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Legend

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| style="border: ; background:; padding:0 2px;" |primordial element

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| style="border: ; background:;" |element from decay

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| style="border:; background:;" | Synthetic

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alt=Halogens|thumb|Astatine is here represented by [[Uraninite (uranium ore mostly composed of uranium dioxide (UO<sub>2</sub>), which contains <sup>238</sup>U, which then decays into astatine among other things. Chlorine and iodine samples were prepared in the lab, while all other samples were obtained commercially from online sellers (e.g. Luciteria, NovaElements, etc...).]]

The halogens () 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 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.

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.

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. 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

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. In 1886, Henri Moissan, a chemist in Paris, performed electrolysis on potassium bifluoride dissolved in anhydrous hydrogen fluoride, and successfully isolated fluorine.

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.

Etymology

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. Davy's name for the element prevailed. 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.

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".

{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"

|+ style="margin-bottom: 5px;" | Halogen bond energies (kJ/mol)

|-

! X

! X<sub>2</sub>

! HX

! BX<sub>3</sub>

! AlX<sub>3</sub>

! CX<sub>4</sub>

|-

! scope="row" style="background:#ff9;"| F

| style="background:#ff9;"| 159

| style="background:#ff9;"| 574

| style="background:#ff9;"| 645

| style="background:#ff9;"| 582

| style="background:#ff9;"| 456

|-

! scope="row" | Cl

|243

|428

|444

|427

|327

|-

! scope="row" | Br

|193

|363

|368

|360

|272

|-

! scope="row" | 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 (SiF<sub>4</sub>). 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

Diatomic halogen molecules

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".

{| class="wikitable" style="margin: 1em auto 1em auto; text-align:center;"

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! halogen || molecule || structure || model || d(X−X) / pm<br />(gas phase) || d(X−X) / pm<br />(solid phase)

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| fluorine || F<sub>2</sub> || class=skin-invert|45px || 45px || 143 || 149

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| chlorine || Cl<sub>2</sub> || class=skin-invert|70px || 63px || 199 || 198

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| bromine || Br<sub>2</sub> || class=skin-invert|80px || 72px || 228 || 227

|-

| iodine || I<sub>2</sub> || class=skin-invert|70px || 84px || 266 || 272

<!--don't add astatine; it's not proven to form At2 molecules-->

|}

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

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:

: H<sub>2</sub> + X<sub>2</sub> → 2HX

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

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.

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. 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. 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. Hydrogen iodide, like other hydrogen halides, is toxic.

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:

:

Interhalogens are typically more reactive than all diatomic halogen molecules except F<sub>2</sub> 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 BrF<sub>3</sub>, and many iodine-containing interhalogens are solids.

Polyhalogenated compounds

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

Reactions with water

Fluorine reacts vigorously with water to produce oxygen (O<sub>2</sub>) and hydrogen fluoride (HF):

:

Chlorine has maximum solubility of ca. 7.1 g Cl<sub>2</sub> per kg of water at ambient temperature (21&nbsp;°C). Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid, a solution that can be used as a disinfectant or bleach:

:

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

:

Iodine, however, is minimally soluble in water (0.03 g/100 g water at 20&nbsp;°C) and does not react with it. 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

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.

{| class="wikitable" style="text-align:center"

!Halogen

!Standard atomic weight(Da)

!Melting point(K)

!Melting point(°C)

!Boiling point(K)

!Boiling point(°C)

|-

| Fluorine || 18.9984032(5) || 53.53 || −219.62 || 85.03 || −188.12 || 0.0017 || 3.98 || 1681.0 || 71

|-

| Chlorine || [35.446; 35.457]