Halogen - Wikipedia

Group of chemical elements This article is about the chemical series. For other uses, see Halogen (disambiguation).

Halogens
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson chalcogens ←    → noble gases
IUPAC group number 17 Name by element fluorine group Trivial name halogens CAS group number(US, pattern A-B-A) VIIA old IUPAC number(Europe, pattern A-B) VIIB

↓ Period
2	Fluorine (F)9 Halogen
3	Chlorine (Cl)17 Halogen
4	Bromine (Br)35 Halogen
5	Iodine (I)53 Halogen
6	Astatine (At)85 Halogen
7	Tennessine (Ts)117 Halogen
Legend primordial element element from decay Synthetic

The halogens (/ˈhælədʒən,ˈheɪ-,-loʊ-,-ˌdʒɛn/[1][2][3]) are a group in the periodic table consisting of six chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and the radioactive elements astatine (At) and tennessine (Ts), though some authors[4] would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of gallium. In the modern IUPAC nomenclature, this group is known as group 17.[5]

The word "halogen" means "salt former" or "salt maker". When halogens react with metals, they produce a wide range of salts, including calcium fluoride, sodium chloride (common table salt), silver bromide, and potassium iodide.[6]

The group of halogens is the only periodic table group that contains elements in three of the main states of matter at standard temperature and pressure, though not far above room temperature the same becomes true of groups 1 and 15, assuming white phosphorus is taken as the standard state.[n 1] All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from minerals or salts. The middle halogens—chlorine, bromine, and iodine—are often used as disinfectants. Organobromides are the most important class of flame retardants, while elemental halogens are dangerous and can be toxic.

History

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The fluorine mineral fluorspar was known as early as 1529. Early chemists realized that fluorine compounds contain an undiscovered element, but were unable to isolate it. In 1869, George Gore, an English chemist, ran a current of electricity through hydrofluoric acid and probably produced fluorine, but he was unable to prove his results at the time.[7] In 1886, Henri Moissan, a chemist in Paris, performed electrolysis on potassium bifluoride dissolved in anhydrous hydrogen fluoride, and successfully isolated fluorine.[8]

Hydrochloric acid was known to alchemists and early chemists. However, elemental chlorine was not produced until 1774, when Carl Wilhelm Scheele heated hydrochloric acid with manganese dioxide. Scheele called the element "dephlogisticated muriatic acid", which is how chlorine was known for 33 years. In 1807, Humphry Davy investigated chlorine and discovered that it is an actual element. Chlorine gas was used as a poisonous gas during World War I. It displaced oxygen in contaminated areas and replaced common oxygenated air with the toxic chlorine gas. The gas would burn human tissue externally and internally, especially the lungs, making breathing difficult or impossible depending on the level of contamination.[8]

Bromine was discovered in the 1820s by Antoine Jérôme Balard. Balard discovered bromine by passing chlorine gas through a sample of brine. He originally proposed the name muride for the new element, but the French Academy changed the element's name to bromine.[8]

Iodine was discovered by Bernard Courtois, who was using seaweed ash as part of a process for saltpeter manufacture. Courtois typically boiled the seaweed ash with water to generate potassium chloride. However, in 1811, Courtois added sulfuric acid to his process and found that his process produced purple fumes that condensed into black crystals. Suspecting that these crystals were a new element, Courtois sent samples to other chemists for investigation. Iodine was proven to be a new element by Joseph Gay-Lussac.[8]

In 1931, Fred Allison claimed to have discovered element 85 with a magneto-optical machine, and named the element Alabamine, but was mistaken. In 1937, Rajendralal De claimed to have discovered element 85 in minerals, and called the element dakine, but he was also mistaken. An attempt at discovering element 85 in 1939 by Horia Hulubei and Yvette Cauchois via spectroscopy was also unsuccessful, as was an attempt in the same year by Walter Minder, who discovered an iodine-like element resulting from beta decay of polonium. Element 85, now named astatine, was produced successfully in 1940 by Dale R. Corson, K.R. Mackenzie, and Emilio G. Segrè, who bombarded bismuth with alpha particles.[8]

In 2010, a team led by nuclear physicist Yuri Oganessian involving scientists from the JINR, Oak Ridge National Laboratory, Lawrence Livermore National Laboratory, and Vanderbilt University successfully bombarded berkelium-249 atoms with calcium-48 atoms to make tennessine.[9]

Etymology

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In 1811, the German chemist Johann Schweigger proposed that the name "halogen" – meaning "salt producer", from αλς [hals] "salt" and γενειν [genein] "to beget" – replace the name "chlorine", which had been proposed by the English chemist Humphry Davy.[10] Davy's name for the element prevailed.[11] However, in 1826, the Swedish chemist Baron Jöns Jacob Berzelius proposed the term "halogen" for the elements fluorine, chlorine, and iodine, which produce a sea-salt-like substance when they form a compound with an alkaline metal.[12][13]

The English names of these elements all have the ending -ine. Fluorine's name comes from the Latin word fluere, meaning "to flow", because it was derived from the mineral fluorite, which was used as a flux in metalworking. Chlorine's name comes from the Greek word chloros, meaning "greenish-yellow". Bromine's name comes from the Greek word bromos, meaning "stench". Iodine's name comes from the Greek word iodes, meaning "violet". Astatine's name comes from the Greek word astatos, meaning "unstable".[8] Tennessine is named after the US state of Tennessee, where it was synthesized.

