a major
chemical in industry as well as in the laboratory, both as a gas and
dissolved in water as hydrochloric acid. It is often produced by
burning hydrogen gas in chlorine gas, or as a byproduct of
chlorinating hydrocarbons. Another approach is to treat sodium
chloride with concentrated sulfuric acid to produce hydrochloric acid,
also known as the "salt-cake" process:

:NaCl + H2SO4  NaHSO4 + HCl
:NaCl + NaHSO4  Na2SO4 + HCl

In the laboratory, hydrogen chloride gas may be made by drying the
acid with concentrated sulfuric acid. Deuterium chloride, DCl, may be
produced by reacting benzoyl chloride with heavy water (D2O).

At room temperature, hydrogen chloride is a colourless gas, like all
the hydrogen halides apart from hydrogen fluoride, since hydrogen
cannot form strong hydrogen bonds to the larger electronegative
chlorine atom; however, weak hydrogen bonding is present in solid
crystalline hydrogen chloride at low temperatures, similar to the
hydrogen fluoride structure, before disorder begins to prevail as the
temperature is raised. Hydrochloric acid is a strong acid (p'K'a = −7)
because the hydrogen-chlorine bonds are too weak to inhibit
dissociation. The HCl/H2O system has many hydrates HCl·'n'H2O for 'n'
= 1, 2, 3, 4, and 6. Beyond a 1:1 mixture of HCl and H2O, the system
separates completely into two separate liquid phases. Hydrochloric
acid forms an azeotrope with boiling point 108.58 °C at 20.22 g HCl
per 100 g solution; thus hydrochloric acid cannot be concentrated
beyond this point by distillation.

Unlike hydrogen fluoride, anhydrous liquid hydrogen chloride is
difficult to work with as a solvent, because its boiling point is low,
it has a small liquid range, its dielectric constant is low and it
does not dissociate appreciably into H2Cl+ and  ions - the latter, in
any case, are much less stable than the bifluoride ions () due to the
very weak hydrogen bonding between hydrogen and chlorine, though its
salts with very large and weakly polarising cations such as Cs+ and
quaternary ammonium cation (R = Me, Et, Bu'n') may still be isolated.
Anhydrous hydrogen chloride is a poor solvent, only able to dissolve
small molecular compounds such as nitrosyl chloride and phenol, or
salts with very low lattice energies such as tetraalkylammonium
halides. It readily protonates nucleophiles containing lone-pairs or π
bonds. Solvolysis, ligand replacement reactions, and oxidations are
well-characterised in hydrogen chloride solution:
:Ph3SnCl + HCl ⟶ Ph2SnCl2 + PhH (solvolysis)
:Ph3COH + 3 HCl ⟶  + H3O+Cl− (solvolysis)
: + BCl3 ⟶  + HCl (ligand replacement)
:PCl3 + Cl2 + HCl ⟶  (oxidation)


Other binary chlorides
========================
Nearly all elements in the periodic table form binary chlorides. The
exceptions are decidedly in the minority and stem in each case from
one of three causes: extreme inertness and reluctance to participate
in chemical reactions (the noble gases, with the exception of xenon in
the highly unstable XeCl2 and XeCl4); extreme nuclear instability
hampering chemical investigation before decay and transmutation (many
of the heaviest elements beyond bismuth); and having an
electronegativity higher than chlorine's (oxygen and fluorine) so that
the resultant binary compounds are formally not chlorides but rather
oxides or fluorides of chlorine. Even though nitrogen in NCl3 is
bearing a negative charge, the compound is usually called nitrogen
trichloride.

Chlorination of metals with Cl2 usually leads to a higher oxidation
state than bromination with Br2 when multiple oxidation states are
available, such as in MoCl5 and MoBr3. Chlorides can be made by
reaction of an element or its oxide, hydroxide, or carbonate with
hydrochloric acid, and then dehydrated by mildly high temperatures
combined with either low pressure or anhydrous hydrogen chloride gas.
These methods work best when the chloride product is stable to
hydrolysis; otherwise, the possibilities include high-temperature
oxidative chlorination of the element with chlorine or hydrogen
chloride, high-temperature chlorination of a metal oxide or other
halide by chlorine, a volatile metal chloride, carbon tetrachloride,
or an organic chloride. For instance, zirconium dioxide reacts with
chlorine at standard conditions to produce zirconium tetrachloride,
and uranium trioxide reacts with hexachloropropene when heated under
reflux to give uranium tetrachloride. The second example also involves
a reduction in oxidation state, which can also be achieved by reducing
a higher chloride using hydrogen or a metal as a reducing agent. This
may also be achieved by thermal decomposition or disproportionation as
follows:
: EuCl3 +  H2 ⟶ EuCl2 + HCl
: ReCl5  ReCl3 + Cl2
: AuCl3  AuCl + Cl2

Most metal chlorides with the metal in low oxidation states (+1 to +3)
are ionic. Nonmetals tend to form covalent molecular chlorides, as do
metals in high oxidation states from +3 and above. Both ionic and
covalent chlorides are known for metals in oxidation state +3 (e.g.
scandium chloride is mostly ionic, but aluminium chloride is not).
Silver chloride is very insoluble in water and is thus often used as a
qualitative test for chlorine.


