Nonmetal

In chemistry, a nonmetal (or non-metal) is a chemical element that mostly lacks the characteristics of a metal. Physically, a nonmetal tends to have a relatively low melting point, boiling point, and density. A nonmetal is typically brittle when solid and usually has poor thermal conductivity and electrical conductivity. Chemically, nonmetals tend to have relatively high ionization energy, electron affinity, and electronegativity. They gain or share electrons when they react with other elements and chemical compounds. Seventeen elements are generally classified as nonmetals: most are gases (hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon and radon); one is a liquid (bromine); and a few are solids (carbon, phosphorus, sulfur, selenium, and iodine). Metalloids such as boron, silicon, and germanium are sometimes counted as nonmetals.

Nonmetals (and metalloids) in the periodic table:


Metalloid	Reactive nonmetal	Noble gas

Metalloids are included in the legend as they generally behave chemically like nonmetals and are sometimes counted as such.

Apart from hydrogen, nonmetals are located in the p-block. Helium, as an s-block element, would normally be placed next to hydrogen and above beryllium. However, since it is a noble gas, it is instead placed above neon (in the p-block).

The nonmetals are divided into two categories reflecting their relative propensity to form chemical compounds: reactive nonmetals and noble gases. The reactive nonmetals vary in their nonmetallic character. The less electronegative of them, such as carbon and sulfur, mostly have weak to moderately strong nonmetallic properties and tend to form covalent compounds with metals. The more electronegative of the reactive nonmetals, such as oxygen and fluorine, are characterised by stronger nonmetallic properties and a tendency to form predominantly ionic compounds with metals. The noble gases are distinguished by their great reluctance to form compounds with other elements.

The distinction between categories is not absolute. Boundary overlaps, including with the metalloids, occur as outlying elements in each category show or begin to show less-distinct, hybrid-like, or atypical properties.

Although five times more elements are metals than nonmetals, two of the nonmetals—hydrogen and helium—make up over 99 percent of the observable universe.[1] Another nonmetal, oxygen, makes up almost half of the Earth's crust, oceans, and atmosphere.[2] Living organisms are composed almost entirely of nonmetals: hydrogen, oxygen, carbon, and nitrogen.[3] Nonmetals form many more compounds than metals.[4]

Definition and applicable elements

There is no rigorous definition of a nonmetal. Broadly, any element lacking a preponderance of metallic properties can be regarded as a nonmetal.

The elements generally classified as nonmetals include one element in group 1 (hydrogen); one in group 14 (carbon); two in group 15 (nitrogen and phosphorus); three in group 16 (oxygen, sulfur and selenium); most of group 17 (fluorine, chlorine, bromine and iodine); and all of group 18 (with the possible exception of oganesson).

As there is no widely agreed definition of a nonmetal, elements in the periodic table vicinity of where the metals meet the nonmetals are inconsistently classified by different authors. Elements sometimes also classified as nonmetals are the metalloids boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), and astatine (At).[5] The nonmetal selenium (Se) is sometimes instead classified as a metalloid, particularly in environmental chemistry.[6]

Properties

The marvelous variety and infinite subtlety of the non-metallic elements, their compounds, structures and reactions, is not sufficiently acknowledged in the current teaching of chemistry.

JJ Zuckerman and FC Nachod
In Steudel's Chemistry of the non-metals (1977, preface)

Nonmetals show more variability in their properties than do metals.[7] These properties are largely determined by the interatomic bonding strengths and molecular structures of the nonmetals involved, both of which are subject to variation as the number of valence electrons in each nonmetal varies. Metals, in contrast, have more homogenous structures and their properties are more easily reconciled.[8]

Physically, they largely exist as diatomic or monatomic gases, with the remainder having more substantial (open-packed) forms, unlike metals, which are nearly all solid and close-packed. If solid, they have a submetallic appearance (with the exception of sulfur) and are mostly brittle, as opposed to metals, which are lustrous, and generally ductile or malleable; they usually have lower densities than metals; are mostly poorer conductors of heat and electricity; and tend to have significantly lower melting points and boiling points than those of metals.

Scatter plot of electronegativity values and standard electrode potentials of chemically active nonmetallic elements, showing a rough correlation between the two properties. The higher the standard electrode potential, the greater is the capacity to act as an oxidizing agent.[9] The chart shows that oxygen and the nonmetallic halogens are the strongest oxidising agents and that, for the most part, the elements commonly recognised as metalloids are the weakest. The electrode potentials are for the reduction of the elements to monatomic anions (X→X⁻; X = F, Cl, Br, I, or H) or to their protonated forms (for example, O₂→H₂O; N₂→NH₃.[10]
Hydrogen and nitrogen have anomalous standard electrode potentials due to their reluctance to form anions.
A broad progression in nonmetallic character is seen, with the metalloids in the lower left, and oxygen and the nonmetallic halogens in the upper right.
Trend lines are shown with and without the anomalous hydrogen and nitrogen values. The R² values show how close each trend line fits its data points. Values range from 0.0 (indicating no fit) and 1.0 (a very good fit).

Chemically the nonmetals mostly have high ionisation energies, high electron affinities (nitrogen and the noble gases have negative electron affinities) and high electronegativity values[n 1] noting that, in general, the higher an element's ionisation energy, electron affinity, and electronegativity, the more nonmetallic that element is.[11] Nonmetals (including – to a limited extent – xenon and probably radon) usually exist as anions or oxyanions in aqueous solution; they generally form ionic or covalent compounds when combined with metals (unlike metals, which mostly form alloys with other metals); and have acidic oxides whereas the common oxides of nearly all metals are basic.

Complicating the chemistry of the nonmetals is the first row anomaly seen particularly in hydrogen, (boron), carbon, nitrogen, oxygen and fluorine; and the alternation effect seen in (arsenic), selenium and bromine.[12] The first row anomaly largely arises from the electron configurations of the elements concerned.

Hydrogen is noted for the different ways it forms bonds. It most commonly forms covalent bonds.[13] It can lose its single valence electron in aqueous solution, leaving behind a bare proton with tremendous polarising power. This subsequently attaches itself to the lone electron pair of an oxygen atom in a water molecule, thereby forming the basis of acid-base chemistry.[14] Under certain conditions a hydrogen atom in a molecule can form a second, weaker, bond with an atom or group of atoms in another molecule. Such bonding, "helps give snowflakes their hexagonal symmetry, binds DNA into a double helix; shapes the three-dimensional forms of proteins; and even raises water's boiling point high enough to make a decent cup of tea."[15]

From (boron) to neon, since the 2p subshell has no inner analogue and experiences no electron repulsion effects it consequently has a relatively small radius, unlike the 3p, 4p and 5p subshells of heavier elements[16] (a similar effect is seen in the 1s elements, hydrogen and helium). Ionisation energies and electronegativities among these elements are consequently higher than would otherwise be expected, having regard to periodic trends. The small atomic radii of carbon, nitrogen, and oxygen facilitates the formation of triple or double bonds.[17] The larger atomic radii, which enable higher coordination numbers, and lower electronegativities, which better tolerate higher positive charges, of the heavier group 15–18 nonmetals means they are able to exhibit valences other than the lowest for their group (that is, 3, 2, 1, or 0) for example in PCl₅, SF₆, IF₇, and XeF₂.[18] Period four elements immediately after the first row of the transition metals, such as selenium and bromine, have unusually small atomic radii because the 3d electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity.[19]

Comparison of properties

Characteristic and other properties of metalloids, reactive nonmetals, and noble gases are summarized in the following table. Metalloids have been included in light of their generally nonmetallic chemistry. Physical properties are listed in loose order of ease of determination; chemical properties run from general to specific, and then to descriptive.

