Editing Metalloid (section)

==Elements commonly recognised as metalloids==
:''Properties noted in this section refer to the elements in their most thermodynamically stable forms under ambient conditions.''

===Boron===
{{Main|Boron}}
[[File:Boron R105.jpg|thumb|right|Boron, shown here in the form of its β-[[rhombohedral]] phase (its most thermodynamically stable [[allotrope]])<ref>[[#VanSetten2007|Van Setten et al. 2007, pp.&nbsp;2460–61]]; [[#Oganov2009|Oganov et al. 2009, pp.&nbsp;863–64]]</ref>|alt=Several dozen small angular stone like shapes, grey with scattered silver flecks and highlights.]] Pure boron is a shiny, silver-grey crystalline solid.<ref>[[#Housecroft2008|Housecroft & Sharpe 2008, p.&nbsp;331]]; [[#Oganov2010|Oganov 2010, p.&nbsp;212]]</ref> It is less dense than aluminium (2.34 vs. 2.70 g/cm<sup>3</sup>), and is hard and brittle. It is barely reactive under normal conditions, except for attack by [[fluorine]],<ref>[[#Housecroft2008|Housecroft & Sharpe 2008, p.&nbsp;333]]</ref> and has a melting point of 2076&nbsp;°C (cf. steel ~1370&nbsp;°C).<ref>[[#Kross|Kross 2011]]</ref> Boron is a semiconductor;<ref>[[#Berger1997|Berger 1997, p.&nbsp;37]]</ref> its room temperature electrical conductivity is 1.5 × 10<sup>−6</sup> [[Siemens (unit)|S]]•cm<sup>−1</sup><ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;144]]</ref> (about 200 times less than that of tap water)<ref>[[#Kopp|Kopp, Lipták & Eren 2003, p.&nbsp;221]]</ref> and it has a band gap of about 1.56&nbsp;eV.<ref>[[#Prudenziati1977|Prudenziati 1977, p.&nbsp;242]]</ref>{{refn|1=Boron, at 1.56 eV, has the largest band gap amongst the commonly recognised (semiconducting) metalloids. Of nearby elements in periodic table terms, selenium has the next highest band gap (close to 1.8 eV) followed by white phosphorus (around 2.1 eV).<ref>[[#Berger1997|Berger 1997, pp.&nbsp;84, 87]]</ref>|group=n}} Mendeleev commented that, "Boron appears in a free state in several forms which are intermediate between the metals and the nonmmetals."<ref>[[#Mendeléeff1897a|Mendeléeff 1897, p.&nbsp;57]]</ref>
The structural chemistry of boron is dominated by its small atomic size, and relatively high ionization energy. With only three valence electrons per boron atom, simple covalent bonding cannot fulfil the octet rule.<ref name="Rayner-Canham 2006, p. 291">[[#Rayner2006|Rayner-Canham & Overton 2006, p.&nbsp;291]]</ref> Metallic bonding is the usual result among the heavier congenors of boron but this generally requires low ionization energies.<ref>[[#Siekierski2002|Siekierski & Burgess 2002, p.&nbsp;63]]</ref> Instead, because of its small size and high ionization energies, the basic structural unit of boron (and nearly all of its allotropes){{refn|1=The synthesis of B<sub>40</sub> [[borospherene]], a "distorted fullerene with a hexagonal hole on the top and bottom and four heptagonal holes around the waist" was announced in 2014.<ref>[[#Wogan|Wogan 2014]]</ref>|group=n}} is the icosahedral B<sub>12</sub> cluster. Of the 36 electrons associated with 12 boron atoms, 26 reside in 13 delocalized molecular orbitals; the other 10 electrons are used to form two- and three-centre covalent bonds between icosahedra.<ref>[[#Siekierski2002|Siekierski & Burgess 2002, p.&nbsp;86]]</ref> The same motif can be seen, as are [[deltahedron|deltahedral]] variants or fragments, in metal borides and hydride derivatives, and in some halides.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;141]]; [[#Henderson2000|Henderson 2000, p.&nbsp;58]]; [[#Housecroft2008|Housecroft & Sharpe 2008, pp.&nbsp;360–72]]</ref>

The bonding in boron has been described as being characteristic of behaviour intermediate between metals and nonmetallic [[covalent network]] solids (such as [[diamond]]).<ref>[[#Parry1970|Parry et al. 1970, pp.&nbsp;438, 448–51]]</ref> The energy required to transform B, C, N, Si, and P from nonmetallic to metallic states has been estimated as 30, 100, 240, 33, and 50 kJ/mol, respectively. This indicates the proximity of boron to the metal-nonmetal borderline.<ref name=Fehlner1990>[[#Fehlner1990|Fehlner 1990, p.&nbsp;202]]</ref>