Polychlorine compounds
========================
Although dichlorine is a strong oxidising agent with a high first
ionisation energy, it may be oxidised under extreme conditions to form
the  cation. This is very unstable and has only been characterised by
its electronic band spectrum when produced in a low-pressure discharge
tube. The yellow  cation is more stable and may be produced as
follows:
:
This reaction is conducted in the oxidising solvent arsenic
pentafluoride. The trichloride anion, , has also been characterised;
it is analogous to triiodide.


Chlorine fluorides
====================
The three fluorides of chlorine form a subset of the interhalogen
compounds, all of which are diamagnetic. Some cationic and anionic
derivatives are known, such as , , , and Cl2F+. Some pseudohalides of
chlorine are also known, such as cyanogen chloride (ClCN, linear),
chlorine cyanate (ClNCO), chlorine thiocyanate (ClSCN, unlike its
oxygen counterpart), and chlorine azide (ClN3).

Chlorine monofluoride (ClF) is extremely thermally stable, and is sold
commercially in 500-gram steel lecture bottles. It is a colourless gas
that melts at −155.6 °C and boils at −100.1 °C. It may be produced by
the reaction of its elements at 225 °C, though it must then be
separated and purified from chlorine trifluoride and its reactants.
Its properties are mostly intermediate between those of chlorine and
fluorine. It will react with many metals and nonmetals from room
temperature and above, fluorinating them and liberating chlorine. It
will also act as a chlorofluorinating agent, adding chlorine and
fluorine across a multiple bond or by oxidation: for example, it will
attack carbon monoxide to form carbonyl chlorofluoride, COFCl. It will
react analogously with hexafluoroacetone, (CF3)2CO, with a potassium
fluoride catalyst to produce heptafluoroisopropyl hypochlorite,
(CF3)2CFOCl; with nitriles RCN to produce RCF2NCl2; and with the
sulfur oxides SO2 and SO3 to produce ClSO2F and ClOSO2F respectively.
It will also react exothermically with compounds containing -OH and
-NH groups, such as water:
:H2O + 2 ClF ⟶ 2 HF + Cl2O

Chlorine trifluoride (ClF3) is a volatile colourless molecular liquid
which melts at −76.3 °C and boils at 11.8  °C. It may be formed by
directly fluorinating gaseous chlorine or chlorine monofluoride at
200-300 °C. One of the most reactive chemical compounds known, the
list of elements it sets on fire is diverse, containing hydrogen,
potassium, phosphorus, arsenic, antimony, sulfur, selenium, tellurium,
bromine, iodine, and powdered molybdenum, tungsten, rhodium, iridium,
and iron.  It will also ignite water, along with many substances which
in ordinary circumstances would be considered chemically inert such as
asbestos, concrete, glass, and sand. When heated, it will even corrode
noble metals as palladium, platinum, and gold, and even the noble
gases xenon and radon do not escape fluorination. An impermeable
fluoride layer is formed by sodium, magnesium, aluminium, zinc, tin,
and silver, which may be removed by heating. Nickel, copper, and steel
containers are usually used due to their great resistance to attack by
chlorine trifluoride, stemming from the formation of an unreactive
layer of metal fluoride. Its reaction with hydrazine to form hydrogen
fluoride, nitrogen, and chlorine gases was used in experimental rocket
engine, but has problems largely stemming from its extreme
hypergolicity resulting in ignition without any measurable delay.
Today, it is mostly used in nuclear fuel processing, to oxidise
uranium to uranium hexafluoride for its enriching and to separate it
from plutonium, as well as in the semiconductor industry, where it is
used to clean chemical vapor deposition chambers. It can act as a
fluoride ion donor or acceptor (Lewis base or acid), although it does
not dissociate appreciably into  and  ions.

Chlorine pentafluoride (ClF5) is made on a large scale by direct
fluorination of chlorine with excess fluorine gas at 350 °C and 250
atm, and on a small scale by reacting metal chlorides with fluorine
gas at 100-300 °C. It melts at −103 °C and boils at −13.1 °C. It is a
very strong fluorinating agent, although it is still not as effective
as chlorine trifluoride. Only a few specific stoichiometric reactions
have been characterised. Arsenic pentafluoride and antimony
pentafluoride form ionic adducts of the form [ClF4]+[MF6]− (M = As,
Sb) and water reacts vigorously as follows:
:2 H2O + ClF5 ⟶ 4 HF + FClO2

The product, chloryl fluoride, is one of the five known chlorine oxide
fluorides. These range from the thermally unstable FClO to the
chemically unreactive perchloryl fluoride (FClO3), the other three
being FClO2, F3ClO, and F3ClO2. All five behave similarly to the
chlorine fluorides, both structurally and chemically, and may act as
Lewis acids or bases by gaining or losing fluoride ions respectively
or as very strong oxidising and fluorinating agents.


