Metalloid

Elements recognized as metalloids
	13	14	15	16	17
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 (93%): B, Si, Ge, As, Sb, Te   Irregularly recognized (44%): Po, At   Less commonly recognized (24%): Se   Rarely recognized (9%): C, Al   Arbitrary metal-metalloid dividing line: between Be and B, Al and Si, Ge and As, Sb and Te, Po and At
Recognition status, as metalloids, of some elements in the p-block of the periodic table. Percentages are median appearance frequencies in the list of metalloid lists.[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 with properties in between, or that are a mixture of, those of metals and nonmetals. There is neither a standard definition of a metalloid nor complete agreement on the elements appropriately classified as such. Despite the lack of specificity, the term remains in use in the literature of chemistry.

The six commonly recognised metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Five elements are less frequently so classified: carbon, aluminium, selenium, polonium, and astatine. On a standard periodic table, all eleven are in a diagonal area in the p-block extending from boron at the upper left to astatine at lower right, along the dividing line between metals and nonmetals shown on some periodic tables.

Typical metalloids have a metallic appearance, but they are brittle and only fair conductors of electricity. Chemically, they behave mostly as nonmetals. They can form alloys with metals. Most of their other physical and chemical properties are intermediate in nature. Metalloids are usually too brittle to have any structural uses. They and their compounds are used in alloys, biological agents, catalysts, flame retardants, glasses, optical storage and optoelectronics, pyrotechnics, 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 sometimes are called semimetals, a practice that has been discouraged,^[2] as the term semimetal has a different meaning in physics than in chemistry. In physics it more specifically refers to the electronic band structure of a substance.

Definitions

Judgement-based

A metalloid is an element with properties in between, or that are a mixture of, those of metals and nonmetals, and which is therefore hard to classify as either a metal or a nonmetal. This is a generic definition that draws on metalloid attributes consistently cited in the literature.^{[n 2]} Difficulty of categorisation is a key attribute. Most elements have a mixture of metallic and nonmetallic properties,^[9] and can be classified according to which set of properties is more pronounced.^{[10][n 3]} Only the elements at or near the margins, lacking a sufficiently clear preponderance of either metallic or nonmetallic properties, are classified as metalloids.^[14]

Boron, silicon, germanium, arsenic, antimony, and tellurium are recognised commonly as metalloids.^{[15][n 4]} Depending on the author, one or more from selenium, polonium, or astatine sometimes are added to the list.^[17] Boron sometimes is excluded, by itself, or with silicon.^[18] Sometimes tellurium is not regarded as a metalloid.^[19] The inclusion of antimony, polonium, and astatine as metalloids also has been questioned.^[20]

Other elements occasionally are classified as metalloids. These elements include^[21] hydrogen,^[22] beryllium,^[23] nitrogen,^[24] phosphorus,^[25] sulfur,^[26] zinc,^[27] gallium,^[28] tin, iodine,^[29] lead,^[30] bismuth,^[19] and radon.^[31] The term metalloid also has 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.^[32] The p-block metals,^[33] and nonmetals (such as carbon or nitrogen) that can form alloys with metals^[34] or modify their properties^[35] also have occasionally been considered as metalloids.

Criteria-based

Element	IE (kcal/mol)	IE (kJ/mol)	EN	Band structure
Boron	191	801	2.04	semiconductor
Silicon	188	787	1.90	semiconductor
Germanium	182	762	2.01	semiconductor
Arsenic	226	944	2.18	semimetal
Antimony	199	831	2.05	semimetal
Tellurium	208	869	2.10	semiconductor
average	199	832	2.05
The elements commonly recognised as metalloids, and their ionization energies (IE);[36] electronegativities (EN, revised Pauling scale); and electronic band structures[37] (most thermodynamically-stable forms under ambient conditions).

No widely accepted definition of a metalloid exists, nor any division of the periodic table into metals, metalloids and nonmetals;^[38] Hawkes^[39] questioned the feasibility of establishing a specific definition, noting that anomalies can be found in several attempted constructs. Classifying an element as a metalloid has been described by Sharp^[40] as "arbitrary".

The number and identities of metalloids depend on what classification criteria are used. Emsley^[41] recognised four metalloids (germanium, arsenic, antimony and tellurium); James et al.^[42] listed twelve (Emsley's plus boron, carbon, silicon, selenium, bismuth, polonium, ununpentium and livermorium). On average, seven elements are included in such lists; individual classification arrangements tend to share common ground and vary in the ill-defined^[43] margins.^{[n 5][n 6]}

A single quantitative criterion such as electronegativity is commonly used,^[46] metalloids having electronegativity values from 1.8 or 1.9 to 2.2.^[47] Further examples include packing efficiency (the fraction of volume in a crystal structure occupied by atoms) and the Goldhammer-Herzfeld criterion ratio.^[48] The commonly recognised metalloids have packing efficiencies of between 34% and 41%.^{[n 7]} The Goldhammer-Herzfeld ratio, roughly equal to the cube of the atomic radius divided by the molar volume,^{[56][n 8]} is a simple measure of how metallic an element is, the recognised metalloids having ratios from around 0.85 to 1.1 and averaging 1.0.^{[58][n 9]} Other authors have relied on, for example, atomic conductance^{[n 10][62]} or bulk coordination number.^[63]

Jones, writing on the role of classification in science, observed that "[classes] are usually defined by more than two attributes".^[64] Masterton and Slowinski^[65] used three criteria to describe the six elements commonly recognised as metalloids: metalloids have ionization energies around 200 kcal/mol (837 kJ/mol) and electronegativity values close to 2.0. They also said that metalloids are typically semiconductors, though antimony and arsenic (semimetals from a physics perspective) have electrical conductivities approaching those of metals. Selenium and polonium are suspected as not in this scheme, while astatine's status is uncertain.^{[n 11]}

Periodic table territory

Location

Distribution and recognition status of elements classified as metalloids
	1	2		12	13	14	15	16	17	18
	H									He
	Li	Be			B	C	N	O	F	Ne
	Na	Mg			Al	Si	P	S	Cl	Ar
	K	Ca		Zn	Ga	Ge	As	Se	Br	Kr
	Rb	Sr		Cd	In	Sn	Sb	Te	I	Xe
	Cs	Ba		Hg	Tl	Pb	Bi	Po	At	Rn
	Fr	Ra		Cn	Uut	Fl	Uup	Lv	Uus	Uuo
	Commonly (93%) to rarely (9%) recognised as a metalloid: B, C, Al, Si, Ge, As, Se, Sb, Te, Po, At   Very rarely (1–5%): H, Be, P, S, Ga, Sn, I, Pb, Bi, Fl, Uup, Lv, Uus   Sporadically: N, Zn, Rn   Metal–nonmetal dividing line: between H and Li, Be and B, Al and Si, Ge and As, Sb and Te, Po and At, and Uus and Uuo
Periodic table extract showing groups 1–2 and 12–18, and a dividing line between metals and nonmetals. Percentages are median appearance frequencies in the list of metalloid lists. Sporadically recognised elements show that the metalloid net is sometimes cast very widely; although they do not appear in the list of metalloid lists, isolated references to their designation as metalloids can be found in the literature (as cited in this article).

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.^[68] 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.^[69]

The diagonal positioning of the metalloids represents an exception to the observation that elements with similar properties tend to occur in vertical groups.^[70] A related effect can be seen in other diagonal similarities between some elements and their lower right neighbours, specifically lithium-magnesium, beryllium-aluminium, and boron-silicon. Rayner-Canham^[71] has argued that these similarities extend to carbon-phosphorus, nitrogen-sulfur, and into three d-block series.

This exception arises due to competing horizontal and vertical trends in the nuclear charge. Going along a period, the nuclear charge increases with atomic number as do the number of electrons. The additional pull on outer electrons as nuclear charge increases generally outweighs the screening effect of having more electrons. With some irregularities, atoms therefore become smaller, ionization energy increases, and there is a gradual change in character, across a period, from strongly metallic, to weakly metallic, to weakly nonmetallic, to strongly nonmetallic elements.^[72] 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. Atoms generally become larger, ionization energy falls, and metallic character increases.^[73] The net effect is that the location of the metal–nonmetal transition zone shifts to the right in going down a group,^[70] and analogous diagonal similarities are seen elsewhere in the periodic table, as noted.^[74]

Alternative treatments

Depictions of metalloids vary according to the author. Some 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.^[75] Metalloids are variously grouped with metals,^[76] regarded as nonmetals^[77] or treated as a sub-category of nonmetals.^{[78][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".^[80] Some periodic tables distinguish elements that are metalloids and display no formal dividing line between metals and nonmetals. Metalloids are shown as occurring in a diagonal band^[81] or diffuse region.^[82]

Properties

Metalloids usually look like metals but behave largely like nonmetals. Physically, they are shiny, brittle solids with intermediate to relatively good electrical conductivity and the electronic band structure of a semimetal or semiconductor. Chemically, they mostly behave as (weak) nonmetals, have intermediate ionization energies and electronegativity values, and amphoteric or weakly acidic oxides. They can form alloys with metals. Most of their other physical and chemical properties are intermediate in nature.

Compared to metals and nonmetals

Characteristic properties of metals, metalloids and nonmetals are summarized in the table.^[83] Physical properties are listed in order of ease of determination; chemical properties run from general to specific, and then to descriptive.

Properties of metals, metalloids and nonmetals

Physical property	Metals	Metalloids	Nonmetals
Form	solid; a few liquid at or near room temperature (Ga, Hg, Rb, Cs, Fr)[84][n 13]	solid[86]	majority gaseous[87]
Appearance	lustrous (at least when freshly fractured)	lustrous[86]	several colourless; others coloured, or metallic grey to black
Elasticity	typically elastic, ductile, malleable (when solid)	brittle[88]	brittle, if solid
Electrical conductivity	good to high[n 14]	intermediate[90] 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[94]	semiconductor or insulator[95]
Chemical property	Metals	Metalloids	Nonmetals
General chemical behaviour	metallic	nonmetallic[96]	nonmetallic
Ionization energy	relatively low	intermediate ionization energies,[97] usually falling between those of metals and nonmetals[98]	relatively high
Electronegativity	usually low	have electronegativity values close to 2[99] (revised Pauling scale) or within the range of 1.9–2.2 (Allen scale)[16][n 17]	high
When mixed with metals	give alloys	can form alloys[102]	ionic or interstitial compounds formed
Oxides	lower oxides basic; higher oxides increasingly acidic	amphoteric or weakly acidic[103]	acidic

The above table reflects the hybrid nature of metalloids. The properties of form, appearance, and behaviour when mixed with metals are more like metals. Elasticity and general chemical behaviour are more like nonmetals. Electrical conductivity, band structure, ionization energy, electronegativity, and oxides are intermediate between the two.

Common applications

The focus of this section is on the recognised metalloids. Elements less often recognised as metalloids are ordinarily classified as either metals or nonmetals; some of these are included here for comparative purposes.

Metalloids are too brittle to have any structural uses in their pure forms.^[104] They and their compounds are used as (or in) alloying components, biological agents (toxicological, nutritional and medicinal), catalysts, flame retardants, glasses (oxide and metallic), optical storage media and optoelectronics, pyrotechnics, semiconductors and electronics.^{[n 18]}

Alloys

Copper-germanium alloy pellets, likely ~84% Cu; 16% Ge.^[106] When combined with silver the result is a tarnish resistant sterling silver. Also shown are two silver pellets.

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.^[107] Phillips and Williams^[108] suggested that compounds of silicon, germanium, arsenic and antimony with B metals, "are probably best classed as alloys".

