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Definitions

See also: Lists of metalloids

Judgment-based

A metalloid is an element that possesses a preponderance of properties in between, or that are a mixture of, those of 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.⁴ Difficulty of categorisation is a key attribute. Most elements have a mixture of metallic and nonmetallic properties,⁵ and can be classified according to which set of properties is more pronounced.⁶⁷ Only the elements at or near the margins, lacking a sufficiently clear preponderance of either metallic or nonmetallic properties, are classified as metalloids.⁸

Boron, silicon, germanium, arsenic, antimony, and tellurium are commonly recognised as metalloids.⁹¹⁰ Depending on the author, one or more from selenium, polonium, or astatine are sometimes added to the list.¹¹ Boron sometimes is excluded, by itself, or with silicon.¹² Sometimes tellurium is not regarded as a metalloid.¹³ The inclusion of antimony, polonium, and astatine as metalloids has been questioned.¹⁴

Other elements are occasionally classified as metalloids. These elements include¹⁵ hydrogen,¹⁶ beryllium,¹⁷ nitrogen,¹⁸ phosphorus,¹⁹ sulfur,²⁰ zinc,²¹ gallium,²² tin, iodine,²³ lead,²⁴ bismuth,²⁵ and radon.²⁶ The term metalloid has also been used for elements that exhibit metallic lustre and electrical conductivity, and that are amphoteric, such as arsenic, antimony, vanadium, chromium, molybdenum, tungsten, tin, lead, and aluminium.²⁷ The p-block metals,²⁸ and nonmetals (such as carbon or nitrogen) that can form alloys with metals²⁹ or modify their properties³⁰ have also occasionally been considered as metalloids.

Criteria-based

No widely accepted definition of a metalloid exists, nor any division of the periodic table into metals, metalloids, and nonmetals;³³ Hawkes³⁴ 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 Sharp³⁵ as "arbitrary".

The number and identities of metalloids depend on what classification criteria are used. Emsley³⁶ recognised four metalloids (germanium, arsenic, antimony, and tellurium); James et al.³⁷ listed twelve (Emsley's plus boron, carbon, silicon, selenium, bismuth, polonium, moscovium, and livermorium). On average, seven elements are included in such lists; individual classification arrangements tend to share common ground and vary in the ill-defined³⁸ margins.³⁹⁴⁰

A single quantitative criterion such as electronegativity is commonly used,⁴¹ metalloids having electronegativity values from 1.8 or 1.9 to 2.2.⁴² Further examples include packing efficiency (the fraction of volume in a crystal structure occupied by atoms) and the Goldhammer–Herzfeld criterion ratio.⁴³ The commonly recognised metalloids have packing efficiencies of between 34% and 41%.⁴⁴ The Goldhammer–Herzfeld ratio, roughly equal to the cube of the atomic radius divided by the molar volume,⁴⁵⁴⁶ is a simple measure of how metallic an element is, the recognised metalloids having ratios from around 0.85 to 1.1 and averaging 1.0.⁴⁷⁴⁸ Other authors have relied on, for example, atomic conductance⁴⁹⁵⁰ or bulk coordination number.⁵¹

Jones, writing on the role of classification in science, observed that "[classes] are usually defined by more than two attributes".⁵² Masterton and Slowinski⁵³ 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.⁵⁴

In this context, Vernon proposed that a metalloid is a chemical element that, in its standard state, has (a) the electronic band structure of a semiconductor or a semimetal; and (b) an intermediate first ionization potential "(say 750−1,000 kJ/mol)"; and (c) an intermediate electronegativity (1.9–2.2).⁵⁵

Periodic table territory

Location

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.⁵⁶ 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.⁵⁷

The diagonal positioning of the metalloids represents an exception to the observation that elements with similar properties tend to occur in vertical groups.⁵⁸ 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-Canham⁵⁹ 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.⁶⁰ 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.⁶¹ The net effect is that the location of the metal–nonmetal transition zone shifts to the right in going down a group,⁶² and analogous diagonal similarities are seen elsewhere in the periodic table, as noted.⁶³

Alternative treatments

Elements bordering the metal–nonmetal dividing line are not always classified as metalloids, noting a binary classification can facilitate the establishment of rules for determining bond types between metals and nonmetals.⁶⁴ In such cases, the authors concerned focus on one or more attributes of interest to make their classification decisions, rather than being concerned about the marginal nature of the elements in question. Their considerations may or not be made explicit and may, at times, seem arbitrary.⁶⁵⁶⁶ Metalloids may be grouped with metals;⁶⁷ or regarded as nonmetals;⁶⁸ or treated as a sub-category of nonmetals.⁶⁹⁷⁰ Other authors have suggested classifying some elements as metalloids "emphasizes that properties change gradually rather than abruptly as one moves across or down the periodic table".⁷¹ Some periodic tables distinguish elements that are metalloids and display no formal dividing line between metals and nonmetals. Metalloids are instead shown as occurring in a diagonal band⁷² or diffuse region.⁷³ The key consideration is to explain the context for the taxonomy in use.

Properties

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

Compared to metals and nonmetals

Main article: Properties of metals, metalloids and nonmetals

Characteristic properties of metals, metalloids, and nonmetals are summarized in the table.⁷⁴ Physical properties are listed in order of ease of determination; chemical properties run from general to specific, and then to descriptive.

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

Metalloids are too brittle to have any structural uses in their pure forms.⁹⁵ They and their compounds are used in alloys, biological agents (toxicological, nutritional, and medicinal), catalysts, flame retardants, glasses (oxide and metallic), optical storage media and optoelectronics, pyrotechnics, semiconductors, and electronics.⁹⁶

Alloys

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.⁹⁷ Phillips and Williams⁹⁸ 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 MnB, if n > 2.⁹⁹ 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.¹⁰⁰

The heavier metalloids continue the theme. Arsenic can form alloys with metals, including platinum and copper;¹⁰¹ it is also added to copper and its alloys to improve corrosion resistance¹⁰² and appears to confer the same benefit when added to magnesium.¹⁰³ 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).¹⁰⁴ Tellurium readily alloys with iron, as ferrotellurium (50–58% tellurium), and with copper, in the form of copper tellurium (40–50% tellurium).¹⁰⁵ Ferrotellurium is used as a stabilizer for carbon in steel casting.¹⁰⁶ 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.¹⁰⁷

Biological agents

All six of the elements commonly recognised as metalloids have toxic, dietary or medicinal properties.¹⁰⁸ 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 insecticides¹⁰⁹ and herbicides.¹¹⁰ It is an essential trace element.¹¹¹ As boric acid, it has antiseptic, antifungal, and antiviral properties.¹¹²

Silicon is present in silatrane, a highly toxic rodenticide.¹¹³ Long-term inhalation of silica dust causes silicosis, a fatal disease of the lungs. Silicon is an essential trace element.¹¹⁴ Silicone gel can be applied to badly burned patients to reduce scarring.¹¹⁵

Salts of germanium are potentially harmful to humans and animals if ingested on a prolonged basis.¹¹⁶ There is interest in the pharmacological actions of germanium compounds but no licensed medicine as yet.¹¹⁷

Arsenic is notoriously poisonous and may also be an essential element in ultratrace amounts.¹¹⁸ 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."¹¹⁹ It has been used as a pharmaceutical agent since antiquity, including for the treatment of syphilis before the development of antibiotics.¹²⁰ 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.¹²¹ Arsenic in drinking water, which causes lung and bladder cancer, has been associated with a reduction in breast cancer mortality rates.¹²²

Metallic antimony is relatively non-toxic, but most antimony compounds are poisonous.¹²³ Two antimony compounds, sodium stibogluconate and stibophen, are used as antiparasitical drugs.¹²⁴

Elemental tellurium is not considered particularly toxic; two grams of sodium tellurate, if administered, can be lethal.¹²⁵ People exposed to small amounts of airborne tellurium exude a foul and persistent garlic-like odour.¹²⁶ Tellurium dioxide has been used to treat seborrhoeic dermatitis; other tellurium compounds were used as antimicrobial agents before the development of antibiotics.¹²⁷ In the future, such compounds may need to be substituted for antibiotics that have become ineffective due to bacterial resistance.¹²⁸

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.¹²⁹ 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.¹³⁰ 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.¹³¹ Gallium nitrate is used to treat the side effects of cancer;¹³² gallium citrate, a radiopharmaceutical, facilitates imaging of inflamed body areas.¹³³ Selenium sulfide is used in medicinal shampoos and to treat skin infections such as tinea versicolor.¹³⁴ Iodine is used as a disinfectant in various forms. Bismuth is an ingredient in some antibacterials.¹³⁵

Catalysts

Boron trifluoride and trichloride are used as homogeneous catalysts in organic synthesis and electronics; the tribromide is used in the manufacture of diborane.¹³⁶ Non-toxic boron ligands could replace toxic phosphorus ligands in some transition metal catalysts.¹³⁷ Silica sulfuric acid (SiO2OSO3H) is used in organic reactions.¹³⁸ Germanium dioxide is sometimes used as a catalyst in the production of PET plastic for containers;¹³⁹ cheaper antimony compounds, such as the trioxide or triacetate, are more commonly employed for the same purpose¹⁴⁰ despite concerns about antimony contamination of food and drinks.¹⁴¹ 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.¹⁴² Selenium acts as a catalyst in some microorganisms.¹⁴³ Tellurium, its dioxide, and its tetrachloride are strong catalysts for air oxidation of carbon above 500 °C.¹⁴⁴ Graphite oxide can be used as a catalyst in the synthesis of imines and their derivatives.¹⁴⁵ Activated carbon and alumina have been used as catalysts for the removal of sulfur contaminants from natural gas.¹⁴⁶ Titanium doped aluminium has been suggested as a substitute for noble metal catalysts used in the production of industrial chemicals.¹⁴⁷