Chlorine oxides
=================
The chlorine oxides are well-studied in spite of their instability
(all of them are endothermic compounds). They are important because
they are produced when chlorofluorocarbons undergo photolysis in the
upper atmosphere and cause the destruction of the ozone layer. None of
them can be made from directly reacting the elements.

Dichlorine monoxide (Cl2O) is a brownish-yellow gas (red-brown when
solid or liquid) which may be obtained by reacting chlorine gas with
yellow mercury(II) oxide. It is very soluble in water, in which it is
in equilibrium with hypochlorous acid (HOCl), of which it is the
anhydride. It is thus an effective bleach and is mostly used to make
hypochlorites. It explodes on heating or sparking or in the presence
of ammonia gas.

Chlorine dioxide (ClO2) was the first chlorine oxide to be discovered
in 1811 by Humphry Davy. It is a yellow paramagnetic gas (deep-red as
a solid or liquid), as expected from its having an odd number of
electrons: it is stable towards dimerisation due to the delocalisation
of the unpaired electron. It explodes above −40 °C as a liquid and
under pressure as a gas and therefore must be made at low
concentrations for wood-pulp bleaching and water treatment. It is
usually prepared by reducing a chlorate as follows:
: + Cl− + 2 H+ ⟶ ClO2 +  Cl2 + H2O
Its production is thus intimately linked to the redox reactions of the
chlorine oxoacids. It is a strong oxidising agent, reacting with
sulfur, phosphorus, phosphorus halides, and potassium borohydride. It
dissolves exothermically in water to form dark-green solutions that
very slowly decompose in the dark. Crystalline clathrate hydrates
ClO2·'n'H2O ('n' ≈ 6-10) separate out at low temperatures. However, in
the presence of light, these solutions rapidly photodecompose to form
a mixture of chloric and hydrochloric acids. Photolysis of individual
ClO2 molecules result in the radicals ClO and ClOO, while at room
temperature mostly chlorine, oxygen, and some ClO3 and Cl2O6 are
produced. Cl2O3 is also produced when photolysing the solid at −78 °C:
it is a dark brown solid that explodes below 0 °C. The ClO radical
leads to the depletion of atmospheric ozone and is thus
environmentally important as follows:
:Cl• + O3 ⟶ ClO• + O2
:ClO• + O• ⟶ Cl• + O2

Chlorine perchlorate (ClOClO3) is a pale yellow liquid that is less
stable than ClO2 and decomposes at room temperature to form chlorine,
oxygen, and dichlorine hexoxide (Cl2O6). Chlorine perchlorate may also
be considered a chlorine derivative of perchloric acid (HOClO3),
similar to the thermally unstable chlorine derivatives of other
oxoacids: examples include chlorine nitrate (ClONO2, vigorously
reactive and explosive), and chlorine fluorosulfate (ClOSO2F, more
stable but still moisture-sensitive and highly reactive). Dichlorine
hexoxide is a dark-red liquid that freezes to form a solid which turns
yellow at −180 °C: it is usually made by reaction of chlorine dioxide
with oxygen. Despite attempts to rationalise it as the dimer of ClO3,
it reacts more as though it were chloryl perchlorate, [ClO2]+[ClO4]−,
which has been confirmed to be the correct structure of the solid. It
hydrolyses in water to give a mixture of chloric and perchloric acids:
the analogous reaction with anhydrous hydrogen fluoride does not
proceed to completion.

Dichlorine heptoxide (Cl2O7) is the anhydride of perchloric acid
(HClO4) and can readily be obtained from it by dehydrating it with
phosphoric acid at −10 °C and then distilling the product at −35 °C
and 1 mmHg. It is a shock-sensitive, colourless oily liquid. It is the
least reactive of the chlorine oxides, being the only one to not set
organic materials on fire at room temperature. It may be dissolved in
water to regenerate perchloric acid or in aqueous alkalis to
regenerate perchlorates. However, it thermally decomposes explosively
by breaking one of the central Cl-O bonds, producing the radicals ClO3
and ClO4 which immediately decompose to the elements through
intermediate oxides.


Chlorine oxoacids and oxyanions
=================================
Standard reduction potentials for aqueous Cl species
!! (acid)!!!! (base)
|Cl2/Cl−      +1.358  |Cl2/Cl−      +1.358
|HOCl/Cl−     +1.484  ClO−/Cl−    +0.890
|/Cl− +1.459         
|HOCl/Cl2       +1.630  ClO−/Cl2      +0.421
|HClO2/Cl2      +1.659         
|/Cl2   +1.468         
|/Cl2   +1.277         
|HClO2/HOCl     +1.701  /ClO− +0.681
|           /ClO− +0.488
|/HClO2 +1.181  /       +0.295
|/      +1.201  /       +0.374

