Metalloid: Difference between revisions
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===Biological agents=== |
===Biological agents=== |
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[[File:Arsenic trioxide.jpg|thumb|left|[[Arsenic trioxide]] or ''white arsenic,'' one of the most toxic and prevalent forms of arsenic. The 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 (with germanium and tellurium being thought to have potential). |
[[File:Arsenic trioxide.jpg|thumb|left|[[Arsenic trioxide]] or ''white arsenic,'' one of the most toxic and prevalent forms of arsenic. The 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 (with germanium and tellurium being thought to have potential). |
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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. |
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. |
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===Semiconductors and electronics=== |
===Semiconductors and electronics=== |
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[[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|compound]] semiconductor constituents (Si, Ge or [[GaAs]], for example) or as [[doping (semiconductor)|doping agents]] (B, Sb, 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.]] |
[[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|compound]] semiconductor constituents (Si, Ge or [[GaAs]], for example) or as [[doping (semiconductor)|doping agents]] (B, Sb, 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.]] |
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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 crystals are relatively hard to obtain, and introducing and retaining controlled impurities is difficult.<ref>[[#Rochow1966|Rochow 1966, p. 41]]; [[#Berger1997|Berger 1997, pp. 42–3]]</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 W. [[Wolfgang Pauli|Pauli]] and I. [[Isidor Isaac Rabi|Rabi]], to comment derogatorily on the 'Physics of Dirt' '; [[#Hoddeson2007|Hoddeson 2007, pp. 25–34 (29)]]</ref> Silicon has largely replaced germanium 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 however 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 comprises one-atom thick sheets of hydrogen-terminated germanium atoms, |
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 crystals are relatively hard to obtain, and introducing and retaining controlled impurities is difficult.<ref>[[#Rochow1966|Rochow 1966, p. 41]]; [[#Berger1997|Berger 1997, pp. 42–3]]</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 W. [[Wolfgang Pauli|Pauli]] and I. [[Isidor Isaac Rabi|Rabi]], to comment derogatorily on the 'Physics of Dirt' '; [[#Hoddeson2007|Hoddeson 2007, pp. 25–34 (29)]]</ref> Silicon has largely replaced germanium 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 however 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 comprises 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> |
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Although arsenic and antimony are not semiconductors in their [[standard state]]s, 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. These compounds are preferred for some special applications.<ref>[[#Russell2005|Russell & Lee 2005, pp. 421–2, 424]]</ref> Tellurium, which is a semiconductor in its standard state, is used mainly as a component in type [[List of semiconductor materials|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 diodes; light emitting diodes; photodetectors; and solar cells.<ref>[[#Jaskula|Jaskula 2013]]</ref> Selenium is used in the production of [[solar cell]]s<ref>[[#GES|German Energy Society 2008, p. 43–44]]</ref> and in high-energy [[surge protector]]s.<ref>[[#Patel|Patel 2012, p. 248]]</ref> |
Although arsenic and antimony are not semiconductors in their [[standard state]]s, 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. These compounds are preferred for some special applications.<ref>[[#Russell2005|Russell & Lee 2005, pp. 421–2, 424]]</ref> Tellurium, which is a semiconductor in its standard state, is used mainly as a component in type [[List of semiconductor materials|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 diodes; light emitting diodes; photodetectors; and solar cells.<ref>[[#Jaskula|Jaskula 2013]]</ref> Selenium is used in the production of [[solar cell]]s<ref>[[#GES|German Energy Society 2008, p. 43–44]]</ref> and in high-energy [[surge protector]]s.<ref>[[#Patel|Patel 2012, p. 248]]</ref> |
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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>{{#tag:ref|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]] (EoV) greater than ~42 kJ/mol are metallic when liquid.<ref name="Rao & Ganguly 1986">[[#Rao1986|Rao & Ganguly 1986]]</ref> Such elements include boron,{{#tag:ref|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 the EoV of diatomic astatine are 50 kJ/mol or higher;<ref>[[#Vasáros1985|Vasáros & Berei 1985, p. 117]]</ref> diatomic iodine, with an EoV of 41.71,<ref>[[#Kaye1973|Kaye & Laby 1973, p. 228]]</ref> falls just short of the threshold figure. '[L]ike 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>{{#tag:ref|Korenman<ref>[[#Korenman1959|Korenman 1959, p. 1368]]</ref> similarly noted that 'the ability precipitate with hydrogen sulfide distinguishes astatine from other halogens and brings it closer to bismuth and other heavy metals.'|group=n}} Further indications of a tendency for astatine to behave like a (heavy) metal are: '<span style="white-space: nowrap">...</span>the formation of [[pseudohalide]] compounds<span style="white-space: nowrap">...</span>complexes of astatine cations<span style="white-space: nowrap">...</span>complex anions of trivalent astatine<span style="white-space: nowrap">...</span>as well as complexes with a variety of organic solvents'.<ref>[[#Rossler1985|Rossler 1985, pp. 143–4]]</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> |
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>{{#tag:ref|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]] (EoV) greater than ~42 kJ/mol are metallic when liquid.<ref name="Rao & Ganguly 1986">[[#Rao1986|Rao & Ganguly 1986]]</ref> Such elements include boron,{{#tag:ref|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 the EoV of diatomic astatine are 50 kJ/mol or higher;<ref>[[#Vasáros1985|Vasáros & Berei 1985, p. 117]]</ref> diatomic iodine, with an EoV of 41.71,<ref>[[#Kaye1973|Kaye & Laby 1973, p. 228]]</ref> falls just short of the threshold figure. '[L]ike 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>{{#tag:ref|Korenman<ref>[[#Korenman1959|Korenman 1959, p. 1368]]</ref> similarly noted that 'the ability precipitate with hydrogen sulfide distinguishes astatine from other halogens and brings it closer to bismuth and other heavy metals.'|group=n}} Further indications of a tendency for astatine to behave like a (heavy) metal are: '<span style="white-space: nowrap">...</span>the formation of [[pseudohalide]] compounds<span style="white-space: nowrap">...</span>complexes of astatine cations<span style="white-space: nowrap">...</span>complex anions of trivalent astatine<span style="white-space: nowrap">...</span>as well as complexes with a variety of organic solvents'.<ref>[[#Rossler1985|Rossler 1985, pp. 143–4]]</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> |
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Some of astatine's reported properties are nonmetallic. It has the narrow liquid range ordinarily associated with nonmetals (mp 302 |
Some of astatine's reported properties are nonmetallic. It has the narrow liquid range ordinarily associated with nonmetals (mp 302 °C; bp 337 °C).<ref>[[#Borst1982|Borst 1982, pp. 465, 473]]</ref> 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 and conduction bands 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 [[interhalogen compound]]s.<ref>[[#Eberle1985|Eberle1985, pp. 213–16, 222–7]]</ref> |
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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. |
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. |
Revision as of 09:12, 25 February 2014
13 | 14 | 15 | 16 | 17 | ||
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2 | B Boron |
C Carbon |
N Nitrogen |
O Oxygen |
F Fluorine | |
3 | Al Aluminium |
Si Silicon |
P Phosphorus |
S Sulfur |
Cl Chlorine | |
4 | Ga Gallium |
Ge Germanium |
As Arsenic |
Se Selenium |
Br Bromine | |
5 | In Indium |
Sn Tin |
Sb Antimony |
Te Tellurium |
I Iodine | |
6 | Tl Thallium |
Pb Lead |
Bi Bismuth |
Po Polonium |
At Astatine | |
Commonly recognized (86–99%): B, Si, Ge, As, Sb, Te
Irregularly recognized (40–49%): Po, At
Less commonly recognized (24%): Se
Rarely recognized (8–10%): C, Al
(All other elements cited in less than 6% of sources)
Arbitrary metal-nonmetal dividing line: between Be and B, Al and Si, Ge and As, Sb and Te, Po and At
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Recognition status, as metalloids, of some elements in the p-block of the periodic table. Percentages are median appearance frequencies in the lists of metalloids.[n 1] The staircase-shaped line is a typical example of the arbitrary metal–nonmetal dividing line found on some periodic tables. |
A metalloid is a chemical element that has properties in between those of metals and nonmetals. There is no standard definition of a metalloid, nor is there agreement as to which elements are appropriately classified as such. Despite this lack of specificity the term remains in use in chemistry literature.