Most of the chemistry of boron is nonmetallic in nature.<ref name=Fehlner1990/> Unlike its heavier congeners, it is not known to form a simple B<sup>3+</sup> or hydrated [B(H<sub>2</sub>O)<sub>4</sub>]<sup>3+</sup> cation.<ref>[[#Owen|Owen & Brooker 1991, p.&nbsp;59]]; [[#Wiberg2001|Wiberg 2001, p.&nbsp;936]]</ref> The small size of the boron atom enables the preparation of many [[interstitial compound|interstitial]] alloy-type borides.<ref name=Greenwood145>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;145]]</ref> Analogies between boron and transition metals have been noted in the formation of [[complex (chemistry)|complexes]],<ref>[[#Houghton1979|Houghton 1979, p.&nbsp;59]]</ref> and [[adduct]]s (for example, BH<sub>3</sub> + [[Carbon monoxide|CO]] →BH<sub>3</sub>CO and, similarly, Fe(CO)<sub>4</sub> + CO →Fe(CO)<sub>5</sub>),{{refn|1=The BH<sub>3</sub> and Fe(CO<sub>4</sub>) species in these reactions are short-lived [[reaction intermediate]]s.<ref>[[#Fehlner1990|Fehlner 1990, p.&nbsp;205]]</ref>|group=n}} as well as in the geometric and electronic structures of [[cluster compound|cluster species]] such as [B<sub>6</sub>H<sub>6</sub>]<sup>2−</sup> and [Ru<sub>6</sub>(CO)<sub>18</sub>]<sup>2−</sup>.<ref>[[#Fehlner1990|Fehlner 1990, pp.&nbsp;204–05, 207]]</ref>{{refn|1=On the analogy between boron and metals, Greenwood<ref>[[#Greenwood2001|Greenwood 2001, p.&nbsp;2057]]</ref> commented that: "The extent to which metallic elements mimic boron (in having fewer electrons than orbitals available for bonding) has been a fruitful cohering concept in the development of metalloborane chemistry&nbsp;... Indeed, metals have been referred to as "honorary boron atoms" or even as "flexiboron atoms". The converse of this relationship is clearly also valid&nbsp;..."|group=n}} The aqueous chemistry of boron is characterised by the formation of many different [[Borate#Polymeric ions|polyborate anions]].<ref>[[#Salentine1987|Salentine 1987, pp.&nbsp;128–32]]; [[#MacKay2002|MacKay, MacKay & Henderson 2002, pp.&nbsp;439–40]]; [[#Kneen1972|Kneen, Rogers & Simpson 1972, p.&nbsp;394]]; [[#Hiller1960|Hiller & Herber 1960, inside front cover; p.&nbsp;225]]</ref> Given its high charge-to-size ratio, boron bonds covalently in nearly all of its compounds;<ref>[[#Sharp1983|Sharp 1983, p.&nbsp;56]]</ref> the exceptions are the [[boride]]s as these include, depending on their composition, covalent, ionic, and metallic bonding components.<ref>[[#Fokwa|Fokwa 2014, p.&nbsp;10]]</ref>{{refn|1=The bonding in [[boron trifluoride]], a gas, has been referred to as predominately ionic<ref name=Gillespie1998>[[#Gillespie1998|Gillespie 1998]]</ref> a description which was subsequently described as misleading.<ref name=Haaland>[[#Haaland|Haaland et al. 2000]]</ref>|group=n}} Simple binary compounds, such as [[boron trichloride]] are [[Lewis acid]]s as the formation of three covalent bonds leaves a hole in the [[octet rule|octet]] which can be filled by an electron-pair donated by a [[Lewis base]].<ref name="Rayner-Canham 2006, p. 291"/> Boron has a strong affinity for [[oxygen]] and a duly extensive [[borate]] chemistry.<ref name=Greenwood145/> The oxide B<sub>2</sub>O<sub>3</sub> is [[polymeric]] in structure,<ref name=Pudd59>[[#Puddephatt1989|Puddephatt & Monaghan 1989, p.&nbsp;59]]</ref> weakly acidic,<ref>[[#Mahan1965|Mahan 1965, p.&nbsp;485]]</ref>{{refn|1=Boron trioxide B<sub>2</sub>O<sub>3</sub> is sometimes described as being (weakly) [[amphoteric]].<ref>[[#Danaith|Danaith 2008, p.&nbsp;81]].</ref> It reacts with [[alkali]]es to give various borates.<ref>[[#Lidin|Lidin 1996, p.&nbsp;28]]</ref> In its [[hydrated]] form (as H<sub>3</sub>BO<sub>3</sub>, [[boric acid]]) it reacts with [[sulfur trioxide]], the [[anhydride]] of [[sulfuric acid]], to form a [[bisulfate]] B(HSO<sub>3</sub>) <sub>4</sub>.<ref>[[#Kondratev|Kondrat'ev & Mel'nikova 1978]]</ref> In its pure (anhydrous) form it reacts with [[phosphoric acid]] to form a "[[phosphate]]" BPO<sub>4</sub>.<ref>[[#Holderness|Holderness & Berry 1979, p.&nbsp;111]]; [[#Wiberg2001|Wiberg 2001, p.&nbsp;980]]</ref> The latter compound may be regarded as a [[mixed oxide]] of B<sub>2</sub>O<sub>3</sub> and [[P2O5|P<sub>2</sub>O<sub>5</sub>]].<ref>[[#Toy|Toy 1975, p.&nbsp;506]]</ref>|group=n}} and a glass former.<ref name=Rao22>[[#Rao2002|Rao 2002, p.&nbsp;22]]</ref> [[Organometallic chemistry|Organometallic compounds]] of boron{{refn|1=Organic derivatives of metalloids are traditionally counted as organometallic compounds.<ref>[[#Fehlner|Fehlner 1992, p.&nbsp;1]]</ref>|group=n}} have been known since the 19th century (see [[organoboron chemistry]]).<ref>[[#Haiduc1985|Haiduc & Zuckerman 1985, p.&nbsp;82]]</ref>