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.¹⁴⁸ 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.¹⁴⁹ 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.¹⁵⁰ Antimony trioxide is a flame retardant.¹⁵¹ Aluminium hydroxide has been used as a wood-fibre, rubber, plastic, and textile flame retardant since the 1890s.¹⁵² 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.¹⁵³

Glass formation

The oxides B2O3, SiO2, GeO2, As2O3, and Sb2O3 readily form glasses. TeO2 forms a glass but this requires a "heroic quench rate"¹⁵⁴ or the addition of an impurity; otherwise the crystalline form results.¹⁵⁵ These compounds are used in chemical, domestic, and industrial glassware¹⁵⁶ and optics.¹⁵⁷ Boron trioxide is used as a glass fibre additive,¹⁵⁸ and is also a component of borosilicate glass, widely used for laboratory glassware and domestic ovenware for its low thermal expansion.¹⁵⁹ Most ordinary glassware is made from silicon dioxide.¹⁶⁰ Germanium dioxide is used as a glass fibre additive, as well as in infrared optical systems.¹⁶¹ Arsenic trioxide is used in the glass industry as a decolourizing and fining agent (for the removal of bubbles),¹⁶² as is antimony trioxide.¹⁶³ Tellurium dioxide finds application in laser and nonlinear optics.¹⁶⁴

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.¹⁶⁵¹⁶⁶ Aside from thin films deposited at very low temperatures, the first known metallic glass was an alloy of composition Au75Si25 reported in 1960.¹⁶⁷ A metallic glass having a strength and toughness not previously seen, of composition Pd82.5P6Si9.5Ge2, was reported in 2011.¹⁶⁸

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.¹⁶⁹ Selenium compounds can be used both as decolourising agents and to add a red colour to glass.¹⁷⁰ 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.¹⁷¹ Lead-based glasses have also been extensively used in electronic components, enamelling, sealing and glazing materials, and solar cells. Bismuth based oxide glasses have emerged as a less toxic replacement for lead in many of these applications.¹⁷²

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.¹⁷³ 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."¹⁷⁴

Pyrotechnics

The recognised metalloids have either pyrotechnic applications or associated properties. Boron and silicon are commonly encountered;¹⁷⁵ they act somewhat like metal fuels.¹⁷⁶ Boron is used in pyrotechnic initiator compositions (for igniting other hard-to-start compositions), and in delay compositions that burn at a constant rate.¹⁷⁷ Boron carbide has been identified as a possible replacement for more toxic barium or hexachloroethane mixtures in smoke munitions, signal flares, and fireworks.¹⁷⁸ Silicon, like boron, is a component of initiator and delay mixtures.¹⁷⁹ Doped germanium can act as a variable speed thermite fuel.¹⁸⁰ Arsenic trisulfide As2S3 was used in old naval signal lights; in fireworks to make white stars;¹⁸¹ in yellow smoke screen mixtures; and in initiator compositions.¹⁸² Antimony trisulfide Sb2S3 is found in white-light fireworks and in flash and sound mixtures.¹⁸³ Tellurium has been used in delay mixtures and in blasting cap initiator compositions.¹⁸⁴

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.¹⁸⁵¹⁸⁶ Aluminium is a common pyrotechnic ingredient,¹⁸⁷ and is widely employed for its capacity to generate light and heat,¹⁸⁸ including in thermite mixtures.¹⁸⁹ Phosphorus can be found in smoke and incendiary munitions, paper caps used in toy guns, and party poppers.¹⁹⁰ Selenium has been used in the same way as tellurium.¹⁹¹

Semiconductors and electronics

All the elements commonly recognised as metalloids (or their compounds) have been used in the semiconductor or solid-state electronic industries.¹⁹²

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.¹⁹³

Silicon is the leading commercial semiconductor; it forms the basis of modern electronics (including standard solar cells)¹⁹⁴ and information and communication technologies.¹⁹⁵ 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.¹⁹⁶

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.¹⁹⁷ 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.¹⁹⁸ The synthesis of gram-scale quantities of semiconducting germanane was reported in 2013. This consists of one-atom thick sheets of hydrogen-terminated germanium atoms, analogous to graphane. It conducts electrons more than ten times faster than silicon and five times faster than germanium, and is thought to have potential for optoelectronic and sensing applications.¹⁹⁹ The development of a germanium-wire based anode that more than doubles the capacity of lithium-ion batteries was reported in 2014.²⁰⁰ In the same year, Lee et al. reported that defect-free crystals of graphene large enough to have electronic uses could be grown on, and removed from, a germanium substrate.²⁰¹

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, but they have direct band gaps. These compounds are preferred for optical applications.²⁰² Antimony nanocrystals may enable lithium-ion batteries to be replaced by more powerful sodium ion batteries.²⁰³

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.²⁰⁴ 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.²⁰⁵ Bismuth telluride (Bi2Te3), alloyed with selenium and antimony, is a component of thermoelectric devices used for refrigeration or portable power generation.²⁰⁶

Five metalloids – boron, silicon, germanium, arsenic, and antimony – can be found in cell phones (along with at least 39 other metals and nonmetals).²⁰⁷ Tellurium is expected to find such use.²⁰⁸ 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.²⁰⁹ 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.²¹⁰ Selenium is used in the production of solar cells²¹¹ and in high-energy surge protectors.²¹²

Boron, silicon, germanium, antimony, and tellurium,²¹³ as well as heavier metals and metalloids such as Sm, Hg, Tl, Pb, Bi, and Se,²¹⁴ can be found in topological insulators. These are alloys²¹⁵ or compounds which, at ultracold temperatures or room temperature (depending on their composition), are metallic conductors on their surfaces but insulators through their interiors.²¹⁶ 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.²¹⁷ These two classes of material are thought to have potential quantum computing applications.²¹⁸

Nomenclature and history

Derivation and other names

Several names are sometimes used synonymously although some of these have other meanings that are not necessarily interchangeable: amphoteric element,²¹⁹ boundary element,²²⁰ half-way element,²²¹ near metal,²²² meta-metal,²²³ semiconductor,²²⁴ semimetal²²⁵ and submetal.²²⁶ "Amphoteric element" is sometimes used more broadly to include transition metals capable of forming oxyanions, such as chromium and manganese.²²⁷ "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.²²⁸ These metals tend to have distorted crystalline structures, electrical conductivity values at the lower end of those of metals, and amphoteric (weakly basic) oxides.²²⁹ The names amphoteric element and semiconductor are problematic as some elements referred to as metalloids do not show marked amphoteric behaviour (bismuth, for example)²³⁰ or semiconductivity (polonium)²³¹ in their most stable forms.

Origin and usage

Main article: Origin and use of the term metalloid

The origin and usage of the term metalloid is convoluted. The "Manual of Metalloids" published in 1864 divided all elements into either metals or metalloids.²³²: 31  Earlier usage in mineralogy, to describe a mineral having a metallic appearance, can be sourced to as early as 1800.²³³ Since the mid-20th century it has been used to refer to intermediate or borderline chemical elements.²³⁴ The International Union of Pure and Applied Chemistry (IUPAC) previously recommended abandoning the term metalloid, and suggested using the term semimetal instead.²³⁵ Use of this latter term has more recently been discouraged by Atkins et al.²³⁶ as it has a more common meaning that refers to the electronic band structure of a substance rather than the overall classification of an element. The most recent IUPAC publications on nomenclature and terminology do not include any recommendations on the usage of the terms metalloid or semimetal.²³⁷

Elements commonly recognised as metalloids

Boron

Main article: Boron

Pure boron is a shiny, silver-grey crystalline solid.²³⁸ 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,²³⁹ and has a melting point of 2076 °C (cf. steel ~1370 °C).²⁴⁰ Boron is a semiconductor;²⁴¹ its room temperature electrical conductivity is 1.5 × 10−6 S•cm−1²⁴² (about 200 times less than that of tap water)²⁴³ and it has a band gap of about 1.56 eV.²⁴⁴²⁴⁵ Mendeleev commented that, "Boron appears in a free state in several forms which are intermediate between the metals and the nonmmetals."²⁴⁶

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.²⁴⁷ Metallic bonding is the usual result among the heavier congenors of boron but this generally requires low ionization energies.²⁴⁸ Instead, because of its small size and high ionization energies, the basic structural unit of boron (and nearly all of its allotropes)²⁴⁹ 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.²⁵⁰ The same motif can be seen, as are deltahedral variants or fragments, in metal borides and hydride derivatives, and in some halides.²⁵¹

The bonding in boron has been described as being characteristic of behaviour intermediate between metals and nonmetallic covalent network solids (such as diamond).²⁵² 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.²⁵³

Most of the chemistry of boron is nonmetallic in nature.²⁵⁴ Unlike its heavier congeners, it is not known to form a simple B3+ or hydrated [B(H2O)4]3+ cation.²⁵⁵ The small size of the boron atom enables the preparation of many interstitial alloy-type borides.²⁵⁶ Analogies between boron and transition metals have been noted in the formation of complexes,²⁵⁷ and adducts (for example, BH3 + CO →BH3CO and, similarly, Fe(CO)4 + CO →Fe(CO)5),²⁵⁸ as well as in the geometric and electronic structures of cluster species such as [B6H6]2− and [Ru6(CO)18]2−.²⁵⁹²⁶⁰ The aqueous chemistry of boron is characterised by the formation of many different polyborate anions.²⁶¹ Given its high charge-to-size ratio, boron bonds covalently in nearly all of its compounds;²⁶² the exceptions are the borides as these include, depending on their composition, covalent, ionic, and metallic bonding components.²⁶³²⁶⁴ 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.²⁶⁵ Boron has a strong affinity for oxygen and a duly extensive borate chemistry.²⁶⁶ The oxide B2O3 is polymeric in structure,²⁶⁷ weakly acidic,²⁶⁸²⁶⁹ and a glass former.²⁷⁰ Organometallic compounds of boron²⁷¹ have been known since the 19th century (see organoboron chemistry).²⁷²