The six commonly recognised metalloids are boron, silicon, germanium, arsenic, antimony and tellurium. Less commonly recognised metalloids include carbon, aluminium, selenium, polonium and astatine. On a standard periodic table all of these elements can be found in a diagonal region of the p-block, with its main axis anchored by boron at one end and astatine at the other. Some periodic tables include a dividing line between metals and nonmetals and it is generally the elements adjacent to this line or, less often, one or more of the elements adjacent to those elements, which are identified as metalloids.
Metalloids usually have a metallic appearance but they are brittle and only fair conductors of electricity; chemically, they mostly behave as (weak) nonmetals. They can form alloys with metals. Most of the other physical and chemical properties of metalloids are intermediate in nature.
Metalloids are too brittle to have any structural uses. They and their compounds are used in alloys, biological agents, flame retardants, glasses, optical storage, semiconductors and electronics. The electrical properties of silicon and germanium enabled the establishment of the semiconductor industry in the 1950s and the development of solid-state electronics from the early 1960s.[1]
The term metalloid originally referred to nonmetals. 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;[2] the term semimetal has a different meaning in physics, one that more specifically refers to the electronic band structure of a substance rather than the overall classification of a chemical element.
Definition
Metalloids are usually regarded as a third category of chemical elements occupying a fuzzy 'buffer zone' between metals and nonmetals.[3][n 2] There is no universally agreed, rigorous definition of a metalloid.[8] The feasibility of establishing a specific definition has been questioned; anomalies can be found in several such attempted constructs.[9] Classifying an element as a metalloid has been described as 'arbitrary'.[10]
Generic
A metalloid is an element with properties that are in between or a mixture of those of metals and nonmetals, and that is considered to be 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.[n 3] Difficulty of categorisation is a key attribute. Although most other elements have a mixture of metallic and nonmetallic properties,[17] they can be classified according to which set of properties is more pronounced.[18][n 4] Only the elements at or near the margins, ordinarily those regarded as lacking a sufficiently clear preponderance of either metallic or nonmetallic properties, are classified as metalloids.[22]
Boron, silicon, germanium, arsenic, antimony and tellurium are commonly recognised as metalloids.[23][n 5] Depending on the author, one or more from selenium, polonium or astatine are sometimes added to the list.[25] Boron is sometimes excluded, by itself or with silicon.[26] Tellurium is sometimes not regarded as a metalloid.[27] The inclusion of antimony, polonium and astatine as metalloids has also been questioned.[28]
Other elements are occasionally classified as metalloids. These elements include[29] hydrogen,[30] beryllium,[31] nitrogen,[32] phosphorus,[33] sulfur,[34] zinc,[35] gallium,[36] tin, iodine,[37] lead,[38] bismuth[27] and radon.[39] The term metalloid has also been used for elements that exhibit metallic lustre and electrical conductivity, and that are amphoteric; such as arsenic, antimony, vanadium, chromium, molybdenum, tungsten, tin, lead and aluminium.[40] The poor metals,[41] and nonmetals (such as carbon or nitrogen) that can form alloys with metals[42] or modify their properties[43] have also occasionally been considered as metalloids.
Specific
Element | Band structure | ||||
Boron | 191 | 2.04 | semiconductor | ||
Silicon | 187 | 1.90 | semiconductor | ||
Germanium | 182 | 2.01 | semiconductor | ||
Arsenic | 225 | 2.18 | semimetal | ||
Antimony | 198 | 2.05 | semimetal | ||
Tellurium | 207 | 2.10 | semiconductor | ||
average | 198 | 2.05 | |||
The elements commonly recognised as metalloids, and their ionization energies (kcal/mol);[44] electronegativities (revised Pauling scale); and electronic band structures[45] (most thermodynamically stable forms under ambient conditions). | |||||
Metalloids tend to be characterised in terms of generalities or a few broadly indicative physical or chemical properties.[46] A single quantitative criterion, such as electronegativity is sometimes used.[n 6] In contrast, Jones[55] (writing on the role of classification in science) observed that, 'Classes are usually defined by more than two attributes.'
Masterton and Slowinski[56] offered a more specific treatment. They wrote that metalloids have ionization energies around 200 kcal/mol, and electronegativity values close to 2.0. They also said that metalloids are typically semiconductors, though antimony and arsenic (semimetals in the physics sense) have electrical conductivities that approach those of metals. Their description, using these three properties, includes the six elements commonly recognised as metalloids (see table, this section). Selenium and polonium are probably excluded from this scheme; astatine may or may not be included.[n 7]
The commonly recognised metalloids can also be quantitatively described as having packing efficiencies between 34% and 41% and Goldhammer-Herzfeld criterion metallization ratios between ~0.85 and 1.1; average 1.0.[59][n 8] The packing efficiency of boron is 38%; silicon and germanium 34; arsenic 38.5; antimony 41; and tellurium 36.4.[61] These values are lower than in most metals (80% of which have a packing efficiency of at least 68%)[62][n 9] but higher than those of elements usually classified as nonmetals. Packing efficiencies for nonmetals are: graphite 17%,[66] sulfur 19.2,[67] iodine 23.9,[67] selenium 24.2,[67] and black phosphorus 28.5.[64] The Goldhammer-Herzfeld criterion ratio values of the recognised metalloids are lower than those of representative and transition metals and higher than those of nonmetals.[n 10]
Periodic table territory
Location
Template:Periodic table (metalloid border) Metalloids lie on either side of the dividing line between metals and nonmetals. This can be found, in varying configurations, on some periodic tables. Elements to the lower left of the line generally display increasing metallic behaviour; elements to the upper right display increasing nonmetallic behaviour.[71] When presented as a regular stairstep, elements with the highest critical temperature for their groups (Li, Be, Al, Ge, Sb, Po) lie just below the line.[72]
The diagonal positioning of the metalloids represents an exception to the observation that elements with similar properties tend to occur in vertical groups.[73] Going along a period, the nuclear charge increases with atomic number as there is as an increase in electrons. The additional 'pull' on outer electrons with increasing nuclear charge generally outweighs the screening efficacy 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.[74] Going down a main group, the effect of increasing nuclear charge is generally outweighed by the effect of additional electrons being further away from the nucleus. With some irregularities, atoms become larger, ionization energy falls, and metallic character increases.[75] The combined effect of these competing horizontal and vertical trends is that the location of the metal-nonmetal transition zone shifts to the right in going down a group.[73] A related effect can be seen in other diagonal similarities between some elements and their lower right neighbours, such as lithium-magnesium, beryllium-aluminium, carbon-phosphorus, and nitrogen-sulfur.[76]
Number, composition and alternative treatments
How many and which elements are metalloids depends on the classification criteria being used. Emsley[77] recognised four metalloids: germanium, arsenic, antimony and tellurium; James et al.[78] listed twelve: boron, carbon, silicon, germanium, arsenic, selenium, antimony, tellurium, bismuth, polonium, ununpentium and livermorium. On average, seven elements are included in such lists. Although there is no standardized division of the elements into metals, metalloids and nonmetals, individual classification arrangements tend to share common ground, with most variations occurring around the indistinct[79] margins.[n 11]
Some authors do not classify elements bordering the metal-nonmetal dividing line as metalloids, noting that a binary classification can facilitate the establishment of rules for determining bond types between metals and nonmetals.[3] Metalloids are grouped with metals,[81] regarded as nonmetals[82] or treated as a sub-category of nonmetals.[83][n 12] Other authors have suggested that classifying some elements as metalloids 'emphasizes that properties change gradually rather than abruptly as one moves across or down the periodic table'.[85] Some periodic tables distinguish elements that are metalloids in the absence of any formal dividing line between metals and nonmetals. Metalloids are shown as occurring in a diagonal band[86] or diffuse region.[87]
Properties of metalloids
Physical and chemical
Metalloids are usually characterised as metallic-looking brittle solids with intermediate to relatively good electrical conductivity values, and each having the electronic band structure of a semimetal or semiconductor. Chemically, they mostly behave as (weak) nonmetals, have intermediate ionization energies and electronegativity values, and have amphoteric or weakly acidic oxides. They can form alloys with metals. Most of the other physical and chemical properties of metalloids are intermediate in nature.