Some properties of metalloids, reactive nonmetals, and noble gases
Physical property	Metalloid	Reactive nonmetal	Noble gas
Form	solid	solid: C, P, S, Se, I liquid: Br gaseous: H, N, O, F, Cl	gaseous
Appearance	metallic	metallic, coloured, or translucent	translucent
Elasticity	brittle	brittle if solid	soft and easily crushed when frozen
Atomic structure	close-packed* or polyatomic	polyatomic: C, P, S, Se diatomic: H, N, O, F, Cl, Br, I	monatomic
Bulk coordination number	12*, 6, 4, 3, or 2	3, 2, or 1	0
Allotropes	most form	known for C, P, O, S, Se	none form
Electrical conductivity	moderate	poor to moderate	poor
Volatility	low: B, Si, Ge, Sb, Te moderate: As, At?	low: C moderate: P, S, Se, Br, I high: H, N, O, F, Cl	high
Electronic structure	metallic* to semiconductor	semimetallic, semiconductor, or insulator	insulator
Outer s and p electrons	3–7	1, 4–7	2, 8
Crystal structure	rhombohedral: B, As, Sb cubic: Si, Ge, At? hexagonal: Te	cubic: P, O, F hexagonal: H, C, N, Se orthorhombic: S, Cl, Br, I	cubic: Ne, Ar, Kr, Xe, Rn hexagonal: He
Chemical property	Metalloid	Reactive nonmetal	Noble gas
General chemical behaviour	nonmetallic to incipient metallic		• inert to nonmetallic • Rn shows some cationic behaviour[40]
Ionization energy	low	moderate to high	high to very high
Electron affinity	low to high	moderate to high (exception: N is negative)	negative
Electronegativity	moderate: • Si < Ge ≈ B ≈ Sb < Te <As ≈ At	moderate to high: • P < Se ≈ C < S < I < Br < N < Cl < O < F	moderate to very high
Non-zero oxidation states	• negative oxidation states known for all, but for H this is an unstable state • positive oxidation states known for all but F, and only exceptionally for O • from −5 for B to +7 for Cl, Br, I, and At		• only positive oxidation states known, and only for heavier noble gases • from +2 for Kr, Xe, and Rn to +8 for Xe
Oxidising power	low (exception: At is moderate)	low to high	n/a
Catenation	marked tendency	marked tendency: C, P, S, Se less tendency: H, N, O, F, Cl, Br, I	least inclination
Compounds with metals	tend to form alloys or inter-metallic compounds	mainly covalent: H†, C, N, P, S, Se mainly ionic: O, F, Cl, Br, I	none form simple compounds
Oxides	• polymeric in structure[41] • B, Si, Ge, As, Sb, Te[42] are glass formers • tend to be amphoteric or weakly acidic[43][44]	• C, P, S, Se, and I are known in at least one polymeric form • P, S, Se are glass formers; CO₂ forms a glass at 40 GPa • acidic, or neutral (H₂O, CO, NO, N₂O)	• XeO₂ is polymeric;[45] other noble gas oxides are molecular • no glass formers • stable xenon oxides (XeO₃, XeO₄) are acidic
Sulfates	most form	some form	not known

* Bulk astatine has been predicted to have a metallic face-centred cubic structure
† Hydrogen can also form alloy-like hydrides

Properties of nonmetals (and metalloids) by group

Abbreviations used in this section are: AR Allred-Rochow; CN coordination number; and MH Moh's hardness

Group 1

Hydrogen in an electrical discharge tube

Hydrogen is a colourless, odourless, and comparatively unreactive diatomic gas with a density of 8.988 × 10⁻⁵ g/cm³ and is about 14 times lighter than air. It condenses to a colourless liquid −252.879 °C and freezes into an ice- or snow-like solid at −259.16 °C. The solid form has a hexagonal crystalline structure and is soft and easily crushed. Hydrogen is an insulator in all of its forms. It has a high ionisation energy (1312.0 kJ/mol), moderate electron affinity (73 kJ/mol), and moderate electronegativity (2.2). Hydrogen is a poor oxidising agent (H₂ + 2e⁻ → 2H^– = –2.25 V at pH 0). Its chemistry, most of which is based around its tendency to acquire the electron configuration of the noble gas helium, is largely covalent in nature, noting it can form ionic hydrides with highly electropositive metals, and alloy-like hydrides with some transition metals. The common oxide of hydrogen (H₂O) is a neutral oxide.[n 4]

Group 13

Boron

Boron is a lustrous, barely reactive solid with a density 2.34 g/cm³ (cf. aluminium 2.70), and is hard (MH 9.3) and brittle. It melts at 2076 °C (cf. steel ~1370 °C) and boils at 3927 °C. Boron has a complex rhombohedral crystalline structure (CN 5+). It is a semiconductor with a band gap of about 1.56 eV. Boron has a moderate ionisation energy (800.6 kJ/mol), low electron affinity (27 kJ/mol), and moderate electronegativity (2.04). Being a metalloid, most of its chemistry is nonmetallic in nature. Boron is a poor oxidizing agent (B₁₂ + 3e → BH₃ = –0.15 V at pH 0). While it bonds covalently in nearly all of its compounds, it can form intermetallic compounds and alloys with transition metals of the composition M_nB, if n > 2. The common oxide of boron (B₂O₃) is weakly acidic.

Group 14

Carbon, as graphite

Carbon (as graphite, its most thermodynamically stable form) is a lustrous and comparatively unreactive solid with a density of 2.267 g/cm³, and is soft (MH 0.5) and brittle. It sublimes to vapour at 3642 C°. Carbon has a hexagonal crystalline structure (CN 3). It is a semimetal in the direction of its planes, with an electrical conductivity exceeding that of some metals, and behaves as a semiconductor in the direction perpendicular to its planes. It has a high ionisation energy (1086.5 kJ/mol), moderate electron affinity (122 kJ/mol), and high electronegativity (2.55). Carbon is a poor oxidising agent (C + 4e⁻ → CH₄ = 0.13 V at pH 0). Its chemistry is largely covalent in nature, noting it can form salt-like carbides with highly electropositive metals. The common oxide of carbon (CO₂) is a medium-strength acidic oxide.

Silicon has a blue-grey metallic lustre.

Silicon is a metallic-looking relatively unreactive solid with a density of 2.3290 g/cm³, and is hard (MH 6.5) and brittle. It melts at 1414 °C (cf. steel ~1370 °C) and boils at 3265 °C. Silicon has a diamond cubic structure (CN 4). It is a semiconductor with a band gap of about 1.11 eV. Silicon has a moderate ionisation energy (786.5 kJ/mol), moderate electron affinity (134 kJ/mol), and moderate electronegativity (1.9). It is a poor oxidising agent (Si + 4e → Si₄ = –0.147 at pH 0). As a metalloid the chemistry of silicon is largely covalent in nature, noting it can form alloys with metals such as iron and copper. The common oxide of silicon (SiO₂) is weakly acidic.

Germanium

Germanium is a shiny, mostly unreactive grey-white solid with a density of 5.323 g/cm³ (about two-thirds that of iron), and is hard (MH 6.0) and brittle. It melts at 938.25 °C (cf. silver 961.78 °C) and boils at 2833 °C. Germanium has a diamond cubic structure (CN 4). It is a semiconductor with a band gap of about 0.67 eV. Germanium has a moderate ionisation energy (762 kJ/mol), moderate electron affinity (119 kJ/mol), and moderate electronegativity (2.01). It is a poor oxidising agent (Ge + 4e → GeH₄ = –0.294 at pH 0). As a metalloid the chemistry of germanium is largely covalent in nature, noting it can form alloys with metals such as aluminium and gold. Most alloys of germanium with metals lack metallic or semimetallic conductivity. The common oxide of germanium (GeO₂) is amphoteric.

Group 15

Liquid nitrogen

Nitrogen is a colourless, odourless, and relatively inert diatomic gas with a density of 1.251 × 10⁻³ g/cm³ (marginally heavier than air). It condenses to a colourless liquid at −195.795 °C and freezes into an ice- or snow-like solid at −210.00 °C. The solid form (density 0.85 g/cm³; cf. lithium 0.534) has a hexagonal crystalline structure and is soft and easily crushed. Nitrogen is an insulator in all of its forms. It has a high ionisation energy (1402.3 kJ/mol), low electron affinity (–6.75 kJ/mol), and high electronegativity (3.04). The latter property manifests in the capacity of nitrogen to form usually strong hydrogen bonds, and its preference for forming complexes with metals having low electronegativities, small cationic radii, and often high charges (+3 or more). Nitrogen is a poor oxidising agent (N₂ + 6e⁻ → 2NH₃ = −0.057 V at pH 0). Only when it is in a positive oxidation state, that is, in combination with oxygen or fluorine, are its compounds good oxidising agents, for example, 2NO₃⁻ → N₂ = 1.25 V. Its chemistry is largely covalent in nature; anion formation is energetically unfavourable owing to strong inter electron repulsions associated with having three unpaired electrons in its outer valence shell, hence its negative electron affinity. The common oxide of nitrogen (NO) is weakly acidic. Many compounds of nitrogen are less stable than diatomic nitrogen, so nitrogen atoms in compounds seek to recombine if possible and release energy and nitrogen gas in the process, which can be leveraged for explosive purposes.