Characteristics

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Chemical

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The halogens fluorine, chlorine, bromine, and iodine are nonmetals; the chemical properties of astatine and tennessine, two heaviest group 17 members, have not been conclusively investigated. The halogens show trends in chemical bond energy moving from top to bottom of the periodic table column with fluorine deviating slightly. It follows a trend in having the highest bond energy in compounds with other atoms, but it has very weak bonds within the diatomic F2 molecule. This means that further down group 17 in the periodic table, the reactivity of elements decreases because of the increasing size of the atoms.[14]

Halogen bond energies (kJ/mol)[15]

X	X2	HX	BX3	AlX3	CX4
F	159	574	645	582	456
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. This high reactivity is due to the high electronegativity of the atoms due to their high effective nuclear charge. Because the halogens have seven valence electrons in their outermost energy level, they can gain an electron by reacting with atoms of other elements to satisfy the octet rule. Fluorine is the most reactive of all elements; it is the only element more electronegative than oxygen, it attacks otherwise-inert materials such as glass, and it forms compounds with the usually inert noble gases. It is a corrosive and highly toxic gas. The reactivity of fluorine is such that, if used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Thus, fluorine must be handled with substances such as Teflon (which is itself an organofluorine compound), extremely dry glass, or metals such as copper or steel, which form a protective layer of fluoride on their surface.

The high reactivity of fluorine allows some of the strongest bonds possible, especially to carbon. For example, Teflon is fluorine bonded with carbon and is extremely resistant to thermal and chemical attacks and has a high melting point.

Molecules

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Diatomic halogen molecules

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The stable halogens form homonuclear diatomic molecules. Due to relatively weak intermolecular forces, chlorine and fluorine form part of the group known as "elemental gases".

halogen	molecule	structure	model	d(X−X) / pm(gas phase)	d(X−X) / pm(solid phase)
fluorine	F2			143	149
chlorine	Cl2			199	198
bromine	Br2			228	227
iodine	I2			266	272

The elements become less reactive and have higher melting points as the atomic number increases. The higher melting points are caused by stronger London dispersion forces resulting from more electrons.

Compounds

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Hydrogen halides

[edit] Main article: Hydrogen halides

All of the halogens have been observed to react with hydrogen to form hydrogen halides. For fluorine, chlorine, and bromine, this reaction is in the form of:

H2 + X2 → 2HX

However, hydrogen iodide and hydrogen astatide can split back into their constituent elements.[16]

The hydrogen-halogen reactions get gradually less reactive toward the heavier halogens. A fluorine-hydrogen reaction is explosive even when it is dark and cold. A chlorine-hydrogen reaction is also explosive, but only in the presence of light and heat. A bromine-hydrogen reaction is even less explosive; it is explosive only when exposed to flames. Iodine and astatine only partially react with hydrogen, forming equilibria.[16]

All halogens form binary compounds with hydrogen known as the hydrogen halides: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAt). All of these compounds form acids when mixed with water. Hydrogen fluoride is the only hydrogen halide that forms hydrogen bonds. Hydrochloric acid, hydrobromic acid, hydroiodic acid, and hydroastatic acid are all strong acids, but hydrofluoric acid is a weak acid.[17]

All of the hydrogen halides are irritants. Hydrogen fluoride and hydrogen chloride are highly acidic. Hydrogen fluoride is used as an industrial chemical, and is highly toxic, causing pulmonary edema and damaging cells.[18] Hydrogen chloride is also a dangerous chemical. Breathing in gas with more than fifty parts per million of hydrogen chloride can cause death in humans.[19] Hydrogen bromide is even more toxic and irritating than hydrogen chloride. Breathing in gas with more than thirty parts per million of hydrogen bromide can be lethal to humans.[20] Hydrogen iodide, like other hydrogen halides, is toxic.[21]

Metal halides

[edit] Main article: Metal halides

All the halogens are known to react with sodium to form sodium fluoride, sodium chloride, sodium bromide, sodium iodide, and sodium astatide. Heated sodium's reaction with halogens produces bright-orange flames. Sodium's reaction with chlorine is in the form of:

2Na + Cl2 → 2NaCl[16]

Iron reacts with fluorine, chlorine, and bromine to form iron(III) halides. These reactions are in the form of:

2Fe + 3X2 → 2FeX3[16]

However, when iron reacts with iodine, it forms only iron(II) iodide.

Fe + I2 → FeI2

Iron wool can react rapidly with fluorine to form the white compound iron(III) fluoride even in cold temperatures. When chlorine comes into contact with a heated iron, they react to form the black iron(III) chloride. However, if the reaction conditions are moist, this reaction will instead result in a reddish-brown product. Iron can also react with bromine to form iron(III) bromide. This compound is reddish-brown in dry conditions. Iron's reaction with bromine is less reactive than its reaction with fluorine or chlorine. A hot iron can also react with iodine, but it forms iron(II) iodide. This compound may be gray, but the reaction is always contaminated with excess iodine, so it is not known for sure. Iron's reaction with iodine is less vigorous than its reaction with the lighter halogens.[16]

Interhalogen compounds

[edit] Main article: Interhalogen

Interhalogen compounds are in the form of XYn where X and Y are halogens and n is one, three, five, or seven. Interhalogen compounds contain at most two different halogens. Large interhalogens, such as ClF3 can be produced by a reaction of a pure halogen with a smaller interhalogen such as ClF. All interhalogens except IF7 can be produced by directly combining pure halogens in various conditions.[22]

Interhalogens are typically more reactive than all diatomic halogen molecules except F2 because interhalogen bonds are weaker. However, the chemical properties of interhalogens are still roughly the same as those of diatomic halogens. Many interhalogens consist of one or more atoms of fluorine bonding to a heavier halogen. Chlorine and bromine can bond with up to five fluorine atoms, and iodine can bond with up to seven fluorine atoms. Most interhalogen compounds are covalent gases. However, some interhalogens are liquids, such as BrF3, and many iodine-containing interhalogens are solids.[22]

Organohalogen compounds

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Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen atoms; these are known as halogenated compounds or organic halides. Chlorine is by far the most abundant of the halogens in seawater, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in brain function by mediating the action of the inhibitory transmitter GABA and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of thyroid hormones such as thyroxine. Organohalogens are also synthesized through the nucleophilic abstraction reaction.[23]

Polyhalogenated compounds

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Polyhalogenated compounds are industrially created compounds substituted with multiple halogens. Many of them are very toxic and bioaccumulate in humans, and have a very wide application range. They include PCBs, PBDEs, and perfluorinated compounds (PFCs), as well as numerous other compounds.

