Iodine

53

tellurium ← iodine → xenon


Br
↑
I
↓
At




periodic table




General


Name, Symbol, Number

iodine, I, 53


Chemical series
halogens


Group, Period, Block

17, 5, p


Appearance
violet-dark gray, lustrous 



Atomic mass
126.90447(3) g/mol


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/cm³


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


Heat of vaporization
(I2) 41.57 kJ/mol


Heat capacity
(25 °C) (I2) 54.44 J/(mol·K)




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 (Pauling scale)


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·K)


Bulk modulus
7.7 GPa


CAS registry number
7553-56-2


Notable isotopes




Main article: Isotopes of iodine


iso

NA

half-life

DM

DE (MeV)

DP


127I

100%

I is stable with 74 neutrons


129I

syn

1.57×107y

Beta-

0.194

129Xe


131I

syn

8.02070 d

Beta-

0.971

131Xe


Iodine (chemical symbol I, atomic number 53) is a nonmetal that belongs to a group of chemical elements known as halogens. At ordinary temperatures and pressures, it is a dark-gray/purple-black solid that readily sublimes—that is, it goes directly from the solid phase to the gas phase. The gas is purple-pink in color and has an irritating odor. The name iodine was coined from the Greek word iodes, meaning "violet."
Elemental iodine is corrosive on the skin and toxic if ingested. In the form of iodide ions, however, iodine is required as a trace element for most living organisms. In humans, the deficiency or excess of iodide ions can lead to swelling and malfunctioning of the thyroid gland.

Contents

1 Occurrence
2 Discovery
3 Isolation
4 Notable characteristics

4.1 Isotopes


5 Compounds
6 Biological role of iodine

6.1 Iodine deficiency
6.2 Dietary intake
6.3 Toxicity of iodine


7 Applications

7.1 Uses of nonradioactive iodine
7.2 Uses of radioactive iodine


8 See also
9 References
10 External links
11 Credits



Iodine and its compounds have a variety of applications. For instance, tincture of iodine is used to disinfect wounds and sanitize water for drinking. Silver iodide is used in photography, and tungsten iodide is used to stabilize filaments in light bulbs. A number of organic compounds containing iodine are useful in the preparation of pharmaceuticals and dyes. The radioactive isotopes iodine-123 and iodine-125 are used as probes for imaging the thyroid and evaluating its health. An artificial radioactive isotope, iodine-131, is used for the treatment of thyroid cancer.

Occurrence
Iodine occurs in nature in the form of iodide ions, chiefly in solution in seawater but also in some minerals and soils. Although the element is quite rare, it is concentrated in kelp and some other plants, which help introduce the element into the food chain and keep its cost down.

Discovery
Iodine was discovered in 1811 by the Frenchman Bernard Courtois, when he was working with his father to manufacture saltpeter (potassium nitrate). At the time, Napoleon's army was engaged in war and saltpeter, a key component of gunpowder, was in great demand. The process of producing saltpeter from French niter beds required sodium carbonate, which could be isolated from the ashes of burned seaweed. The remaining waste was destroyed by adding sulfuric acid. One day, Courtois added too much sulfuric acid and a cloud of purple vapor arose. The vapor condensed on cold surfaces to produce dark crystals.
Courtois performed a few experiments that led him to suspect that this was a new element, but he lacked the funds to pursue his observations. He therefore gave samples of the material to his friends, Charles Bernard Désormes (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 November 29, 1813, Désormes and Clément made public Courtois’ discovery, describing 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). When Davy experimented with the substance, he 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 major argument erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the substance.

Isolation
There are several methods of isolating iodine.