===Silicon===
{{Main|Silicon}}
[[File:SiliconCroda.jpg|thumb|left|[[Silicon]] has a blue-grey metallic [[lustre (mineralogy)|lustre]].|alt=A lustrous blue grey potato-shaped lump with an irregular corrugated surface.]]
Silicon is a crystalline solid with a blue-grey metallic lustre.<ref name=Greenwood331>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;331]]</ref> Like boron, it is less dense (at 2.33 g/cm<sup>3</sup>) than aluminium, and is hard and brittle.<ref>[[#Wiberg2001|Wiberg 2001, p.&nbsp;824]]</ref> It is a relatively unreactive element.<ref name=Greenwood331/> According to Rochow,<ref>[[#Rochow1973|Rochow 1973, pp.&nbsp;1337‒38]]</ref> the massive crystalline form (especially if pure) is "remarkably inert to all acids, including [[hydrofluoric acid|hydrofluoric]]".{{refn|1=In air, silicon forms a thin coating of amorphous silicon dioxide, 2 to 3 nm thick.<ref name=R393/> This coating is dissolved by [[hydrogen fluoride]] at a very low pace – on the order of two to three hours per nanometre.<ref>[[#Zhang|Zhang 2002, p.&nbsp;70]]</ref> Silicon dioxide, and silicate glasses (of which silicon dioxide is a major component), are otherwise readily attacked by hydrofluoric acid.<ref>[[#Sacks|Sacks 1998, p.&nbsp;287]]</ref>|group=n}} Less pure silicon, and the powdered form, are variously susceptible to attack by strong or heated acids, as well as by steam and fluorine.<ref>[[#Rochow1973|Rochow 1973, pp.&nbsp;1337, 1340]]</ref> Silicon dissolves in hot aqueous [[alkali]]s with the evolution of [[hydrogen]], as do metals<ref>[[#Allen1968|Allen & Ordway 1968, p.&nbsp;152]]</ref> such as beryllium, aluminium, zinc, gallium or indium.<ref>[[#Eagleson1994|Eagleson 1994, pp.&nbsp;48, 127, 438, 1194]]; [[#Massey2000|Massey 2000, p.&nbsp;191]]</ref> It melts at 1414&nbsp;°C. Silicon is a semiconductor with an electrical conductivity of 10<sup>−4</sup>&nbsp;S•cm<sup>−1</sup><ref>[[#Orton2004|Orton 2004, p.&nbsp;7]]. This is a typical value for high-purity silicon.</ref> and a band gap of about 1.11&nbsp;eV.<ref name=R393>[[#Russell2005|Russell & Lee 2005, p.&nbsp;393]]</ref> When it melts, silicon becomes a reasonable metal<ref>[[#Coles1976|Coles & Caplin 1976, p.&nbsp;106]]</ref> with an electrical conductivity of 1.0–1.3 × 10<sup>4</sup>&nbsp;S•cm<sup>−1</sup>, similar to that of liquid mercury.<ref>[[#Glazov1969|Glazov, Chizhevskaya & Glagoleva 1969, pp.&nbsp;59–63]]; [[#Allen1987|Allen & Broughton 1987, p.&nbsp;4967]]</ref>

The chemistry of silicon is generally nonmetallic (covalent) in nature.<ref>[[#Cotton1995|Cotton, Wilkinson & Gaus 1995, p.&nbsp;393]]</ref> It is not known to form a cation.<ref>[[#Wiberg2001|Wiberg 2001, p.&nbsp;834]]</ref>{{refn|1=The bonding in [[silicon tetrafluoride]], a gas, has been referred to as predominately ionic<ref name=Gillespie1998/> a description which was subsequently described as misleading.<ref name=Haaland/>|group=n}} Silicon can form alloys with metals such as iron and copper.<ref>[[#Partington1944|Partington 1944, p.&nbsp;723]]</ref> It shows fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox>[[#Cox2004|Cox 2004, p.&nbsp;27]]</ref> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225>[[#Hiller1960|Hiller & Herber 1960, inside front cover; p.&nbsp;225]]</ref> The high strength of the [[silicon–oxygen bond]] dominates the chemical behaviour of silicon.<ref>[[#Kneen1972|Kneen, Rogers and Simpson 1972, p.&nbsp;384]]</ref> Polymeric silicates, built up by tetrahedral SiO<sub>4</sub> units sharing their oxygen atoms, are the most abundant and important compounds of silicon.<ref name="Bailar513"/> The polymeric borates, comprising linked trigonal and tetrahedral BO<sub>3</sub> or BO<sub>4</sub> units, are built on similar structural principles.<ref>[[#Cotton1995|Cotton, Wilkinson & Gaus 1995, pp.&nbsp;319, 321]]</ref> The oxide SiO<sub>2</sub> is polymeric in structure,<ref name=Pudd59/> weakly acidic,<ref>[[#Smith1990|Smith 1990, p.&nbsp;175]]</ref>{{refn|1=Although SiO<sub>2</sub> is classified as an acidic oxide, and hence reacts with alkalis to give silicates, it reacts with phosphoric acid to yield a silicon oxide orthophosphate Si<sub>5</sub>O(PO<sub>4</sub>)<sub>6</sub>,<ref>[[#Poojary1993|Poojary, Borade & Clearfield 1993]]</ref> and with hydrofluoric acid to give [[hexafluorosilicic acid]] H<sub>2</sub>SiF<sub>6</sub>.<ref>[[#Wiberg2001|Wiberg 2001, pp.&nbsp;851, 858]]</ref> The latter reaction "is sometimes quoted as evidence of basic [that is, metallic] properties".<ref>[[#Barnett|Barmett & Wilson 1959, p.&nbsp;332]]</ref>|group=n}} and a glass former.<ref name=Rao22/> Traditional organometallic chemistry includes the carbon compounds of silicon (see [[organosilicon]]).<ref>[[#Powell1988|Powell 1988, p.&nbsp;1]]</ref>

===Germanium===
{{Main|Germanium}}
[[File:Polycrystalline-germanium.jpg|thumb|right|[[Germanium]] is sometimes described as a [[metal]]|alt=Greyish lustrous block with uneven cleaved surface.]]
Germanium is a shiny grey-white solid.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;371]]</ref> It has a density of 5.323 g/cm<sup>3</sup> and is hard and brittle.<ref>[[#Cusack1967|Cusack 1967, p.&nbsp;193]]</ref> It is mostly unreactive at room temperature{{refn|1=Temperatures above 400&nbsp;°C are required to form a noticeable surface oxide layer.<ref>[[#Russell2005|Russell & Lee 2005, pp.&nbsp;399–400]]</ref>|group=n}} but is slowly attacked by hot concentrated [[sulfuric acid|sulfuric]] or [[nitric acid]].<ref name=Greenwood373>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;373]]</ref> Germanium also reacts with molten [[sodium hydroxide|caustic soda]] to yield sodium germanate Na<sub>2</sub>GeO<sub>3</sub> and hydrogen gas.<ref>[[#Moody1991|Moody 1991, p.&nbsp;273]]</ref> It melts at 938&nbsp;°C. Germanium is a semiconductor with an electrical conductivity of around 2 × 10<sup>−2</sup>&nbsp;S•cm<sup>−1</sup><ref name=Greenwood373/> and a band gap of 0.67&nbsp;eV.<ref>[[#Russell2005|Russell & Lee 2005, p.&nbsp;399]]</ref> Liquid germanium is a metallic conductor, with an electrical conductivity similar to that of liquid mercury.<ref>[[#Berger1997|Berger 1997, pp.&nbsp;71–72]]</ref>

