recommended daily allowance

Iodine

[ahy-uh-dahyn, -din; in Chem. also ahy-uh-deen]

Iodine (or /ˈaɪədiːn/; from 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 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.

Properties

Iodine under standard conditions is a dark-purple/dark-brown solid. It can be seen apparently sublimating 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 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 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.

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 of the Napoleonic Wars, France was at war and 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 physicist 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 Courtois as the first to isolate the chemical element.

Applications

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.
Iodine, as a heavy element, is quite radio-opaque. Organic compounds of a certain type (typically iodine-substituted benzene derivatives) are thus used in medicine as X-ray radiocontrast agents for intravenous injection. This is often in conjunction with advanced X-ray techniques such as angiography and CT scanning
Silver iodide is used in photography.
Tungsten iodide is used to stabilize the filaments in light bulbs.
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 a 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 other algae have some ability to concentrate iodine, which helps introduce the element into the food chain.

Sources

Iodine is found in the mineral caliche, found in Chile, between the Andes and the sea. It can also be found in some seaweeds (algae), as well as extracted from seawater.

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,a kind of brown alga. 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

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)

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). 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:

2I⁻ + 4H⁺ + MnO₂ → I₂ + 2H₂O + Mn²⁺

Iodine is reduced to hydroiodic acid by hydrogen sulfide:

I₂ + H₂S → 2HI + S↓

or by hydrazine:

2I₂ + N₂H₄ → 4HI + N₂

Iodine is oxidized to iodate by nitric acid:

I₂ + 10HNO₃ → 2HIO₃ + 10NO₂ + 4H₂O

or by chlorates:

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):

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

Notable inorganic iodine compounds

Stable iodine in biology

Iodine is an essential trace element, the heaviest element known to be needed by living organisms. Its main role in animal biology is 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 an iodine-containing enzyme 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.

Thyroid hormones play a 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 minor exceptions) the biologically active hormone.

Iodine accounts for 65% of the molecular weight of T4 and 59% of the T3. 15-20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of the body's iodine is distributed in other tissues, including mammary glands, eyes, gastric mucosa, the cervix, and salivary glands. Its role in mammary tissue is related to fetal and neonatal development, but its role in the other tissues is unknown. It has been shown to act as an antioxidant in these tissues.

Iodine may have a relationship with selenium, and iodine supplementation in selenium-deficient populations may pose risks for thyroid function.

Human dietary intake

The United States Recommended Daily Allowance (RDA) is 150 micrograms per day (μg/day) for both men and women, with a Tolerable Upper Intake Level (UL) for adults is 1,100 μg/day (1.1 mg/day). The tolerable upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.

Natural sources of iodine include sea life, such as kelp and certain seafood, as well as plants grown on iodine-rich soil. Iodized salt is fortified with iodine.

As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and 190 to 210 μg/day for women. In Japan, consumption is much higher due to the frequent consumption of seaweed or kombu kelp. Estimates range from 5,280 to 13,800 μg/day. Although some Chinese data associates excess iodine with autoimmune thyroiditis and hypothyroidism, these effects have not been observed in Japanese populations, and a protective effect on breast cancer has been hypothesized.

Donald W. Miller believes that the daily FDA intake recommendation may be 100 times too low.

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.

Iodine deficiency is the leading cause of preventable mental retardation, a result which occurs primarily when babies or small children are rendered hypothyroidic by a lack of the element. The addition of iodine to table salt has largely eliminated this problem in the wealthier nations, but as of March 2006, iodine deficiency remained a serious public health problem in the developing world. Iodine deficiency is also a problem in certain areas of Europe. In Germany it has been estimated to cause a billion dollars in healthcare costs per year.

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.

Iodine-123 and iodine-125 are used in medicine as tracers for imaging and evaluating the function of the thyroid.
Noncombined (elemental) iodine is mildly toxic to all living things.

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. 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.

Iodine-129 ¹²⁹I (half-life 15.7 million years) is a product of cosmic ray spallation on various isotopes of xenon in the atmosphere, in cosmic ray muon interaction with tellurium-130, and also 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 groundwater tracer and as an indicator of nuclear waste dispersion into the natural environment.

Radioiodine and the kidney

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

Isotopes

There are 37 isotopes of iodine, but 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 halflife 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" 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 85 million years of solar system evolution.

Effects of various radioiodine isotopes in biology are discussed above.

Toxicity

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 are consumed, it is fatal to humans. Iodides are similar in toxicity to bromides.

Precautions

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.

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.