Nonmetal

The seven metalloids are boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), and astatine (At). On a standard periodic table, they occupy a diagonal area in the p-block extending from boron at the upper left to astatine at lower right, along the dividing line between metals and nonmetals shown on some periodic tables. They are called metalloids mainly in light of their physical resemblance to metals.

While they each have a metallic appearance, they are brittle and only fair conductors of electricity. Boron, silicon, germanium, tellurium are semiconductors. Arsenic and antimony have the electronic band structures of semimetals although both have less stable semiconducting allotropes. Astatine has been predicted to have a metallic crystalline structure.

Chemically the metalloids generally behave like (weak) nonmetals. They have moderate ionisation energies, low to high electron affinities, moderate electronegativity values, are poor to moderately strong oxidising agents, and demonstrate a tendency to form alloys with metals.

The reactive nonmetals have a diverse range of individual physical and chemical properties. In periodic table terms they largely occupy a position between the weakly nonmetallic metalloids to the left and the noble gases to the right.

Physically, five are solids, one is a liquid (bromine), and five are gases. Of the solids, graphite carbon, selenium, and iodine are metallic-looking, whereas S₈ sulfur has a pale-yellow appearance. Ordinary white phosphorus has a yellowish-white appearance but the black allotrope, which is the most stable form of phosphorus, has a metallic-looking appearance. Bromine is a reddish-brown liquid. Of the gases, fluorine and chlorine are coloured pale yellow, and yellowish green. Electrically, most are insulators whereas graphite is a semimetal and black phosphorus, selenium, and iodine are semiconductors.

Chemically, they tend to have moderate to high ionisation energies, electron affinities, and electronegativity values, and be relatively strong oxidising agents. Collectively, the highest values of these properties are found among oxygen and the nonmetallic halogens. Manifestations of this status include oxygen's major association with the ubiquitous processes of corrosion and combustion, and the intrinsically corrosive nature of the nonmetallic halogens. All five of these nonmetals exhibit a tendency to form predominately ionic compounds with metals whereas the remaining nonmetals tend to form predominately covalent compounds with metals.

Six nonmetals are categorised as noble gases: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn). In periodic table terms they occupy the outermost right column. They are called noble gases in light of their characteristically very low chemical reactivity.

They have very similar properties, all being colorless, odorless, and nonflammable. With their closed valence shells the noble gases have feeble interatomic forces of attraction resulting in very low melting and boiling points.[22] That is why they are all gases under standard conditions, even those with atomic masses larger than many normally solid elements.[23]

Chemically, the noble gases have relatively high ionization energies, negative electron affinities, and relatively high electronegativities. Compounds of the noble gases number less than half a thousand, with most of these occurring via oxygen or fluorine combining with either krypton, xenon or radon.

The status of the period 7 congener of the noble gases, oganesson (Og), is not known—it may or may not be a noble gas. It was originally predicted to be a noble gas[24] but may instead be a fairly reactive solid with an anomalously low first ionisation potential, and a positive electron affinity, due to relativistic effects.[25] On the other hand, if relativistic effects peak in period 7 at element 112, copernicium, oganesson may turn out to be a noble gas after all,[26] albeit more reactive than either xenon or radon. While oganesson could be expected to be the most metallic of the group 18 elements, credible predictions on its status as either a metal or a nonmetal (or a metalloid) appear to be absent.

The nonmetals are sometimes instead divided according to either (1) the relative homogeneity of the halogens; (2) physical form; (3) electronegativity; (4) molecular structure; (5) the peculiar nature of hydrogen, and the relative homogeneity of the halogens; (6) their analogous categories among the metals; or (7) the uniqueness of hydrogen and the treatment of the metalloids as nonmetallic analogues of the post-transition metals.

In scheme (1), the halogens are in a category of their own; astatine is classed as a nonmetal, rather than a metalloid; and the remaining nonmetals are referred to as other nonmetals.[27] If selenium is counted as a metalloid rather than another nonmetal, the resulting set of less active nonmetals (H, C, N, P, O, S) are sometimes instead referred to or categorised as organogens,[28] CHONPS elements[29] or biogens.[30] Collectively these six nonmetals comprise the bulk of life on Earth;[31] a rough estimate of the composition of the biosphere is C₁₄₅₀H₃₀₀₀O₁₄₅₀N₁₅P₁S₁.[32]

In scheme (2), the nonmetals can simply be divided based on their physical forms at room temperature and pressure. The fluid nonmetals (bromine and the gaseous nonmetals) have the highest ionisation energy and electronegativity values among the elements, with the exception of hydrogen which tends to be anomalous in whichever category it is placed in. The solid nonmetals are collectively the most metallic of the nonmetallic elements, apart from the metalloids.