Most of the chemistry of germanium is characteristic of a nonmetal.<ref>[[#Jolly1966|Jolly 1966, pp.&nbsp;125–6]]</ref> Whether or not germanium forms a cation is unclear, aside from the reported existence of the Ge<sup>2+</sup> ion in a few esoteric compounds.{{refn|1=Sources mentioning germanium cations include: Powell & Brewer<ref>[[#Powell|Powell & Brewer 1938]]</ref> who state that the [[cadmium iodide]] CdI<sub>2</sub> structure of [[germanous iodide]] GeI<sub>2</sub> establishes the existence of the Ge<sup>++</sup> ion (the CdI<sub>2</sub> structure being found, according to Ladd,<ref>[[#Ladd|Ladd 1999, p.&nbsp;55]]</ref> in "many metallic halides, hydroxides, and chalcides"); Everest<ref>[[#Everest|Everest 1953, p.&nbsp;4120]]</ref> who comments that, "it seems probable that the Ge<sup>++</sup> ion can also occur in other crystalline germanous salts such as the [[germanous phosphite|phosphite]], which is similar to the salt-like [[stannous phosphite]] and [[germanous phosphate]], which resembles not only the stannous phosphates, but the [[manganous phosphate]]s also"; Pan, Fu & Huang<ref>[[#Pan|Pan, Fu and Huang 1964, p.&nbsp;182]]</ref> who presume the formation of the simple Ge<sup>++</sup> ion when Ge(OH)<sub>2</sub> is dissolved in a [[perchloric acid]] solution, on the basis that, "ClO4<sup>−</sup> has little tendency to enter [[coordination complex|complex]] formation with a cation"; Monconduit et al.<ref>[[#Monconduit|Monconduit et al. 1992]]</ref> who prepared the layer compound or phase Nb<sub>3</sub>Ge<sub>x</sub>Te<sub>6</sub> (x ≃ 0.9), and reported that this contained a Ge<sup>II</sup> cation; Richens<ref>[[#Richens|Richens 1997, p.&nbsp;152]]</ref> who records that, "Ge<sup>2+</sup> (aq) or possibly Ge(OH)<sup>+</sup>(aq) is said to exist in dilute air-free aqueous suspensions of the yellow hydrous monoxide…however both are unstable with respect to the ready formation of GeO<sub>2</sub>.''n''H<sub>2</sub>O"; Rupar et al.<ref>[[#Rupar|Rupar et al. 2008]]</ref> who synthesized a [[cryptand]] compound containing a Ge<sup>2+</sup> cation; and Schwietzer and Pesterfield<ref>[[#Schwietzer2010|Schwietzer & Pesterfield 2010, p.&nbsp;190]]</ref> who write that, "the monoxide [[GeO]] dissolves in dilute acids to give Ge<sup>+2</sup> and in dilute bases to produce GeO<sub>2</sub><sup>−2</sup>, all three entities being unstable in water". Sources dismissing germanium cations or further qualifying their presumed existence include: Jolly and Latimer<ref>[[#Jolly|Jolly & Latimer 1951, p.&nbsp;2]]</ref> who assert that, "the germanous ion cannot be studied directly because no germanium (II) species exists in any appreciable concentration in noncomplexing aqueous solutions"; Lidin<ref>[[#Lidin|Lidin 1996, p.&nbsp;140]]</ref> who says that, "[germanium] forms no aquacations"; Ladd<ref>[[#Ladd|Ladd 1999, p.&nbsp;56]]</ref> who notes that the CdI<sub>2</sub> structure is "intermediate in type between ionic and molecular compounds"; and Wiberg<ref>[[#Wiberg2001|Wiberg 2001, p.&nbsp;896]]</ref> who states that, "no germanium cations are known".|group=n}} It can form alloys with metals such as aluminium and [[gold]].<ref>[[#Schwartz2002|Schwartz 2002, p.&nbsp;269]]</ref> It shows fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox/> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225/> Germanium generally forms tetravalent (IV) compounds, and it can also form less stable divalent (II) compounds, in which it behaves more like a metal.<ref name="ReferenceC">[[#Eggins1972|Eggins 1972, p.&nbsp;66]]; [[#Wiberg2001|Wiberg 2001, p.&nbsp;895]]</ref> Germanium analogues of all of the major types of silicates have been prepared.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;383]]</ref> The metallic character of germanium is also suggested by the formation of various [[oxoacid]] salts. A phosphate [(HPO<sub>4</sub>)<sub>2</sub>Ge·H<sub>2</sub>O] and highly stable trifluoroacetate Ge(OCOCF<sub>3</sub>)<sub>4</sub> have been described, as have Ge<sub>2</sub>(SO<sub>4</sub>)<sub>2</sub>, Ge(ClO<sub>4</sub>)<sub>4</sub> and GeH<sub>2</sub>(C<sub>2</sub>O<sub>4</sub>)<sub>3</sub>.<ref>[[#Glockling1969|Glockling 1969, p.&nbsp;38]]; [[#Wells1984|Wells 1984, p.&nbsp;1175]]</ref> The oxide GeO<sub>2</sub> is polymeric,<ref name=Pudd59/> amphoteric,<ref>[[#Cooper1968|Cooper 1968, pp.&nbsp;28–29]]</ref> and a glass former.<ref name=Rao22/> The dioxide is soluble in acidic solutions (the monoxide GeO, is even more so), and this is sometimes used to classify germanium as a metal.<ref>[[#Steele1966|Steele 1966, pp.&nbsp;178, 188–89]]</ref> Up to the 1930s germanium was considered to be a poorly conducting metal;<ref>[[#Haller 2006|Haller 2006, p.&nbsp;3]]</ref> it has occasionally been classified as a metal by later writers.<ref>[[#Walker|See, for example, Walker & Tarn 1990, p.&nbsp;590]]</ref> As with all the elements commonly recognised as metalloids, germanium has an established organometallic chemistry (see [[Organogermanium chemistry]]).<ref>[[#Wiberg2001|Wiberg 2001, p.&nbsp;742]]</ref>

