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53 telluriumiodinexenon


Name, Symbol, Number iodine, I, 53
Chemical series halogens
Group, Period, Block 17, 5, p
Appearance violet-dark gray, lustrous
Standard atomic weight 126.90447(3)  g·mol−1
Electron configuration [Kr] 4d10 5s2 5p5
Electrons per shell 2, 8, 18, 18, 7
Physical properties
Phase solid
Density (near r.t.) 4.933  g·cm−3
Melting point 386.85 K
(113.7 °C, 236.66 °F)
Boiling point 457.4 K
(184.3 °C, 363.7 °F)
Critical point 819 K, 11.7 MPa
Heat of fusion (I2) 15.52  kJ·mol−1
Heat of vaporization (I2) 41.57  kJ·mol−1
Vapor pressure (rhombic)
at T(K) 260 282 309 342 381 457
Atomic properties
Crystal structure orthorhombic
Oxidation states ±1, 5, 7
(strongly acidic oxide)
Electronegativity 2.66 (scale Pauling)
Ionization energies 1st: 1008.4 kJ/mol
2nd: 1845.9 kJ/mol
3rd: 3180 kJ/mol
Atomic radius 140pm
Atomic radius (calc.) 115  pm
Covalent radius 133  pm
Van der Waals radius 198 pm
Magnetic ordering nonmagnetic
Electrical resistivity (0 °C) 1.3×107 Ω·m
Bulk modulus 7.7  GPa
CAS registry number 7553-56-2
Selected isotopes
iso NA half-life DM DE (MeV) DP
127I 100% I is stable with 74 neutrons
129I syn 15.7×106y β- 0.194 129Xe
131I syn 8.02070 d β- 0.971 131Xe
CAS number 7553-56-2
R-phrases R21 R23 R25 R34
S-phrases S23 S25
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references



Iodine (IPA: /ˈaɪədaɪn, ˈaɪədɪn/, or /ˈaɪədiːn/; from Greek: ιώδης iodes "violet"), is a chemical element that has the symbol I and atomic number 53. Naturally-occurring iodine is a single isotope with 74 neutrons.

Chemically, iodine is the least reactive of the halogens, and the most electropositive halogen after astatine. However, the element does not occur in the free state in nature. As with all other halogens (members of Group VII in the Periodic Table), when freed from its compounds iodine forms diatomic molecules (I2).

Iodine and its compounds are primarily used in medicine, photography and in dyes. Although it is rare in the solar system and Earth's crust, the iodides are very soluble in water, and the element is concentrated in seawater. This mechanism helps to explain how the element came to be required in trace amounts by all animals and some plants, being by far the heaviest element known to be necessary to living organisms.


Iodine under standard conditions is a dark-purple/dark-brown solid. It can be seen apparently subliming at standard temperatures into a violet-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less reactive than the other members of its Group VII (halogens) and has some metallic light reflectance.

Elemental iodine dissolves easily in chloroform and carbon tetrachloride. The solubility of elementary iodine in water can be vastly increased by the addition of potassium iodide. The molecular iodine reacts reversibly with the negative ion, creating the triiodide anion, I3, which dissolves well in water. This is also the formulation of some types of medicinal (antiseptic) iodine, although tincture of iodine classically dissolves the element in alcohol. The deep blue color of starch-iodine complexes is produced only by the free element.

Students who have seen the classroom demonstration in which iodine crystals are gently heated in a test tube to violet vapor, may gain the impression that liquid iodine does not exist at atmospheric pressure. This misconception arises because the small amount of vapor produced has such a deep colour that the liquid appears not to form. In fact, if iodine crystals are heated carefully to just above their melting point of 113.7 °C, the crystals melt into a liquid which is present under a dense blanket of the vapor.





Iodine is used in pharmaceuticals, antiseptics, medicine, food supplements, dyes, catalysts, halogen lights, photography, water purifying, and starch detection.

  • Tincture of iodine (10% elemental iodine in ethanol base) is an essential component of any emergency survival kit, used both to disinfect wounds and to sanitize surface water for drinking (3 drops per litre, let stand for 30 minutes). Alcohol-free iodine solutions such as Lugol's iodine, as well as other iodophor type antiseptics, are also available as effective elemental iodine sources for this purpose.
  • Iodine compounds are important in the field of organic chemistry
  • Silver iodide is used in photography.
  • Iodine crystals are used in the process to make NI3 or nitrogen triiodide. This compound is a shock-sensitive explosive when dry. It has commonly been used for pranks, but because of its extreme touch sensitivity, is not useful commercially.

Occurrence on earth

Iodine naturally occurs in the environment chiefly as dissolved iodide in seawater, although it is also found in some minerals and soils. The element may be prepared in an ultrapure form through the reaction of potassium iodide with copper(II) sulfate. There are also a few other methods of isolating this element in the laboratory-- for example the method used to isolate other halogens: oxidation of the iodide in hydroiodic acid (often made in situ with an iodide and sulfuric acid) by manganese dioxide (see below in Descriptive chemistry). Although the element is actually quite rare, kelp and certain plants and algae have some ability to concentrate iodine, which helps introduce the element into the food chain.


Iodine output in 2005


Extraction from seawater involves electrolysis. The brine is first purified and acidified using sulphuric acid and is then reacted with chlorine. An iodine solution is produced but it is yet too dilute and has to be concentrated. To do this air is blown into the solution which causes the iodine to evaporate, then it is passed into an absorbing tower containing acid where sulfur dioxide is added to reduce the iodine. The solution is then added to chlorine again to concentrate the solution more, and the final solution is at a level of about 99%.