Among the lighter metalloids, alloys with transition metals are well-represented. Boron can form intermetallic compounds and alloys with such metals of the composition M_nB, if n > 2.^[109] 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.^[110]

The heavier metalloids continue the theme. Arsenic can form alloys with metals, including platinum and copper;^[111] it is also added to copper and its alloys to improve corrosion resistance^[112] and appears to confer the same benefit when added to magnesium.^[113] 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).^[114] Tellurium readily alloys with iron, as ferrotellurium (50–58% tellurium), and with copper, in the form of copper tellurium (40–50% tellurium).^[115] Ferrotellurium is used as a stabilizer for carbon in steel casting.^[116] Of the non-metallic elements less often recognised as metalloids, selenium—in the form of ferroselenium (50–58% selenium)—is used to improve the machinability of stainless steels.^[117]

Biological agents

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.^[118]

All six of the elements commonly recognised as metalloids have toxic, dietary or medicinal properties.^[119] Arsenic and antimony compounds are especially toxic; boron, silicon, and possibly arsenic, are essential trace elements. Boron, silicon, arsenic and antimony have medical applications, and germanium and tellurium are thought to have potential.

Boron is used in insecticides^[120] and herbicides.^[121] It is an essential trace element.^[122] As boric acid, it has antiseptic, antifungal, and antiviral properties.^[123]

Silicon is present in silatrane, a highly toxic rodenticide.^[124] Long-term inhalation of silica dust causes silicosis, a fatal disease of the lungs. Silicon is an essential trace element.^[122] Silicone gel can be applied to badly burned patients to reduce scarring.^[125]

Salts of germanium are potentially harmful to humans and animals if ingested on a prolonged basis.^[126] There is interest in the pharmacological actions of germanium compounds but no licensed medicine as yet.^[127]

Arsenic is notoriously poisonous and may also be an essential element in ultratrace amounts.^[128] During World War I, both sides used "arsenic-based sneezing and vomiting agents…to force enemy soldiers to remove their gas masks before firing mustard or phosgene at them in a second salvo.^[129] It has been used as a pharmaceutical agent since antiquity, including for the treatment of syphilis before the development of antibiotics.^[130] 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.^[130] Arsenic in drinking water, which causes lung and bladder cancer, has been associated with a reduction in breast cancer mortality rates.^[131]

Metallic antimony is relatively non-toxic, but most antimony compounds are poisonous.^[132] Two antimony compounds, sodium stibogluconate and stibophen, are used as antiparasitical drugs.^[133]

Elemental tellurium is not considered particularly toxic; two grams of sodium tellurate, if administered, can be lethal.^[134] People exposed to small amounts of airborne tellurium exude a foul and persistent garlic-like odour.^[135] Tellurium dioxide has been used to treat seborrhoeic dermatitis; other tellurium compounds were used as antimicrobial agents before the development of antibiotics.^[136] In the future, such compounds may need to be substituted for antibiotics that have become ineffective due to bacterial resistance.^[137]

Of the elements less often recognised as metalloids, beryllium and lead are noted for their toxicity; lead arsenate has been extensively used as an insecticide.^[138] Sulfur is one of the oldest of the fungicides and pesticides. Phosphorus, sulfur, zinc, selenium and iodine are essential nutrients, and aluminium, tin and lead may be.^[128] Sulfur, gallium, selenium, 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.^[139] Gallium nitrate is used to treat the side effects of cancer;^[140] gallium citrate, a radiopharmaceutical, facilitates imaging of inflamed body areas.^[141] Selenium sulfide is used in medicinal shampoos and to treat skin infections such as tinea versicolor.^[142] Iodine is used as a disinfectant in various forms. Bismuth is an ingredient in some antibacterials.^[143]

Catalysts

Boron trifluoride and trichloride are used as catalysts in organic synthesis and electronics; the tribromide is used in the manufacture of diborane.^[144] Non-toxic boron ligands can replace toxic phosphorus ligands in transition metal catalysts.^[145] Silica sulfuric acid (SiO₂OSO₃H) is used in organic reactions.^[146] Germanium dioxide is sometimes used as a catalyst in the production of PET plastic for containers;^[147] cheaper antimony compounds, such as the trioxide or triacetate, are more commonly employed for the same purpose^[148] despite concerns about antimony contamination of food and drinks.^[149] Arsenic trioxide has been used in the production of natural gas, to boost the removal of carbon dioxide, as have selenous acid and tellurous acid.^[150] Selenium acts as a catalyst in some microorganisms.^[151] Tellurium, and its dioxide and tetrachloride, are strong catalysts for air oxidation of carbon above 500 °C.^[152] Graphite oxide can be used as a catalyst in the synthesis of imines and their derivatives.^[153] Activated carbon and alumina have been used as catalysts for the removal of sulfur contaminants from natural gas.^[154] Titanium doped aluminium has been identified as a substitute for expensive noble metal catalysts used in the production of industrial chemicals.^[155]

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.^[156] 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.^[157] Arsenic compounds such as sodium arsenite or sodium arsenate are effective flame retardants for wood but have been less frequently used due to their toxicity.^[158] Antimony trioxide is a flame retardant.^[159] Aluminium hydroxide has been used as a wood-fibre, rubber, plastic and textile flame retardant since the 1890s.^[160] 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. These employ boron, antimony or halogenated hydrocarbon compounds.^[161]

Glass formation

Optical fibres, usually made of pure silicon dioxide glass, with additives such as boron trioxide or germanium dioxide for increased sensitivity

The oxides B₂O₃, SiO₂, GeO₂, As₂O₃ and Sb₂O₃ readily form glasses. TeO₂ forms a glass but this requires a "heroic quench rate"^[162] or the addition of an impurity, otherwise the crystalline form results.^[162] These compounds are used in chemical, domestic and industrial glassware^[163] and optics.^[164] Boron trioxide is used as a glass fibre additive,^[165] and is also a component of borosilicate glass, widely used for laboratory glassware and domestic ovenware for its low thermal expansion.^[166] Most ordinary glassware is made from silicon dioxide.^[167] Germanium dioxide is used as a glass fibre additive, as well as in infrared optical systems.^[168] Arsenic trioxide is used in the glass industry as a decolourizing and fining agent (for the removal of bubbles),^[169] as is antimony trioxide.^[170] Tellurium dioxide finds application in laser and nonlinear optics.^[171]

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.^{[172][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.^[174] A metallic glass having a strength and toughness not previously seen, of composition Pd_82.5P₆Si_9.5Ge₂, was reported in 2011.^[175]

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. It is used, for example, to make sodium lamps.^[176] Selenium compounds can be used both as decolourising agents and to add a red colour to glass.^[177] 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.^[178] 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.^[179]

Optical storage and optoelectronics

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.^[180] Future applications for GeSbTe may include, "ultrafast, entirely solid-state displays with nanometre-scale pixels, semi-transparent "smart" glasses, "smart" contact lenses and artificial retina devices."^[181]

Pyrotechnics

Archaic blue light signal, fuelled by a mixture of sodium nitrate, sulfur and (red) arsenic trisulfide^[182]

The recognised metalloids have either pyrotechnic applications or associated properties. Boron and silicon are commonly encountered;^[183] they act somewhat like metal fuels.^[184] Boron is used in pyrotechnic initiator compositions (for igniting other hard-to-start compositions), and in delay compositions that burn at a constant rate.^[185] Boron carbide has been identified as a possible replacement for more toxic barium or hexachloroethane mixtures in smoke munitions, signal flares and fireworks.^[186] Silicon, like boron, is a component of initiator and delay mixtures.^[185] Doped germanium can act as a variable speed thermite fuel.^{[n 20]} Arsenic trisulfide As₂S₃ was used in old naval signal lights; in fireworks to make white stars;^[188] in yellow smoke screen mixtures; and in initiator compositions.^[189] Antimony trisulfide Sb₂S₃ is found in white-light fireworks and in flash and sound mixtures.^[190] Tellurium has been used in delay mixtures and in blasting cap initiator compositions.^[191]

Carbon, aluminium, phosphorus and selenium continue the theme. Carbon, in black powder, is a constituent of fireworks rocket propellants, bursting charges, and effects mixtures, and military delay fuses and igniters.^{[192][n 21]} Aluminium is a common pyrotechnic ingredient,^[183] and is widely employed for its capacity to generate light and heat,^[194] including in thermite mixtures.^[195] Phosphorus can be found in smoke and incendiary munitions, paper caps used in toy guns, and party poppers.^[196] Selenium has been used in the same way as tellurium.^[191]

Semiconductors and electronics

Semiconductor-based electronic components. From left to right: a transistor, an integrated circuit and an LED. The elements commonly recognised as metalloids find widespread use in such devices, as elemental or compound semiconductor constituents (Si, Ge or GaAs, for example) or as doping agents (B, Sb, Te, for example).

All the elements commonly recognised as metalloids (or their compounds) have been used in the semiconductor or solid-state electronic industries.^[197]

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.^[198]

Silicon is the leading commercial semiconductor; it forms the basis of modern electronics (including standard solar cells)^[199] and information and communication technologies.^[200] 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.^[201]

Germanium has largely been replaced by silicon in semiconducting devices, being cheaper, more resilient at higher operating temperatures, and easier to work during the microelectronic fabrication process.^[106] Germanium is still a constituent of semiconducting silicon-germanium "alloys" and these have been growing in use, particularly for wireless communication devices; such alloys exploit the higher carrier mobility of germanium.^[106] 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.^[202] The development of a germanium-wire based anode that more than doubles the capacity of lithium-ion batteries was reported in 2014.^[203] In the same year, Lee at al. reported that defect-free crystals of graphene large enough to have electronic uses could be grown on, and removed from, a germanium substrate.^[204]

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.^[205] Antimony nanocrystals may enable lithium-ion batteries to be replaced by more powerful sodium ion batteries.^[206]

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.^[207] 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.^[199] Bismuth telluride (Bi₂Te₃), alloyed with selenium and antimony, is a component of thermoelectric devices used for refrigeration or portable power generation.^[208]

Five metalloids—boron, silicon, germanium, arsenic and antimony—can be found in cell phones (along with at least 39 other metals and nonmetals).^[209] Tellurium is expected to find such use.^[210] 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.^[211] 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.^[212] Selenium is used in the production of solar cells^[213] and in high-energy surge protectors.^[214]

Boron, silicon, germanium, antimony and tellurium,^[215] as well as heavier metals and metalloids such as Sm, Hg, Tl, Pb, Bi and Se,^[216] can be found in topological insulators. These are alloys^[217] or compounds which, at ultracold temperatures or room temperature (depending on their composition), are metallic conductors on their surfaces but insulators through their interiors.^[218] Cadmium arsenide Cd₃As₂, at about 1 K, is a Dirac-semimetal—a bulk electronic analogue of graphene—in which electrons travel effectively as massless particles.^[219] These two classes of material are thought to have potential quantum computing applications.^[220]

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").^[221] Several names are sometimes used synonymously although some of these have other meanings that are not necessarily interchangeable: amphoteric element,^[222] boundary element,^[223] half-metal,^[224] half-way element,^[225] near metal,^[226] meta-metal,^[227] semiconductor,^[228] semimetal^[229] and submetal.^[230] "Amphoteric element" is sometimes used more broadly to include transition metals capable of forming oxyanions, such as chromium and manganese.^[231] "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.^[224] These metals are mostly diamagnetic^[232] and tend to have distorted crystalline structures, electrical conductivity values at the lower end of those of metals, and amphoteric (weakly basic) oxides.^[233] "Semimetal" sometimes refers, loosely or explicitly, to metals with incomplete metallic character in crystalline structure, electrical conductivity or electronic structure. Examples include gallium,^[234] ytterbium,^[235] bismuth^[236] and neptunium.^[237] The names amphoteric element and semiconductor are problematic as some elements referred to as metalloids do not show marked amphoteric behaviour (bismuth, for example)^[238] or semiconductivity (polonium)^[239] 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.^[240] Since the mid-20th century it has been used to refer to intermediate or borderline chemical elements.^{[241][n 22]} The International Union of Pure and Applied Chemistry (IUPAC) previously recommended abandoning the term metalloid, and suggested using the term semimetal instead.^[243] Use of this latter term has more recently been discouraged by Atkins et al.^[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.^[244]

Elements commonly recognised as metalloids

Properties noted in this section refer to the elements in their most thermodynamically stable forms under ambient conditions.