Silicon

Main article: Silicon

Silicon is a crystalline solid with a blue-grey metallic lustre.²⁷³ Like boron, it is less dense (at 2.33 g/cm3) than aluminium, and is hard and brittle.²⁷⁴ It is a relatively unreactive element.²⁷⁵ According to Rochow,²⁷⁶ the massive crystalline form (especially if pure) is "remarkably inert to all acids, including hydrofluoric".²⁷⁷ Less pure silicon, and the powdered form, are variously susceptible to attack by strong or heated acids, as well as by steam and fluorine.²⁷⁸ Silicon dissolves in hot aqueous alkalis with the evolution of hydrogen, as do metals²⁷⁹ such as beryllium, aluminium, zinc, gallium or indium.²⁸⁰ It melts at 1414 °C. Silicon is a semiconductor with an electrical conductivity of 10−4 S•cm−1²⁸¹ and a band gap of about 1.11 eV.²⁸² When it melts, silicon becomes a reasonable metal²⁸³ with an electrical conductivity of 1.0–1.3 × 104 S•cm−1, similar to that of liquid mercury.²⁸⁴

The chemistry of silicon is generally nonmetallic (covalent) in nature.²⁸⁵ It is not known to form a cation.²⁸⁶²⁸⁷ Silicon can form alloys with metals such as iron and copper.²⁸⁸ It shows fewer tendencies to anionic behaviour than ordinary nonmetals.²⁸⁹ Its solution chemistry is characterised by the formation of oxyanions.²⁹⁰ The high strength of the silicon–oxygen bond dominates the chemical behaviour of silicon.²⁹¹ Polymeric silicates, built up by tetrahedral SiO4 units sharing their oxygen atoms, are the most abundant and important compounds of silicon.²⁹² The polymeric borates, comprising linked trigonal and tetrahedral BO3 or BO4 units, are built on similar structural principles.²⁹³ The oxide SiO2 is polymeric in structure,²⁹⁴ weakly acidic,²⁹⁵²⁹⁶ and a glass former.²⁹⁷ Traditional organometallic chemistry includes the carbon compounds of silicon (see organosilicon).²⁹⁸

Germanium

Main article: Germanium

Germanium is a shiny grey-white solid.²⁹⁹ It has a density of 5.323 g/cm3 and is hard and brittle.³⁰⁰ It is mostly unreactive at room temperature³⁰¹ but is slowly attacked by hot concentrated sulfuric or nitric acid.³⁰² Germanium also reacts with molten caustic soda to yield sodium germanate Na2GeO3 and hydrogen gas.³⁰³ It melts at 938 °C. Germanium is a semiconductor with an electrical conductivity of around 2 × 10−2 S•cm−1³⁰⁴ and a band gap of 0.67 eV.³⁰⁵ Liquid germanium is a metallic conductor, with an electrical conductivity similar to that of liquid mercury.³⁰⁶

Most of the chemistry of germanium is characteristic of a nonmetal.³⁰⁷ Whether or not germanium forms a cation is unclear, aside from the reported existence of the Ge2+ ion in a few esoteric compounds.³⁰⁸ It can form alloys with metals such as aluminium and gold.³⁰⁹ It shows fewer tendencies to anionic behaviour than ordinary nonmetals.³¹⁰ Its solution chemistry is characterised by the formation of oxyanions.³¹¹ Germanium generally forms tetravalent (IV) compounds, and it can also form less stable divalent (II) compounds, in which it behaves more like a metal.³¹² Germanium analogues of all of the major types of silicates have been prepared.³¹³ 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)3.³¹⁴ The oxide GeO2 is polymeric,³¹⁵ amphoteric,³¹⁶ and a glass former.³¹⁷ 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.³¹⁸ Up to the 1930s germanium was considered to be a poorly conducting metal;³¹⁹ it has occasionally been classified as a metal by later writers.³²⁰ As with all the elements commonly recognised as metalloids, germanium has an established organometallic chemistry (see Organogermanium chemistry).³²¹

Arsenic

Main article: Arsenic

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).³²² 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.³²³ Arsenic sublimes at 615 °C. The vapour is lemon-yellow and smells like garlic.³²⁴ Arsenic only melts under a pressure of 38.6 atm, at 817 °C.³²⁵ It is a semimetal with an electrical conductivity of around 3.9 × 104 S•cm−1³²⁶ and a band overlap of 0.5 eV.³²⁷³²⁸ Liquid arsenic is a semiconductor with a band gap of 0.15 eV.³²⁹

The chemistry of arsenic is predominately nonmetallic.³³⁰ Whether or not arsenic forms a cation is unclear.³³¹ Its many metal alloys are mostly brittle.³³² It shows fewer tendencies to anionic behaviour than ordinary nonmetals.³³³ Its solution chemistry is characterised by the formation of oxyanions.³³⁴ Arsenic generally forms compounds in which it has an oxidation state of +3 or +5.³³⁵ The halides, and the oxides and their derivatives are illustrative examples.³³⁶ In the trivalent state, arsenic shows some incipient metallic properties.³³⁷ The halides are hydrolysed by water but these reactions, particularly those of the chloride, are reversible with the addition of a hydrohalic acid.³³⁸ The oxide is acidic but, as noted below, (weakly) amphoteric. The higher, less stable, pentavalent state has strongly acidic (nonmetallic) properties.³³⁹ Compared to phosphorus, the stronger metallic character of arsenic is indicated by the formation of oxoacid salts such as AsPO4, As2(SO4)3³⁴⁰ and arsenic acetate As(CH3COO)3.³⁴¹ The oxide As2O3 is polymeric,³⁴² amphoteric,³⁴³³⁴⁴ and a glass former.³⁴⁵ Arsenic has an extensive organometallic chemistry (see Organoarsenic chemistry).³⁴⁶

Antimony

Main article: Antimony

Antimony is a silver-white solid with a blue tint and a brilliant lustre.³⁴⁷ 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).³⁴⁸ 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)3.³⁴⁹ It is not affected by molten alkali.³⁵⁰ Antimony is capable of displacing hydrogen from water, when heated: 2 Sb + 3 H2O → Sb2O3 + 3 H2.³⁵¹ It melts at 631 °C. Antimony is a semimetal with an electrical conductivity of around 3.1 × 104 S•cm−1³⁵² and a band overlap of 0.16 eV.³⁵³³⁵⁴ Liquid antimony is a metallic conductor with an electrical conductivity of around 5.3 × 104 S•cm−1.³⁵⁵

Most of the chemistry of antimony is characteristic of a nonmetal.³⁵⁶ Antimony has some definite cationic chemistry,³⁵⁷ SbO+ and Sb(OH)2+ being present in acidic aqueous solution;³⁵⁸³⁵⁹ the compound Sb8(GaCl4)2, which contains the homopolycation, Sb82+, was prepared in 2004.³⁶⁰ It can form alloys with one or more metals such as aluminium,³⁶¹ iron, nickel, copper, zinc, tin, lead, and bismuth.³⁶² Antimony has fewer tendencies to anionic behaviour than ordinary nonmetals.³⁶³ Its solution chemistry is characterised by the formation of oxyanions.³⁶⁴ Like arsenic, antimony generally forms compounds in which it has an oxidation state of +3 or +5.³⁶⁵ The halides, and the oxides and their derivatives are illustrative examples.³⁶⁶ 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.³⁶⁷ Tripositive antimony is amphoteric; pentapositive antimony is (predominately) acidic.³⁶⁸ Consistent with an increase in metallic character down group 15, antimony forms salts including an acetate Sb(CH3CO2)3, phosphate SbPO4, sulfate Sb2(SO4)3 and perchlorate Sb(ClO4)3.³⁶⁹ 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+2.³⁷⁰ The oxide Sb2O3 is polymeric,³⁷¹ amphoteric,³⁷² and a glass former.³⁷³ Antimony has an extensive organometallic chemistry (see Organoantimony chemistry).³⁷⁴

Tellurium

Main article: Tellurium

Tellurium is a silvery-white shiny solid.³⁷⁵ 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.³⁷⁶ 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 H2.³⁷⁷ It reacts (to varying degrees) with nitric, sulfuric, and hydrochloric acids to give compounds such as the sulfoxide TeSO3 or tellurous acid H2TeO3,³⁷⁸ the basic nitrate (Te2O4H)+(NO3)−,³⁷⁹ or the oxide sulfate Te2O3(SO4).³⁸⁰ 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.³⁸¹

At higher temperatures tellurium is sufficiently plastic to extrude.³⁸² 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.³⁸³ Tellurium is a semiconductor with an electrical conductivity of around 1.0 S•cm−1³⁸⁴ and a band gap of 0.32 to 0.38 eV.³⁸⁵ Liquid tellurium is a semiconductor, with an electrical conductivity, on melting, of around 1.9 × 103 S•cm−1.³⁸⁶ Superheated liquid tellurium is a metallic conductor.³⁸⁷

Most of the chemistry of tellurium is characteristic of a nonmetal.³⁸⁸ It shows some cationic behaviour. The dioxide dissolves in acid to yield the trihydroxotellurium(IV) Te(OH)3+ ion;³⁸⁹³⁹⁰ the red Te42+ and yellow-orange Te62+ ions form when tellurium is oxidized in fluorosulfuric acid (HSO3F), or liquid sulfur dioxide (SO2), respectively.³⁹¹ It can form alloys with aluminium, silver, and tin.³⁹² Tellurium shows fewer tendencies to anionic behaviour than ordinary nonmetals.³⁹³ Its solution chemistry is characterised by the formation of oxyanions.³⁹⁴ Tellurium generally forms compounds in which it has an oxidation state of −2, +4 or +6. The +4 state is the most stable.³⁹⁵ Tellurides of composition XxTey 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.³⁹⁶ 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·HXO4.³⁹⁷ Tellurium forms a polymeric,³⁹⁸ amphoteric,³⁹⁹ glass-forming oxide⁴⁰⁰ TeO2. It is a "conditional" glass-forming oxide – it forms a glass with a very small amount of additive.⁴⁰¹ Tellurium has an extensive organometallic chemistry (see Organotellurium chemistry).⁴⁰²