Another source is from kelp. This source was used in the 18th and 19th centuries but is no longer economically viable.

In 2005, Chile was the top producer of iodine with almost two-thirds world share followed by Japan and the USA, reports the British Geological Survey.

Descriptive chemistry


I2+ H2O ↔ H+ + I + HIO   (K = 2.0×10−13)[1]

Solubility in water is greatly improved if the solution contains dissolved iodides such as hydroiodic acid, potassium iodide, or sodium iodide; this extra solubility results from the high solubility of the I3- ion. Dissolved bromides also improve water solubility of iodine. Iodine is soluble in a number of organic solvents, including ethanol (20.5 g/100 ml at 15 °C, 21.43 g/100 ml at 25 °C), diethyl ether (20.6 g/100 ml at 17 °C, 25.20 g/100 ml at 25 °C), chloroform, acetic acid, glycerol, benzene (14.09 g/100 ml at 25 °C), carbon tetrachloride (2.603 g/100 ml at 35 °C), and carbon disulfide (16.47 g/100 ml at 25 °C).[2] Aqueous and ethanol solutions are brown. Solutions in chloroform, carbon tetrachloride, and carbon disulfide are violet.

Elemental iodine can be prepared by oxidizing iodides with chlorine:

2I + Cl2 → I2 + 2Cl

or with manganese dioxide in acid solution:[1]



I2 + H2S → 2HI + S↓

or by hydrazine:

2I2 + N2H4 → 4HI + N2

Iodine is oxidized to iodate by nitric acid:[4]

I2 + 10HNO3 → 2HIO3 + 10NO2 + 4H2O

or by chlorates:[4]

I2 + 2ClO3 → 2IO3 + Cl2


3IO → 2I + IO3 (K = 1020)

Notable inorganic iodine compounds

LiI BeI2 BI3 CI4 NI3 I2O4 I2O5 I4O9 IF IF3 IF5 IF7 Ne
NaI MgI2 AlI3 SiI4 PI3 S ICl ICl3 Ar
KI CaI2 Sc TiI4 V Cr MnI2 Fe Co NiI2 CuI ZnI2 Ga2I6 GeI2 GeI4 As Se IBr Kr
RbI SrI2 Y ZrI4 Nb Mo Tc Ru Rh Pd AgI CdI2 InI3 SnI4 SnI2 SbI3 TeI4 I Xe
CsI BaI2 Hf Ta W Re Os Ir Pt AuI HgI2 TlI PbI2 Bi Po At Rn
Ac ThI4 Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

See also iodine compounds

Stable iodine in biology

Iodine is an essential trace element, the heaviest-known element needed by living organisms. Its only known roles in animal biology are as constituents of the thyroid hormones, thyroxine (T4) and triiodothyronine (T3). These are made from addition condensation products of the amino acid tyrosine, and are stored prior to release in a protein-like iodine-containing molecule called thyroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. The thyroid gland actively absorbs iodide from the blood to make and release these hormones into the blood, actions which are regulated by a second hormone TSH from the pituitary. Thyroid hormones are phylogenetically very old molecules which are synthesized by most multicellular organisms, and which even have some effect on unicellular organisms.


Human dietary intake

The United States Food and Drug Administration recommends 150 micrograms of iodine per day for both men and women.[5] This is necessary for proper production of thyroid hormone.[5] Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil.[6][7] Salt for human consumption is often fortified with iodine and is referred to as iodized salt.[7]

Iodine deficiency

Main article: Iodine deficiency



Radioiodine and biology

Radioiodine and the thyroid

Human exposure to radioactive iodine will cause thyroid uptake, as with all iodine, leading to elevated chances of thyroid cancer. Isotopes with shorter half-lives such as I131 present a greater risk than those with longer half-lives since they generate more radiation per unit of time. Taking large amounts of regular iodine will saturate the thyroid and prevent uptake. Iodine pills are sometimes distributed to persons living close to nuclear establishments, for use in case of accidents that could lead to releases of radioactive iodine.


The artificial radioisotope 131I (a beta emitter), has a half-life of 8.0207 days. Also known as radioiodine, 131I has been used in treating cancer and other pathologies of the thyroid glands. 123I is the radioisotope most often used in nuclear imaging of the kidney and thyroid as well as thyroid uptake scans (used for the evaluation of Graves' Disease). The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO3).

Potassium iodide (KI tablets, or "SSKI" = "Saturated Solution of KI" liquid drops) can be given to people in a nuclear disaster area when fission has taken place, to block the uptake of iodine-131 by the thyroid. The protective effect of KI lasts approximately 24 hours, so it should be dosed daily until a risk of significant exposure to radioiodines no longer exists.[9][10] The exposure can be reduced by evacuation, sheltering, and by control of the food supply. Iodine-131 also decays rapidly, with a half-life of 8 days, so that 99.95% of the original radioiodine is gone after three months.


Radioiodine and the kidney



Main article: isotopes of iodine



Excesses of stable 129Xe in meteorites have been shown to result from decay of "primordial" Iodine-129 produced newly by the supernovas which created the dust and gas from which the solar system formed. 129I was the first extinct radionuclide to be identified as present in the early solar system. Its decay is the basis of the I-Xe Iodine-xenon radiometric dating scheme, which covers the first 83 million years of solar system evolution.

Effects of various radioiodine isotopes in biology are discussed above.

Toxicity of iodine

Excess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms are abnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole. Elemental iodine, I2, is a deadly poison if taken in larger amounts; if 2-3 grams of it is consumed, it is fatal to humans. Iodides are similar in toxicity to bromides.

Precautions for stable iodine


Clandestine use


See also


External links

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