Chlorine forms four oxoacids: hypochlorous acid (HOCl), chlorous acid
(HOClO), chloric acid (HOClO2), and perchloric acid (HOClO3). As can
be seen from the redox potentials given in the adjacent table,
chlorine is much more stable towards disproportionation in acidic
solutions than in alkaline solutions:
: Cl2 + H2O       HOCl + H+ + Cl−      'K'ac = 4.2 × 10−4 mol2 l−2
Cl2 + 2 OH−     OCl− + H2O + Cl−   'K'alk = 7.5 × 1015 mol−1 l

The hypochlorite ions also disproportionate further to produce
chloride and chlorate (3 ClO−  2 Cl− + ) but this reaction is quite
slow at temperatures below 70 °C in spite of the very favourable
equilibrium constant of 1027. The chlorate ions may themselves
disproportionate to form chloride and perchlorate (4   Cl− + 3 ) but
this is still very slow even at 100 °C despite the very favourable
equilibrium constant of 1020. The rates of reaction for the chlorine
oxyanions increases as the oxidation state of chlorine decreases. The
strengths of the chlorine oxyacids increase very quickly as the
oxidation state of chlorine increases due to the increasing
delocalisation of charge over more and more oxygen atoms in their
conjugate bases.

Most of the chlorine oxoacids may be produced by exploiting these
disproportionation reactions. Hypochlorous acid (HOCl) is highly
reactive and quite unstable; its salts are mostly used for their
bleaching and sterilising abilities. They are very strong oxidising
agents, transferring an oxygen atom to most inorganic species.
Chlorous acid (HOClO) is even more unstable and cannot be isolated or
concentrated without decomposition: it is known from the decomposition
of aqueous chlorine dioxide. However, sodium chlorite is a stable salt
and is useful for bleaching and stripping textiles, as an oxidising
agent, and as a source of chlorine dioxide. Chloric acid (HOClO2) is a
strong acid that is quite stable in cold water up to 30%
concentration, but on warming gives chlorine and chlorine dioxide.
Evaporation under reduced pressure allows it to be concentrated
further to about 40%, but then it decomposes to perchloric acid,
chlorine, oxygen, water, and chlorine dioxide. Its most important salt
is sodium chlorate, mostly used to make chlorine dioxide to bleach
paper pulp. The decomposition of chlorate to chloride and oxygen is a
common way to produce oxygen in the laboratory on a small scale.
Chloride and chlorate may comproportionate to form chlorine as
follows:
: + 5 Cl− + 6 H+ ⟶ 3 Cl2 + 3 H2O

Perchlorates and perchloric acid (HOClO3) are the most stable
oxo-compounds of chlorine, in keeping with the fact that chlorine
compounds are most stable when the chlorine atom is in its lowest (−1)
or highest (+7) possible oxidation states. Perchloric acid and aqueous
perchlorates are vigorous and sometimes violent oxidising agents when
heated, in stark contrast to their mostly inactive nature at room
temperature due to the high activation energies for these reactions
for kinetic reasons. Perchlorates are made by electrolytically
oxidising sodium chlorate, and perchloric acid is made by reacting
anhydrous sodium perchlorate or barium perchlorate with concentrated
hydrochloric acid, filtering away the chloride precipitated and
distilling the filtrate to concentrate it. Anhydrous perchloric acid
is a colourless mobile liquid that is sensitive to shock that explodes
on contact with most organic compounds, sets hydrogen iodide and
thionyl chloride on fire and even oxidises silver and gold. Although
it is a weak ligand, weaker than water, a few compounds involving
coordinated  are known. The Table below presents typical oxidation
states for chlorine element as given in the secondary schools or
colleges. There are more complex chemical compounds, the structure of
which can only be explained using modern quantum chemical methods, for
example, cluster technetium chloride [(CH3)4N]3[Tc6Cl14], in which 6
of the 14 chlorine atoms are formally divalent, and oxidation states
are fractional. In addition, all the above chemical regularities are
valid for "normal" or close to normal conditions, while at ultra-high
pressures (for example, in the cores of large planets), chlorine can
form a Na3Cl compound with sodium, which does not fit into traditional
concepts of chemistry.

Chlorine oxidation state        −1    +1      +3      +5      +7
Name    chloride        hypochlorite    chlorite        chlorate        perchlorate
Formula Cl−   ClO−
Structure       The chloride ion        The hypochlorite ion    The chlorite ion        The
chlorate ion    The perchlorate ion


Organochlorine compounds
==========================
Like the other carbon-halogen bonds, the C-Cl bond is a common
functional group that forms part of core organic chemistry. Formally,
compounds with this functional group may be considered organic
derivatives of the chloride anion. Due to the difference of
electronegativity between chlorine (3.16) and carbon (2.55), the
carbon in a C-Cl bond is electron-deficient and thus electrophilic.
Chlorination modifies the physical properties of hydrocarbons in
several ways: chlorocarbons are typically denser than water due to the
higher atomic weight of chlorine versus hydrogen, and aliphatic
organochlorides are alkylating agents because chloride is a leaving
group.