Distinctive
Brittleness,[88] semiconductivity[89] or both[90] have been used as distinguishing indicators of metalloid status. Not all elements classified in the literature as metalloids display semiconductivity, although most do.[91] Metallic lustre along with dualistic chemical behaviour—for example, amphoterism—has also been cited as a benchmark.[92]
Metalloids are all solid,[93] and have metallic lustre, but their other properties vary.[94] Given that metallic character (for example) is a combination of several properties, it has been suggested that metalloid status be judged separately for each element. This could be done based on the extent to which an element exhibits properties relevant to such status.[95]
Compared to metals and nonmetals
Characteristic properties of metals, metalloids and nonmetals are summarized in the table.[96] Physical properties are listed in order of ease of determination; chemical properties run from general to specific, and then to descriptive.
Physical property | Metals | Metalloids | Nonmetals |
---|---|---|---|
Form | solid; a few liquid at or near room temperature (Ga, Hg, Rb, Cs, Fr)[97][n 13] | solid[99] | mostly gases[100] |
Appearance | lustrous (at least when freshly fractured) | lustrous[99] | several colourless; others coloured, or metallic grey to black |
Elasticity | typically elastic, ductile, malleable (when solid) | brittle[101] | brittle, if solid |
Electrical conductivity | good to high[n 14] | intermediate[103] to good[n 15] | poor to good[n 16] |
Band structure | metallic (Bi = semimetallic) | are semiconductors or, if not (As, Sb = semimetallic), exist in semiconducting forms[107] | semiconductor or insulator[108] |
Chemical property | Metals | Metalloids | Nonmetals |
General chemical behaviour | metallic | nonmetallic[109] | nonmetallic |
Ionization energy | relatively low | intermediate ionization energies,[110] usually falling between those of metals and nonmetals[111] | relatively high |
Electronegativity | usually low | have electronegativity values close to 2[112] (revised Pauling scale) or within the range of 1.9–2.2 (Allen scale)[24][n 17] | high |
When mixed with metals |
give alloys | can form alloys[115] | ionic or interstitial compounds formed |
Oxides | lower oxides basic; higher oxides increasingly acidic | amphoteric or weakly acidic[116] | acidic |
The metalloid properties for 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.
Typical or shared applications
Metalloids are too brittle to have any structural uses in their pure forms.[117] They and their compounds are used as (or in) alloying components, biological agents (toxicological, nutritional and medicinal), flame retardants, glasses (oxide and metallic), optical storage media, semiconductors and electronics.[n 18]
Alloys
![Several dozen metallic pellets, reddish-brown with a highly polished appearance, as if they had a cellophane coating.](http://upload.wikimedia.org/wikipedia/commons/thumb/7/74/Copper_germanium.jpg/220px-Copper_germanium.jpg)
Writing early in the history of 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.[120] Compounds of silicon, germanium, arsenic and antimony with the poor metals, it has been suggested, 'are probably best classed as alloys.'[121]
Alloys with transition metals are well-represented. Boron can form intermetallic compounds and alloys with such metals, of the composition MnB, if n > 2.[122] 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 coinage metals.[123] Arsenic can form alloys with metals, including platinum and copper;[124] it is also added to copper and its alloys to improve corrosion resistance[125] and appears to confer the same benefit when added to magnesium.[126] 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).[127] Tellurium readily alloys with iron, as ferrotellurium (50–58% tellurium), and with copper, in the form of copper tellurium (40–50% tellurium).[128] Ferrotellurium is used as a stabilizer for carbon in steel casting.[129] Of the non-metallic elements that are less often recognised as metalloids, selenium—in the form of ferroselenium (50–58% selenium)—is used to improve the workability of stainless steels.[130]
Biological agents
![A clear glass dish on which is a small mound of a white crystalline powder.](http://upload.wikimedia.org/wikipedia/commons/thumb/a/a7/Arsenic_trioxide.jpg/220px-Arsenic_trioxide.jpg)
All six of the elements commonly recognised as metalloids have toxic, dietary or medicinal properties.[132] Arsenic, and antimony compounds are especially toxic; boron, silicon, and possibly arsenic, are essential trace elements. Boron, silicon, arsenic and antimony have medical applications (with germanium and tellurium being thought to have potential).
Boron is used in insecticides[133] and herbicides.[134] It is an essential trace element.[135] As boric acid, it has antiseptic, antifungal, and antiviral properties.
Silicon is present in silatrane, a highly toxic rodenticide.[136] Long-term inhalation of silica dust causes silicosis, a fatal disease of the lungs. Silicon is an essential trace element.[135] As a silicone gel it can be applied to badly burned patients to reduce scarring.[137]
Salts of germanium are potentially harmful to humans and animals if ingested on a prolonged basis.[138] Although there is interest in the pharmacological actions of germanium compounds, there is (as yet) no licensed medicine.[139]
Arsenic is notoriously poisonous and may also be an essential element in ultratrace amounts.[140] It has been used as a pharmaceutical agent since antiquity, including for the treatment of syphilis prior to the development of antibiotics.[141] 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.[141]
Metallic antimony is relatively non-toxic but most antimony compounds are poisonous.[142] Compounds of antimony are used as antiprotozoan drugs, and in some veterinary preparations.
Tellurium is not considered particularly toxic although as little as two grams of sodium tellurate, if administered, can be lethal.[143] People exposed to small amounts of airborne tellurium exude a foul and persistent garlic-like odour.[144] Tellurium dioxide has been used to treat seborrhoeic dermatitis; other tellurium compounds were used as antimicrobial agents before the development of antibiotics.[145] In future, such compounds may need to be substituted for antibiotics that have become ineffective due to bacterial resistance.[146]
Of the elements less often recognised as metalloids, beryllium and lead are noted for their toxicity, with lead arsenate having been extensively used as an insecticide.[147] Sulfur is one of the oldest of the fungicides and pesticides. Phosphorus, sulfur, zinc and iodine are essential nutrients, as are possibly aluminium, tin and lead.[140] Sulfur, gallium, iodine and bismuth have medicinal applications. Sulfur is a constituent of sulfonamide drugs, still widely used for conditions such as acne and urinary tract infections.[148] Gallium nitrate is used to treat the side effects of cancer;[149] gallium citrate, a radiopharmaceutical, facilitates imaging of inflamed body areas.[150] Iodine is used as a disinfectant in various forms. Bismuth is an ingredient in some antibacterials.[151]
Flame retardants
Compounds of boron, silicon, arsenic and antimony have been used as flame retardants. Boron, in the form of borax, has been used as a textile flame retardant since at least the 18th century.[152] Silicon compounds such as silicones, silanes, silsesquioxane, silica and silicates, some of which were developed as alternatives to more toxic halogenated products, can considerably improve the flame retardancy of plastic materials.[153] Arsenic compounds such as sodium arsenite or sodium arsenate are effective flame retardants for wood but were less frequently used due to their toxicity.[154] Antimony trioxide is a flame retardant.[155] Aluminium hydroxide has been used as a wood-fibre, rubber, plastic and textile flame retardant since the 1890s.[156] Apart from aluminium hydroxide, use of phosphorus based flame-retardants in the form of, for example, organophosphates, now exceeds that of any of the other main retardant types, which employ boron, antimony or halogenated hydrocarbon compounds.[157]
Glass formation
![A bunch of pale yellow semi-transparent thin strands, with bright points of white light at their tips.](http://upload.wikimedia.org/wikipedia/commons/thumb/4/42/Fibreoptic4.jpg/220px-Fibreoptic4.jpg)
The oxides B2O3, SiO2, GeO2, As2O3 and Sb2O3 readily form glasses. TeO2 forms a glass but this requires a 'heroic quench rate' or the addition of an impurity, otherwise the crystalline form results.[158] These compounds are used in chemical, domestic and industrial glassware[159] and optics.[160] Boron trioxide is used as a glass fibre additive;[161] it is also a component of borosilicate glass, which is widely used for laboratory glassware and domestic ovenware.[162] Most ordinary glassware is made from silicon dioxide.[163] Germanium dioxide is used as a glass fibre additive, as well as in infrared optical systems.[164] Arsenic trioxide is used in the glass industry as a decolourizing and fining agent, as is antimony trioxide.[165] Tellurium dioxide finds application in laser and nonlinear optics.[166]
Amorphous metallic glasses are generally most easily prepared if one of the components is a metalloid or 'near metalloid' such as boron, carbon, silicon, phosphorus or germanium.[167][n 19] Aside from thin films deposited at very low temperatures, the first known metallic glass was an alloy of composition Au75Si25 reported in 1960.[169] A metallic glass having a strength and toughness not previously seen, of composition Pd82.5P6Si9.5Ge2, was reported in 2011.[170]
