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ITU-T

The International Telecommunication Union Telecommunication Standardization Sector (ITU-T) is one of the three Sectors (branches) of the International Telecommunication Union (ITU). It is responsible for coordinating standards for telecommunications and Information Communication Technology, such as X.509 for cybersecurity, Y.3172 and Y.3173 for machine learning, and H.264/MPEG-4 AVC for video compression, between its Member States, Private Sector Members, and Academia Members.

The World Telecommunication Standardization Assembly (WTSA), the sector's governing conference, convenes every four years.

ITU-T has a permanent secretariat called the Telecommunication Standardization Bureau (TSB), which is based at the ITU headquarters in Geneva, Switzerland. The current director of the TSB is Seizo Onoe (of Japan), whose 4-year term commenced on 1 January 2023. Seizo Onoe succeeded Chaesub Lee of South Korea, who was director from 1 January 2015 until 31 December 2022.

Primary function
The ITU-T mission is to ensure the efficient and timely production of standards covering all fields of telecommunications and Information Communication Technology (ICTs) on a worldwide basis, as well as defining tariff and accounting principles for international telecommunication services.

The international standards that are produced by the ITU-T are referred to as "Recommendations" (with the word capitalized to distinguish its meaning from the common parlance sense of the word "recommendation"), as they become mandatory only when adopted as part of a national law.

Since the ITU-T is part of the ITU, which is a United Nations specialized agency, its standards carry more formal international weight than those of most other standards development organizations that publish technical specifications of a similar form.

History
At the initiative of Napoleon III, the French government invited international participants to a conference in Paris in 1865 to facilitate and regulate international telegraph services. A result of the conference was the founding of the forerunner of the modern ITU.

At the 1925 Paris conference, the ITU created two consultative committees to deal with the complexities of the international telephone services, known as CCIF (as the French acronym) and with long-distance telegraphy CCIT (Comité Consultatif International des Communications Téléphoniques a grande distance).

In view of the basic similarity of many of the technical problems faced by the CCIF and CCIT, a decision was taken in 1956 to merge them into a single entity, the International Telegraph and Telephone Consultative Committee (CCITT, in Comité Consultatif International Téléphonique et Télégraphique). The first Plenary Assembly of the new organization was held in Geneva, Switzerland in December 1956.

In 1992, the Plenipotentiary Conference (the top policy-making conference of ITU) saw a reform of ITU, giving the Union greater flexibility to adapt to an increasingly complex, interactive and competitive environment. The CCITT was renamed the Telecommunication Standardization Sector (ITU-T), as one of three Sectors of the Union alongside the Radiocommunication Sector (ITU-R) and the Telecommunication Development Sector (ITU-D).

Historically, the Recommendations of the CCITT were presented at plenary assemblies for endorsement, held every four years, and the full set of Recommendations were published after each plenary assembly. However, the delays in producing texts, and translating them into other working languages, did not suit the fast pace of change in the telecommunications industry.

"Real time" standardization
The rise of the personal computer industry in the early 1980s created a new common practice among both consumers and businesses of adopting "bleeding edge" communications technology even if it was not yet standardized. Thus, standards organizations had to put forth standards much faster, or find themselves ratifying de facto standards after the fact. One of the most prominent examples of this was the Open Document Architecture project, which began in 1985 when a profusion of software firms around the world were still furiously competing to shape the future of the electronic office, and was completed in 1999 long after Microsoft Office's then-secret binary file formats had become established as the global de facto standard.

The ITU-T now operates under much more streamlined processes. The time between an initial proposal of a draft document by a member company and the final approval of a full-status ITU-T Recommendation can now be as short as a few months (or less in some cases). This makes the standardization approval process in the ITU-T much more responsive to the needs of rapid technology development than in the ITU's historical past. New and updated Recommendations are published on an almost daily basis, and nearly all of the library of over 3,270 Recommendations is now free of charge online. (About 30 specifications jointly maintained by the ITU-T and ISO/IEC are not available for free to the public. )

ITU-T has moreover tried to facilitate cooperation between the various forums and standard-developing organizations (SDOs). This collaboration is necessary to avoid duplication of work and the consequent risk of conflicting standards in the market place.

In the work of standardization, ITU-T cooperates with other SDOs, e.g., the International Organization for Standardization (ISO) and the Internet Engineering Task Force (IETF).

Development of Recommendations
Most of the work of ITU-T is carried out by its Sector Members and Associates, while the Telecommunication Standardization Bureau (TSB) is the executive arm of ITU-T and coordinator for a number of workshops and seminars to progress existing work areas and explore new ones. The events cover a wide array of topics in the field of information and communication technologies (ICT) and attract high-ranking experts as speakers, and attendees from engineers to high-level management from all industry sectors.

The technical work, the development of Recommendations, of ITU-T is managed by Study Groups (SGs), such as Study Group 13 for network standards, Study Group 16 for multimedia standards, and Study Group 17 for security standards, which are created by the World Telecommunication Standardization Assembly (WTSA) which is held every four years. As part of the deliberations, WTSA has instructed ITU to hold the Global Standards Symposium, which unlike WTSA is open to public for participation. The people involved in these SGs are experts in telecommunications from all over the world. There are currently 11 SGs. Study groups meet face to face (or virtually under exceptional circumstances) according to a calendar issued by the TSB. SGs are augmented by Focus Groups (FGs), an instrument created by ITU-T, providing a way to quickly react to ICT standardization needs and allowing great flexibility in terms of participation and working methods. The key difference between SGs and FGs is that the latter have greater freedom to organize and finance themselves, and to involve non-members in their work, but they do not have the authority to approve Recommendations. Focus Groups can be created very quickly, are usually short-lived and can choose their own working methods, leadership, financing, and types of deliverables. Current Focus Groups include the ITU-WHO Focus Group on Artificial Intelligence for Health (FG-AI4H) as well as Machine Learning for 5G (which developed Y.3172), Quantum Information Technologies for Networks, and Artificial Intelligence for Assisted and Autonomous Driving.

Alternative Approval Process
The Alternative Approval Process (AAP) is a fast-track approval procedure that was developed to allow standards to be brought to market in the timeframe that industry now demands. The AAP is defined in ITU-T Recommendation A.8.

This dramatic overhaul of standards-making by streamlining approval procedures was implemented in 2001 and is estimated to have cut the time involved in this critical aspect of the standardization process by 80 to 90 percent. This means that an average standard that took around four years to approve and publish until the mid nineties, and two years until 1997, can now be approved in an average of two months, or as little as five weeks.

Besides streamlining the underlying procedures involved in the approval process, an important contributory factor to the use of AAP is electronic document handling. Once the approval process has begun the rest of the process can be completed electronically, in the vast majority of cases, with no further physical meetings.

The introduction of AAP also formalizes public/private partnership in the approval process by providing equal opportunities for both sector members and member states in the approval of technical standards.

A panel of SG experts drafts a proposal that is then forwarded at an SG meeting to the appropriate body which decides if it is sufficiently ready to be designated a draft text and thus gives its consent for further review at the next level.

After this Consent has been given, TSB announces the start of the AAP procedure by posting the draft text to the ITU-T website and calling for comments. This gives the opportunity for all members to review the text. This phase, called last call, is a four-week period in which comments can be submitted by member states and sector members.

If no comments other than editorial corrections are received, the Recommendation is considered approved since no issues were identified that might need any further work. However, if there are any comments, the SG chairman, in consultation with TSB, sets up a comment resolution process by the concerned experts. The revised text is then posted on the web for an additional review period of three weeks.

Similar to the last call phase, in additional review the Recommendation is considered as approved if no comments are received. If comments are received, it is apparent that there are some issues that still need more work, and the draft text and all comments are sent to the next Study Group meeting for further discussion and possible approval.

Those Recommendations considered as having policy or regulatory implications are approved through what is known as the Traditional Approval Process (TAP), which allows a longer period for reflection and commenting by member states. TAP Recommendations are also translated into the six working languages of ITU (Arabic, Chinese, English, French, Russian, and Spanish).

Series and Recommendations
ITU-T Recommendations are the names given to telecommunications and computer protocol specification documents published by ITU-T.