Alkanes and aryl alkanes may be chlorinated under free-radical
conditions, with UV light. However, the extent of chlorination is
difficult to control: the reaction is not regioselective and often
results in a mixture of various isomers with different degrees of
chlorination, though this may be permissible if the products are
easily separated. Aryl chlorides may be prepared by the Friedel-Crafts
halogenation, using chlorine and a Lewis acid catalyst. The haloform
reaction, using chlorine and sodium hydroxide, is also able to
generate alkyl halides from methyl ketones, and related compounds.
Chlorine adds to the multiple bonds on alkenes and alkynes as well,
giving di- or tetrachloro compounds. However, due to the expense and
reactivity of chlorine, organochlorine compounds are more commonly
produced by using hydrogen chloride, or with chlorinating agents such
as phosphorus pentachloride (PCl5) or thionyl chloride (SOCl2). The
last is very convenient in the laboratory because all side products
are gaseous and do not have to be distilled out.

Many organochlorine compounds have been isolated from natural sources
ranging from bacteria to humans. Chlorinated organic compounds are
found in nearly every class of biomolecules including alkaloids,
terpenes, amino acids, flavonoids, steroids, and fatty acids.
Organochlorides, including dioxins, are produced in the high
temperature environment of forest fires, and dioxins have been found
in the preserved ashes of lightning-ignited fires that predate
synthetic dioxins. In addition, a variety of simple chlorinated
hydrocarbons including dichloromethane, chloroform, and carbon
tetrachloride have been isolated from marine algae. A majority of the
chloromethane in the environment is produced naturally by biological
decomposition, forest fires, and volcanoes.

Some types of organochlorides, though not all, have significant
toxicity to plants or animals, including humans. Dioxins, produced
when organic matter is burned in the presence of chlorine, and some
insecticides, such as DDT, are persistent organic pollutants which
pose dangers when they are released into the environment. For example,
DDT, which was widely used to control insects in the mid 20th century,
also accumulates in food chains, and causes reproductive problems
(e.g., eggshell thinning) in certain bird species. Due to the ready
homolytic fission of the C-Cl bond to create chlorine radicals in the
upper atmosphere, chlorofluorocarbons have been phased out due to the
harm they do to the ozone layer.


                             Occurrence
======================================================================
Chlorine is too reactive to occur as the free element in nature but is
very abundant in the form of its chloride salts. It is the 20th most
abundant element in Earth's crust and makes up 126 parts per million
of it, through the large deposits of chloride minerals, especially
sodium chloride, that have been evaporated from water bodies. All of
these pale in comparison to the reserves of chloride ions in seawater:
smaller amounts at higher concentrations occur in some inland seas and
underground brine wells, such as the Great Salt Lake in Utah and the
Dead Sea in Israel.

Small batches of chlorine gas are prepared in the laboratory by
combining hydrochloric acid and manganese dioxide, but the need rarely
arises due to its ready availability. In industry, elemental chlorine
is usually produced by the electrolysis of sodium chloride dissolved
in water. This method, the chloralkali process industrialized in 1892,
now provides most industrial chlorine gas. Along with chlorine, the
method yields hydrogen gas and sodium hydroxide, which is the most
valuable product. The process proceeds according to the following
chemical equation:

:2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH


                             Production
======================================================================
Chlorine is primarily produced by the chloralkali process, although
non-chloralkali processes exist. Global 2022 production was estimated
to be 97 million tonnes. The most visible use of chlorine is in water
disinfection. 35-40 % of chlorine produced is used to make poly(vinyl
chloride) through ethylene dichloride and vinyl chloride.  The
chlorine produced is available in cylinders from sizes ranging from
450 g to 70 kg, as well as drums (865 kg), tank wagons (15 tonnes on
roads; 27-90 tonnes by rail), and barges (600-1200 tonnes).
Due to the difficulty and hazards in transporting elemental chlorine,
production is typically located near where it is consumed. As
examples, vinyl chloride producers such as Westlake Chemical and
Formosa Plastics have integrated chloralkali assets.


Chloralkali processes
=======================
The electrolysis of chloride solutions all proceed according to the
following equations:

:Cathode: 2 H2O + 2 e− → H2 + 2 OH−
:Anode: 2 Cl− → Cl2 + 2 e−

In the conventional case where sodium chloride is electrolyzed, sodium
hydroxide and chlorine are coproducts.

Industrially, there are three chloralkali processes:
* The Castner-Kellner process that utilizes a mercury electrode
* The diaphragm cell process that utilizes an asbestos diaphragm that
separates the cathode and anode
* The membrane cell process that uses an ion exchange membrane in
place of the diaphragm

The Castner-Kellner process was the first method used at the end of
the nineteenth century to produce chlorine on an industrial scale.
Mercury (that is toxic) was used as an electrode to amalgamate the
sodium product, preventing undesirable side reactions.