In scheme (3), the nonmetals are divided based on a loose correlation between electronegativity and oxidizing power.[33] Very electronegative nonmetals have electronegativity values over 2.8; electronegative nonmetals have values of 1.9 to 2.8.

In scheme (4), the nonmetals are distinguished based on the molecular structures of their most thermodynamically stable forms in ambient conditions.[34] Polyatomic nonmetals form structures or molecules in which each atom has two or three nearest neighbours (C_x, P₄, S₈, Se_x); diatomic nonmetals form molecules in which each atom has one nearest neighbour (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂); and the monatomic noble gases exist as isolated atoms (He, Ne, Ar, Kr, Xe, Rn) with no fixed nearest neighbour. This gradual reduction in the number of nearest neighbours corresponds (approximately) to a reduction in metallic character. A similar progression is seem among the metals. Metallic bonding tends to involve close-packed centrosymmetric structures with a high number of nearest neighbours. Post-transition metals and metalloids, sandwiched between the true metals and the nonmetals, tend to have more complex structures with an intermediate number of nearest neighbours.

In scheme (5), hydrogen is placed by itself on account of it being "so different from all other elements".[35] The remaining nonmetals are divided into nonmetals, halogens, and noble gases, with the unnamed category being distinguished by including nonmetals with relatively strong interatomic bonding, and the metalloids being effectively treated as a third super-category alongside metals and nonmetals.

In scheme (6) the nonmetals are divided into four classes that complement a four-fold division of the metals, with the noble metals treated as a subset of the transition metals. The metalloids are treated as chemically weak nonmetals, in a manner analogous to their chemically weak frontier metal counterparts.[36]

In scheme (7), hydrogen is again placed by itself on account of its uniqueness. The remaining nonmetals are divided into metalloids, quintessential nonmetals, halogens, and noble gases. Since the metalloids abut the post-transition or "poor" metals, they might be renamed the "poor non-metals".[37]

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]

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.

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.

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[44]

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.

Liquid oxygen (boiling)

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".[45] On the other hand, oxygen can react with many inorganic and organic compounds either spontaneously or under the right conditions,[46] (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.[47]

Sulfur

Sulfur is a bright-yellow moderately reactive[48] 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[48] 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,[48] 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.

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[45] 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".[49] 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.

Astatine is expected to have properties intermediate between iodine, a nonmetal with incident metallic properties, and tennessine, which is predicted to be a metal. Astatine has not so far been synthesised in sufficient quantities to enable a determination of its bulk properties. A macro-sized sample of astatine would vaporise itself due to radioactive heating; it is not known if such a phenomenon could be prevented with sufficient cooling. Many of the properties of astatine have nevertheless been predicted. It is expected to have a metallic appearance, a density of 6.35±0.15 g/cm³, a melting point of 302 °C, a boiling point of 337 °C(?), and a face-centred cubic crystalline structure. It has a moderate ionisation energy (899.003 kJ/mol), and is expected to have a high electron affinity (222 kJ/mol), and moderate electronegativity (2.2). Astatine is a weak oxidizing agent (At + e → At⁻ = 0.3 V at pH 0).

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.

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

Hydrogen in group 1, and carbon in group 14, show some out-of-group similarities.[50] 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.

Just as the metalloids cluster along a diagonal path, similar diagonal relationships occur between carbon and phosphorus, and between nitrogen and sulfur.

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.[51] Further illustrating the theme is the "extraordinary" similarity between low coordinate phosphorus compounds and unsaturated carbon compounds, and related research into organophosphorus chemistry.[52] 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.[53]

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.[51]

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

"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."[54]

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.

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.

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.

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

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.

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.

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.

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.

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;[70] standard electrode potentials are from the 1989 compilation by Steven Bratsch.[71]