Phosphorus, selenium and lead, which are less often recognised as metalloids, are also used in glasses. Phosphate glass has a substrate of phosphorus pentoxide (P2O5), rather than the silica (SiO2) of conventional silicate glasses and is used, for example, to make sodium lamps.[171] Selenium compounds can be used both as decolourising agents and to add a red colour to glass.[172] 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.[173] Lead-based glasses have also been extensively used in electronics 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.[174]
Optical storage
Varying compositions of GeSbTe ("GST alloys") and Ag- and In- doped Sb2Te ("AIST alloys"), being examples of phase-change materials, are widely used in rewritable optical discs 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.[175]
Semiconductors and electronics
![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.](http://upload.wikimedia.org/wikipedia/commons/thumb/8/8d/Semiconductor-1.jpg/220px-Semiconductor-1.jpg)
All the elements commonly recognised as metalloids (or their compounds) have been used in the semiconductor or solid-state electronic industries.[176] Some properties of boron have limited its use as a semiconductor. It has a high melting point, single crystals are relatively hard to obtain, and introducing and retaining controlled impurities is difficult.[177] Silicon is the leading commercial semiconductor; it forms the basis of modern electronics (including standard solar cells)[178] and information and communication technologies.[179] 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.[180] Silicon has largely replaced germanium in semiconducting devices, being cheaper, more resilient at higher operating temperatures, and easier to work during the microelectronic fabrication process.[119] Germanium is however 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.[119] The synthesis of gram-scale quantities of semiconducting germanane was reported in 2013. This comprises 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.[181]
Although arsenic and antimony are not semiconductors in their standard states, both form 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. These compounds are preferred for some special applications.[182] Tellurium, which is a semiconductor in its standard state, is used mainly as a component in type II/VI semiconducting-chalcogenides; these have applications in electro-optics and electronics.[183] 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.[178] Bismuth telluride (Bi2Te3), alloyed with selenium and antimony, is a component of thermoelectric devices, used for refrigeration or portable power generation.[184] Five metalloids—boron, silicon, germanium, arsenic and antimony—can be found in cell phones (along with at least 39 other metals and nonmetals).[185] Tellurium is expected to find such use.[186] 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 semiconductors.[187] The commercial use of gallium compounds is dominated by semiconductor applications—in integrated circuits; cell phones; laser diodes; light emitting diodes; photodetectors; and solar cells.[188] Selenium is used in the production of solar cells[189] and in high-energy surge protectors.[190]
Nomenclature and history
Derivation and other names
The word metalloid comes from the Latin metallum ("metal") and the Greek oeides ("resembling in form or appearance").[191] The terms amphoteric element,[192] boundary element,[193] half-metal,[194] half-way element,[195] near metal,[196] meta-metal,[197] semiconductor,[198] semimetal[199] and submetal[200] are sometimes used synonymously although some of these have other meanings that may not be interchangeable. 'Amphoteric element' is sometimes used more broadly to include transition metals capable of forming oxyanions, such as chromium and manganese.[201] 'Half-metal' is used in physics to refer to a compound (such as chromium dioxide) or alloy that can act as a conductor and an insulator. 'Meta-metal' is sometimes used instead to refer to certain metals (Be, Zn, Cd, Hg, In, Tl, β-Sn, Pb) located just to the left of the metalloids on standard periodic tables.[194] These metals are mostly diamagnetic[202] and tend to have distorted crystalline structures, electrical conductivity values at the lower end of those of metals, and amphoteric (weakly basic) oxides.[203] 'Semimetal' sometimes refers, loosely or explicitly, to metals with incomplete metallic character in crystalline structure, electrical conductivity or electronic structure. Examples include gallium,[204] ytterbium,[205] bismuth[206] and neptunium.[207] The terms amphoteric element and semiconductor are problematic as some elements referred to as metalloids do not show marked amphoteric behaviour or semiconductivity in their most stable forms.
Origin and usage
The origin and usage of the term 'metalloid' is convoluted. Its origin lies in attempts, dating from antiquity, to describe metals and to distinguish between typical and less typical forms. It was first applied in the early 19th century to metals that floated on water (sodium and potassium), and then more popularly to nonmetals. Earlier usage in mineralogy, to describe a mineral having a metallic appearance, can be sourced to as early as 1800.[208] Since the mid-20th century it has been used to refer to intermediate or borderline chemical elements.[46] The International Union of Pure and Applied Chemistry (IUPAC) previously recommended abandoning the term metalloid, and suggested using the term 'semimetal' instead.[209] Use of this latter term has more recently been discouraged[2] as it has a different meaning in physics—one that more specifically 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'.[210]
Elements commonly recognised as metalloids
- Properties noted in this section refer to the elements in their most thermodynamically stable forms under ambient conditions.
Boron
![Several dozen small angular stone like shapes, grey with scattered silver flecks and highlights.](http://upload.wikimedia.org/wikipedia/commons/thumb/1/19/Boron_R105.jpg/220px-Boron_R105.jpg)
Pure boron is a shiny, silver-grey crystalline solid.[212] It is less dense than aluminium (2.34 v 2.70 g/cm3), and is hard and brittle. It is barely reactive under normal conditions, except for attack by fluorine,[213] and has a melting point of 2076 °C (cf. steel ~1370 °C).[214] Boron is a semiconductor,[215] with a room temperature electrical conductivity of 1.5 × 10−6 S•cm−1[216] (about 200 times less than that of tap water)[217] and a band gap of about 1.56 eV.[218][n 20]
The chemistry of boron is dominated by its small atomic size, relatively high ionization energy, and having fewer valence electrons (three) than atomic orbitals (four) available for bonding. With only three valence electrons, simple covalent bonding is electron deficient with respect to the octet rule.[220] Elements in this situation usually adopt metallic bonding, but the small size and high ionization energies of boron tends to instead result in delocalized covalent bonding,[221] in which three atoms share two electrons. The associated structural component that pervades the various allotropes of boron is the icosahedral B12 unit. This also occurs, as do deltahedral variants or fragments, in several metal borides, certain hydrides, and some halides.[222] The bonding in boron has been described as being characteristic of behaviour intermediate between metals and nonmetallic covalent network solids (such as diamond).[223] 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 how close boron is to the metal-nonmetal borderline.[224]
Most of the chemistry of boron is nonmetallic in nature.[224] The small size of the boron atom enables the preparation of many interstitial alloy-type borides.[225] Analogies between boron and transition metals have been noted in the formation of complexes,[226] and adducts (for example, BH3 + CO →BH3CO and, similarly, Fe(CO)4 + CO →Fe(CO)5), as well as in the geometric and electronic structures of cluster species such as [B6H6]2– and [Ru6(CO)18]2–.[227][n 21] The aqueous chemistry of boron is characterised by the formation of many different polyborate anions.[229] Given its high charge-to-size ratio, nearly all compounds of boron are covalent, with a few complexed anionic and cationic species.[230] Boron has a strong affinity for oxygen and a duly extensive borate chemistry.[225] The oxide B2O3 is polymeric in structure,[231] weakly acidic,[232] and a glass former.[233] Organometallic compounds of boron have been known since the 19th century (see organoboron chemistry).[234]
Silicon
![A lustrous blue grey potato shaped lump with an irregular corrugated surface.](http://upload.wikimedia.org/wikipedia/commons/thumb/e/e9/SiliconCroda.jpg/220px-SiliconCroda.jpg)
Silicon is a shiny crystalline solid, with a blue-grey metallic lustre.[235] Like boron, it is less dense than aluminium (2.33 v 2.70 g/cm3), and is hard and brittle.[236] It is a relatively unreactive element,[235] the massive, crystalline form (especially if pure) being 'remarkably inert to all acids, including hydrofluoric'.[237] Less pure silicon, and the powdered form, are variously susceptible to attack by strong or heated acids, as well as by steam and fluorine.[238] Silicon dissolves in hot aqueous alkalis with the evolution of hydrogen, as do metals[239] such as beryllium, aluminium, zinc, gallium or indium.[240] It melts at 1414 °C. Silicon is a semiconductor with an electrical conductivity of 10−4 S•cm−1[241] and a band gap of about 1.11 eV.[242] When it melts, silicon becomes a reasonable metal[243] with an electrical conductivity of 1.0–1.3 × 104 S•cm−1, similar to that of liquid mercury.[244]
