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'''Metaloid''' je [[chemical element]] which has a preponderance of [[material property|properties]] in between, or that are a mixture of, those of [[metal]]s and [[Nonmetal (chemistry)|nonmetal]]s. The word metalloid comes from the [[Latin language|Latin]] ''metallum'' ("metal") and the [[Greek language|Greek]] ''oeides'' ("resembling in form or appearance").<ref>[[Metalloid#OED1989|''Oxford English Dictionary'' 1989, 'metalloid']]; [[Metalloid#GGH2003|Gordh, Gordh & Headrick 2003, p. 753]]</ref> There is no standard definition of a metalloid and no complete agreement on which elements are metalloids. Despite the lack of specificity, the term remains in use in the literature.
The six commonly recognised metalloids are [[boron]], [[silicon]], [[germanium]], [[arsenic]], [[antimony]] and [[tellurium]]. Five elements are less frequently so classified: [[carbon]], [[aluminium]], [[selenium]], [[polonium]] and [[astatine]]. On a standard periodic table, all eleven elements are in a diagonal region of the [[p-block]] extending from boron at the upper left to astatine at lower right. Some periodic tables include a [[dividing line between metals and nonmetals]], and the metalloids may be found close to this line.
Typical metalloids have a metallic appearance, may be brittle and are only fair [[electrical conductor|conductors of electricity]]. They can form [[alloy]]s with [[metal]]s, and many of their other [[physical property|physical properties]] and [[chemical property|chemical properties]] are intermediate between those of metallic and nonmetallic elements. They and their compounds are used in alloys, biological agents, [[catalyst]]s, [[flame retardant]]s, [[glass]]es, [[optical storage]] and [[optoelectronics]], [[pyrotechnics]], [[semiconductor]]s, and electronics.
The term ''metalloid'' originally referred to [[nonmetal]]s. Its more recent meaning, as a category of elements with intermediate or hybrid properties, became widespread in 1940–1960. Metalloids are sometimes called semimetals, a practice that has been discouraged,<ref name=Atkins2010p20/> as the term ''[[semimetal]]'' has a more common usage as a specific kind of [[electronic band structure]] of a substance. In this context, only [[arsenic]] and [[antimony]] are semimetals, and commonly recognised as metalloids.
==Definitions==
{{See also|Lists of metalloids}}
===Judgment-based===
A metalloid is an element that possesses a preponderance of properties in between, or that are a mixture of, those of [[metal]]s and [[nonmetal]]s, and which is therefore hard to classify as either a [[metal]] or a [[nonmetal]]. This is a generic definition that draws on metalloid attributes consistently cited in the literature.{{refn|1=Definitions and extracts by different authors, illustrating aspects of the generic definition, follow:
*"In chemistry a metalloid is an element with properties intermediate between those of [[metal]]s and [[nonmetal]]s."<ref>[[#Cusack1987|Cusack 1987, p. 360]]</ref>
*"Between the [[metal]]s and [[nonmetal]]s in the periodic table we find elements ... [that] share some of the characteristic properties of both the [[metal]]s and [[Nonmetal|nonmetals]], making it difficult to place them in either of these two main categories"<ref>[[#Kelter2009|Kelter, Mosher & Scott 2009, p. 268]]</ref>
*"Chemists sometimes use the name metalloid ... for these elements which are difficult to classify one way or the other."<ref name="Hill 2000, p. 41">[[#Hill2000|Hill & Holman 2000, p. 41]]</ref>
*"Because the traits distinguishing [[metal]]s and [[nonmetal]]s are qualitative in nature, some elements do not fall unambiguously in either category. These elements ... are called metalloids ..."<ref>[[#King1979|King 1979, p. 13]]</ref>
More broadly, metalloids have been referred to as:
*"elements that ... are somewhat of a cross between [[metal]]s and [[nonmetal]]s";<ref>[[#Moore2011|Moore 2011, p. 81]]</ref> or
*"weird in-between elements".<ref>[[#Gray2010|Gray 2010]]</ref>|group=n}} Difficulty of categorisation is a key attribute. Most elements have a mixture of metallic and nonmetallic properties,<ref name=Hopkins>[[#Hopkins1956|Hopkins & Bailar 1956, p. 458]]</ref> and can be classified according to which set of properties is more pronounced.<ref>[[#Glinka1965|Glinka 1965, p. 77]]</ref>{{refn|1=[[Gold]], for example, has mixed properties but is still recognised as "king of [[metal]]s". Besides metallic behaviour (such as high electrical conductivity, and [[cation]] formation), gold shows nonmetallic behaviour:
*It has the [[table of standard electrode potentials|highest]] [http://www.chemguide.co.uk/physical/redoxeqia/introduction.html electrode potential]
*It has the third-highest [[ionization energy]] among the [[metal]]s (after [[zinc]] and [[mercury (element)|mercury]])
*It has the highest [[electron affinity]]
*Its [[electronegativity]] of 2.54 is highest among the [[metal]]s and exceeds that of some [[nonmetal]]s ([[hydrogen]] 2.2; [[phosphorus]] 2.19; and [[radon]] 2.2)
*It forms the Au<sup>−</sup> auride [[anion]], acting in this way like a [[halogen]]
*It sometimes has a tendency, known as "[[aurophilicity]]", to bond to itself.<ref>[[#Wiberg2001|Wiberg 2001, p. 1279]]</ref>
On halogen character, see also Belpassi et al.,<ref>[[#Belpassi2006|Belpassi et al. 2006, pp. 4543–44]]</ref> who conclude that in the aurides MAu (M = [[alkali metal|Li–Cs]]) gold "behaves as a halogen, intermediate between [[bromine|Br]] and [[iodine|I]]"; on aurophilicity, see also Schmidbaur and Schier.<ref>[[#Schmidbaur2008|Schmidbaur & Schier 2008, pp. 1931–51]]</ref>|group=n}} Only the elements at or near the margins, lacking a sufficiently clear preponderance of either metallic or nonmetallic properties, are classified as metalloids.<ref>[[#TM1987|Tyler Miller 1987, p. 59]]</ref>
[[Boron]], [[silicon]], [[germanium]], [[arsenic]], [[antimony]], and [[tellurium]] are commonly recognised as metalloids.<ref>[[#Goldsmith1982|Goldsmith 1982, p. 526]]; [[#Kotz2009|Kotz, Treichel & Weaver 2009, p. 62]]; [[#Bettelheim|Bettelheim et al. 2010, p. 46]]</ref>{{refn|1=Mann et al.<ref name=Mann/> refer to these elements as "the recognized metalloids".|group=n}} Depending on the author, one or more from [[selenium]], [[polonium]], or [[astatine]] are sometimes added to the list.<ref>[[#Hawkes2001|Hawkes 2001, p. 1686]]; [[#Segal1989|Segal 1989, p. 965]]; [[#McMurray2009|McMurray & Fay 2009, p. 767]]</ref> [[Boron]] sometimes is excluded, by itself, or with [[silicon]].<ref>[[#Bucat1983|Bucat 1983, p. 26]]; [[#Brown2007|Brown c. 2007]]</ref> Sometimes [[tellurium]] is not regarded as a metalloid.<ref name="Swift1962,100">[[#Swift1962|Swift & Schaefer 1962, p. 100]]</ref> The inclusion of [[antimony]], [[polonium]], and [[astatine]] as metalloids has been questioned.<ref>[[#Hawkes2001|Hawkes 2001, p. 1686]]; [[#Hawkes2010|Hawkes 2010]]; [[#Holt2007|Holt, Rinehart & Wilson c. 2007]]</ref>
Other elements are occasionally classified as metalloids. These elements include<ref>[[#Dunstan1968|Dunstan 1968, pp. 310, 409]]. Dunstan lists Be, Al, Ge (maybe), As, Se (maybe), Sn, Sb, Te, Pb, Bi, and Po as metalloids (pp. 310, 323, 409, 419).</ref> hydrogen,<ref>[[#Tilden1876|Tilden 1876, pp. 172, 198–201]]; [[#Smith1994|Smith 1994, p. 252]]; [[#Bodner1993|Bodner & Pardue 1993, p. 354]]</ref> [[beryllium]],<ref>[[#Bassett1966|Bassett et al. 1966, p. 127]]</ref> [[nitrogen]],<ref name=rausch>[[#Rausch1960|Rausch 1960]]</ref> [[phosphorus]],<ref>[[#Thayer1977|Thayer 1977, p. 604]]; [[#Warren1981|Warren & Geballe 1981]]; [[#M&E|Masters & Ela 2008, p. 190]]</ref> [[sulfur]],<ref>[[#Warren1981|Warren & Geballe 1981]]; [[#Chalmers1959|Chalmers 1959, p. 72]]; [[#United1965|US Bureau of Naval Personnel 1965, p. 26]]</ref> [[zinc]],<ref>[[#Siebring1967|Siebring 1967, p. 513]]</ref> [[gallium]],<ref>[[#Wiberg2001|Wiberg 2001, p. 282]]</ref> [[tin]], [[iodine]],<ref>[[#Rausch1960|Rausch 1960]]; [[#Friend1953|Friend 1953, p. 68]]</ref> [[lead]],<ref>[[#Murray1928|Murray 1928, p. 1295]]</ref> [[bismuth]],<ref name="Swift1962,100"/> and radon.<ref>[[#Hampel&H1966|Hampel & Hawley 1966, p. 950]]; [[#Stein1985|Stein 1985]]; [[#Stein1987|Stein 1987, pp. 240, 247–48]]</ref> The term metalloid has also been used for elements that exhibit metallic [[Lustre (mineralogy)|lustre]] and [[Electrical resistivity and conductivity|electrical conductivity]], and that are [[amphoterism|amphoteric]], such as [[arsenic]], [[antimony]], [[vanadium]], [[chromium]], [[molybdenum]], [[tungsten]], [[tin]], [[lead]], and [[aluminium]].<ref>[[#Hatcher1949|Hatcher 1949, p. 223]]; [[#Secrist|Secrist & Powers 1966, p. 459]]</ref> The [[Post-transition metal#p-block metals|p-block metals]],<ref>[[#Taylor1960|Taylor 1960, p. 614]]</ref> and nonmetals (such as carbon or nitrogen) that can form [[alloy]]s with metals<ref>[[#Considine1984|Considine & Considine 1984, p. 568]]; [[#Cegielski1998|Cegielski 1998, p. 147]]; [[#TheAmerican2005|''The American heritage science dictionary 2005'', p. 397]]</ref> or modify their properties<ref>[[#Woodward1948|Woodward 1948, p. 1]]</ref> have also occasionally been considered as metalloids.
===Criteria-based===
{| class="wikitable floatright" style="width: 75px;"
|-
! Element
! IE<br/>(kcal/mol)
! IE<br/>(kJ/mol)
! EN
! nowrap|[[electronic band structure|Band structure]]
|-
| Boron
| style="text-align:center;"| 191
| style="text-align:center;"| 801
|style="padding-left:1em; padding-right:1em;"| 2.04
|style="padding-left:1em; padding-right:1em;"| [[semiconductor]]
|-
| Silicon
| style="text-align:center;"| 188
| style="text-align:center;"| 787
|style="padding-left:1em; padding-right:1em;"| 1.90
|style="padding-left:1em; padding-right:1em;"| semiconductor
|-
| Germanium
| style="text-align:center;"| 182
| style="text-align:center;"| 762
|style="padding-left:1em; padding-right:1em;"| 2.01
|style="padding-left:1em; padding-right:1em;"| semiconductor
|-
| Arsenic
| style="text-align:center;"| 226
| style="text-align:center;"| 944
|style="padding-left:1em; padding-right:1em;"| 2.18
|style="padding-left:1em; padding-right:1em;"| [[semimetal]]
|-
| Antimony
| style="text-align:center;"| 199
| style="text-align:center;"| 831
|style="padding-left:1em; padding-right:1em;"| 2.05
|style="padding-left:1em; padding-right:1em;"| semimetal
|-
| Tellurium
| style="text-align:center;"| 208
| style="text-align:center;"| 869
|style="padding-left:1em; padding-right:1em;"| 2.10
|style="padding-left:1em; padding-right:1em;"| semiconductor
|-
| style="text-align: right"| ''average''
| style="text-align:center;"| 199
| style="text-align:center;"| 832
|style="padding-left:1em; padding-right:1em;"| 2.05
|
|-
| colspan="5" style="text-align: left; font-size: 90%" |The elements commonly recognised as metalloids, and their [[ionization energy|ionization energies]] (IE);<ref>[[#NIST2010|NIST 2010]]. Values shown in the above table have been converted from the NIST values, which are given in eV.</ref> electronegativities (EN, revised Pauling scale); and electronic band structures<ref>[[#Berger1997|Berger 1997]]; [[#Lovett1977|Lovett 1977, p. 3]]</ref> (most thermodynamically stable forms under ambient conditions).
|}
No widely accepted definition of a metalloid exists, nor any division of the periodic table into [[Metal|metals]], metalloids, and [[Nonmetal|nonmetals]];<ref>[[#Goldsmith1982|Goldsmith 1982, p. 526]]; [[#Hawkes2001|Hawkes 2001, p. 1686]]</ref> Hawkes<ref name=H1687>[[#Hawkes2001|Hawkes 2001, p. 1687]]</ref> questioned the feasibility of establishing a specific definition, noting that anomalies can be found in several attempted constructs. Classifying an element as a metalloid has been described by Sharp<ref name="Sharp1981">[[#Sharp1981|Sharp 1981, p. 299]]</ref> as "arbitrary".