Recommendation categorization
ITU-T assigns each Recommendation a name based on the series and Recommendation number. The name starts with the letter of the series the Recommendation belongs to. Each series encompasses a broad category of Recommendations, such as "H-Series Recommendations: Audiovisual and multimedia systems". The series letter is followed by a period and the Recommendation number, which uniquely identifies the Recommendation within the series. Often, a range of related Recommendations are further grouped within the series and given adjacent numbers, such as "H.200-H.499: Infrastructure of audiovisual services" or "H.260-H.279: Coding of moving video". Many numbers are "skipped" to give room for future Recommendations to be adjacent to related Recommendations. Recommendations can be revised or "superseded" and keep their existing Recommendation number.

Key standards published by ITU

 * ASN.1 (Abstract Syntax Notation One)
 * Coding of audio: G.711 and G.72x series
 * Coding of still images: JPEG T.8x, JPEG 2000 T.80x and JPEG XR T.83x series
 * Coding of video: H.120, H.261, H.262 (MPEG-2 Video), H.263, H.264 (MPEG-4 AVC), H.265 (HEVC), H.266 (VVC)
 * Construction, installation and protection of cables and other elements of outside plant, L-series
 * Data communication over the telephone network, V-series
 * Fax standards T.2 – T.4, T.30, T.37, T.38
 * F.780.2
 * G.114 is an ITU Recommendation that addresses acceptable delays for voice applications, is oriented to national telecommunications, and is more stringent than what is normally applied in private voice networks. One of the issues it addresses is voice delay. It states that 150 milliseconds is acceptable, while 400 milliseconds is unacceptable.
 * G.191: Software tools for speech and audio coding standardization
 * G.651.1 Multi-mode optical fiber (MMF)
 * G.652, G.655, G.657 Single-mode optical fiber (SMF)
 * G.8201 Error performance parameters and objectives for multi-operator international paths within optical transport networks
 * G.hn (Next generation wired home networking over powerlines, phonelines and coaxial cable)
 * Green ICT standards and supplements (L.1000 series)
 * H.323 family of standards for multimedia and VoIP
 * H.810 Continua Design Guidelines
 * H.870 safe listening standard (audio) developed together with the World Health Organization
 * Interconnection rate harmonization, D-series (Secretary: Simao Campos)
 * International Emergency Preference Scheme E.106
 * International telecommunication charge card E.118
 * IMSI codes used in SIM cards E.212
 * ISDN and PSTN/3G videoconferencing systems, H.320 and H.324
 * ISDN (Integrated Services Digital Network) Q.931
 * Objective voice quality measurements P.861, P.862, P.863
 * Open Systems Interconnection
 * Optical Transport Network (OTN) G.709, G.798, G.872
 * Passive optical networks (PON) G.983, G.984, G.987
 * Public key infrastructure (PKI) X.509
 * Public telecommunication numbering plan, E.164
 * Security framework X.805
 * Signalling System 7 Q.7xx series
 * Standards relating to quality of service (QoS) and quality of experience (QoE)
 * Specification and Description Language
 * Synchronous Digital Hierarchy (SDH) G.707, G.783, G.803
 * T.192, Cooperative document handling
 * TrustworthyAI
 * Wavelength-division multiplexing (WDM)
 * X.25, packet switching
 * X.400 email
 * Y.3800 Quantum Key Distribution
 * (x)DSL (Digital Subscriber Line) series of standards for broadband telecoms
 * G.8013/Y.1731 Operations, administration and maintenance (OAM) functions and mechanisms for Ethernet-based networks
 * Machine Learning for 5G and future networks (IMT2020):
 * Y.3172
 * Y.3173
 * Y.3176
 * Y.3181

International Telecommunication Regulations (ITRs)
In addition to the ITU-T Recommendations, which have non-mandatory status unless they are adopted in national laws, ITU-T is also the custodian of a binding international treaty, the International Telecommunication Regulations. The ITRs go back to the earliest days of the ITU when there were two separate treaties, dealing with telegraph and telephone. The ITRs were adopted, as a single treaty, at the World Administrative Telegraphy and Telephone Conference held in Melbourne, 1988 (WATTC-88).

The ITRs comprise ten articles which deal, inter alia, with the definition of international telecommunication services, cooperation between countries and national administrations, safety of life and priority of telecommunications and charging and accounting principles. The adoption of the ITRs in 1988 is often taken as the start of the wider liberalization process in international telecommunications, though a few countries, including United States and United Kingdom, had made steps to liberalize their markets before 1988.

The Constitution and Convention of ITU provides for the amendment of ITRs through a World Conference on International Telecommunications (WCIT). Accordingly, in 1998 there began a process of review of the ITRs; and in 2009 extensive preparations began for such a conference, WCIT-12. In addition to "regional preparatory meetings", the ITU Secretariat developed 13 "Background Briefs on key issues" that were expected to be discussed at the conference. Convened by former ITU secretary-general Hamadoun Touré, the Conference, WCIT-12, was then held in Dubai, United Arab Emirates, during the period 3–14 December 2014.

AI for Good
The Standardization Sector of ITU also organizes AI for Good, the United Nations platform for the sustainable development of Artificial Intelligence.

Hot topics

 * ITU-T has expressed a commitment to "bridging the standardization gap" – disparities in the ability of developing countries, relative to developed ones, to access, implement, contribute to and influence international ICT standards.
 * The ICT Security Standards Roadmap has been developed to assist in the development of security standards by bringing together information about existing standards and current standards work in key standards development organizations.
 * The Next Generation Networks (NGN) concept takes into consideration new realities in the telecommunication industry characterized by factors such as; the need to converge and optimize the operating networks and the extraordinary expansion of digital traffic (i.e., increasing demand for new multimedia services, mobility, etc.).

title: ITU-T


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Indian

Indian or Indians may refer to:
 * Something or someone of, from, or associated with the nation of India

 * Something or someone of, from, or associated with the indigenous peoples of the Americas.

South Asia

 * Indian people, people of Indian nationality, or people who have an Indian ancestor
 * Non-resident Indian, a citizen of India who has temporarily emigrated to another country
 * South Asian ethnic groups, referring to people of the Indian subcontinent, as well as the greater South Asia region prior to the 1947 partition of India
 * Anglo-Indians, people with mixed Indian and British ancestry, or people of British descent born or living in the Indian subcontinent
 * Bombay East Indians, a Christian community in India

Europe

 * British Indians, British people of Indian origin

The Americas

 * Indo-Canadians, Canadian people of Indian origin
 * Indo-Caribbeans, Caribbean people of Indian origin; mainly:
 * Indo-Trinidadian and Tobagonian
 * Indo-Guyanese
 * Indo-Surinamese
 * Indo-Jamaicans
 * Indian Americans, American people of Indian origin
 * Indigenous peoples of the Americas, the pre-Columbian inhabitants of the Americas and their descendants
 * Plains Indians, the common name for the Native Americans who lived on the Great Plains of North America
 * Native Americans in the United States, the indigenous people in the United States
 * Native American tribes, specific groups of Native Americans
 * Indigenous peoples in Canada
 * First Nations in Canada, the various Indigenous peoples in Canada who are neither Inuit nor Métis
 * Indigenous peoples of South America, peoples living in South America in the pre-Columbian era and their descendants
 * Native Mexicans, indigenous people of Mexico
 * Indigenous peoples of Central America
 * Indigenous peoples of the Caribbean
 * West Indians, people from the Caribbean region and the Lucayan Archipelago
 * Mardi Gras Indians, African-American Carnival revelers in New Orleans, Louisiana, whose suits are influenced by Native American ceremonial apparel

Australia

 * Aboriginal Australians, called "Indians" until the 19th century
 * Indian Australians

Languages

 * Languages of India, including Indo-Aryan languages and Dravidian languages
 * Indian English, a dialect of the English language used in India
 * Indigenous languages of the Americas, spoken by indigenous peoples from Alaska and Greenland to the southern tip of South America

Places

 * Indian, West Virginia, a former unincorporated community in Kanawha County
 * The Indians, an islet group in the British Virgin Islands
 * Indian Creek (disambiguation)
 * Indian Island (disambiguation)
 * Indian Peak, a mountain in British Columbia
 * Indian River (disambiguation), several rivers and communities
 * Indian Run (disambiguation), streams in the U.S. states of Pennsylvania and West Virginia
 * Indian subcontinent
 * Indian Ocean