Phosphorus, as black phosphorus

Phosphorus in its most thermodynamically stable black form, is a lustrous and comparatively unreactive solid with a density of 2.69 g/cm³, and is soft (MH 2.0) and has a flaky comportment. It sublimes at 620 °C. Black phosphorus has an orthorhombic crystalline structure (CN 3). It is a semiconductor with a band gap of 0.3 eV. It has a high ionisation energy (1086.5 kJ/mol), moderate electron affinity (72 kJ/mol), and moderate electronegativity (2.19). In comparison to nitrogen, phosphorus usually forms weak hydrogen bonds, and prefers to form complexes with metals having high electronegativities, large cationic radii, and often low charges (usually +1 or +2. Phosphorus is a poor oxidising agent (P₄ + 3e⁻ → PH₃^– = −0.046 V at pH 0 for the white form, −0.088 V for the red). Its chemistry is largely covalent in nature, noting it can form salt-like phosphides with highly electropositive metals. Compared to nitrogen, electrons have more space on phosphorus, which lowers their mutual repulsion and results in anion formation requiring less energy. The common oxide of phosphorus (P₂O₅) is a medium-strength acidic oxide.

White phosphorus, stored under water to prevent its oxidation[46]

When assessing periodicity in the properties of the elements it needs to be borne in mind that the quoted properties of phosphorus tend to be those of its least stable white form rather than, as is the case with all other elements, the most stable form. White phosphorus is the most common, industrially important, and easily reproducible allotrope. For those reasons it is the standard state of the element. Paradoxically, it is also thermodynamically the least stable, as well as the most volatile and reactive form. It gradually changes to red phosphorus. This transformation is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and, accordingly, appear yellow. For this reason, white phosphorus that is aged or otherwise impure is also called yellow phosphorus. When exposed to oxygen, white phosphorus glows in the dark with a very faint tinge of green and blue. It is highly flammable and pyrophoric (self-igniting) upon contact with air. White phosphorus has a density of 1.823 g/cm³, is soft (MH 0.5) as wax, pliable and can be cut with a knife. It melts at 44.15 °C and, if heated rapidly, boils at 280.5 °C; it otherwise remains solid and transforms to violet phosphorus at 550 °C. It has a body-centred cubic structure, analogous to that of manganese, with unit cell comprising 58 P₄ molecules. It is an insulator with a band gap of about 3.7 eV.

Arsenic, sealed in a container to prevent tarnishing

Arsenic is a grey, metallic looking solid which is stable in dry air but develops a golden bronze patina in moist air, which blackens on further exposure. It has a density of 5.727 g/cm³, and is brittle and moderately hard (MH 3.5; more than aluminium; less than iron). Arsenic sublimes at 615 °C. It has a rhombohedral polyatomic crystalline structure (CN 3). Arsenic is a semimetal, with an electrical conductivity of around 3.9 × 10⁴ S•cm⁻¹ and a band overlap of 0.5 eV. It has a moderate ionisation energy (947 kJ/mol), moderate electron affinity (79 kJ/mol), and moderate electronegativity (2.18). Arsenic is a poor oxidising agent (As + 3e → AsH₃ = –0.22 at pH 0). As a metalloid, its chemistry is largely covalent in nature, noting it can form brittle alloys with metals, and has an extensive organometallic chemistry. Most alloys of arsenic with metals lack metallic or semimetallic conductivity. The common oxide of arsenic (As₂O₃) is acidic but weakly amphoteric.

Antimony, showing its brilliant lustre

Antimony is a silver-white solid with a blue tint and a brilliant lustre. It is stable in air and moisture at room temperature. Antimony has a density of 6.697 g/cm³, and is moderately hard (MH 3.0; about the same as copper). It has a rhombohedral crystalline structure (CN 3). Antimony melts at 630.63 °C and boils at 1635 °C. It is a semimetal, with an electrical conductivity of around 3.1 × 10⁴ S•cm⁻¹ and a band overlap of 0.16 eV. Antimony has a moderate ionisation energy (834 kJ/mol), moderate electron affinity (101 kJ/mol), and moderate electronegativity (2.05). It is a poor oxidising agent (Sb + 3e → SbH₃ = –0.51 at pH 0). As a metalloid, its chemistry is largely covalent in nature, noting it can form alloys with one or more metals such as aluminium, iron, nickel, copper, zinc, tin, lead and bismuth, and has an extensive organometallic chemistry. Most alloys of antimony with metals have metallic or semimetallic conductivity. The common oxide of antimony (Sb₂O₃) is amphoteric.

Group 16

Liquid oxygen (boiling)

In the United States alone, more than $10 billion is lost each year to corrosion...Much of this corrosion is the rusting of iron and steel...The oxidizing agent causing all of this corrosion is usually oxygen.

MD Joesten, L Hogg, and ME Castellion
In The world of chemistry (2007, p. 217)

Oxygen is a colourless, odourless, and unpredictably reactive diatomic gas with a gaseous density of 1.429 × 10⁻³ g/cm³ (marginally heavier than air). It is generally unreactive at room temperature. Thus, sodium metal will "retain its metallic lustre for days in the presence of absolutely dry air and can even be melted (m.p. 97.82 °C) in the presence of dry oxygen without igniting".[47] On the other hand, oxygen can react with many inorganic and organic compounds either spontaneously or under the right conditions,[48] (such as a flame or a spark) [or ultra-violet light?]. It condenses to pale blue liquid −182.962 °C and freezes into a light blue solid at −218.79 °C. The solid form (density 0.0763 g/cm³) has a cubic crystalline structure and is soft and easily crushed. Oxygen is an insulator in all of its forms. It has a high ionisation energy (1313.9 kJ/mol), high electron affinity (141 kJ/mol), and high electronegativity (3.44). Oxygen is a strong oxidising agent (O₂ + 4e → 2H₂O = 1.23 V at pH 0). Metal oxides are largely ionic in nature.[49]

Sulfur

Sulfur is a bright-yellow moderately reactive[50] solid. It has a density of 2.07 g/cm³ and is soft (MH 2.0) and brittle. It melts to a light yellow liquid 95.3 °C and boils at 444.6 °C. Sulfur has an abundance on earth one-tenth that of oxygen. It has an orthorhombic polyatomic (CN 2) crystalline structure, and is brittle. Sulfur is an insulator with a band gap of 2.6 eV, and a photoconductor meaning its electrical conductivity increases a million-fold when illuminated. Sulfur has a moderate ionisation energy (999.6 kJ/mol), moderate electron affinity (200 kJ/mol), and high electronegativity (2.58). It is a poor oxidising agent (S₈ + 2e⁻ → H₂S = 0.14 V at pH 0). The chemistry of sulfur is largely covalent in nature, noting it can form ionic sulfides with highly electropositive metals. The common oxide of sulfur (SO₃) is strongly acidic.

Selenium

Selenium is a metallic-looking, moderately reactive[50] solid with a density of 4.81 g/cm³ and is soft (MH 2.0) and brittle. It melts at 221 °C to a black liquid and boils at 685 °C to a dark yellow vapour. Selenium has a hexagonal polyatomic (CN 2) crystalline structure. It is a semiconductor with a band gap of 1.7 eV, and a photoconductor meaning its electrical conductivity increases a million-fold when illuminated. Selenium has a moderate ionisation energy (941.0 kJ/mol), high electron affinity (195 kJ/mol), and high electronegativity (2.55). It is a poor oxidising agent (Se + 2e⁻ → H₂Se = −0.082 V at pH 0). The chemistry of selenium is largely covalent in nature, noting it can form ionic selenides with highly electropositive metals. The common oxide of selenium (SeO₃) is strongly acidic.