Reactions

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Reactions with water

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Fluorine reacts vigorously with water to produce oxygen (O2) and hydrogen fluoride (HF):[24]

2 F2(g) + 2 H2O(l) → O2(g) + 4 HF(aq)

Chlorine has maximum solubility of ca. 7.1 g Cl2 per kg of water at ambient temperature (21 °C).[25] Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid, a solution that can be used as a disinfectant or bleach:

Cl2(g) + H2O(l) → HCl(aq) + HClO(aq)

Bromine has a solubility of 3.41 g per 100 g of water,[26] but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO):

Br2(g) + H2O(l) → HBr(aq) + HBrO(aq)

Iodine, however, is minimally soluble in water (0.03 g/100 g water at 20 °C) and does not react with it.[27] However, iodine will form an aqueous solution in the presence of iodide ion, such as by addition of potassium iodide (KI), because the triiodide ion is formed.

Physical and atomic

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The table below is a summary of the key physical and atomic properties of the halogens. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations.

Halogen	Standard atomic weight(Da)[n 2][29]	Melting point(K)	Melting point(°C)	Boiling point(K)[30]	Boiling point(°C)[30]	Density(g/cm3 at 25 °C)	Electronegativity(Pauling)	First ionization energy(kJ·mol−1)	Covalent radius(pm)[31]
Fluorine	18.9984032(5)	53.53	−219.62	85.03	−188.12	0.0017	3.98	1681.0	71
Chlorine	[35.446; 35.457][n 3]	171.6	−101.5	239.11	−34.04	0.0032	3.16	1251.2	99
Bromine	79.904(1)	265.8	−7.3	332.0	58.8	3.1028	2.96	1139.9	114
Iodine	126.90447(3)	386.85	113.7	457.4	184.3	4.933	2.66	1008.4	133
Astatine	[210][n 4]	575	302	? 610	? 337	? 6.2–6.5[32]	2.2	899.0[33]	? 145[34]
Tennessine	[294][n 4]	? 623-823[35]	? 350-550[35]	? 883[35]	? 610[35]	? 7.1-7.3[35]	-	? 743[36]	? 157[35]

Z	Element	Electrons per shell
9	fluorine	2, 7
17	chlorine	2, 8, 7
35	bromine	2, 8, 18, 7
53	iodine	2, 8, 18, 18, 7
85	astatine	2, 8, 18, 32, 18, 7
117	tennessine	2, 8, 18, 32, 32, 18, 7 (predicted)[37]

Sublimation or boiling point (°C) of halogens at various pressures[38]

Tmelt (оС)		−100.7	−7.3	112.9
log(P[Pa])	mmHg	Cl2	Br2	I2
2.12490302	1	−118	−48.7	38.7
2.82387302	5	−106.7	−32.8	62.2
3.12490302	10	−101.6	−25	73.2
3.42593302	20	−93.3	−16.8	84.7
3.72696301	40	−84.5	−8	97.5
3.90305427	60	−79	−0.6	105.4
4.12490302	100	−71.7	9.3	116.5
4.42593302	200	−60.2	24.3	137.3
4.72696301	400	−47.3	41	159.8
5.00571661	760	−33.8	58.2	183
log(P[Pa])	atm	Cl2	Br2	I2
5.00571661	1	−33.8	58.2	183
5.30674661	2	−16.9	78.8
5.70468662	5	10.3	110.3
6.00571661	10	35.6	139.8
6.30674661	20	65	174
6.48283787	30	84.8	197
6.6077766	40	101.6	215
6.70468662	50	115.2	230
6.78386786	60	127.1	243.5

Isotopes

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Fluorine has one stable and naturally occurring isotope, fluorine-19. However, there are trace amounts in nature of the radioactive isotope fluorine-23, which occurs via cluster decay of protactinium-231. A total of eighteen isotopes of fluorine have been discovered, with atomic masses ranging from 13 to 31.

Chlorine has two stable and naturally occurring isotopes, chlorine-35 and chlorine-37. However, there are trace amounts in nature of the isotope chlorine-36, which occurs via spallation of argon-36. A total of 24 isotopes of chlorine have been discovered, with atomic masses ranging from 28 to 51.[8]

There are two stable and naturally occurring isotopes of bromine, bromine-79 and bromine-81. A total of 33 isotopes of bromine have been discovered, with atomic masses ranging from 66 to 98.

There is one stable and naturally occurring isotope of iodine, iodine-127. However, there are trace amounts in nature of the radioactive isotope iodine-129, which occurs via spallation and from the radioactive decay of uranium in ores. Several other radioactive isotopes of iodine have also been created naturally via the decay of uranium. A total of 38 isotopes of iodine have been discovered, with atomic masses ranging from 108 to 145.[8]

There are no stable isotopes of astatine. However, there are four naturally occurring radioactive isotopes of astatine produced via radioactive decay of uranium, neptunium, and plutonium. These isotopes are astatine-215, astatine-217, astatine-218, and astatine-219. A total of 31 isotopes of astatine have been discovered, with atomic masses ranging from 191 to 227.[8]

There are no stable isotopes of tennessine. Tennessine has only two known synthetic radioisotopes, tennessine-293 and tennessine-294.