The number and identities of metalloids depend on what classification criteria are used. Emsley<ref>[[#Emsley1971|Emsley 1971, p. 1]]</ref> recognised four metalloids (germanium, arsenic, antimony, and tellurium); James et al.<ref>[[#James2000|James et al. 2000, p. 480]]</ref> listed twelve (Emsley's plus boron, carbon, silicon, selenium, bismuth, polonium, [[moscovium]], and [[livermorium]]). On average, seven elements are included in [[lists of metalloids|such lists]]; individual classification arrangements tend to share common ground and vary in the ill-defined<ref>[[#Chatt1951|Chatt 1951, p. 417]] "The boundary between metals and metalloids is indefinite ..."; [[#Burrows2009|Burrows et al. 2009, p. 1192]]: "Although the elements are conveniently described as metals, metalloids, and nonmetals, the transitions are not exact ..."</ref> margins.{{refn|1=Jones<ref>[[#Jones2010|Jones 2010, p. 170]]</ref> writes: "Though classification is an essential feature in all branches of science, there are always hard cases at the boundaries. Indeed, the boundary of a class is rarely sharp."|group=n}}{{refn|1=The lack of a standard division of the elements into metals, metalloids, and nonmetals is not necessarily an issue. There is more or less, a continuous progression from the metallic to the nonmetallic. A specified subset of this continuum could serve its particular purpose as well as any other.<ref>[[#Kneen1972|Kneen, Rogers & Simpson 1972, pp. 218–20]]</ref>|group=n}}
A single quantitative criterion such as [[electronegativity]] is commonly used,<ref>[[#Rochow1966|Rochow 1966, pp. 1, 4–7]]</ref> metalloids having electronegativity values from 1.8 or 1.9 to 2.2.<ref>[[#Rochow1977|Rochow 1977, p. 76]]; [[#Mann2000|Mann et al. 2000, p. 2783]]</ref> Further examples include [[atomic packing factor|packing efficiency]] (the fraction of volume in a [[crystal structure]] occupied by atoms) and the Goldhammer–Herzfeld criterion ratio.<ref>[[#Askeland|Askeland, Phulé & Wright 2011, p. 69]]</ref> The commonly recognised metalloids have packing efficiencies of between 34% and 41%.{{refn|1=The packing efficiency of boron is 38%; silicon and germanium 34; arsenic 38.5; antimony 41; and tellurium 36.4.<ref>[[#VanSetten2007|Van Setten et al. 2007, pp. 2460–61]]; [[#Russell2005|Russell & Lee 2005, p. 7]] (Si, Ge); [[#Pearson1972|Pearson 1972, p. 264]] (As, Sb, Te; also black P)</ref> These values are lower than in most metals (80% of which have a packing efficiency of at least 68%),<ref>[[#Russell2005|Russell & Lee 2005, p. 1]]</ref> but higher than those of elements usually classified as nonmetals. (Gallium is unusual, for a metal, in having a packing efficiency of just 39%.)<ref>[[#Russell2005|Russell & Lee 2005, pp. 6–7, 387]]</ref> Other notable values for metals are 42.9 for bismuth<ref name="ReferenceB">[[#Pearson1972|Pearson 1972, p. 264]]</ref> and 58.5 for liquid mercury.<ref>[[#Okajima1972|Okajima & Shomoji 1972, p. 258]]</ref>) Packing efficiencies for nonmetals are: graphite 17%,<ref>[[#Kitaĭgorodskiĭ1961|Kitaĭgorodskiĭ 1961, p. 108]]</ref> sulfur 19.2,<ref name="Neuburger">[[#Neuburger1936|Neuburger 1936]]</ref> iodine 23.9,<ref name="Neuburger"/> selenium 24.2,<ref name="Neuburger"/> and black phosphorus 28.5.<ref name="ReferenceB"/>|group=n}} The Goldhammer–Herzfeld ratio, roughly equal to the cube of the atomic radius divided by the [[molar volume]],<ref>[[#Edwards1983|Edwards & Sienko 1983, p. 693]]</ref>{{refn|1=More specifically, the <span id="Gold"></span>''Goldhammer–[[Karl Herzfeld|Herzfeld]] criterion'' is the ratio of the force holding an individual atom's [[valence electron]]s in place with the forces on the same electrons from interactions ''between'' the atoms in the solid or liquid element. When the interatomic forces are greater than, or equal to, the atomic force, valence electron itinerancy is indicated and metallic behaviour is predicted.<ref>[[#Herzfeld|Herzfeld 1927]]; [[#Edwards2000|Edwards 2000, pp. 100–03]]</ref> Otherwise nonmetallic behaviour is anticipated.|group=n}} is a simple measure of how metallic an element is, the recognised metalloids having ratios from around 0.85 to 1.1 and averaging 1.0.<ref>[[#Edwards1983|Edwards & Sienko 1983, p. 695]]; [[#Edwards2010|Edwards et al. 2010]]</ref>{{refn|1=As the ratio is based on classical arguments<ref>[[#Edwards1999|Edwards 1999, p. 416]]</ref> it does not accommodate the finding that polonium, which has a value of ~0.95, adopts a metallic (rather than [[covalent]]) [[crystalline structure]], on [[relativistic quantum chemistry|relativistic]] grounds.<ref>[[#Steurer2007|Steurer 2007, p. 142]]; [[#Pyykkö|Pyykkö 2012, p. 56]]</ref> Even so it offers a [[wikt:first-order|first order]] rationalization for the occurrence of metallic character amongst the elements.<ref name=edwards695>[[#Edwards1983|Edwards & Sienko 1983, p. 695]]</ref>|group=n}}
Other authors have relied on, for example, atomic conductance{{refn|1=Atomic conductance is the electrical conductivity of one mole of a substance. It is equal to electrical conductivity divided by molar volume.<ref name="Hill 2000, p. 41"/>|group=n}}<ref>[[#Hill2000|Hill & Holman 2000, p. 160]]. They characterise metalloids (in part) on the basis that they are "poor conductors of electricity with atomic conductance usually less than 10<sup>−3</sup> but greater than 10<sup>−5</sup> ohm<sup>−1</sup> cm<sup>−4</sup>".</ref> or [[coordination number#Usage in quasicrystal, liquid and other disordered systems|bulk coordination number]].<ref>[[#Bond2005|Bond 2005, p. 3]]: "One criterion for distinguishing semi-metals from true metals under normal conditions is that the [[coordination number#Usage in quasicrystal, liquid and other disordered systems|bulk coordination number]] of the former is never greater than eight, while for metals it is usually twelve (or more, if for the body-centred cubic structure one counts next-nearest neighbours as well)."</ref>
Jones, writing on the role of classification in science, observed that "[classes] are usually defined by more than two attributes".<ref>[[#Jones2010|Jones 2010, p. 169]]</ref> Masterton and Slowinski<ref>[[#Masterton1977|Masterton & Slowinski 1977, p. 160]] list B, Si, Ge, As, Sb, and Te as metalloids, and comment that Po and At are ordinarily classified as metalloids but add that this is arbitrary as so little is known about them.</ref> used three criteria to describe the six elements commonly recognised as metalloids: metalloids have [[ionization energy|ionization energies]] around 200 kcal/mol (837 kJ/mol) and electronegativity values close to 2.0. They also said that metalloids are typically semiconductors, though antimony and arsenic (semimetals from a physics perspective) have electrical conductivities approaching those of metals. Selenium and polonium are suspected as not in this scheme, while astatine's status is uncertain.{{refn|1=Selenium has an ionization energy (IE) of 225 kcal/mol (941 kJ/mol) and is sometimes described as a semiconductor. It has a relatively high 2.55 electronegativity (EN). Polonium has an IE of 194 kcal/mol (812 kJ/mol) and a 2.0 EN, but has a metallic band structure.<ref>[[#Kraig2004|Kraig, Roundy & Cohen 2004, p. 412]]; [[#Alloul2010|Alloul 2010, p. 83]]</ref> Astatine has an IE of 215 kJ/mol (899 kJ/mol) and an EN of 2.2.<ref>[[#Vernon|Vernon 2013, p. 1704]]</ref> Its electronic band structure is not known with any certainty.|group=n}}
In this context, Vernon proposed that a metalloid is a chemical element that, in its standard state, has (a) the electronic band structure of a semiconductor or a semimetal; and (b) an intermediate first ionization potential "(say 750−1,000 kJ/mol)"; and (c) an intermediate electronegativity (1.9–2.2).<ref>[[#Vernon|Vernon 2013, p. 1703]]</ref>
{{clear}}
==Periodic table territory==
{| cellspacing=0 cellpadding="1" style="float:right; text-align:center; border:1px solid grey; margin-left:1.2em; margin-bottom:1.2em; width:23.5em; {{{style|}}};"
|- <!-- title -->
! colspan=12 style="text-align:center; background:{{element color|table title}}" | Distribution and recognition status<br>of elements classified as metalloids
|- <!-- groups, header -->
|
| style="line-height:100%;" | [[Alkali metal|<small>1</small>]]
| [[Alkaline earth metal|<small>2</small>]]
|
| [[Group 12 element|<small>12</small>]]
| [[Boron group|<small>13</small>]]
| [[Carbon group|<small>14</small>]]
| [[Pnictogen|<small>15</small>]]
| [[Chalcogen|<small>16</small>]]
| [[Halogen|<small>17</small>]]
| [[Noble gas|<small>18</small>]]
|
|-
|
|-
|
| style="width:2.5em; background:#faebd7; border-right:1px solid black; border-top:1px solid black; border-left:1px solid black; border-bottom:2px solid grey" | <big>[[Hydrogen|H]]</big>
| style="width:2.5em; border-bottom:2px solid grey" |
| style="width:1em;" |
| style="width:2.5em" |
| style="width:2.5em; border-bottom:1px solid black" |
| style="width:2.5em; border-bottom:1px solid black" |
| style="width:2.5em; border-bottom:1px solid black" |
| style="width:2.5em; border-bottom:1px solid black" |
| style="width:2.5em; border-right:1px solid black; border-bottom:1px solid black"|
| style="width:2.5em; border-right:1px solid black; border-top:1px solid black; border-bottom:1px solid black" | <big>He</big>
|
|-
|
| style="border-right:1px solid black; border-top:2px solid grey; border-left:1px solid black; border-bottom:1px solid black" | <big>Li</big>
| style="background:#faebd7; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Beryllium|Be]]</big>
| style="border-left:2px solid grey" |
| style="border-right:2px solid grey" |
| style="background:#dcdcdc; border-right:1px solid black; border-left:2px solid grey; border-bottom:2px solid grey" | <big>[[Boron|B]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Carbon|C]]</big>
| style="background:#f0f8ff; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Nitrogen|N]]</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>O</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>F</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Ne</big>
|
|-
|
| style="border-right:1px solid black; border-left:1px solid black; border-bottom:1px solid black" | <big>Na</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Mg</big>
|
| style="border-right:1px solid black; border-bottom:1px solid black" |
| style="background:#dcdcdc; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Aluminium|Al]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-left:2px solid grey; border-bottom:2px solid grey" | <big>[[Silicon|Si]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Phosphorus|P]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Sulfur|S]]</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Cl</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Ar</big>
|
|-
|
| style="border-right:1px solid black; border-left:1px solid black; border-bottom:1px solid black" | <big>K</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Ca</big>
| style="border-right:1px solid black" |
| style="background:#f0f8ff; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Zinc|Zn]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Gallium|Ga]]</big>
| style="background:#dcdcdc; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Germanium|Ge]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-left:2px solid grey; border-bottom:2px solid grey" | <big>[[Arsenic|As]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Selenium|Se]]</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Br</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Kr</big>
|
|-
|
| style="border-right:1px solid black; border-left:1px solid black; border-bottom:1px solid black" | <big>Rb</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Sr</big>
| style="border-right:1px solid black" |
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Cd</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>In</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Tin|Sn]]</big>
| style="background:#dcdcdc; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Antimony|Sb]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-left:2px solid grey; border-bottom:2px solid grey" | <big>[[Tellurium|Te]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Iodine|I]]</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Xe</big>
|
|-
|
| style="border-right:1px solid black; border-left:1px solid black; border-bottom:1px solid black" | <big>Cs</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Ba</big>
| style="border-right:1px solid black"|
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Hg</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Tl</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Lead|Pb]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Bismuth|Bi]]</big>
| style="background:#dcdcdc; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Polonium|Po]]</big>
| style="background:#dcdcdc; border-right:1px solid black; border-left:2px solid grey; border-bottom:2px solid grey" | <big>[[Astatine|At]]</big>
| style="background:#f0f8ff; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Radon|Rn]]</big>
|
|-
|
| style="border-right:1px solid black; border-left:1px solid black; border-bottom:1px solid black" | <big>Fr</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Ra</big>
| style="border-right:1px solid black"|
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Cn</big>
| style="border-right:1px solid black; border-bottom:1px solid black" | <big>Nh</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Flerovium|Fl]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Moscovium|Mc]]</big>
| style="background:#faebd7; border-right:1px solid black; border-bottom:1px solid black" | <big>[[Livermorium|Lv]]</big>
| style="background:#faebd7; border-right:2px solid grey; border-top:2px solid grey; border-bottom:1px solid black" | <big>[[Tennessine|Ts]]</big>
| style="border-right:1px solid black; border-left:2px solid grey; border-bottom:1px solid black" | <big>Og</big>
|
|- <!-- PT caption -->
|
| colspan="10" style="text-align:center; font-size:20%" |
|
|- <!-- Legend --> style="text-align:left; font-size:90%"
|
| colspan=10 style="padding-left:0.5em; background:#f8f8f8;" |
{{legend |#dcdcdc |Commonly (93%) to rarely (9%) recognised as a<br>metalloid: B, C, Al, Si, Ge, As, Se, Sb, Te, Po, At}}
{{legend |#faebd7 |Very rarely (1–5%): H, Be, P, S, Ga, Sn, I, Pb, Bi, Fl, Mc, Lv, Ts}}
{{legend |#f0f8ff |Sporadically: N, Zn, Rn}}
<div class="legend"><span style="display:inline-block; width:1.5em; height:1.5em; margin:1px 0; border:0px; border-top:4px solid grey; border-right:4px solid grey; background-color:#e8e8e8;"> </span> [[Dividing line between metals and nonmetals|Metal–nonmetal dividing line]]: between {{nowrap|H and Li}}, {{nowrap|Be and B}}, {{nowrap|Al and Si}}, {{nowrap|Ge and As}}, {{nowrap|Sb and Te}}, {{nowrap|Po and At}}, and {{nowrap|Ts and Og}}</div><!-- code from {{legend}}, reduced -->
|
|- <!--Additional text --> style="text-align:left; padding-left:0.5em; padding-right:2px; font-size:90%"
| colspan="11" style="padding-left:0.5em; padding-right:2px"|
Periodic table extract showing groups 1–2 and 12–18, and a dividing line between metals and nonmetals. Percentages are median appearance frequencies in the [[list of metalloid lists]]. Sporadically recognised elements show that the metalloid net is sometimes cast very widely; although they do not appear in the list of metalloid lists, isolated references to their designation as metalloids can be found in the literature (as cited in this article).
|
|}
===Location===
Metalloids lie on either side of the [[dividing line between metals and nonmetals]]. This can be found, in varying configurations, on some [[periodic table]]s. Elements to the lower left of the line generally display increasing metallic behaviour; elements to the upper right display increasing nonmetallic behaviour.<ref name="Hamm 1969, p.653"/> When presented as a regular stairstep, elements with the highest [[critical point (thermodynamics)|critical temperature]] for their groups (Li, Be, Al, Ge, Sb, Po) lie just below the line.<ref>[[#Horvath1973|Horvath 1973, p. 336]]</ref>
The diagonal positioning of the metalloids represents an exception to the observation that elements with similar properties tend to occur in vertical [[group (periodic table)|groups]].<ref name="Gray91">[[#Gray2009|Gray 2009, p. 9]]</ref> A related effect can be seen in other [[diagonal relationship|diagonal similarities]] between some elements and their lower right neighbours, specifically lithium-magnesium, beryllium-aluminium, and boron-silicon. Rayner-Canham<ref name="Rayner2011">[[#Rayner2011|Rayner-Canham 2011]]</ref> has argued that these similarities extend to carbon-phosphorus, nitrogen-sulfur, and into three [[d-block]] series.
This exception arises due to competing horizontal and vertical trends in the [[nuclear charge]]. Going along a [[period (periodic table)|period]], the [[effective nuclear charge|nuclear charge]] increases with [[atomic number]] as do the number of electrons. The additional pull on outer electrons as nuclear charge increases generally outweighs the screening effect of having more electrons. With some irregularities, atoms therefore become smaller, ionization energy increases, and there is a gradual change in character, across a period, from strongly metallic, to weakly metallic, to weakly nonmetallic, to strongly nonmetallic elements.<ref>[[#Booth1972|Booth & Bloom 1972, p. 426]]; [[#Cox2004|Cox 2004, pp. 17, 18, 27–28]]; [[#Silberberg2006|Silberberg 2006, pp. 305–13]]</ref> Going down a [[main-group element|main group]], the effect of increasing nuclear charge is generally outweighed by the effect of additional electrons being further away from the nucleus. Atoms generally become larger, ionization energy falls, and metallic character increases.<ref>[[#Cox2004|Cox 2004, pp. 17–18, 27–28]]; [[#Silberberg2006|Silberberg 2006, pp. 305–13]]</ref> The net effect is that the location of the metal–nonmetal transition zone shifts to the right in going down a group,<ref name=Gray91/> and analogous diagonal similarities are seen elsewhere in the periodic table, as noted.<ref>[[#Rodgers|Rodgers 2011, pp. 232–33; 240–41]]</ref>
===Alternative treatments===
Elements bordering the metal–nonmetal dividing line are not always classified as metalloids, noting a binary classification can facilitate the establishment of rules for determining bond types between metals and nonmetals.<ref name=roher>[[#Roher2001|Roher 2001, pp. 4–6]]</ref> In such cases, the authors concerned focus on one or more attributes of interest to make their classification decisions, rather than being concerned about the marginal nature of the elements in question. Their considerations may or not be made explicit and may, at times, seem arbitrary.<ref name=Sharp1981/>{{refn|Jones (2010, pp. 169–71): "Though classification is an essential feature of all branches of science, there are always hard cases at the boundaries. The boundary of a class is rarely sharp…Scientists should not lose sleep over the hard cases. As long as a classification system is beneficial to economy of description, to structuring knowledge and to our understanding, and hard cases constitute a small minority, then keep it. If the system becomes less than useful, then scrap it and replace it with a system based on different shared characteristics."|group=n}} Metalloids may be grouped with metals;<ref>[[#Tyler1948|Tyler 1948, p. 105]]; [[#Reilly2002|Reilly 2002, pp. 5–6]]</ref> or regarded as nonmetals;<ref>[[#Hampel1976|Hampel & Hawley 1976, p. 174]];</ref> or treated as a sub-category of nonmetals.<ref>[[#Goodrich1844|Goodrich 1844, p. 264]]; [[#TheChemical1897|''The Chemical News'' 1897, p. 189]]; [[#Hampel1976|Hampel & Hawley 1976, p. 191]]; [[#Lewis1993|Lewis 1993, p. 835]]; [[#Hérold2006|Hérold 2006, pp. 149–50]]</ref>{{refn|1=Oderberg<ref>[[#Oderberg2007|Oderberg 2007, p. 97]]</ref> argues on [[ontology|ontological]] grounds that anything not a metal is therefore a nonmetal, and that this includes semi-metals (i.e. metalloids).|group=n}} Other authors have suggested classifying some elements as metalloids "emphasizes that properties change gradually rather than abruptly as one moves across or down the periodic table".<ref name=brown>[[#Brown2006|Brown & Holme 2006, p. 57]]</ref> Some periodic tables distinguish elements that are metalloids and display no formal dividing line between metals and nonmetals. Metalloids are instead shown as occurring in a diagonal band<ref>[[#Wiberg2001|Wiberg 2001, p. 282]]; [[#Simple2005|Simple Memory Art c. 2005]]</ref> or diffuse region.<ref>[[#Chedd1969|Chedd 1969, pp. 12–13]]</ref> The key consideration is to explain the context for the taxonomy in use.