Arts, entertainment, and media

 * Indian cinema

Films

 * Indian (film series), an Indian Tamil-language film series starring Kamal Haasan
 * Indian (1996 film)
 * Indian 2, 2024 sequel
 * Indian 3, 2025 sequel
 * Indian (2001 film), an Indian Hindi film

Music

 * Indians (musician), moniker of Danish singer Søren Løkke Juul accompanied by some musicians also collectively known as Indians
 * "Indian" (song), by Sturm und Drang
 * Indian (soundtrack), an album from the 1996 film
 * "Indians" (song), by Anthrax
 * Indians, a song by Gojira from their 2003 album The Link

Other arts, entertainment, and media

 * Indian (card game), a simple card game that involves strategy
 * Indian soap opera, soap operas written, produced, and filmed in India
 * Indians (play), a 1968 play by Arthur Kopit
 * Indians (sculpture), a name for The Bowman and The Spearman, sculptures by Ivan Meštrović

Businesses

 * Indian (airline), a now-defunct state-owned airline of India, merged with Air India
 * Indian Motocycle Manufacturing Company, an American company later called Indian Manufacturing Company, that produced the brand "Indian motorcycles" from 1901 to 1953

Sport

 * Akwesasne Indians, a Mohawk lacrosse team
 * Calcutta Indians, pre independent football team of India that briefly played against international clubs and teams.
 * Cleveland Indians, former name of the Cleveland Guardians, an American professional baseball team based in Cleveland, Ohio, US
 * Frölunda HC or Frölunda Indians, a Swedish professional ice hockey club based in Gothenburg, Sweden
 * Hannover Indians, a German ice hockey club based in Hannover, Germany
 * Indian Arrows, the developmental football team of All India Football Federation
 * Indianapolis Indians, an American minor league baseball team based in Indianapolis, Indiana, US
 * Indios de Mayagüez, a baseball team in the Puerto Rican Professional Baseball League
 * Mumbai Indians, a franchise cricket team representing the city of Mumbai, India in the Indian Premier League
 * Springfield Indians, an American minor professional ice hockey franchise, originally based in West Springfield, Massachusetts, US

Other uses

 * Indian cuisine, a wide variety of regional cuisines native to India
 * Indian Head cent, a coin produced by the United States Mint from 1859 to 1909

license: Creative Commons - Attribution Share-Alike - https://creativecommons.org/licenses/by-sa/4.0/

title: Indian


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Ionic

Ionic or Ionian may refer to:

Arts and entertainment

 * Ionic meter, a poetic metre in ancient Greek and Latin poetry
 * Ionian mode, a musical mode or a diatonic scale

Places and peoples

 * Ionian, of or from Ionia, an ancient region in western Anatolia
 * Ionians, one of four major tribes of the ancient Greeks
 * Ionian Sea, part of the Mediterranean Sea
 * Ionian Islands, a group of islands in Greece

Science

 * Ionian, of or relating to Io, a moon of the planet Jupiter
 * Ionian stage, a proposed name for the now-defined Chibanian stage in stratigraphy.
 * Ionic, of or relating to an ion, an atom or molecule with a net electric charge
 * Ionic bonding, a type chemical bonding
 * Ionic compound, a chemical compound involving ionic bonding

Other uses

 * Ionic Greek, an ancient dialect of the Greek language
 * Ionic (mobile app framework), a software development kit
 * Ionic order, one of the orders of classical architecture
 * Ionian Technologies, an American biotechnology company
 * Hull Ionians, an English rugby club
 * Ionic No. 5, a typeface
 * SS Ionic, the name of two ships of the White Star Line.

title: Ionic


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Indium

Indium is a chemical element; it has symbol In and atomic number 49. It is a silvery-white post-transition metal and one of the softest elements. Chemically, indium is similar to gallium and thallium, and its properties are largely intermediate between the two. It was discovered in 1863 by Ferdinand Reich and Hieronymous Theodor Richter by spectroscopic methods and named for the indigo blue line in its spectrum.

Indium is a technology-critical element used primarily in the production of flat-panel displays as indium tin oxide (ITO), a transparent and conductive coating applied to glass. Indium is also used in the semiconductor industry, in low-melting-point metal alloys such as solders and soft-metal high-vacuum seals. It is produced exclusively as a by-product during the processing of the ores of other metals, chiefly from sphalerite and other zinc sulfide ores.

Indium has no biological role and its compounds are toxic when inhaled or injected into the bloodstream, although they are poorly absorbed following ingestion.

Etymology
The name comes from the Latin word indicum meaning violet or indigo.

Physical
Indium is a shiny silvery-white, highly ductile post-transition metal with a bright luster. It is so soft (Mohs hardness 1.2) that it can be cut with a knife and leaves a visible line like a pencil when rubbed on paper. It is a member of group 13 on the periodic table and its properties are mostly intermediate between its vertical neighbors gallium and thallium. Like tin, a high-pitched cry is heard when indium is bent – a crackling sound due to crystal twinning. Like gallium, indium is able to wet glass. Like both, indium has a low melting point, 156.60 °C (313.88 °F); higher than its lighter homologue, gallium, but lower than its heavier homologue, thallium, and lower than tin. The boiling point is 2072 °C (3762 °F), higher than that of thallium, but lower than gallium, conversely to the general trend of melting points, but similarly to the trends down the other post-transition metal groups because of the weakness of the metallic bonding with few electrons delocalized.

The density of indium, 7.31 g/cm3, is also greater than gallium, but lower than thallium. Below the critical temperature, 3.41 K, indium becomes a superconductor. Indium crystallizes in the body-centered tetragonal crystal system in the space group I4/mmm (lattice parameters: a = 325 pm, c = 495 pm): this is a slightly distorted face-centered cubic structure, where each indium atom has four neighbours at 324 pm distance and eight neighbours slightly further (336 pm). Indium has greater solubility in liquid mercury than any other metal (more than 50 mass percent of indium at 0 °C). Indium displays a ductile viscoplastic response, found to be size-independent in tension and compression. However it does have a size effect in bending and indentation, associated to a length-scale of order 50–100 μm, significantly large when compared with other metals.

Chemical
Indium has 49 electrons, with an electronic configuration of [Kr]4d105s25p1. In compounds, indium most commonly donates the three outermost electrons to become indium(III), In3+. In some cases, the pair of 5s-electrons are not donated, resulting in indium(I), In+. The stabilization of the monovalent state is attributed to the inert pair effect, in which relativistic effects stabilize the 5s-orbital, observed in heavier elements. Thallium (indium's heavier homolog) shows an even stronger effect, causing oxidation to thallium(I) to be more probable than to thallium(III), whereas gallium (indium's lighter homolog) commonly shows only the +3 oxidation state. Thus, although thallium(III) is a moderately strong oxidizing agent, indium(III) is not, and many indium(I) compounds are powerful reducing agents. While the energy required to include the s-electrons in chemical bonding is lowest for indium among the group 13 metals, bond energies decrease down the group so that by indium, the energy released in forming two additional bonds and attaining the +3 state is not always enough to outweigh the energy needed to involve the 5s-electrons. Indium(I) oxide and hydroxide are more basic and indium(III) oxide and hydroxide are more acidic.

A number of standard electrode potentials, depending on the reaction under study, are reported for indium, reflecting the decreased stability of the +3 oxidation state:

 * In2+ + e−|| ⇌ In+ || E0 = −0.40 V
 * In3+ + e−|| ⇌ In2+ || E0 = −0.49 V
 * In3+ + 2 e−|| ⇌ In+ || E0 = −0.443 V
 * In3+ + 3 e−|| ⇌ In || E0 = −0.3382 V
 * In+ + e−|| ⇌ In || E0 = −0.14 V
 * }
 * In3+ + 3 e−|| ⇌ In || E0 = −0.3382 V
 * In+ + e−|| ⇌ In || E0 = −0.14 V
 * }
 * In+ + e−|| ⇌ In || E0 = −0.14 V
 * }

Indium metal does not react with water, but it is oxidized by stronger oxidizing agents such as halogens to give indium(III) compounds. It does not form a boride, silicide, or carbide, and the hydride InH3 has at best a transitory existence in ethereal solutions at low temperatures, being unstable enough to spontaneously polymerize without coordination. Indium is rather basic in aqueous solution, showing only slight amphoteric characteristics, and unlike its lighter homologs aluminium and gallium, it is insoluble in aqueous alkaline solutions.