Elements less commonly recognised as metalloids

Carbon

Main article: Carbon

Carbon is ordinarily classified as a nonmetal⁴⁰³ 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 appearance⁴⁰⁶ and is a fairly good electrical conductor.⁴⁰⁷ Graphite has a layered structure. Each layer consists of carbon atoms bonded to three other carbon atoms in a hexagonal lattice arrangement. The layers are stacked together and held loosely by van der Waals 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;⁴⁰⁹⁴¹⁰ it has the electronic band structure of a semimetal.⁴¹¹ The allotropes of carbon, including graphite, can accept foreign atoms or compounds into their structures via substitution, intercalation, or doping. The resulting materials are sometimes referred to as "carbon alloys".⁴¹² Carbon can form ionic salts, including a hydrogen sulfate, perchlorate, and nitrate (C+24X−.2HX, where X = HSO4, ClO4; and C+24NO–3.3HNO3).⁴¹³⁴¹⁴ 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

Graphite is an established solid lubricant and behaves as a semiconductor in a direction perpendicular to its planes.⁴¹⁶ Most of its chemistry is nonmetallic;⁴¹⁷ it has a relatively high ionization energy⁴¹⁸ 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 H2CO3.⁴²¹⁴²²

Aluminium

Main article: Aluminium

Aluminium is ordinarily classified as a metal.⁴²³ It is lustrous, malleable and ductile, and has high electrical and thermal conductivity. Like most metals it has a close-packed crystalline structure,⁴²⁴ and forms a cation in aqueous solution.⁴²⁵

It has some properties that are unusual for a metal; taken together,⁴²⁶ these are sometimes used as a basis to classify aluminium as a metalloid.⁴²⁷ Its crystalline structure shows some evidence of directional bonding.⁴²⁸ Aluminium bonds covalently in most compounds.⁴²⁹ The oxide Al2O3 is amphoteric⁴³⁰ and a conditional glass-former.⁴³¹ Aluminium can form anionic aluminates,⁴³² such behaviour being considered nonmetallic in character.⁴³³

Classifying aluminium as a metalloid has been disputed⁴³⁴ 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.⁴³⁵⁴³⁶

Stott⁴³⁷ labels aluminium as a weak metal. It has the physical properties of a metal but some of the chemical properties of a nonmetal. Steele⁴³⁸ 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". Moody⁴³⁹ says that, "aluminium is on the 'diagonal borderland' between metals and non-metals in the chemical sense."

Selenium

Main article: Selenium

Selenium shows borderline metalloid or nonmetal behaviour.⁴⁴⁰⁴⁴¹

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 oleums⁴⁴⁷ (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 brittleness⁴⁴⁹ and the low electrical conductivity (~10−9 to 10−12 S•cm−1) of its highly purified form.⁴⁵⁰ This is comparable to or less than that of bromine (7.95×10–12 S•cm−1),⁴⁵¹ a nonmetal. Selenium has the electronic band structure of a semiconductor⁴⁵² and retains its semiconducting properties in liquid form.⁴⁵³ It has a relatively high⁴⁵⁴ electronegativity (2.55 revised Pauling scale). Its reaction chemistry is mainly that of its nonmetallic anionic forms Se2−, SeO2−3 and SeO2−4.⁴⁵⁵

Selenium is commonly described as a metalloid in the environmental chemistry literature.⁴⁵⁶ It moves through the aquatic environment similarly to arsenic and antimony;⁴⁵⁷ its water-soluble salts, in higher concentrations, have a similar toxicological profile to that of arsenic.⁴⁵⁸

Polonium

Main article: Polonium

Polonium is "distinctly metallic" in some ways.⁴⁵⁹ Both of its allotropic forms are metallic conductors.⁴⁶⁰ It is soluble in acids, forming the rose-coloured Po2+ cation and displacing hydrogen: Po + 2 H+ → Po2+ + H2.⁴⁶¹ Many polonium salts are known.⁴⁶² The oxide PoO2 is predominantly basic in nature.⁴⁶³ Polonium is a reluctant oxidizing agent, unlike its lightest congener oxygen: highly reducing conditions are required for the formation of the Po2− anion in aqueous solution.⁴⁶⁴

Whether polonium is ductile or brittle is unclear. It is predicted to be ductile based on its calculated elastic constants.⁴⁶⁵ 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".⁴⁶⁶

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).⁴⁶⁷ Many metal polonides, obtained by heating the elements together at 500–1,000 °C, and containing the Po2− anion, are also known.⁴⁶⁸

Astatine

Main article: Astatine

As a halogen, astatine tends to be classified as a nonmetal.⁴⁶⁹ It has some marked metallic properties⁴⁷⁰ and is sometimes instead classified as either a metalloid⁴⁷¹ or (less often) as a metal.⁴⁷² Immediately following its production in 1940, early investigators considered it a metal.⁴⁷³ In 1949 it was called the most noble (difficult to reduce) nonmetal as well as being a relatively noble (difficult to oxidize) metal.⁴⁷⁴ In 1950 astatine was described as a halogen and (therefore) a reactive nonmetal.⁴⁷⁵ In 2013, on the basis of relativistic modelling, astatine was predicted to be a monatomic metal, with a face-centred cubic crystalline structure.⁴⁷⁶

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.⁴⁷⁷⁴⁷⁸ Astatine may be metallic in the liquid state on the basis that elements with an enthalpy of vaporization (∆Hvap) greater than ~42 kJ/mol are metallic when liquid.⁴⁷⁹ Such elements include boron,⁴⁸⁰ silicon, germanium, antimony, selenium, and tellurium. Estimated values for ∆Hvap of diatomic astatine are 50 kJ/mol or higher;⁴⁸¹ diatomic iodine, with a ∆Hvap of 41.71,⁴⁸² 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."⁴⁸³⁴⁸⁴ 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".⁴⁸⁵ It has also been argued that astatine demonstrates cationic behaviour, by way of stable At+ and AtO+ forms, in strongly acidic aqueous solutions.⁴⁸⁶

Some of astatine's reported properties are nonmetallic. It has been extrapolated to have the narrow liquid range ordinarily associated with nonmetals (mp 302 °C; bp 337 °C),⁴⁸⁷ although experimental indications suggest a lower boiling point of about 230±3 °C. Batsanov gives a calculated band gap energy for astatine of 0.7 eV;⁴⁸⁸ this is consistent with nonmetals (in physics) having separated valence and conduction bands and thereby being either semiconductors or insulators.⁴⁸⁹ The chemistry of astatine in aqueous solution is mainly characterised by the formation of various anionic species.⁴⁹⁰ Most of its known compounds resemble those of iodine,⁴⁹¹ which is a halogen and a nonmetal.⁴⁹² Such compounds include astatides (XAt), astatates (XAtO3), and monovalent interhalogen compounds.⁴⁹³

Restrepo et al.⁴⁹⁴ 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

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-metalloids⁴⁹⁵ or noted for their metalloidal character. To the left of the metal–nonmetal dividing line, such elements include gallium,⁴⁹⁶ tin⁴⁹⁷ 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 selenium⁵⁰⁴ and iodine.⁵⁰⁵ They exhibit metallic lustre, semiconducting properties⁵⁰⁶ 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;⁵⁰⁷ and selenium as grey selenium.

Allotropes

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.⁵⁰⁸ The existence of such allotropes can complicate the classification of the elements involved.⁵⁰⁹

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).⁵¹⁰ 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).⁵¹¹ Grey tin has the same crystalline structure as that of diamond. It behaves as a semiconductor (as if it had a band gap of 0.08 eV), but has the electronic band structure of a semimetal.⁵¹² It has been referred to as either a very poor metal,⁵¹³ a metalloid,⁵¹⁴ a nonmetal⁵¹⁵ or a near metalloid.⁵¹⁶

The diamond allotrope of carbon is clearly nonmetallic, being translucent and having a low electrical conductivity of 10−14 to 10−16 S·cm−1.⁵¹⁷ Graphite has an electrical conductivity of 3 × 104 S·cm−1,⁵¹⁸ a figure more characteristic of a metal. Phosphorus, sulfur, arsenic, selenium, antimony, and bismuth also have less stable allotropes that display different behaviours.⁵¹⁹

Abundance, extraction, and cost

Abundance

The table gives crustal abundances of the elements commonly to rarely recognised as metalloids.⁵²⁰ 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.⁵²¹

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.⁵²² 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%.⁵²³ 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 CO2.⁵²⁴ 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;⁵²⁵ Na2TeO3 + H2O → Te + 2 NaOH + O2.⁵²⁶ Another option is reduction of the oxide by roasting with carbon: TeO2 + C → Te + CO2.⁵²⁷

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 the 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.⁵²⁸

Cost

The recognised metalloids and their closer neighbours mostly cost less than silver; only polonium and astatine are more expensive than gold, on account of their significant radioactivity. As of 5 April 2014, prices for small samples (up to 100 g) of silicon, antimony and tellurium, and graphite, aluminium and selenium, average around one third the cost of silver (US$1.5 per gram or about $45 an ounce). Boron, germanium, and arsenic samples average about three-and-a-half times the cost of silver.⁵²⁹ Polonium is available for about $100 per microgram.⁵³⁰ Zalutsky and Pruszynski⁵³¹ 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).⁵³²

Notes

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• Zingaro RA 1994, 'Arsenic: Inorganic Chemistry', in RB King (ed.) 1994, Encyclopedia of Inorganic Chemistry, John Wiley & Sons, Chichester, pp. 192–218, ISBN 0-471-93620-0