Boron

Boron, shown here in the form of its β-rhombohedral phase (its most thermodynamically stable allotrope)^[245]

Pure boron is a shiny, silver-grey crystalline solid.^[246] It is less dense than aluminium (2.34 vs. 2.70 g/cm³), and is hard and brittle. It is barely reactive under normal conditions, except for attack by fluorine,^[247] and has a melting point of 2076 °C (cf. steel ~1370 °C).^[248] Boron is a semiconductor;^[249] its room temperature electrical conductivity is 1.5 × 10⁻⁶ S•cm^−1[250] (about 200 times less than that of tap water)^[251] and it has a band gap of about 1.56 eV.^{[252][n 23]}

The structural chemistry of boron is dominated by its small atomic size, and relatively high ionization energy. With only three valence electrons per boron atom, simple covalent bonding cannot fulfil the octet rule.^[254] Metallic bonding is the usual result among the heavier congenors of boron but this generally requires low ionization energies.^[255] Instead, because of its small size and high ionization energies, the basic structural unit of boron (and nearly all of its allotropes)^{[n 24]} is the icosahedral B₁₂ cluster. Of the 36 electrons associated with 12 boron atoms, 26 reside in 13 delocalized molecular orbitals; the other 10 electrons are used to form two- and three-centre covalent bonds between icosahedra.^[257] The same motif can be seen, as are deltahedral variants or fragments, in metal borides and hydride derivatives, and in some halides.^[258]

The bonding in boron has been described as being characteristic of behaviour intermediate between metals and nonmetallic covalent network solids (such as diamond).^[259] The energy required to transform B, C, N, Si and P from nonmetallic to metallic states has been estimated as 30, 100, 240, 33 and 50 kJ/mol, respectively. This indicates the proximity of boron to the metal-nonmetal borderline.^[260]

Most of the chemistry of boron is nonmetallic in nature.^[260] Unlike its heavier congeners, it is not known to form a simple B³⁺ or hydrated [B(H₂O)₄]³⁺ cation.^[261] The small size of the boron atom enables the preparation of many interstitial alloy-type borides.^[262] Analogies between boron and transition metals have been noted in the formation of complexes,^[263] and adducts (for example, BH₃ + CO →BH₃CO and, similarly, Fe(CO)₄ + CO →Fe(CO)₅),^{[n 25]} as well as in the geometric and electronic structures of cluster species such as [B₆H₆]²⁻ and [Ru₆(CO)₁₈]²⁻.^{[265][n 26]} The aqueous chemistry of boron is characterised by the formation of many different polyborate anions.^[267] Given its high charge-to-size ratio, boron bonds covalently in nearly all of its compounds;^[268] the exceptions are the borides as these include, depending on their composition, covalent, ionic and metallic bonding components.^[269] Simple binary compounds, such as boron trichloride are Lewis acids as the formation of three covalent bonds leaves a hole in the octet which can be filled by an electron-pair donated by a Lewis base.^[254] Boron has a strong affinity for oxygen and a duly extensive borate chemistry.^[262] The oxide B₂O₃ is polymeric in structure,^[270] weakly acidic,^{[271][n 27]} and a glass former.^[277] Organometallic compounds of boron^{[n 28]} have been known since the 19th century (see organoboron chemistry).^[279]

Silicon

Silicon has a blue-grey metallic lustre.

Silicon is a crystalline solid with a blue-grey metallic lustre.^[280] Like boron, it is less dense (at 2.33 g/cm³) than aluminium, and is hard and brittle.^[281] It is a relatively unreactive element.^[280] According to Rochow,^[282] the massive crystalline form (especially if pure) is "remarkably inert to all acids, including hydrofluoric".^{[n 29]} Less pure silicon, and the powdered form, are variously susceptible to attack by strong or heated acids, as well as by steam and fluorine.^[286] Silicon dissolves in hot aqueous alkalis with the evolution of hydrogen, as do metals^[287] such as beryllium, aluminium, zinc, gallium or indium.^[288] It melts at 1414 °C. Silicon is a semiconductor with an electrical conductivity of 10⁻⁴ S•cm^−1[289] and a band gap of about 1.11 eV.^[283] When it melts, silicon becomes a reasonable metal^[290] with an electrical conductivity of 1.0–1.3 × 10⁴ S•cm⁻¹, similar to that of liquid mercury.^[291]

The chemistry of silicon is generally nonmetallic (covalent) in nature.^[292] It is not known to form a cation.^[293] Silicon can form alloys with metals such as iron and copper.^[294] It shows fewer tendencies to anionic behaviour than ordinary nonmetals.^[295] Its solution chemistry is characterised by the formation of oxyanions.^[296] The high strength of the silicon-oxygen bond dominates the chemical behaviour of silicon.^[297] Polymeric silicates, built up by tetrahedral SiO₄ units sharing their oxygen atoms, are the most abundant and important compounds of silicon.^[298] The polymeric borates, comprising linked trigonal and tetrahedral BO₃ or BO₄ units, are built on similar structural principles.^[299] The oxide SiO₂ is polymeric in structure,^[270] weakly acidic,^{[300][n 30]} and a glass former.^[277] Traditional organometallic chemistry includes the carbon compounds of silicon (see organosilicon).^[304]

Germanium

Germanium is sometimes described as a metal

Germanium is a shiny grey-white solid.^[305] It has a density of 5.323 g/cm³ and is hard and brittle.^[306] It is mostly unreactive at room temperature^{[n 31]} but is slowly attacked by hot concentrated sulfuric or nitric acid.^[308] Germanium also reacts with molten caustic soda to yield sodium germanate Na₂GeO₃ and hydrogen gas.^[309] It melts at 938 °C. Germanium is a semiconductor with an electrical conductivity of around 2 × 10⁻² S•cm^−1[308] and a band gap of 0.67 eV.^[310] Liquid germanium is a metallic conductor, with an electrical conductivity similar to that of liquid mercury.^[311]

Most of the chemistry of germanium is characteristic of a nonmetal.^[312] Whether or not germanium forms a cation is unclear, aside from the reported existence of the Ge²⁺ ion in a few esoteric compounds.^{[n 32]} It can form alloys with metals such as aluminium and gold.^[325] It shows fewer tendencies to anionic behaviour than ordinary nonmetals.^[295] Its solution chemistry is characterised by the formation of oxyanions.^[296] Germanium generally forms tetravalent (IV) compounds, and it can also form less stable divalent (II) compounds, in which it behaves more like a metal.^[326] Germanium analogues of all of the major types of silicates have been prepared.^[327] 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.^[328] The oxide GeO₂ is polymeric,^[270] amphoteric,^[329] and a glass former.^[277] The dioxide is soluble in acidic solutions (the monoxide GeO, is even more so), and this is sometimes used to classify germanium as a metal.^[330] Up to the 1930s germanium was considered to be a poorly conducting metal;^[331] it has occasionally been classified as a metal by later writers.^[332] As with all the elements commonly recognised as metalloids, germanium has an established organometallic chemistry (see organogermanium chemistry).^[333]

Arsenic

Arsenic, sealed in a container to prevent tarnishing

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).^[334] 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 and hydrogen gas.^[335] Arsenic sublimes at 615 °C. The vapour is lemon-yellow and smells like garlic.^[336] Arsenic only melts under a pressure of 38.6 atm, at 817 °C.^[337] It is a semimetal with an electrical conductivity of around 3.9 × 10⁴ S•cm^−1[338] and a band overlap of 0.5 eV.^{[339][n 33]} Liquid arsenic is a semiconductor with a band gap of 0.15 eV.^[341]

The chemistry of arsenic is predominately nonmetallic.^[342] Whether or not arsenic forms a cation is unclear.^{[n 34]} Its many metal alloys are mostly brittle.^[350] It shows fewer tendencies to anionic behaviour than ordinary nonmetals.^[295] Its solution chemistry is characterised by the formation of oxyanions.^[296] Arsenic generally forms compounds in which it has an oxidation state of +3 or +5.^[351] The halides, and the oxides and their derivatives are illustrative examples.^[298] In the trivalent state, arsenic shows some incipient metallic properties.^[352] The halides are hydrolysed by water but these reactions, particularly those of the chloride, are reversible with the addition of a hydrohalic acid.^[353] The oxide is acidic but, as noted below, (weakly) amphoteric. The higher, less stable, pentavalent state has strongly acidic (nonmetallic) properties.^[354] Compared to phosphorus, the stronger metallic character of arsenic is indicated by the formation of oxoacid salts such as AsPO₄, As₂(SO₄)₃^{[n 35]} and arsenic acetate As(CH₃COO)₃.^[357] The oxide As₂O₃ is polymeric,^[270] amphoteric,^{[358][n 36]} and a glass former.^[277] Arsenic has an extensive organometallic chemistry (see organoarsenic chemistry).^[361]

Antimony

Antimony, showing its brilliant lustre

Antimony is a silver-white solid with a blue tint and a brilliant lustre.^[335] 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).^[334] 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.^[362] It is not affected by molten alkali.^[363] Antimony is capable of displacing hydrogen from water, when heated: 2 Sb + 3 H₂O → Sb₂O₃ + 3 H₂.^[364] It melts at 631 °C. Antimony is a semimetal with an electrical conductivity of around 3.1 × 10⁴ S•cm^−1[365] and a band overlap of 0.16 eV.^{[339][n 37]} Liquid antimony is a metallic conductor with an electrical conductivity of around 5.3 × 10⁴ S•cm⁻¹.^[367]

Most of the chemistry of antimony is characteristic of a nonmetal.^[368] Antimony has some definite cationic chemistry,^[369] SbO⁺ and Sb(OH)₂⁺ being present in acidic aqueous solution;^{[370][n 38]} the compound Sb₈(GaCl₄)₂, which contains the homopolycation, Sb₈²⁺, was prepared in 2004.^[372] It can form alloys with one or more metals such as aluminium,^[373] iron, nickel, copper, zinc, tin, lead and bismuth.^[374] Antimony has fewer tendencies to anionic behaviour than ordinary nonmetals.^[295] Its solution chemistry is characterised by the formation of oxyanions.^[296] Like arsenic, antimony generally forms compounds in which it has an oxidation state of +3 or +5.^[351] The halides, and the oxides and their derivatives are illustrative examples.^[298] 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 (being a heavier atom) to oxidize more easily partially offsets the effect of its 4d¹⁰ shell.^[375] Tripositive antimony is amphoteric; pentapositive antimony is (predominately) acidic.^[376] 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.^[377] 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.^[378] The oxide Sb₂O₃ is polymeric,^[270] amphoteric,^[379] and a glass former.^[277] Antimony has an extensive organometallic chemistry (see organoantimony chemistry).^[380]

Tellurium

Tellurium, described by Dmitri Mendeleev as forming a transition between metals and nonmetals^[381]

Tellurium is a silvery-white shiny solid.^[382] 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.^[334] 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.^[383] It reacts (to varying degrees) with nitric, sulfuric and hydrochloric acids to give compounds such as the sulfoxide TeSO₃ or tellurous acid H₂TeO₃,^[384] the basic nitrate (Te₂O₄H)⁺(NO₃)⁻,^[385] or the oxide sulfate Te₂O₃(SO₄).^[386] It dissolves in boiling alkalis, to give 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.^[387]

At higher temperatures tellurium is sufficiently plastic to extrude.^[388] It melts at 449.51 °C. Crystalline tellurium has a structure consisting of parallel infinite spiral chains. The bonding between adjacent atoms in a chain is covalent, but there is evidence of a weak metallic interaction between the neighbouring atoms of different chains.^[389] Tellurium is a semiconductor with an electrical conductivity of around 1.0 S•cm^−1[390] and a band gap of 0.32 to 0.38 eV.^[391] Liquid tellurium is a semiconductor, with an electrical conductivity, on melting, of around 1.9 × 10³ S•cm^−1[391] Superheated liquid tellurium is a metallic conductor.^[392]