===Arsenic===
{{Main|Arsenic}}
[[File:Arsen 1a.jpg|thumb|left|[[Arsenic]], sealed in a container to prevent [[tarnishing]]|alt=Two dull silver clusters of crystalline shards.]]
Arsenic is a grey, metallic looking solid. It has a density of 5.727 g/cm<sup>3</sup> and is brittle, and moderately hard (more than aluminium; less than [[iron]]).<ref name="GWM2011">[[#Gray2011|Gray, Whitby & Mann 2011]]</ref> It is stable in dry air but develops a golden bronze patina in moist air, which blackens on further exposure. Arsenic is attacked by nitric acid and concentrated sulfuric acid. It reacts with fused caustic soda to give the arsenate Na<sub>3</sub>AsO<sub>3</sub> and hydrogen gas.<ref name="Greenwood 2002, p. 552">[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;552]]</ref> Arsenic [[sublimation (phase transition)|sublimes]] at 615&nbsp;°C. The vapour is lemon-yellow and smells like garlic.<ref>[[#Parkes1943|Parkes & Mellor 1943, p.&nbsp;740]]</ref> Arsenic only melts under a pressure of 38.6 [[Atmosphere (unit)|atm]], at 817&nbsp;°C.<ref>[[#Russell2005|Russell & Lee 2005, p.&nbsp;420]]</ref> It is a semimetal with an electrical conductivity of around 3.9 × 10<sup>4</sup>&nbsp;S•cm<sup>−1</sup><ref name="Carapella1968p30">[[#Carapella1968|Carapella 1968, p.&nbsp;30]]</ref> and a band overlap of 0.5&nbsp;eV.<ref name="Barfuß 1981, p. 967">[[#Barfuß1981|Barfuß et al. 1981, p.&nbsp;967]]</ref>{{refn|1=Arsenic also exists as a naturally occurring (but rare) allotrope ''(arsenolamprite),'' a crystalline semiconductor with a band gap of around 0.3&nbsp;eV or 0.4&nbsp;eV. It can also be prepared in a semiconducting [[amorphous solid|amorphous]] form, with a band gap of around 1.2–1.4&nbsp;eV.<ref>[[#Greaves1974|Greaves, Knights & Davis 1974, p.&nbsp;369]]; [[#Madelung2004|Madelung 2004, pp.&nbsp;405, 410]]</ref>|group=n}} Liquid arsenic is a semiconductor with a band gap of 0.15&nbsp;eV.<ref>[[#Bailar1973|Bailar & Trotman-Dickenson 1973, p.&nbsp;558]]; [[#Li1990|Li 1990]]</ref>

