Iodine

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53 tellurium ← iodine → xenon
Br
↑
I
↓
At
Periodic Table - Extended Periodic Table
General
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
Heat capacity (25 °C) (I2) 54.44  J·mol−1·K−1
Vapor pressure (rhombic)
P(Pa) 1 10 100 1 k 10 k 100 k
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 140  pm
Atomic radius (calc.) 115  pm
Covalent radius 133  pm
Van der Waals radius 198 pm
Miscellaneous
Magnetic ordering nonmagnetic
Electrical resistivity (0 °C) 1.3×107 Ω·m
Thermal conductivity (300 K) 0.449  W·m−1·K−1
Bulk modulus 7.7  GPa
CAS registry number 7553-56-2
Selected isotopes
Main article: Isotopes of iodine
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
References
For the record label, see Iodine Recordings.

Iodine (IPA: [ˈaɪəˌdaɪn], /ˈaɪəˌdɪn/, or /ˈaɪəˌdiːn/; from Greek: iodes "violet"), is a chemical element that has the symbol I and atomic number 53. Chemically, iodine is the least reactive of the halogens, and the most electropositive halogen after astatine. Iodine is primarily used in medicine, photography and dyes. It is required in trace amounts by most living organisms.

As with all other halogens (members of Group VII in the Periodic Table), iodine forms diatomic molecules, and hence has the molecular formula of I2.

Contents

  • 1 Occurrence on earth
  • 2 Uses
  • 3 Isotopes
  • 4 Notable characteristics
  • 5 Sources
  • 6 Descriptive chemistry
  • 7 History
  • 8 Notable inorganic iodine compounds
  • 9 Stable iodine in biology
    • 9.1 Human dietary intake
    • 9.2 Iodine deficiency
    • 9.3 Toxicity of iodine
  • 10 Radioiodine and biology
    • 10.1 Radioiodine and the thyroid
    • 10.2 Radioiodine and the kidney
  • 11 Non-hormone-related applications of iodine
  • 12 Precautions for stable iodine
  • 13 Clandestine use
  • 14 See also
  • 15 References
  • 16 External links

[edit] 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. Although the element is actually quite rare, kelp and certain other plants have the ability to concentrate iodine, which helps introduce the element into the food chain as well as keeping its cost down.

[edit] Uses

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

[edit] Isotopes

Main article: isotopes of iodine

There are 37 isotopes of iodine and only one, 127I, is stable.

In many ways, 129I is similar to 36Cl. It is a soluble halogen, fairly non-reactive, exists mainly as a non-sorbing anion, and is produced by cosmogenic, thermonuclear, and in-situ reactions. In hydrologic studies, 129I concentrations are usually reported as the ratio of 129I to total I (which is virtually all 127I). As is the case with 36Cl/Cl, 129I/I ratios in nature are quite small, 10−14 to 10−10 (peak thermonuclear 129I/I during the 1960s and 1970s reached about 10−7). 129I differs from 36Cl in that its half-life is longer (15.7 vs. 0.301 million years), it is highly biophilic, and occurs in multiple ionic forms (commonly, I− and IO3−) which have different chemical behaviors. This makes it fairly easy for 129I to enter the biosphere as it becomes incorporated into vegetation, soil, milk, animal tissue, etc.

Excesses of stable 129Xe in meteorites have been shown to result from decay of "primordial" 129I 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 radiometric dating scheme, which covers the first 83 million years of solar system evolution.

Effects of various radioiodine isotopes in biology are discussed below.

[edit] Notable characteristics

Iodine is a dark-gray/purple-black solid that sublimes at standard temperatures into a purple-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less active than the other members of its Group VII (halogens) and has some metallic-like properties. Iodine dissolves easily in chloroform, carbon tetrachloride, or carbon disulphide to form purple solutions. It is only slightly soluble in water, giving a yellow solution. 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 I3− which dissolves well in water. This is also the formulation of medicinal iodine of old. The deep blue color of starch-iodine complexes is produced only by the free element.

Many students who have seen the classroom demonstration where iodine crystals are gently heated in a test tube come away with the impression that liquid iodine cannot exist at atmospheric pressure. This misconception arises because sublimation occurs without the intermediacy of liquid. The truth is that if iodine crystals are heated carefully to their melting point of 113.7 °C, the crystals will fuse into a liquid, which will be present under a dense blanket of the vapour.

[edit] Sources

Iodine output in 2005

Iodine is found in the mineral caliche, found in Chile, between the Andes and the sea. It can also be found in some seaweeds as well as extracted from seawater, however extracting iodine from the mineral is the only economical way to extract the substance.[citation needed]

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%.[citation needed]

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.

[edit] Descriptive chemistry

Elemental iodine is poorly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C. By contrast with chlorine, the formation of the hypohalite ion (IO–) in neutral aqueous solutions of iodine is negligible.