Tellurium

Tellurium is a silvery-white, moderately reactive,[50] shiny solid, that has a density of 6.24 g/cm³ and is soft (MH 2.25) and brittle. It is the softest of the commonly recognised metalloids. 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. It has a melting point of 450 °C and a boiling point of 988 °C. Tellurium has a polyatomic (CN 2) hexagonal crystalline structure. It is a semiconductor with a band gap of 0.32 to 0.38 eV. Tellurium has a moderate ionisation energy (869.3 kJ/mol), high electron affinity (190 kJ/mol), and moderate electronegativity (2.1). It is a poor oxidising agent (Te + 2e⁻ → H₂Te = −0.45 V at pH 0). The chemistry of tellurium is largely covalent in nature, noting it has an extensive organometallic chemistry and that many tellurides can be regarded as metallic alloys. The common oxide of tellurium (TeO₂) is amphoteric.

Group 17

Liquid fluorine, in a cryogenic bath

Fluorine is an extremely toxic and reactive pale yellow diatomic gas that, with a gaseous density of 1.696 × 10⁻³ g/cm³, is about 40% heavier than air. Its extreme reactivity is such that it was not isolated (via electrolysis) until 1886 and was not isolated chemically until 1986. Its occurrence in an uncombined state in nature was first reported in 2012, but is contentious. Fluorine condenses to a pale yellow liquid at −188.11 °C and freezes into a colourless solid[47] at −219.67 °C. The solid form (density 1.7 g/cm⁻³) has a cubic crystalline structure and is soft and easily crushed. Fluorine is an insulator in all of its forms. It has a high ionisation energy (1681 kJ/mol), high electron affinity (328 kJ/mol), and high electronegativity (3.98). Fluorine is a powerful oxidising agent (F₂ + 2e → 2HF = 2.87 V at pH 0); "even water, in the form of steam, will catch fire in an atmosphere of fluorine".[51] Metal fluorides are generally ionic in nature.

Chlorine gas

Chlorine is an irritating green-yellow diatomic gas that is extremely reactive, and has a gaseous density of 3.2 × 10⁻³ g/cm³ (about 2.5 times heavier than air). It condenses at −34.04 °C to an amber-coloured liquid and freezes at −101.5 °C into a yellow crystalline solid. The solid form (density 1.9 g/cm⁻³) has an orthorhombic crystalline structure and is soft and easily crushed. Chlorine is an insulator in all of its forms. It has a high ionisation energy (1251.2 kJ/mol), high electron affinity (349 kJ/mol; higher than fluorine), and high electronegativity (3.16). Chlorine is a strong oxidising agent (Cl₂ + 2e → 2HCl = 1.36 V at pH 0). Metal chlorides are largely ionic in nature. The common oxide of chlorine (Cl₂O₇) is strongly acidic.

Liquid bromine

Bromine is a deep brown diatomic liquid that is quite reactive, and has a liquid density of 3.1028 g/cm³. It boils at 58.8 °C and solidifies at −7.3 °C to an orange crystalline solid (density 4.05 g/cm⁻³). It is the only element, apart from mercury, known to be a liquid at room temperature. The solid form, like chlorine, has an orthorhombic crystalline structure and is soft and easily crushed. Bromine is an insulator in all of its forms. It has a high ionisation energy (1139.9 kJ/mol), high electron affinity (324 kJ/mol), and high electronegativity (2.96). Bromine is a strong oxidising agent (Br₂ + 2e → 2HBr = 1.07 V at pH 0). Metal bromides are largely ionic in nature. The unstable common oxide of bromine (Br₂O₅) is strongly acidic.

Iodine crystals

Iodine, the rarest of the nonmetallic halogens, is a metallic looking solid that is moderately reactive, and has a density of 4.933 g/cm³. It melts at 113.7 °C to a brown liquid and boils at 184.3 °C to a violet-coloured vapour. It has an orthorhombic crystalline structure with a flaky habit. Iodine is semiconductor in the direction of its planes, with a band gap of about 1.3 eV and a conductivity of 1.7 × 10⁻⁸ S•cm⁻¹ at room temperature. This is higher than selenium but lower than boron, the least electrically conducting of the recognised metalloids. Iodine is an insulator in the direction perpendicular to its planes. It has a high ionisation energy (1008.4 kJ/mol), high electron affinity (295 kJ/mol), and high electronegativity (2.66). Iodine is a moderately strong oxidising agent (I₂ + 2e → 2I⁻ = 0.53 V at pH 0). Metal iodides are predominantly ionic in nature. The only stable oxide of iodine (I₂O₅) is strongly acidic.

Group 18

Liquified helium

Helium has a density of 1.785 × 10⁻⁴ g/cm³ (cf. air 1.225 × 10⁻³ g/cm³), liquifies at −268.928 °C, and cannot be solidified at normal pressure. It has the lowest boiling point of all of the elements. Liquid helium exhibits super-fluidity, superconductivity, and near-zero viscosity; its thermal conductivity is greater than that of any other known substance (more than 1,000 times that of copper). Helium can only be solidified at −272.20 °C under a pressure of 2.5 MPa. It has a very high ionisation energy (2372.3 kJ/mol), low electron affinity (estimated at −50 kJ/mol), and very high electronegativity (5.5 AR). No normal compounds of helium have so far been synthesised.

Neon in an electrical discharge tube

Neon has a density of 9.002 × 10⁻⁴ g/cm³, liquifies at −245.95 °C, and solidifies at −248.45 °C. It has the narrowest liquid range of any element and, in liquid form, has over 40 times the refrigerating capacity of liquid helium and three times that of liquid hydrogen. Neon has a very high ionisation energy (2080.7 kJ/mol), low electron affinity (estimated at −120 kJ/mol), and very high electronegativity (4.84 AR). It is the least reactive of the noble gases; no normal compounds of neon have so far been synthesised.

A small piece of rapidly melting solid argon

Argon has a density of 1.784 × 10⁻³ g/cm³, liquifies at −185.848 °C, and solidifies at −189.34 °C. Although non-toxic, it is 38% denser than air and therefore considered a dangerous asphyxiant in closed areas. It is difficult to detect because (like all the noble gases) it is colourless, odourless, and tasteless. Argon has a high ionisation energy (1520.6 kJ/mol), low electron affinity (estimated at −96 kJ/mol), and high electronegativity (3.2 AR). One interstitial compound of argon, Ar₁C₆₀ is a stable solid at room temperature.

A Kr-shaped krypton discharge tube

Krypton has a density of 3.749 × 10⁻³ g/cm³, liquifies at −153.415 °C, and solidifies at −157.37 °C. It has a high ionisation energy (1350.8 kJ/mol), low electron affinity (estimated at −60 kJ/mol), and high electronegativity (2.94 AR). Krypton can be reacted with fluorine to form the difluoride, KrF₂. The reaction of KrF
₂ with B(OTeF
₅)
₃ produces an unstable compound, Kr(OTeF
₅)
₂, that contains a krypton-oxygen bond.

Pressurized xenon gas encapsulated in an acrylic cube

Xenon has a density of 5.894 × 10⁻³ g/cm³, liquifies at −161.4 °C, and solidifies at −165.051 °C. It is non-toxic, and belongs to a select group of substances that penetrate the blood–brain barrier, causing mild to full surgical anesthesia when inhaled in high concentrations with oxygen. Xenon has a high ionisation energy (1170.4 kJ/mol), low electron affinity (estimated at −80 kJ/mol), and high electronegativity (2.4 AR). It forms a relatively large number of compounds, mostly containing fluorine or oxygen. An unusual ion containing xenon is the tetraxenonogold(II) cation, AuXe²⁺
₄, which contains Xe–Au bonds. This ion occurs in the compound AuXe
₄(Sb
₂F
₁₁)
₂, and is remarkable in having direct chemical bonds between two notoriously unreactive atoms, xenon and gold, with xenon acting as a transition metal ligand. The compound Xe
₂Sb
₂F
₁₁ contains a Xe–Xe bond, the longest element-element bond known (308.71 pm = 3.0871 Å). The most common oxide of xenon (XeO₃) is strongly acidic.