Isotopes
Indium has 39 known isotopes, ranging in mass number from 97 to 135. Only two isotopes occur naturally as primordial nuclides: indium-113, the only stable isotope, and indium-115, which has a half-life of 4.41 years, four orders of magnitude greater than the age of the Universe and nearly 30,000 times greater than half life of thorium-232. The half-life of 115In is very long because the beta decay to 115Sn is spin-forbidden. Indium-115 makes up 95.7% of all indium. Indium is one of three known elements (the others being tellurium and rhenium) of which the stable isotope is less abundant in nature than the long-lived primordial radioisotopes.

The stablest artificial isotope is indium-111, with a half-life of approximately 2.8 days. All other isotopes have half-lives shorter than 5 hours. Indium also has 47 meta states, among which indium-114m1 (half-life about 49.51 days) is the most stable, more stable than the ground state of any indium isotope other than the primordial. All decay by isomeric transition. The indium isotopes lighter than 113In predominantly decay through electron capture or positron emission to form cadmium isotopes, while the indium isotopes heavier than 113In predominantly decay through beta-minus decay to form tin isotopes.

Indium(III)
Indium(III) oxide, In2O3, forms when indium metal is burned in air or when the hydroxide or nitrate is heated. In2O3 adopts a structure like alumina and is amphoteric, that is able to react with both acids and bases. Indium reacts with water to reproduce soluble indium(III) hydroxide, which is also amphoteric; with alkalis to produce indates(III); and with acids to produce indium(III) salts:

 * In(OH)3 + 3 HCl → InCl3 + 3 H2O

The analogous sesqui-chalcogenides with sulfur, selenium, and tellurium are also known. Indium forms the expected trihalides. Chlorination, bromination, and iodination of In produce colorless InCl3, InBr3, and yellow InI3. The compounds are Lewis acids, somewhat akin to the better known aluminium trihalides. Again like the related aluminium compound, InF3 is polymeric.

Direct reaction of indium with the pnictogens produces the gray or semimetallic III–V semiconductors. Many of them slowly decompose in moist air, necessitating careful storage of semiconductor compounds to prevent contact with the atmosphere. Indium nitride is readily attacked by acids and alkalis.

Indium(I)
Indium(I) compounds are not common. The chloride, bromide, and iodide are deeply colored, unlike the parent trihalides from which they are prepared. The fluoride is known only as an unstable gaseous compound. Indium(I) oxide black powder is produced when indium(III) oxide decomposes upon heating to 700 °C.

Other oxidation states
Less frequently, indium forms compounds in oxidation state +2 and even fractional oxidation states. Usually such materials feature In–In bonding, most notably in the halides In2X4 and [In2X6]2−, and various subchalcogenides such as In4Se3. Several other compounds are known to combine indium(I) and indium(III), such as InI6(InIIICl6)Cl3, InI5(InIIIBr4)2(InIIIBr6), and InIInIIIBr4.

Organoindium compounds
Organoindium compounds feature In–C bonds. Most are In(III) derivatives, but cyclopentadienylindium(I) is an exception. It was the first known organoindium(I) compound, and is polymeric, consisting of zigzag chains of alternating indium atoms and cyclopentadienyl complexes. Perhaps the best-known organoindium compound is trimethylindium, In(CH3)3, used to prepare certain semiconducting materials.

History
In 1863, the German chemists Ferdinand Reich and Hieronymous Theodor Richter were testing ores from the mines around Freiberg, Saxony. They dissolved the minerals pyrite, arsenopyrite, galena and sphalerite in hydrochloric acid and distilled raw zinc chloride. Reich, who was color-blind, employed Richter as an assistant for detecting the colored spectral lines. Knowing that ores from that region sometimes contain thallium, they searched for the green thallium emission spectrum lines. Instead, they found a bright blue line. Because that blue line did not match any known element, they hypothesized a new element was present in the minerals. They named the element indium, from the indigo color seen in its spectrum, after the Latin indicum, meaning 'of India'.

Richter went on to isolate the metal in 1864. An ingot of 0.5 kg was presented at the World Fair 1867. Reich and Richter later fell out when the latter claimed to be the sole discoverer.

Occurrence
Indium is created by the long-lasting (up to thousands of years) s-process (slow neutron capture) in low-to-medium-mass stars (range in mass between 0.6 and 10 solar masses). When a silver-109 atom captures a neutron, it transmutes into silver-110, which then undergoes beta decay to become cadmium-110. Capturing further neutrons, it becomes cadmium-115, which decays to indium-115 by another beta decay. This explains why the radioactive isotope is more abundant than the stable one. The stable indium isotope, indium-113, is one of the p-nuclei, the origin of which is not fully understood; although indium-113 is known to be made directly in the s- and r-processes (rapid neutron capture), and also as the daughter of very long-lived cadmium-113, which has a half-life of about eight quadrillion years, this cannot account for all indium-113.

Indium is the 68th most abundant element in Earth's crust at approximately 50 ppb. This is similar to the crustal abundance of silver, bismuth and mercury. It very rarely forms its own minerals, or occurs in elemental form. Fewer than 10 indium minerals such as roquesite (CuInS2) are known, and none occur at sufficient concentrations for economic extraction. Instead, indium is usually a trace constituent of more common ore minerals, such as sphalerite and chalcopyrite. From these, it can be extracted as a by-product during smelting. While the enrichment of indium in these deposits is high relative to its crustal abundance, it is insufficient, at current prices, to support extraction of indium as the main product.

Different estimates exist of the amounts of indium contained within the ores of other metals. However, these amounts are not extractable without mining of the host materials (see Production and availability). Thus, the availability of indium is fundamentally determined by the rate at which these ores are extracted, and not their absolute amount. This is an aspect that is often forgotten in the current debate, e.g. by the Graedel group at Yale in their criticality assessments, explaining the paradoxically low depletion times some studies cite.

Production and availability
Indium is produced exclusively as a by-product during the processing of the ores of other metals. Its main source material are sulfidic zinc ores, where it is mostly hosted by sphalerite. Minor amounts are also extracted from sulfidic copper ores. During the roast-leach-electrowinning process of zinc smelting, indium accumulates in the iron-rich residues. From these, it can be extracted in different ways. It may also be recovered directly from the process solutions. Further purification is done by electrolysis. The exact process varies with the mode of operation of the smelter.

Its by-product status means that indium production is constrained by the amount of sulfidic zinc (and copper) ores extracted each year. Therefore, its availability needs to be discussed in terms of supply potential. The supply potential of a by-product is defined as that amount which is economically extractable from its host materials per year under current market conditions (i.e. technology and price). Reserves and resources are not relevant for by-products, since they cannot be extracted independently from the main-products. Recent estimates put the supply potential of indium at a minimum of 1,300 t/yr from sulfidic zinc ores and 20 t/yr from sulfidic copper ores. These figures are significantly greater than current production (655 t in 2016). Thus, major future increases in the by-product production of indium will be possible without significant increases in production costs or price. The average indium price in 2016 was US$240/kg, down from US$705/kg in 2014.

China is a leading producer of indium (290 tonnes in 2016), followed by South Korea (195 t), Japan (70 t) and Canada (65 t). The Teck Resources refinery in Trail, British Columbia, is a large single-source indium producer, with an output of 32.5 tonnes in 2005, 41.8 tonnes in 2004 and 36.1 tonnes in 2003.

The primary consumption of indium worldwide is LCD production. Demand rose rapidly from the late 1990s to 2010 with the popularity of LCD computer monitors and television sets, which now account for 50% of indium consumption. Increased manufacturing efficiency and recycling (especially in Japan) maintain a balance between demand and supply. According to the UNEP, indium's end-of-life recycling rate is less than 1%.

Applications
In 1924, indium was found to have a valued property of stabilizing non-ferrous metals, and that became the first significant use for the element. The first large-scale application for indium was coating bearings in high-performance aircraft engines during World War II, to protect against damage and corrosion; this is no longer a major use of the element. New uses were found in fusible alloys, solders, and electronics. In the 1950s, tiny beads of indium were used for the emitters and collectors of PNP alloy-junction transistors. In the middle and late 1980s, the development of indium phosphide semiconductors and indium tin oxide thin films for liquid-crystal displays (LCD) aroused much interest. By 1992, the thin-film application had become the largest end use.

