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User:InterstellarGamer12321/Chlorine compounds

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Halogen bond energies (kJ/mol)[1]
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.[1]

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.[2]

Hydrogen chloride edit

 
Structure of solid deuterium chloride, with D···Cl hydrogen bonds

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:[3]

NaCl + H2SO4 150 °C⟶  NaHSO4 + HCl
NaCl + NaHSO4 540–600 °C⟶  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).[3]

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.[3] 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.[4]

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 HCl−
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:[5]

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 edit

 
Hydrated nickel(II) chloride, NiCl2(H2O)6.

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.[6] 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:[6]

EuCl3 + 1/2 H2 ⟶ EuCl2 + HCl
ReCl5 at "bp"⟶  ReCl3 + Cl2
AuCl3 160 °C⟶  AuCl + Cl2

Most of the chlorides the metals in groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states, are mostly ionic, while nonmetals tend to form covalent molecular chlorides, as do metals in high oxidation states from +3 and above. Silver chloride is very insoluble in water and is thus often used as a qualitative test for chlorine.[6]

Polychlorine compounds edit

Although dichlorine is a strong oxidising agent with a high first ionisation energy, it may be oxidised under extreme conditions to form the Cl+
2 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 Cl+
3 cation is more stable and may be produced as follows:[7]

Cl2 + ClF + AsF5 −78 °C⟶  Cl+
3AsF−
6

This reaction is conducted in the oxidising solvent arsenic pentafluoride. The trichloride anion, Cl−
3, has also been characterised; it is analogous to triiodide.[8]

Chlorine fluorides edit

The three fluorides of chlorine form a subset of the interhalogen compounds, all of which are diamagnetic.[8] Some cationic and anionic derivatives are known, such as ClF−
2, ClF−
4, ClF+
2, and Cl2F+.[9] 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).[8]

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 direction 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:[8]

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.[10] 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.[11]

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:[12]

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.[13]

Chlorine oxides edit

 
Yellow chlorine dioxide (ClO2) gas above a solution containing chlorine dioxide.[clarification needed]
 
Structure of dichlorine heptoxide, Cl2O7, the most stable of the 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.[14]

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.[14]

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:[14]

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 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·nH2O (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:[14]

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).[14] 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).[15] 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.[14]

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.[14]

Chlorine oxoacids and oxyanions edit

Standard reduction potentials for aqueous Cl species[2]
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:[2]

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.[2]

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:[16]

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.[16]

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 edit

Main article: Organochlorine compound
 
Suggested mechanism for the chlorination of a carboxylic acid by phosphorus pentachloride to form an acyl chloride

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.[17]

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.[17] 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 tetra-chloro 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.[17]

Many organochlorine compounds have been isolated from natural sources ranging from bacteria to humans.[18][19] Chlorinated organic compounds are found in nearly every class of biomolecules including alkaloids, terpenes, amino acids, flavonoids, steroids, and fatty acids.[18][20] 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.[21] In addition, a variety of simple chlorinated hydrocarbons including dichloromethane, chloroform, and carbon tetrachloride have been isolated from marine algae.[22] A majority of the chloromethane in the environment is produced naturally by biological decomposition, forest fires, and volcanoes.[23]

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.[24] 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.[14]

See also edit

References edit

  1. ^ a b Cite error: The named reference Greenwood804 was invoked but never defined (see the help page).
  2. ^ a b c d Greenwood & Earnshaw 1997, pp. 853–856
  3. ^ a b c Greenwood & Earnshaw 1997, pp. 809–812
  4. ^ Greenwood & Earnshaw 1997, pp. 812–816
  5. ^ Greenwood & Earnshaw 1997, pp. 818–819
  6. ^ a b c Greenwood & Earnshaw 1997, pp. 821–844
  7. ^ Greenwood & Earnshaw 1997, pp. 842–844
  8. ^ a b c d Greenwood & Earnshaw 1997, pp. 824–828
  9. ^ Greenwood & Earnshaw 1997, pp. 835–842
  10. ^ "In Situ Cleaning of CVD Chambers". Semiconductor International. 1 June 1999.[permanent dead link]
  11. ^ Greenwood & Earnshaw 1997, pp. 828–831
  12. ^ Greenwood & Earnshaw 1997, pp. 832–835
  13. ^ Greenwood & Earnshaw 1997, pp. 875–880
  14. ^ a b c d e f g h Greenwood & Earnshaw 1997, pp. 844–850
  15. ^ Greenwood & Earnshaw 1997, pp. 883–885
  16. ^ a b Greenwood & Earnshaw 1997, pp. 856–870
  17. ^ a b c M. Rossberg et al. "Chlorinated Hydrocarbons" in Ullmann's Encyclopedia of Industrial Chemistry 2006, Wiley-VCH, Weinheim. doi:10.1002/14356007.a06_233.pub2
  18. ^ a b Gordon W. Gribble (1998). "Naturally Occurring Organohalogen Compounds". Acc. Chem. Res. 31 (3): 141–52. doi:10.1021/ar9701777.
  19. ^ Gordon W. Gribble (1999). "The diversity of naturally occurring organobromine compounds". Chemical Society Reviews. 28 (5): 335–46. doi:10.1039/a900201d.
  20. ^ Kjeld C. Engvild (1986). "Chlorine-Containing Natural Compounds in Higher Plants". Phytochemistry. 25 (4): 7891–91. doi:10.1016/0031-9422(86)80002-4.
  21. ^ Gribble, G. W. (1994). "The Natural production of chlorinated compounds". Environmental Science and Technology. 28 (7): 310A–319A. Bibcode:1994EnST...28..310G. doi:10.1021/es00056a712. PMID 22662801.
  22. ^ Gribble, G. W. (1996). "Naturally occurring organohalogen compounds – A comprehensive survey". Progress in the Chemistry of Organic Natural Products. 68 (10): 1–423. doi:10.1021/np50088a001. PMID 8795309.
  23. ^ Public Health Statement – Chloromethane Archived 2007-09-27 at the Wayback Machine, Centers for Disease Control, Agency for Toxic Substances and Disease Registry
  24. ^ Connell, D.; et al. (1999). Introduction to Ecotoxicology. Blackwell Science. p. 68. ISBN 978-0-632-03852-7.

Category:Chlorine Category:Chlorine compounds Category:Chemical compounds by element

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