Radon, which is radioactive, has a density of 9.73 × 10⁻³ g/cm³, liquifies at −61.7 °C, and solidifies at −71 °C. It has a high ionisation energy (1037 kJ/mol), low electron affinity (estimated at −70 kJ/mol), and moderate electronegativity (2.06 AR). The only confirmed compounds of radon, which is the rarest of the naturally occurring noble gases, are the difluoride RnF₂, and trioxide, RnO₃. It has been reported that radon is capable of forming a simple Rn²⁺ cation in halogen fluoride solution, which is highly unusual behaviour for a nonmetal, and a noble gas at that. Radon trioxide (RnO₃) is expected to be acidic.

Oganesson, the heaviest element on the periodic table, has only recently been synthesized. Owing to its short half-life, its chemical properties have not yet been investigated. Due to the significant relativistic destabilisation of the 7p_3/2 orbitals, it is expected to be significantly reactive and behave more similarly to the group 14 elements, as it effectively has four valence electrons outside a pseudo-noble gas core. Its boiling point is expected to be about 80±30 °C, so that it is probably neither noble nor a gas; as a liquid it is expected to have a density of about 5 g/cm³. It is expected to have a barely positive electron affinity (estimated as 5 kJ/mol) and a moderate ionisation energy of about 860 kJ/mol, which is rather low for a nonmetal and close to those of the metalloids tellurium and astatine. The oganesson fluorides OgF₂ and OgF₄ are expected to show significant ionic character, suggesting that oganesson may have at least incipient metallic properties. The oxides of oganesson, OgO and OgO₂, are predicted to be amphoteric.

Cross-cutting relationships

Periodic table extract showing some relationships among the nonmetals. The dashed line around H denotes that H is normally positioned on the far left of the periodic table, above Li in group 1. The red arrows denote that, as with the metalloids, the most stable forms of C, P, Se, and I each have a metallic appearance. The white arrow denotes that N, S, and Br are a gas, solid, and liquid, respectively. That leaves the triangle of O, F, and Cl representing the most corrosive nonmetals. Not shown here are At (a metalloid, predicted to be a post-transition metal), Rn (a noble gas, showing incipient metallic behavior), and Og (possibly a metalloid).

Some pairs of nonmetals show additional relationships, beyond those associated with group membership.

H and C. Hydrogen in group 1, and carbon in group 14, show some out-of-group similarities.[52] These include proximity in ionization energies, electron affinities and electronegativity values; half-filled valence shells; and correlations between the chemistry of H–H and C–H bonds.

H and N. Both are relatively unreactive colourless diatomic gases, with comparably high ionization energies (1312.0 and 1402.3 kJ/mol), each having half-valence subshells, 1s and 2p respectively. Like the reactive azide N₃⁻ anion, inter-electron repulsions in the H⁻ hydride anion (with its single nuclear charge) make ionic hydrides highly reactive. Unusually for nonmetals, the two elements are known in cationic forms. In water the H⁺ "cation" exists as an H₁₃O₆⁺ ion, with a delocalised proton in a central OHO group.[53] Nitrogen forms an N₅⁺ pentazenium cation; bulk quantities of the salt N₅⁺SbF₆⁻ can be prepared. Coincidentally, the NH₄⁺ ammonium cation behaves in many respects as an alkali metal anion.[54]

C and P. Carbon and phosphorus represent an example of a less-well known diagonal relationship, especially in organic chemistry. "Spectacular" evidence of this relationship was provided in 1987 with the synthesis of a ferrocene-like molecule in which six of the carbon atoms were replaced by phosphorus atoms.[54] Further illustrating the theme is the "extraordinary" similarity between low coordinate phosphorus compounds and unsaturated carbon compounds, and related research into organophosphorus chemistry.[55] In 2020, the first compound containing three carbon atoms and one phosphorus arranged in a tetrahedron, tri-tert-butyl phosphatetrahedrane, (PC₃)(C₄H₉)₃ was synthesised. While plain all-carbon tetrahedrane (CH)₄ has never been isolated, phosphorus was selected in light of its capacity to form tetrahedral molecules, and the similarity of some of its properties to those of carbon.[56]

C and N. With nitrogen, carbon forms an extensive series of nitride compounds including those with high N:C ratios, and with structures that are simple (CN₁₂); chain-like (C₆N₂ for example); graphitic (linked C₆N₇ units); fullerenic (C₄₈N₁₂) or polymeric (C₃N₃ units). Most of the compounds prepared to date also contain quantities of hydrogen.[57]

N and P. Like nitrogen, the chemistry of phosphorus is that of the covalent bond; the two nonmetals rarely form anions. Despite them being in the same group, and the composition of some of their compounds resembling one another, the individual chemistries of nitrogen and phosphorus are very different.[58] That said, the two elements form an extensive series of phosphorus–nitrogen compounds having chain, ring and cage structures; the P–N repeat unit in these structures bears a strong resemblance to the S–N repeat unit found in the wide range of sulfur–nitrogen compounds, discussed next.[59]

N and S. Nitrogen and sulfur have a less-well known diagonal relationship, manifested in like charge densities and electronegativities (the latter are identical if only the p electrons are counted; see Hinze and Jaffe 1962) especially when sulfur is bonded to an electron-withdrawing group. They are able to form an extensive series of seemingly interchangeable sulfur nitrides, the most famous of which, polymeric sulfur nitride, is metallic, and a superconductor below 0.26 K. The aromatic nature of the S₃N₂²⁺ ion, in particular, serves as an "exemplar" of the similarity of electronic energies between the two nonmetals.[54]

N and O. Nitrogen and oxygen represent the main parts of air. They both become toxic under pressure thus, nitrogen narcosis; oxygen narcosis. They react readily with one another. Nitrogen forms several oxides, including nitrous oxide, N₂O, in which nitrogen is in the +1 oxidation state; nitric oxide, NO, in which it is in the +2 state; and nitrogen dioxide, NO₂, in which it is in the +4 state.

Many of the nitrogen oxides are extremely volatile; they are prime sources of pollution in the atmosphere. Nitrous oxide, also known as laughing gas, is sometimes used as an anaesthetic; when inhaled it produces mild hysteria. Nitric oxide reacts rapidly with oxygen to form brown nitrogen dioxide, an intermediate in the manufacture of nitric acid and a powerful oxidizing agent utilized in chemical processes and rocket fuels.

More generally nitrogen resembles oxygen with its high electronegativity and concomitant capability for hydrogen bonding and the ability to form coordination complexes by donating its lone pairs of electrons. There are some parallels between the chemistry of ammonia NH₃ and water H₂O. For example, the capacity of both compounds to be pronated to give NH₄⁺ and H₃O⁺ or deprotonated to give NH₂⁻ and OH⁻, with all of these able to be isolated in solid compounds.

O and S. Oxygen and sulfur react readily with one another, forming lower sulfur oxides (S_nO, S₇O₂ and S₆O₂); sulfur monoxide (SO) and its dimer, disulfur dioxide (S₂O₂); sulfur dioxide (SO₂); sulfur trioxide (SO₃); higher sulfur oxides (SO₃ and SO₄ and polymeric condensates of them); and disulfur monoxide (S₂O). The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO₂) that reacts with atmospheric water and oxygen to produce sulfuric acid (H₂SO₄) and sulfurous acid (H₂SO₃). These acids are components of acid rain, lowering the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. In most oxygen-containing organic molecules, the oxygen atoms can be replaced by sulfur atoms.

O and Cl. "Chlorination reactions have many similarities to oxidation reactions. They tend not to be limited to thermodynamic equilibrium and often go to complete chlorination. The reactions are often highly exothermic. Chlorine, like oxygen, forms flammable mixtures with organic compounds."[60]

O and F. Fluorine and oxygen share the ability to often bring out the highest oxidation states among the elements.

P and S (Se). Phosphorus reacts with sulfur and selenium (and oxygen) to form a large number of compounds. These compounds are characterized by structural analogies derived from the white phosphorus P₄ tetrahedron.[61]

S and Se Commonalties between sulfur and selenium are abundantly obvious. For example, selenium is found in metal sulfide ores, where it partially replaces sulfur; both elements are photoconductors—their electrical conductivities increase by up to six orders of magnitude when exposed to light.[62]

I and Xe. The chemistry of iodine in its oxidation states of +1, +3, +5, and +7 is analogous to that of xenon in an immediately higher oxidation state.