Indium(III) oxide and indium tin oxide (ITO) are used as a transparent conductive coating on glass substrates in electroluminescent panels. Indium tin oxide is used as a light filter in low-pressure sodium-vapor lamps. The infrared radiation is reflected back into the lamp, which increases the temperature within the tube and improves the performance of the lamp.

Indium has many semiconductor-related applications. Some indium compounds, such as indium antimonide and indium phosphide, are semiconductors with useful properties: one precursor is usually trimethylindium (TMI), which is also used as the semiconductor dopant in II–VI compound semiconductors. InAs and InSb are used for low-temperature transistors and InP for high-temperature transistors. The compound semiconductors InGaN and InGaP are used in light-emitting diodes (LEDs) and laser diodes. Indium is used in photovoltaics as the semiconductor copper indium gallium selenide (CIGS), also called CIGS solar cells, a type of second-generation thin-film solar cell. Indium is used in PNP bipolar junction transistors with germanium: when soldered at low temperature, indium does not stress the germanium.

Indium wire is used as a vacuum seal and a thermal conductor in cryogenics and ultra-high-vacuum applications, in such manufacturing applications as gaskets that deform to fill gaps. Owing to its great plasticity and adhesion to metals, Indium sheets are sometimes used for cold-soldering in microwave circuits and waveguide joints, where direct soldering is complicated. Indium is an ingredient in the gallium–indium–tin alloy galinstan, which is liquid at room temperature and replaces mercury in some thermometers. Other alloys of indium with bismuth, cadmium, lead, and tin, which have higher but still low melting points (between 50 and 100 °C), are used in fire sprinkler systems and heat regulators.

Indium is one of many substitutes for mercury in alkaline batteries to prevent the zinc from corroding and releasing hydrogen gas. Indium is added to some dental amalgam alloys to decrease the surface tension of the mercury and allow for less mercury and easier amalgamation.

Indium's high neutron-capture cross-section for thermal neutrons makes it suitable for use in control rods for nuclear reactors, typically in an alloy of 80% silver, 15% indium, and 5% cadmium. In nuclear engineering, the (n,n') reactions of 113In and 115In are used to determine magnitudes of neutron fluxes.

In 2009, Professor Mas Subramanian and former graduate student Andrew Smith at Oregon State University discovered that indium can be combined with yttrium and manganese to form an intensely blue, non-toxic, inert, fade-resistant pigment, YInMn blue, the first new inorganic blue pigment discovered in 200 years.

Biological role and precautions
Indium has no metabolic role in any organism. In a similar way to aluminium salts, indium(III) ions can be toxic to the kidney when given by injection. Indium tin oxide and indium phosphide harm the pulmonary and immune systems, predominantly through ionic indium, though hydrated indium oxide is more than forty times as toxic when injected, measured by the quantity of indium introduced. Radioactive indium-111 (in very small amounts on a chemical basis) is used in nuclear medicine tests, as a radiotracer to follow the movement of labeled proteins and white blood cells in the body. Indium compounds are mostly not absorbed upon ingestion and are only moderately absorbed on inhalation; they tend to be stored temporarily in the muscles, skin, and bones before being excreted, and the biological half-life of indium is about two weeks in humans.

People can be exposed to indium in the workplace by inhalation, ingestion, skin contact, and eye contact. Indium lung is a lung disease characterized by pulmonary alveolar proteinosis and pulmonary fibrosis, first described by Japanese researchers in 2003. , 10 cases had been described, though more than 100 indium workers had documented respiratory abnormalities. The National Institute for Occupational Safety and Health has set a recommended exposure limit (REL) of 0.1 mg/m3 over an eight-hour workday.

title: Indium


text:
Iodine

Iodine is a chemical element; it has symbol I and atomic number 53. The heaviest of the stable halogens, it exists at standard conditions as a semi-lustrous, non-metallic solid that melts to form a deep violet liquid at 114 C, and boils to a violet gas at 184 C. The element was discovered by the French chemist Bernard Courtois in 1811 and was named two years later by Joseph Louis Gay-Lussac, after the Ancient Greek Ιώδης, meaning 'violet'.

Iodine occurs in many oxidation states, including iodide (I−), iodate, and the various periodate anions. As the heaviest essential mineral nutrient, iodine is required for the synthesis of thyroid hormones. Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.

The dominant producers of iodine today are Chile and Japan. Due to its high atomic number and ease of attachment to organic compounds, it has also found favour as a non-toxic radiocontrast material. Because of the specificity of its uptake by the human body, radioactive isotopes of iodine can also be used to treat thyroid cancer. Iodine is also used as a catalyst in the industrial production of acetic acid and some polymers.

It is on the World Health Organization's List of Essential Medicines.

History
In 1811, iodine was discovered by French chemist Bernard Courtois, who was born to a family of manufacturers of saltpetre (an essential component of gunpowder). At the time of the Napoleonic Wars, saltpetre was in great demand in France. Saltpetre produced from French nitre beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding sulfuric acid. Courtois once added excessive sulfuric acid and a cloud of violet vapour rose. He noted that the vapour crystallised on cold surfaces, making dark black crystals. Courtois suspected that this material was a new element but lacked funding to pursue it further.

Courtois gave samples to his friends, Charles Bernard Desormes (1777–1838) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to chemist Joseph Louis Gay-Lussac (1778–1850), and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Desormes and Clément made Courtois' discovery public by describing the substance to a meeting of the Imperial Institute of France. On 6 December 1813, Gay-Lussac found and announced that the new substance was either an element or a compound of oxygen and he found that it is an element. Gay-Lussac suggested the name "iode" (anglicized as "iodine"), from the Ancient Greek Ιώδης (, "violet"), because of the colour of iodine vapor. Ampère had given some of his sample to British chemist Humphry Davy (1778–1829), who experimented on the substance and noted its similarity to chlorine and also found it as an element. Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element called iodine. Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists found that both of them identified iodine first and also knew that Courtois is the first one to isolate the element.

In 1873, the French medical researcher Casimir Davaine (1812–1882) discovered the antiseptic action of iodine. Antonio Grossich (1849–1926), an Istrian-born surgeon, was among the first to use sterilisation of the operative field. In 1908, he introduced tincture of iodine as a way to rapidly sterilise the human skin in the surgical field.

In early periodic tables, iodine was often given the symbol J, for Jod, its name in German.

Properties
Iodine is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; it is the heaviest stable member of its group. (The fifth and sixth halogens, the radioactive astatine and tennessine, are not well-studied due to their expense and inaccessibility in large quantities, but appear to show various unusual properties for the group due to relativistic effects.) Iodine has an electron configuration of [Kr]4d105s25p5, with the seven electrons in the fifth and outermost shell being its valence electrons. Like the other halogens, it is one electron short of a full octet and is hence an oxidising agent, reacting with many elements in order to complete its outer shell, although in keeping with periodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowest electronegativity among them, just 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatine continues the trend with an electronegativity of 2.2). Elemental iodine hence forms diatomic molecules with chemical formula I2, where two iodine atoms share a pair of electrons in order to each achieve a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of iodine atoms. Similarly, the iodide anion, I−, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic I2. (Astatine goes further, being indeed unstable as At− and readily oxidised to At0 or At+.)

The halogens darken in colour as the group is descended: fluorine is a very pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, and iodine is violet.

Elemental iodine is slightly soluble in water, with one gram dissolving in 3450 mL at 20 °C and 1280 mL at 50 °C; potassium iodide may be added to increase solubility via formation of triiodide ions, among other polyiodides. Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility. Polar solutions, such as aqueous solutions, are brown, reflecting the role of these solvents as Lewis bases; on the other hand, nonpolar solutions are violet, the color of iodine vapour. Charge-transfer complexes form when iodine is dissolved in polar solvents, hence changing the colour. Iodine is violet when dissolved in carbon tetrachloride and saturated hydrocarbons but deep brown in alcohols and amines, solvents that form charge-transfer adducts.

