Chlorine
Chlorine | ||||||||||||||||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
Pronunciation | /ˈklɔːriːn, -aɪn/ | |||||||||||||||||||||||||||
Appearance | pale yellow-green gas | |||||||||||||||||||||||||||
Standard atomic weight Ar°(Cl) | ||||||||||||||||||||||||||||
Chlorine in the periodic table | ||||||||||||||||||||||||||||
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kJ/mol | ||||||||||||||||||||||||||||
Heat of vaporisation | (Cl2) 20.41 kJ/mol | |||||||||||||||||||||||||||
Molar heat capacity | (Cl2) 33.949 J/(mol·K) | |||||||||||||||||||||||||||
Vapour pressure
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Atomic properties | ||||||||||||||||||||||||||||
Discovery and first isolation | Carl Wilhelm Scheele (1774) | |||||||||||||||||||||||||||
Recognized as an element by | Humphry Davy (1808) | |||||||||||||||||||||||||||
Isotopes of chlorine | ||||||||||||||||||||||||||||
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Chlorine is a
Chlorine played an important role in the experiments conducted by medieval
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 twenty-first most abundant chemical element in Earth's crust. These crustal deposits are nevertheless dwarfed by the huge reserves of chloride in seawater.
Elemental chlorine is commercially produced from
In the form of chloride
History
The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that
Early discoveries
Around 900, the authors of the Arabic writings attributed to
Isolation
The element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, and he is credited with the discovery.[15][16] Scheele produced chlorine by reacting MnO2 (as the mineral pyrolusite) with HCl:[14]
- 4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2
Scheele observed several of the properties of chlorine: the bleaching effect on
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
In 1809,
In 1810,
Later uses
Chlorine gas was first used by French chemist
Elemental chlorine solutions dissolved in
Chlorine gas was first used as a weapon on April 22, 1915, at the
Properties
Chlorine is the second
All four stable halogens experience intermolecular
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.[38]
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
The most stable chlorine radioisotope is 36Cl. The primary decay mode of isotopes lighter than 35Cl is
Chemistry and compounds
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
Given that E°(1/2O2/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.[44]
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:[45]
- 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).[45]
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.[45] Hydrochloric acid is a strong acid (pKa = −7) because the hydrogen bonds to chlorine are too weak to inhibit dissociation. The HCl/H2O system has many hydrates HCl·nH2O 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.[46]
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
2 ions – the latter, in any case, are much less stable than the bifluoride ions (HF−
2) 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 NR+
4 (R = Me, Et, Bun) 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 electrophiles containing lone-pairs or π bonds. Solvolysis, ligand replacement reactions, and oxidations are well-characterised in hydrogen chloride solution:[47]
- Ph3SnCl + HCl ⟶ Ph2SnCl2 + PhH (solvolysis)
- Ph3COH + 3 HCl ⟶ Ph
3C+
HCl−
2 + H3O+Cl− (solvolysis) - Me
4N+
HCl−
2 + BCl3 ⟶ Me
4N+
BCl−
4 + HCl (ligand replacement) - PCl3 + Cl2 + HCl ⟶ PCl+
4HCl−
2 (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.[48] 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
- EuCl3 + 1/2 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.[48]
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 [Cl2]+ 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 [Cl3]+ cation is more stable and may be produced as follows:[49]
- Cl2 + ClF + AsF5 [Cl3]+[AsF6]−
This reaction is conducted in the oxidising solvent arsenic pentafluoride. The trichloride anion, [Cl3]−, has also been characterised; it is analogous to triiodide.[50]
Chlorine fluorides
The three fluorides of chlorine form a subset of the
2, ClF−
4, ClF+
2, and Cl2F+.[51] 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).[50]
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:[50]
- 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.[52] It can act as a fluoride ion donor or acceptor (Lewis base or acid), although it does not dissociate appreciably into ClF+
2 and ClF−
4 ions.[53]
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:[54]
- 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.[55]
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.[56]
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.[56]
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:[56]
- ClO−
3 + Cl− + 2 H+ ⟶ ClO2 + 1/2 Cl2 + H2O
Its production is thus intimately linked to the redox reactions of the chlorine oxoacids. It is a strong oxidising agent, reacting with
- 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).[56] 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).[57] 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.[56]
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.[56]
Chlorine oxoacids and oxyanions
E°(couple) | a(H+) = 1 (acid) |
E°(couple) | a(OH−) = 1 (base) |
---|---|---|---|
Cl2/Cl− | +1.358 | Cl2/Cl− | +1.358 |
HOCl/Cl− | +1.484 | ClO−/Cl− | +0.890 |
ClO− 3/Cl− |
+1.459 | ||
HOCl/Cl2 | +1.630 | ClO−/Cl2 | +0.421 |
HClO2/Cl2 | +1.659 | ||
ClO− 3/Cl2 |
+1.468 | ||
ClO− 4/Cl2 |
+1.277 | ||
HClO2/HOCl | +1.701 | ClO− 2/ClO− |
+0.681 |
ClO− 3/ClO− |
+0.488 | ||
ClO− 3/HClO2 |
+1.181 | ClO− 3/ClO− 2 |
+0.295 |
ClO− 4/ClO− 3 |
+1.201 | ClO− 4/ClO− 3 |
+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:[44]
Cl2 + H2O ⇌ HOCl + H+ + Cl− Kac = 4.2 × 10−4 mol2 l−2 Cl2 + 2 OH− ⇌ OCl− + H2O + Cl− Kalk = 7.5 × 1015 mol−1 l
The hypochlorite ions also disproportionate further to produce chloride and chlorate (3 ClO− ⇌ 2 Cl− + ClO−
3) 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 ClO−
3 ⇌ Cl− + 3 ClO−
4) 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.[44]
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:[58]
