radioisotope, it has uses in biological assays, nuclear medicine imaging and in radiation therapy as brachytherapy to treat a number of conditions, including prostate cancer, uveal melanomas, and brain tumours. finally, iodine-131, with a half-life of eight days, beta decays to an excited state of stable xenon-131 that then converts to the ground state by emitting gamma radiation. it is a common fission product and thus is present in high levels in radioactive fallout. it may then be absorbed through contaminated food, and will also accumulate in the thyroid. as it decays, it may cause damage to the thyroid. the primary risk from exposure to high levels of iodine-131 is the chance occurrence of radiogenic thyroid cancer in later life. other risks include the possibility of non-cancerous growths and thyroiditis. the usual means of protection against the negative effects of iodine-131 is by saturating the thyroid gland with stable iodine-127 in the form of potassium iodide tablets, taken daily for optimal prophylaxis. however, iodine-131 may also be used for medicinal purposes in radiation therapy for this very reason, when tissue destruction is desired after iodine uptake by the tissue. iodine-131 is also used as a radioactive tracer. 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 iodine is quite reactive, but it is much less reactive than the other halogens. for example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide (to phosgene, nitrosyl chloride, and sulfuryl chloride respectively), iodine will not do so. furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. by the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride. charge-transfer complexes the iodine molecule, i2, dissolves in ccl4 and aliphatic hydrocarbons to give bright violet solutions. in these solvents the absorption band maximum occurs in the 520 – 540 nm region and is assigned to a π* to σ* transition. when i2 reacts with lewis bases in these solvents a blue shift in i2 peak is seen and the new peak (230 – 330 nm) arises that is due to the formation of adducts, which are referred to as charge-transfer complexes. hydrogen iodide the simplest compound of iodine is hydrogen iodide, hi. it is a colourless gas that reacts with oxygen to give water and iodine. although it is useful in iodination reactions in the laboratory, it does not have large-scale industrial uses, unlike the other hydrogen halides. commercially, it is usually made by reacting iodine with hydrogen sulfide or hydrazine: 2 i2 + n2h4 h2o⟶ 4 hi + n2 at room temperature, it is a colourless gas, like all of the hydrogen halides except hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative iodine atom. it melts at −51.0 °c and boils at −35.1 °c. it is an endothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless a catalyst is present: the reaction between hydrogen and iodine at room temperature to give hydrogen iodide does not proceed to completion. the h–i bond dissociation energy is likewise the smallest of the hydrogen halides, at 295 kj/mol. aqueous hydrogen iodide is known as hydroiodic acid, which is a strong acid. hydrogen iodide is exceptionally soluble in water: one litre of water will dissolve 425 litres of hydrogen iodide, and the saturated solution has only four water molecules per molecule of hydrogen iodide. commercial so-called "concentrated" hydroiodic acid usually contains 48–57% hi by mass; the solution forms an azeotrope with boiling point 126.7 °c at 56.7 g hi per 100 g solution. hence hydroiodic acid cannot be concentrated past this point by evaporation of water. unlike hydrogen fluoride, anhydrous liquid hydrogen iodide 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 h2i+ and hi−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 iodine, 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 iodide is a poor solvent, able to dissolve only small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides. other binary iodine compounds with the exception of the noble gases, nearly all elements on the periodic table up to einsteinium (esi3 is known) are known to form binary compounds with iodine. until 1990, nitrogen triiodide was only known as an ammonia adduct. ammonia-free ni3 was found to be isolable at –196 °c but spontaneously decomposes at 0 °c. for thermodynamic reasons related to electronegativity of the elements, neutral sulfur and selenium iodides that are stable at room temperature are also nonexistent, although s2i2 and si2 are stable up to 183 and 9 k, respectively. as of 2022, no neutral binary selenium iodide has been unambiguously identified (at any temperature). sulfur- and selenium-iodine polyatomic cations (e.g., 2 and 2) have been prepared and characterized crystallographically. given the large size of the iodide anion and iodine's weak oxidising power, high oxidation states are difficult to achieve in binary iodides, the maximum known being in the pentaiodides of niobium, tantalum, and protactinium. iodides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen iodide gas. these methods work best when the iodide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative iodination of the element with iodine or hydrogen iodide, high-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, carbon tetraiodide, or an organic iodide. for example, molybdenum(iv) oxide reacts with aluminium(iii) iodide at 230 °c to give molybdenum(ii) iodide. an example involving halogen exchange is given below, involving the reaction of tantalum(v) chloride with excess aluminium(iii) iodide at 400 °c to give tantalum(v) iodide: 3 tacl 5 + 5 ali 3 ( excess ) ⟶ 3 tai 5 + 5 alcl 3 -> 3tai5 + 5alcl3}}} lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example: tai 5 + ta → 630 ∘ c ⟶ 575 ∘ c thermal gradient ta 6 i 14 +ta->ta6i14}}} most metal iodides with the metal in low oxidation states (+1 to +3) are ionic. nonmetals tend to form covalent molecular iodides, as do metals in high oxidation states from +3 and above. both ionic and covalent iodides are known for metals in oxidation state +3 (e.g. scandium iodide is mostly ionic, but aluminium iodide is not). ionic iodides min tend to have the lowest melting and boiling points among the halides mxn of the same element, because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion. in contrast, covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element, since iodine is the most polarisable of the halogens and, having the most electrons among them, can contribute the most to van der waals forces. naturally, exceptions abound in intermediate iodides where one trend gives way to the other. similarly, solubilities in water of predominantly ionic iodides (e.g. potassium and calcium) are the greatest among ionic halides of that element, while those of covalent iodides (e.g. silver) are the lowest of that element. in particular, silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine. iodine halides the halogens form many binary, diamagnetic interhalogen compounds with stoichiometries xy, xy3, xy5, and xy7 (where x is heavier than y), and iodine is no exception. iodine forms all three possible diatomic interhalogens, a trifluoride and trichloride, as well as a pentafluoride and, exceptionally among the halogens, a heptafluoride. numerous cationic and anionic derivatives are also characterised, such as the wine-red or bright orange compounds of icl+2 and the dark brown or purplish black compounds of i2cl+. apart from these, some pseudohalides are also known, such as cyanogen iodide (icn), iodine thiocyanate (iscn), and iodine azide (in3). iodine monochloride iodine monofluoride (if) is unstable at room temperature and disproportionates very readily and irreversibly to iodine and iodine pentafluoride, and thus cannot be obtained pure. it can be synthesised from the reaction of iodine with fluorine gas in trichlorofluoromethane at −45 °c, with iodine trifluoride in trichlorofluoromethane at −78 °c, or with silver(i) fluoride at 0 °c. iodine monochloride (icl) and iodine monobromide (ibr), on the other hand, are moderately stable. the former, a volatile red-brown compound, was discovered independently by joseph louis gay-lussac and humphry davy in 1813–1814 not long after the discoveries of chlorine and iodine, and it mimics the intermediate halogen bromine so well that justus von liebig was misled into mistaking bromine (which he had found) for iodine monochloride. iodine monochloride and iodine monobromide may be prepared simply by reacting iodine with chlorine or bromine at room temperature and purified by fractional crystallisation. both are quite reactive and attack even platinum and gold, though not boron, carbon, cadmium, lead, zirconium, niobium, molybdenum, and tungsten. their reaction with organic compounds depends on conditions. iodine chloride vapour tends to chlorinate phenol and salicyclic acid, since when iodine chloride undergoes homolytic dissociation, chlorine and iodine are produced and the former is more reactive. however, iodine chloride in tetrachloromethane solution results in iodination being the main reaction, since now heterolytic fission of the i–cl bond occurs and i+ attacks phenol as an electrophile. however, iodine monobromide tends to brominate phenol even in tetrachloromethane solution because it tends to dissociate into its elements in solution, and bromine is more reactive than iodine. when liquid, iodine monochloride and iodine monobromide dissociate into i2x+ and ix−2 anions (x = cl, br); thus they are significant conductors of electricity and can be used as ionising solvents. iodine trifluoride (if3) is an unstable yellow solid that decomposes above −28 °c. it is thus little-known. it is difficult to produce because fluorine gas would tend to oxidise iodine all the way to the pentafluoride; reaction at low temperature with xenon difluoride is necessary. iodine trichloride, which exists in the solid state as the planar dimer i2cl6, is a bright yellow solid, synthesised by reacting iodine with liquid chlorine at −80 °c; caution is necessary during purification because it easily dissociates to iodine monochloride and chlorine and hence can act as a strong chlorinating agent. liquid iodine trichloride conducts electricity, possibly indicating dissociation to icl+2 and icl−4 ions. iodine pentafluoride (if5), a colourless, volatile liquid, is the most thermodynamically stable iodine fluoride, and can be made by reacting iodine with fluorine gas at room temperature. it is a fluorinating agent, but is mild enough to store in glass apparatus. again, slight electrical conductivity is present in the liquid state because of dissociation to if+4 and if−6. the pentagonal bipyramidal iodine heptafluoride (if7) is an extremely powerful fluorinating agent, behind only chlorine trifluoride, chlorine pentafluoride, and bromine pentafluoride among the interhalogens: it reacts with almost all the elements even at low temperatures, fluorinates pyrex glass to form iodine(vii) oxyfluoride (iof5), and sets carbon monoxide on fire. iodine oxides and oxoacids structure of iodine pentoxide iodine oxides are the most stable of all the halogen oxides, because of the strong i–o bonds resulting from the large electronegativity difference between iodine and oxygen, and they have been known for the longest time. the stable, white, hygroscopic iodine pentoxide (i2o5) has been