The melting and boiling points of iodine are the highest among the halogens, conforming to the increasing trend down the group, since iodine has the largest electron cloud among them that is the most easily polarised, resulting in its molecules having the strongest Van der Waals interactions among the halogens. Similarly, iodine is the least volatile of the halogens, though the solid still can be observed to give off purple vapor. Due to this property iodine is commonly used to demonstrate sublimation directly from solid to gas, which gives rise to a misconception that it does not melt in atmospheric pressure. Because it has the largest atomic radius among the halogens, iodine has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the halogens.

The interhalogen bond in diiodine is the weakest of all the halogens. As such, 1% of a sample of gaseous iodine at atmospheric pressure is dissociated into iodine atoms at 575 °C. Temperatures greater than 750 °C are required for fluorine, chlorine, and bromine to dissociate to a similar extent. Most bonds to iodine are weaker than the analogous bonds to the lighter halogens. Gaseous iodine is composed of I2 molecules with an I–I bond length of 266.6 pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5 pm) in solid orthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbour xenon: the Xe–Xe bond length is 308.71 pm.) As such, within the iodine molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk iodine, to a shiny appearance and semiconducting properties. Iodine is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.

Isotopes
Of the thirty-seven known isotopes of iodine, only one occurs in nature, iodine-127. The others are radioactive and have half-lives too short to be primordial. As such, iodine is both monoisotopic and mononuclidic and its atomic weight is known to great precision, as it is a constant of nature.

The longest-lived of the radioactive isotopes of iodine is iodine-129, which has a half-life of 15.7 million years, decaying via beta decay to stable xenon-129. Some iodine-129 was formed along with iodine-127 before the formation of the Solar System, but it has by now completely decayed away, making it an extinct radionuclide that is nevertheless still useful in dating the history of the early Solar System or very old groundwaters, due to its mobility in the environment. Its former presence may be determined from an excess of its daughter xenon-129. Traces of iodine-129 still exist today, as it is also a cosmogenic nuclide, formed from cosmic ray spallation of atmospheric xenon: these traces make up 10−14 to 10−10 of all terrestrial iodine. It also occurs from open-air nuclear testing, and is not hazardous because of its very long half-life, the longest of all fission products. At the peak of thermonuclear testing in the 1960s and 1970s, iodine-129 still made up only about 10−7 of all terrestrial iodine. Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.

The other iodine radioisotopes have much shorter half-lives, no longer than days. Some of them have medical applications involving the thyroid gland, where the iodine that enters the body is stored and concentrated. Iodine-123 has a half-life of thirteen hours and decays by electron capture to tellurium-123, emitting gamma radiation; it is used in nuclear medicine imaging, including single photon emission computed tomography (SPECT) and X-ray computed tomography (X-Ray CT) scans. Iodine-125 has a half-life of fifty-nine days, decaying by electron capture to tellurium-125 and emitting low-energy gamma radiation; the second-longest-lived iodine 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.

Protection usually used 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
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 $\pi$* 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 $H_{2}O ⟶$ 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 gaseous hydrogen iodide, hydroiodic acid has major industrial use in the manufacture of acetic acid by the Cativa process.

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 permittivity is low and it does not dissociate appreciably into H2I+ and ions – the latter, in any case, are much less stable than the bifluoride ions  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 quaternary ammonium cation (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., [S2I42+][AsF6–]2 and [Se2I42+][Sb2F11–]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. Other syntheses 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:

Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:

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 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 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 salicylic acid, since when iodine chloride undergoes homolytic fission, chlorine and iodine are produced and the former is more reactive. However, iodine chloride in carbon tetrachloride 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 carbon tetrachloride 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 and  ions (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 and  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 and. 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
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 known since its formation in 1813 by Gay-Lussac and Davy. It is most easily made by the dehydration of iodic acid (HIO3), of which it is the anhydride. It will quickly oxidise carbon monoxide completely to carbon dioxide at room temperature, and is thus a useful reagent in determining carbon monoxide concentration. It also oxidises nitrogen oxide, ethylene, and hydrogen sulfide. It reacts with sulfur trioxide and peroxydisulfuryl difluoride (S2O6F2) to form salts of the iodyl cation, [IO2]+, and is reduced by concentrated sulfuric acid to iodosyl salts involving [IO]+. It may be fluorinated by fluorine, bromine trifluoride, sulfur tetrafluoride, or chloryl fluoride, resulting iodine pentafluoride, which also reacts with iodine pentoxide, giving iodine(V) oxyfluoride, IOF3. A few other less stable oxides are known, notably I4O9 and I2O4; their structures have not been determined, but reasonable guesses are IIII(IVO3)3 and [IO]+[IO3]− respectively.

More important are the four oxoacids: hypoiodous acid (HIO), iodous acid (HIO2), iodic acid (HIO3), and periodic acid (HIO4 or H5IO6). When iodine dissolves in aqueous solution, the following reactions occur:

 * I2 + H2O || HIO + H+ + I− || Kac = 2.0 × 10−13 mol2 L−2
 * I2 + 2 OH− || IO− + H2O + I− || Kalk = 30 mol2 L−2
 * }
 * I2 + 2 OH− || IO− + H2O + I− || Kalk = 30 mol2 L−2
 * }

Hypoiodous acid is unstable to disproportionation. The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate:

 * 3 IO− 2 I− +  || K = 1020
 * }
 * }

Iodous acid and iodite are even less stable and exist only as a fleeting intermediate in the oxidation of iodide to iodate, if at all. Iodates are by far the most important of these compounds, which can be made by oxidising alkali metal iodides with oxygen at 600 °C and high pressure, or by oxidising iodine with chlorates. Unlike chlorates, which disproportionate very slowly to form chloride and perchlorate, iodates are stable to disproportionation in both acidic and alkaline solutions. From these, salts of most metals can be obtained. Iodic acid is most easily made by oxidation of an aqueous iodine suspension by electrolysis or fuming nitric acid. Iodate has the weakest oxidising power of the halates, but reacts the quickest.

Many periodates are known, including not only the expected tetrahedral, but also square-pyramidal , octahedral orthoperiodate , [IO3(OH)3]2−, [I2O8(OH2)]4−, and. They are usually made by oxidising alkaline sodium iodate electrochemically (with lead(IV) oxide as the anode) or by chlorine gas:
 * + 6 OH− → + 3 H2O + 2 e−
 * + 6 OH− + Cl2 → + 2 Cl− + 3 H2O

They are thermodymically and kinetically powerful oxidising agents, quickly oxidising Mn2+ to permanganate, and cleaving glycols, α-diketones, α-ketols, α-aminoalcohols, and α-diamines. Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of −5. Orthoperiodic acid, H5IO6, is stable, and dehydrates at 100 °C in a vacuum to Metaperiodic acid, HIO4. Attempting to go further does not result in the nonexistent iodine heptoxide (I2O7), but rather iodine pentoxide and oxygen. Periodic acid may be protonated by sulfuric acid to give the cation, isoelectronic to Te(OH)6 and, and giving salts with bisulfate and sulfate.

Polyiodine compounds
When iodine dissolves in strong acids, such as fuming sulfuric acid, a bright blue paramagnetic solution including cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine with antimony pentafluoride:
 * 2 I2 + 5 SbF5 $$ 2 I2Sb2F11 + SbF3

The salt I2Sb2F11 is dark blue, and the blue tantalum analogue I2Ta2F11 is also known. Whereas the I–I bond length in I2 is 267 pm, that in is only 256 pm as the missing electron in the latter has been removed from an antibonding orbital, making the bond stronger and hence shorter. In fluorosulfuric acid solution, deep-blue reversibly dimerises below −60 °C, forming red rectangular diamagnetic. Other polyiodine cations are not as well-characterised, including bent dark-brown or black and centrosymmetric C2h green or black, known in the  and  salts among others.

The only important polyiodide anion in aqueous solution is linear triiodide,. Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution:
 * I2 + I−  (Keq = ~700 at 20 °C)

Many other polyiodides may be found when solutions containing iodine and iodide crystallise, such as, , , and , whose salts with large, weakly polarising cations such as Cs+ may be isolated.

Organoiodine compounds
Organoiodine compounds have been fundamental in the development of organic synthesis, such as in the Hofmann elimination of amines, the Williamson ether synthesis, the Wurtz coupling reaction, and in Grignard reagents.

