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 clf+2 and clf−4 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 +− (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 yellow chlorine dioxide (clo2) gas above a solution containing chlorine dioxide. 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. 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: 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: 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, +−, 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 e°(couple)a(h+) = 1(acid)e°(couple)a(oh−) = 1(base) cl2/cl−+1.358cl2/cl−+1.358 hocl/cl−+1.484clo−/cl−+0.890 clo−3/cl−+1.459 hocl/cl2+1.630clo−/cl2+0.421 hclo2/cl2+1.659 clo−3/cl2+1.468 clo−4/cl2+1.277 hclo2/hocl+1.701clo−2/clo−+0.681 clo−3/clo−+0.488 clo−3/hclo2+1.181clo−3/clo−2+0.295 clo−4/clo−3+1.201clo−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: 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. 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: 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. 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 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. 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 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. 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