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 (t1/2 = 3.0×105 y) and 38cl (t1/2 = 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°(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. hydrogen chloride 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: 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). 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 (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. 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: ph3sncl + hcl ⟶ ph2sncl2 + phh (solvolysis) ph3coh + 3 hcl ⟶ ph3c+hcl−2 + h3o+cl− (solvolysis) me4n+hcl−2 + bcl3 ⟶ me4n+bcl−4 + hcl (ligand replacement) pcl3 + cl2 + hcl ⟶ pcl+4hcl−2 (oxidation) other binary chlorides 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. 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 + 1/2 h2 ⟶ eucl2 + hcl recl5 at "bp"⟶ recl3 + cl2 aucl3 160 °c⟶ 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: cl2 + clf + asf5 −78 °c⟶ +− 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 clf−2, clf−4, clf+2, 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 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: 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,