Most of the chemistry of tellurium is characteristic of a nonmetal.^[393] It shows some cationic behaviour. The dioxide dissolves in acid to yield the trihydroxotellurium (IV) Te(OH)₃⁺ ion;^{[394][n 39]} the red Te₄²⁺ and yellow-orange Te₆²⁺ ions form when tellurium is oxidized in fluorosulfuric acid (HSO₃F), or liquid sulfur dioxide (SO₂), respectively.^[397] It can form alloys with aluminium, silver and tin.^[398] Tellurium shows fewer tendencies to anionic behaviour than ordinary nonmetals.^[295] Its solution chemistry is characterised by the formation of oxyanions.^[296] Tellurium generally forms compounds in which it has an oxidation state of −2, +4 or +6. The +4 state is the most stable.^[383] Tellurides of composition X_xTe_y are easily formed with most other elements and represent the most common tellurium minerals. Nonstoichiometry is pervasive, especially with transition metals. Many tellurides can be regarded as metallic alloys.^[399] 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.^[400] Tellurium forms a polymeric,^[270] amphoteric,^[379] glass-forming oxide^[277] TeO₂. It is a "conditional" glass-forming oxide—it forms a glass with a very small amount of additive.^[277] Tellurium has an extensive organometallic chemistry (see organotellurium chemistry).^[401]

Elements less commonly recognised as metalloids

Carbon

Carbon (as graphite). Delocalized valence electrons within the layers of graphite give it a metallic appearance.^[402]

Carbon is ordinarily classified as a nonmetal^[403] but has some metallic properties and is occasionally classified as a metalloid.^[404] Hexagonal graphitic carbon (graphite) is the most thermodynamically stable allotrope of carbon under ambient conditions.^[405] It has a lustrous appearance^[406] and is a fairly good electrical conductor.^[407] 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.^[408]

Like a metal, the conductivity of graphite in the direction of its planes decreases as the temperature is raised;^{[409][n 40]} it has the electronic band structure of a semimetal.^[409] The allotropes of carbon, including graphite, can accept foreign atoms or compounds into their structures via substitution, intercalation or doping. The resulting materials are referred to as "carbon alloys".^[413] Carbon can form ionic salts, including a hydrogen sulfate, perchlorate, and nitrate (C+
24X⁻.2HX, where X = HSO₄, ClO₄; and C+
24NO–
3.3HNO₃).^{[414][n 41]} 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.^[415]

Carbon is brittle,^[416] and behaves as a semiconductor in a direction perpendicular to its planes.^[409] Most of its chemistry is nonmetallic;^[417] it has a relatively high ionization energy^[418] and, compared to most metals, a relatively high electronegativity.^[419] 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.^[420] Its oxide CO₂ forms carbonic acid H₂CO₃.^{[421][n 42]}

Aluminium

High purity aluminium is very much softer than its familiar alloys. People who handle it for the first time often ask if it is the real thing.^[423]

Aluminium is ordinarily classified as a metal.^[424] It is lustrous, malleable and ductile, and has high electrical and thermal conductivity. Like most metals it has a close-packed crystalline structure,^[425] and forms a cation in aqueous solution.^[426]

It has some properties that are unusual for a metal; taken together,^[427] these are sometimes used as a basis to classify aluminium as a metalloid.^[428] Its crystalline structure shows some evidence of directional bonding.^[429] Aluminium bonds covalently in most compounds.^[430] The oxide Al₂O₃ is amphoteric,^[431] and a conditional glass-former.^[277] Aluminium can form anionic aluminates,^[427] such behaviour being considered nonmetallic in character.^[68]

Classifying aluminium as a metalloid has been disputed^[432] given its many metallic properties. It is therefore, arguably, an exception to the mnemonic that elements adjacent to the metal–nonmetal dividing line are metalloids.^{[433][n 43]}

Stott^[435] labels aluminium as a weak metal. It has the physical properties of a metal but some of the chemical properties of a nonmetal. Steele^[436] notes the paradoxical chemical behaviour of aluminium: "It resembles a weak metal in its amphoteric oxide and in the covalent character of many of its compounds ... Yet it is a highly electropositive metal ... [with] a high negative electrode potential". Moody^[437] says that, "aluminium is on the 'diagonal borderland' between metals and non-metals in the chemical sense."

Selenium

Grey selenium, being a photoconductor, conducts electricity around 1,000 times better when light falls on it, a property used since the mid-1870s in various light-sensing applications^[438]

Selenium shows borderline metalloid or nonmetal behaviour.^{[439][n 44]}

Its most stable form, the grey trigonal allotrope, is sometimes called "metallic" selenium because its electrical conductivity is several orders of magnitude greater than that of the red monoclinic form.^[442] The metallic character of selenium is further shown by its lustre,^[443] and its crystalline structure, which is thought to include weakly "metallic" interchain bonding.^[444] Selenium can be drawn into thin threads when molten and viscous.^[445] It shows reluctance to acquire "the high positive oxidation numbers characteristic of nonmetals".^[446] It can form cyclic polycations (such as Se2+
8) when dissolved in oleums^[447] (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.^[448]

The nonmetallic character of selenium is shown by its brittleness^[443] and the low electrical conductivity (~10⁻⁹ to 10⁻¹² S•cm⁻¹) of its highly purified form.^[92] This is comparable to or less than that of bromine (7.95×10^–12 S•cm⁻¹),^[449] a nonmetal. Selenium has the electronic band structure of a semiconductor^[450] and retains its semiconducting properties in liquid form.^[450] It has a relatively high^[451] electronegativity (2.55 revised Pauling scale). Its reaction chemistry is mainly that of its nonmetallic anionic forms Se²⁻, SeO2−
3 and SeO2−
4.^[452]

Selenium is commonly described as a metalloid in the environmental chemistry literature.^[453] It moves through the aquatic environment similarly to arsenic and antimony;^[454] its water-soluble salts, in higher concentrations, have a similar toxicological profile to that of arsenic.^[455]

Polonium

Polonium is "distinctly metallic" in some ways.^[239] Both of its allotropic forms are metallic conductors.^[239] It is soluble in acids, forming the rose-coloured Po²⁺ cation and displacing hydrogen: Po + 2 H⁺ → Po²⁺ + H₂.^[456] Many polonium salts are known.^[457] The oxide PoO₂ is predominantly basic in nature.^[458] 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.^[459]

Whether polonium is ductile or brittle is unclear. It is predicted to be ductile based on its calculated elastic constants.^[460] It has a simple cubic crystalline structure. Such a structure has few slip systems and "leads to very low ductility and hence low fracture resistance".^[461]

Polonium shows nonmetallic character in its halides, and by the existence of polonides. The halides have properties generally characteristic of nonmetal halides (being volatile, easily hydrolyzed, and soluble in organic solvents).^[462] Many metal polonides, obtained by heating the elements together at 500–1,000 °C, and containing the Po²⁻ anion, are also known.^[463]

Astatine

As a halogen, astatine tends to be classified as a nonmetal.^[464] It has some marked metallic properties^[465] and is sometimes instead classified as either a metalloid^[466] or (less often) as a metal.^{[n 45]} Immediately following its production in 1940, early investigators considered it a metal.^[468] In 1949 it was called the most noble (difficult to reduce) nonmetal as well as being a relatively noble (difficult to oxidize) metal.^[469] In 1950 astatine was described as a halogen and (therefore) a reactive nonmetal.^[470] In 2013, on the basis of relativistic modelling, astatine was predicted to be a monatomic metal, with a face-centred cubic crystalline structure.^[471]

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.^{[472][n 46]} Astatine may be metallic in the liquid state on the basis that elements with an enthalpy of vaporization (∆H_vap) greater than ~42 kJ/mol are metallic when liquid.^[474] Such elements include boron,^{[n 47]} silicon, germanium, antimony, selenium and tellurium. Estimated values for ∆H_vap of diatomic astatine are 50 kJ/mol or higher;^[478] diatomic iodine, with a ∆H_vap of 41.71,^[479] falls just short of the threshold figure.

"Like typical metals, it [astatine] is precipitated by hydrogen sulfide even from strongly acid solutions and is displaced in a free form from sulfate solutions; it is deposited on the cathode on electrolysis."^{[480][n 48]} 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".^[482] It has also been argued that astatine demonstrates cationic behaviour, by way of stable At⁺ and AtO⁺ forms, in strongly acidic aqueous solutions.^[483]

Some of astatine's reported properties are nonmetallic. It has the narrow liquid range ordinarily associated with nonmetals (mp 302 °C; bp 337 °C).^[484] Batsanov gives a calculated band gap energy for astatine of 0.7 eV;^[485] this is consistent with nonmetals (in physics) having separated valence and conduction bands and thereby being either semiconductors or insulators.^[486] The chemistry of astatine in aqueous solution is mainly characterised by the formation of various anionic species.^[487] Most of its known compounds resemble those of iodine,^[488] which is a halogen and a nonmetal.^[489] Such compounds include astatides (XAt), astatates (XAtO₃), and monovalent interhalogen compounds.^[490]

Restrepo et al.^[491] 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.

Related concepts

Near metalloids

Iodine crystals, showing a metallic lustre. Iodine is a semiconductor in the direction of its planes, with a band gap of ~1.3 eV. It has an electrical conductivity of 1.7 × 10⁻⁸ S•cm⁻¹ at room temperature.^[492] This is higher than selenium but lower than boron, the least electrically conducting of the recognised metalloids.^{[n 49]}

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^[495] or noted for their metalloidal character. To the left of the metal–nonmetal dividing line, such elements include gallium,^[496] tin^[497] and bismuth.^[498] They show unusual packing structures,^[499] marked covalent chemistry (molecular or polymeric),^[500] and amphoterism.^[501] To the right of the dividing line are carbon,^[502] phosphorus,^[503] selenium^[504] and iodine.^[505] They exhibit metallic lustre, semiconducting properties^{[n 50]} 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 51]} and selenium as grey selenium.

Allotropes

White tin (left) and grey tin (right). Both forms have a metallic appearance.

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.^[511] The existence of such allotropes can complicate the classification of the elements involved.^[512]

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).^[513] 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).^[514] 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.^[515] It has been referred to as either a very poor metal,^[516] a metalloid,^[517] a nonmetal^[518] or a near metalloid.^[498]

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.^[519] Graphite has an electrical conductivity of 3 × 10⁴ S·cm⁻¹,^[520] a figure more characteristic of a metal. Phosphorus, sulfur, arsenic, selenium, antimony and bismuth also have less stable allotropes that display different behaviours.^[521]

Abundance, extraction and cost

Abundance

Z	Element	Grams /tonne
8	Oxygen	461,000
14	Silicon	282,000
13	Aluminium	82,300
26	Iron	56,300
6	Carbon	200
29	Copper	60
5	Boron	10
33	Arsenic	1.8
32	Germanium	1.5
47	Silver	0.075
34	Selenium	0.05
51	Antimony	0.02
79	Gold	0.004
52	Tellurium	0.001
75	Rhenium	7×10^−10
54	Xenon	3×10^−11
84	Polonium	2×10^−16
85	Astatine	3×10^−20

The table gives crustal abundances of the elements commonly to rarely recognised as metalloids.^[522] Some other elements are included for comparison: oxygen and xenon (the most and least abundant elements with stable isotopes); iron and the coinage metals copper, silver and gold; and rhenium, the least abundant stable metal (aluminium is normally the most abundant metal). Various abundance estimates have been published; these often disagree to some extent.^[523]