==Properties==
Metalloids usually look like metals but behave largely like nonmetals. Physically, they are shiny, brittle solids with intermediate to relatively good electrical conductivity and the electronic band structure of a semimetal or semiconductor. Chemically, they mostly behave as (weak) nonmetals, have intermediate ionization energies and electronegativity values, and amphoteric or weakly acidic [[oxide]]s. Most of their other physical and chemical properties are [[properties of metals, metalloids and nonmetals|intermediate in nature]].
===Compared to metals and nonmetals===
{{Main|Properties of metals, metalloids and nonmetals}}
Characteristic properties of metals, metalloids, and nonmetals are summarized in the table.<ref>[[#Kneen1972|Kneen, Rogers & Simpson, 1972, p. 263]]. Columns 2 and 4 are sourced from this reference unless otherwise indicated.</ref> Physical properties are listed in order of ease of determination; chemical properties run from general to specific, and then to descriptive.
{|class="wikitable"
|+Properties of metals, metalloids and nonmetals
|- valign=top
! scope="col" style="width:10em;" | Physical property
! scope="col" style="width:20em;" | Metals
! scope="col" style="width:20em;" | Metalloids
! scope="col" style="width:20em;" | Nonmetals
|- valign=top
| scope="row"| Form
| solid; a few liquid at or near room temperature ([[gallium|Ga]], [[mercury (element)|Hg]], [[rubidium|Rb]], [[caesium|Cs]], [[francium|Fr]])<ref>[[#Stoker2010|Stoker 2010, p. 62]]; [[#Chang2002|Chang 2002, p. 304]]. Chang speculates that the melting point of francium would be about 23 °C.</ref>{{refn|1=[[Copernicium]] is reportedly the only metal thought to be a gas at room temperature.<ref>[[#NS1975|New Scientist 1975]]; [[#Soverna2004|Soverna 2004]]; [[#Eichler2007|Eichler et al. 2007]]; [[#Austen2012|Austen 2012]]</ref>|group=n}}
| solid<ref name="Rochow 1966, p.4">[[#Rochow1966|Rochow 1966, p. 4]]</ref>
| majority gaseous<ref>[[#Hunt2000|Hunt 2000, p. 256]]</ref>
|- valign=top
| scope="row"| Appearance
| lustrous (at least when freshly fractured)
| lustrous<ref name="Rochow 1966, p.4"/>
| several colourless; others coloured, or metallic grey to black
|- valign=top
| scope="row"| [[Plasticity (physics)|Plasticity]]
| typically elastic, ductile, malleable
| often brittle<ref name=McQuarrie85>[[#McQuarrie1987|McQuarrie & Rock 1987, p. 85]]</ref>
| often brittle
|- valign=top
| scope="row"| [[Electrical conductivity]]
| good to high{{refn|1=Metals have electrical conductivity values of from 6.9 × 10<sup>3</sup> S•cm<sup>−1</sup> for [[manganese]] to 6.3 × 10<sup>5</sup> for [[silver]].<ref>[[#Desai1984|Desai, James & Ho 1984, p. 1160]]; [[#Matula1979|Matula 1979, p. 1260]]</ref>|group=n}}
| intermediate<ref>[[#Choppin1972|Choppin & Johnsen 1972, p. 351]]</ref> to good{{refn|1=Metalloids have electrical conductivity values of from 1.5 × 10<sup>−6</sup> S•cm<sup>−1</sup> for boron to 3.9 × 10<sup>4</sup> for arsenic.<ref>[[#Schaefer1968|Schaefer 1968, p. 76]]; [[#Carapella1968|Carapella 1968, p. 30]]</ref> If selenium is included as a metalloid the applicable conductivity range would start from ~10<sup>−9</sup> to 10<sup>−12</sup> S•cm<sup>−1</sup>.<ref name="Kozyrev"/>|group=n}}
| poor to good{{refn|1=Nonmetals have electrical conductivity values of from ~10<sup>−18</sup> S•cm<sup>−1</sup> for the elemental gases to 3 × 10<sup>4</sup> in graphite.<ref>[[#Bogoroditskii1967|Bogoroditskii & Pasynkov 1967, p. 77]]; [[#Jenkins1976|Jenkins & Kawamura 1976, p. 88]]</ref>|group=n}}
|- valign=top
| scope="row"| [[Band structure]]
| metallic ([[bismuth|Bi]] = semimetallic)
| are semiconductors or, if not ([[arsenic|As]], [[antimony|Sb]] = semimetallic), exist in semiconducting forms<ref>[[#Hampel1976|Hampel & Hawley 1976, p. 191]]; [[#Wulfsberg2000|Wulfsberg 2000, p. 620]]</ref>
| semiconductor or [[Insulator (electricity)|insulator]]<ref name=Swalin>[[#Swalin1962|Swalin 1962, p. 216]]</ref>
|- valign=top
! scope="col" style="width:10em;" | Chemical property
! scope="col" style="width:20em;" | Metals
! scope="col" style="width:20em;" | Metalloids
! scope="col" style="width:20em;" | Nonmetals
|- valign=top
| General chemical behaviour
| metallic
| nonmetallic<ref>[[#Bailar1989|Bailar et al. 1989, p. 742]]</ref>
| nonmetallic
|- valign=top
| scope="row" |[[Ionization energy]]
| relatively low
| intermediate ionization energies,<ref>[[#Metcalfe1974|Metcalfe, Williams & Castka 1974, p. 86]]</ref> usually falling between those of metals and nonmetals<ref>[[#Chang2002|Chang 2002, p. 306]]</ref>
| relatively high
|- valign=top
| scope="row" |[[Electronegativity]]
| usually low
| have electronegativity values close to 2<ref>[[#Pauling1988|Pauling 1988, p. 183]]</ref> (revised Pauling scale) or within the range of 1.9–2.2 (Allen scale)<ref name="Mann">[[#Mann2000|Mann et al. 2000, p. 2783]]</ref>{{refn|1=Chedd<ref>[[#Chedd1969|Chedd 1969, pp. 24–25]]</ref> defines metalloids as having electronegativity values of 1.8 to 2.2 ([[Allred-Rochow scale]]). He included boron, silicon, germanium, arsenic, antimony, tellurium, polonium, and [[astatine]] in this category. In reviewing Chedd's work, Adler<ref>[[#Adler1969|Adler 1969, pp. 18–19]]</ref> described this choice as arbitrary, as other elements whose electronegativities lie in this range include [[copper]], silver, phosphorus, mercury, and bismuth. He went on to suggest defining a metalloid as "a semiconductor or semimetal" and to include bismuth and selenium in this category.|group=n}}
| high
|- valign=top
| scope="row" |When mixed<br/>with metals
| give [[alloy]]s
| can form alloys<ref>[[#Hultgren1966|Hultgren 1966, p. 648]]; [[#Young2000|Young & Sessine 2000, p. 849]]; [[#Bassett1966|Bassett et al. 1966, p. 602]]</ref>
| ionic or [[interstitial compound]]s formed
|- valign=top
| scope="row" |[[Oxide]]s
| lower oxides [[base (chemistry)|basic]]; higher oxides increasingly [[acid]]ic
| amphoteric or weakly acidic<ref>[[#Rochow1966|Rochow 1966, p. 4]]; [[#Atkins2006|Atkins et al. 2006, pp. 8, 122–23]]</ref>
| acidic
|}
The above table reflects the hybrid nature of metalloids. The properties of ''form, appearance'', and ''behaviour when mixed with metals'' are more like metals. ''Elasticity'' and ''general chemical behaviour'' are more like nonmetals. ''Electrical conductivity, band structure, ionization energy, electronegativity,'' and ''oxides'' are intermediate between the two.
==Common applications==
:''The focus of this section is on the recognised metalloids. Elements less often recognised as metalloids are ordinarily classified as either metals or nonmetals; some of these are included here for comparative purposes.''
Metalloids are too brittle to have any structural uses in their pure forms.<ref>[[#Russell2005|Russell & Lee 2005, pp. 421, 423]]; [[#Gray2009|Gray 2009, p. 23]]</ref> They and their compounds are used in alloys, biological agents (toxicological, nutritional, and medicinal), catalysts, flame retardants, glasses (oxide and metallic), optical storage media and optoelectronics, pyrotechnics, semiconductors, and electronics.{{refn|1=Olmsted and Williams<ref>[[#Olmsted1997|Olmsted & Williams 1997, p. 975]]</ref> commented that, "Until quite recently, chemical interest in the metalloids consisted mainly of isolated curiosities, such as the poisonous nature of arsenic and the mildly therapeutic value of borax. With the development of metalloid semiconductors, however, these elements have become among the most intensely studied".|group=n}}
===Alloys===
[[File:Copper germanium.jpg|thumb|right|[[Copper-germanium alloy]] pellets, likely ~84% Cu; 16% Ge.<ref name="Russell2005401"/> When combined with [[silver]] the result is a [[argentium sterling silver|tarnish resistant sterling silver]]. Also shown are two silver pellets.|alt=Several dozen metallic pellets, reddish-brown. They have a highly polished appearance, as if they had a cellophane coating.]]
Writing early in the history of [[intermetallic|intermetallic compounds]], the British metallurgist Cecil Desch observed that "certain non-metallic elements are capable of forming compounds of distinctly metallic character with metals, and these elements may therefore enter into the composition of alloys". He associated silicon, arsenic, and tellurium, in particular, with the alloy-forming elements.<ref>[[#Desch1914|Desch 1914, p. 86]]</ref> Phillips and Williams<ref>[[#Phillips1965|Phillips & Williams 1965, p. 620]]</ref> suggested that compounds of silicon, germanium, arsenic, and antimony with [[other metal|B metals]], "are probably best classed as alloys".
Among the lighter metalloids, alloys with [[transition metal]]s are well-represented. Boron can form intermetallic compounds and alloys with such metals of the composition M<sub>''n''</sub>B, if ''n'' > 2.<ref>[[#Vanderput1998|Van der Put 1998, p. 123]]</ref> Ferroboron (15% boron) is used to introduce boron into [[steel]]; nickel-boron alloys are ingredients in welding alloys and [[case hardening]] compositions for the engineering industry. Alloys of silicon with [[iron]] and with aluminium are widely used by the steel and automotive industries, respectively. Germanium forms many alloys, most importantly with the [[Group 11 element|coinage metals]].<ref>[[#Klug1958|Klug & Brasted 1958, p. 199]]</ref>
The heavier metalloids continue the theme. Arsenic can form alloys with metals, including [[platinum]] and [[copper]];<ref>[[#Good1813|Good et al. 1813]]</ref> it is also added to copper and its alloys to improve corrosion resistance<ref>[[#Sequeira|Sequeira 2011, p. 776]]</ref> and appears to confer the same benefit when added to magnesium.<ref>[[#Gary|Gary 2013]]</ref> Antimony is well known as an alloy-former, including with the coinage metals. Its alloys include [[pewter]] (a tin alloy with up to 20% antimony) and [[type metal]] (a lead alloy with up to 25% antimony).<ref>[[#Russell2005|Russell & Lee 2005, pp. 405–06; 423–34]]</ref> Tellurium readily alloys with iron, as ferrotellurium (50–58% tellurium), and with copper, in the form of [[Tellurium Copper|copper tellurium]] (40–50% tellurium).<ref>[[#Davidson1973|Davidson & Lakin 1973, p. 627]]</ref> Ferrotellurium is used as a stabilizer for carbon in steel casting.<ref>[[#Wiberg2001|Wiberg 2001, p. 589]]</ref> Of the non-metallic elements less often recognised as metalloids, selenium – in the form of ferroselenium (50–58% selenium) – is used to improve the [[machinability]] of stainless steels.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p. 749]]; [[#Schwartz2002|Schwartz 2002, p. 679]]</ref>
===Biological agents===
[[File:Arsenic trioxide.jpg|thumb|right|[[Arsenic trioxide]] or ''white arsenic'', one of the most toxic and prevalent forms of [[arsenic]]. The [[antileukemic drug|antileukaemic]] properties of white arsenic were first reported in 1878.<ref>[[#Antman|Antman 2001]]</ref>|alt=A clear glass dish on which is a small mound of a white crystalline powder.]] All six of the elements commonly recognised as metalloids have toxic, dietary or medicinal properties.<ref>[[#Řezanka|Řezanka & Sigler 2008]]; [[#Sekhon|Sekhon 2012]]</ref> Arsenic and antimony compounds are especially toxic; boron, silicon, and possibly arsenic, are essential trace elements. Boron, silicon, arsenic, and antimony have medical applications, and germanium and tellurium are thought to have potential.