One commercial method is to treat a suitable batch of natural brine with chlorine gas and to flush the solution with air. The chlorine oxidizes iodide ions in the brine to generate elemental iodine. The reaction can be written as follows.
2I− + Cl2 → I2 + 2Cl−
To obtain iodine on a small scale, solid sodium iodide (NaI) may be reacted with concentrated sulfuric acid (H2SO4). At first, hydrogen iodide (HI) gas is formed, which is oxidized by the acid to produce iodine and sulfur dioxide in gaseous form.
The element iodine may be prepared in an ultrapure form by reacting potassium iodide with copper(II) sulfate.
Notable characteristics
In the periodic table, iodine is located in group 17 (former group 7A), the halogen family, between bromine and astatine. In addition, it lies in period 5, between tellurium and xenon. The molecular formula of iodine is I2.
Chemically, iodine forms compounds with many elements, but it is the least reactive of the halogens. In addition, it is the most electropositive halogen after astatine and has some metallic properties.
It is only slightly soluble in water, giving a yellow solution, but it readily dissolves in chloroform, carbon tetrachloride, or carbon disulphide to form purple solutions. The free element forms a deep blue complex with starch.
Iodine does not react with oxygen or nitrogen, but with ozone it forms an unstable oxide, I4O9. When mixed with water, it reacts to produce hypoiodite ions (OI−). Under appropriate conditions, iodine reacts with other halogens—fluorine, chlorine, and bromine—to produce "interhalogen" compounds, including IF3, IF5, IF7, ICl, I2Cl6, and BrI. When mixed with ammonia, iodine can form nitrogen triiodide, which is extremely sensitive and can explode unexpectedly.
The most common compounds of iodine are the iodides of sodium and potassium (NaI, KI) and the iodates (NaIO3, KIO3).

Isotopes
There are 37 isotopes of iodine, of which only one, 127I, is stable.
The isotope 129I, with a half-life 15.7 million years, is produced in the Earth's atmosphere when the nuclei of 130Xe are struck by high-energy cosmic rays. It is also produced by the fission of uranium and plutonium, in both subsurface rocks and nuclear reactors. Nuclear fuel reprocessing and atmospheric nuclear weapons tests have now swamped the natural signal for this isotope.
In hydrologic studies, 129I concentrations are usually reported as the ratio of 129I to total I (which is virtually all 127I). This ratio in nature is quite small, ranging from 10−14 to 10−10 (peak thermonuclear 129I/I during the 1960s and 1970s reached about 10−7). 129I occurs in multiple ionic forms (commonly, I− and IO3−) and readily enters the biosphere, becoming incorporated into vegetation, soil, milk, and animal tissue.
Excess quantities of stable 129Xe in meteorites appear to have resulted from the decay of "primordial" 129I produced by the supernovas that created the dust and gas from which the solar system formed. The decay of 129I is the basis for the iodine-xenon radiometric dating scheme, which covers the first 50 million years of development of the solar system.

Compounds
A wide range of organic and inorganic compounds contain iodine. In the case of organic compounds, chemists can replace hydrogen atoms with iodine atoms, thus creating many new products.
A list of notable inorganic compounds of iodine is given below, in alphabetical order.

Ammonium iodide (NH4I)
Cesium iodide (CsI)
Copper(I) iodide (CuI)
Hydroiodic acid (HI)
Iodic acid (HIO3)
Iodine cyanide (ICN)
Iodine heptafluoride (IF7)
Iodine pentafluoride (IF5)
Lead(II) iodide (PbI2)
Lithium iodide (LiI)
Nitrogen triiodide (NI3)
Potassium iodate (KIO3)
Potassium iodide (KI)
Sodium iodate (NaIO3)
Sodium iodide (NaI)
Biological role of iodine
Iodine is an essential trace element in the human body. The thyroid hormones thyroxine (T4) and triiodothyronine (T3) contain four and three atoms of iodine per molecule, respectively. The thyroid actively absorbs elemental iodine from the blood to make and release these hormones into the blood, actions that are regulated by a second hormone (thyroid-stimulating hormone, TSH) from the pituitary.
Thyroid hormones are phylogenetically very old, as they are sythesized by most multicellular organisms and even have some effect on unicellular organisms. These hormones play a very basic role in biology, acting on mitochondria to regulate metabolism. T4 acts largely as a precursor to T3, which is (with some minor exceptions) the biologically active hormone.