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. 537–91, ISBN 0-471-06315-0
• Chedd G (1969), Half-way Elements: The Technology of Metalloids, Doubleday, New York[ISBN missing]
• Choppin GR & Johnsen RH (1972), "Group IV and the Metalloids", in Introductory Chemistry, Addison-Wesley, Reading, Massachusetts, pp. 341–57
• 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. 526527, doi:10.1021/ed059p526
• Hawkes SJ (2001), "Semimetallicity", Journal of Chemical Education, vol. 78, no. 12, pp. 1686–87, 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
• Miller JS (2019), "Viewpoint: Metalloids – An Electronic Band Structure Perspective", Chemistry – A European Perspective, preprint version, doi:10.1002/chem.201903167
• 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. 742–75, ISBN 0-15-506492-4
• Parveen N et al. (2020), "Metalloids in plants: A systematic discussion beyond description", Annals of Applied Biology, doi:10.1111/aab.12666of
• Rieske M (1998), "Metalloids", in Encyclopedia of Earth and Physical Sciences, Marshall Cavendish, New York, vol. 6, pp. 758–59, ISBN 0-7614-0551-8 (set)
• Rochow EG (1966), The Metalloids, DC Heath and Company, Boston[ISBN missing]
• Vernon RE (2013), "Which Elements are Metalloids?", Journal of Chemical Education, vol. 90, no. 12, pp. 1703–07, doi:10.1021/ed3008457
• —— (2020,) "Organising the Metals and Nonmetals", Foundations of Chemistry, (open access)

References

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 /wiki/Doi_(identifier) ↩

2. Oxford English Dictionary 1989, 'metalloid'; Gordh, Gordh & Headrick 2003, p. 753 ↩

3. Atkins et al. 2010, p. 20 ↩

4. 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] /wiki/Metal ↩

5. Hopkins & Bailar 1956, p. 458 ↩

6. Glinka 1965, p. 77 ↩

7. 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 highest 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] /wiki/Gold ↩

8. Tyler Miller 1987, p. 59 ↩

9. Goldsmith 1982, p. 526; Kotz, Treichel & Weaver 2009, p. 62; Bettelheim et al. 2010, p. 46 ↩

10. Mann et al.[16] refer to these elements as "the recognized metalloids". ↩

11. Hawkes 2001, p. 1686; Segal 1989, p. 965; McMurray & Fay 2009, p. 767 ↩

12. Bucat 1983, p. 26; Brown c. 2007 ↩

13. Swift & Schaefer 1962, p. 100 ↩

14. Hawkes 2001, p. 1686; Hawkes 2010; Holt, Rinehart & Wilson c. 2007 ↩

15. 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). ↩

16. Tilden 1876, pp. 172, 198–201; Smith 1994, p. 252; Bodner & Pardue 1993, p. 354 ↩

17. Bassett et al. 1966, p. 127 ↩

18. Rausch 1960 ↩

19. Thayer 1977, p. 604; Warren & Geballe 1981; Masters & Ela 2008, p. 190 ↩

20. Warren & Geballe 1981; Chalmers 1959, p. 72; US Bureau of Naval Personnel 1965, p. 26 ↩

21. Siebring 1967, p. 513 ↩

22. Wiberg 2001, p. 282 ↩

23. Rausch 1960; Friend 1953, p. 68 ↩

24. Murray 1928, p. 1295 ↩

25. Swift & Schaefer 1962, p. 100 ↩

26. Hampel & Hawley 1966, p. 950; Stein 1985; Stein 1987, pp. 240, 247–48 ↩

27. Hatcher 1949, p. 223; Secrist & Powers 1966, p. 459 ↩

28. Taylor 1960, p. 614 ↩

29. Considine & Considine 1984, p. 568; Cegielski 1998, p. 147; The American heritage science dictionary 2005, p. 397 ↩

30. Woodward 1948, p. 1 ↩

31. NIST 2010. Values shown in the above table have been converted from the NIST values, which are given in eV. ↩

32. Berger 1997; Lovett 1977, p. 3 ↩

33. Goldsmith 1982, p. 526; Hawkes 2001, p. 1686 ↩

34. Hawkes 2001, p. 1687 ↩

35. Sharp 1981, p. 299 ↩

36. Emsley 1971, p. 1 ↩

37. James et al. 2000, p. 480 ↩

38. 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 ..." ↩

39. 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." ↩

40. The lack of a standard division of the elements into metals, metalloids, and nonmetals is not necessarily an issue. There is more or less, a continuous progression from the metallic to the nonmetallic. A specified subset of this continuum could serve its particular purpose as well as any other.[45] ↩

41. Rochow 1966, pp. 1, 4–7 ↩

42. Rochow 1977, p. 76; Mann et al. 2000, p. 2783 ↩

43. Askeland, Phulé & Wright 2011, p. 69 ↩

44. 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] ↩

45. Edwards & Sienko 1983, p. 693 ↩

46. 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. /wiki/Karl_Herzfeld ↩

47. Edwards & Sienko 1983, p. 695; Edwards et al. 2010 ↩

48. 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] ↩

49. Atomic conductance is the electrical conductivity of one mole of a substance. It is equal to electrical conductivity divided by molar volume.[5] ↩

50. Hill & Holman 2000, p. 160. They characterise metalloids (in part) on the basis that they are "poor conductors of electricity with atomic conductance usually less than 10−3 but greater than 10−5 ohm−1 cm−4". ↩

51. 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)." ↩

52. Jones 2010, p. 169 ↩

53. 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. ↩

54. 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. ↩

55. Vernon 2013, p. 1703 ↩

56. Hamm 1969, p. 653 ↩

57. Horvath 1973, p. 336 ↩

58. Gray 2009, p. 9 ↩

59. Rayner-Canham 2011 ↩

60. Booth & Bloom 1972, p. 426; Cox 2004, pp. 17, 18, 27–28; Silberberg 2006, pp. 305–13 ↩

61. Cox 2004, pp. 17–18, 27–28; Silberberg 2006, pp. 305–13 ↩

62. Gray 2009, p. 9 ↩

63. Rodgers 2011, pp. 232–33; 240–41 ↩

64. Roher 2001, pp. 4–6 ↩

65. Sharp 1981, p. 299 ↩

66. Jones (2010, pp. 169–71): "Though classification is an essential feature of all branches of science, there are always hard cases at the boundaries. The boundary of a class is rarely sharp…Scientists should not lose sleep over the hard cases. As long as a classification system is beneficial to economy of description, to structuring knowledge and to our understanding, and hard cases constitute a small minority, then keep it. If the system becomes less than useful, then scrap it and replace it with a system based on different shared characteristics." ↩

67. Tyler 1948, p. 105; Reilly 2002, pp. 5–6 ↩

68. Hampel & Hawley 1976, p. 174; ↩

69. Goodrich 1844, p. 264; The Chemical News 1897, p. 189; Hampel & Hawley 1976, p. 191; Lewis 1993, p. 835; Hérold 2006, pp. 149–50 ↩

70. Oderberg[80] argues on ontological grounds that anything not a metal is therefore a nonmetal, and that this includes semi-metals (i.e. metalloids). ↩

71. Brown & Holme 2006, p. 57 ↩

72. Wiberg 2001, p. 282; Simple Memory Art c. 2005 ↩

73. Chedd 1969, pp. 12–13 ↩

74. Kneen, Rogers & Simpson, 1972, p. 263. Columns 2 and 4 are sourced from this reference unless otherwise indicated. ↩

75. Stoker 2010, p. 62; Chang 2002, p. 304. Chang speculates that the melting point of francium would be about 23 °C. ↩

76. Copernicium is reportedly the only metal thought to be a gas at room temperature.[86] /wiki/Copernicium ↩

77. Rochow 1966, p. 4 ↩

78. Hunt 2000, p. 256 ↩

79. Rochow 1966, p. 4 ↩

80. McQuarrie & Rock 1987, p. 85 ↩

81. Metals have electrical conductivity values of from 6.9 × 103 S•cm−1 for manganese to 6.3 × 105 for silver.[90] /wiki/Manganese ↩

82. Choppin & Johnsen 1972, p. 351 ↩

83. Metalloids have electrical conductivity values of from 1.5 × 10−6 S•cm−1 for boron to 3.9 × 104 for arsenic.[92] If selenium is included as a metalloid the applicable conductivity range would start from ~10−9 to 10−12 S•cm−1.[93] ↩

84. Nonmetals have electrical conductivity values of from ~10−18 S•cm−1 for the elemental gases to 3 × 104 in graphite.[94] ↩

85. Hampel & Hawley 1976, p. 191; Wulfsberg 2000, p. 620 ↩

86. Swalin 1962, p. 216 ↩

87. Bailar et al. 1989, p. 742 ↩

88. Metcalfe, Williams & Castka 1974, p. 86 ↩

89. Chang 2002, p. 306 ↩

90. Pauling 1988, p. 183 ↩

91. Mann et al. 2000, p. 2783 ↩

92. Chedd[101] 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[102] 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. ↩