In diaphragm cell electrolysis, an asbestos (or polymer-fiber)
diaphragm separates a cathode and an anode, preventing the chlorine
forming at the anode from re-mixing with the sodium hydroxide and the
hydrogen formed at 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 is partially depleted. Diaphragm methods
produce dilute and slightly impure alkali, but they are not burdened
with the problem of mercury disposal and they are more energy
efficient.

Membrane cell electrolysis employs permeable membrane as an ion
exchanger. Saturated sodium (or potassium) chloride solution is passed
through the anode compartment, leaving at a lower concentration. This
method also produces very pure sodium (or potassium) hydroxide but has
the disadvantage of requiring very pure brine at high concentrations.
However, due to the lower energy requirements of the membrane process,
new chlor-alkali installations are now almost exclusively employing
the membrane process. Next to this, the use of large volumes of
mercury is considered undesirable.
Also, older plants are converted into the membrane process.


Non-chloralkali processes
===========================
In the Deacon process, hydrogen chloride recovered from the production
of organochlorine compounds is recovered as chlorine. The process
relies on oxidation using oxygen:
: 4 HCl + O2 → 2 Cl2 + 2 H2O

The reaction requires a catalyst. As introduced by Deacon, early
catalysts were based on copper. Commercial processes, such as the
Mitsui MT-Chlorine Process, have switched to chromium and
ruthenium-based catalysts.


                            Applications
======================================================================
Sodium chloride is the most common chlorine compound, and is the main
source of chlorine for the demand by the chemical industry. About
15000 chlorine-containing compounds are commercially traded, including
such diverse compounds as chlorinated methane, ethanes, vinyl
chloride, polyvinyl chloride (PVC), aluminium trichloride for
catalysis, the chlorides of magnesium, titanium, zirconium, and
hafnium which are the precursors for producing the pure form of those
elements.

Quantitatively, of all elemental chlorine produced, about 63% is used
in the manufacture of organic compounds, and 18% in the manufacture of
inorganic chlorine compounds. About 15,000 chlorine compounds are used
commercially. The remaining 19% of chlorine produced is used for
bleaches and disinfection products. The most significant of organic
compounds 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. The major inorganic compounds
include HCl, Cl2O, HOCl, NaClO3, AlCl3, SiCl4, SnCl4, PCl3, PCl5,
POCl3, AsCl3, SbCl3, SbCl5, BiCl3, and ZnCl2.


Combating putrefaction
========================
In France (as elsewhere), animal intestines were processed to make
musical instrument strings, Goldbeater's skin and other products. This
was done in "gut factories" ('boyauderies'), and it was an odiferous
and unhealthy process. In or about 1820, the Société d'encouragement
pour l'industrie nationale offered a prize for the discovery of a
method, chemical or mechanical, for separating the peritoneal membrane
of animal intestines without putrefaction. The prize was won by
Antoine-Germain Labarraque, a 44-year-old French chemist and
pharmacist who had discovered that Berthollet's chlorinated bleaching
solutions ("'Eau de Javel'") not only destroyed the smell of
putrefaction of animal tissue decomposition, but also actually
retarded the decomposition.

Labarraque's research resulted in the use of chlorides and
hypochlorites of lime (calcium hypochlorite) and of sodium (sodium
hypochlorite) in the 'boyauderies.' The same chemicals were found to
be useful in the routine disinfection and deodorization of latrines,
sewers, markets, abattoirs, anatomical theatres, and morgues. They
were successful in hospitals, lazarets, prisons, infirmaries (both on
land and at sea), magnaneries, stables, cattle-sheds, etc.; and they
were beneficial during exhumations, embalming, outbreaks of epidemic
disease, fever, and blackleg in cattle.


Disinfection
==============
Labarraque's chlorinated lime and soda solutions have been advocated
since 1828 to prevent infection (called "contagious infection",
presumed to be transmitted by "miasmas"), and to treat putrefaction of
existing wounds, including septic wounds. In his 1828 work, Labarraque
recommended that doctors breathe chlorine, wash their hands in
chlorinated lime, and even sprinkle chlorinated lime about the
patients' beds in cases of "contagious infection". In 1828, the
contagion of infections was well known, even though the agency of the
microbe was not discovered until more than half a century later.

During the Paris cholera outbreak of 1832, large quantities of
so-called 'chloride of lime' were used to disinfect the capital. This
was not simply modern calcium chloride, but chlorine gas dissolved in
lime-water (dilute calcium hydroxide) to form calcium hypochlorite
(chlorinated lime). Labarraque's discovery helped to remove the
terrible stench of decay from hospitals and dissecting rooms, and by
doing so, effectively deodorised the Latin Quarter of Paris. These
"putrid miasmas" were thought by many to cause the spread of
"contagion" and "infection" - both words used before the germ theory
of infection. Chloride of lime was used for destroying odors and
"putrid matter". One source claims chloride of lime was used by Dr.
John Snow to disinfect water from the cholera-contaminated well that
was feeding the Broad Street pump in 1854 London, though three other
reputable sources that describe that famous cholera epidemic do not
mention the incident. One reference makes it clear that chloride of
lime was used to disinfect the offal and filth in the streets
surrounding the Broad Street pump - a common practice in
mid-nineteenth century England.