Boron is used in insecticides<ref>[[#Emsley2001|Emsley 2001, p. 67]]</ref> and herbicides.<ref>[[#Zhang2008|Zhang et al. 2008, p. 360]]</ref> It is an essential trace element.<ref name=SLH>[[#SLH|Science Learning Hub 2009]]</ref> As [[boric acid]], it has antiseptic, antifungal, and antiviral properties.<ref>[[#Skinner|Skinner et al. 1979]]; [[#Tom|Tom, Elden & Marsh 2004, p. 135]]</ref>
Silicon is present in [[silatrane]], a highly toxic rodenticide.<ref>[[#Büchel|Büchel 1983, p. 226]]</ref> Long-term inhalation of silica dust causes [[silicosis]], a fatal disease of the lungs. Silicon is an essential trace element.<ref name=SLH/> [[Silicone]] gel can be applied to badly burned patients to reduce scarring.<ref>[[#Emsley2001|Emsley 2001, p. 391]]</ref>
[[Salt (chemistry)|Salts]] of germanium are potentially harmful to humans and animals if ingested on a prolonged basis.<ref>[[#Schauss1991|Schauss 1991]]; [[#Tao1997|Tao & Bolger 1997]]</ref> There is interest in the pharmacological actions of germanium compounds but no licensed medicine as yet.<ref>[[#Eagleson1994|Eagleson 1994, p. 450]]; [[#EVM|EVM 2003, pp. 197‒202]]</ref>
Arsenic is notoriously poisonous and may also be an [[essential element]] in ultratrace amounts.<ref name=Neilsen>[[#Nielsen|Nielsen 1998]]</ref> During [[World War I]], both sides used "arsenic-based sneezing and vomiting [[diphenylchloroarsine|agents]]…to force enemy soldiers to remove their [[WWI gas mask|gas mask]]s before firing [[mustard gas|mustard]] or [[phosgene]] at them in a second [[salvo]]."<ref>[[#MacKenzie|MacKenzie 2015, p. 36]]</ref> It has been used as a pharmaceutical agent since antiquity, including for the treatment of [[syphilis]] before the development of [[antibiotics]].<ref name=Jaouen>[[#Jaouen|Jaouen & Gibaud 2010]]</ref> Arsenic is also a component of [[melarsoprol]], a medicinal drug used in the treatment of human [[African trypanosomiasis]] or sleeping sickness. In 2003, arsenic trioxide (under the trade name [[Trisenox]]) was re-introduced for the treatment of [[acute promyelocytic leukaemia]], a cancer of the blood and bone marrow.<ref name=Jaouen/> Arsenic in drinking water, which causes lung and bladder cancer, has been associated with a reduction in breast cancer mortality rates.<ref>[[#SmithAH|Smith et al. 2014]]</ref>
Metallic antimony is relatively non-toxic, but most antimony compounds are poisonous.<ref>[[#Stevens1990|Stevens & Klarner, p. 205]]</ref>
Two antimony compounds, [[sodium stibogluconate]] and [[stibophen]], are used as [[antiparasitic|antiparasitical drugs]].<ref>[[#Sneader|Sneader 2005, pp. 57–59]]</ref>
Elemental tellurium is not considered particularly toxic; two grams of sodium tellurate, if administered, can be lethal.<ref>[[#Keall1946|Keall, Martin and Tunbridge 1946]]</ref> People exposed to small amounts of airborne tellurium exude a foul and persistent garlic-like odour.<ref>[[#Emsley2001|Emsley 2001, p. 426]]</ref> Tellurium dioxide has been used to treat [[seborrhoeic dermatitis]]; other tellurium compounds were used as [[antimicrobial]] agents before the development of antibiotics.<ref>[[#Oldfield1974|Oldfield et al. 1974, p. 65]]; [[#Turner2011|Turner 2011]]</ref> In the future, such compounds may need to be substituted for antibiotics that have become ineffective due to bacterial resistance.<ref>[[#Ba|Ba et al. 2010]]; [[#Daniel-Hoffmann|Daniel-Hoffmann, Sredni & Nitzan 2012]]; [[#Molina-Quiroz|Molina-Quiroz et al. 2012]]</ref>
Of the elements less often recognised as metalloids, beryllium and lead are noted for their toxicity; [[lead arsenate]] has been extensively used as an insecticide.<ref>[[#Peryea|Peryea 1998]]</ref> Sulfur is one of the oldest of the fungicides and pesticides. Phosphorus, sulfur, zinc, selenium, and iodine are essential nutrients, and aluminium, tin, and lead may be.<ref name=Neilsen/> Sulfur, gallium, selenium, iodine, and bismuth have medicinal applications. Sulfur is a constituent of [[Sulfonamide (medicine)|sulfonamide drugs]], still widely used for conditions such as acne and urinary tract infections.<ref>[[#Hager|Hager 2006, p. 299]]</ref> [[Gallium nitrate]] is used to treat the side effects of cancer;<ref>[[#Apseloff|Apseloff 1999]]</ref> gallium citrate, a [[radiopharmaceutical]], facilitates imaging of inflamed body areas.<ref>[[#Trivedi|Trivedi, Yung & Katz 2013, p. 209]]</ref> [[Selenium sulfide]] is used in medicinal shampoos and to treat skin infections such as [[tinea versicolor]].<ref>[[#Emsley2001|Emsley 2001, p. 382]]; [[#Burkhart|Burkhart, Burkhart & Morrell 2011]]</ref> Iodine is used as a disinfectant in various forms. Bismuth is an ingredient in some [[antibacterial]]s.<ref>[[#Thomas|Thomas, Bialek & Hensel 2013, p. 1]]</ref>
===Catalysts===
[[Boron trifluoride]] and [[boron trichloride|trichloride]] are used as homogeneous [[catalyst]]s in organic synthesis and electronics; the [[boron tribromide|tribromide]] is used in the manufacture of [[diborane]].<ref>[[#Perry|Perry 2011, p. 74]]</ref> Non-toxic boron [[ligand]]s could replace toxic phosphorus ligands in some transition metal catalysts.<ref>[[#UCR|UCR Today 2011]]; [[#Wang|Wang & Robinson 2011]]; [[#Kinjo|Kinjo et al. 2011]]</ref> [[Silica sulfuric acid]] (SiO<sub>2</sub>OSO<sub>3</sub>H) is used in organic reactions.<ref>[[#Kauthale|Kauthale et al. 2015]]</ref> Germanium dioxide is sometimes used as a catalyst in the production of [[Polyethylene terephthalate|PET]] plastic for containers;<ref>[[#Gunn|Gunn 2014, pp. 188, 191]]</ref> cheaper antimony compounds, such as the trioxide or [[antimony triacetate|triacetate]], are more commonly employed for the same purpose<ref>[[#Gupta|Gupta, Mukherjee & Cameotra 1997, p. 280]]; [[#Thomas2012|Thomas & Visakh 2012, p. 99]]</ref> despite concerns about antimony contamination of food and drinks.<ref>[[#Muncke|Muncke 2013]]</ref> Arsenic trioxide has been used in the production of [[natural gas]], to boost the removal of [[carbon dioxide]], as have [[selenous acid]] and [[tellurous acid]].<ref>[[#Mokhatab|Mokhatab & Poe 2012, p. 271]]</ref> Selenium acts as a catalyst in some microorganisms.<ref>[[#Craig|Craig, Eng & Jenkins 2003, p. 25]]</ref> Tellurium, its dioxide, and its [[tellurium tetrachloride|tetrachloride]] are strong catalysts for air oxidation of carbon above 500 °C.<ref>[[#McKee|McKee 1984]]</ref> [[Graphite oxide]] can be used as a catalyst in the synthesis of [[imine]]s and their derivatives.<ref>[[#Hai|Hai et al. 2012]]</ref> [[Activated carbon]] and [[alumina]] have been used as catalysts for the removal of sulfur contaminants from natural gas.<ref>[[#Kohl|Kohl & Nielsen 1997, pp. 699–700]]</ref> [[Titanium]] doped aluminium has been suggested as a substitute for [[noble metal]] catalysts used in the production of industrial chemicals.<ref>[[#Chopra|Chopra et al. 2011]]</ref>
===Flame retardants===
Compounds of boron, silicon, arsenic, and antimony have been used as [[flame retardant]]s. Boron, in the form of [[borax]], has been used as a textile flame retardant since at least the 18th century.<ref>[[#LeBras|Le Bras, Wilkie & Bourbigot 2005, p. v]]</ref> Silicon compounds such as silicones, [[silane]]s, [[silsesquioxane]], [[silica]], and [[silicate]]s, some of which were developed as alternatives to more toxic [[halogenation|halogenated]] products, can considerably improve the flame retardancy of plastic materials.<ref>[[#Wilkie|Wilkie & Morgan 2009, p. 187]]</ref>
Arsenic compounds such as [[sodium arsenite]] or [[sodium arsenate]] are effective flame retardants for wood but have been less frequently used due to their toxicity.<ref>[[#Locke1956|Locke et al. 1956, p. 88]]</ref> Antimony trioxide is a flame retardant.<ref>[[#Carlin|Carlin 2011, p. 6.2]]</ref> [[Aluminium hydroxide]] has been used as a wood-fibre, rubber, plastic, and textile flame retardant since the 1890s.<ref>[[#Evans|Evans 1993, pp. 257–28]]</ref> Apart from aluminium hydroxide, use of phosphorus based flame-retardants – in the form of, for example, [[organophosphate]]s – now exceeds that of any of the other main retardant types. These employ boron, antimony, or [[halogenated hydrocarbon]] compounds.<ref>[[#Corbridge|Corbridge 2013, p. 1149]]</ref>
===Glass formation===
[[File:Fibreoptic4.jpg|thumb|right|[[Optical fibers]], usually made of pure [[silicon dioxide]] glass, with additives such as [[boron trioxide]] or [[germanium dioxide]] for increased sensitivity|alt=A bunch of pale yellow semi-transparent thin strands, with bright points of white light at their tips.]]
The oxides [[boron trioxide|B<sub>2</sub>O<sub>3</sub>]], [[silicon dioxide|SiO<sub>2</sub>]], [[germanium dioxide|GeO<sub>2</sub>]], [[arsenic trioxide|As<sub>2</sub>O<sub>3</sub>]], and [[antimony trioxide|Sb<sub>2</sub>O<sub>3</sub>]] readily form [[glass]]es. [[Tellurium dioxide|TeO<sub>2</sub>]] forms a glass but this requires a "heroic quench rate"<ref name=K2002/> or the addition of an impurity; otherwise the crystalline form results.<ref name=K2002>[[#Kaminow2002|Kaminow & Li 2002, p. 118]]</ref> These compounds are used in chemical, domestic, and industrial glassware<ref>[[#Deming1925|Deming 1925]], pp. 330 (As<sub>2</sub>O<sub>3</sub>), 418 (B<sub>2</sub>O<sub>3</sub>; SiO<sub>2</sub>; Sb<sub>2</sub>O<sub>3</sub>); [[#Witt1968|Witt & Gatos 1968, p. 242]] (GeO<sub>2</sub>)</ref> and optics.<ref>[[#Eagleson1994|Eagleson 1994, p. 421]] (GeO<sub>2</sub>); [[#Rothenberg1976|Rothenberg 1976, 56, 118–19]] (TeO<sub>2</sub>)</ref> Boron trioxide is used as a [[glass fibre]] additive,<ref>[[#Geckeler1987|Geckeler 1987, p. 20]]</ref> and is also a component of [[borosilicate glass]], widely used for laboratory glassware and domestic ovenware for its low thermal expansion.<ref>[[#Kreith2005|Kreith & Goswami 2005, pp. 12–109]]</ref> Most ordinary glassware is made from silicon dioxide.<ref>[[#Russell2005|Russell & Lee 2005, p. 397]]</ref> Germanium dioxide is used as a glass fibre additive, as well as in infrared optical systems.<ref>[[#Butterman2005|Butterman & Jorgenson 2005, pp. 9–10]]</ref> Arsenic trioxide is used in the glass industry as a [[glass coloring and color marking|decolourizing]] and fining agent (for the removal of bubbles),<ref>[[#Shelby|Shelby 2005, p. 43]]</ref> as is antimony trioxide.<ref>[[#Butterman2004|Butterman & Carlin 2004, p. 22]]; [[#Russell2005|Russell & Lee 2005, p. 422]]</ref> Tellurium dioxide finds application in laser and [[nonlinear optics]].<ref>[[#Träger2007|Träger 2007, pp. 438, 958]]; [[#Eranna2011|Eranna 2011, p. 98]]</ref>
[[Amorphous]] [[metallic glass]]es are generally most easily prepared if one of the components is a metalloid or "near metalloid" such as boron, carbon, silicon, phosphorus or germanium.<ref>[[#Rao2002|Rao 2002, p. 552]]; [[#Loffler|Löffler, Kündig & Dalla Torre 2007, p. 17–11]]</ref>{{refn|1=Research published in 2012 suggests that metal-metalloid glasses can be characterised by an interconnected atomic packing scheme in which metallic and [[covalent]] bonding structures coexist.<ref>[[#Guan|Guan et al. 2012]]; [[#World|WPI-AIM 2012]]</ref>|group=n}} Aside from thin films deposited at very low temperatures, the first known metallic glass was an alloy of composition Au<sub>75</sub>Si<sub>25</sub> reported in 1960.<ref>[[#Klement|Klement, Willens & Duwez 1960]]; [[#Wanga|Wanga, Dongb & Shek 2004, p. 45]]</ref> A metallic glass having a strength and toughness not previously seen, of composition Pd<sub>82.5</sub>P<sub>6</sub>Si<sub>9.5</sub>Ge<sub>2</sub>, was reported in 2011.<ref>[[#Demetriou|Demetriou et al. 2011]]; [[#Oliwenstein|Oliwenstein 2011]]</ref>
Phosphorus, selenium, and lead, which are less often recognised as metalloids, are also used in glasses. [[Phosphate glass]] has a substrate of phosphorus pentoxide (P<sub>2</sub>O<sub>5</sub>), rather than the silica (SiO<sub>2</sub>) of conventional silicate glasses. It is used, for example, to make [[sodium lamp]]s.<ref>[[#Karabulut|Karabulut et al. 2001, p. 15]]; [[#Haynes|Haynes 2012, pp. 4–26]]</ref> Selenium compounds can be used both as decolourising agents and to add a red colour to glass.<ref>[[#Schwartz2002|Schwartz 2002, pp. 679–80]]</ref> Decorative glassware made of traditional [[lead glass]] contains at least 30% [[lead(II) oxide]] (PbO); lead glass used for radiation shielding may have up to 65% PbO.<ref>[[#Carter|Carter & Norton 2013, p. 403]]</ref> Lead-based glasses have also been extensively used in electronic components, enamelling, sealing and glazing materials, and solar cells. Bismuth based oxide glasses have emerged as a less toxic replacement for lead in many of these applications.<ref>[[#Maeder|Maeder 2013, pp. 3, 9–11]]</ref>
===Optical storage and optoelectronics===
Varying compositions of [[GeSbTe]] ("GST alloys") and [[AgInSbTe|Ag- and In- doped Sb<sub>2</sub>Te]] ("AIST alloys"), being examples of [[phase-change material]]s, are widely used in rewritable [[optical disc]]s and [[phase-change memory]] devices. By applying heat, they can be switched between amorphous (glassy) and [[crystalline]] states. The change in optical and electrical properties can be used for information storage purposes.<ref>[[#Tominaga2006|Tominaga 2006, pp. 327–28]]; [[#Chung2010|Chung 2010, pp. 285–86]]; [[#Kolobov 2012|Kolobov & Tominaga 2012, p. 149]]</ref> Future applications for GeSbTe may include, "ultrafast, entirely solid-state displays with nanometre-scale pixels, semi-transparent 'smart' glasses, 'smart' contact lenses, and artificial retina devices."<ref>[[#NS2014|New Scientist 2014]]; [[#Hosseini|Hosseini, Wright & Bhaskaran 2014]]; [[#Farandos|Farandos et al. 2014]]</ref>
===Pyrotechnics===
[[File:Blue Light.JPG|thumb|right|upright|Archaic [[blue light (pyrotechnic signal)|blue light signal]], fuelled by a mixture of [[sodium nitrate]], [[sulfur]], and (red) [[arsenic trisulfide]]<ref>[[#OO|Ordnance Office 1863, p. 293]]</ref>|alt=A man is standing in the dark. He is holding out a short stick at mid-chest level. The end of the stick is alight, burning very brightly, and emitting smoke.]]
The recognised metalloids have either pyrotechnic applications or associated properties. Boron and silicon are commonly encountered;<ref name=Kos>[[#Kosanke|Kosanke 2002, p. 110]]</ref> they act somewhat like metal fuels.<ref>[[#Ellern|Ellern 1968, pp. 246, 326–27]]</ref> Boron is used in [[pyrotechnic initiator]] compositions (for igniting other hard-to-start compositions), and in [[delay composition]]s that burn at a constant rate.<ref name=Conkling82>[[#Conkling|Conkling & Mocella 2010, p. 82]]</ref> [[Boron carbide]] has been identified as a possible replacement for more toxic [[barium]] or [[hexachloroethane]] mixtures in smoke munitions, signal flares, and fireworks.<ref>[[#Crow|Crow 2011]]; [[#Daily|Mainiero 2014]]</ref> Silicon, like boron, is a component of initiator and delay mixtures.<ref name=Conkling82/> Doped germanium can act as a variable speed [[thermite]] fuel.{{refn|1=The reaction involved is Ge + 2 [[MoO3|MoO<sub>3</sub>]] → GeO<sub>2</sub> + 2 [[MoO2|MoO<sub>2</sub>]]. Adding arsenic or antimony ([[extrinsic semiconductor#n-type semiconductors|n-type]] electron donors) increases the rate of reaction; adding gallium or indium ([[intrinsic semiconductor#p-type semiconductors|p-type]] electron acceptors) decreases it.<ref>[[#Schwab|Schwab & Gerlach 1967]]; [[#Yetter|Yetter 2012, p. 81]]; [[#Lipscomb|Lipscomb 1972, pp. 2–3, 5–6, 15]]</ref>|group=n}} [[Arsenic trisulfide]] As<sub>2</sub>S<sub>3</sub> was used in old [[blue light (pyrotechnic signal)|naval signal lights]]; in fireworks to make white stars;<ref>[[#Ellern|Ellern 1968, p. 135]]; [[#Weingart|Weingart 1947, p. 9]]</ref> in yellow [[smoke screen]] mixtures; and in initiator compositions.<ref>[[#Conkling|Conkling & Mocella 2010, p. 83]]</ref> [[stibnite|Antimony trisulfide]] Sb<sub>2</sub>S<sub>3</sub> is found in white-light fireworks and in [[flash powder|flash and sound]] mixtures.<ref>[[#Conkling|Conkling & Mocella 2010, pp. 181, 213]]</ref> Tellurium has been used in delay mixtures and in [[blasting cap]] initiator compositions.<ref name=Ellern>[[#Ellern|Ellern 1968, pp. 209–10, 322]]</ref>
Carbon, aluminium, phosphorus, and selenium continue the theme. Carbon, in [[black powder]], is a constituent of fireworks rocket propellants, bursting charges, and effects mixtures, and military delay fuses and igniters.<ref>[[#RussellF|Russell 2009, pp. 15, 17, 41, 79–80]]</ref>{{refn|1=Ellern, writing in ''Military and Civilian Pyrotechnics'' (1968), comments that [[carbon black]] "has been specified for and used in a nuclear air-burst simulator."<ref>[[#Ellern|Ellern 1968, p. 324]]</ref>|group=n}} Aluminium is a common pyrotechnic ingredient,<ref name=Kos/> and is widely employed for its capacity to generate light and heat,<ref>[[#Ellern|Ellern 1968, p. 328]]</ref> including in thermite mixtures.<ref>[[#Conkling|Conkling & Mocella 2010, p. 171]]</ref> Phosphorus can be found in smoke and incendiary munitions, [[Armstrong's mixture|paper caps]] used in [[cap gun|toy guns]], and [[party popper]]s.<ref>[[#Conkling|Conkling & Mocella 2011, pp. 83–84]]</ref> Selenium has been used in the same way as tellurium.<ref name=Ellern/>
===Semiconductors and electronics===
[[File:Semiconductor-1.jpg|thumb|left|[[Semiconductor]]-based electronic components. From left to right: a [[transistor]], an [[integrated circuit]], and an [[LED]]. The elements commonly recognised as metalloids find widespread use in such devices, as elemental or [[compound semiconductor]] constituents ([[silicon|Si]], [[germanium|Ge]] or [[GaAs]], for example) or as [[doping (semiconductor)|doping agents]] ([[boron|B]], [[antimony|Sb]], [[tellurium|Te]], for example).|alt=A small square plastic piece with three parallel wire protrusions on one side; a larger rectangular plastic chip with multiple plastic and metal pin-like legs; and a small red light globe with two long wires coming out of its base.]]