93. Hultgren 1966, p. 648; Young & Sessine 2000, p. 849; Bassett et al. 1966, p. 602 ↩

94. Rochow 1966, p. 4; Atkins et al. 2006, pp. 8, 122–23 ↩

95. Russell & Lee 2005, pp. 421, 423; Gray 2009, p. 23 ↩

96. Olmsted and Williams[106] 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". ↩

97. Desch 1914, p. 86 ↩

98. Phillips & Williams 1965, p. 620 ↩

99. Van der Put 1998, p. 123 ↩

100. Klug & Brasted 1958, p. 199 ↩

101. Good et al. 1813 ↩

102. Sequeira 2011, p. 776 ↩

103. Gary 2013 ↩

104. Russell & Lee 2005, pp. 405–06; 423–34 ↩

105. Davidson & Lakin 1973, p. 627 ↩

106. Wiberg 2001, p. 589 ↩

107. Greenwood & Earnshaw 2002, p. 749; Schwartz 2002, p. 679 ↩

108. Řezanka & Sigler 2008; Sekhon 2012 ↩

109. Emsley 2001, p. 67 ↩

110. Zhang et al. 2008, p. 360 ↩

111. Science Learning Hub 2009 ↩

112. Skinner et al. 1979; Tom, Elden & Marsh 2004, p. 135 ↩

113. Büchel 1983, p. 226 ↩

114. Science Learning Hub 2009 ↩

115. Emsley 2001, p. 391 ↩

116. Schauss 1991; Tao & Bolger 1997 ↩

117. Eagleson 1994, p. 450; EVM 2003, pp. 197‒202 ↩

118. Nielsen 1998 ↩

119. MacKenzie 2015, p. 36 ↩

120. Jaouen & Gibaud 2010 ↩

121. Jaouen & Gibaud 2010 ↩

122. Smith et al. 2014 ↩

123. Stevens & Klarner, p. 205 ↩

124. Sneader 2005, pp. 57–59 ↩

125. Keall, Martin and Tunbridge 1946 ↩

126. Emsley 2001, p. 426 ↩

127. Oldfield et al. 1974, p. 65; Turner 2011 ↩

128. Ba et al. 2010; Daniel-Hoffmann, Sredni & Nitzan 2012; Molina-Quiroz et al. 2012 ↩

129. Peryea 1998 ↩

130. Nielsen 1998 ↩

131. Hager 2006, p. 299 ↩

132. Apseloff 1999 ↩

133. Trivedi, Yung & Katz 2013, p. 209 ↩

134. Emsley 2001, p. 382; Burkhart, Burkhart & Morrell 2011 ↩

135. Thomas, Bialek & Hensel 2013, p. 1 ↩

136. Perry 2011, p. 74 ↩

137. UCR Today 2011; Wang & Robinson 2011; Kinjo et al. 2011 ↩

138. Kauthale et al. 2015 ↩

139. Gunn 2014, pp. 188, 191 ↩

140. Gupta, Mukherjee & Cameotra 1997, p. 280; Thomas & Visakh 2012, p. 99 ↩

141. Muncke 2013 ↩

142. Mokhatab & Poe 2012, p. 271 ↩

143. Craig, Eng & Jenkins 2003, p. 25 ↩

144. McKee 1984 ↩

145. Hai et al. 2012 ↩

146. Kohl & Nielsen 1997, pp. 699–700 ↩

147. Chopra et al. 2011 ↩

148. Le Bras, Wilkie & Bourbigot 2005, p. v ↩

149. Wilkie & Morgan 2009, p. 187 ↩

150. Locke et al. 1956, p. 88 ↩

151. Carlin 2011, p. 6.2 ↩

152. Evans 1993, pp. 257–28 ↩

153. Corbridge 2013, p. 1149 ↩

154. Kaminow & Li 2002, p. 118 ↩

155. Kaminow & Li 2002, p. 118 ↩

156. Deming 1925, pp. 330 (As2O3), 418 (B2O3; SiO2; Sb2O3); Witt & Gatos 1968, p. 242 (GeO2) ↩

157. Eagleson 1994, p. 421 (GeO2); Rothenberg 1976, 56, 118–19 (TeO2) ↩

158. Geckeler 1987, p. 20 ↩

159. Kreith & Goswami 2005, pp. 12–109 ↩

160. Russell & Lee 2005, p. 397 ↩

161. Butterman & Jorgenson 2005, pp. 9–10 ↩

162. Shelby 2005, p. 43 ↩

163. Butterman & Carlin 2004, p. 22; Russell & Lee 2005, p. 422 ↩

164. Träger 2007, pp. 438, 958; Eranna 2011, p. 98 ↩

165. Rao 2002, p. 552; Löffler, Kündig & Dalla Torre 2007, p. 17–11 ↩

166. 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.[174] /wiki/Covalent ↩

167. Klement, Willens & Duwez 1960; Wanga, Dongb & Shek 2004, p. 45 ↩

168. Demetriou et al. 2011; Oliwenstein 2011 ↩

169. Karabulut et al. 2001, p. 15; Haynes 2012, pp. 4–26 ↩

170. Schwartz 2002, pp. 679–80 ↩

171. Carter & Norton 2013, p. 403 ↩

172. Maeder 2013, pp. 3, 9–11 ↩

173. Tominaga 2006, pp. 327–28; Chung 2010, pp. 285–86; Kolobov & Tominaga 2012, p. 149 ↩

174. New Scientist 2014; Hosseini, Wright & Bhaskaran 2014; Farandos et al. 2014 ↩

175. Kosanke 2002, p. 110 ↩

176. Ellern 1968, pp. 246, 326–27 ↩

177. Conkling & Mocella 2010, p. 82 ↩

178. Crow 2011; Mainiero 2014 ↩

179. Conkling & Mocella 2010, p. 82 ↩

180. 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.[188] /wiki/MoO3 ↩

181. Ellern 1968, p. 135; Weingart 1947, p. 9 ↩

182. Conkling & Mocella 2010, p. 83 ↩

183. Conkling & Mocella 2010, pp. 181, 213 ↩

184. Ellern 1968, pp. 209–10, 322 ↩

185. Russell 2009, pp. 15, 17, 41, 79–80 ↩

186. Ellern, writing in Military and Civilian Pyrotechnics (1968), comments that carbon black "has been specified for and used in a nuclear air-burst simulator."[194] /wiki/Carbon_black ↩

187. Kosanke 2002, p. 110 ↩

188. Ellern 1968, p. 328 ↩

189. Conkling & Mocella 2010, p. 171 ↩

190. Conkling & Mocella 2011, pp. 83–84 ↩

191. Ellern 1968, pp. 209–10, 322 ↩

192. Berger 1997, p. 91; Hampel 1968, passim ↩

193. Rochow 1966, p. 41; Berger 1997, pp. 42–43 ↩

194. Bomgardner 2013, p. 20 ↩

195. Russell & Lee 2005, p. 395; Brown et al. 2009, p. 489 ↩

196. 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) ↩

197. Russell & Lee 2005, p. 401; Büchel, Moretto & Woditsch 2003, p. 278 ↩

198. Russell & Lee 2005, p. 401; Büchel, Moretto & Woditsch 2003, p. 278 ↩

199. Bianco et al. 2013 ↩

200. University of Limerick 2014; Kennedy et al. 2014 ↩

201. Lee et al. 2014 ↩

202. Russell & Lee 2005, pp. 421–22, 424 ↩

203. He et al. 2014 ↩

204. Berger 1997, p. 91 ↩

205. Bomgardner 2013, p. 20 ↩

206. ScienceDaily 2012 ↩

207. Reardon 2005; Meskers, Hagelüken & Van Damme 2009, p. 1131 ↩

208. The Economist 2012 ↩

209. Whitten 2007, p. 488 ↩

210. Jaskula 2013 ↩

211. German Energy Society 2008, pp. 43–44 ↩

212. Patel 2012, p. 248 ↩

213. Moore 2104; University of Utah 2014; Xu et al. 2014 ↩

214. Yang et al. 2012, p. 614 ↩

215. Moore 2010, p. 195 ↩

216. Moore 2011 ↩

217. Liu 2014 ↩

218. Bradley 2014; University of Utah 2014 ↩

219. Foster 1936, pp. 212–13; Brownlee et al. 1943, p. 293 ↩

220. Calderazzo, Ercoli & Natta 1968, p. 257 ↩

221. Walters 1982, pp. 32–33 ↩

222. Tyler 1948, p. 105 ↩

223. 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." ↩

224. Slade 2006, p. 16 ↩

225. Corwin 2005, p. 80 ↩

226. Barsanov & Ginzburg 1974, p. 330 ↩

227. Bradbury et al. 1957, pp. 157, 659 ↩

228. Klemm 1950, pp. 133–42; Reilly 2004, p. 4 ↩

229. King 2004, pp. 196–98; Ferro & Saccone 2008, p. 233 ↩

230. Lister 1965, p. 54 ↩

231. Cotton et al. 1999, p. 502 ↩

232. Apjohn, J. (1864). Manual of the Metalloids. United Kingdom: Longman. ↩

233. Pinkerton 1800, p. 81 ↩

234. Goldsmith 1982, p. 526 ↩

235. Friend 1953, p. 68; IUPAC 1959, p. 10; IUPAC 1971, p. 11 ↩

236. Atkins et al. 2010, p. 20 ↩

237. IUPAC 2005; IUPAC 2006– ↩

238. Housecroft & Sharpe 2008, p. 331; Oganov 2010, p. 212 ↩

239. Housecroft & Sharpe 2008, p. 333 ↩

240. Kross 2011 ↩

241. Berger 1997, p. 37 ↩

242. Greenwood & Earnshaw 2002, p. 144 ↩

243. Kopp, Lipták & Eren 2003, p. 221 ↩

244. Prudenziati 1977, p. 242 ↩

245. 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).[248] ↩

246. Mendeléeff 1897, p. 57 ↩

247. Rayner-Canham & Overton 2006, p. 291 ↩

248. Siekierski & Burgess 2002, p. 63 ↩

249. 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.[252] /wiki/Borospherene ↩

250. Siekierski & Burgess 2002, p. 86 ↩

251. Greenwood & Earnshaw 2002, p. 141; Henderson 2000, p. 58; Housecroft & Sharpe 2008, pp. 360–72 ↩

252. Parry et al. 1970, pp. 438, 448–51 ↩

253. Fehlner 1990, p. 202 ↩

254. Fehlner 1990, p. 202 ↩

255. Owen & Brooker 1991, p. 59; Wiberg 2001, p. 936 ↩

256. Greenwood & Earnshaw 2002, p. 145 ↩

257. Houghton 1979, p. 59 ↩

258. The BH3 and Fe(CO4) species in these reactions are short-lived reaction intermediates.[260] /wiki/Reaction_intermediate ↩

259. Fehlner 1990, pp. 204–05, 207 ↩

260. On the analogy between boron and metals, Greenwood[262] 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 ..." ↩