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. 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]

2I– + 4H+ + MnO2 → I2 + 2H2O + Mn2+

Iodine is reduced to hydroiodic acid by hydrogen sulfide:[3]

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

Iodine is converted in a two stage reaction to iodide and iodate in solutions of alkali hydroxides (such as sodium hydroxide):[1]

I2 + 2OH– → I– + IO– + H2O (K = 30)
3IO– → 2I– + IO3– (K = 1020)

[edit] History

Iodine was discovered by Bernard Courtois in 1811. He was born to a manufacturer of saltpeter (a vital part of gunpowder). At the time France was at war, saltpeter was in great demand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash then washed with water. The remaining waste was destroyed by adding sulfuric acid. One day Courtois added too much sulfuric acid and a cloud of purple vapor rose. Courtois noted that the vapor crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element but lacked the money to pursue his observations.

However he gave samples to his friends, Charles Bernard Desormes (1777 - 1862) and Nicolas Clément (1779 - 1841), to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778 - 1850), a well-known chemist at that time, and to André-Marie Ampère (1775 - 1836). On 29 November 1813, Dersormes and Clément made public Courtois’ discovery. They described the substance to a meeting of the Imperial Institute of France. On December 6, Gay-Lussac announced that the new substance was either an element or a compound of oxygen. Ampère had given some of his sample to Humphry Davy (1778 - 1829). Davy did some experiments on the substance and noted its similarity to chlorine. Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element. A large argument erupted between Davy and Gay-Lussac over who identified iodine first but both scientists acknowledged Barnard Courtois as the first to isolate the chemical element.

[edit] Notable inorganic iodine compounds

See also iodine compounds

[edit] Stable iodine in biology

Iodine is an essential trace element; its only known roles in 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 molecule called thryroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. The thyroid gland actively absorbs iodide ion 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.

Thyroid hormones play a very basic role in biology, acting on gene transcription to regulate the basal metabolic rate. The total deficiency of thyroid hormones can reduce basal metabolic rate up to 50%, while in excessive production of thyroid hormones the basal metabolic rate can be increased by 100%. T4 acts largely as a precursor to T3, which is (with some minor exceptions) the biologically active hormone.

[edit] 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.[citation needed] Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil.[citation needed] Salt for human consumption is often fortified with iodine and is referred to as iodized salt.

[edit] Iodine deficiency

Main article: Iodine deficiency

In areas where there is little iodine in the diet—typically remote inland areas and semi-arid equatorial climates where no marine foods are eaten—iodine deficiency gives rise to hypothyroidism, symptoms of which are extreme fatigue, goitre, mental slowing, depression, weight gain, and low basal body temperatures.[citation needed]

Iodine deficiency is also the leading cause of preventable mental retardation, an effect which happens primarily when babies and small children are made hypothyroid by lack of the element. The addition of iodine to table salt has largely eliminated this problem in the wealthier nations, but iodine deficiency remains a serious public health problem in the developing world.[citation needed]

[edit] 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.

[edit] Radioiodine and biology

[edit] Radioiodine and the thyroid

The artificial radioisotope 131I (a beta emitter), also known as radioiodine which has a half-life of 8.0207 days, 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 Grave's disease). The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO3).

129I (half-life 15.7 million years) is a product of 130Xe spallation in the atmosphere and uranium and plutonium fission, both in subsurface rocks and nuclear reactors. Nuclear processes, in particular nuclear fuel reprocessing and atmospheric nuclear weapons tests have now swamped the natural signal for this isotope. 129I was used in rainwater studies following the Chernobyl accident. It also has been used as a ground-water tracer and as an indicator of nuclear waste dispersion into the natural environment.

If humans are exposed to radioactive iodine, the thyroid gland will absorb it as if it were non-radioactive iodine, leading to elevated chances of thyroid cancer. Isotopes with shorter half-lives such as 131I 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.

[edit] Radioiodine and the kidney

In the 1970s imaging techniques were developed in California to utilize radioiodine in diagnostics for renal hypertension.

[edit] Non-hormone-related applications of iodine

[edit] Precautions for stable iodine

Direct contact with skin can cause lesions, so it should be handled with care. Iodine vapor is very irritating to the eye and to mucous membranes. Concentration of iodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average). When mixed with ammonia, it can form nitrogen triiodide which is extremely sensitive and can explode unexpectedly.

[edit] Clandestine use

In the United States, the Drug Enforcement Agency (DEA) regards iodine and compounds containing iodine (ionic iodides, iodoform, ethyl iodide, and so on) as reagents useful for the clandestine manufacture of methamphetamine. Persons who attempt to purchase significant quantities of such chemicals without establishing a legitimate use are likely to find themselves the target of a DEA investigation. Persons selling such compounds without doing due diligence to establish that the materials are not being diverted to clandestine use may be subject to stiff penalties, such as expensive fines or even imprisonment.[6][7]