Allotropes

Some allotropes of carbon

Many nonmetals have less stable allotropes, with either nonmetallic or metallic properties. Graphite, the standard state of carbon, has a lustrous appearance and is a fairly good electrical conductor. The diamond allotrope of carbon is clearly nonmetallic, however, being translucent and having a relatively poor electrical conductivity. Carbon is also known in several other allotropic forms, including semiconducting buckminsterfullerene (C₆₀). Nitrogen can form gaseous tetranitrogen (N₄), an unstable polyatomic molecule with a lifetime of about one microsecond.[63] Oxygen is a diatomic molecule in its standard state; it also exists as ozone (O₃), an unstable nonmetallic allotrope with a half-life of around half an hour.[64] Phosphorus, uniquely, exists in several allotropic forms that are more stable than that of its standard state as white phosphorus (P₄). The red and black allotropes are probably the best known; both are semiconductors. Phosphorus is also known as diphosphorus (P₂), an unstable diatomic allotrope.[65] Sulfur has more allotropes than any other element;[66] all of these, except plastic sulfur (a metastable ductile mixture of allotropes)[67] have nonmetallic properties. Selenium has several nonmetallic allotropes, all of which are much less electrically conducting than its standard state of grey "metallic" selenium.[68] Iodine is also known in a semiconducting amorphous form.[69] Under sufficiently high pressures, just over half of the nonmetals, starting with phosphorus at 1.7 GPa,[70] have been observed to form metallic allotropes.

Most metalloids, like the less electronegative nonmetals, form allotropes. Boron is known in several crystalline and amorphous forms. The discovery of a quasispherical allotropic molecule borospherene (B₄₀) was announced in July 2014. Silicon was most recently known only in its crystalline and amorphous forms. Silicene, a two-dimensional allotrope of silicon, with a hexagonal honeycomb structure similar to that of graphene, was observed in 2010. The synthesis of an orthorhombic allotrope Si₂₄, was subsequently reported in 2014. At pressure of ~10–11 GPa, germanium transforms to a metallic phase with the same tetragonal structure as tin; when decompressed—and depending on the speed of pressure release—metallic germanium forms a series of allotropes that are metastable at ambient condition. Germanium also forms a graphene analogue, germanene. Arsenic and antimony form several well known allotropes (yellow, grey, and black). Tellurium is known only in its crystalline and amorphous forms; astatine is not known to have any allotropes.

Abundance and extraction

Fluorite, a source of fluorine

Hydrogen and helium are estimated to make up approximately 99 per cent of all ordinary matter in the universe. Less than five percent of the Universe is believed to be made of ordinary matter, represented by stars, planets and living beings. The balance is made of dark energy and dark matter, both of which are poorly understood at present.[71]

Hydrogen, carbon, nitrogen, and oxygen constitute the great bulk of the Earth's atmosphere, oceans, crust, and biosphere; the remaining nonmetals have abundances of 0.5 per cent or less. In comparison, 35 per cent of the crust is made up of the metals sodium, magnesium, aluminium, potassium and iron; together with a metalloid, silicon. All other metals and metalloids have abundances within the crust, oceans or biosphere of 0.2 per cent or less.[72]

Nonmetals, and metalloids, in their elemental forms are extracted from:[73] brine: Cl, Br, I; liquid air: N, O, Ne, Ar, Kr, Xe; minerals: B (borate minerals); C (coal; diamond; graphite); F (fluorite); Si (silica) P (phosphates); Sb (stibnite, tetrahedrite); I (in sodium iodate NaIO₃ and sodium iodide NaI); natural gas: H, He, S; and from ores, as processing byproducts: Ge (zinc ores); As (copper and lead ores); Se, Te (copper ores); and Rn (uranium bearing ores). Astatine is produced in minute quantities by irradiating bismuth.

Applications in common

For prevalent and speciality applications of individual nonmetals see the main article for each element.

A high-voltage circuit-breaker employing sulfur hexafluoride SF₆ as its inert (air replacement) interrupting medium[74]

Nonmetals do not have any universal or near-universal applications. This is not the case with metals, most of which have structural uses; nor the metalloids, the typical uses of which extend to (for example) oxide glasses, alloying components, and semiconductors.

Shared applications of different subsets of the nonmetals instead encompass their presence in, or specific uses in the fields of cryogenics and refrigerants: H, He, N, O, F and Ne; fertilisers: H, N, P, S, Cl (as a micronutrient) and Se; household accoutrements: H (primary constituent of water), He (party balloons), C (in pencils, as graphite), N (beer widgets), O (as peroxide, in detergents), F (as fluoride, in toothpaste), Ne (lighting), P (matches), S (garden treatments), Cl (bleach constituent), Ar (insulated windows), Se (glass; solar cells), Br (as bromide, for purification of spa water), Kr (energy saving fluorescent lamps), I (in antiseptic solutions), Xe (in plasma TV display cells, a technology subsequently made redundant by low cost OLED displays), while Rn also sometimes occurs, but then as an unwanted, potentially hazardous indoor pollutant;[75] industrial acids: C, N, F, P, S and Cl; inert air replacements: N, Ne, S (in sulfur hexafluoride SF₆), Ar, Kr and Xe; lasers and lighting: He, C (in carbon dioxide lasers, CO₂), N, O (in a chemical oxygen iodine laser), F (in a hydrogen fluoride laser, HF), Ne, S (in a sulfur lamp), Ar, Kr and Xe; and medicine and pharmaceuticals: He, O, F, Cl, Br, I, Xe and Rn.

The number of compounds formed by nonmetals is vast.[76] The first nine places in a "top 20" table of elements most frequently encountered in 8,427,300 compounds, as listed in the Chemical Abstracts Service register for July 1987, were occupied by nonmetals. Hydrogen, carbon, oxygen and nitrogen were found in the majority (greater than 64 per cent) of compounds. Silicon, a metalloid, was in 10th place. The highest rated metal, with an occurrence frequency of 2.3 per cent, was iron, in 11th place.[77]

Discovery

Antiquity: C, S, (Sb)

Carbon, sulfur, and antimony were known in antiquity. The earliest known use of charcoal dates to around 3750 BCE. The Egyptians and Sumerians employed it for the reduction of copper, zinc, and tin ores in the manufacture of bronze. Diamonds were probably known from as early as 2500 BCE. The first true chemical analyses were made in the 18th century; Lavoisier recognized carbon as an element in 1789. Sulfur usage dates from before 2500 BCE; it was recognized as an element by Antoine Lavoisier in 1777. Antimony usage was concurrent with that of sulfur; the Louvre holds a 5,000 year old vase made of almost pure antimony.

13th century: (As)

Albertus Magnus (Albert the Great, 1193–1280) is believed to have been the first to isolate the element from a compound in 1250, by heating soap together with arsenic trisulfide. If so, it was the first element to be chemically discovered.

17th century: P

Phosphorus was prepared from urine, by Hennig Brand, in 1669.

18th century: H, O, N, (Te), Cl

Hydrogen: Cavendish, in 1766, was the first to distinguish hydrogen from other gases, although Paracelsus around 1500, Robert Boyle (1670), and Joseph Priestley (?) had observed its production by reacting strong acids with metals. Lavoisier named it in 1793. Oxygen: Carl Wilhelm Scheele obtained oxygen by heating mercuric oxide and nitrates in 1771, but did not publish his findings until 1777. Priestley also prepared this new "air" by 1774, but only Lavoisier recognized it as a true element; he named it in 1777. Nitrogen: Rutherford discovered nitrogen while he was studying at the University of Edinburgh. He showed that the air in which animals breathed, after removal of exhaled carbon dioxide, was no longer able to burn a candle. Scheele, Henry Cavendish, and Priestley also studied this element at about the same time; Lavoisier named it in 1775 or 1776. Tellurium: In 1783, Franz-Joseph Müller von Reichenstein, who was then serving as the Austrian chief inspector of mines in Transylvania, concluded that a new element was present in a gold ore from the mines in Zlatna, near today's city of Alba Iulia, Romania. In 1789, a Hungarian scientist, Pál Kitaibel, discovered the element independently in an ore from Deutsch-Pilsen that had been regarded as argentiferous molybdenite, but later he gave the credit to Müller. In 1798, it was named by Martin Heinrich Klaproth, who had earlier isolated it from the mineral calaverite. Chlorine: In 1774, Scheele obtained chlorine from hydrochloric acid but thought it was an oxide. Only in 1808 did Humphry Davy recognize it as an element.