261. 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 ↩

262. Sharp 1983, p. 56 ↩

263. Fokwa 2014, p. 10 ↩

264. The bonding in boron trifluoride, a gas, has been referred to as predominately ionic[266] a description which was subsequently described as misleading.[267] /wiki/Boron_trifluoride ↩

265. Rayner-Canham & Overton 2006, p. 291 ↩

266. Greenwood & Earnshaw 2002, p. 145 ↩

267. Puddephatt & Monaghan 1989, p. 59 ↩

268. Mahan 1965, p. 485 ↩

269. Boron trioxide B2O3 is sometimes described as being (weakly) amphoteric.[270] It reacts with alkalies to give various borates.[271] In its hydrated form (as H3BO3, boric acid) it reacts with sulfur trioxide, the anhydride of sulfuric acid, to form a bisulfate B(HSO3) 4.[272] In its pure (anhydrous) form it reacts with phosphoric acid to form a "phosphate" BPO4.[273] The latter compound may be regarded as a mixed oxide of B2O3 and P2O5.[274] /wiki/Amphoteric ↩

270. Rao 2002, p. 22 ↩

271. Organic derivatives of metalloids are traditionally counted as organometallic compounds.[276] ↩

272. Haiduc & Zuckerman 1985, p. 82 ↩

273. Greenwood & Earnshaw 2002, p. 331 ↩

274. Wiberg 2001, p. 824 ↩

275. Greenwood & Earnshaw 2002, p. 331 ↩

276. Rochow 1973, pp. 1337‒38 ↩

277. In air, silicon forms a thin coating of amorphous silicon dioxide, 2 to 3 nm thick.[281] This coating is dissolved by hydrogen fluoride at a very low pace – on the order of two to three hours per nanometre.[282] Silicon dioxide, and silicate glasses (of which silicon dioxide is a major component), are otherwise readily attacked by hydrofluoric acid.[283] ↩

278. Rochow 1973, pp. 1337, 1340 ↩

279. Allen & Ordway 1968, p. 152 ↩

280. Eagleson 1994, pp. 48, 127, 438, 1194; Massey 2000, p. 191 ↩

281. Orton 2004, p. 7. This is a typical value for high-purity silicon. ↩

282. Russell & Lee 2005, p. 393 ↩

283. Coles & Caplin 1976, p. 106 ↩

284. Glazov, Chizhevskaya & Glagoleva 1969, pp. 59–63; Allen & Broughton 1987, p. 4967 ↩

285. Cotton, Wilkinson & Gaus 1995, p. 393 ↩

286. Wiberg 2001, p. 834 ↩

287. The bonding in silicon tetrafluoride, a gas, has been referred to as predominately ionic[266] a description which was subsequently described as misleading.[267] /wiki/Silicon_tetrafluoride ↩

288. Partington 1944, p. 723 ↩

289. Cox 2004, p. 27 ↩

290. Hiller & Herber 1960, inside front cover; p. 225 ↩

291. Kneen, Rogers and Simpson 1972, p. 384 ↩

292. Bailar, Moeller & Kleinberg 1965, p. 513 ↩

293. Cotton, Wilkinson & Gaus 1995, pp. 319, 321 ↩

294. Puddephatt & Monaghan 1989, p. 59 ↩

295. Smith 1990, p. 175 ↩

296. 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,[299] and with hydrofluoric acid to give hexafluorosilicic acid H2SiF6.[300] The latter reaction "is sometimes quoted as evidence of basic [that is, metallic] properties".[301] ↩

297. Rao 2002, p. 22 ↩

298. Powell 1988, p. 1 ↩

299. Greenwood & Earnshaw 2002, p. 371 ↩

300. Cusack 1967, p. 193 ↩

301. Temperatures above 400 °C are required to form a noticeable surface oxide layer.[305] ↩

302. Greenwood & Earnshaw 2002, p. 373 ↩

303. Moody 1991, p. 273 ↩

304. Greenwood & Earnshaw 2002, p. 373 ↩

305. Russell & Lee 2005, p. 399 ↩

306. Berger 1997, pp. 71–72 ↩

307. Jolly 1966, pp. 125–6 ↩

308. Sources mentioning germanium cations include: Powell & Brewer[311] 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,[312] in "many metallic halides, hydroxides, and chalcides"); Everest[313] 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[314] 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.[315] who prepared the layer compound or phase Nb3GexTe6 (x ≃ 0.9), and reported that this contained a GeII cation; Richens[316] 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.[317] who synthesized a cryptand compound containing a Ge2+ cation; and Schwietzer and Pesterfield[318] 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[319] 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[320] who says that, "[germanium] forms no aquacations"; Ladd[321] who notes that the CdI2 structure is "intermediate in type between ionic and molecular compounds"; and Wiberg[322] who states that, "no germanium cations are known". ↩

309. Schwartz 2002, p. 269 ↩

310. Cox 2004, p. 27 ↩

311. Hiller & Herber 1960, inside front cover; p. 225 ↩

312. Eggins 1972, p. 66; Wiberg 2001, p. 895 ↩

313. Greenwood & Earnshaw 2002, p. 383 ↩

314. Glockling 1969, p. 38; Wells 1984, p. 1175 ↩

315. Puddephatt & Monaghan 1989, p. 59 ↩

316. Cooper 1968, pp. 28–29 ↩

317. Rao 2002, p. 22 ↩

318. Steele 1966, pp. 178, 188–89 ↩

319. Haller 2006, p. 3 ↩

320. See, for example, Walker & Tarn 1990, p. 590 ↩

321. Wiberg 2001, p. 742 ↩

322. Gray, Whitby & Mann 2011 ↩

323. Greenwood & Earnshaw 2002, p. 552 ↩

324. Parkes & Mellor 1943, p. 740 ↩

325. Russell & Lee 2005, p. 420 ↩

326. Carapella 1968, p. 30 ↩

327. Barfuß et al. 1981, p. 967 ↩

328. 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.[338] /wiki/Amorphous_solid ↩

329. Bailar & Trotman-Dickenson 1973, p. 558; Li 1990 ↩

330. Bailar, Moeller & Kleinberg 1965, p. 477 ↩

331. Sources mentioning cationic arsenic include: Gillespie & Robinson[341] 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.[342] 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[343] 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,[344] 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[345] 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[346] 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.[347] 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". ↩

332. Eagleson 1994, p. 91 ↩

333. Cox 2004, p. 27 ↩

334. Hiller & Herber 1960, inside front cover; p. 225 ↩

335. Massey 2000, p. 267 ↩

336. Bailar, Moeller & Kleinberg 1965, p. 513 ↩

337. Timm 1944, p. 454 ↩

338. Partington 1944, p. 641; Kleinberg, Argersinger & Griswold 1960, p. 419 ↩

339. Morgan 1906, p. 163; Moeller 1954, p. 559 ↩

340. 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 consists of AsO3 pyramids and PO4 tetrahedra, joined together by all their corner atoms to form a continuous polymeric network.[353] As2(SO4)3 has a structure in which each SO4 tetrahedron is bridged by two AsO3 trigonal pyramida.[354] ↩

341. Zingaro 1994, p. 197; Emeléus & Sharpe 1959, p. 418; Addison & Sowerby 1972, p. 209; Mellor 1964, p. 337 ↩

342. Puddephatt & Monaghan 1989, p. 59 ↩

343. Pourbaix 1974, p. 521; Eagleson 1994, p. 92; Greenwood & Earnshaw 2002, p. 572 ↩

344. As2O3 is usually regarded as being amphoteric but a few sources say it is (weakly)[357] 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)[358] ↩

345. Rao 2002, p. 22 ↩

346. Krannich & Watkins 2006 ↩

347. Greenwood & Earnshaw 2002, p. 552 ↩

348. Gray, Whitby & Mann 2011 ↩

349. Greenwood & Earnshaw 2002, p. 553 ↩

350. Dunstan 1968, p. 433 ↩

351. Parise 1996, p. 112 ↩

352. Carapella 1968a, p. 23 ↩

353. Barfuß et al. 1981, p. 967 ↩

354. Antimony can also be prepared in an amorphous semiconducting black form, with an estimated (temperature-dependent) band gap of 0.06–0.18 eV.[364] /wiki/Amorphous_solid ↩

355. Dupree, Kirby & Freyland 1982, p. 604; Mhiaoui, Sar, & Gasser 2003 ↩

356. Kotz, Treichel & Weaver 2009, p. 62 ↩

357. Cotton et al. 1999, p. 396 ↩

358. King 1994, p. 174 ↩

359. Lidin[369] 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]+. ↩

360. Lindsjö, Fischer & Kloo 2004 ↩

361. Friend 1953, p. 87 ↩

362. Fesquet 1872, pp. 109–14 ↩

363. Cox 2004, p. 27 ↩

364. Hiller & Herber 1960, inside front cover; p. 225 ↩

365. Massey 2000, p. 267 ↩

366. Bailar, Moeller & Kleinberg 1965, p. 513 ↩

367. Greenwood & Earnshaw 2002, p. 553; Massey 2000, p. 269 ↩

368. King 1994, p. 171 ↩

369. Turova 2011, p. 46 ↩

370. Pourbaix 1974, p. 530 ↩

371. Puddephatt & Monaghan 1989, p. 59 ↩

372. Wiberg 2001, p. 764 ↩

373. Rao 2002, p. 22 ↩

374. House 2008, p. 497 ↩

375. Emsley 2001, p. 428 ↩

376. Gray, Whitby & Mann 2011 ↩

377. Kudryavtsev 1974, p. 78 ↩

378. Bagnall 1966, pp. 32–33, 59, 137 ↩

379. Swink et al. 1966; Anderson et al. 1980 ↩

380. Ahmed, Fjellvåg & Kjekshus 2000 ↩

381. Chizhikov & Shchastlivyi 1970, p. 28 ↩

382. Kudryavtsev 1974, p. 77 ↩

383. Stuke 1974, p. 178; Donohue 1982, pp. 386–87; Cotton et al. 1999, p. 501 ↩

384. Becker, Johnson & Nussbaum 1971, p. 56 ↩

385. Berger 1997, p. 90 ↩

386. Berger 1997, p. 90 ↩

387. Chizhikov & Shchastlivyi 1970, p. 16 ↩

388. Jolly 1966, pp. 66–67 ↩

389. Schwietzer & Pesterfield 2010, p. 239 ↩

390. Cotton et al.[393] note that TeO2 appears to have an ionic lattice; Wells[394] suggests that the Te–O bonds have "considerable covalent character". ↩