Iodine deficiency
Iodine deficiency is a serious problem in various parts of the globe. It particularly affects people in places 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 leads to goiter. It is also the leading cause of preventable mental retardation.

Dietary intake
The U.S. Food and Drug Administration recommends an intake of 150 micrograms of iodine per day for both men and women. This is necessary for proper production and functioning of thyroid hormones. Natural sources of iodine include seaweed and seafood.

Toxicity of iodine
If elemental iodine (I2) comes in direct contact with the skin, it can cause lesions; so it should be handled with care. In addition, iodine vapor is very irritating to the eyes and mucous membranes. Concentration of iodine in the air should not exceed 1 milligram per cubic meter.
Elemental iodine is mildly toxic if ingested in small amounts and extremely poisonous if taken in high doses. Consumption of 2–3 grams of it is fatal for humans.
In the human body, excess iodine (in the form of iodide ions) produces symptoms similar to those of iodine deficiency. Common symptoms are abnormal growth of the thyroid gland and disorders in the growth and functioning of the organism as a whole.
If a person is exposed to radioactive iodine (radioiodine), the thyroid gland absorbs it as if it were nonradioactive iodine, raising the chances of thyroid cancer. Radioactive isotopes with shorter half-lives (such as 131I) generate more radiation per unit time and present a greater risk than those with longer half-lives. By taking relatively large amounts of regular iodine, one can saturate the thyroid and prevent uptake of radioiodine.
Applications
Uses of nonradioactive iodine
To combat iodine deficiency, table salt is often enriched with iodine, by adding small amounts of sodium iodide, potassium iodide, or potassium iodate. The product is referred to as iodized salt. Compounds of iodine may also be added to other foodstuffs, such as flour.
Tincture of iodine (3 percent elemental iodine in a water/ethanol base) is an essential component of emergency survival kits, used to disinfect wounds and sanitize surface water for drinking. (To sanitize water for drinking, add 3 drops of the tincture per liter of water and let stand for 30 minutes.) Alcohol-free iodine solutions, such as Lugol's iodine, and other iodine-providing antiseptics (iodophors) are also available as effective sources of elemental iodine for this purpose.
Silver iodide is used in photography.
Tungsten iodide is used to stabilize filaments in light bulbs.
Organic compounds containing iodine are useful in the preparation of pharmaceuticals and dyes.
Potassium iodide (in the form of KI tablets or "super-saturated KI" liquid drops) can be given to people in a nuclear disaster area, to flush out radioactive iodine-131 (a fission product) from the body. Alternatively, iodine pills may be distributed to persons living close to nuclear establishments, for use in case of accidents that could lead to the release of radioiodine.
Uses of radioactive iodine
Iodine-123 and iodine-125 are used in medicine as tracers for imaging the thyroid gland and evaluating how well it is functioning.
Iodine-129 was used in rainwater studies following the nuclear reactor accident at Chernobyl. It has also been used as a groundwater tracer and an indicator of nuclear waste dispersion into the natural environment.
Iodine-131 is an artificial radioisotope used for the treatment of thyroid cancer and other diseases of the thyroid gland.
See also
Chemical element
Periodic table
Periodic table, main group elements
ReferencesISBN links support NWE through referral fees
U. S. Dept. of Engergy Los Alamos National Laboratory - Iodine Retrieved November 6, 2007.
Mark Winter. The University of SheffieldWebElements.com - Iodine Retrieved November 6, 2007.
The Wellness Directory The History of Iodine Retrieved November 6, 2007.
21 CFR 101.9 (c)(8)(iv) (Text PDF) — Food and Drug Administration nutritional facts label information for vitamins and minerals. Retrieved November 6, 2007.
External links
All links retrieved March 5, 2018.

ChemicalElements.com - Iodine
who.int - WHO Global Database on Iodine Deficiency Iodine Status Worldwide.
Network for Sustained Elimination of Iodine Deficiency
Organic Chemistry portal Oxidizing Agents > Iodine

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