The chemistry of silicon is generally nonmetallic (covalent) in nature.[245] It can form alloys with metals such as iron and copper.[246] Silicon shows fewer tendencies to anionic behaviour than ordinary nonmetals.[247] Its solution chemistry is characterised by the formation of oxyanions.[248] The high strength of the silicon-oxygen bond dominates the chemical behaviour of silicon.[249] Polymeric silicates, built up by tetrahedral SiO4 units sharing their oxygen atoms, are the most abundant and important compounds of silicon.[250] The polymeric borates, comprising linked trigonal and tetrahedral BO3 or BO4 units, are built on similar structural principles.[251] The oxide SiO2 is polymeric in structure,[231] weakly acidic,[252][n 22] and a glass former.[233] Traditional organometallic chemistry includes the carbon compounds of silicon (see organosilicon).[255]
Germanium
![Grayish lustrous block with uneven cleaved surface.](http://upload.wikimedia.org/wikipedia/commons/thumb/0/08/Polycrystalline-germanium.jpg/220px-Polycrystalline-germanium.jpg)
Germanium is a shiny grey-white solid.[256] It has a density of 5.323 g/cm3 and is hard and brittle.[257] It is mostly unreactive at room temperature[n 23] but is slowly attacked by hot concentrated sulfuric or nitric acid.[259] Germanium also reacts with molten caustic soda to yield sodium germanate Na2GeO3, together with the evolution of hydrogen.[260] It melts at 938 °C. Germanium is a semiconductor with an electrical conductivity of around 2 × 10−2 S•cm−1[259] and a band gap of 0.67 eV.[261] Liquid germanium is a metallic conductor, with an electrical conductivity on par with that of liquid mercury.[262]
Most of the chemistry of germanium is characteristic of a nonmetal.[263] It can form alloys with metals such as aluminium and gold.[264] Germanium shows fewer tendencies to anionic behaviour than ordinary nonmetals.[247] Its solution chemistry is characterised by the formation of oxyanions.[248] Germanium generally forms tetravalent (IV) compounds, and it can also form less stable divalent (II) compounds, in which it behaves more like a metal.[265] Germanium analogues of all of the major types of silicates have been prepared.[266] The metallic character of germanium is also suggested by the formation of various oxoacid salts. A phosphate [(HPO4)2Ge·H2O] and highly stable trifluoroacetate Ge(OCOCF3)4 have been described, as have Ge2(SO4)2, Ge(ClO4)4 and GeH2(C2O4)3.[267] The oxide GeO2 is polymeric,[231] amphoteric,[268] and a glass former.[233] The dioxide is soluble in acidic solutions (the monoxide GeO, is even more so), and this is sometimes used as a basis to classify germanium as a metal.[269] Up to the 1930s germanium was in fact considered to be a poorly conducting metal rather than a nonmetal.[270] As is the case with all the elements commonly recognised as metalloids, germanium has an established organometallic chemistry (see organogermanium chemistry).[271]
Arsenic
![Two dull silver clusters of crystalline shards.](http://upload.wikimedia.org/wikipedia/commons/thumb/7/7b/Arsen_1a.jpg/220px-Arsen_1a.jpg)
Arsenic is a grey, metallic looking solid. It has a density of 5.727 g/cm3 and is brittle, and moderately hard (more than aluminium; less than iron).[272] 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 Na3AsO3, together with the evolution of hydrogen.[273] Arsenic sublimes at 615 °C. The vapour is lemon-yellow and smells like garlic.[274] Arsenic only melts under a pressure of 38.6 atm, at 817 °C.[275] It is a semimetal with an electrical conductivity of around 3.9 × 104 S•cm−1[276] and a band overlap of 0.5 eV.[277][n 24] Liquid arsenic is a semiconductor with a band gap of 0.15 eV.[279]
The chemistry of arsenic is predominately nonmetallic.[280] It can form alloys with many metals; most of these are brittle.[281] Arsenic shows fewer tendencies to anionic behaviour than ordinary nonmetals.[247] Its solution chemistry is characterised by the formation of oxyanions.[248] Arsenic generally forms compounds in which it has an oxidation state of +3 or +5.[282] The halides, and the oxides and their derivatives are illustrative examples.[250] In the trivalent state, arsenic shows some incipient metallic properties.[283] The halides are hydrolysed by water but these reactions, particularly those of the chloride, are reversible with the addition of a hydrohalic acid.[284] The oxide is acidic but, as noted below, (weakly) amphoteric. The higher, less stable, pentavalent state has strongly acidic (nonmetallic) properties.[285] Compared to phosphorus, the stronger metallic character of arsenic is indicated by the formation of oxoacid salts such as AsPO4, As2(SO4)3[n 25] and arsenic acetate As(CH3COO)3.[289] The oxide As2O3 is polymeric,[231] amphoteric,[290][n 26] and a glass former.[233] Arsenic has an extensive organometallic chemistry (see organoarsenic chemistry).[293]
Antimony
![A glistening silver rock-like chunk, with a blue tint, and roughly parallel furrows.](http://upload.wikimedia.org/wikipedia/commons/thumb/5/5c/Antimony-4.jpg/220px-Antimony-4.jpg)
Antimony is a silver-white solid with a blue tint and a brilliant lustre.[273] It has a density of 6.697 g/cm3 and is brittle, and moderately hard (more so than arsenic; less so than iron; about the same as copper).[272] It is stable in air and moisture at room temperature. It is attacked by concentrated nitric acid, yielding the hydrated pentoxide Sb2O5. Aqua regia gives the pentachloride SbCl5 and hot concentrated sulfuric acid results in the sulfate Sb2(SO4)3.[294] It is not affected by molten alkali.[295] Antimony is capable of displacing hydrogen from water, when heated: 2 Sb + 3 H2O → Sb2O3 + 3 H2.[296] It melts at 631 °C. Antimony is a semimetal with an electrical conductivity of around 3.1 × 104 S•cm−1[297] and a band overlap of 0.16 eV.[277][n 27] Liquid antimony is a metallic conductor with an electrical conductivity of around 5.3 × 104 S•cm−1.[299]
Most of the chemistry of antimony is characteristic of a nonmetal.[300] It can form alloys with one or more metals such as aluminium,[301] iron, nickel, copper, zinc, tin, lead and bismuth.[302] Antimony has fewer tendencies to anionic behaviour than ordinary nonmetals.[247] Its solution chemistry is characterised by the formation of oxyanions.[248] Like arsenic, antimony generally forms compounds in which it has an oxidation state of +3 or +5.[282] The halides, and the oxides and their derivatives are illustrative examples.[250] 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 3d10 electrons. In comparison, the tendency of antimony to oxidize more easily partially offsets the effect of its 4d10 shell.[303] Tripositive antimony is amphoteric; pentapositive antimony is (predominately) acidic.[304] Consistent with an increase in metallic character down group 15, antimony forms salts or salt-like compounds including a nitrate Sb(NO3)3, phosphate SbPO4, sulfate Sb2(SO4)3 and perchlorate Sb(ClO4)3.[305] The otherwise acidic pentoxide Sb2O5 shows some basic (metallic) behaviour in that it can be dissolved in very acidic solutions, with the formation of the oxycation SbO+
2.[306] The oxide Sb2O3 is a polymeric,[231] amphoteric,[307] and a glass former.[233] Antimony has an extensive organometallic chemistry (see organoantimony chemistry).[308]
Tellurium
![A shiny silver-white medallion with a striated surface, irregular around the outside, with a square spiral-like pattern in the middle.](http://upload.wikimedia.org/wikipedia/commons/thumb/c/c1/Tellurium2.jpg/220px-Tellurium2.jpg)
Tellurium is a silvery-white shiny solid.[310] It has a density of 6.24 g/cm3, is brittle, and is the softest of the commonly recognised metalloids, being marginally harder than sulfur.[272] 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 H2O → TeO2 + 2 H2.[311] It reacts (to varying degrees) with nitric, sulfuric and hydrochloric acids to give compounds such as the sulfoxide TeSO3 or tellurous acid H2TeO3,[312] the basic nitrate (Te2O4H)+(NO3)–,[313] or the oxide sulfate Te2O3(SO4).[314] It dissolves in boiling alkalis, with the formation of the tellurite and telluride: 3 Te + 6 KOH = K2TeO3 + 2 K2Te + 3 H2O, a reaction that proceeds or is reversible with increasing or decreasing temperature.[315] At higher temperatures tellurium is sufficiently plastic to extrude.[316] It melts at 449.51 °C. Crystalline tellurium has a structure consisting of parallel infinite spiral chains. Whereas the bonding between adjacent atoms in a chain is covalent, there is evidence of a weak metallic interaction between the neighbouring atoms of different chains.[317] Tellurium is a semiconductor with an (intrinsic) electrical conductivity of around 1.0 S•cm−1[318] and a band gap of 0.32 to 0.38 eV.[319] Liquid tellurium is a semiconductor, with an electrical conductivity, on melting, of around 1.9 × 103 S•cm−1[319] Superheated liquid tellurium is a metallic conductor.[320]