Production

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Approximately six million metric tons of the fluorine mineral fluorite are produced each year. Four hundred-thousand metric tons of hydrofluoric acid are made each year. Fluorine gas is made from hydrofluoric acid produced as a by-product in phosphoric acid manufacture. Approximately 15,000 metric tons of fluorine gas are made per year.[8]

The mineral halite is the mineral that is most commonly mined for chlorine, but the minerals carnallite and sylvite are also mined for chlorine. Forty million metric tons of chlorine are produced each year by the electrolysis of brine.[8]

Approximately 450,000 metric tons of bromine are produced each year. Fifty percent of all bromine produced is produced in the United States, 35% in Israel, and most of the remainder in China. Historically, bromine was produced by adding sulfuric acid and bleaching powder to natural brine. However, in modern times, bromine is produced by electrolysis, a method invented by Herbert Dow. It is also possible to produce bromine by passing chlorine through seawater and then passing air through the seawater.[8]

In 2003, 22,000 metric tons of iodine were produced. Chile produces 40% of all iodine produced, Japan produces 30%, and smaller amounts are produced in Russia and the United States. Until the 1950s, iodine was extracted from kelp. However, in modern times, iodine is produced in other ways. One way that iodine is produced is by mixing sulfur dioxide with nitrate ores, which contain some iodates. Iodine is also extracted from natural gas fields.[8]

Even though astatine is naturally occurring, it is usually produced by bombarding bismuth with alpha particles.[8]

Tennessine is made by using a cyclotron, fusing berkelium-249 and calcium-48 to make tennessine-293 and tennessine-294.

Applications

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Disinfectants

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Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms through a process known as sterilization. Their reactivity is also put to use in bleaching. Sodium hypochlorite, which is produced from chlorine, is the active ingredient of most fabric bleaches, and chlorine-derived bleaches are used in the production of some paper products.

Lighting

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Halogen lamps are a type of incandescent lamp using a tungsten filament in bulbs that have small amounts of a halogen, such as iodine or bromine added. This enables the production of lamps that are much smaller than non-halogen incandescent lightbulbs at the same wattage. The gas reduces the thinning of the filament and blackening of the inside of the bulb resulting in a bulb that has a much greater life. Halogen lamps glow at a higher temperature (2800 to 3400 kelvin) with a whiter colour than other incandescent bulbs. However, this requires bulbs to be manufactured from fused quartz rather than silica glass to reduce breakage.[39]

Drug components

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In drug discovery, the incorporation of halogen atoms into a lead drug candidate results in analogues that are usually more lipophilic and less water-soluble.[40] As a consequence, halogen atoms are used to improve penetration through lipid membranes and tissues. It follows that there is a tendency for some halogenated drugs to accumulate in adipose tissue.

The chemical reactivity of halogen atoms depends on both their point of attachment to the lead and the nature of the halogen. Aromatic halogen groups are far less reactive than aliphatic halogen groups, which can exhibit considerable chemical reactivity. For aliphatic carbon-halogen bonds, the C-F bond is the strongest and usually less chemically reactive than aliphatic C-H bonds. The other aliphatic-halogen bonds are weaker, their reactivity increasing down the periodic table. They are usually more chemically reactive than aliphatic C-H bonds. As a consequence, the most common halogen substitutions are the less reactive aromatic fluorine and chlorine groups.

Biological role

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Fluoride anions are found in ivory, bones, teeth, blood, eggs, urine, and hair of organisms. Fluoride anions in very small amounts may be essential for humans.[41] There are 0.5 milligrams of fluorine per liter of human blood. Human bones contain 0.2 to 1.2% fluorine. Human tissue contains approximately 50 parts per billion of fluorine. A typical 70-kilogram human contains 3 to 6 grams of fluorine.[8]

Chloride anions are essential to a large number of species, humans included. The concentration of chlorine in the dry weight of cereals is 10 to 20 parts per million, while in potatoes the concentration of chloride is 0.5%. Plant growth is adversely affected by chloride levels in the soil falling below 2 parts per million. Human blood contains an average of 0.3% chlorine. Human bone typically contains 900 parts per million of chlorine. Human tissue contains approximately 0.2 to 0.5% chlorine. There is a total of 95 grams of chlorine in a typical 70-kilogram human.[8]

Some bromine in the form of the bromide anion is present in all organisms. A biological role for bromine in humans has not been proven, but some organisms contain organobromine compounds. Humans typically consume 1 to 20 milligrams of bromine per day. There are typically 5 parts per million of bromine in human blood, 7 parts per million of bromine in human bones, and 7 parts per million of bromine in human tissue. A typical 70-kilogram human contains 260 milligrams of bromine.[8]

Humans typically consume less than 100 micrograms of iodine per day. Iodine deficiency can cause intellectual disability. Organoiodine compounds occur in humans in some of the glands, especially the thyroid gland, as well as the stomach, epidermis, and immune system. Foods containing iodine include cod, oysters, shrimp, herring, lobsters, sunflower seeds, seaweed, and mushrooms. However, iodine is not known to have a biological role in plants. There are typically 0.06 milligrams per liter of iodine in human blood, 300 parts per billion of iodine in human bones, and 50 to 700 parts per billion of iodine in human tissue. There are 10 to 20 milligrams of iodine in a typical 70-kilogram human.[8]

Astatine, although very scarce, has been found in micrograms in the earth.[8] It has no known biological role because of its high radioactivity, extreme rarity, and has a half-life of just about 8 hours for the most stable isotope.

Tennessine is purely man-made and has no other roles in nature.