The chemistry of arsenic is predominately nonmetallic.<ref>[[#Bailar1965|Bailar, Moeller & Kleinberg 1965, p.&nbsp;477]]</ref> Whether or not arsenic forms a cation is unclear.{{refn|1=Sources mentioning cationic arsenic include: Gillespie & Robinson<ref>[[#Gillespie|Gillespie & Robinson 1963, p.&nbsp;450]]</ref> who find that, "in very dilute solutions in 100% sulphuric acid, arsenic (III) oxide forms arsonyl (III) hydrogen sulphate, AsO.HO<sub>4</sub>, which is partly ionized to give the AsO<sup>+</sup> cation. Both these species probably exist mainly in solvated forms, e.g., As(OH)(SO<sub>4</sub>H)<sub>2</sub>, and As(OH)(SO<sub>4</sub>H)<sup>+</sup> respectively"; Paul et al.<ref>[[#Pauletal|Paul et al. 1971]]; see also [[#Ahmeda|Ahmeda & Rucka 2011, pp.&nbsp;2893–94]]</ref> who reported spectroscopic evidence for the presence of As<sub>4</sub><sup>2+</sup> and As<sub>2</sub><sup>2+</sup> cations when arsenic was oxidized with [[peroxydisulfuryl difluoride]] S<sub>2</sub>O<sub>6</sub>F<sub>2</sub> in highly acidic media (Gillespie and Passmore<ref>[[#GillespieP|Gillespie & Passmore 1972, p.&nbsp;478]]</ref> noted the spectra of these species were very similar to S<sub>4</sub><sup>2+</sup> and S<sub>8</sub><sup>2+</sup> and concluded that, "at present" there was no reliable evidence for any homopolycations of arsenic); Van Muylder and Pourbaix,<ref>[[#Van Muylder|Van Muylder & Pourbaix 1974, p.&nbsp;521]]</ref> who write that, "As<sub>2</sub>O<sub>3</sub> is an amphoteric oxide which dissolves in water and in solutions of pH between 1 and 8 with the formation of undissociated [[arsenious acid]] HAsO<sub>2</sub>; the solubility…increases at pH's below 1 with the formation of 'arsenyl' ions AsO<sup>+</sup>…"; Kolthoff and Elving<ref>[[#Kolthoff|Kolthoff & Elving 1978, p.&nbsp;210]]</ref> who write that, "the As<sup>3+</sup> cation exists to some extent only in strongly acid solutions; under less acid conditions the tendency is toward [[hydrolysis]], so that the anionic form predominates"; Moody<ref>[[#Moody|Moody 1991, pp.&nbsp;248–49]]</ref> who observes that, "arsenic trioxide, As<sub>4</sub>O<sub>6</sub>, and arsenious acid, H<sub>3</sub>AsO<sub>3</sub>, are apparently amphoteric but no cations, As<sup>3+</sup>, As(OH)<sup>2+</sup> or As(OH)<sub>2</sub><sup>+</sup> are known"; and Cotton et al.<ref>[[#Cotton1999|Cotton & Wilkinson 1999, pp.&nbsp;396, 419]]</ref> who write that (in aqueous solution) the simple arsenic cation As<sup>3+</sup> "may occur to some slight extent [along with the AsO<sup>+</sup> cation]" and that, "Raman spectra show that in acid solutions of As<sub>4</sub>O<sub>6</sub> the only detectable species is the pyramidal As(OH)<sub>3</sub>".|group=n}} Its many metal alloys are mostly brittle.<ref>[[#Eagleson1994|Eagleson 1994, p.&nbsp;91]]</ref> It shows fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox/> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225/> Arsenic generally forms compounds in which it has an oxidation state of +3 or +5.<ref name="Massey267">[[#Massey2000|Massey 2000, p.&nbsp;267]]</ref> The halides, and the oxides and their derivatives are illustrative examples.<ref name="Bailar513">[[#Bailar1965|Bailar, Moeller & Kleinberg 1965, p.&nbsp;513]]</ref> In the trivalent state, arsenic shows some incipient metallic properties.<ref>[[#Timm1944|Timm 1944, p.&nbsp;454]]</ref> The halides are [[hydrolysed]] by water but these reactions, particularly those of the chloride, are reversible with the addition of a [[hydrohalic acid]].<ref>[[#Partington1944|Partington 1944, p.&nbsp;641]]; [[#Kleinberg1960|Kleinberg, Argersinger & Griswold 1960, p.&nbsp;419]]</ref> The oxide is acidic but, as noted below, (weakly) amphoteric. The higher, less stable, pentavalent state has strongly acidic (nonmetallic) properties.<ref>[[#Morgan1906|Morgan 1906, p.&nbsp;163]]; [[#Moeller1954|Moeller 1954, p.&nbsp;559]]</ref> Compared to phosphorus, the stronger metallic character of arsenic is indicated by the formation of oxoacid salts such as AsPO<sub>4</sub>, As<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>{{refn|1=The formulae of AsPO<sub>4</sub> and As<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> suggest straightforward ionic formulations, with As<sup>3+</sup>, but this is not the case. AsPO<sub>4</sub>, "which is virtually a covalent oxide", has been referred to as a double oxide, of the form As<sub>2</sub>O<sub>3</sub>·P<sub>2</sub>O<sub>5</sub>. It consists of AsO<sub>3</sub> pyramids and PO<sub>4</sub> tetrahedra, joined together by all their corner atoms to form a continuous polymeric network.<ref>[[#Corbridge|Corbridge 2013, pp.&nbsp;122, 215]]</ref> As<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> has a structure in which each SO<sub>4</sub> tetrahedron is bridged by two AsO<sub>3</sub> trigonal pyramida.<ref>[[#Douglade|Douglade 1982]]</ref>|group=n}} and arsenic acetate As(CH<sub>3</sub>COO)<sub>3</sub>.<ref>[[#Zingaro1994|Zingaro 1994, p.&nbsp;197]]; [[#Emeléus1959|Emeléus & Sharpe 1959, p.&nbsp;418]]; [[#Addison1972|Addison & Sowerby 1972, p.&nbsp;209]]; [[#Mellor1964|Mellor 1964, p.&nbsp;337]]</ref> The oxide As<sub>2</sub>O<sub>3</sub> is polymeric,<ref name=Pudd59/> amphoteric,<ref>[[#Pourbaix1974|Pourbaix 1974, p.&nbsp;521]]; [[#Eagleson1994|Eagleson 1994, p.&nbsp;92]]; [[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;572]]</ref>{{refn|1=As<sub>2</sub>O<sub>3</sub> is usually regarded as being amphoteric but a few sources say it is (weakly)<ref>[[#Wiberg2001|Wiberg 2001, pp.&nbsp;750, 975]]; [[#Silberberg2006|Silberberg 2006, p.&nbsp;314]]</ref> acidic. They describe its "basic" properties (its reaction with concentrated [[hydrochloric acid]] to form arsenic trichloride) as being alcoholic, in analogy with the formation of covalent alkyl chlorides by covalent alcohols (e.g., R-OH + HCl <big>→</big> RCl + H<sub>2</sub>O)<ref>[[#Sidgwick1950|Sidgwick 1950, p.&nbsp;784]]; [[#Moody1991|Moody 1991, pp.&nbsp;248–9, 319]]</ref>|group=n}} and a glass former.<ref name=Rao22/> Arsenic has an extensive organometallic chemistry (see [[Organoarsenic chemistry]]).<ref>[[#Krannich2006|Krannich & Watkins 2006]]</ref>

===Antimony===
{{Main|Antimony}}
[[File:Antimony-4.jpg|thumb|right|[[Antimony]], showing its brilliant [[lustre (mineralogy)|lustre]]|alt=A glistening silver rock-like chunk, with a blue tint, and roughly parallel furrows.]]
Antimony is a silver-white solid with a blue tint and a brilliant lustre.<ref name="Greenwood 2002, p. 552"/> It has a density of 6.697 g/cm<sup>3</sup> and is brittle, and moderately hard (more so than arsenic; less so than iron; about the same as copper).<ref name="GWM2011"/> It is stable in air and moisture at room temperature. It is attacked by concentrated nitric acid, yielding the hydrated pentoxide Sb<sub>2</sub>O<sub>5</sub>. [[Aqua regia]] gives the pentachloride SbCl<sub>5</sub> and hot concentrated sulfuric acid results in the [[antimony sulfate|sulfate]] Sb<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>.<ref name="Greenwood 2002, p. 553">[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;553]]</ref> It is not affected by molten alkali.<ref>[[#Dunstan1968|Dunstan 1968, p.&nbsp;433]]</ref> Antimony is capable of displacing hydrogen from water, when heated:  2 Sb + 3 H<sub>2</sub>O → Sb<sub>2</sub>O<sub>3</sub> + 3 H<sub>2</sub>.<ref>[[#Parise1996|Parise 1996, p.&nbsp;112]]</ref> It melts at 631&nbsp;°C. Antimony is a semimetal with an electrical conductivity of around 3.1 × 10<sup>4</sup>&nbsp;S•cm<sup>−1</sup><ref>[[#Carapella1968a|Carapella 1968a, p.&nbsp;23]]</ref> and a band overlap of 0.16&nbsp;eV.<ref name="Barfuß 1981, p. 967"/>{{refn|1=Antimony can also be prepared in an [[amorphous solid|amorphous]] semiconducting black form, with an estimated (temperature-dependent) band gap of 0.06–0.18&nbsp;eV.<ref>[[#Moss1952|Moss 1952, pp.&nbsp;174, 179]]</ref>|group=n}} Liquid antimony is a metallic conductor with an electrical conductivity of around 5.3 × 10<sup>4</sup>&nbsp;S•cm<sup>−1</sup>.<ref>[[#Dupree1982|Dupree, Kirby & Freyland 1982, p.&nbsp;604]]; [[#Mhiaoui2003|Mhiaoui, Sar, & Gasser 2003]]</ref>