Extraction

The recognised metalloids can be obtained by chemical reduction of either their oxides or their sulfides. Simpler or more complex extraction methods may be employed depending on the starting form and economic factors.^[524] Boron is routinely obtained by reducing the trioxide with magnesium: B₂O₃ + 3 Mg → 2 B + 3MgO; after secondary processing the resulting brown powder has a purity of up to 97%.^[525] Boron of higher purity (> 99%) is prepared by heating volatile boron compounds, such as BCl₃ or BBr₃, either in a hydrogen atmosphere (2 BX₃ + 3 H₂ → 2 B + 6 HX) or to the point of thermal decomposition. Silicon and germanium are obtained from their oxides by heating the oxide with carbon or hydrogen: SiO₂ + C → Si + CO₂; GeO₂ + 2 H₂ → Ge + 2 H₂O. Arsenic is isolated from its pyrite (FeAsS) or arsenical pyrite (FeAs₂) by heating; alternatively, it can be obtained from its oxide by reduction with carbon: 2 As₂O₃ + 3 C → 2 As + 3 CO₂.^[526] Antimony is derived from its sulfide by reduction with iron: Sb₂S₃ → 2 Sb + 3 FeS. Tellurium is prepared from its oxide by dissolving it in aqueous NaOH, yielding tellurite, then by electrolytic reduction: TeO₂ + 2 NaOH → Na₂TeO₃ + H₂O;^[527] Na₂TeO₃ + H₂O → Te + 2 NaOH + O₂.^[528] Another option is reduction of the oxide by roasting with carbon: TeO₂ + C → Te + CO₂.^[529]

Production methods for the elements less frequently recognised as metalloids involve natural processing, electrolytic or chemical reduction, or irradiation. Carbon (as graphite) occurs naturally and is extracted by crushing the parent rock and floating the lighter graphite to the surface. Aluminium is extracted by dissolving its oxide Al₂O₃ in molten cryolite Na₃AlF₆ and then by high temperature electrolytic reduction. Selenium is produced by roasting its coinage metal selenides X₂Se (X = Cu, Ag, Au) with soda ash to give the selenite: X₂Se + O₂ + Na₂CO₃ → Na₂SeO₃ + 2 X + CO₂; the selenide is neutralized by sulfuric acid H₂SO₄ to give selenous acid H₂SeO₃; this is reduced by bubbling with SO₂ to yield elemental selenium. Polonium and astatine are produced in minute quantities by irradiating bismuth.^[530]

Cost

The recognised metalloids and their closer neighbours mostly cost less than silver; only polonium and astatine are more expensive than gold. As of 5 April 2014, prices for small samples (up to 100 g) of silicon, antimony and tellurium, and graphite, aluminium and selenium, average around one third the cost of silver (US$1.5 per gram or about $45 an ounce). Boron, germanium and arsenic samples average about three-and-a-half times the cost of silver.^{[n 52]} Polonium is available for about $100 per microgram, which is $100,000,000 a gram.^[531] Zalutsky and Pruszynski^[532] estimate a similar cost for producing astatine. Prices for the applicable elements traded as commodities tend to range from two to three times cheaper than the sample price (Ge), to nearly three thousand times cheaper (As).^{[n 53]}

See also

• Properties of metals, metalloids and nonmetals

Notes

1. 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
2. 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."^[3]
   • "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"^[4]
   • "Chemists sometimes use the name metalloid ... for these elements which are difficult to classify one way or the other."^[5]
   • "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 ..."^[6]
   More broadly, metalloids have been referred to as:
   • "elements that ... are somewhat of a cross between metals and nonmetals";^[7] or
   • "weird in-between elements".^[8]
3. 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 highest electrode potential
   • It has the third-highest ionization energy among the metals (after zinc and mercury)
   • It has the lowest electron affinity
   • 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 forms the Au⁻ auride anion, acting in this way like a halogen
   • It sometimes has a tendency, known as "aurophilicity", to bond to itself;^[11]
   on halogen character, see also Belpassi et al,^[12] who conclude that in the aurides MAu (M = Li–Cs) gold "behaves as a halogen, intermediate between Br and I"; on aurophilicity, see also Schmidbaur and Schier.^[13]
4. Mann et al.^[16] refer to these elements as "the recognized metalloids".
5. Jones^[44] 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."
6. The lack of a standard division of the elements into metals, metalloids and nonmetals is not necessarily an issue. There is more or less, a continuous progression from the metallic to the nonmetallic. Potentially, a specified subset of this continuum can serve its particular purpose as well as any other.^[45]
7. The packing efficiency of boron is 38%; silicon and germanium 34; arsenic 38.5; antimony 41; and tellurium 36.4.^[49] These values are lower than in most metals (80% of which have a packing efficiency of at least 68%),^[50] but higher than those of elements usually classified as nonmetals. (Gallium is unusual, for a metal, in having a packing efficiency of just 39%.^[51] Other notable values for metals are 42.9 for bismuth^[52] and 58.5 for liquid mercury.^[53]) Packing efficiencies for nonmetals are: graphite 17%,^[54] sulfur 19.2,^[55] iodine 23.9,^[55] selenium 24.2,^[55] and black phosphorus 28.5.^[52]
8. More specifically, the Goldhammer-Herzfeld criterion is the ratio of the force holding an individual atom's valence electrons in place with the forces on the same electrons from interactions between the atoms in the solid or liquid element. When the interatomic forces are greater than, or equal to, the atomic force, valence electron itinerancy is indicated and metallic behaviour is predicted.^[57] Otherwise nonmetallic behaviour is anticipated.
9. As the ratio is based on classical arguments^[59] 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.^[60] Even so it offers a first order rationalization for the occurrence of metallic character amongst the elements.^[61]
10. Atomic conductance is the electrical conductivity of one mole of a substance. It is equal to electrical conductivity divided by molar volume.^[5]
11. Selenium has an ionization energy (IE) of 225 kcal/mol (941 kJ/mol) and is sometimes described as a semiconductor. It has a relatively high 2.55 electronegativity (EN). Polonium has an IE of 194 kcal/mol (812 kJ/mol) and a 2.0 EN, but has a metallic band structure.^[66] Astatine has an IE of 215 kJ/mol (899 kJ/mol) and an EN of 2.2.^[67] Its electronic band structure is not known with any certainty.
12. Oderberg^[79] argues on ontological grounds that anything not a metal is therefore a nonmetal, and that this includes semi-metals (i.e. metalloids).
13. Copernicium is reportedly the only metal thought to be a gas at room temperature.^[85]
14. Metals have electrical conductivity values of from 6.9 × 10³ S•cm⁻¹ for manganese to 6.3 × 10⁵ for silver.^[89]
15. Metalloids have electrical conductivity values of from 1.5 × 10⁻⁶ S•cm⁻¹ for boron to 3.9 × 10⁴ for arsenic.^[91] If selenium is included as a metalloid the applicable conductivity range would start from ~10⁻⁹ to 10⁻¹² S•cm⁻¹.^[92]
16. Nonmetals have electrical conductivity values of from ~10⁻¹⁸ S•cm⁻¹ for the elemental gases to 3 × 10⁴ in graphite.^[93]
17. Chedd^[100] 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^[101] described this choice as arbitrary, as other elements whose electronegativities lie in this range include copper, silver, phosphorus, mercury and bismuth. He went on to suggest defining a metalloid as "a semiconductor or semimetal" and to include bismuth and selenium in this category.
18. Olmsted and Williams^[105] 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".
19. 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.^[173]
20. The reaction involved is Ge + 2 MoO₃ → GeO₂ + 2 MoO₂. Adding arsenic or antimony (n-type electron donors) increases the rate of reaction; adding gallium or indium (p-type electron acceptors) decreases it.^[187]
21. Ellern, writing in Military and Civilian Pyrotechnics (1968), comments that carbon black "has been specified for and used in a nuclear air-burst simulator."^[193]
22. For a post-1960 example of the former use of the term metalloid to refer to nonmetals see Zhdanov,^[242] who divides the elements into metals; intermediate elements (H, B, C, Si, Ge, Se, Te); and metalloids (of which the most typical are given as O, F and Cl).
23. 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).^[253]
24. The synthesis of B₄₀ borospherene, a "distorted fullerene with a hexagonal hole on the top and bottom and four heptagonal holes around the waist" was announced in 2014.^[256]
25. The BH₃ and Fe(CO₄) species in these reactions are short-lived reaction intermediates.^[264]
26. On the analogy between boron and metals, Greenwood^[266] 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 ..."
27. Boron trioxide B₂O₃ is sometimes described as being (weakly) amphoteric.^[272] It reacts with alkalies to give various borates.^[273] In its hydrated form (as H₃BO₃, boric acid) it reacts with sulfur trioxide, the anhydride of sulphuric acid, to form a bisulfate B(HSO₃) ₄.^[274] In its pure (anhydrous) form it reacts with phosphoric acid to form a "phosphate" BPO₄.^[275] The latter compound may be regarded as a mixed oxide of B₂O₃ and P₂O₅.^[276]
28. Organic derivatives of metalloids are traditionally counted as organometallic compounds.^[278]
29. In air, silicon forms a thin coating of amorphous silicon dioxide, 2 to 3 nm thick.^[283] This coating is dissolved by hydrogen fluoride at a very low pace—on the order of two to three hours per nanometre.^[284] Silicon dioxide, and silicate glasses (of which silicon dioxide is a major component), are otherwise readily attacked by hydrofluoric acid.^[285]
30. 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,^[301] and with hydrofluoric acid to give hexafluorosilicic acid H₂SiF₆.^[302] The latter reaction "is sometimes quoted as evidence of basic [that is, metallic] properties".^[303]
31. Temperatures above 400 °C are required to form a noticeable surface oxide layer.^[307]
32. Sources mentioning germanium cations include: Powell & Brewer^[313] who state that the cadmium iodide CdI₂ structure of germanous iodide GeI₂ establishes the existence of the Ge⁺⁺ ion (the CdI₂ structure being found, according to Ladd,^[314] in "many metallic halides, hydroxides and chalcides"); Everest^[315] who comments that, "it seems probable that the Ge⁺⁺ ion can also occur in other crystalline germanous salts such as the phosphite, which is similar to the salt-like stannous phosphite and germanous phosphate, which resembles not only the stannous phosphates, but the manganous phosphates also"; Pan, Fu & Huang^[316] who presume the formation of the simple Ge⁺⁺ ion when Ge(OH)₂ is dissolved in a perchloric acid solution, on the basis that, "ClO4⁻ has little tendency to enter complex formation with a cation"; Monconduit et al.^[317] who prepared the layer compound or phase Nb₃Ge_xTe₆ (x ≃ 0.9), and reported that this contained a Ge^II cation; Richens^[318] who records that, "Ge²⁺ (aq) or possibly Ge(OH)⁺(aq) is said to exist in dilute air-free aqueous suspensions of the yellow hydrous monoxide…however both are unstable with respect to the ready formation of GeO₂.nH₂O"; Rupar et al.^[319] who synthesized a cryptand compound containing a Ge²⁺ cation; and Schwietzer and Pesterfield^[320] who write that, "the monoxide GeO dissolves in dilute acids to give Ge⁺² and in dilute bases to produce GeO₂⁻², all three entities being unstable in water". Sources dismissing germanium cations or further qualifying their presumed existence include: Jolly and Latimer^[321] who assert that, "the germanous ion cannot be studied directly because no germanium (II) species exists in any appreciable concentration in noncomplexing aqueous solutions"; Lidin^[322] who says that, "[germanium] forms no aquacations"; Ladd^[323] who notes that the CdI₂ structure is "intermediate in type between ionic and molecular compounds"; and Wiberg^[324] who states that, "no germanium cations are known".
33. 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.^[340]
34. Sources mentioning cationic arsenic include: Gillespie & Robinson^[343] who find that, "in very dilute solutions in 100% sulphuric acid, arsenic (III) oxide forms arsonyl (III) hydrogen sulphate, AsO.HO₄, which is partly ionized to give the AsO⁺ cation. Both these species probably exist mainly in solvated forms, e.g., As(OH)(SO₄H)₂, and As(OH)(SO₄H)⁺ respectively"; Paul et al.^[344] who reported spectroscopic evidence for the presence of As₄²⁺ and As₂²⁺ cations when arsenic was oxidized with peroxydisulfuryl difluoride S₂O₆F₂ in highly acidic media (Gillespie and Passmore^[345] noted the spectra of these species were very similar to S₄²⁺ and S₈²⁺ and concluded that, "at present" there was no reliable evidence for any homopolycations of arsenic); Van Muylder and Pourbaix,^[346] who write that, "As₂O₃ is an amphoteric oxide which dissolves in water and in solutions of pH between 1 and 8 with the formation of undissociated arsenious acid HAsO₂; the solubility…increases at pH’s below 1 with the formation of 'arsenyl' ions AsO⁺…"; Kolthoff and Elving^[347] who write that, "the As³⁺ cation exists to some extent only in strongly acid solutions; under less acid conditions the tendency is toward hydrolysis, so that the anionic form predominates"; Moody^[348] who observes that, "arsenic trioxide, As₄O₆, and arsenious acid, H₃AsO₃, are apparently amphoteric but no cations, As³⁺, As(OH)²⁺ or As(OH)₂⁺ are known"; and Cotton et al.^[349] who write that (in aqueous solution) the simple arsenic cation As³⁺ "may occur to some slight extent [along with the AsO⁺ cation]" and that, "Raman spectra show that in acid solutions of As₄O₆ the only detectable species is the pyramidal As(OH)₃".
35. The formulae of AsPO₄ and As₂(SO₄)₃ suggest straightforward ionic formulations, with As³⁺, but this is not the case. 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.^[355] As₂(SO₄)₃ has a structure in which each SO₄ tetrahedron is bridged by two AsO₃ trigonal pyramida.^[356]
36. As₂O₃ is usually regarded as being amphoteric but a few sources say it is (weakly)^[359] acidic. They describe its "basic" properties (its reaction with concentrated hydrochloric acid to form arsenic trichloride) as being alcoholic, in analogy with the formation of covalent alkyl chlorides by covalent alcohols (e.g., R-OH + HCl → RCl + H₂O)^[360]
37. Antimony can also be prepared in an amorphous semiconducting black form, with an estimated (temperature-dependent) band gap of 0.06–0.18 eV.^[366]
38. Lidin^[371] asserts that SbO⁺ does not exist and that the stable form of Sb(III) in aqueous solution is an incomplete hydrocomplex [Sb(H₂O)₄(OH)₂]⁺.
39. Cotton et al.^[395] note that TeO₂ appears to have an ionic lattice; Wells^[396] suggests that the Te–O bonds have "considerable covalent character".
40. Liquid carbon may^[410] or may not^[411] be a metallic conductor, depending on pressure and temperature; see also.^[412]
41. 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.
42. 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.^[422]
43. A mnemonic that captures the elements commonly recognised as metalloids goes: Up, up-down, up-down, up ... are the metalloids!^[434]
44. Rochow,^[440] who later wrote his 1966 monograph The metalloids,^[441] commented that, "In some respects selenium acts like a metalloid and tellurium certainly does".
45. A further option is to include astatine both as a nonmetal and as a metalloid.^[467]
46. A visible piece of astatine would be immediately and completely vaporized because of the heat generated by its intense radioactivity.^[473]
47. The literature is contradictory as to whether boron exhibits metallic conductivity in liquid form. Krishnan et al.^[475] found that liquid boron behaved like a metal. Glorieux et al.^[476] characterised liquid boron as a semiconductor, on the basis of its low electrical conductivity. Millot et al.^[477] reported that the emissivity of liquid boron was not consistent with that of a liquid metal.
48. Korenman^[481] similarly noted that "the ability to precipitate with hydrogen sulfide distinguishes astatine from other halogens and brings it closer to bismuth and other heavy metals".
49. 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).^[493] 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.^[494]
50. For example: intermediate electrical conductivity;^[506] a relatively narrow band gap;^[507] light sensitivity.^[506]
51. White phosphorus is the least stable and most reactive form.^[508] It is also the most common, industrially important,^[509] and easily reproducible allotrope, and for these three reasons is regarded as the standard state of the element.^[510]
52. Sample prices of gold, in comparison, start at roughly thirty-five times that of silver. Based on sample prices for B, C, Al, Si, Ge, As, Se, Ag, Sb, Te and Au available on-line from Alfa Aesa; Goodfellow; Metallium; and United Nuclear Scientific.
53. Based on spot prices for Al, Si, Ge, As, Sb, Se, and Te available on-line from FastMarkets: Minor Metals; Fast Markets: Base Metals; EnergyTrend: PV Market Status, Polysilicon; and Metal-Pages: Arsenic metal prices, news and information.