Early 19th century: (B) I, Se, (Si), Br

Boron was identified by Sir Humphry Davy in 1808 but not isolated in a pure form until 1909, by the American chemist Ezekiel Weintraub. Iodine was discovered in 1811 by Courtois from the ashes of seaweed. Selenium: In 1817, when Berzelius and Johan Gottlieb Gahn were working with lead they discovered a substance that was similar to tellurium. After more investigation Berzelius concluded that it was a new element, related to sulfur and tellurium. Because tellurium had been named for the Earth, Berzelius named the new element "selenium", after the moon. Silicon: In 1823, Berzelius prepared amorphous silicon by reducing potassium fluorosilicate with molten potassium metal. Bromine: Balard and Gmelin both discovered bromine in the autumn of 1825 and published their results in the following year.

Late 19th century: He, F, (Ge), Ar, Kr, Ne, Xe

Helium: In 1868, Janssen and Lockyer independently observed a yellow line in the solar spectrum that did not match that of any other element. In 1895, in each case at around the same time, Ramsay, Cleve, and Langlet independently observed helium trapped in cleveite. Fluorine: André-Marie Ampère predicted an element analogous to chlorine obtainable from hydrofluoric acid, and between 1812 and 1886 many researchers tried to obtain it. Fluorine was eventually isolated in 1886 by Moissan. Germanium: In mid-1885, at a mine near Freiberg, Saxony, a new mineral was discovered and named argyrodite because of its silver content. The chemist Clemens Winkler analyzed this new mineral, which proved to be a combination of silver, sulfur, and a new element, germanium, which he was able to isolate in 1886. Argon: Lord Rayleigh and Ramsay discovered argon in 1894 by comparing the molecular weights of nitrogen prepared by liquefaction from air, and nitrogen prepared by chemical means. It was the first noble gas to be isolated. Krypton, neon, and xenon: In 1898, within a period of three weeks, Ramsay and Travers successively separated krypton, neon and xenon from liquid argon by exploiting differences in their boiling points.

20th century: Rn, (At)

In 1898, Friedrich Ernst Dorn discovered a radioactive gas resulting from the radioactive decay of radium; Ramsay and Robert Whytlaw-Gray subsequently isolated radon in 1910. Astatine was synthesised in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè. They bombarded bismuth-209 with alpha particles in a cyclotron to produce, after emission of two neutrons, astatine-211.

Notes

An ionisation energy of less than 750 kJ/mol is taken to be low, 750–1000 is moderate, and > 1000 is high (> 2000 is very high); an electron affinity of less than 70 kJ/mol is taken to be low, 70–140 is moderate, and > 140 is high; an electronegativity of less than 1.8 is taken to be low; 1.8–2.2 is moderate; and > than 2.2 is high (> 4.0 is very high).
Revised Pauling values are used for the metalloids, and reactive nonmetals; Allred-Rochow values for the noble gases
The nonmetallic halogens (F, Cl, Br, I) readily form anions including in aqueous solution; the oxide ion O²⁻ is unstable in aqueous solution—its affinity for H⁺ is so great that it abstracts a proton from a solvent H₂O molecule (O²⁻ + H₂O → 2 OH⁻)—but is found in an extensive series of metal oxides
The common oxide is the most stable oxide for that element

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References

Data sources

Unless otherwise stated, melting points, boiling points, densities, crystalline structures, ionisation energies, electron affinities, and electronegativity values are from the CRC Handbook of Physics and Chemistry;[78] standard electrode potentials are from the 1989 compilation by Steven Bratsch.[79]

Citations

Sukys 1999, p. 60.
Bettelheim et al. 2016, p. 33.
Schulze-Makuch & Irwin 2008, p. 89.
Steurer 2007, p. 7.
Cox 2004, p. 26
Meyer et al. 2005, p. 284; Manahan 2001, p. 911; Szpunar et al. 2004, p. 17
Brown & Rogers 1987, p. 40
Kneen, Rogers & Simpson 1972, p. 262
Greenwood & Earnshaw 2002, p. 434
Bratsch 1989; Bard, Parsons & Jordan 1985, p. 133
Yoder, Suydam & Snavely 1975, p. 58
Kneen, Rogers & Simpson 1972, p. 360
Lee 1996, p. 240
Greenwood & Earnshaw 2002, p. 43
Cressey 2010
Siekierski & Burgess 2002, p. 24–25
Siekierski & Burgess 2002, p. 23
Cox 2004, p. 146
Kneen, Rogers & Simpson 1972, p. 362
Bailar et al. 1989, p. 742
Stein 1983, p. 165
Jolly 1966, p. 20
Clugston & Flemming 2000, pp. 100–1, 104–5, 302
Seaborg 1969, p. 626
Nash 2005
Scerri 2013, pp. 204–8
Challoner 2014, p. 5; Government of Canada 2015; Gargaud et al. 2006, p. 447
Ivanenko et al. 2011, p. 784
Catling 2013, p. 12
Crawford 1968, p. 540
Berkowitz 2012, p. 293
Jørgensen & Mitsch 1983, p. 59
Wulfsberg 1987, p. 159–160
Bettelheim et al. 2016, p. 33—34
Field & Gray 2011, p. 12
Dinwiddle et al. 2018, pp. 34–35
Vernon 2020
Dingle 2017, pp. 9, 101, 179
Myers, Oldham & Tocci 2004, pp. 120–121
Stein 1969; Pitzer 1975; Schrobilgen 2011
Brasted 1974, p. 814
Sidorov 1960
Rochow 1966, p. 4
Atkins 2006 et al., pp. 8, 122–23
Ritter 2011, p. 10
Wiberg 2001, p. 680
Wiberg 2001, p. 403
Greenwood & Earnshaw 2002, p. 612
Moeller 1952, p. 208
Cotton 2003, p. 205
Wulfsberg 1987, p. 159
Cronyn 2003
Stoyanov et al.
Rayner-Canham 2011, p. 126
Dillon, Mathey & Nixon 1998
Martin-Louis et al. 2020
Miller et al.
Wiberg 2001, p. 686
Roy et al. 1994
Kent 2007, p. 104
Monteil & Vincent 1976
Moss 1952
Cacace, de Petris & Troiani 2002
Koziel 2002, p. 18
Piro et al. 2006
Steudel & Eckert 2003, p. 1
Greenwood & Earnshaw 2002, pp. 659–660
Moss 1952, p. 192; Greenwood & Earnshaw 2002, p. 751
Shanabrook, Lannin & Hisatsune 1981
Yousuf 1998, p. 425
Ostriker & Steinhardt 2001
Nelson 1987, p. 732
Emsley 2001, p. 428
Bolin 2012, p. 2-1
Maroni 1995
King & Caldwell 1954, p. 17; Brady & Senese 2009, p. 69
Nelson 1987, p. 735
Lide 2003
Bratsch 1989

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Monographs

Emsley J 1971, The inorganic chemistry of the non-metals, Methuen Educational, London, ISBN 0423861204
Johnson RC 1966, Introductory descriptive chemistry: selected nonmetals, their properties, and behavior, WA Benjamin, New York
Jolly WL 1966, The chemistry of the non-metals, Prentice-Hall, Englewood Cliffs, New Jersey
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Sherwin E & Weston GJ 1966, Chemistry of the non-metallic elements, Pergamon Press, Oxford
Steudel R 1977, Chemistry of the non-metals: with an introduction to atomic structure and chemical bonding, English edition by FC Nachod & JJ Zuckerman, Berlin, Walter de Gruyter, ISBN 3110048825

External links

Media related to Nonmetals at Wikimedia Commons

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[11] An ionisation energy of less than 750 kJ/mol is taken to be low, 750–1000 is moderate, and > 1000 is high (> 2000 is very high); an electron affinity of less than 70 kJ/mol is taken to be low, 70–140 is moderate, and > 140 is high; an electronegativity of less than 1.8 is taken to be low; 1.8–2.2 is moderate; and > than 2.2 is high (> 4.0 is very high).