- ClO−
3 + 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 ClO−
4 are known.[58] The Table below presents typical oxidation states for chlorine element as given in the secondary schools or colleges. Anyhow in university chemistry courses it should be pointed out that 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 [1].[59] 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 exhibit an oxidation state of -3, forming a Na3Cl compound with sodium, which does not fit into traditional concepts of chemistry.[60]
Chlorine oxidation state | −1 | +1 | +3 | +5 | +7 |
---|---|---|---|---|---|
Name | chloride | hypochlorite | chlorite | chlorate | perchlorate |
Formula | Cl− | ClO− | ClO− 2 |
ClO− 3 |
ClO− 4 |
Structure |
Organochlorine compounds
Like the other carbon–halogen bonds, the C–Cl bond is a common functional group that forms part of core
Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.[62][63] Chlorinated organic compounds are found in nearly every class of biomolecules including alkaloids, terpenes, amino acids, flavonoids, steroids, and fatty acids.[62][64] 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.[65] In addition, a variety of simple chlorinated hydrocarbons including dichloromethane, chloroform, and carbon tetrachloride have been isolated from marine algae.[66] A majority of the chloromethane in the environment is produced naturally by biological decomposition, forest fires, and volcanoes.[67]
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.[68] 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.[56]
Occurrence and production
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 twenty-first most abundant element in Earth's crust and makes up 126
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.[31] 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:[70]
- 2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH
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 diaphragm cell electrolysis, an
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.[72]
In the
- 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.[73] 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).[74]
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
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.
Sanitation, disinfection, and antisepsis
Combating putrefaction
In France (as elsewhere),
Labarraque's research resulted in the use of chlorides and hypochlorites of lime (
Disinfection
Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", presumed to be transmitted by "
During the
Semmelweis and experiments with antisepsis
Perhaps the most famous application of Labarraque's chlorine and
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).[88]
Public sanitation
The first continuous application of chlorination to drinking U.S. water was installed in
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.[91] 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.[92] Increasingly, monochloramine itself is being directly added to drinking water for purposes of disinfection, a process known as chloramination.[93]
It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include
Use as a weapon
World War I
Chlorine gas, also known as bertholite, was first
Middle east
Chlorine gas was also used during the
On 23 October 2014, it was reported that the
Another country in the middle east, Syria, has used chlorine as a chemical weapon[104] delivered from barrel bombs and rockets.[105][106] In 2016, the OPCW-UN Joint Investigative Mechanism concluded that the Syrian government used chlorine as a chemical weapon in three separate attacks.[107] 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.[108]
Biological role
The
Hazards
Hazards | |
---|---|
GHS labelling:[113] | |
Danger | |
H270, H315, H319, H330, H335, H400 | |
P220, P233, P244, P261, P304, P312, P340, P403, P410 | |
NFPA 704 (fire diamond) |
Chlorine is a toxic gas that attacks the respiratory system, eyes, and skin.[115] 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.[116][117]
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.
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
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.[118]
In the home, accidents occur when hypochlorite bleach solutions come into contact with certain acidic drain-cleaners to produce chlorine gas.[122] Hypochlorite bleach (a popular laundry additive) combined with ammonia (another popular laundry additive) produces chloramines, another toxic group of chemicals.[123]
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
Chlorine-iron fire
The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.[126][127] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework that carries chlorine gas is made of steel.[126][127]
See also
References
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Explanatory notes
- ^ van Helmont, Joannis Baptistae (1682). Opera omnia [All Works] (in Latin). Frankfurt-am-Main, (Germany): Johann Just Erythropel. From "Complexionum atque mistionum elementalium figmentum." (Formation of combinations and of mixtures of elements), §37, p. 105: Archived 2023-12-30 at the Wayback Machine "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), p. 408 Archived 2023-12-30 at the Wayback Machine: "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:- Helmont, Johannes (Joan) Baptista Van, Encyclopedia.Com Archived 2021-12-18 at the Wayback Machine: "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."
General bibliography
- ISBN 978-0-08-037941-8.
External links
- Chlorine at The Periodic Table of Videos(University of Nottingham)
- Agency for Toxic Substances and Disease Registry: Chlorine
- Electrolytic production
- Production and liquefaction of chlorine
- Chlorine Production Using Mercury, Environmental Considerations and Alternatives
- National Pollutant Inventory – Chlorine
- National Institute for Occupational Safety and Health – Chlorine Page
- Chlorine Institute Archived 2012-08-27 at the Wayback Machine – Trade association representing the chlorine industry
- Chlorine Online – the web portal of Eurochlor – the business association of the European chlor-alkali industry
- Encyclopædia Britannica. Vol. 6 (11th ed.). 1911. pp. 254–56. .