Most of the chemistry of tellurium is characteristic of a nonmetal.[321] It can form alloys with aluminium, silver and tin.[322] Tellurium shows fewer tendencies to anionic behaviour than ordinary nonmetals.[247] Its solution chemistry is characterised by the formation of oxyanions.[248] Tellurium generally forms compounds in which it has an oxidation state of −2, +4 or +6, with the tetrapositive state being the most stable.[311] It combines easily with most other elements to form binary tellurides XxTey these representing the most common mineral form. Non-stoichiometry is frequently encountered. This is particularly so with the transition metals, where electronegativity differences are small and irregular valence is favoured. Many of the associated tellurides can be treated as metallic alloys.[323] The increase in metallic character evident in tellurium, as compared to the lighter chalcogens, is further reflected in the reported formation of various other oxyacid salts, such as a basic selenate 2TeO2·SeO3 and an analogous perchlorate and periodate 2TeO2·HXO4.[324] Tellurium forms a polymeric,[231] amphoteric,[307] glass-forming oxide[233] TeO2. The latter is a 'conditional' glass-forming oxide—it forms a glass with a very small amount of additive.[233] Tellurium has an extensive organometallic chemistry (see organotellurium chemistry).[325]
Elements less commonly recognised as metalloids
Carbon
![A shiny grey-black cuboid nugget with a rough surface.](http://upload.wikimedia.org/wikipedia/commons/thumb/4/48/Graphite2.jpg/220px-Graphite2.jpg)
Carbon is ordinarily classified as a nonmetal[327] but has some metallic properties and is occasionally classified as a metalloid.[328] Hexagonal graphitic carbon is the most thermodynamically stable allotrope of carbon under ambient conditions.[329] It has a lustrous appearance[330] and is a fairly good electrical conductor.[331] Graphite has a layered structure. Each layer comprises carbon atoms bonded to three other carbon atoms in a honeycomb lattice arrangement. The layers are stacked together and held loosely by van der Waals forces and delocalized valence electrons.[332] Like a metal, the conductivity of graphite in the direction of its planes decreases as the temperature is raised;[333][n 28] it has the electronic band structure of a semimetal.[333] The allotropes of carbon, including graphite, can accept foreign atoms or compounds into their structures via substitution, intercalation or doping (interstitial or intrastitial) with the resulting materials being referred to as 'carbon alloys'.[337] Carbon can form ionic salts, including a sulfate, perchlorate, nitrate, hydrogen selenate, and hydrogen phosphate.[338][n 29] In organic chemistry, carbon can form complex cations—termed carbocations—in which the positive charge is on the carbon atom; examples are CH+
3 and CH+
5, and their derivatives.[339]
Carbon is brittle,[340] and behaves as a semiconductor in a direction perpendicular to its planes.[333] Most of its chemistry is nonmetallic;[341] it has a relatively high ionization energy[342] and, compared to most metals, a relatively high electronegativity.[343] Carbon can form anions such as C4– (methanide), C2–
2 (acetylide) and C3–
4 (sesquicarbide or allylenide), in compounds with metals of main groups 1–3, and with the lanthanides and actinides.[344] Its oxide CO2 forms a medium-strength carbonic acid H2CO3.[345][n 30]
Aluminium
![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.](http://upload.wikimedia.org/wikipedia/commons/thumb/5/5d/Aluminium-4.jpg/220px-Aluminium-4.jpg)
Aluminium is classified as a metal. It is lustrous, malleable and ductile, has high electrical and thermal conductivity and a close-packed crystalline structure.[348]
It has some properties that are unusual for a metal; taken together,[349] these properties are sometimes used as a basis to classify aluminium as a metalloid.[350] Its crystalline structure shows some evidence of directional bonding.[351] Although it forms an Al3+ cation in some compounds, aluminium bonds covalently in most others.[352] Its oxide is amphoteric,[353] and a conditional glass-former.[233] Aluminium can form anionic aluminates,[349] such behaviour being considered nonmetallic in character.[71]
Stott[354] labels aluminium as a weak metal. It has the physical properties of a metal but some of the chemical properties of a nonmetal. Steele[355] notes the paradoxical chemical behaviour of aluminium. It resembles a weak metal with its amphoteric oxide and the covalent character of many of its compounds. Yet it is also a strongly electropositive metal, with a high negative electrode potential.
Classifying aluminium as a metalloid is disputed[356] as it has many metallic properties. It is therefore, arguably, an exception to the mnemonic that elements adjacent to the metal-nonmetal dividing line are metalloids.[357][n 31]
Selenium
![A small glass jar filled with small dull grey concave buttons. The pieces of selenium look like tiny mushrooms without their stems.](http://upload.wikimedia.org/wikipedia/commons/thumb/d/d3/Selenium_black_%28cropped%29.jpg/170px-Selenium_black_%28cropped%29.jpg)
Selenium shows borderline metalloid or nonmetal behaviour.[360][n 32]
Its most stable form, the grey trigonal allotrope, is sometimes called 'metallic' selenium. This is because its electrical conductivity is several orders of magnitude greater than that of the red monoclinic form.[363] The metallic character of selenium is further shown by its lustre[364] and its crystalline structure, the latter of which is thought to include weakly 'metallic' interchain bonding.[365] Selenium can be drawn into thin threads when molten.[366] It shows reluctance to acquire 'the high positive oxidation numbers characteristic of nonmetals'.[367] It can form cyclic polycations (such as Se2+
8) when dissolved in oleums[368] (an attribute it shares with sulfur and tellurium), and a hydrolysed cationic salt in the form of trihydroxoselenium (IV) perchlorate [Se(OH)3]+·ClO–
4.[369]
The nonmetallic character of selenium is shown by its brittleness[364] and the low electrical conductivity (~10−9 to 10−12 S•cm−1) of its highly purified form.[105] This is comparable to or less than that of bromine (7.95×10–12 S•cm−1),[370] a nonmetal. Selenium has the electronic band structure of a semiconductor[371] and retains its semiconducting properties in liquid form.[371] It has a relatively high[372] electronegativity (2.55 revised Pauling scale). Its reaction chemistry is mainly that of its nonmetallic anionic forms Se2–, SeO2−
3 and SeO2−
4.[373]
Selenium is commonly described as a metalloid in the environmental chemistry literature.[374] It moves through the aquatic environment like arsenic and antimony do.[375] Its water-soluble salts, in higher concentrations, have a toxicological profile similar to that of arsenic.[376]
Polonium
![A thin film of a bluish-grey metal on a stainless steel disc.](http://upload.wikimedia.org/wikipedia/en/thumb/6/66/Polonium.jpg/220px-Polonium.jpg)
Polonium is 'distinctly metallic' in some ways.[377] Both of its allotropic forms are metallic conductors.[377] It is soluble in acids, forming the rose-coloured Po2+ cation and displacing hydrogen: Po + 2 H+ → Po2+ + H2.[378] Many polonium salts are known.[379] The oxide PoO2 is predominantly basic in nature.[380] Polonium is a reluctant oxidizing agent, unlike its lighter congener oxygen: highly reducing conditions are required for the formation of the Po2– anion in aqueous solution.[381]
Polonium also shows nonmetallic character in its halides, and by way of the existence of polonides. The halides have properties generally characteristic of nonmetal halides (being volatile, easily hydrolyzed, and soluble in organic solvents).[382] Many metal polonides, obtained by heating the elements together at 500–1,000 °C, and containing the Po2– anion, are also known.[383]
Astatine
Astatine, which is ordinarily classified as a nonmetal[384] has some marked metallic properties[385] and may be more appropriately classified as either a metalloid[386] or a metal.[n 33] Immediately following its production in 1940, early investigators considered it a metal.[388] In 1949 it was called the most noble (difficult to reduce) nonmetal as well as being a relatively noble (difficult to oxidize) metal.[389] In 1950 astatine was described as a halogen and (therefore) a reactive nonmetal.[390] In 2013, on the basis of relativistic modelling, astatine was predicted to be a monatomic metal, with a face-centred cubic crystalline structure.[391]
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.[392][n 34] Astatine may be metallic in the liquid state on the basis that elements with an enthalpy of vaporization (EoV) greater than ~42 kJ/mol are metallic when liquid.[394] Such elements include boron,[n 35] silicon, germanium, antimony, selenium and tellurium. Estimated values for the EoV of diatomic astatine are 50 kJ/mol or higher;[398] diatomic iodine, with an EoV of 41.71,[399] falls just short of the threshold figure. '[L]ike 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'.[400][n 36] Further indications of a tendency for astatine to behave like a (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'.[402] It has also been argued that astatine demonstrates cationic behaviour, by way of stable At+ and AtO+ forms, in strongly acidic aqueous solutions.[403]
Some of astatine's reported properties are nonmetallic. It has the narrow liquid range ordinarily associated with nonmetals (mp 302 °C; bp 337 °C).[404] Batsanov gives a calculated band gap energy for astatine of 0.7 eV;[405] this is consistent with nonmetals (in physics) having separated valence and conduction bands and thereby being either semiconductors or insulators.[406] The chemistry of astatine in aqueous solution is mainly characterised by the formation of various anionic species.[407] Most of its known compounds resemble those of iodine,[408] which is a halogen and a nonmetal.[409] Such compounds include astatides (XAt), astatates (XAtO3), and monovalent interhalogen compounds.[410]
Restrepo et al.[411] 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.