Toxicity

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The halogens tend to decrease in toxicity towards the heavier halogens.[42]

Fluorine gas is extremely toxic; breathing in fluorine at a concentration of 25 parts per million is potentially lethal. Hydrofluoric acid is also toxic, being able to penetrate skin and cause highly painful burns. In addition, fluoride anions are toxic, but not as toxic as pure fluorine. Fluoride can be lethal in amounts of 5 to 10 grams. Prolonged consumption of fluoride above concentrations of 1.5 mg/L is associated with a risk of dental fluorosis, an aesthetic condition of the teeth.[43] At concentrations above 4 mg/L, there is an increased risk of developing skeletal fluorosis, a condition in which bone fractures become more common due to the hardening of bones. Current recommended levels in water fluoridation, a way to prevent dental caries, range from 0.7 to 1.2 mg/L to avoid the detrimental effects of fluoride while at the same time reaping the benefits.[44] People with levels between normal levels and those required for skeletal fluorosis tend to have symptoms similar to arthritis.[8]

Chlorine gas is highly toxic. Breathing in chlorine at a concentration of 3 parts per million can rapidly cause a toxic reaction. Breathing in chlorine at a concentration of 50 parts per million is highly dangerous. Breathing in chlorine at a concentration of 500 parts per million for a few minutes is lethal. In addition, breathing in chlorine gas is highly painful because of its corrosive properties. Hydrochloric acid is the acid of chlorine, while relatively nontoxic, it is highly corrosive and releases very irritating and toxic hydrogen chloride gas in open air.[42]

Pure bromine is somewhat toxic but less toxic than fluorine and chlorine. One hundred milligrams of bromine is lethal.[8] Bromide anions are also toxic, but less so than bromine. Bromide has a lethal dose of 30 grams.[8]

Iodine is somewhat toxic, being able to irritate the lungs and eyes, with a safety limit of 1 milligram per cubic meter. When taken orally, 3 grams of iodine can be lethal. Iodide anions are mostly nontoxic, but these can also be deadly if ingested in large amounts.[8]

Astatine is radioactive and thus highly dangerous, but it has not been produced in macroscopic quantities and hence it is most unlikely that its toxicity will be of much relevance to the average individual.[8]

Tennessine cannot be chemically investigated due to how short its half-life is, although its radioactivity would make it very dangerous.

Superhalogen

[edit] Main article: Superatom

Certain aluminium clusters have superatom properties. These aluminium clusters are generated as anions (Al−n with n = 1, 2, 3, ... ) in helium gas and reacted with a gas containing iodine. When analyzed by mass spectrometry one main reaction product turns out to be Al13I−.[45] These clusters of 13 aluminium atoms with an extra electron added do not appear to react with oxygen when it is introduced in the same gas stream. Assuming each atom liberates its 3 valence electrons, this means 40 electrons are present, which is one of the magic numbers for sodium and implies that these numbers are a reflection of the noble gases.

Calculations show that the additional electron is located in the aluminium cluster at the location directly opposite from the iodine atom. The cluster must therefore have a higher electron affinity for the electron than iodine and therefore the aluminium cluster is called a superhalogen (i.e., the vertical electron detachment energies of the moieties that make up the negative ions are larger than those of any halogen atom).[46] The cluster component in the Al13I− ion is similar to an iodide ion or a bromide ion. The related Al13I−2 cluster is expected to behave chemically like the triiodide ion.[47][48]

See also

[edit] Look up halogen in Wiktionary, the free dictionary.

• Halogen bond
• Halogen addition reaction
• Halogen lamp
• Halogenation
• Interhalogen
• Pseudohalogen

Notes

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1. ^ This could also be the case for group 12, although copernicium's melting and boiling points are still uncertain.
2. ^ The number given in parentheses refers to the measurement uncertainty. This uncertainty applies to the least significant figure(s) of the number prior to the parenthesized value (i.e., counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, while 1.00794(72) stands for 1.00794±0.00072.[28]
3. ^ The average atomic weight of this element changes depending on the source of the chlorine, and the values in brackets are the upper and lower bounds.[29]
4. ^ a b The element does not have any stable nuclides, and the value in brackets indicates the mass number of the longest-lived isotope of the element.[29]