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2. Atkins et al. 2010, p. 20
3. Cusack 1987, p. 360
4. Kelter, Mosher & Scott 2009, p. 268
5. Hill & Holman 2000, p. 41
6. King 1979, p. 13
7. Moore 2011, p. 81
8. Gray 2010
9. Hopkins & Bailar 1956, p. 458
10. Glinka 1965, p. 77
11. Wiberg 2001, p. 1279
12. Belpassi et al. 2006, pp. 4543–4
13. Schmidbaur & Schier 2008, pp. 1931–51
14. Tyler Miller 1987, p. 59
15. Goldsmith 1982, p. 526; Kotz, Treichel & Weaver 2009, p. 62; Bettelheim et al. 2010, p. 46
16. Mann et al. 2000, p. 2783
17. Hawkes 2001, p. 1686; Segal 1989, p. 965; McMurray & Fay 2009, p. 767
18. Bucat 1983, p. 26; Brown c. 2007
19. Swift & Schaefer 1962, p. 100
20. Hawkes 2001, p. 1686; Hawkes 2010; Holt, Rinehart & Wilson c. 2007
21. 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).
22. Tilden 1876, pp. 172, 198–201; Smith 1994, p. 252; Bodner & Pardue 1993, p. 354
23. Bassett et al. 1966, p. 127
24. Rausch 1960
25. Thayer 1977, p. 604; Warren & Geballe 1981; Masters & Ela 2008, p. 190
26. Warren & Geballe 1981; Chalmers 1959, p. 72; US Bureau of Naval Personnel 1965, p. 26
27. Siebring 1967, p. 513
28. Wiberg 2001, p. 282
29. Rausch 1960; Friend 1953, p. 68
30. Murray 1928, p. 1295
31. Hampel & Hawley 1966, p. 950; Stein 1985; Stein 1987, pp. 240, 247–8
32. Hatcher 1949, p. 223; Secrist & Powers 1966, p. 459
33. Taylor 1960, p. 614
34. Considine & Considine 1984, p. 568; Cegielski 1998, p. 147; The American heritage science dictionary 2005 p. 397
35. Woodward 1948, p. 1
36. NIST 2010. Values shown in the above table have been converted from the NIST values, which are given in eV.
37. Berger 1997; Lovett 1977, p. 3
38. Goldsmith 1982, p. 526; Hawkes 2001, p. 1686
39. Hawkes 2001, p. 1687
40. Sharp 1981, p. 299
41. Emsley 1971, p. 1
42. James et al. 2000, p. 480
43. Chatt 1951, p. 417 "The boundary between metals and metalloids is indefinite ..."; Burrows et al. 2009, p. 1192: "Although the elements are conveniently described as metals, metalloids, and nonmetals, the transitions are not exact ..."
44. Jones 2010, p. 170
45. Kneen, Rogers & Simpson 1972, pp. 218–220
46. Rochow 1966, pp. 1, 4–7
47. Rochow 1977, p. 76; Mann et al. 2000, p. 2783
48. Askeland, Phulé & Wright 2011, p. 69
49. 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)
50. Russell & Lee 2005, p. 1
51. Russell & Lee 2005, pp. 6–7, 387
52. Pearson 1972, p. 264
53. Okajima & Shomoji 1972, p. 258
54. Kitaĭgorodskiĭ 1961, p. 108
55. Neuburger 1936
56. Edwards & Sienko 1983, p. 693
57. Herzfeld 1927; Edwards 2000, pp. 100–3
58. Edwards & Sienko 1983, p. 695; Edwards et al. 2010
59. Edwards 1999, p. 416
60. Steurer 2007, p. 142; Pyykkö 2012, p. 56
61. Edwards & Sienko 1983, p. 695
62. Hill & Holman 2000, p. 41. They 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".
63. Bond 2005, p. 3: "One criterion for distinguishing semi-metals from true metals under normal conditions is that the bulk coordination number of the former is never greater than eight, while for metals it is usually twelve (or more, if for the body-centred cubic structure one counts next-nearest neighbours as well)."
64. Jones 2010, p. 169
65. 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.
66. Kraig, Roundy & Cohen 2004, p. 412; Alloul 2010, p. 83
67. Vernon 2013, pp. 1704
68. Hamm 1969, p. 653
69. Horvath 1973, p. 336
70. Gray 2009, p.  9
71. Rayner-Canham 2011
72. Booth & Bloom 1972, p. 426; Cox 2004, pp. 17, 18, 27–8; Silberberg 2006, pp. 305–13
73. Cox 2004, pp. 17–18, 27–8; Silberberg 2006, p. 305–13
74. Rodgers 2011, pp. 232–3; 240–1
75. Roher 2001, pp. 4–6
76. Tyler 1948, p. 105; Reilly 2002, pp. 5–6
77. Hampel & Hawley 1976, p. 174
78. Goodrich 1844, p. 264; The Chemical News 1897, p. 189; Hampel & Hawley 1976, p. 191; Lewis 1993, p. 835; Hérold 2006, pp. 149–50
79. Oderberg 2007, p. 97
80. Brown & Holme 2006, p. 57
81. Wiberg 2001, p. 282; Simple Memory Art c. 2005
82. Chedd 1969, pp. 12–13
83. Kneen, Rogers & Simpson, 1972, p. 263. Columns 2 and 4 are sourced from this reference unless otherwise indicated.
84. Stoker 2010, p. 62; Chang 2002, p. 304. Chang speculates that the melting point of francium would be about 23 °C.
85. New Scientist 1975; Soverna 2004; Eichler et al. 2007; Austen 2012
86. Rochow 1966, p. 4
87. Hunt 2000, p. 256
88. McQuarrie & Rock 1987, p. 85
89. Desai, James & Ho 1984, p. 1160; Matula 1979, p. 1260
90. Choppin & Johnsen 1972, p. 351
91. Schaefer 1968, p. 76; Carapella 1968, p. 30
92. Kozyrev 1959, p. 104; Chizhikov & Shchastlivyi 1968, p. 25; Glazov, Chizhevskaya & Glagoleva 1969, p. 86
93. Bogoroditskii & Pasynkov 1967, p. 77; Jenkins & Kawamura 1976, p. 88
94. Hampel & Hawley 1976, p. 191; Wulfsberg 2000, p. 620
95. Swalin 1962, p. 216
96. Bailar et al. 1989, p. 742
97. Metcalfe, Williams & Castka 1974, p. 86
98. Chang 2002, p. 306
99. Pauling 1988, p. 183
100. Chedd 1969, pp. 24–5
101. Adler 1969, pp. 18–19
102. Hultgren 1966, p. 648; Young & Sessine 2000, p. 849; Bassett et al. 1966, p. 602
103. Rochow 1966, p. 4; Atkins et al. 2006, pp. 8, 122–3
104. Russell & Lee 2005, pp. 421, 423; Gray 2009, p. 23
105. Olmsted & Williams 1997, p. 975
106. Russell & Lee 2005, p. 401; Büchel, Moretto & Woditsch 2003, p. 278
107. Desch 1914, p. 86
108. Phillips & Williams 1965, p. 620
109. Van der Put 1998, p. 123
110. Klug & Brasted 1958, p. 199
111. Good et al. 1813
112. Sequeira 2011, p. 776
113. Gary 2013
114. Russell & Lee 2005, pp. 423–4; 405–6
115. Davidson & Lakin 1973, p. 627
116. Wiberg 2001, p. 589
117. Greenwood & Earnshaw 2002, p. 749; Schwartz 2002, p. 679
118. Antman 2001
119. Řezanka & Sigler 2008; Sekhon 2012
120. Emsley 2001, p. 67
121. Zhang et al. 2008, p. 360
122. Science Learning Hub 2009
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125. Emsley 2001, p. 391
126. Schauss 1991; Tao & Bolger 1997
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128. Nielsen 1998
129. MacKenzie 2015, p. 36
130. Jaouen & Gibaud 2010
131. Smith et al. 2014
132. Stevens & Klarner, p. 205
133. Sneader 2005, pp. 57–59
134. Keall, Martin and Tunbridge 1946
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163. Deming 1925, pp. 330 (As₂O₃), 418 (B₂O₃; SiO₂; Sb₂O₃); Witt & Gatos 1968, p. 242 (GeO₂)
164. Eagleson 1994, p. 421 (GeO₂); Rothenberg 1976, 56, 118–19 (TeO₂)
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167. Russell & Lee 2005, p. 397
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190. Conkling & Mocella 2010, pp. 181, 213
191. Ellern 1968, pp. 209–10; 322
192. Russell 2009, pp. 15, 17, 41, 79–80
193. Ellern 1968, p. 324
194. Ellern 1968, p. 328
195. Conkling & Mocella 2010, p. 171
196. Conkling & Mocella 2011, pp. 83–4
197. Berger 1997, p. 91; Hampel 1968, passim
198. Rochow 1966, p. 41; Berger 1997, pp. 42–3
199. Bomgardner 2013, p. 20
200. Russell & Lee 2005, p. 395; Brown et al. 2009, p. 489
201. 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. Pauli and I. Rabi, to comment derogatorily on the 'Physics of Dirt'."; Hoddeson 2007, pp. 25–34 (29)
202. Bianco et. al. 2013
203. University of Limerick 2014; Kennedy et al. 2014
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205. Russell & Lee 2005, pp. 421–2, 424
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208. ScienceDaily 2012
209. Reardon 2005; Meskers, Hagelüken & Van Damme 2009, p. 1131
210. The Economist 2012
211. Whitten 2007, p. 488
212. Jaskula 2013
213. German Energy Society 2008, p. 43–44
214. Patel 2012, p. 248
215. Moore 2104; University of Utah 2014; Xu et al. 2014
216. Yang et al. 2012, p. 614
217. Moore 2010, p. 195
218. Moore 2011
219. Liu 2014
220. Bradley 2014; University of Utah 2014
221. Oxford English Dictionary 1989, 'metalloid'; Gordh, Gordh & Headrick 2003, p. 753
222. Foster 1936, pp. 212–13; Brownlee et al. 1943, p. 293
223. Calderazzo, Ercoli & Natta 1968, p. 257
224. Klemm 1950, pp. 133–42; Reilly 2004, p. 4
225. Walters 1982, pp. 32–3
226. Tyler 1948, p. 105
227. Foster & Wrigley 1958, p. 218: "The elements may be grouped into two classes: those that are metals and those that are nonmetals. There is also an intermediate group known variously as metalloids, meta-metals, semiconductors, or semimetals."
228. Slade 2006, p. 16
229. Corwin 2005, p. 80
230. Barsanov & Ginzburg 1974, p. 330
231. Bradbury et al. 1957, pp. 157, 659
232. Miller, Lee & Choe 2002, p. 21
233. King 2004, pp. 196–8; Ferro & Saccone 2008, p. 233
234. Pashaey & Seleznev 1973, p. 565; Gladyshev & Kovaleva 1998, p. 1445; Eason 2007, p. 294
235. Johansen & Mackintosh 1970, pp. 121–4; Divakar, Mohan & Singh 1984, p. 2337; Dávila et al. 2002, p. 035411-3
236. Jezequel & Thomas 1997, pp. 6620–6
237. Hindman 1968, p. 434: "The high values obtained for the [electrical] resistivity indicate that the metallic properties of neptunium are closer to the semimetals than the true metals. This is also true for other metals in the actinide series."; Dunlap et al. 1970, pp. 44, 46: "... α-Np is a semimetal, in which covalency effects are believed to also be of importance ... For a semimetal having strong covalent bonding, like α-Np ..."
238. Lister 1965, p. 54
239. Cotton et al. 1999, p. 502
240. Pinkerton 1800, p. 81
241. Goldsmith 1982, p. 526
242. Zhdanov 1965, pp. 74–5
243. Friend 1953, p. 68; IUPAC 1959, p. 10; IUPAC 1971, p. 11
244. IUPAC 2005; IUPAC 2006–
245. Van Setten et al. 2007, pp. 2460–1; Oganov et al. 2009, pp. 863–4
246. Housecroft & Sharpe 2008, p. 331; Oganov 2010, p. 212
247. Housecroft & Sharpe 2008, p. 333
248. Kross 2011
249. Berger 1997, p. 37
250. Greenwood & Earnshaw 2002, p. 144
251. Kopp, Lipták & Eren 2003, p. 221
252. Prudenziati 1977, p. 242
253. Berger 1997, pp. 87, 84
254. Rayner-Canham & Overton 2006, p. 291
255. Siekierski & Burgess 2002, p. 63
256. Wogan 2014
257. Siekierski & Burgess 2002, p. 86
258. Greenwood & Earnshaw 2002, p. 141; Henderson 2000, p. 58; Housecroft & Sharpe 2008, pp. 360–72
259. Parry et al. 1970, pp. 438, 448–51
260. Fehlner 1990, p. 202
261. Owen & Brooker 1991, p. 59; Wiberg 2001, p. 936
262. Greenwood & Earnshaw 2002, p. 145
263. Houghton 1979, p. 59
264. Fehlner 1990, pp. 205
265. Fehlner 1990, pp. 204–205, 207
266. Greenwood 2001, p. 2057
267. Salentine 1987, pp. 128–32; MacKay, MacKay & Henderson 2002, pp. 439–40; Kneen, Rogers & Simpson 1972, p. 394; Hiller & Herber 1960, inside front cover; p. 225
268. Sharp 1983, p. 56
269. Fokwa 2014, p. 10
270. Puddephatt & Monaghan 1989, p. 59
271. Mahan 1965, p. 485
272. Danaith 2008, p. 81.
273. Lidin 1996, p. 28
274. Kondrat'ev & Mel'nikova 1978
275. Holderness & Berry 1979, p. 111; Wiberg 2001, p. 980
276. Toy 1975, p. 506
277. Rao 2002, p. 22
278. Fehlner 1992, p. 1
279. Haiduc & Zuckerman 1985, p. 82
280. Greenwood & Earnshaw 2002, p. 331
281. Wiberg 2001, p. 824
282. Rochow 1973, p. 1337‒38
283. Russell & Lee 2005, p. 393
284. Zhang 2002, p. 70