Honourable mentions
Elements near the metalloids
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[415] or noted for their metalloidal character. To the left of the metal-nonmetal dividing line, such elements include gallium,[416] tin[417] and bismuth.[418] They show unusual packing structures,[419] marked covalent chemistry (molecular or polymeric),[420] and amphoterism.[421] To the right of the dividing line are carbon,[422] phosphorus,[423] selenium[424] and iodine.[425] They exhibit metallic lustre, semiconducting properties[n 38] 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;[n 39] and selenium as grey selenium.
Allotropes
![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.](http://upload.wikimedia.org/wikipedia/commons/thumb/2/2b/Sn-Alpha-Beta.jpg/220px-Sn-Alpha-Beta.jpg)
Different crystalline forms of an element are called 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.[431] The existence of such allotropes can complicate the classification of the elements involved.[432]
Tin, for example, has two allotropes: 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 9.17×104 S·cm−1 (~1/6th that of copper).[433] 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)×102 S·cm−1 (~1/250th that of white tin).[434] Grey tin has the same crystalline structure as that of diamond. It behaves as a semiconductor (with a band gap of 0.08 eV), but has the electronic band structure of a semimetal.[435] It has been referred to as either a very poor metal,[436] a metalloid,[437] a nonmetal[438] or a near metalloid.[418]
The diamond allotrope of carbon is clearly nonmetallic, being translucent and having a low electrical conductivity of 10−14 to 10−16 S·cm−1.[439] Graphite has an electrical conductivity of 3×104 S·cm−1,[440] a figure more characteristic of a metal. Phosphorus, sulfur, arsenic, selenium, antimony and bismuth also have less stable allotropes that display different behaviours.[441]
Notes
- ^ For a related commentary see also: Vernon RE 2013, 'Which Elements Are Metalloids?', Journal of Chemical Education, vol. 90, no. 12, pp. 1703–1707, doi:10.1021/ed3008457
- ^ On the fuzziness of metalloids see, for example: Rouvray;[4] Cobb & Fetterolf;[5] and Fellet.[6] For the 'buffer zone' terminology see Rochow.[7]
- ^ 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 metals and nonmetals.'[11]
- 'Between the metals and nonmetals in the periodic table we find elements…[that] share some of the characteristic properties of both the metals and nonmetals, making it difficult to place them in either of these two main categories.'[12]
- 'Chemists sometimes use the name metalloid…for these elements which are difficult to classify one way or the other.'[13]
- 'Because the traits distinguishing metals and nonmetals are qualitative in nature, some elements do not fall unambiguously in either category. These elements…are called metalloids…'.[14]
- ^ Gold, for example, has mixed properties but is still recognised as 'king of metals.' Besides metallic behaviour (such as high electrical conductivity, and cation formation), gold shows nonmetallic behaviour:
- It has the most positive electrode potential.
- Its electronegativity of 2.54 is highest among the metals and exceeds that of some nonmetals (hydrogen 2.2; phosphorus 2.19; and radon 2.2).
- It has the most negative electron affinity.
- It has the third-highest ionization energy among the metals (after zinc and mercury).
- It forms the Au– auride anion thereby behaving analogously to the halogens.
- It sometimes has a tendency, known as 'aurophilicity', to bond to itself.[19]
- ^ Mann et al.[24] refer to these elements as 'the recognized metalloids'.
- ^ Rochow[47] concluded that no single measurement indicates exactly which elements are properly classified as metalloids, and that students and teachers usually agree to use electronegativity as a compromise. He described metalloids as a collection of in between elements, of electronegativity 1.8 to 2.2 (classical Pauling scale), that "...resemble metals, yet are not completely metallic either in appearance or in properties," and "...are neither metals nor nonmetals." In mentioning 'appearance', Rochow referred to the intermediate reflectivity values of the commonly recognised metalloids,[48] compared to the intermediate to typically high[49] values of the metals,[50] and the zero or low (mostly)[51] to intermediate reflectivity values of the non-metals.[52] See also, for example:
- Hill and Hollman,[13] who characterise metalloids (in part) on the basis that they are 'poor conductors of electricity with atomic conductance usually less than 10−3 but greater than 10−5 ohm−1 cm−4'.
- Bond,[53] who suggests that 'one criterion for distinguishing semi-metals from true metals under normal conditions is that the co-ordination number of the former is never greater than eight'.
- Edwards et al.,[54] who state that, 'Using the Goldhammer-Herzfeld criterion with measured atomic electronic polarizabilities and condensed phase molar volumes allows one to readily predict which elements are metallic, which are nonmetallic, and which are borderline when in their condensed phases (solid or liquid).'
- ^ Selenium has an ionization energy (IE) of ~226 kcal/mol and is sometimes described as a semiconductor. It has a relatively high 2.55 electronegativity (EN). Polonium has an IE of ~196 kcal/mol and a 2.0 EN, but has a metallic band structure.[57] Astatine has an IE of 215 kJ/mol and an EN of 2.2.[58] Its electronic band structure is not known with any certainty.
- ^ The Goldhammer-Herzfeld criterion is a ratio that compares the force holding an individual atom's valence electrons in place with the forces, acting on the same electrons, arising 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. Metallic behaviour is then predicted.[60] Otherwise nonmetallic behaviour is anticipated.
- ^ Gallium is unusual (for a metal) in having a packing efficiency of just 39%.[63] Other notable values are 42.9 for bismuth[64] and 58.5 for liquid mercury.[65]
- ^ As the ratio is based on classical arguments[68] 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 grounds.[69] Even so it offers a first order rationalization for the occurrence of metallic character amongst the elements.[70]
- ^ Jones[80] 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'.
- ^ Oderberg[84] argues on ontological grounds that anything not a metal is therefore a nonmetal, and that this includes semi-metals (i.e. metalloids).
- ^ Copernicium is reportedly the only metal known to be a gas at room temperature.[98]
- ^ Metals have electrical conductivity values of from 6.9 × 103 S•cm−1 for manganese to 6.3 × 105 for silver.[102]
- ^ Metalloids have electrical conductivity values of from 1.5 × 10−6 S•cm−1 for boron to 3.9 × 104 for arsenic.[104] If selenium is included as a metalloid the applicable conductivity range would start from ~10−9 to 10−12 S•cm−1.[105]
- ^ Nonmetals have electrical conductivity values of from ~10−18 S•cm−1 for the elemental gases to 3 × 104 in graphite.[106]
- ^ Chedd[113] 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[114] described this choice as arbitrary, as other elements with electronegativities 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.
- ^ Olmsted and Williams[118] 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'.
- ^ 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.[168]
- ^ 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).[219]
- ^ On the analogy between boron and metals, Greenwood[228] 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...'.