Semmelweis and experiments with antisepsis
============================================
Perhaps the most famous application of Labarraque's chlorine and
chemical base solutions was in 1847, when Ignaz Semmelweis used
chlorine-water (chlorine dissolved in pure water, which was cheaper
than chlorinated lime solutions) to disinfect the hands of Austrian
doctors, which Semmelweis noticed still carried the stench of
decomposition from the dissection rooms to the patient examination
rooms. Long before the germ theory of disease, Semmelweis theorized
that "cadaveric particles" were transmitting decay from fresh medical
cadavers to living patients, and he used the well-known "Labarraque's
solutions" as the only known method to remove the smell of decay and
tissue decomposition (which he found that soap did not). The solutions
proved to be far more effective antiseptics than soap (Semmelweis was
also aware of their greater efficacy, but not the reason), and this
resulted in Semmelweis's celebrated success in stopping the
transmission of childbed fever ("puerperal fever") in the maternity
wards of Vienna General Hospital in Austria in 1847.

Much later, during World War I in 1916, a standardized and diluted
modification of Labarraque's solution containing hypochlorite (0.5%)
and boric acid as an acidic stabilizer was developed by Henry Drysdale
Dakin (who gave full credit to Labarraque's prior work in this area).
Called Dakin's solution, the method of wound irrigation with
chlorinated solutions allowed antiseptic treatment of a wide variety
of open wounds, long before the modern antibiotic era. A modified
version of this solution continues to be employed in wound irrigation
in modern times, where it remains effective against bacteria that are
resistant to multiple antibiotics (see Century Pharmaceuticals).


Public sanitation
===================
The first continuous application of chlorination to drinking U.S.
water was installed in Jersey City, New Jersey, in 1908. By 1918, the
US Department of Treasury called for all drinking water to be
disinfected with chlorine. Chlorine is presently an important chemical
for water purification (such as in water treatment plants), in
disinfectants, and in bleach. Even small water supplies are now
routinely chlorinated.

Chlorine is usually used (in the form of hypochlorous acid) to kill
bacteria and other microbes in drinking water supplies and public
swimming pools. In most private swimming pools, chlorine itself is not
used, but rather sodium hypochlorite, formed from chlorine and sodium
hydroxide, or solid tablets of chlorinated isocyanurates. The drawback
of using chlorine in swimming pools is that the chlorine reacts with
the amino acids in proteins in human hair and skin. Contrary to
popular belief, the distinctive "chlorine aroma" associated with
swimming pools is not the result of elemental chlorine itself, but of
chloramine, a chemical compound produced by the reaction of free
dissolved chlorine with amines in organic substances including those
in urine and sweat. As a disinfectant in water, chlorine is more than
three times as effective against 'Escherichia coli' as bromine, and
more than six times as effective as iodine. Increasingly,
monochloramine itself is being directly added to drinking water for
purposes of disinfection, a process known as chloramination.

It is often impractical to store and use poisonous chlorine gas for
water treatment, so alternative methods of adding chlorine are used.
These include hypochlorite solutions, which gradually release chlorine
into the water, and compounds like sodium dichloro-s-triazinetrione
(dihydrate or anhydrous), sometimes referred to as "dichlor", and
trichloro-s-triazinetrione, sometimes referred to as "trichlor". These
compounds are stable while solid and may be used in powdered,
granular, or tablet form. When added in small amounts to pool water or
industrial water systems, the chlorine atoms hydrolyze from the rest
of the molecule, forming hypochlorous acid (HOCl), which acts as a
general biocide, killing germs, microorganisms, algae, and so on.


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 the distinctive smell of a
mixture of pepper and pineapple. It also tasted metallic and stung the
back of the throat and chest. Chlorine reacts with water in the mucosa
of the lungs to form hydrochloric acid, destructive to living tissue
and potentially lethal. Human respiratory systems can be protected
from chlorine gas by gas masks with activated charcoal or other
filters, which makes chlorine gas much less lethal than 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, which
developed methods for discharging chlorine gas against an entrenched
enemy. After its first use, both sides in the conflict used chlorine
as a chemical weapon, but it was soon replaced by the more deadly
phosgene and mustard gas.


Middle east
=============
Chlorine gas was also used during the Iraq War in Anbar Province in
2007, with insurgents packing truck bombs with mortar shells and
chlorine tanks. The attacks killed two people from the explosives and
sickened more than 350. 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. In
some bombings, over a hundred civilians were hospitalized due to
breathing difficulties. The Iraqi authorities tightened security for
elemental chlorine, which is essential for providing safe drinking
water to the population.

On 23 October 2014, it was reported that the Islamic State of Iraq and
the Levant had used chlorine gas in the town of Duluiyah, Iraq.
Laboratory analysis of clothing and soil samples confirmed the use of
chlorine gas against Kurdish Peshmerga Forces in a vehicle-borne
improvised explosive device attack on 23 January 2015 at the Highway
47 Kiske Junction near Mosul.