References

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1. ^ Jones, Daniel (2017) [1917]. Peter Roach; James Hartmann; Jane Setter (eds.). English Pronouncing Dictionary. Cambridge: Cambridge University Press. ISBN 978-3-12-539683-8.
2. ^ "Halogen". Merriam-Webster.com Dictionary. Merriam-Webster.
3. ^ "Halogen". Dictionary.com Unabridged (Online). n.d.
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5. ^ Union internationale de chimie pure et appliquée, ed. (2005). Nomenclature of inorganic chemistry: IUPAC Recommendations 2005. Cambridge: Royal Society of Chemistry. p. 51. ISBN 978-0-85404-438-2.
6. ^ "Chemical properties of the halogens - Group 17 - the halogens - Edexcel - GCSE Combined Science Revision - Edexcel". BBC Bitesize. Retrieved 2022-03-21.
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9. ^ Oganessian, Yu.Ts.; Abdullin, F.Sh.; Bailey, P.D.; Benker, D.E.; Bennett, M.E.; Dmitriev, S.N.; et al. (2010). "Synthesis of a new element with atomic number Z = 117". Physical Review Letters. 104 (14): 142502. Bibcode:2010PhRvL.104n2502O. doi:10.1103/PhysRevLett.104.142502. PMID 20481935. S2CID 3263480.{{cite journal}}: CS1 maint: article number as page number (link)
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11. ^ Snelders, H. A. M. (1971). "J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter". Isis. 62 (3): 328–338. doi:10.1086/350763. JSTOR 229946. S2CID 170337569.
12. ^ In 1826, Berzelius coined the terms Saltbildare (salt-formers) and Corpora Halogenia (salt-making substances) for the elements chlorine, iodine, and fluorine. See: Berzelius, Jacob (1826). "Årsberättelser om Framstegen i Physik och Chemie" [Annual Report on Progress in Physics and Chemistry]. Arsb. Vetensk. Framsteg (in Swedish). 6. Stockholm, Sweden: P.A. Norstedt & Söner: 187. From p. 187: "De förre af dessa, d. ä. de electronegativa, dela sig i tre klasser: 1) den första innehåller kroppar, som förenade med de electropositiva, omedelbart frambringa salter, hvilka jag derför kallar Saltbildare (Corpora Halogenia). Desse utgöras af chlor, iod och fluor *)." (The first of them [i.e., elements], the electronegative [ones], are divided into three classes: 1) The first includes substances which, [when] united with electropositive [elements], immediately produce salts, and which I therefore name "salt-formers" (salt-producing substances). These are chlorine, iodine, and fluorine *).)
13. ^ The word "halogen" appeared in English as early as 1832 (or earlier). See, for example: Berzelius, J.J. with A.D. Bache, trans., (1832) "An essay on chemical nomenclature, prefixed to the treatise on chemistry," The American Journal of Science and Arts, 22: 248–276 ; see, for example p. 263.
14. ^ Page 43, Edexcel International GCSE chemistry revision guide, Curtis 2011
15. ^ Greenwood & Earnshaw 1997, p. 804.
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17. ^ Jim Clark (2002). "THE ACIDITY OF THE HYDROGEN HALIDES". Retrieved February 24, 2013.
18. ^ "Facts about hydrogen fluoride". 2005. Archived from the original on 2013-02-01. Retrieved 2017-10-28.
19. ^ "Hydrogen chloride". Retrieved February 24, 2013.
20. ^ "Hydrogen bromide". Retrieved February 24, 2013.
21. ^ "Poison Facts:Low Chemicals: Hydrogen Iodid". Retrieved 2015-04-12.
22. ^ a b Saxena, P. B (2007). Chemistry Of Interhalogen Compounds. Discovery Publishing House. ISBN 978-81-8356-243-0. Retrieved February 27, 2013.
23. ^ Gribble, G. W (2009). Naturally Occurring Organohalogen Compounds - A Comprehensive Update. Springer. ISBN 978-3-211-99322-4. Retrieved April 23, 2022.
24. ^ "The Oxidising Ability of the Group 7 Elements". Chemguide.co.uk. Retrieved 2011-12-29.
25. ^ "Solubility of chlorine in water". Resistoflex.com. Archived from the original on 2012-04-23. Retrieved 2011-12-29.
26. ^ "Properties of bromine". bromaid.org. Archived from the original on December 8, 2007.
27. ^ "Iodine MSDS". Hazard.com. 1998-04-21. Retrieved 2011-12-29.
28. ^ "Standard Uncertainty and Relative Standard Uncertainty". CODATA reference. National Institute of Standards and Technology. Retrieved 26 September 2011.
29. ^ a b c Wieser, Michael E.; Coplen, Tyler B. (2011). "Atomic weights of the elements 2009 (IUPAC Technical Report)" (PDF). Pure Appl. Chem. 83 (2): 359–396. doi:10.1351/PAC-REP-10-09-14. S2CID 95898322. Retrieved 5 December 2012.
30. ^ a b Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press.
31. ^ Slater, J. C. (1964). "Atomic Radii in Crystals". Journal of Chemical Physics. 41 (10): 3199–3205. Bibcode:1964JChPh..41.3199S. doi:10.1063/1.1725697.
32. ^ Bonchev, Danail; Kamenska, Verginia (1981). "Predicting the properties of the 113–120 transactinide elements". The Journal of Physical Chemistry. 85 (9): 1177–86. Bibcode:1981JPhCh..85.1177B. doi:10.1021/j150609a021.
33. ^ Rothe, S.; Andreyev, A. N.; Antalic, S.; Borschevsky, A.; Capponi, L.; Cocolios, T. E.; De Witte, H.; Eliav, E.; et al. (2013). "Measurement of the First Ionization Potential of Astatine by Laser Ionization Spectroscopy". Nature Communications. 4 1835: 1–6. Bibcode:2013NatCo...4.1835R. doi:10.1038/ncomms2819. PMC 3674244. PMID 23673620.
34. ^ "Get Facts About the Element Astatine". www.thoughtco.com. Retrieved November 12, 2021.
35. ^ a b c d e f "How Much Do You Know About the Element Tennessine?". www.thoughtco.com. Retrieved November 12, 2021.
36. ^ "WebElements Periodic Table » Tennessine » properties of free atoms". www.webelements.com. Retrieved November 12, 2021.
37. ^ Morss, Lester R.; Edelstein, Norman M.; Fuger, Jean (2011). Morss, Lester R; Edelstein, Norman M; Fuger, Jean (eds.). The Chemistry of the Actinide and Transactinide Elements. Dordrecht, The Netherlands: Springer Science+Business Media. Bibcode:2011tcot.book.....M. doi:10.1007/978-94-007-0211-0. ISBN 978-94-007-0210-3.
38. ^ "Краткий справочник физико-химических величин Равделя, Л.: Химия, 1974 г. – 200 стр. \\ стр 67 табл. 24" (PDF).
39. ^ "The Halogen Lamp". Edison Tech Center. Retrieved 2014-09-05.
40. ^ Thomas, G. (2000). Medicinal Chemistry an Introduction. John Wiley & Sons, West Sussex, UK. ISBN 978-0-470-02597-0.
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42. ^ a b Gray, Theodore (2010). The Elements. Running Press. ISBN 978-1-57912-895-1.
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Bibliography

[edit]

• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.

v t e Periodic table
Periodic table forms	Alternatives Extended periodic table
Sets of elements	By periodic table structure Groups 1 (Hydrogen and alkali metals) 2 (Alkaline earth metals) 3 4 5 6 7 8 9 10 11 12 13 (Triels) 14 (Tetrels) 15 (Pnictogens) 16 (Chalcogens) 17 (Halogens) 18 (Noble gases) Periods 1 2 3 4 5 6 7 8+ Aufbau Fricke Pyykkö Blocks Aufbau principle By metallicity Metals Lanthanides Actinides Transition metals Post-transition metals Metalloids Lists of metalloids by source Dividing line Nonmetals Noble gases Other sets Platinum-group metals (PGM) Rare-earth elements Refractory metals Precious metals Coinage metals Noble metals Heavy metals Native metals Transuranium elements Superheavy elements Major actinides Minor actinides
Elements	Lists By: Abundance (in humans) Atomic properties Nuclear stability Symbol Properties Aqueous chemistry Crystal structure Electron configuration Electronegativity Goldschmidt classification Term symbol Data pages Abundance Atomic radius Boiling point Critical point Density Elasticity Electrical resistivity Electron affinity Electron configuration Electronegativity Hardness Heat capacity Heat of fusion Heat of vaporization Ionization energy Melting point Oxidation state Speed of sound Thermal conductivity Thermal expansion coefficient Vapor pressure
History	Element discoveries Dmitri Mendeleev 1871 table 1869 predictions Naming etymology controversies for places for people in East Asian languages
See also	IUPAC nomenclature systematic element name Trivial name Dmitri Mendeleev
Category WikiProject