All the elements commonly recognised as metalloids (or their compounds) have been used in the semiconductor or solid-state electronic industries.<ref>[[#Berger1997|Berger 1997, p. 91]]; [[#Hampel1968|Hampel 1968, passim]]</ref>
Some properties of boron have limited its use as a semiconductor. It has a high melting point, single [[crystal]]s are relatively hard to obtain, and introducing and retaining controlled impurities is difficult.<ref>[[#Rochow1966|Rochow 1966, p. 41]]; [[#Berger1997|Berger 1997, pp. 42–43]]</ref>
Silicon is the leading commercial semiconductor; it forms the basis of modern electronics (including standard solar cells)<ref name=Bom>[[#Bomgardner|Bomgardner 2013, p. 20]]</ref> and information and communication technologies.<ref>[[#Russell2005|Russell & Lee 2005, p. 395]]; [[#Brown2009|Brown et al. 2009, p. 489]]</ref> This was despite the study of semiconductors, early in the 20th century, having been regarded as the "physics of dirt" and not deserving of close attention.<ref>[[#Haller 2006|Haller 2006, p. 4]]: "The study and understanding of the physics of semiconductors progressed slowly in the 19th and early 20th centuries ... Impurities and defects ... could not be controlled to the degree necessary to obtain reproducible results. This led influential physicists, including [[Wolfgang Pauli|W. Pauli]] and [[Isidor Isaac Rabi|I. Rabi]], to comment derogatorily on the 'Physics of Dirt'."; [[#Hoddeson2007|Hoddeson 2007, pp. 25–34 (29)]]</ref>
Germanium has largely been replaced by silicon in semiconducting devices, being cheaper, more resilient at higher operating temperatures, and easier to work during the microelectronic fabrication process.<ref name=Russell2005401>[[#Russell2005|Russell & Lee 2005, p. 401]]; [[#Büchel2003|Büchel, Moretto & Woditsch 2003, p. 278]]</ref> Germanium is still a constituent of semiconducting [[silicon-germanium]] "alloys" and these have been growing in use, particularly for wireless communication devices; such alloys exploit the higher carrier mobility of germanium.<ref name=Russell2005401/> The synthesis of gram-scale quantities of semiconducting [[germanane]] was reported in 2013. This consists of one-atom thick sheets of hydrogen-terminated germanium atoms, analogous to [[graphane]]. It conducts electrons more than ten times faster than silicon and five times faster than germanium, and is thought to have potential for optoelectronic and sensing applications.<ref>[[#Bianco2013|Bianco et al. 2013]]</ref> The development of a germanium-wire based anode that more than doubles the capacity of [[lithium-ion battery|lithium-ion batteries]] was reported in 2014.<ref>[[#Limerick|University of Limerick 2014]]; [[#Kennedy|Kennedy et al. 2014]]</ref> In the same year, Lee et al. reported that defect-free crystals of [[graphene]] large enough to have electronic uses could be grown on, and removed from, a germanium substrate.<ref>[[#Lee|Lee et al. 2014]]</ref>
Arsenic and antimony are not semiconductors in their [[standard state#Liquids and solids|standard states]]. Both form [[compound semiconductor|type III-V semiconductors]] (such as GaAs, [[AlSb]] or GaInAsSb) in which the average number of valence electrons per atom is the same as that of [[Group 14]] elements, but they have [[direct band gap]]s. These compounds are preferred for optical applications.<ref>[[#Russell2005|Russell & Lee 2005, pp. 421–22, 424]]</ref> Antimony nanocrystals may enable [[lithium-ion batteries]] to be replaced by more powerful [[sodium-ion battery|sodium ion batteries]].<ref>[[#He|He et al. 2014]]</ref>
Tellurium, which is a semiconductor in its standard state, is used mainly as a component in [[list of semiconductor materials|type II/VI]] semiconducting-[[chalcogenide]]s; these have applications in electro-optics and electronics.<ref>[[#Berger1997|Berger 1997, p. 91]]</ref> [[Cadmium telluride]] (CdTe) is used in solar modules for its high conversion efficiency, low manufacturing costs, and large [[band gap]] of 1.44 eV, letting it absorb a wide range of wavelengths.<ref name=Bom/> [[Bismuth telluride]] (Bi<sub>2</sub>Te<sub>3</sub>), alloyed with selenium and antimony, is a component of [[thermoelectric materials|thermoelectric devices]] used for refrigeration or portable power generation.<ref>[[#ScienceDaily|ScienceDaily 2012]]</ref>
Five metalloids – boron, silicon, germanium, arsenic, and antimony – can be found in cell phones (along with at least 39 other metals and nonmetals).<ref>[[#Reardon2005|Reardon 2005]]; [[#Meskers|Meskers, Hagelüken & Van Damme 2009, p. 1131]]</ref> Tellurium is expected to find such use.<ref>[[#The Economist|The Economist 2012]]</ref> Of the less often recognised metalloids, phosphorus, gallium (in particular) and selenium have semiconductor applications. Phosphorus is used in trace amounts as a [[dopant]] for [[n-type semiconductor]]s.<ref>[[#Whitten2007|Whitten 2007, p. 488]]</ref> The commercial use of gallium compounds is dominated by semiconductor applications – in integrated circuits, cell phones, [[laser diode]]s, [[light-emitting diode]]s, [[photodetector]]s, and [[solar cell]]s.<ref>[[#Jaskula|Jaskula 2013]]</ref> Selenium is used in the production of solar cells<ref>[[#GES|German Energy Society 2008, pp. 43–44]]</ref> and in high-energy [[surge protector]]s.<ref>[[#Patel|Patel 2012, p. 248]]</ref>
Boron, silicon, germanium, antimony, and tellurium,<ref>[[#Moore2014|Moore 2104]]; [[#Utah|University of Utah 2014]]; [[#Xu|Xu et al. 2014]]</ref> as well as heavier metals and metalloids such as Sm, Hg, Tl, Pb, Bi, and Se,<ref>[[#Yang|Yang et al. 2012, p. 614]]</ref> can be found in [[topological insulator]]s. These are alloys<ref>[[#Moore2010|Moore 2010, p. 195]]</ref> or compounds which, at ultracold temperatures or room temperature (depending on their composition), are metallic conductors on their surfaces but insulators through their interiors.<ref>[[#Moore2011|Moore 2011]]</ref> [[Cadmium arsenide]] Cd<sub>3</sub>As<sub>2</sub>, at about 1 K, is a Dirac-semimetal – a bulk electronic analogue of graphene – in which electrons travel effectively as massless particles.<ref>[[#Liu|Liu 2014]]</ref> These two classes of material are thought to have potential [[topological quantum computing|quantum computing]] applications.<ref>[[#Bradley|Bradley 2014]]; [[#Utah|University of Utah 2014]]</ref>
{{clear}}
==Nomenclature and history==
===Derivation and other names===
Several names are sometimes used synonymously although some of these have other meanings that are not necessarily interchangeable: ''amphoteric element,''<ref>[[#Foster1936|Foster 1936, pp. 212–13]]; [[#Brownlee1936|Brownlee et al. 1943, p. 293]]</ref> ''boundary element,''<ref>[[#Calderazzo|Calderazzo, Ercoli & Natta 1968, p. 257]]</ref> ''half-way element,''<ref>[[#Walters1982|Walters 1982, pp. 32–33]]</ref> ''near metal,''<ref name=tyler>[[#Tyler1948|Tyler 1948, p. 105]]</ref> ''meta-metal,''<ref>[[#Foster1958|Foster & Wrigley 1958, p. 218]]: "The elements may be grouped into two classes: those that are ''metals'' and those that are ''nonmetals.'' There is also an intermediate group known variously as ''metalloids,'' ''meta-metals,'' ''semiconductors''."</ref> ''semiconductor,''<ref>[[#Slade2006|Slade 2006, p. 16]]</ref> ''semimetal''<ref>[[#Corwin2005|Corwin 2005, p. 80]]</ref> and ''submetal''.<ref>[[#Barsanov1974|Barsanov & Ginzburg 1974, p. 330]]</ref> "Amphoteric element" is sometimes used more broadly to include transition metals capable of forming [[oxyanion]]s, such as chromium and [[manganese]].<ref>[[#Bradbury1957|Bradbury et al. 1957, pp. 157, 659]]</ref> "Meta-metal" is sometimes used instead to refer to certain metals ([[beryllium|Be]], [[zinc|Zn]], [[cadmium|Cd]], [[mercury (element)|Hg]], [[indium|In]], [[thallium|Tl]], [[Tin#Physical properties|β-Sn]], [[lead|Pb]]) located just to the left of the metalloids on standard periodic tables.<ref name="Klemm">[[Metalloid#Klemm1950|Klemm 1950, pp. 133–42]]; [[Metalloid#Reilly2004|Reilly 2004, p. 4]]</ref> These metals tend to have distorted crystalline structures, electrical conductivity values at the lower end of those of metals, and amphoteric (weakly basic) oxides.<ref>[[#King2004|King 2004, pp. 196–98]]; [[#Ferro2008|Ferro & Saccone 2008, p. 233]]</ref> The names ''amphoteric element'' and ''semiconductor'' are problematic as some elements referred to as metalloids do not show marked amphoteric behaviour (bismuth, for example)<ref>[[#Lister|Lister 1965, p. 54]]</ref> or semiconductivity (polonium)<ref name="Cotton FA 1999, p.502"/> in their most stable forms.
===Origin and usage===
{{Main|Origin and use of the term metalloid}}
The origin and usage of the term ''metalloid'' is convoluted. The "Manual of Metalloids" published in 1864 divided all elements into either metals or metalloids.<ref>Apjohn, J. (1864). Manual of the Metalloids. United Kingdom: Longman.</ref>{{rp|31}} Earlier usage in [[mineralogy]], to describe a mineral having a metallic appearance, can be sourced to as early as 1800.<ref>[[#Pinkerton1800|Pinkerton 1800, p. 81]]</ref> Since the mid-20th century it has been used to refer to intermediate or borderline chemical elements.<ref name="ReferenceA">[[#Goldsmith1982|Goldsmith 1982, p. 526]]</ref> The [[International Union of Pure and Applied Chemistry]] (IUPAC) previously recommended abandoning the term metalloid, and suggested using the term ''semimetal'' instead.<ref>[[#Friend1953|Friend 1953, p. 68]]; [[#IUPAC1959|IUPAC 1959, p. 10]]; [[#IUPAC1971|IUPAC 1971, p. 11]]</ref> Use of this latter term has more recently been discouraged by Atkins et al.<ref name=Atkins2010p20>[[#Atkins2010|Atkins et al. 2010, p. 20]]</ref> as it has a more common meaning that refers to the [[electronic band structure]] of a substance rather than the overall classification of an element. The most recent IUPAC publications on nomenclature and terminology do not include any recommendations on the usage of the terms metalloid or semimetal.<ref>[[#IUPAC2005|IUPAC 2005]]; [[#IUPAC2006|IUPAC 2006–]]</ref>
==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. 2460–61]]; [[#Oganov2009|Oganov et al. 2009, pp. 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. 331]]; [[#Oganov2010|Oganov 2010, p. 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. 333]]</ref> and has a melting point of 2076 °C (cf. steel ~1370 °C).<ref>[[#Kross|Kross 2011]]</ref> Boron is a semiconductor;<ref>[[#Berger1997|Berger 1997, p. 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. 144]]</ref> (about 200 times less than that of tap water)<ref>[[#Kopp|Kopp, Lipták & Eren 2003, p. 221]]</ref> and it has a band gap of about 1.56 eV.<ref>[[#Prudenziati1977|Prudenziati 1977, p. 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. 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. 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. 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. 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. 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. 141]]; [[#Henderson2000|Henderson 2000, p. 58]]; [[#Housecroft2008|Housecroft & Sharpe 2008, pp. 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. 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. 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. 59]]; [[#Wiberg2001|Wiberg 2001, p. 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. 145]]</ref> Analogies between boron and transition metals have been noted in the formation of [[complex (chemistry)|complexes]],<ref>[[#Houghton1979|Houghton 1979, p. 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. 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. 204–05, 207]]</ref>{{refn|1=On the analogy between boron and metals, Greenwood<ref>[[#Greenwood2001|Greenwood 2001, p. 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 ... Indeed, metals have been referred to as "honorary boron atoms" or even as "flexiboron atoms". The converse of this relationship is clearly also valid ..."|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. 128–32]]; [[#MacKay2002|MacKay, MacKay & Henderson 2002, pp. 439–40]]; [[#Kneen1972|Kneen, Rogers & Simpson 1972, p. 394]]; [[#Hiller1960|Hiller & Herber 1960, inside front cover; p. 225]]</ref> Given its high charge-to-size ratio, boron bonds covalently in nearly all of its compounds;<ref>[[#Sharp1983|Sharp 1983, p. 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. 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. 59]]</ref> weakly acidic,<ref>[[#Mahan1965|Mahan 1965, p. 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. 81]].</ref> It reacts with [[alkali]]es to give various borates.<ref>[[#Lidin|Lidin 1996, p. 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. 111]]; [[#Wiberg2001|Wiberg 2001, p. 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. 506]]</ref>|group=n}} and a glass former.<ref name=Rao22>[[#Rao2002|Rao 2002, p. 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. 1]]</ref>|group=n}} have been known since the 19th century (see [[organoboron chemistry]]).<ref>[[#Haiduc1985|Haiduc & Zuckerman 1985, p. 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. 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. 824]]</ref> It is a relatively unreactive element.<ref name=Greenwood331/> According to Rochow,<ref>[[#Rochow1973|Rochow 1973, pp. 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. 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. 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. 1337, 1340]]</ref> Silicon dissolves in hot aqueous [[alkali]]s with the evolution of [[hydrogen]], as do metals<ref>[[#Allen1968|Allen & Ordway 1968, p. 152]]</ref> such as beryllium, aluminium, zinc, gallium or indium.<ref>[[#Eagleson1994|Eagleson 1994, pp. 48, 127, 438, 1194]]; [[#Massey2000|Massey 2000, p. 191]]</ref> It melts at 1414 °C. Silicon is a semiconductor with an electrical conductivity of 10<sup>−4</sup> S•cm<sup>−1</sup><ref>[[#Orton2004|Orton 2004, p. 7]]. This is a typical value for high-purity silicon.</ref> and a band gap of about 1.11 eV.<ref name=R393>[[#Russell2005|Russell & Lee 2005, p. 393]]</ref> When it melts, silicon becomes a reasonable metal<ref>[[#Coles1976|Coles & Caplin 1976, p. 106]]</ref> with an electrical conductivity of 1.0–1.3 × 10<sup>4</sup> S•cm<sup>−1</sup>, similar to that of liquid mercury.<ref>[[#Glazov1969|Glazov, Chizhevskaya & Glagoleva 1969, pp. 59–63]]; [[#Allen1987|Allen & Broughton 1987, p. 4967]]</ref>