The carbon–iodine bond is a common functional group that forms part of core organic chemistry; formally, these compounds may be thought of as organic derivatives of the iodide anion. The simplest organoiodine compounds, alkyl iodides, may be synthesised by the reaction of alcohols with phosphorus triiodide; these may then be used in nucleophilic substitution reactions, or for preparing Grignard reagents. The C–I bond is the weakest of all the carbon–halogen bonds due to the minuscule difference in electronegativity between carbon (2.55) and iodine (2.66). As such, iodide is the best leaving group among the halogens, to such an extent that many organoiodine compounds turn yellow when stored over time due to decomposition into elemental iodine; as such, they are commonly used in organic synthesis, because of the easy formation and cleavage of the C–I bond. They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine. A few organic oxidising agents like the iodanes contain iodine in a higher oxidation state than −1, such as 2-iodoxybenzoic acid, a common reagent for the oxidation of alcohols to aldehydes, and iodobenzene dichloride (PhICl2), used for the selective chlorination of alkenes and alkynes. One of the more well-known uses of organoiodine compounds is the so-called iodoform test, where iodoform (CHI3) is produced by the exhaustive iodination of a methyl ketone (or another compound capable of being oxidised to a methyl ketone), as follows:

 * [[Image:Iodoform synthesis.svg|450px]]

Some drawbacks of using organoiodine compounds as compared to organochlorine or organobromine compounds is the greater expense and toxicity of the iodine derivatives, since iodine is expensive and organoiodine compounds are stronger alkylating agents. For example, iodoacetamide and iodoacetic acid denature proteins by irreversibly alkylating cysteine residues and preventing the reformation of disulfide linkages.

Halogen exchange to produce iodoalkanes by the Finkelstein reaction is slightly complicated by the fact that iodide is a better leaving group than chloride or bromide. The difference is nevertheless small enough that the reaction can be driven to completion by exploiting the differential solubility of halide salts, or by using a large excess of the halide salt. In the classic Finkelstein reaction, an alkyl chloride or an alkyl bromide is converted to an alkyl iodide by treatment with a solution of sodium iodide in acetone. Sodium iodide is soluble in acetone and sodium chloride and sodium bromide are not. The reaction is driven toward products by mass action due to the precipitation of the insoluble salt.

Occurrence and production
Iodine is the least abundant of the stable halogens, comprising only 0.46 parts per million of Earth's crustal rocks (compare: fluorine: 544 ppm, chlorine: 126 ppm, bromine: 2.5 ppm) making it the 60th most abundant element. Iodide minerals are rare, and most deposits that are concentrated enough for economical extraction are iodate minerals instead. Examples include lautarite, Ca(IO3)2, and dietzeite, 7Ca(IO3)2·8CaCrO4. These are the minerals that occur as trace impurities in the caliche, found in Chile, whose main product is sodium nitrate. In total, they can contain at least 0.02% and at most 1% iodine by mass. Sodium iodate is extracted from the caliche and reduced to iodide by sodium bisulfite. This solution is then reacted with freshly extracted iodate, resulting in comproportionation to iodine, which may be filtered off.

The caliche was the main source of iodine in the 19th century and continues to be important today, replacing kelp (which is no longer an economically viable source), but in the late 20th century brines emerged as a comparable source. The Japanese Minami Kantō gas field east of Tokyo and the American Anadarko Basin gas field in northwest Oklahoma are the two largest such sources. The brine is hotter than 60 °C from the depth of the source. The brine is first purified and acidified using sulfuric acid, then the iodide present is oxidised to iodine with chlorine. An iodine solution is produced, but is dilute and must be concentrated. Air is blown into the solution to evaporate the iodine, which is passed into an absorbing tower, where sulfur dioxide reduces the iodine. The hydrogen iodide (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.

 * 2 HI + Cl2 → I2↑ + 2 HCl
 * I2 + 2 H2O + SO2 → 2 HI + H2SO4
 * 2 HI + Cl2 → I2↓ + 2 HCl

These sources ensure that Chile and Japan are the largest producers of iodine today. Alternatively, the brine may be treated with silver nitrate to precipitate out iodine as silver iodide, which is then decomposed by reaction with iron to form metallic silver and a solution of iron(II) iodide. The iodine is then liberated by displacement with chlorine.

Applications
About half of all produced iodine goes into various organoiodine compounds, another 15% remains as the pure element, another 15% is used to form potassium iodide, and another 15% for other inorganic iodine compounds. Among the major uses of iodine compounds are catalysts, animal feed supplements, stabilisers, dyes, colourants and pigments, pharmaceutical, sanitation (from tincture of iodine), and photography; minor uses include smog inhibition, cloud seeding, and various uses in analytical chemistry.

Chemical analysis
The iodide and iodate anions are often used for quantitative volumetric analysis, for example in iodometry. Iodine and starch form a blue complex, and this reaction is often used to test for either starch or iodine and as an indicator in iodometry. The iodine test for starch is still used to detect counterfeit banknotes printed on starch-containing paper.

The iodine value is the mass of iodine in grams that is consumed by 100 grams of a chemical substance typically fats or oils. Iodine numbers are often used to determine the amount of unsaturation in fatty acids. This unsaturation is in the form of double bonds, which react with iodine compounds.

Potassium tetraiodomercurate(II), K2HgI4, is also known as Nessler's reagent. It is often used as a sensitive spot test for ammonia. Similarly, Mayer's reagent (potassium tetraiodomercurate(II) solution) is used as a precipitating reagent to test for alkaloids. Aqueous alkaline iodine solution is used in the iodoform test for methyl ketones.

Spectroscopy
The spectrum of the iodine molecule, I2, consists of (not exclusively) tens of thousands of sharp spectral lines in the wavelength range 500–700 nm. It is therefore a commonly used wavelength reference (secondary standard). By measuring with a spectroscopic Doppler-free technique while focusing on one of these lines, the hyperfine structure of the iodine molecule reveals itself. A line is now resolved such that either 15 components (from even rotational quantum numbers, Jeven), or 21 components (from odd rotational quantum numbers, Jodd) are measurable.

Caesium iodide and thallium-doped sodium iodide are used in crystal scintillators for the detection of gamma rays. The efficiency is high and energy dispersive spectroscopy is possible, but the resolution is rather poor.

Spacecraft propulsion
Ion thruster propulsion systems employing iodine as the reaction mass can be built more compactly, with less mass (and cost), and operate more efficiently than the gridded ion thrusters that were utilized to propel previous spacecraft, such as Japan's Hayabusa probes, ESA's GOCE satellite, or NASA's DART mission, all of which used xenon for this purpose. Iodine's atomic weight is only 3.3% less than that of xenon, while its first two ionisation energies average 12% less; together, these make iodine ions a promising substitute. However, iodine introduces chemical reactivity issues not present in xenon plasmas.

Use of iodine should allow more widespread application of ion-thrust technology, particularly with smaller-scale space vehicles. According to the European Space Agency, "This small but potentially disruptive innovation could help to clear the skies of space junk, by enabling tiny satellites to self-destruct cheaply and easily at the end of their missions, by steering themselves into the atmosphere where they would burn up."

In early 2021, the French group ThrustMe performed an in-orbit demonstration of an electric-powered ion thruster for spacecraft, where iodine was used in lieu of xenon as the source of plasma, in order to generate thrust by accelerating ions with an electrostatic field.

Medicine
Elemental iodine (I2) is used as an antiseptic in medicine. A number of water-soluble compounds, from triiodide (I3−, generated in situ by adding iodide to poorly water-soluble elemental iodine) to various iodophors, slowly decompose to release I2 when applied.

Salts of iodide and iodate are used at low doses in iodised salt. A saturated solution of potassium iodide is used to treat acute thyrotoxicosis. It is also used to block uptake of iodine-131 in the thyroid gland (see isotopes section above), when this isotope is used as part of radiopharmaceuticals (such as iobenguane) that are not targeted to the thyroid or thyroid-type tissues.

As an element with high electron density and atomic number, iodine absorbs X-rays weaker than 33.3 keV due to the photoelectric effect of the innermost electrons. Organoiodine compounds are used with intravenous injection as X-ray radiocontrast agents. This application is often in conjunction with advanced X-ray techniques such as angiography and CT scanning. At present, all water-soluble radiocontrast agents rely on iodine-containing compounds.