Another country in the middle east, Syria, has used chlorine as a
chemical weapon delivered from barrel bombs and rockets. In 2016, the
OPCW-UN Joint Investigative Mechanism concluded that the Syrian
government used chlorine as a chemical weapon in three separate
attacks. Later investigations from the OPCW's Investigation and
Identification Team concluded that the Syrian Air Force was
responsible for chlorine attacks in 2017 and 2018.


                          Biological role
======================================================================
The chloride anion is an essential nutrient for metabolism. Chlorine
is needed for the production of hydrochloric acid in the stomach and
in cellular pump functions. The main dietary source is table salt, or
sodium chloride. Overly low or high concentrations of chloride in the
blood are examples of electrolyte disturbances. Hypochloremia (having
too little chloride) rarely occurs in the absence of other
abnormalities. It is sometimes associated with hypoventilation. It can
be associated with chronic respiratory acidosis. Hyperchloremia
(having too much chloride) usually does not produce symptoms. When
symptoms do occur, they tend to resemble those of hypernatremia
(having too much sodium). Reduction in blood chloride leads to
cerebral dehydration; symptoms are most often caused by rapid
rehydration which results in cerebral edema. Hyperchloremia can affect
oxygen transport.


                              Hazards
======================================================================
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Chlorine is a toxic gas that attacks the respiratory system, eyes, and
skin. Because it is denser 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 with measuring devices in concentrations as low
as 0.2 parts per million (ppm), and by smell at 3 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. The IDLH
(immediately dangerous to life and health) concentration is 10 ppm.
Breathing lower concentrations can aggravate the respiratory system
and exposure to the gas can irritate the eyes. When chlorine is
inhaled at concentrations greater than 30 ppm, it reacts with water
within the lungs, producing hydrochloric acid (HCl) and hypochlorous
acid (HOCl).

When used at specified levels for water disinfection, the reaction of
chlorine with water is not a major concern for human health. Other
materials present in the water may generate disinfection by-products
that are associated with negative effects on human health.

In the United States, the Occupational Safety and Health
Administration (OSHA) has set the permissible exposure limit for
elemental chlorine at 1 ppm, or 3 mg/m3. The National Institute for
Occupational Safety and Health has designated a recommended exposure
limit of 0.5 ppm over 15 minutes.

In the home, accidents occur when hypochlorite bleach solutions come
into contact with certain acidic drain-cleaners to produce chlorine
gas. Hypochlorite bleach (a popular laundry additive) combined with
ammonia (another popular laundry additive) produces chloramines,
another toxic group of chemicals.


Chlorine-induced cracking in structural materials
===================================================
Chlorine is widely used for purifying water, especially potable water
supplies and water used in swimming pools. Several catastrophic
collapses of swimming pool ceilings have occurred from
chlorine-induced stress corrosion cracking of stainless steel
suspension rods. Some polymers are also sensitive to attack, including
acetal resin and polybutene. Both materials were used in hot and cold
water domestic plumbing, and stress corrosion cracking caused
widespread failures in the US in the 1980s and 1990s.


Chlorine-iron fire
====================
The element iron can combine with chlorine at high temperatures in a
strong exothermic reaction, creating a 'chlorine-iron fire'.
Chlorine-iron fires are a risk in chemical process plants, where much
of the pipework that carries chlorine gas is made of steel.


                              See also
======================================================================
* 2022 Aqaba toxic gas leak
* Chlorine cycle
* Chlorine gas poisoning
* Industrial gas
* Polymer degradation
* Reductive dechlorination


                           External links
======================================================================
* [https://www.periodicvideos.com/videos/017.htm Chlorine] at 'The
Periodic Table of Videos' (University of Nottingham)
* Agency for Toxic Substances and Disease Registry:
[https://web.archive.org/web/20100607045216/http://www.atsdr.cdc.gov/substances/toxsubstance.asp?toxid=36
Chlorine]
*
[https://web.archive.org/web/20090518141937/http://electrochem.cwru.edu/encycl/art-b01-brine.htm
Electrolytic production]
*
[https://web.archive.org/web/20100429205421/http://www.amazingrust.com/Experiments/how_to/Liquid_Cl2.html
Production and liquefaction of chlorine]
*
[https://web.archive.org/web/20080706020039/http://oceana.org/chlorine
Chlorine Production Using Mercury, Environmental Considerations and
Alternatives]
*
[https://web.archive.org/web/20091029202733/http://www.npi.gov.au/database/substance-info/profiles/20.html
National Pollutant Inventory - Chlorine]
* [https://www.cdc.gov/niosh/topics/chlorine/ National Institute for
Occupational Safety and Health - Chlorine Page]
* [http://www.chlorineinstitute.org/ Chlorine Institute]  - Trade
association representing the chlorine industry
* [https://www.eurochlor.org/ Chlorine Online] - the web portal of
Eurochlor - the business association of the European chlor-alkali
industry


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