The chemistry of silicon is generally nonmetallic (covalent) in nature.<ref>[[#Cotton1995|Cotton, Wilkinson & Gaus 1995, p. 393]]</ref> It is not known to form a cation.<ref>[[#Wiberg2001|Wiberg 2001, p. 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. 723]]</ref> It shows fewer tendencies to anionic behaviour than ordinary nonmetals.<ref name=Cox>[[#Cox2004|Cox 2004, p. 27]]</ref> Its solution chemistry is characterised by the formation of oxyanions.<ref name=Hiller225>[[#Hiller1960|Hiller & Herber 1960, inside front cover; p. 225]]</ref> The high strength of the [[silicon–oxygen bond]] dominates the chemical behaviour of silicon.<ref>[[#Kneen1972|Kneen, Rogers and Simpson 1972, p. 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. 319, 321]]</ref> The oxide SiO<sub>2</sub> is polymeric in structure,<ref name=Pudd59/> weakly acidic,<ref>[[#Smith1990|Smith 1990, p. 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. 851, 858]]</ref> The latter reaction "is sometimes quoted as evidence of basic [that is, metallic] properties".<ref>[[#Barnett|Barmett & Wilson 1959, p. 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. 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. 371]]</ref> It has a density of 5.323 g/cm<sup>3</sup> and is hard and brittle.<ref>[[#Cusack1967|Cusack 1967, p. 193]]</ref> It is mostly unreactive at room temperature{{refn|1=Temperatures above 400 °C are required to form a noticeable surface oxide layer.<ref>[[#Russell2005|Russell & Lee 2005, pp. 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. 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. 273]]</ref> It melts at 938 °C. Germanium is a semiconductor with an electrical conductivity of around 2 × 10<sup>−2</sup> S•cm<sup>−1</sup><ref name=Greenwood373/> and a band gap of 0.67 eV.<ref>[[#Russell2005|Russell & Lee 2005, p. 399]]</ref> Liquid germanium is a metallic conductor, with an electrical conductivity similar to that of liquid mercury.<ref>[[#Berger1997|Berger 1997, pp. 71–72]]</ref>
Most of the chemistry of germanium is characteristic of a nonmetal.<ref>[[#Jolly1966|Jolly 1966, pp. 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. 55]]</ref> in "many metallic halides, hydroxides, and chalcides"); Everest<ref>[[#Everest|Everest 1953, p. 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. 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. 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. 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. 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. 140]]</ref> who says that, "[germanium] forms no aquacations"; Ladd<ref>[[#Ladd|Ladd 1999, p. 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. 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. 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. 66]]; [[#Wiberg2001|Wiberg 2001, p. 895]]</ref> Germanium analogues of all of the major types of silicates have been prepared.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p. 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. 38]]; [[#Wells1984|Wells 1984, p. 1175]]</ref> The oxide GeO<sub>2</sub> is polymeric,<ref name=Pudd59/> amphoteric,<ref>[[#Cooper1968|Cooper 1968, pp. 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. 178, 188–89]]</ref> Up to the 1930s germanium was considered to be a poorly conducting metal;<ref>[[#Haller 2006|Haller 2006, p. 3]]</ref> it has occasionally been classified as a metal by later writers.<ref>[[#Walker|See, for example, Walker & Tarn 1990, p. 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. 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. 552]]</ref> Arsenic [[sublimation (phase transition)|sublimes]] at 615 °C. The vapour is lemon-yellow and smells like garlic.<ref>[[#Parkes1943|Parkes & Mellor 1943, p. 740]]</ref> Arsenic only melts under a pressure of 38.6 [[Atmosphere (unit)|atm]], at 817 °C.<ref>[[#Russell2005|Russell & Lee 2005, p. 420]]</ref> It is a semimetal with an electrical conductivity of around 3.9 × 10<sup>4</sup> S•cm<sup>−1</sup><ref name="Carapella1968p30">[[#Carapella1968|Carapella 1968, p. 30]]</ref> and a band overlap of 0.5 eV.<ref name="Barfuß 1981, p. 967">[[#Barfuß1981|Barfuß et al. 1981, p. 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 eV or 0.4 eV. It can also be prepared in a semiconducting [[amorphous solid|amorphous]] form, with a band gap of around 1.2–1.4 eV.<ref>[[#Greaves1974|Greaves, Knights & Davis 1974, p. 369]]; [[#Madelung2004|Madelung 2004, pp. 405, 410]]</ref>|group=n}} Liquid arsenic is a semiconductor with a band gap of 0.15 eV.<ref>[[#Bailar1973|Bailar & Trotman-Dickenson 1973, p. 558]]; [[#Li1990|Li 1990]]</ref>
The chemistry of arsenic is predominately nonmetallic.<ref>[[#Bailar1965|Bailar, Moeller & Kleinberg 1965, p. 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. 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. 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. 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. 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. 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. 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. 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. 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. 267]]</ref> The halides, and the oxides and their derivatives are illustrative examples.<ref name="Bailar513">[[#Bailar1965|Bailar, Moeller & Kleinberg 1965, p. 513]]</ref> In the trivalent state, arsenic shows some incipient metallic properties.<ref>[[#Timm1944|Timm 1944, p. 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. 641]]; [[#Kleinberg1960|Kleinberg, Argersinger & Griswold 1960, p. 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. 163]]; [[#Moeller1954|Moeller 1954, p. 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. 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. 197]]; [[#Emeléus1959|Emeléus & Sharpe 1959, p. 418]]; [[#Addison1972|Addison & Sowerby 1972, p. 209]]; [[#Mellor1964|Mellor 1964, p. 337]]</ref> The oxide As<sub>2</sub>O<sub>3</sub> is polymeric,<ref name=Pudd59/> amphoteric,<ref>[[#Pourbaix1974|Pourbaix 1974, p. 521]]; [[#Eagleson1994|Eagleson 1994, p. 92]]; [[#Greenwood2002|Greenwood & Earnshaw 2002, p. 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. 750, 975]]; [[#Silberberg2006|Silberberg 2006, p. 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. 784]]; [[#Moody1991|Moody 1991, pp. 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. 553]]</ref> It is not affected by molten alkali.<ref>[[#Dunstan1968|Dunstan 1968, p. 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. 112]]</ref> It melts at 631 °C. Antimony is a semimetal with an electrical conductivity of around 3.1 × 10<sup>4</sup> S•cm<sup>−1</sup><ref>[[#Carapella1968a|Carapella 1968a, p. 23]]</ref> and a band overlap of 0.16 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 eV.<ref>[[#Moss1952|Moss 1952, pp. 174, 179]]</ref>|group=n}} Liquid antimony is a metallic conductor with an electrical conductivity of around 5.3 × 10<sup>4</sup> S•cm<sup>−1</sup>.<ref>[[#Dupree1982|Dupree, Kirby & Freyland 1982, p. 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. 62]]</ref> Antimony has some definite cationic chemistry,<ref>[[#Cotton1999|Cotton et al. 1999, p. 396]]</ref> SbO<sup>+</sup> and Sb(OH)<sub>2</sub><sup>+</sup> being present in acidic aqueous solution;<ref>[[#King1994|King 1994, p. 174]]</ref>{{refn|1=Lidin<ref>[[#Lidin|Lidin 1996, p. 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. 87]]</ref> iron, [[nickel]], copper, zinc, tin, lead, and bismuth.<ref>[[#Fesquet1872|Fesquet 1872, pp. 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. 553]]; [[#Massey2000|Massey 2000, p. 269]]</ref> Tripositive antimony is amphoteric; [[penta-|pentapositive]] antimony is (predominately) acidic.<ref>[[#King1994|King 1994, p. 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. 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. 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. 764]]</ref> and a glass former.<ref name=Rao22/> Antimony has an extensive organometallic chemistry (see [[Organoantimony chemistry]]).<ref>[[#House2008|House 2008, p. 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. 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. 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 °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. 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. 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. 28]]</ref>
At higher temperatures tellurium is sufficiently plastic to extrude.<ref>[[#Kudryavtsev1974|Kudryavtsev 1974, p. 77]]</ref> It melts at 449.51 °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. 178]]; [[#Donohue1982|Donohue 1982, pp. 386–87]]; [[#Cotton1999|Cotton et al. 1999, p. 501]]</ref> Tellurium is a semiconductor with an electrical conductivity of around 1.0 S•cm<sup>−1</sup><ref>[[#Becker1971|Becker, Johnson & Nussbaum 1971, p. 56]]</ref> and a band gap of 0.32 to 0.38 eV.<ref name=Berger90>[[#Berger1997|Berger 1997, p. 90]]</ref> Liquid tellurium is a semiconductor, with an electrical conductivity, on melting, of around 1.9 × 10<sup>3</sup> S•cm<sup>−1</sup>.<ref name=Berger90/> [[Superheated]] liquid tellurium is a metallic conductor.<ref>[[#Chizhikov1970|Chizhikov & Shchastlivyi 1970, p. 16]]</ref>
Most of the chemistry of tellurium is characteristic of a nonmetal.<ref>[[#Jolly1966|Jolly 1966, pp. 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. 239]]</ref>{{refn|1=Cotton et al.<ref>[[#Cotton1999|Cotton et al. 1999, p. 498]]</ref> note that TeO<sub>2</sub> appears to have an ionic lattice; Wells<ref>[[#Wells1984|Wells 1984, p. 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. 588]]</ref> It can form alloys with aluminium, [[silver]], and tin.<ref>[[#Mellor1964a|Mellor 1964a, p. 30]]; [[#Wiberg2001|Wiberg 2001, p. 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. 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. 134–51]]; [[#Greenwood2002|Greenwood & Earnshaw 2002, p. 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. 1–2]]</ref>
==Elements less commonly recognised as metalloids==
===Carbon===
{{Main|Carbon}}
[[File:Graphite2.jpg|thumb|right|[[Carbon]] (as [[graphite]]). [[Delocalized electron|Delocalized valence electrons]] within the layers of graphite give it a metallic appearance.<ref>[[#Hill2000|Hill & Holman 2000, p. 124]]</ref>|alt=A shiny grey-black cuboid nugget with a rough surface.]]
Carbon is ordinarily classified as a nonmetal<ref>[[#Chang2002|Chang 2002, p. 314]]</ref> but has some metallic properties and is occasionally classified as a metalloid.<ref>[[#Kent1950|Kent 1950, pp. 1–2]]; [[#Clark1960|Clark 1960, p. 588]]; [[#Warren1981|Warren & Geballe 1981]]</ref> [[Graphite|Hexagonal graphitic carbon]] (graphite) is the most thermodynamically stable [[allotrope]] of carbon under ambient conditions.<ref>[[#Housecroft2008|Housecroft & Sharpe 2008, p. 384]]; [[#IUPAC2006|IUPAC 2006–, rhombohedral graphite entry]]</ref> It has a lustrous appearance<ref>[[#Mingos1998|Mingos 1998, p. 171]]</ref> and is a fairly good electrical conductor.<ref>[[#Wiberg2001|Wiberg 2001, p. 781]]</ref> Graphite has a layered structure. Each layer consists of carbon atoms bonded to three other carbon atoms in a [[hexagonal lattice]] arrangement. The layers are stacked together and held loosely by [[van der Waals force]]s and [[delocalized electron|delocalized valence electrons]].<ref>[[#Charlier|Charlier, Gonze & Michenaud 1994]]</ref>
Like a metal, the conductivity of graphite in the direction of its planes decreases as the temperature is raised;<ref name="Atkins320">[[#Atkins2006|Atkins et al. 2006, pp. 320–21]]</ref>{{refn|1=Liquid carbon may<ref>[[#Savvatimskiy2005|Savvatimskiy 2005, p. 1138]]</ref> or may not<ref>[[#Togaya2000|Togaya 2000]]</ref> be a metallic conductor, depending on pressure and temperature; see also.<ref>[[#Savvatimskiy2009|Savvatimskiy 2009]]</ref>|group=n}} it has the electronic band structure of a semimetal.<ref name=Atkins320/> The allotropes of carbon, including graphite, can accept foreign atoms or compounds into their structures via substitution, [[intercalation (chemistry)|intercalation]], or [[dopant|doping]]. The resulting materials are sometimes referred to as "carbon alloys".<ref>[[#Inagaki2000|Inagaki 2000, p. 216]]; [[#Yasuda2003|Yasuda et al. 2003, pp. 3–11]]</ref> Carbon can form ionic salts, including a hydrogen sulfate, perchlorate, and nitrate (C{{su|b=24|p=+}}X<sup>−</sup>.2HX, where X = HSO<sub>4</sub>, ClO<sub>4</sub>; and C{{su|b=24|p=+}}NO{{su|b=3|p=–}}.3HNO<sub>3</sub>).<ref>[[#O'Hare|O'Hare 1997, p. 230]]</ref>{{refn|1=For the sulfate, the method of preparation is (careful) direct oxidation of graphite in concentrated sulfuric acid by an [[oxidising agent]], such as [[nitric acid]], [[chromium trioxide]] or [[ammonium persulfate]]; in this instance the concentrated sulfuric acid is acting as an [[inorganic nonaqueous solvent]].|group=n}} In [[organic chemistry]], carbon can form complex cations{{snd}}termed [[carbocation|''carbocations'']]{{snd}}in which the positive charge is on the carbon atom; examples are [[carbenium ion|{{chem|CH|3|+}}]] and [[carbonium ion|{{chem|CH|5|+}}]], and their derivatives.<ref>[[#Traynham1989|Traynham 1989, pp. 930–31]]; [[#Prakash1997|Prakash & Schleyer 1997]]</ref>
Graphite is an established solid lubricant and behaves as a semiconductor in a direction perpendicular to its planes.<ref name=Atkins320/> Most of its chemistry is nonmetallic;<ref>[[#Bailar1989|Bailar et al. 1989, p. 743]]</ref> it has a relatively high ionization energy<ref>[[#Moore1985|Moore et al. 1985]]</ref> and, compared to most metals, a relatively high electronegativity.<ref>[[#House2010|House & House 2010, p. 526]]</ref> Carbon can form anions such as C<sup>4−</sup> ([[methanide]]), C{{su|b=2|p=2–}} ([[acetylide]]), and C{{su|b=4|p=3–}} ([[Sesquicarbide|sesquicarbide or allylenide]]), in compounds with metals of main groups 1–3, and with the [[lanthanide]]s and [[actinide]]s.<ref>[[#Wiberg2001|Wiberg 2001, p. 798]]</ref> Its oxide [[carbon dioxide|CO<sub>2</sub>]] forms [[carbonic acid]] H<sub>2</sub>CO<sub>3</sub>.<ref>[[#Eagleson1994|Eagleson 1994, p. 175]]</ref>{{refn|1=Only a small fraction of dissolved CO<sub>2</sub> is present in water as carbonic acid so, even though H<sub>2</sub>CO<sub>3</sub> is a medium-strong acid, solutions of carbonic acid are only weakly acidic.<ref>[[#Atkins2006|Atkins et al. 2006, p. 121]]</ref>|group=n}}
===Aluminium===
{{Main|Aluminium}}
[[File:Aluminium-4.jpg|thumb|left|High purity [[aluminium]] is much softer than its familiar [[aluminium alloys|alloys]]. People who handle it for the first time often ask if it is the real thing.<ref>[[#Russell2005|Russell & Lee 2005, pp. 358–59]]</ref>|alt=A silvery white steam-iron shaped lump with semi-circular striations along the width of its top surface and rough furrows in the middle portion of its left edge.]]