285. Sacks 1998, p. 287
286. Rochow 1973, p. 1337, 1340
287. Allen & Ordway 1968, p. 152
288. Eagleson 1994, pp. 48, 127, 438, 1194; Massey 2000, p. 191
289. Orton 2004, p. 7. This is a typical value for high-purity silicon.
290. Coles & Caplin 1976, p. 106
291. Glazov, Chizhevskaya & Glagoleva 1969, pp. 59–63; Allen & Broughton 1987, p. 4967
292. Cotton, Wilkinson & Gaus 1995, p. 393
293. Wiberg 2001, p. 834
294. Partington 1944, p. 723
295. Cox 2004, p. 27
296. Hiller & Herber 1960, inside front cover; p. 225
297. Kneen, Rogers and Simpson 1972, p. 384
298. Bailar, Moeller & Kleinberg 1965, p. 513
299. Cotton, Wilkinson & Gaus 1995, pp. 319, 321
300. Smith 1990, p. 175
301. Poojary, Borade & Clearfield 1993
302. Wiberg 2001, pp. 851, 858
303. Barmett & Wilson 1959, p. 332
304. Powell 1988, p. 1
305. Greenwood & Earnshaw 2002, p. 371
306. Cusack 1967, p. 193
307. Russell & Lee 2005, pp. 399–400
308. Greenwood & Earnshaw 2002, p. 373
309. Moody 1991, p. 273
310. Russell & Lee 2005, p. 399
311. Berger 1997, pp. 71–2
312. Jolly 1966, pp. 125–6
313. Powell & Brewer 1938
314. Ladd 1999, p. 55
315. Everest 1953, p. 4120
316. Pan, Fu and Huang 1964, p. 182
317. Monconduit et al. 1992
318. Richens 1997, p. 152
319. Rupar et al. 2008
320. Schwietzer & Pesterfield 2010, pp. 190
321. Jolly & Latimer, p. 2
322. Lidin 1996, p. 140
323. Ladd 1999, p. 56
324. Wiberg 2001, p. 896
325. Schwartz 2002, p. 269
326. Eggins 1972, p. 66; Wiberg 2001, p. 895
327. Greenwood & Earnshaw 2002, p. 383
328. Glockling 1969, p. 38; Wells 1984, p. 1175
329. Cooper 1968, pp. 28–9
330. Steele 1966, pp. 178, 188–9
331. Haller 2006, p. 3
332. See, for example, Walker & Tarn 1990, p. 590
333. Wiberg 2001, p. 742
334. Gray, Whitby & Mann 2011
335. Greenwood & Earnshaw 2002, p. 552
336. Parkes & Mellor 1943, p. 740
337. Russell & Lee 2005, p. 420
338. Carapella 1968, p. 30
339. Barfuß et al. 1981, p. 967
340. Greaves, Knights & Davis 1974, p. 369; Madelung 2004, pp. 405, 410
341. Bailar & Trotman-Dickenson 1973, p. 558; Li 1990
342. Bailar, Moeller & Kleinberg 1965, p. 477
343. Gillespie & Robinson 1963, p. 450
344. Paul et al. 1971; see also Ahmeda & Rucka 2011, pp. 2893, 2894
345. Gillespie & Passmore 1972, p. 478
346. Van Muylder & Pourbaix 1974, p. 521
347. Kolthoff & Elving 1978, p. 210
348. Moody 1991, p. 248–249
349. Cotton & Wilkinson 1999, pp. 396, 419
350. Eagleson 1994, p. 91
351. Massey 2000, p. 267
352. Timm 1944, p. 454
353. Partington 1944, p. 641; Kleinberg, Argersinger & Griswold 1960, p. 419
354. Morgan 1906, p. 163; Moeller 1954, p. 559
355. Corbridge 2013, pp. 122, 215
356. Douglade 1982
357. Zingaro 1994, p. 197; Emeléus & Sharpe 1959, p. 418; Addison & Sowerby 1972, p. 209; Mellor 1964, p. 337
358. Pourbaix 1974, p. 521; Eagleson 1994, p. 92; Greenwood & Earnshaw 2002, p. 572
359. Wiberg 2001, pp. 750, 975; Silberberg 2006, p. 314
360. Sidgwick 1950, p. 784; Moody 1991, pp. 248–9, 319
361. Krannich & Watkins 2006
362. Greenwood & Earnshaw 2002, p. 553
363. Dunstan 1968, p. 433
364. Parise 1996, p. 112
365. Carapella 1968a, p. 23
366. Moss 1952, pp. 174, 179
367. Dupree, Kirby & Freyland 1982, p. 604; Mhiaoui, Sar, & Gasser 2003
368. Kotz, Treichel & Weaver 2009, p. 62
369. Cotton et al. 1999, p. 396
370. King 1994, p. 174
371. Lidin 1996, p. 372
372. Lindsjö, Fischer & Kloo 2004
373. Friend 1953, p. 87
374. Fesquet 1872, pp. 109–14
375. Greenwood & Earnshaw 2002, p. 553; Massey 2000, p. 269
376. King 1994, p.171
377. Turova 2011, p. 46
378. Pourbaix 1974, p. 530
379. Wiberg 2001, p. 764
380. House 2008, p. 497
381. Mendeléeff 1897, p. 274
382. Emsley 2001, p. 428
383. Kudryavtsev 1974, p. 78
384. Bagnall 1966, pp. 32–3, 59, 137
385. Swink et al. 1966; Anderson et al. 1980
386. Ahmed, Fjellvåg & Kjekshus 2000
387. Chizhikov & Shchastlivyi 1970, p. 28
388. Kudryavtsev 1974, p. 77
389. Stuke 1974, p. 178; Donohue 1982, pp. 386–7; Cotton et al. 1999, p. 501
390. Becker, Johnson & Nussbaum 1971, p. 56
391. Berger 1997, p. 90
392. Chizhikov & Shchastlivyi 1970, p. 16
393. Jolly 1966, pp. 66–7
394. Schwietzer & Pesterfield 2010, p. 239
395. Cotton et al. 1999, p. 498
396. Wells 1984, p. 715
397. Wiberg 2001, p. 588
398. Mellor 1964a, p.  30; Wiberg 2001, p. 589
399. Greenwood & Earnshaw 2002, p. 765–6
400. Bagnall 1966, p. 134–51; Greenwood & Earnshaw 2002, p. 786
401. Detty & O'Regan 1994, pp. 1–2
402. Hill & Holman 2000, p. 124
403. Chang 2002, p. 314
404. Kent 1950, pp. 1–2; Clark 1960, p. 588; Warren & Geballe 1981
405. Housecroft & Sharpe 2008, p. 384; IUPAC 2006–, rhombohedral graphite entry
406. Mingos 1998, p. 171
407. Wiberg 2001, p. 781
408. Charlier, Gonze & Michenaud 1994
409. Atkins et al. 2006, pp. 320–1
410. Savvatimskiy 2005, p. 1138
411. Togaya 2000
412. Savvatimskiy 2009
413. Inagaki 2000, p. 216; Yasuda et al. 2003, pp. 3–11
414. O'Hare 1997, p. 230
415. Traynham 1989, pp. 930–1; Prakash & Schleyer 1997
416. Olmsted & Williams 1997, p. 436
417. Bailar et al. 1989, p. 743
418. Moore et al. 1985
419. House & House 2010, p. 526
420. Wiberg 2001, p. 798
421. Eagleson 1994, p. 175
422. Atkins et al. 2006, p. 121
423. Russell & Lee 2005, pp. 358–9
424. Keevil 1989, p. 103
425. Russell & Lee 2005, pp. 358–60 et seq
426. Harding, Janes & Johnson 2002, pp. 118
427. Metcalfe, Williams & Castka 1974, p. 539
428. Cobb & Fetterolf 2005, p. 64; Metcalfe, Williams & Castka 1974, p. 539
429. Ogata, Li & Yip 2002; Boyer et al. 2004, p. 1023; Russell & Lee 2005, p. 359
430. Cooper 1968, p. 25; Henderson 2000, p. 5; Silberberg 2006, p. 314
431. Wiberg 2001, p. 1014
432. Daub & Seese 1996, pp. 70, 109: "Aluminum is not a metalloid but a metal because it has mostly metallic properties."; Denniston, Topping & Caret 2004, p. 57: "Note that aluminum (Al) is classified as a metal, not a metalloid."; Hasan 2009, p. 16: "Aluminum does not have the characteristics of a metalloid but rather those of a metal."
433. Holt, Rinehart & Wilson c. 2007
434. Tuthill 2011
435. Stott 1956, p. 100
436. Steele 1966, p. 60
437. Moody 1991, p. 303
438. Emsley 2001, p. 382
439. Young et al. 2010, p. 9; Craig & Maher 2003, p. 391. Selenium is "near metalloidal".
440. Rochow 1957
441. Rochow 1966, p. 224
442. Moss 1952, p. 192
443. Glinka 1965, p. 356
444. Evans 1966, pp. 124–5
445. Regnault 1853, p. 208
446. Scott & Kanda 1962, p. 311
447. Cotton et al. 1999, pp. 496, 503–4
448. Arlman 1939; Bagnall 1966, pp. 135, 142–3
449. Chao & Stenger 1964
450. Berger 1997, pp. 86–7
451. Snyder 1966, p. 242
452. Fritz & Gjerde 2008, p. 235
453. Meyer et al. 2005, p. 284; Manahan 2001, p. 911; Szpunar et al. 2004, p. 17
454. US Environmental Protection Agency 1988, p. 1; Uden 2005, pp. 347‒8
455. De Zuane 1997, p. 93; Dev 2008, pp. 2‒3
456. Wiberg 2001, p. 594
457. Greenwood & Earnshaw 2002, p. 786; Schwietzer & Pesterfield 2010, pp. 242–3
458. Bagnall 1966, p. 41; Nickless 1968, p. 79
459. Bagnall 1990, pp. 313–14; Lehto & Hou 2011, p. 220; Siekierski & Burgess 2002, p. 117: "The tendency to form X²⁻ anions decreases down the Group [16 elements] ..."
460. Legit, Friák & Šob 2010, p. 214118-18
461. Manson & Halford 2006, pp. 378, 410
462. Bagnall 1957, p. 62; Fernelius 1982, p. 741
463. Bagnall 1966, p. 41; Barrett 2003, p. 119
464. Hawkes 2010; Holt, Rinehart & Wilson c. 2007; Hawkes 1999, p. 14; Roza 2009, p. 12
465. Keller 1985
466. Harding, Johnson & Janes 2002, p. 61
467. Long & Hentz 1986, p. 58
468. Vasáros & Berei 1985, p. 109
469. Haissinsky & Coche 1949, p. 400
470. Brownlee et al. 1950, p. 173
471. Hermann, Hoffmann & Ashcroft 2013
472. Siekierski & Burgess 2002, pp. 65, 122
473. Emsley 2001, p. 48
474. Rao & Ganguly 1986
475. Krishnan et al. 1998
476. Glorieux, Saboungi & Enderby 2001
477. Millot et al. 2002
478. Vasáros & Berei 1985, p. 117
479. Kaye & Laby 1973, p. 228
480. Samsonov 1968, p. 590
481. Korenman 1959, p. 1368
482. Rossler 1985, pp. 143–4
483. Champion et al. 2010
484. Borst 1982, pp. 465, 473
485. Batsanov 1971, p. 811
486. Swalin 1962, p. 216; Feng & Lin 2005, p. 157
487. Schwietzer & Pesterfield 2010, pp. 258–60
488. Hawkes 1999, p. 14
489. Olmsted & Williams 1997, p. 328; Daintith 2004, p. 277
490. Eberle1985, pp. 213–16, 222–7
491. Restrepo et al. 2004, p. 69; Restrepo et al. 2006, p. 411
492. Greenwood & Earnshaw 2002, p. 804
493. Greenwood & Earnshaw 2002, p. 803
494. Wiberg 2001, p. 416
495. Craig & Maher 2003, p. 391; Schroers 2013, p. 32; Vernon 2013, pp. 1704–1705
496. Cotton et al. 1999, p. 42
497. Marezio & Licci 2000, p. 11
498. Vernon 2013, p. 1705
499. Russell & Lee 2005, p. 5
500. Parish 1977, pp. 178, 192–3
501. Eggins 1972, p. 66; Rayner-Canham & Overton 2006, pp. 29–30
502. Atkins et al. 2006, pp. 320–1; Bailar et al. 1989, p. 742–3
503. Rochow 1966, p. 7; Taniguchi et al. 1984, p. 867: "... black phosphorus ... [is] characterized by the wide valence bands with rather delocalized nature."; Morita 1986, p. 230; Carmalt & Norman 1998, p. 7: "Phosphorus ... should therefore be expected to have some metalloid properties."; Du et al. 2010. Interlayer interactions in black phosphorus, which are attributed to van der Waals-Keesom forces, are thought to contribute to the smaller band gap of the bulk material (calculated 0.19 eV; observed 0.3 eV) as opposed to the larger band gap of a single layer (calculated ~0.75 eV).
504. Stuke 1974, p. 178; Cotton et al. 1999, p. 501; Craig & Maher 2003, p. 391
505. Steudel 1977, p. 240: "... considerable orbital overlap must exist, to form intermolecular, many-center ... [sigma] bonds, spread through the layer and populated with delocalized electrons, reflected in the properties of iodine (lustre, color, moderate electrical conductivity)."; Segal 1989, p. 481: "Iodine exhibits some metallic properties ..."
506. Lutz et al. 2011, p. 17
507. Yacobi & Holt 1990, p. 10; Wiberg 2001, p. 160
508. Greenwood & Earnshaw 2002, pp. 479, 482
509. Eagleson 1994, p. 820
510. Oxtoby, Gillis & Campion 2008, p. 508
511. Brescia et al. 1980, pp. 166–71
512. Fine & Beall 1990, p. 578
513. Wiberg 2001, p. 901
514. Berger 1997, p. 80
515. Lovett 1977, p. 101
516. Cohen & Chelikowsky 1988, p. 99
517. Taguena-Martinez, Barrio & Chambouleyron 1991, p. 141
518. Ebbing & Gammon 2010, p. 891
519. Asmussen & Reinhard 2002, p. 7
520. Deprez & McLachan 1988
521. Addison 1964 (P, Se, Sn); Marković, Christiansen & Goldman 1998 (Bi); Nagao et al. 2004
522. Lide 2005; Wiberg 2001, p. 423: At
523. Cox 1997, pp. 182‒86
524. MacKay, MacKay & Henderson 2002, p. 204
525. Baudis 2012, pp. 207–8
526. Wiberg 2001, p. 741
527. Chizhikov & Shchastlivyi 1968, p. 96
528. Greenwood & Earnshaw 2002, pp. 140–1, 330, 369, 548–9, 749: B, Si, Ge, As, Sb, Te
529. Kudryavtsev 1974, p. 158
530. Greenwood & Earnshaw 2002, pp. 271, 219, 748–9, 886: C, Al, Se, Po, At; Wiberg 2001, p. 573: Se
531. United Nuclear 2013
532. Zalutsky & Pruszynski 2011, p. 181