Most of the chemistry of antimony is characteristic of a nonmetal.<ref>[[#Kotz2009|Kotz, Treichel & Weaver 2009, p.&nbsp;62]]</ref> Antimony has some definite cationic chemistry,<ref>[[#Cotton1999|Cotton et al. 1999, p.&nbsp;396]]</ref> SbO<sup>+</sup> and Sb(OH)<sub>2</sub><sup>+</sup> being present in acidic aqueous solution;<ref>[[#King1994|King 1994, p.&nbsp;174]]</ref>{{refn|1=Lidin<ref>[[#Lidin|Lidin 1996, p.&nbsp;372]]</ref> asserts that SbO<sup>+</sup> does not exist and that the stable form of Sb(III) in aqueous solution is an incomplete hydrocomplex [Sb(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]<sup>+</sup>.|group=n}} the compound Sb<sub>8</sub>(GaCl<sub>4</sub>)<sub>2</sub>, which contains the homopolycation, Sb<sub>8</sub><sup>2+</sup>, was prepared in 2004.<ref>[[#Lindsjö|Lindsjö, Fischer & Kloo 2004]]</ref> It can form alloys with one or more metals such as aluminium,<ref>[[#Friend1953|Friend 1953, p.&nbsp;87]]</ref> iron, [[nickel]], copper, zinc, tin, lead, and bismuth.<ref>[[#Fesquet1872|Fesquet 1872, pp.&nbsp;109–14]]</ref> Antimony has fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox/> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225/> Like arsenic, antimony generally forms compounds in which it has an oxidation state of +3 or +5.<ref name=Massey267/> The halides, and the oxides and their derivatives are illustrative examples.<ref name=Bailar513/> The +5 state is less stable than the +3, but relatively easier to attain than with arsenic. This is explained by the poor shielding afforded the arsenic nucleus by its [[d electron count|3d<sup>10</sup> electrons]]. In comparison, the tendency of antimony (being a heavier atom) to [[redox|oxidize]] more easily partially offsets the effect of its 4d<sup>10</sup> shell.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;553]]; [[#Massey2000|Massey 2000, p.&nbsp;269]]</ref> Tripositive antimony is amphoteric; [[penta-|pentapositive]] antimony is (predominately) acidic.<ref>[[#King1994|King 1994, p.&nbsp;171]]</ref> Consistent with an increase in metallic character down [[pnictogen|group 15]], antimony forms salts including an [[acetate]] Sb(CH<sub>3</sub>CO<sub>2</sub>)<sub>3</sub>, [[phosphate]] SbPO<sub>4</sub>, sulfate Sb<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub> and [[perchlorate]] Sb(ClO<sub>4</sub>)<sub>3</sub>.<ref>[[#Turova2011|Turova 2011, p.&nbsp;46]]</ref> The otherwise acidic pentoxide Sb<sub>2</sub>O<sub>5</sub> shows some basic (metallic) behaviour in that it can be dissolved in very acidic solutions, with the formation of the [[oxycation]] SbO{{su|b=2|p=+}}.<ref>[[#Pourbaix1974|Pourbaix 1974, p.&nbsp;530]]</ref> The oxide Sb<sub>2</sub>O<sub>3</sub> is polymeric,<ref name=Pudd59/> amphoteric,<ref name="Wiberg2001p764">[[#Wiberg2001|Wiberg 2001, p.&nbsp;764]]</ref> and a glass former.<ref name=Rao22/> Antimony has an extensive organometallic chemistry (see [[Organoantimony chemistry]]).<ref>[[#House2008|House 2008, p.&nbsp;497]]</ref>