Aluminium is ordinarily classified as a metal.<ref>[[#Keevil|Keevil 1989, p. 103]]</ref> It is lustrous, malleable and ductile, and has high electrical and thermal conductivity. Like most metals it has a [[close-packed]] crystalline structure,<ref>[[#Russell2005|Russell & Lee 2005, pp. 358–60 et seq]]</ref> and forms a cation in aqueous solution.<ref>[[#Harding|Harding, Janes & Johnson 2002, p. 118]]</ref>
It has some properties that are unusual for a metal; taken together,<ref name="Metcalfe et al. 1974, p.539">[[#Metcalfe1974|Metcalfe, Williams & Castka 1974, p. 539]]</ref> these are sometimes used as a basis to classify aluminium as a metalloid.<ref>[[#Cobb2005|Cobb & Fetterolf 2005, p. 64]]; [[#Metcalfe1974|Metcalfe, Williams & Castka 1974, p. 539]]</ref> Its crystalline structure shows some evidence of [[Bonding in solids#Properties|directional bonding]].<ref>[[#Ogata2002|Ogata, Li & Yip 2002]]; [[#Boyer2004|Boyer et al. 2004, p. 1023]]; [[#Russell2005|Russell & Lee 2005, p. 359]]</ref> Aluminium bonds covalently in most compounds.<ref>[[#Cooper1968|Cooper 1968, p. 25]]; [[#Henderson2000|Henderson 2000, p. 5]]; [[#Silberberg2006|Silberberg 2006, p. 314]]</ref> The oxide [[aluminium oxide|Al<sub>2</sub>O<sub>3</sub>]] is amphoteric<ref>[[#Wiberg2001|Wiberg 2001, p. 1014]]</ref> and a conditional glass-former.<ref name=Rao22/> Aluminium can form anionic [[aluminate]]s,<ref name="Metcalfe et al. 1974, p.539"/> such behaviour being considered nonmetallic in character.<ref name="Hamm 1969, p.653">[[#Hamm1969|Hamm 1969, p. 653]]</ref>
Classifying aluminium as a metalloid has been disputed<ref>[[#Daub1996|Daub & Seese 1996, pp. 70, 109]]: "Aluminum is not a metalloid but a metal because it has mostly metallic properties."; [[#Denniston2004|Denniston, Topping & Caret 2004, p. 57]]: "Note that aluminum (Al) is classified as a metal, not a metalloid."; [[#Hasan2009|Hasan 2009, p. 16]]: "Aluminum does not have the characteristics of a metalloid but rather those of a metal."</ref> given its many metallic properties. It is therefore, arguably, an exception to the mnemonic that elements adjacent to the metal–nonmetal dividing line are metalloids.<ref>[[#Holt2007|Holt, Rinehart & Wilson c. 2007]]</ref>{{refn|1=A mnemonic that captures the elements commonly recognised as metalloids goes: ''Up, up-down, up-down, up ... are the metalloids!''<ref>[[#Tuthill2011|Tuthill 2011]]</ref>|group=n}}
Stott<ref>[[#Stott1956|Stott 1956, p. 100]]</ref> labels aluminium as a weak metal. It has the physical properties of a metal but some of the chemical properties of a nonmetal. Steele<ref>[[#Steele1966|Steele 1966, p. 60]]</ref> notes the paradoxical chemical behaviour of aluminium: "It resembles a weak metal in its amphoteric oxide and in the covalent character of many of its compounds ... Yet it is a highly [[electronegativity#Electropositivity|electropositive]] metal ... [with] a [[table of standard electrode potentials|high negative]] electrode potential". Moody<ref>[[#Moody|Moody 1991, p. 303]]</ref> says that, "aluminium is on the 'diagonal borderland' between metals and non-metals in the chemical sense."
===Selenium===
{{Main|Selenium}}
[[File:Selenium black (cropped).jpg|thumb|right|upright|Grey [[selenium]], being a [[photoconductor]], conducts electricity around 1,000 times better when light falls on it, a property used since the mid-1870s in various light-sensing applications<ref>[[#Emsley2001|Emsley 2001, p. 382]]</ref>|alt=A small glass jar filled with small dull grey concave buttons. The pieces of selenium look like tiny mushrooms without their stems.]]
Selenium shows borderline metalloid or nonmetal behaviour.<ref>[[#Young2010|Young et al. 2010, p. 9]]; [[#Craig2003|Craig & Maher 2003, p. 391]]. Selenium is "near metalloidal".</ref>{{refn|1=[[Eugene G. Rochow|Rochow]],<ref>[[#Rochow1957|Rochow 1957]]</ref> who later wrote his 1966 monograph ''The metalloids'',<ref>[[#Rochow1966|Rochow 1966, p. 224]]</ref> commented that, "In some respects selenium acts like a metalloid and tellurium certainly does".|group=n}}
Its most stable form, the grey [[trigonal crystal system|trigonal]] allotrope, is sometimes called "metallic" selenium because its electrical conductivity is several orders of magnitude greater than that of the red [[monoclinic crystal system|monoclinic]] form.<ref>[[#Moss1952|Moss 1952, p. 192]]</ref> The metallic character of selenium is further shown by its lustre,<ref name="Glinka 1965, p.356">[[#Glinka1965|Glinka 1965, p. 356]]</ref> and its crystalline structure, which is thought to include weakly "metallic" interchain bonding.<ref>[[#Evans1966|Evans 1966, pp. 124–25]]</ref> Selenium can be drawn into thin threads when molten and viscous.<ref>[[#Regnault1853|Regnault 1853, p. 208]]</ref> It shows reluctance to acquire "the high positive oxidation numbers characteristic of nonmetals".<ref>[[#Scott1962|Scott & Kanda 1962, p. 311]]</ref> It can form cyclic polycations (such as Se{{su|b=8|p=2+}}) when dissolved in [[oleum]]s<ref>[[#Cotton1999|Cotton et al. 1999, pp. 496, 503–04]]</ref> (an attribute it shares with sulfur and tellurium), and a hydrolysed cationic salt in the form of trihydroxoselenium(IV) perchlorate <span style="white-space: nowrap">[Se(OH)<sub>3</sub>]<sup>+</sup>·ClO{{su|b=4|p=–}}.<ref>[[#Arlman1939|Arlman 1939]]; [[#Bagnall1966|Bagnall 1966, pp. 135, 142–43]]</ref></span>
The nonmetallic character of selenium is shown by its brittleness<ref name="Glinka 1965, p.356"/> and the low electrical conductivity (~10<sup>−9</sup> to 10<sup>−12</sup> S•cm<sup>−1</sup>) of its highly purified form.<ref name="Kozyrev">[[#Kozyrev1959|Kozyrev 1959, p. 104]]; [[#Chizhikov1968|Chizhikov & Shchastlivyi 1968, p. 25]]; [[#Glazov1969|Glazov, Chizhevskaya & Glagoleva 1969, p. 86]]</ref> This is comparable to or less than that of [[bromine]] (7.95{{e|–12}} S•cm<sup>−1</sup>),<ref>[[#Chao1964|Chao & Stenger 1964]]</ref> a nonmetal. Selenium has the electronic band structure of a [[semiconductor]]<ref name="Berger 1997, pp.86–7">[[#Berger1997|Berger 1997, pp. 86–87]]</ref> and retains its semiconducting properties in liquid form.<ref name="Berger 1997, pp.86–7"/> It has a relatively high<ref>[[#Snyder1966|Snyder 1966, p. 242]]</ref> electronegativity (2.55 revised Pauling scale). Its reaction chemistry is mainly that of its nonmetallic anionic forms Se<sup>2−</sup>, SeO{{su|b=3|p=2−}} and SeO{{su|b=4|p=2−}}.<ref>[[#Fritz2008|Fritz & Gjerde 2008, p. 235]]</ref>
Selenium is commonly described as a metalloid in the [[environmental chemistry]] literature.<ref>[[#Meyer2005|Meyer et al. 2005, p. 284]]; [[#Manahan|Manahan 2001, p. 911]]; [[#Szpunar|Szpunar et al. 2004, p. 17]]</ref> It moves through the aquatic environment similarly to arsenic and antimony;<ref>[[#USEPA1988|US Environmental Protection Agency 1988, p. 1]]; [[#Uden|Uden 2005, pp. 347‒48]]</ref> its water-soluble salts, in higher concentrations, have a similar [[toxicology|toxicological profile]] to that of arsenic.<ref>[[#DeZuane|De Zuane 1997, p. 93]]; [[#Dev|Dev 2008, pp. 2‒3]]</ref>
===Polonium===
{{Main|Polonium}}
Polonium is "distinctly metallic" in some ways.<ref name="Cotton FA 1999, p.502">[[#Cotton1999|Cotton et al. 1999, p. 502]]</ref> Both of its allotropic forms are metallic conductors.<ref name="Cotton FA 1999, p.502"/> It is soluble in acids, forming the rose-coloured Po<sup>2+</sup> cation and displacing hydrogen: Po + 2 H<sup>+</sup> → Po<sup>2+</sup> + H<sub>2</sub>.<ref>[[#Wiberg2001|Wiberg 2001, p. 594]]</ref> Many polonium salts are known.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p. 786]]; [[#Schwietzer2010|Schwietzer & Pesterfield 2010, pp. 242–43]]</ref> The oxide [[polonium dioxide|PoO<sub>2</sub>]] is predominantly basic in nature.<ref name=Bagnall1966p41>[[#Bagnall1966|Bagnall 1966, p. 41]]; [[#Nickless1968|Nickless 1968, p. 79]]</ref> Polonium is a reluctant oxidizing agent, unlike its lightest congener oxygen: highly [[reducing agent|reducing conditions]] are required for the formation of the Po<sup>2−</sup> anion in aqueous solution.<ref>[[#Bagnall1990|Bagnall 1990, pp. 313–14]]; [[#Lehto2011|Lehto & Hou 2011, p. 220]]; [[#Siekierski2002|Siekierski & Burgess 2002, p. 117]]: "The tendency to form X<sup>2−</sup> anions decreases down the Group [16 elements] ..."</ref>
Whether polonium is ductile or brittle is unclear. It is predicted to be ductile based on its calculated [[Young's modulus#Relation among elastic constants|elastic constants]].<ref>[[#Legit|Legit, Friák & Šob 2010, pp. 214118–18]]</ref> It has a simple [[cubic crystal system|cubic crystalline structure]]. Such a structure has few [[Slip (materials science)#slip systems|slip systems]] and "leads to very low ductility and hence low fracture resistance".<ref>[[#Halford|Manson & Halford 2006, pp. 378, 410]]</ref>
Polonium shows nonmetallic character in its halides, and by the existence of [[polonide]]s. The halides have properties generally characteristic of nonmetal halides (being volatile, easily hydrolyzed, and soluble in [[organic solvent]]s).<ref>[[#Bagnall1957|Bagnall 1957, p. 62]]; [[#Fernelius1982|Fernelius 1982, p. 741]]</ref> Many metal polonides, obtained by heating the elements together at 500–1,000 °C, and containing the Po<sup>2−</sup> anion, are also known.<ref>[[#Bagnall1966|Bagnall 1966, p. 41]]; [[#Barrett2003|Barrett 2003, p. 119]]</ref>
===Astatine===
{{Main|Astatine}}
As a [[halogen]], astatine tends to be classified as a nonmetal.<ref>[[#Hawkes2010|Hawkes 2010]]; [[#Holt2007|Holt, Rinehart & Wilson c. 2007]]; [[#Hawkes1999|Hawkes 1999, p. 14]]; [[#Roza2009|Roza 2009, p. 12]]</ref> It has some marked metallic properties<ref>[[#Keller1985|Keller 1985]]</ref> and is sometimes instead classified as either a metalloid<ref>[[#Harding2002|Harding, Johnson & Janes 2002, p. 61]]</ref> or (less often) as a metal.{{refn|1=A further option is to include astatine both as a nonmetal and as a metalloid.<ref>[[#Long1986|Long & Hentz 1986, p. 58]]</ref>|group=n}} Immediately following its production in 1940, early investigators considered it a metal.<ref>[[#Vasáros1985|Vasáros & Berei 1985, p. 109]]</ref> In 1949 it was called the most noble (difficult to [[redox|reduce]]) nonmetal as well as being a relatively noble (difficult to oxidize) metal.<ref>[[#Haissinsky1949|Haissinsky & Coche 1949, p. 400]]</ref> In 1950 astatine was described as a halogen and (therefore) a [[reactivity (chemistry)|reactive]] nonmetal.<ref>[[#Brownlee1950|Brownlee et al. 1950, p. 173]]</ref> In 2013, on the basis of [[relativistic quantum chemistry|relativistic]] modelling, astatine was predicted to be a monatomic metal, with a [[Face-centred cubic|face-centred cubic crystalline structure]].<ref>[[#Hermann|Hermann, Hoffmann & Ashcroft 2013]]</ref>
Several authors have commented on the metallic nature of some of the properties of astatine. Since iodine is a semiconductor in the direction of its planes, and since the halogens become more metallic with increasing atomic number, it has been presumed that astatine would be a metal if it could form a condensed phase.<ref>[[#Siekierski2002|Siekierski & Burgess 2002, pp. 65, 122]]</ref>{{refn|1=A visible piece of astatine would be immediately and completely vaporized because of the heat generated by its intense radioactivity.<ref>[[#Emsley2001|Emsley 2001, p. 48]]</ref>|group=n}} Astatine may be metallic in the liquid state on the basis that elements with an [[enthalpy of vaporization]] (∆H<sub>vap</sub>) greater than ~42 kJ/mol are metallic when liquid.<ref name="Rao & Ganguly 1986">[[#Rao1986|Rao & Ganguly 1986]]</ref> Such elements include boron,{{refn|1=The literature is contradictory as to whether boron exhibits metallic conductivity in liquid form. Krishnan et al.<ref>[[#Krishnan1998|Krishnan et al. 1998]]</ref> found that liquid boron behaved like a metal. Glorieux et al.<ref>[[#Glorieux2001|Glorieux, Saboungi & Enderby 2001]]</ref> characterised liquid boron as a semiconductor, on the basis of its low electrical conductivity. Millot et al.<ref>[[#Millot2002|Millot et al. 2002]]</ref> reported that the emissivity of liquid boron was not consistent with that of a liquid metal.|group=n}} silicon, germanium, antimony, selenium, and tellurium. Estimated values for ∆H<sub>vap</sub> of [[Diatomic molecule|diatomic]] astatine are 50 kJ/mol or higher;<ref>[[#Vasáros1985|Vasáros & Berei 1985, p. 117]]</ref> diatomic iodine, with a ∆H<sub>vap</sub> of 41.71,<ref>[[#Kaye1973|Kaye & Laby 1973, p. 228]]</ref> falls just short of the threshold figure.
"Like typical metals, it [astatine] is precipitated by [[hydrogen sulfide]] even from strongly acid solutions and is displaced in a free form from sulfate solutions; it is deposited on the [[cathode]] on [[electrolysis]]."<ref>[[#Samsonov1968|Samsonov 1968, p. 590]]</ref>{{refn|1=Korenman<ref>[[#Korenman1959|Korenman 1959, p. 1368]]</ref> similarly noted that "the ability to precipitate with hydrogen sulfide distinguishes astatine from other halogens and brings it closer to bismuth and other [[heavy metal (chemistry)|heavy metals]]".|group=n}} Further indications of a tendency for astatine to behave like a [[heavy metal (chemistry)|(heavy) metal]] are: "... the formation of [[pseudohalide]] compounds ... complexes of astatine cations ... complex anions of trivalent astatine ... as well as complexes with a variety of organic solvents".<ref>[[#Rossler1985|Rossler 1985, pp. 143–44]]</ref> It has also been argued that astatine demonstrates cationic behaviour, by way of stable At<sup>+</sup> and AtO<sup>+</sup> forms, in strongly acidic aqueous solutions.<ref>[[#Champion2010|Champion et al. 2010]]</ref>
Some of astatine's reported properties are nonmetallic. It has been extrapolated to have the narrow liquid range ordinarily associated with nonmetals (mp 302 °C; bp 337 °C),<ref>[[#Borst1982|Borst 1982, pp. 465, 473]]</ref> although experimental indications suggest a lower boiling point of about 230±3 °C. Batsanov gives a calculated band gap energy for astatine of 0.7 eV;<ref>[[#Batsanov1971|Batsanov 1971, p. 811]]</ref> this is consistent with nonmetals (in physics) having separated [[valence band|valence]] and [[conduction band]]s and thereby being either semiconductors or insulators.<ref>[[#Swalin1962|Swalin 1962, p. 216]]; [[#Feng2005|Feng & Lin 2005, p. 157]]</ref> The chemistry of astatine in aqueous solution is mainly characterised by the formation of various anionic species.<ref>[[#Schwietzer2010|Schwietzer & Pesterfield 2010, pp. 258–60]]</ref> Most of its known compounds resemble those of iodine,<ref>[[#Hawkes1999|Hawkes 1999, p. 14]]</ref> which is a halogen and a nonmetal.<ref>[[#Olmsted1997|Olmsted & Williams 1997, p. 328]]; [[#Daintith2004|Daintith 2004, p. 277]]</ref> Such compounds include astatides (XAt), astatates (XAtO<sub>3</sub>), and [[monovalent ion|monovalent]] [[interhalogen compound]]s.<ref>[[#Eberle1985|Eberle1985, pp. 213–16, 222–27]]</ref>
Restrepo et al.<ref>[[#Restrepo2004|Restrepo et al. 2004, p. 69]]; [[#Restrepo2006|Restrepo et al. 2006, p. 411]]</ref> reported that astatine appeared to be more polonium-like than halogen-like. They did so on the basis of detailed comparative studies of the known and interpolated properties of 72 elements.