[23] Revised Pauling values are used for the metalloids, and reactive nonmetals; Allred-Rochow values for the noble gases

[24] The nonmetallic halogens (F, Cl, Br, I) readily form anions including in aqueous solution; the oxide ion O²⁻ is unstable in aqueous solution—its affinity for H⁺ is so great that it abstracts a proton from a solvent H₂O molecule (O²⁻ + H₂O → 2 OH⁻)—but is found in an extensive series of metal oxides

[49] The common oxide is the most stable oxide for that element

[1] Sukys 1999, p. 60.

[2] Bettelheim et al. 2016, p. 33.

[3] Schulze-Makuch & Irwin 2008, p. 89.

[4] Steurer 2007, p. 7.

[Cox-5] Cox 2004, p. 26

[6] Meyer et al. 2005, p. 284; Manahan 2001, p. 911; Szpunar et al. 2004, p. 17

[7] Brown & Rogers 1987, p. 40

[8] Kneen, Rogers & Simpson 1972, p. 262

[9] Greenwood & Earnshaw 2002, p. 434

[10] Bratsch 1989; Bard, Parsons & Jordan 1985, p. 133

[12] Yoder, Suydam & Snavely 1975, p. 58

[13] Kneen, Rogers & Simpson 1972, p. 360

[14] Lee 1996, p. 240

[15] Greenwood & Earnshaw 2002, p. 43

[16] Cressey 2010

[17] Siekierski & Burgess 2002, p. 24–25

[18] Siekierski & Burgess 2002, p. 23

[19] Cox 2004, p. 146

[20] Kneen, Rogers & Simpson 1972, p. 362

[21] Bailar et al. 1989, p. 742

[22] Stein 1983, p. 165

[25] Jolly 1966, p. 20

[26] Clugston & Flemming 2000, pp. 100–1, 104–5, 302

[27] Seaborg 1969, p. 626

[28] Nash 2005

[29] Scerri 2013, pp. 204–8

[30] Challoner 2014, p. 5; Government of Canada 2015; Gargaud et al. 2006, p. 447

[31] Ivanenko et al. 2011, p. 784

[32] Catling 2013, p. 12

[Crawford-33] Crawford 1968, p. 540

[34] Berkowitz 2012, p. 293

[35] Jørgensen & Mitsch 1983, p. 59

[36] Wulfsberg 1987, p. 159–160

[37] Bettelheim et al. 2016, p. 33—34

[38] Field & Gray 2011, p. 12

[39] Dinwiddle et al. 2018, pp. 34–35

[40] Vernon 2020

[41] Dingle 2017, pp. 9, 101, 179

[42] Myers, Oldham & Tocci 2004, pp. 120–121

[43] Stein 1969; Pitzer 1975; Schrobilgen 2011

[44] Brasted 1974, p. 814

[45] Sidorov 1960

[46] Rochow 1966, p. 4

[47] Atkins 2006 et al., pp. 8, 122–23

[48] Ritter 2011, p. 10

[50] Wiberg 2001, p. 680

[ReferenceA-51] Wiberg 2001, p. 403

[52] Greenwood & Earnshaw 2002, p. 612

[53] Moeller 1952, p. 208

[Cotton-54] Cotton 2003, p. 205

[55] Wulfsberg 1987, p. 159

[56] Cronyn 2003

[57] Stoyanov et al.

[Rayner-Canham-58] Rayner-Canham 2011, p. 126

[59] Dillon, Mathey & Nixon 1998

[60] Martin-Louis et al. 2020

[61] Miller et al.

[62] Wiberg 2001, p. 686

[63] Roy et al. 1994

[64] Kent 2007, p. 104

[65] Monteil & Vincent 1976

[66] Moss 1952

[67] Cacace, de Petris & Troiani 2002

[68] Koziel 2002, p. 18

[69] Piro et al. 2006

[70] Steudel & Eckert 2003, p. 1

[71] Greenwood & Earnshaw 2002, pp. 659–660

[72] Moss 1952, p. 192; Greenwood & Earnshaw 2002, p. 751

[73] Shanabrook, Lannin & Hisatsune 1981

[74] Yousuf 1998, p. 425

[75] Ostriker & Steinhardt 2001

[76] Nelson 1987, p. 732

[77] Emsley 2001, p. 428

[78] Bolin 2012, p. 2-1

[79] Maroni 1995

[80] King & Caldwell 1954, p. 17; Brady & Senese 2009, p. 69

[81] Nelson 1987, p. 735

[82] Lide 2003

[83] Bratsch 1989

Electronegativity values of metalloids and nonmetals[n 2]
	1	2	13	14	15	16	17	18
								Noble gases
1	H 2.2			Reactive nonmetals				He (5.5)
2			B 2.04	C 2.55	N 3.04	O 3.44	F 3.98	Ne (4.84)
3				Si 1.9	P 2.19	S 2.58	Cl 3.16	Ar (3.2)
4				Ge 2.01	As 2.18	Se 2.55	Br 2.96	Kr (2.94)
5					Sb 2.05	Te 2.1	I 2.66	Xe (2.4)
6					Metalloids			Rn (2.06)

Electronegativity (EN) gives some indication of nonmetallic character. The metalloids have uniformly moderate values (1.8–2.2). Among the reactive nonmetals, hydrogen (2.2) and phosphorus (2.19) have moderate values but they each have higher ionisation energies than the metalloids, and are very rarely classed as such. Oxygen and the nonmetallic halogens have uniformly high EN values; nitrogen has a high EN but a marginally negative electron affinity that makes it a reluctant anion former.[n 3] The noble gases have some of the highest ENs but their complete valence shells and sizeably negative electron affinities render them chemically inert to a large degree.

Nonmetal categorisation and alternative schemes
		Reactive nonmetal			Noble gas
		H, C, N, P, O, S, Se, F, Cl, Br, I			He, Ne, Ar, Kr, Xe, Rn
(1)		Other nonmetal		Halogen	Noble gas
(1)		H, C, N, P, O, S, (Se)		F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn
(2)		Solid	Liquid	Gaseous
(2)		C, P, S, Se, I, At	Br	H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, Rn
(3)		Electronegative nonmetal	Very electronegative nonmetal		Noble gas
(3)		H, C, P, S, Se, I	N, O, F, Cl, Br		He, Ne, Ar, Kr, Xe, Rn
(4)		Polyatomic element	Diatomic element		Monatomic element (noble gas)
(4)		C, P, S, Se	H, N, O, F, Cl, Br, I		He, Ne, Ar, Kr, Xe, Rn
(5)		Hydrogen	Nonmetal	Halogen	Noble gas
(5)		H	C, N, P, O, S, Se	F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn
(6)		Hydrogen	Carbon and other nonmetals	Halogen	Noble gas
(6)		H	C, N, P, O, S, Se	F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn
(7)		Metalloid	Intermediate nonmetal	Corrosive nonmetal	Noble gas
(7)		B, Si, Ge, As, Sb, Te	H, C, N, P, S, Se	O, F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn
(8)	Hydrogen	Metalloid	Nonmetal	Halogen	Noble gas
(8)	H	B, Si, Ge, As, Sb, Te, Po	C, N, P, O, S, Se	F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn
(9)	Hydrogen	Semiconductor	Other nonmetal	Halogen	Noble gas
(9)	H	B, Si, Ge, As, Sb, Te	C, N, P, O, S, Se	F, Cl, Br, I, At	He, Ne, Ar, Kr, Xe, Rn

	Noble gases He, Ne, Ar, Kr, Xe, Rn
Active metals Groups 1–3, Ln, An, (Al)		Corrosive nonmetals O, F, Cl, Br, I
Transition metals Most of them		Intermediate nonmetals H, C, N, P, S, Se
Frontier metals (Al) Ag, Sn, Bi etc.		Metalloids B, Si, Ge, As, Sb, Te
	Noble metals Ru, Rh, Pd, Os, Ir, Pt, Au