Others
The production of ethylenediamine dihydroiodide, provided as a nutritional supplement for livestock, consumes a large portion of available iodine. Another significant use is a catalyst for the production of acetic acid by the Monsanto and Cativa processes. In these technologies, which support the world's demand for acetic acid, hydroiodic acid converts the methanol feedstock into methyl iodide, which undergoes carbonylation. Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid.

Inorganic iodides find specialised uses. Titanium, zirconium, hafnium, and thorium are purified by the Van Arkel–de Boer process, which involves the reversible formation of the tetraiodides of these elements. Silver iodide is a major ingredient to traditional photographic film. Thousands of kilograms of silver iodide are used annually for cloud seeding to induce rain.

The organoiodine compound erythrosine is an important food coloring agent. Perfluoroalkyl iodides are precursors to important surfactants, such as perfluorooctanesulfonic acid.

The iodine clock reaction (in which iodine also serves as a test for starch, forming a dark blue complex), is a popular educational demonstration experiment and example of a seemingly oscillating reaction (it is only the concentration of an intermediate product that oscillates).

$125$I is used as the radiolabel in investigating which ligands go to which plant pattern recognition receptors (PRRs).

Biological role
Iodine is an essential element for life and, at atomic number Z = 53, is the heaviest element commonly needed by living organisms. (Lanthanum and the other lanthanides, as well as tungsten with Z = 74 and uranium with Z = 92, are used by a few microorganisms.  ) It is required for the synthesis of the growth-regulating thyroid hormones tetraiodothyronine and triiodothyronine (T4 and T3 respectively, named after their number of iodine atoms). A deficiency of iodine leads to decreased production of T3 and T4 and a concomitant enlargement of the thyroid tissue in an attempt to obtain more iodine, causing the disease goitre. The major form of thyroid hormone in the blood is tetraiodothyronine (T4), which has a longer life than triiodothyronine (T3). In humans, the ratio of T4 to T3 released into the blood is between 14:1 and 20:1. T4 is converted to the active T3 (three to four times more potent than T4) within cells by deiodinases (5'-iodinase). These are further processed by decarboxylation and deiodination to produce iodothyronamine (T1a) and thyronamine (T0a'). All three isoforms of the deiodinases are selenium-containing enzymes; thus metallic selenium is needed for triiodothyronine and tetraiodothyronine production.

Iodine accounts for 65% of the molecular weight of T4 and 59% of T3. Fifteen to 20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of all iodine in the body is found in other tissues, including mammary glands, eyes, gastric mucosa, thymus, cerebrospinal fluid, choroid plexus, arteries, cervix, salivary glands. During pregnancy, the placenta is able to store and accumulate iodine. In the cells of those tissues, iodine enters directly by sodium-iodide symporter (NIS). The action of iodine in mammal tissues is related to fetal and neonatal development, and in the other tissues, it is known.

Dietary recommendations and intake
The daily levels of intake recommended by the United States National Academy of Medicine are between 110 and 130 μg for infants up to 12 months, 90 μg for children up to eight years, 130 μg for children up to 13 years, 150 μg for adults, 220 μg for pregnant women and 290 μg for lactating women. The Tolerable Upper Intake Level (TUIL) for adults is 1,100 μg/day. This upper limit was assessed by analyzing the effect of supplementation on thyroid-stimulating hormone.

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR; AI and UL are defined the same as in the United States. For women and men ages 18 and older, the PRI for iodine is set at 150 μg/day; the PRI during pregnancy and lactation is 200 μg/day. For children aged 1–17 years, the PRI increases with age from 90 to 130 μg/day. These PRIs are comparable to the U.S. RDAs with the exception of that for lactation.

The thyroid gland needs 70 μg/day of iodine to synthesise the requisite daily amounts of T4 and T3. The higher recommended daily allowance levels of iodine seem necessary for optimal function of a number of body systems, including mammary glands, gastric mucosa, salivary glands, brain cells, choroid plexus, thymus, arteries.

Natural food sources of iodine include seafood which contains fish, seaweeds, kelp, shellfish and other foods which contain dairy products, eggs, meats, vegetables, so long as the animals ate iodine richly, and the plants are grown on iodine-rich soil. Iodised salt is fortified with potassium iodate, a salt of iodine, potassium, oxygen.

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. The general US population has adequate iodine nutrition, with lactating women and pregnant women having a mild risk of deficiency. In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day from wakame and kombu that are eaten, both in the form of kombu and wakame and kombu and wakame umami extracts for soup stock and potato chips. However, new studies suggest that Japan's consumption is closer to 1,000–3,000 μg/day. The adult UL in Japan was last revised to 3,000 μg/day in 2015.

After iodine fortification programs such as iodisation of salt have been done, some cases of iodine-induced hyperthyroidism have been observed (so-called Jod-Basedow phenomenon). The condition occurs mainly in people above 40 years of age, and the risk is higher when iodine deficiency is high and the first rise in iodine consumption is high.

Deficiency
In areas where there is little iodine in the diet, which are remote inland areas and faraway mountainous areas where no iodine rich foods are eaten, iodine deficiency gives rise to hypothyroidism, symptoms of which are extreme fatigue, goitre, mental slowing, depression, low weight gain, and low basal body temperatures. Iodine deficiency is the leading cause of preventable intellectual disability, a result that occurs primarily when babies or small children are rendered hypothyroidic by no iodine. The addition of iodine to salt has largely destroyed this problem in wealthier areas, but iodine deficiency remains a serious public health problem in poorer areas today. Iodine deficiency is also a problem in certain areas of all continents of the world. Information processing, fine motor skills, and visual problem solving are normalised by iodine repletion in iodine-deficient people.

Toxicity
Elemental iodine (I2) is toxic if taken orally undiluted. The lethal dose for an adult human is 30 mg/kg, which is about 2.1–2.4 grams for a human weighing 70 to 80 kg (even when experiments on rats demonstrated that these animals could survive after eating a 14000 mg/kg dose and are still living after that). Excess iodine is more cytotoxic in the presence of selenium deficiency. Iodine supplementation in selenium-deficient populations is problematic for this reason. The toxicity derives from its oxidizing properties, through which it denaturates proteins (including enzymes).

Elemental iodine is also a skin irritant. Direct contact with skin can cause damage, and solid iodine crystals should be handled with care. Solutions with high elemental iodine concentration, such as tincture of iodine and Lugol's solution, are capable of causing tissue damage if used in prolonged cleaning or antisepsis; similarly, liquid Povidone-iodine (Betadine) trapped against the skin resulted in chemical burns in some reported cases.

Occupational exposure
People can be exposed to iodine in the workplace by inhalation, ingestion, skin contact, and eye contact. The Occupational Safety and Health Administration (OSHA) has set the legal limit (Permissible exposure limit) for iodine exposure in the workplace at 0.1 ppm (1 mg/m3) during an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 0.1 ppm (1 mg/m3) during an 8-hour workday. At levels of 2 ppm, iodine is immediately dangerous to life and health.

Allergic reactions
Some people develop a hypersensitivity to products and foods containing iodine. Applications of tincture of iodine or Betadine can cause rashes, sometimes severe. Parenteral use of iodine-based contrast agents (see above) can cause reactions ranging from a mild rash to fatal anaphylaxis. Such reactions have led to the misconception (widely held, even among physicians) that some people are allergic to iodine itself; even allergies to iodine-rich foods have been so construed. In fact, there has never been a confirmed report of a true iodine allergy, as an allergy to iodine or iodine salts is biologically impossible. Hypersensitivity reactions to products and foods containing iodine are apparently related to their other molecular components; thus, a person who has demonstrated an allergy to one food or product containing iodine may not have an allergic reaction to another. Patients with various food allergies (fishes, shellfishes, eggs, milk, seaweeds, kelp, meats, vegetables, kombu, wakame) do not have an increased risk for a contrast medium hypersensitivity. As with all medications, the patient's allergy history should be questioned and consulted before any containing iodine are administered.

US DEA List I status
Phosphorus reduces iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine and pseudoephedrine to methamphetamine. For this reason, iodine was designated by the United States Drug Enforcement Administration as a List I precursor chemical under 21 CFR 1310.02.

title: Iodine