391. Wiberg 2001, p. 588 ↩

392. Mellor 1964a, p.  30; Wiberg 2001, p. 589 ↩

393. Cox 2004, p. 27 ↩

394. Hiller & Herber 1960, inside front cover; p. 225 ↩

395. Kudryavtsev 1974, p. 78 ↩

396. Greenwood & Earnshaw 2002, pp. 765–66 ↩

397. Bagnall 1966, pp. 134–51; Greenwood & Earnshaw 2002, p. 786 ↩

398. Puddephatt & Monaghan 1989, p. 59 ↩

399. Wiberg 2001, p. 764 ↩

400. Rao 2002, p. 22 ↩

401. Rao 2002, p. 22 ↩

402. Detty & O'Regan 1994, pp. 1–2 ↩

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. 320–21 ↩

410. Liquid carbon may[408] or may not[409] be a metallic conductor, depending on pressure and temperature; see also.[410] ↩

411. Atkins et al. 2006, pp. 320–21 ↩

412. Inagaki 2000, p. 216; Yasuda et al. 2003, pp. 3–11 ↩

413. O'Hare 1997, p. 230 ↩

414. 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. /wiki/Oxidising_agent ↩

415. Traynham 1989, pp. 930–31; Prakash & Schleyer 1997 ↩

416. Atkins et al. 2006, pp. 320–21 ↩

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. 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.[419] ↩

423. Keevil 1989, p. 103 ↩

424. Russell & Lee 2005, pp. 358–60 et seq ↩

425. Harding, Janes & Johnson 2002, p. 118 ↩

426. Metcalfe, Williams & Castka 1974, p. 539 ↩

427. Cobb & Fetterolf 2005, p. 64; Metcalfe, Williams & Castka 1974, p. 539 ↩

428. Ogata, Li & Yip 2002; Boyer et al. 2004, p. 1023; Russell & Lee 2005, p. 359 ↩

429. Cooper 1968, p. 25; Henderson 2000, p. 5; Silberberg 2006, p. 314 ↩

430. Wiberg 2001, p. 1014 ↩

431. Rao 2002, p. 22 ↩

432. Metcalfe, Williams & Castka 1974, p. 539 ↩

433. Hamm 1969, p. 653 ↩

434. 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." ↩

435. Holt, Rinehart & Wilson c. 2007 ↩

436. A mnemonic that captures the elements commonly recognised as metalloids goes: Up, up-down, up-down, up ... are the metalloids![431] ↩

437. Stott 1956, p. 100 ↩

438. Steele 1966, p. 60 ↩

439. Moody 1991, p. 303 ↩

440. Young et al. 2010, p. 9; Craig & Maher 2003, p. 391. Selenium is "near metalloidal". ↩

441. Rochow,[437] who later wrote his 1966 monograph The metalloids,[438] commented that, "In some respects selenium acts like a metalloid and tellurium certainly does". /wiki/Eugene_G._Rochow ↩

442. Moss 1952, p. 192 ↩

443. Glinka 1965, p. 356 ↩

444. Evans 1966, pp. 124–25 ↩

445. Regnault 1853, p. 208 ↩

446. Scott & Kanda 1962, p. 311 ↩

447. Cotton et al. 1999, pp. 496, 503–04 ↩

448. Arlman 1939; Bagnall 1966, pp. 135, 142–43 ↩

449. Glinka 1965, p. 356 ↩

450. Kozyrev 1959, p. 104; Chizhikov & Shchastlivyi 1968, p. 25; Glazov, Chizhevskaya & Glagoleva 1969, p. 86 ↩

451. Chao & Stenger 1964 ↩

452. Berger 1997, pp. 86–87 ↩

453. Berger 1997, pp. 86–87 ↩

454. Snyder 1966, p. 242 ↩

455. Fritz & Gjerde 2008, p. 235 ↩

456. Meyer et al. 2005, p. 284; Manahan 2001, p. 911; Szpunar et al. 2004, p. 17 ↩

457. US Environmental Protection Agency 1988, p. 1; Uden 2005, pp. 347‒48 ↩

458. De Zuane 1997, p. 93; Dev 2008, pp. 2‒3 ↩

459. Cotton et al. 1999, p. 502 ↩

460. Cotton et al. 1999, p. 502 ↩

461. Wiberg 2001, p. 594 ↩

462. Greenwood & Earnshaw 2002, p. 786; Schwietzer & Pesterfield 2010, pp. 242–43 ↩

463. Bagnall 1966, p. 41; Nickless 1968, p. 79 ↩

464. 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] ..." ↩

465. Legit, Friák & Šob 2010, pp. 214118–18 ↩

466. Manson & Halford 2006, pp. 378, 410 ↩

467. Bagnall 1957, p. 62; Fernelius 1982, p. 741 ↩

468. Bagnall 1966, p. 41; Barrett 2003, p. 119 ↩

469. Hawkes 2010; Holt, Rinehart & Wilson c. 2007; Hawkes 1999, p. 14; Roza 2009, p. 12 ↩

470. Keller 1985 ↩

471. Harding, Johnson & Janes 2002, p. 61 ↩

472. A further option is to include astatine both as a nonmetal and as a metalloid.[464] ↩

473. Vasáros & Berei 1985, p. 109 ↩

474. Haissinsky & Coche 1949, p. 400 ↩

475. Brownlee et al. 1950, p. 173 ↩

476. Hermann, Hoffmann & Ashcroft 2013 ↩

477. Siekierski & Burgess 2002, pp. 65, 122 ↩

478. A visible piece of astatine would be immediately and completely vaporized because of the heat generated by its intense radioactivity.[470] ↩

479. Rao & Ganguly 1986 ↩

480. The literature is contradictory as to whether boron exhibits metallic conductivity in liquid form. Krishnan et al.[472] found that liquid boron behaved like a metal. Glorieux et al.[473] characterised liquid boron as a semiconductor, on the basis of its low electrical conductivity. Millot et al.[474] reported that the emissivity of liquid boron was not consistent with that of a liquid metal. ↩

481. Vasáros & Berei 1985, p. 117 ↩

482. Kaye & Laby 1973, p. 228 ↩

483. Samsonov 1968, p. 590 ↩

484. Korenman[478] 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". ↩

485. Rossler 1985, pp. 143–44 ↩

486. Champion et al. 2010 ↩

487. Borst 1982, pp. 465, 473 ↩

488. Batsanov 1971, p. 811 ↩

489. Swalin 1962, p. 216; Feng & Lin 2005, p. 157 ↩

490. Schwietzer & Pesterfield 2010, pp. 258–60 ↩

491. Hawkes 1999, p. 14 ↩

492. Olmsted & Williams 1997, p. 328; Daintith 2004, p. 277 ↩

493. Eberle1985, pp. 213–16, 222–27 ↩

494. Restrepo et al. 2004, p. 69; Restrepo et al. 2006, p. 411 ↩

495. Craig & Maher 2003, p. 391; Schroers 2013, p. 32; Vernon 2013, pp. 1704–05 ↩

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, 192–93 ↩

501. Eggins 1972, p. 66; Rayner-Canham & Overton 2006, pp. 29–30 ↩

502. Atkins et al. 2006, pp. 320–21; Bailar et al. 1989, pp. 742–43 ↩

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. For example: intermediate electrical conductivity;[503] a relatively narrow band gap;[504] light sensitivity.[503] ↩

507. White phosphorus is the least stable and most reactive form.[505] It is also the most common, industrially important,[506] and easily reproducible allotrope, and for these three reasons is regarded as the standard state of the element.[507] ↩

508. Brescia et al. 1980, pp. 166–71 ↩

509. Fine & Beall 1990, p. 578 ↩

510. Wiberg 2001, p. 901 ↩

511. Berger 1997, p. 80 ↩

512. Lovett 1977, p. 101 ↩

513. Cohen & Chelikowsky 1988, p. 99 ↩

514. Taguena-Martinez, Barrio & Chambouleyron 1991, p. 141 ↩

515. Ebbing & Gammon 2010, p. 891 ↩

516. Vernon 2013, p. 1705 ↩

517. Asmussen & Reinhard 2002, p. 7 ↩

518. Deprez & McLachan 1988 ↩

519. Addison 1964 (P, Se, Sn); Marković, Christiansen & Goldman 1998 (Bi); Nagao et al. 2004 ↩

520. Lide 2005; Wiberg 2001, p. 423: At ↩

521. Cox 1997, pp. 182‒86 ↩

522. MacKay, MacKay & Henderson 2002, p. 204 ↩

523. Baudis 2012, pp. 207–08 ↩

524. Wiberg 2001, p. 741 ↩

525. Chizhikov & Shchastlivyi 1968, p. 96 ↩

526. Greenwood & Earnshaw 2002, pp. 140–41, 330, 369, 548–59, 749: B, Si, Ge, As, Sb, Te ↩

527. Kudryavtsev 1974, p. 158 ↩

528. Greenwood & Earnshaw 2002, pp. 271, 219, 748–49, 886: C, Al, Se, Po, At; Wiberg 2001, p. 573: Se ↩

529. 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. https://www.alfa.com/en/pure-elements/ ↩

530. United Nuclear 2013 ↩

531. Zalutsky & Pruszynski 2011, p. 181 ↩

532. 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. /wiki/Spot_price ↩