v t e Periodic table
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 1 H He 2 Li Be B C N O F Ne 3 Na Mg Al Si P S Cl Ar 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 7 Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
s-block f-block d-block p-block

v t e Halogens

Fluorine F Atomic Number: 9 Atomic Weight: 18.9984032 Melting Point: 53.63 KBoiling Point: 85.03 KSpecific mass: 0.001696 g/cm3 Electronegativity: 3.98 Chlorine Cl Atomic Number: 17 Atomic Weight: 35.453 Melting Point: 172.31 K Boiling Point: 239.11 KSpecific mass: 0.003214 g/cm3 Electronegativity: 3.16 Bromine Br Atomic Number: 35 Atomic Weight: 79.904 Melting Point: 266.05 K Boiling Point: 332.0 KSpecific mass: 3.122 g/cm3 Electronegativity: 2.96 Iodine I Atomic Number: 53 Atomic Weight: 126.90447 Melting Point: 386.65 K Boiling Point: 475.4 KSpecific mass: 4.93 g/cm3 Electronegativity: 2.66 Astatine At Atomic Number: 85 Atomic Weight: [210] Melting Point: 575.15 K Boiling Point: 610 KSpecific mass: 7 g/cm3 Electronegativity: 2.2 Tennessine Ts Atomic Number: 117 Atomic Weight: [294] Melting Point: ? 573–773 K Boiling Point: ? 823 K Specific mass: ? g/cm3 Electronegativity: ?

v t e Salts and covalent derivatives of the fluoride ion

HF ?HeF2 LiF BeF2 BFBF3B2F4+BO3 CF4CxFy+CO3 NF3FN3N2F2NFN2F4NF2?NF5+N+NO3 OF2O2F2OFO3F2O4F2?OF4 F2 Ne NaF MgF2 AlFAlF3 SiF4 P2F4PF3PF5+PO4 S2F2SF2S2F4SF3SF4S2F10SF6+SO4 ClFClF3ClF5 ?ArF2?ArF4 KF CaFCaF2 ScF3 TiF2TiF3TiF4 VF2VF3VF4VF5 CrF2CrF3CrF4CrF5?CrF6 MnF2MnF3MnF4?MnF5 FeF2FeF3FeF4 CoF2 CoF3 CoF4 NiF2NiF3NiF4 CuFCuF2?CuF3 ZnF2 GaF2GaF3 GeF2GeF4 AsF3AsF5 Se2F2SeF4SeF6+SeO3 BrFBrF3BrF5 KrF2?KrF4?KrF6 RbF SrFSrF2 YF3 ZrF2ZrF3ZrF4 NbF4NbF5 MoF4MoF5MoF6 TcF4TcF5 TcF6 RuF3RuF4RuF5RuF6 RhF3RhF4RhF5RhF6 PdF2Pd[PdF6]PdF4?PdF6 Ag2FAgFAgF2AgF3 CdF2 InFInF3 SnF2SnF4 SbF3SbF5 TeF4?Te2F10TeF6+TeO3 IFIF3IF5IF7+IO3 XeF2XeF4XeF6?XeF8 CsF BaF2   LuF3 HfF4 TaF5 WF4WF5WF6 ReF4ReF5ReF6ReF7 OsF4OsF5OsF6?OsF7?OsF8 IrF2IrF3IrF4IrF5IrF6 PtF2Pt[PtF6]PtF4PtF5PtF6 AuFAuF3Au2F10?AuF6AuF5•F2 Hg2F2HgF2?HgF4 TlFTlF3 PbF2PbF4 BiF3BiF5 PoF2PoF4PoF6 AtF?AtF3?AtF5 RnF2?RnF4?RnF6 FrF RaF2   LrF3 Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og ↓ LaF3 CeF3CeF4 PrF3 PrF4 NdF2 NdF3 NdF4 PmF3 SmF SmF2 SmF3 EuF2 EuF3 GdF3 TbF3 TbF4 DyF2 DyF3 DyF4 HoF3 ErF3 TmF2 TmF3 YbF2 YbF3 AcF3 ThF2ThF3ThF4 PaF4PaF5 UF3UF4UF5UF6 NpF3NpF4NpF5NpF6 PuF3PuF4PuF5PuF6 AmF2AmF3AmF4?AmF6 CmF3CmF4 ?CmF6 BkF3 BkF4 CfF3 CfF4 EsF3 EsF4?EsF6 Fm MdF3 No