- ^ Although SiO2 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 Si5O(PO4)6,[253] and with hydrofluoric acid to give hexafluorosilicic acid H2SiF6.[254]
- ^ Temperatures above 400 °C are required to form a noticeable surface oxide layer.[258]
- ^ 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 form, with a band gap of around 1.2–1.4 eV.[278]
- ^ The formulae of AsPO4 and As2(SO4)3 suggest straightforward ionic formulations, with As3+, but compounds in which arsenic is present as a cation are extremely rare.[286] AsPO4, 'which is virtually a covalent oxide,' has been referred to as a double oxide, of the form As2O3·P2O5. It comprises AsO3 pyramids and PO4 tetrahedra, joined together by all their corner atoms to form a continuous polymeric network.[287] As2(SO4)3 has a structure in which each SO4 tetrahedron is bridged by two AsO3 trigonal pyramida.[288]
- ^ As2O3 is usually regarded as being amphoteric but a few sources say it is (weakly)[291] acidic. They describe its 'basic' properties (that is, its reaction with concentrated hydrochloric acid to form arsenic trichloride) as being alcoholic, in analogy with the formation of covalent alklyl chlorides by covalent alcohols (e.g., R-OH + HCl → RCl + H2O)[292]
- ^ Antimony can also be prepared in an amorphous semiconducting black form, with an estimated (temperature-dependent) band gap of 0.06–0.18 eV.[298]
- ^ Liquid carbon may[334] or may not[335] be a metallic conductor, depending on pressure and temperature; see also.[336]
- ^ 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. Analogous salts are formed in other strong acids, such as perchloric acid.[338]
- ^ Only a small fraction of dissolved CO2 is present in water as carbonic acid so, even though H2CO3 is a medium-strong acid, solutions of carbonic acid are only weakly acidic.[346]
- ^ A mnemonic that captures the elements commonly recognised as metalloids goes: Up, up-down, up-down, up...are the metalloids![358]
- ^ Rochow,[361] who later wrote his 1966 monograph The metalloids,[362] commented that, 'In some respects selenium acts like a metalloid and tellurium certainly does'.
- ^ A third option is to include astatine both as a nonmetal and as a metalloid.[387]
- ^ A visible piece of astatine would be immediately and completely vaporized because of the heat generated by its intense radioactivity.[393]
- ^ The literature is contradictory as to whether boron exhibits metallic conductivity in liquid form. Krishnan et al.[395] found that liquid boron behaved like a metal. Glorieux et al.[396] characterised liquid boron as a semiconductor, on the basis of its low electrical conductivity. Millot et al.[397] reported that the emissivity of liquid boron was not consistent with that of a liquid metal.
- ^ Korenman[401] similarly noted that 'the ability precipitate with hydrogen sulfide distinguishes astatine from other halogens and brings it closer to bismuth and other heavy metals.'
- ^ 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).[413] 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.[414]
- ^ For example: intermediate electrical conductivity;[426] a relatively narrow band gap;[427] light sensitivity.[426]
- ^ White phosphorus is the most common, industrially important,[428] and easily reproducible allotrope. For those reasons it is the standard state of the element.[429] Paradoxically, it is the least thermodynamically stable, and the most volatile and reactive form.[430]
Citations
- ^ Chedd 1969, pp. 58, 78; National Research Council 1984, p. 43
- ^ a b Atkins et al. 2010, p. 20
- ^ a b Roher 2001, pp. 4–6
- ^ Rouvray 1995, p. 546. Rouvray submits that classifying the electrical conductivity of the elements using the overlapping domains of metals, metalloids, and nonmetals better reflects reality than a strictly black or white paradigm.
- ^ Cobb & Fetterolf 2005, p. 64: 'The division between metals and nonmetals is rather fuzzy, so the elements in the immediate vicinity of the zigzag staircase line are called metalloids, which means they don't fit either definition exactly.'
- ^ Fellet 2011: 'Chemistry has all sorts of fuzzy definitions'.
- ^ Rochow 1977, p. 14
- ^ Goldsmith 1982, p. 526; Hawkes 2001, p. 1686
- ^ Hawkes 2001, p. 1687
- ^ Sharp 1981, p. 299
- ^ Cusack 1987, p. 360
- ^ Kelter, Mosher & Scott 2009, p. 268
- ^ a b Hill & Holman 2000, p. 41
- ^ King 1979, p. 13
- ^ Moore 2011, p. 81
- ^ Gray 2010
- ^ Hopkins & Bailar 1956, p. 458
- ^ Glinka 1965, p. 77
- ^ Wiberg 2001, p. 1279
- ^ Belpassi et al. 2006, pp. 4543–4
- ^ Schmidbaur & Schier 2008, pp. 1931–51
- ^ Tyler Miller 1987, p. 59
- ^ Goldsmith 1982, p. 526; Hawkes 2001, p. 1686; Boylan 1962, p. 493; Sherman & Weston 1966, p. 64; Wulfsberg 1991, p. 201; Kotz, Treichel & Weaver 2009, p. 62
- ^ a b Mann et al. 2000, p. 2783
- ^ Hawkes 2001, p. 1686; Segal 1989, p. 965; McMurray & Fay 2009, p. 767
- ^ Bucat 1983, p. 26; Brown c. 2007
- ^ a b Swift & Schaefer 1962, p. 100
- ^ Hawkes 2001, p. 1686; Hawkes 2010; Holt, Rinehart & Wilson c. 2007
- ^ 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).
- ^ Tilden 1876, pp. 172, 198–201; Smith 1994, p. 252; Bodner & Pardue 1993, p. 354
- ^ Bassett et al. 1966, p. 127
- ^ Rausch 1960
- ^ Thayer 1977, p. 604; Warren & Geballe 1981; Masters & Ela 2008, p. 190
- ^ Warren & Geballe 1981; Chalmers 1959, p. 72; US Bureau of Naval Personnel 1965, p. 26
- ^ Siebring 1967, p. 513
- ^ Wiberg 2001, p. 282
- ^ Rausch 1960; Friend 1953, p. 68
- ^ Murray 1928, p. 1295
- ^ Hampel & Hawley 1966, p. 950; Stein 1985; Stein 1987, pp. 240, 247–8
- ^ Hatcher 1949, p. 223; Secrist & Powers 1966, p. 459
- ^ Taylor 1960, p. 614
- ^ Considine & Considine 1984, p. 568; Cegielski 1998, p. 147; The American heritage science dictionary 2005 p. 397
- ^ Woodward 1948, p. 1
- ^ NIST 2010. Values shown in the above table have been converted from the NIST values, which are given in eV.
- ^ Berger 1997; Lovett 1977, p. 3
- ^ a b Goldsmith 1982, p. 526
- ^ Rochow 1966, pp. 1, 4–7
- ^ Lagrenaudie 1953; Rochow 1966, pp. 23, 25
- ^ Askeland, Fulay & Wright 2011, p. 806
- ^ Born & Wolf 1999, p. 746
- ^ Burakowski & Wierzchoń 1999, p. 336
- ^ Olechna & Knox 1965, pp. A991‒2
- ^ Bond 2005, p. 3
- ^ Edwards et al. 2010, p. 958
- ^ Jones 2010, p. 169
- ^ 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.
- ^ Kraig, Roundy & Cohen 2004, p. 412; Alloul 2010, p. 83
- ^ Vernon 2013, pp. 1704
- ^ Edwards & Sienko 1983, p. 695; Edwards et al. 2010
- ^ Herzfeld 1927; Edwards 2000, pp. 100–3
- ^ Van Setten et al. 2007, pp. 2460–1; Russell & Lee 2005, p. 7 (Si, Ge); Pearson 1972, p. 264 (As, Sb, Te; also black P)
- ^ Russell & Lee 2005, p. 1
- ^ Russell & Lee 2005, pp. 6–7, 387
- ^ a b Pearson 1972, p. 264
- ^ Okakjima & Shomoji 1972, p. 258
- ^ Kitaĭgorodskiĭ 1961, p. 108
- ^ a b c Neuburger 1936
- ^ Edwards 1999, p. 416
- ^ Steurer 2007, p. 142; Pyykkö 2012, p. 56
- ^ Edwards & Sienko 1983, p. 695
- ^ a b Hamm 1969, p. 653
- ^ Horvath 1973, p. 336
- ^ a b Gray 2009, p. 9
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- ^ Asmussen & Reinhard 2002, p. 7
- ^ Deprez & McLachan 1988
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