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Further reading

• Brady JE, Humiston GE & Heikkinen H 1980, 'Chemistry of the Representative Elements: Part II, The Metalloids and Nonmetals', in General Chemistry: Principles and Structure, 2nd ed., SI version, John Wiley & Sons, New York, pp. 537–591, ISBN 0-471-06315-0
• Chedd G 1969, Half-way Elements: The Technology of Metalloids, Doubleday, New York
• Choppin GR & Johnsen RH 1972, 'Group IV and the Metalloids,' in Introductory Chemistry, Addison-Wesley, Reading, Massachusetts, pp. 341–357
• Dunstan S 1968, 'The Metalloids', in Principles of Chemistry, D. Van Nostrand Company, London, pp. 407–39
• Goldsmith RH 1982, 'Metalloids', Journal of Chemical Education, vol. 59, no. 6, pp. 526–527, doi:10.1021/ed059p526
• Hawkes SJ 2001, 'Semimetallicity', Journal of Chemical Education, vol. 78, no. 12, pp. 1686–7, doi:10.1021/ed078p1686
• Metcalfe HC, Williams JE & Castka JF 1974, 'Aluminum and the Metalloids', in Modern Chemistry, Holt, Rinehart and Winston, New York, pp. 538–57, ISBN 0-03-089450-6
• Moeller T, Bailar JC, Kleinberg J, Guss CO, Castellion ME & Metz C 1989, 'Carbon and the Semiconducting Elements', in Chemistry, with Inorganic Qualitative Analysis, 3rd ed., Harcourt Brace Jovanovich, San Diego, pp. 751–75, ISBN 0-15-506492-4
• Rieske M 1998, 'Metalloids', in Encyclopedia of Earth and Physical Sciences, Marshall Cavendish, New York, vol. 6, pp. 758–9, ISBN 0-7614-0551-8 (set)
• Rochow EG 1966, The Metalloids, DC Heath and Company, Boston
• Vernon RE 2013, 'Which Elements are Metalloids?', Journal of Chemical Education, vol. 90, no. 12, pp. 1703–7, doi:10.1021/ed3008457