v t e Salts and covalent derivatives of the chloride ion

HCl He LiCl BeCl2 B4Cl4B12Cl12BCl3B2Cl4+BO3 C2Cl2C2Cl4C2Cl6CCl4+C+CO3 NCl3ClN3+N+NO3 ClxOyCl2OCl2O2ClOClO2Cl2O4Cl2O6Cl2O7ClO4+O ClFClF3ClF5 Ne NaCl MgCl2 AlClAlCl3 Si5Cl12Si2Cl6SiCl4 P2Cl4PCl3PCl5+P S2Cl2SCl2SCl4+SO4 Cl2 Ar KCl CaClCaCl2 ScCl3 TiCl2TiCl3TiCl4 VCl2VCl3VCl4VCl5 CrCl2CrCl3CrCl4 MnCl2MnCl3 FeCl2FeCl3 CoCl2CoCl3 NiCl2 CuClCuCl2 ZnCl2 GaClGaCl3 GeCl2GeCl4 AsCl3AsCl5+As Se2Cl2SeCl2SeCl4 BrCl Kr RbCl SrCl2 YCl3 ZrCl2ZrCl3ZrCl4 NbCl3NbCl4NbCl5 MoCl2MoCl3MoCl4MoCl5MoCl6 TcCl3TcCl4 RuCl2RuCl3RuCl4 RhCl3 PdCl2 AgCl CdCl2 InClInCl2InCl3 SnCl2SnCl4 SbCl3SbCl5 Te3Cl2TeCl2TeCl4 IClICl3 XeClXeCl2XeCl4 CsCl BaCl2 * LuCl3177LuCl3 HfCl4 TaCl3TaCl4TaCl5 WCl2WCl3WCl4WCl5WCl6 ReCl3ReCl4ReCl5ReCl6 OsCl2OsCl3OsCl4OsCl5 IrCl2IrCl3IrCl4 PtCl2PtCl4PtCl2−6 AuCl(Au[AuCl4])2AuCl3AuCl−4 Hg2Cl2HgCl2 TlClTlCl3 PbCl2PbCl4 BiCl3 PoCl2PoCl4 AtCl Rn FrCl RaCl2 ** LrCl3 RfCl4 DbCl5 SgO2Cl2 BhO3Cl Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og   * LaCl3 CeCl3 PrCl3 NdCl2NdCl3 PmCl3 SmCl2SmCl3 EuCl2EuCl3 GdCl3 TbCl3 DyCl2DyCl3 HoCl3 ErCl3 TmCl2TmCl3 YbCl2YbCl3 ** AcCl3 ThCl3ThCl4 PaCl4PaCl5 UCl3UCl4UCl5UCl6 NpCl3NpCl4 PuCl3PuCl4PuCl2−6 AmCl2AmCl3 CmCl3 BkCl3 CfCl3CfCl2 EsCl2EsCl3 FmCl2 MdCl2 NoCl2

v t e Salts and covalent derivatives of the bromide ion

HBr He LiBr BeBr2 BBr3+BO3 CBr4+C NBr3BrN3NH4BrNOBr+N Br2OBrO2Br2O3Br2O5 BrFBrF3BrF5 Ne NaBr MgBr2 AlBrAlBr3 SiBr4 PBr3PBr5PBr7+P S2Br2SBr2 BrCl Ar KBr CaBr2 ScBr3 TiBr2TiBr3TiBr4 VBr2VBr3 CrBr2CrBr3CrBr4 MnBr2 FeBr2FeBr3 CoBr2 NiBr2 CuBrCuBr2 ZnBr2 GaBr3 GeBr2GeBr4 AsBr3+As+AsO3 SeBr2SeBr4 Br2 Kr RbBr SrBr2 YBr3 ZrBr2ZrBr3ZrBr4 NbBr5 MoBr2MoBr3MoBr4 TcBr3TcBr4 RuBr3 RhBr3 PdBr2 AgBr CdBr2 InBrInBr3 SnBr2SnBr4 SbBr3+Sb-Sb Te2BrTeBr4+Te IBrIBr3 XeBr2 CsBr BaBr2 * LuBr3 HfBr4 TaBr5 WBr5WBr6 ReBr3 OsBr3OsBr4 IrBr3IrBr4 PtBr2PtBr4 AuBrAuBr3 Hg2Br2 HgBr2 TlBr PbBr2 BiBr3 PoBr2PoBr4 AtBr Rn FrBr RaBr2 ** Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og   * LaBr3 CeBr3 PrBr3 NdBr2 NdBr3 PmBr3 SmBr2 SmBr3 EuBr2 EuBr3 GdBr3 TbBr3 DyBr3 HoBr3 ErBr3 TmBr2 TmBr3 YbBr2 YbBr3 ** AcBr3 ThBr4 PaBr4PaBr5 UBr3UBr4UBr5 NpBr3NpBr4 PuBr3 AmBr2AmBr3 CmBr3 BkBr3 CfBr3 EsBr2EsBr3 Fm Md No

v t e Salts and covalent derivatives of the iodide ion

HI+H He LiI BeI2 BI3+BO3 CI4+C NI3NH4I+N I2O4I2O5I2O6I4O9 IFIF3IF5IF7 Ne NaI MgI2 AlIAlI3 SiI4 PI3P2I4+PPI5 S2I2 IClICl3 Ar KI CaI2 ScI3 TiI2TiI3TiI4 VI2VI3 CrI2CrI3CrI4 MnI2 FeI2FeI3 CoI2 NiI2-Ni CuI ZnI2 GaIGaI3 GeI2GeI4+Ge AsI3As2I4+As Se IBrIBr3 Kr RbIRbI3 SrI2 YI3 ZrI2ZrI3ZrI4 NbI4NbI5 MoI2MoI3 TcI3 RuI3 RhI3 PdI2 AgI CdI2 InIInI3 SnI2SnI4 SbI3+Sb TeI4+Te I−I−3 Xe CsICsI3 BaI2   LuI3 HfI3HfI4 TaI4TaI5 WI2WI3WI4 ReI3ReI4 OsIOsI2OsI3 IrI3IrI4 PtI2PtI4 AuIAuI3 Hg2I2HgI2 TlITlI3 PbI2 BiI3 PoI2PoI4 AtI Rn FrI RaI2   Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og ↓ LaI2LaI3 CeI2CeI3 PrI2PrI3 NdI2NdI3 PmI3 SmI2SmI3 EuI2EuI3 GdI2GdI3 TbI3 DyI2DyI3 HoI3 ErI3 TmI2TmI3 YbI2YbI3 AcI3 ThI2ThI3ThI4 PaI3PaI4PaI5 UI3UI4 NpI3 PuI3 AmI2AmI3 CmI3 BkI3 CfI2CfI3 EsI2EsI3 Fm Md No

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