<span id="SCC"></span>
==Related concepts==
===Near metalloids===
[[File:Iodinecrystals.JPG|thumb|right|[[Iodine]] crystals, showing a metallic [[lustre (mineralogy)|lustre]]. Iodine is a [[semiconductor]] in the direction of its planes, with a [[band gap]] of ~1.3 eV. It has an [[electrical conductivity]] of 1.7 × 10<sup>−8</sup> S•cm<sup>−1</sup> at [[room temperature]].<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p. 804]]</ref> This is higher than selenium but lower than boron, the least electrically conducting of the recognised metalloids.{{refn|1=The separation between molecules in the layers of iodine (350 pm) is much less than the separation between iodine layers (427 pm; cf. twice the van der Waals radius of 430 pm).<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, p. 803]]</ref> This is thought to be caused by electronic interactions between the molecules in each layer of iodine, which in turn give rise to its semiconducting properties and shiny appearance.<ref>[[#Wiberg2001|Wiberg 2001, p. 416]]</ref>|group=n}}|alt=Shiny violet-black coloured crystalline shards.]]
In the periodic table, some of the elements adjacent to the commonly recognised metalloids, although usually classified as either metals or nonmetals, are occasionally referred to as ''near-metalloids''<ref>[[#Craig2003|Craig & Maher 2003, p. 391]]; [[#Schroers2013|Schroers 2013, p. 32]]; [[#Vernon|Vernon 2013, pp. 1704–05]]</ref> or noted for their metalloidal character. To the left of the metal–nonmetal dividing line, such elements include gallium,<ref>[[#Cotton1999|Cotton et al. 1999, p. 42]]</ref> tin<ref>[[#Marezio|Marezio & Licci 2000, p. 11]]</ref> and bismuth.<ref name=Vernon/> They show unusual packing structures,<ref>[[#Russell2005|Russell & Lee 2005, p. 5]]</ref> marked covalent chemistry (molecular or polymeric),<ref>[[#Parish1977|Parish 1977, pp. 178, 192–93]]</ref> and amphoterism.<ref>[[#Eggins1972|Eggins 1972, p. 66]]; [[#Rayner2006|Rayner-Canham & Overton 2006, pp. 29–30]]</ref> To the right of the dividing line are carbon,<ref>[[#Atkins2006|Atkins et al. 2006, pp. 320–21]]; [[#Bailar1989|Bailar et al. 1989, pp. 742–43]]</ref> phosphorus,<ref>[[#Rochow1966|Rochow 1966, p. 7]]; [[#Taniguchi1984|Taniguchi et al. 1984, p. 867]]: "... black phosphorus ... [is] characterized by the wide valence bands with rather delocalized nature."; [[#Morita1986|Morita 1986, p. 230]]; [[#Carmalt|Carmalt & Norman 1998, p. 7]]: "Phosphorus ... should therefore be expected to have some metalloid properties."; [[#Du2010|Du et al. 2010]]. Interlayer interactions in black phosphorus, which are attributed to van der Waals-Keesom forces, are thought to contribute to the smaller band gap of the bulk material (calculated 0.19 eV; observed 0.3 eV) as opposed to the larger band gap of a single layer (calculated ~0.75 eV).</ref> selenium<ref>[[#Stuke1974|Stuke 1974, p. 178]]; [[#Cotton1999|Cotton et al. 1999, p. 501]]; [[#Craig2003|Craig & Maher 2003, p. 391]]</ref> and iodine.<ref>[[#Steudel1977|Steudel 1977, p. 240]]: "... considerable orbital overlap must exist, to form intermolecular, many-center ... [sigma] bonds, spread through the layer and populated with delocalized electrons, reflected in the properties of iodine (lustre, color, moderate electrical conductivity)."; [[#Segal1989|Segal 1989, p. 481]]: "Iodine exhibits some metallic properties ..."</ref> They exhibit metallic lustre, semiconducting properties{{refn|1=For example: intermediate electrical conductivity;<ref name="Lutz 2011, p. 17">[[#Lutz2011|Lutz et al. 2011, p. 17]]</ref> a relatively narrow band gap;<ref>[[#Yacobi1990|Yacobi & Holt 1990, p. 10]]; [[#Wiberg2001|Wiberg 2001, p. 160]]</ref> light sensitivity.<ref name="Lutz 2011, p. 17"/>|group=n}} and bonding or valence bands with delocalized character. This applies to their most thermodynamically stable forms under ambient conditions: carbon as graphite; phosphorus as black phosphorus;{{refn|1=White phosphorus is the least stable and most reactive form.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, pp. 479, 482]]</ref> It is also the most common, industrially important,<ref>[[#Eagleson1994|Eagleson 1994, p. 820]]</ref> and easily reproducible allotrope, and for these three reasons is regarded as the standard state of the element.<ref>[[#Oxtoby2008|Oxtoby, Gillis & Campion 2008, p. 508]]</ref>|group=n}} and selenium as grey selenium.
===Allotropes===
[[File:Sn-Alpha-Beta.jpg|thumb|left|[[White tin]] (left) and [[grey tin]] (right). Both forms have a metallic appearance.|alt=Many small, shiny, silver-coloured spheres on the left; many of the same sized spheres on the right are duller and darker than the ones of the left and have a subdued metallic shininess.]]
Different crystalline forms of an element are called [[allotropy|allotropes]]. Some allotropes, particularly those of elements located (in periodic table terms) alongside or near the notional dividing line between metals and nonmetals, exhibit more pronounced metallic, metalloidal or nonmetallic behaviour than others.<ref>[[#Brescia1980|Brescia et al. 1980, pp. 166–71]]</ref> The existence of such allotropes can complicate the classification of the elements involved.<ref>[[#Fine|Fine & Beall 1990, p. 578]]</ref>
Tin, for example, has two allotropes: [[tetragonal crystal system|tetragonal]] "white" β-tin and cubic "grey" α-tin. White tin is a very shiny, ductile and malleable metal. It is the stable form at or above room temperature and has an electrical conductivity of 9.17 × 10<sup>4</sup> S·cm<sup>−1</sup> (~1/6th that of copper).<ref>[[#Wiberg2001|Wiberg 2001, p. 901]]</ref> Grey tin usually has the appearance of a grey micro-crystalline powder, and can also be prepared in brittle semi-lustrous crystalline or [[polycrystalline]] forms. It is the stable form below 13.2 °C and has an electrical conductivity of between (2–5) × 10<sup>2</sup> S·cm<sup>−1</sup> (~1/250th that of white tin).<ref>[[#Berger1997|Berger 1997, p. 80]]</ref> Grey tin has the same crystalline structure as that of diamond. It behaves as a semiconductor (as if it had a band gap of 0.08 eV), but has the electronic band structure of a semimetal.<ref>[[#Lovett1977|Lovett 1977, p. 101]]</ref> It has been referred to as either a very poor metal,<ref>[[#Cohen|Cohen & Chelikowsky 1988, p. 99]]</ref> a metalloid,<ref>[[#Taguena|Taguena-Martinez, Barrio & Chambouleyron 1991, p. 141]]</ref> a nonmetal<ref>[[#Ebbing|Ebbing & Gammon 2010, p. 891]]</ref> or a near metalloid.<ref name=Vernon>[[#Vernon|Vernon 2013, p. 1705]]</ref>
The diamond allotrope of carbon is clearly nonmetallic, being translucent and having a low electrical conductivity of 10<sup>−14</sup> to 10<sup>−16</sup> S·cm<sup>−1</sup>.<ref>[[#Asmussen|Asmussen & Reinhard 2002, p. 7]]</ref> Graphite has an electrical conductivity of 3 × 10<sup>4</sup> S·cm<sup>−1</sup>,<ref>[[#Deprez1988|Deprez & McLachan 1988]]</ref> a figure more characteristic of a metal. Phosphorus, sulfur, arsenic, selenium, antimony, and bismuth also have less stable allotropes that display different behaviours.<ref>[[#Addison1964|Addison 1964 (P, Se, Sn)]]; [[#Marko1998|Marković, Christiansen & Goldman 1998 (Bi)]]; [[#Nagao2004|Nagao et al. 2004]]</ref>
==Abundance, extraction, and cost==
{| class="wikitable sortable" style="float:right"
|-
! [[atomic number|Z]] !! Element !! Grams<br/>/tonne
|-
| align="right" |8|| Oxygen|| align="right" | 461,000
|-
| align="right" |14|| '''Silicon'''|| align="right" | 282,000
|-
| align="right" |13|| ''Aluminium''|| align="right" | 82,300
|-
| align="right" |26|| Iron|| align="right" | 56,300
|-
| align="right" |6|| ''Carbon''|| align="right" | 200
|-
| align="right" |29|| Copper|| align="right" | 60
|-
| align="right" |5|| '''Boron''' || align="right" | 10
|-
| align="right" |33|| '''Arsenic'''|| align="right" | 1.8
|-
| align="right" |32|| '''Germanium'''|| align="right" | 1.5
|-
| align="right" |47|| Silver|| align="right" | 0.075
|-
| align="right" |34|| ''Selenium''|| align="right" | 0.05
|-
| align="right" |51|| '''Antimony'''|| align="right" | 0.02
|-
| align="right" |79|| Gold|| align="right" | 0.004
|-
| align="right" |52|| '''Tellurium'''|| align="right" | 0.001
|-
| align="right" |75|| Rhenium|| align="right" | <span style="display:none;">0.0000000007</span>7×10<sup>−10</sup>
|-
| align="right" |54|| Xenon|| align="right" | <span style="display:none;">0.00000000003</span>3×10<sup>−11</sup>
|-
| align="right" |84|| ''Polonium''|| align="right" | <span style="display:none;">0.0000000000000002</span>2×10<sup>−16</sup>
|-
| align="right" |85|| ''Astatine''|| align="right" | <span style="display:none;">0.000000000000000003</span>3×10<sup>−20</sup>
|}
===Abundance===
The table gives [[crustal abundance]]s of the elements commonly to rarely recognised as metalloids.<ref>[[#Lide|Lide 2005]]; [[#Wiberg2001|Wiberg 2001, p. 423: At]]</ref> Some other elements are included for comparison: oxygen and xenon (the most and least abundant elements with stable isotopes); iron and the coinage metals copper, silver, and gold; and rhenium, the least abundant stable metal (aluminium is normally the most abundant metal). Various abundance estimates have been published; these often disagree to some extent.<ref>[[#Cox|Cox 1997, pp. 182‒86]]</ref>
===Extraction===
The recognised metalloids can be obtained by [[redox|chemical reduction]] of either their oxides or their [[sulfide]]s. Simpler or more complex extraction methods may be employed depending on the starting form and economic factors.<ref>[[#MacKay2002|MacKay, MacKay & Henderson 2002, p. 204]]</ref> Boron is routinely obtained by reducing the trioxide with magnesium: B<sub>2</sub>O<sub>3</sub> + 3 Mg → 2 B + 3MgO; after secondary processing the resulting brown powder has a purity of up to 97%.<ref>[[#Baudis|Baudis 2012, pp. 207–08]]</ref> Boron of higher purity (> 99%) is prepared by heating volatile boron compounds, such as BCl<sub>3</sub> or BBr<sub>3</sub>, either in a hydrogen atmosphere (2 BX<sub>3</sub> + 3 H<sub>2</sub> → 2 B + 6 HX) or to the point of [[thermal decomposition]]. Silicon and germanium are obtained from their oxides by heating the oxide with carbon or hydrogen: SiO<sub>2</sub> + C → Si + CO<sub>2</sub>; GeO<sub>2</sub> + 2 H<sub>2</sub> → Ge + 2 H<sub>2</sub>O. Arsenic is isolated from its pyrite (FeAsS) or arsenical pyrite (FeAs<sub>2</sub>) by heating; alternatively, it can be obtained from its oxide by reduction with carbon: 2 As<sub>2</sub>O<sub>3</sub> + 3 C → 2 As + 3 CO<sub>2</sub>.<ref>[[#Wiberg2001|Wiberg 2001, p. 741]]</ref> Antimony is derived from its sulfide by reduction with iron: Sb<sub>2</sub>S<sub>3</sub> → 2 Sb + 3 FeS. Tellurium is prepared from its oxide by dissolving it in aqueous NaOH, yielding tellurite, then by [[electrolytic reduction]]: TeO<sub>2</sub> + 2 NaOH → Na<sub>2</sub>TeO<sub>3</sub> + H<sub>2</sub>O;<ref>[[#Chizhikov1968|Chizhikov & Shchastlivyi 1968, p. 96]]</ref> Na<sub>2</sub>TeO<sub>3</sub> + H<sub>2</sub>O → Te + 2 NaOH + O<sub>2</sub>.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, pp. 140–41, 330, 369, 548–59, 749: B, Si, Ge, As, Sb, Te]]</ref> Another option is reduction of the oxide by roasting with carbon: TeO<sub>2</sub> + C → Te + CO<sub>2</sub>.<ref>[[#Kudryavtsev1974|Kudryavtsev 1974, p. 158]]</ref>
Production methods for the elements less frequently recognised as metalloids involve natural processing, electrolytic or chemical reduction, or irradiation. Carbon (as graphite) occurs naturally and is extracted by crushing the parent rock and floating the lighter graphite to the surface. Aluminium is extracted by dissolving its oxide Al<sub>2</sub>O<sub>3</sub> in molten [[cryolite]] Na<sub>3</sub>AlF<sub>6</sub> and then by high temperature electrolytic reduction. Selenium is produced by roasting the coinage metal selenides X<sub>2</sub>Se (X = Cu, Ag, Au) with [[soda ash]] to give the selenite: X<sub>2</sub>Se + O<sub>2</sub> + Na<sub>2</sub>CO<sub>3</sub> → Na<sub>2</sub>SeO<sub>3</sub> + 2 X + CO<sub>2</sub>; the selenide is neutralized by sulfuric acid H<sub>2</sub>SO<sub>4</sub> to give [[selenous acid]] H<sub>2</sub>SeO<sub>3</sub>; this is reduced by bubbling with [[sulfur dioxide|SO<sub>2</sub>]] to yield elemental selenium. Polonium and astatine are produced in minute quantities by irradiating bismuth.<ref>[[#Greenwood2002|Greenwood & Earnshaw 2002, pp. 271, 219, 748–49, 886: C, Al, Se, Po, At]]; [[#Wiberg2001|Wiberg 2001, p. 573: Se]]</ref>
===Cost===
The recognised metalloids and their closer neighbours mostly cost less than silver; only polonium and astatine are more expensive than gold, on account of their significant radioactivity. As of 5 April 2014, prices for small samples (up to 100 g) of silicon, antimony and tellurium, and graphite, aluminium and selenium, average around one third the cost of silver (US$1.5 per gram or about $45 an ounce). Boron, germanium, and arsenic samples average about three-and-a-half times the cost of silver.{{refn|1=Sample prices of gold, in comparison, start at roughly thirty-five times that of silver. Based on sample prices for B, C, Al, Si, Ge, As, Se, Ag, Sb, Te, and Au available on-line from [https://www.alfa.com/en/pure-elements/ Alfa Aesa]; [http://www.goodfellow.com/catalogue/GFCatL.php?ewd_token=d9JJMyYK0a7QlDSc171OUVMREOSDVO&n=FkZgkaaF0WMB5EJhaP0zaLTZIN5ixG&ewd_urlNo=GFCatalogue2&next=Period Goodfellow]; [http://www.elementsales.com/pl_element.htm Metallium]; and [http://www.unitednuclear.com/index.php?main_page=index&cPath=16_17_69 United Nuclear Scientific].|group=n}} Polonium is available for about $100 per [[microgram]].<ref>[[#United|United Nuclear 2013]]</ref> Zalutsky and Pruszynski<ref>[[#Zalutsky|Zalutsky & Pruszynski 2011, p. 181]]</ref> estimate a similar cost for producing astatine. Prices for the applicable elements traded as commodities tend to range from two to three times cheaper than the sample price (Ge), to nearly three thousand times cheaper (As).{{refn|1=Based on [[spot price]]s for Al, Si, Ge, As, Sb, Se, and Te available on-line from [http://www.fastmarkets.com/minor-metals FastMarkets: Minor Metals]; [http://www.fastmarkets.com/base-metals Fast Markets: Base Metals]; [http://pv.energytrend.com/ EnergyTrend: PV Market Status, Polysilicon]; and [http://www.metal-pages.com/metals/arsenic/metal-prices-news-information/ Metal-Pages: Arsenic metal prices, news, and information].|group=n}}
== Napomene ==
{{Reflist|group=n|colwidth=45em}}
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{{refend}}
== Spoljašnje veze ==
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[[Категорија:Металоиди| ]]
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