===Tellurium===
{{Main|Tellurium}}
[[File:Tellurium2.jpg|thumb|left|[[Tellurium]], described by [[Dmitri Mendeleev]] as forming a transition between [[metals]] and [[nonmetals]]<ref>[[#Mendeléeff1897a|Mendeléeff 1897, p.&nbsp;274]]</ref>|alt=A shiny silver-white medallion with a striated surface, irregular around the outside, with a square spiral-like pattern in the middle.]]
Tellurium is a silvery-white shiny solid.<ref>[[#Emsley2001|Emsley 2001, p.&nbsp;428]]</ref> It has a density of 6.24 g/cm<sup>3</sup>, is brittle, and is the softest of the commonly recognised metalloids, being marginally harder than sulfur.<ref name="GWM2011"/> Large pieces of tellurium are stable in air. The finely powdered form is oxidized by air in the presence of moisture. Tellurium reacts with boiling water, or when freshly precipitated even at 50&nbsp;°C, to give the dioxide and hydrogen: Te + 2 H<sub>2</sub>O → TeO<sub>2</sub> + 2 H<sub>2</sub>.<ref name=Kudryavtsev78>[[#Kudryavtsev1974|Kudryavtsev 1974, p.&nbsp;78]]</ref> It reacts (to varying degrees) with nitric, sulfuric, and hydrochloric acids to give compounds such as the [[sulfoxide]] TeSO<sub>3</sub> or [[tellurous acid]] H<sub>2</sub>TeO<sub>3</sub>,<ref>[[#Bagnall1966|Bagnall 1966, pp.&nbsp;32–33, 59, 137]]</ref> the basic nitrate (Te<sub>2</sub>O<sub>4</sub>H)<sup>+</sup>(NO<sub>3</sub>)<sup>−</sup>,<ref>[[#Swink1966|Swink et al. 1966]]; [[#Anderson1980|Anderson et al. 1980]]</ref> or the oxide sulfate Te<sub>2</sub>O<sub>3</sub>(SO<sub>4</sub>).<ref>[[#Ahmed2000|Ahmed, Fjellvåg & Kjekshus 2000]]</ref> It dissolves in boiling alkalis, to give the [[tellurite]] and [[telluride (chemistry)|telluride]]: 3 Te + 6 KOH = K<sub>2</sub>TeO<sub>3</sub> + 2 K<sub>2</sub>Te + 3 H<sub>2</sub>O, a reaction that proceeds or is reversible with increasing or decreasing temperature.<ref>[[#Chizhikov1970|Chizhikov & Shchastlivyi 1970, p.&nbsp;28]]</ref>

At higher temperatures tellurium is sufficiently plastic to extrude.<ref>[[#Kudryavtsev1974|Kudryavtsev 1974, p.&nbsp;77]]</ref> It melts at 449.51&nbsp;°C. Crystalline tellurium has a structure consisting of parallel infinite spiral chains. The bonding between adjacent atoms in a chain is covalent, but there is evidence of a weak metallic interaction between the neighbouring atoms of different chains.<ref name="Stuke1074p178">[[#Stuke1974|Stuke 1974, p.&nbsp;178]]; [[#Donohue1982|Donohue 1982, pp.&nbsp;386–87]]; [[#Cotton1999|Cotton et al. 1999, p.&nbsp;501]]</ref> Tellurium is a semiconductor with an electrical conductivity of around 1.0&nbsp;S•cm<sup>−1</sup><ref>[[#Becker1971|Becker, Johnson & Nussbaum 1971, p.&nbsp;56]]</ref> and a band gap of 0.32 to 0.38&nbsp;eV.<ref name=Berger90>[[#Berger1997|Berger 1997, p.&nbsp;90]]</ref> Liquid tellurium is a semiconductor, with an electrical conductivity, on melting, of around 1.9 × 10<sup>3</sup>&nbsp;S•cm<sup>−1</sup>.<ref name=Berger90/> [[Superheated]] liquid tellurium is a metallic conductor.<ref>[[#Chizhikov1970|Chizhikov & Shchastlivyi 1970, p.&nbsp;16]]</ref>

Most of the chemistry of tellurium is characteristic of a nonmetal.<ref>[[#Jolly1966|Jolly 1966, pp.&nbsp;66–67]]</ref>
It shows some cationic behaviour. The dioxide dissolves in acid to yield the trihydroxotellurium(IV) Te(OH)<sub>3</sub><sup>+</sup> ion;<ref>[[#Schwietzer2010|Schwietzer & Pesterfield 2010, p.&nbsp;239]]</ref>{{refn|1=Cotton et al.<ref>[[#Cotton1999|Cotton et al. 1999, p.&nbsp;498]]</ref> note that TeO<sub>2</sub> appears to have an ionic lattice; Wells<ref>[[#Wells1984|Wells 1984, p.&nbsp;715]]</ref> suggests that the Te–O bonds have "considerable covalent character".|group=n}} the red Te<sub>4</sub><sup>2+</sup> and yellow-orange Te<sub>6</sub><sup>2+</sup> ions form when tellurium is oxidized in [[fluorosulfuric acid]] (HSO<sub>3</sub>F), or liquid [[sulfur dioxide]] (SO<sub>2</sub>), respectively.<ref>[[#Wiberg2001|Wiberg 2001, p.&nbsp;588]]</ref> It can form alloys with aluminium, [[silver]], and tin.<ref>[[#Mellor1964a|Mellor 1964a, p.&nbsp; 30]]; [[#Wiberg2001|Wiberg 2001, p.&nbsp;589]]</ref> Tellurium shows fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox/> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225/> Tellurium generally forms compounds in which it has an oxidation state of −2, +4 or +6. The +4 state is the most stable.<ref name=Kudryavtsev78/> Tellurides of composition X<sub>''x''</sub>Te<sub>''y''</sub> are easily formed with most other elements and represent the most common tellurium minerals. [[Non-stoichiometric compound|Nonstoichiometry]] is pervasive, especially with transition metals. Many tellurides can be regarded as metallic alloys.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, pp.&nbsp;765–66]]</ref> The increase in metallic character evident in tellurium, as compared to the lighter [[chalcogen]]s, is further reflected in the reported formation of various other oxyacid salts, such as a [[basic salt|basic]] selenate 2TeO<sub>2</sub>·SeO<sub>3</sub>  and an analogous perchlorate and [[periodate]] 2TeO<sub>2</sub>·HXO<sub>4</sub>.<ref>[[#Bagnall1966|Bagnall 1966, pp.&nbsp;134–51]]; [[#Greenwood2002|Greenwood & Earnshaw 2002, p.&nbsp;786]]</ref> Tellurium forms a polymeric,<ref name=Pudd59/> amphoteric,<ref name="Wiberg2001p764"/> glass-forming oxide<ref name=Rao22/> TeO<sub>2</sub>. It is a "conditional" glass-forming oxide – it forms a glass with a very small amount of additive.<ref name=Rao22/> Tellurium has an extensive organometallic chemistry (see [[Organotellurium chemistry]]).<ref>[[#Detty1994|Detty & O'Regan 1994, pp.&nbsp;1–2]]</ref>