User:Sandbh/Nonmetal redraft

Source: Wikipedia, the free encyclopedia.
  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

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).]]

Part of a series on the
Periodic table
Periodic table forms
Periodic table history
Sets of elements
By periodic table structure
By metallic classification
  • Nonmetals
By other characteristics
  • transplutonium
    elements
  • trans-
    actinides
Elements
List of chemical elements
Properties of elements
Data pages for elements

In chemistry, a nonmetal (or non-metal) is a chemical element that mostly lacks metallic attributes. Physically, nonmetals tend to have relatively low melting and boiling points, and densities, are mostly brittle if solid, and are usually poor conductors of heat and electricity; chemically, they tend to have relatively high ionization energy, electron affinity, and electronegativity values, and gain or share electrons when they react with other elements or 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, sulphur, selenium, and iodine). Metalloids such as boron, silicon and germanium are sometimes counted as nonmetals.

The nonmetals are divided into two categories reflecting their relative propensity to form chemical compounds namely reactive nonmetals and noble gases. The reactive nonmetals vary in 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 per cent of the observable Universe,[1] and one—oxygen—makes up close to half of the Earth's crust, oceans and atmosphere.[2] Living organisms are composed almost entirely of nonmetals,[3] and nonmetals form many more compounds than metals.[4]

Definition and applicable elements

Nonmetals in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson

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 the remainder having more substantial (open-packed) forms, unlike metals, which are nearly all solid and close-packed. If solid, they have a

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.
The dotted horizontal and vertical guides are intended only to sketch the boundaries of a broad progression in nonmetallic character, from the metalloids towards the lower the left to oxygen and the nonmetallic halogens towards the upper right.
The R2 value on the chart (which excludes H and N from its dataset) is a measure of how close a formula fits a set of data points. Values fall between 0.0 and 1.0, with those near 1.0 indicating a 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 xenon to a limited extent) 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]

Categories

Immediately to the left of most nonmetals on the periodic table are metalloids such as boron, silicon, and germanium, which generally behave chemically like nonmetals,[20] and are included here for comparative purposes. In this sense they can be regarded as the most metallic of nonmetallic elements.

Based on shared attributes, the nonmetals can be divided into the two categories of reactive nonmetal, and noble gas. The metalloids and the two nonmetal categories then span a progression in chemical nature from weakly nonmetallic, to moderately nonmetallic, to strongly nonmetallic (oxygen and the four nonmetallic halogens), to almost inert. Analogous categories occur among the metals in the form of the weakly metallic (the

alkaline earth metals, and the lanthanides and actinides), and the relatively inert (the noble transition metals
).

As with categorisation schemes generally, there is some variation and overlapping of properties within and across each category. One or more of the metalloids are sometimes classified as nonmetals.[5] Among the reactive nonmetals, carbon, phosphorus, selenium, and iodine—which border the metalloids—show some metallic character, as does hydrogen. Among the noble gases, radon is the most metallic and begins to show some cationic behaviour, which is unusual for a nonmetal.[21]

Metalloid

Main article: Metalloid

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.

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 At
2.2 		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.

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.

Reactive nonmetal

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, four are solids, one is a liquid (bromine), and six are gases. Of the solids, carbon, selenium, and iodine are metallic-looking, whereas 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 reddish-brown in colour. Of the gases, fluorine and chlorine are coloured pale yellow, and yellowish green. Electrically, most are insulators whereas carbon 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.

Noble gas

Main article: Noble gas

Six nonmetals are categorised as noble gases:

chemical reactivity
.

They have very similar properties, all being colorless, odourless, and nonflammable. With their closed valence shells the noble gases have feeble

standard conditions, even those with atomic masses larger than many normally solid elements.[23]

Chemically, the noble gases have relatively high ionisation 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,

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.

Alternative categories

Categorisation scheme alternatives
Base
scheme
 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
H,C,N,P,O,S,(Se) F,Cl,Br,I,At He,Ne,Ar,Kr,Xe,Rn
2  Solid Liquid Gaseous
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
H,C,P,S,Se,I N,O,F,Cl,Br He,Ne,Ar,Kr,Xe,Rn
4  Polyatomic
nonmetal 		Diatomic
nonmetal 		Noble gas (monatomic
nonmetal)
C,P,S,Se H,N,O,F,Cl,Br,I He,Ne,Ar,Kr,Xe,Rn
5  Hydrogen Nonmetal Halogen Noble gas
H C,N,P,O,S,Se F,Cl,Br,I,At He,Ne,Ar,Kr,Xe,Rn

The nonmetals are sometimes instead simply divided according to either (1) the relative homogeneity of the halogens; (2) physical form; (3) electronegativity; (4) molecular structure; or (5) the peculiar nature of hydrogen, and the relative homogeneity of the halogens.

In the first scheme, 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.

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 the second scheme, 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 the third scheme, 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 the fourth scheme, the nonmetals are distinguished based on their molecular structures.[34] Polyatomic nonmetals have structures featuring three or two nearest neighbours, diatomic nonmetals have one nearest neighbour, and the monatomic noble gases have none. 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 the fifth scheme, 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.

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[36]
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
Oxidation states • negative oxidation states known for all, but for H this is an unstable state
• positive oxidation states known for all bar 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
amphoteric or weakly acidic[39][40]
 		• 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;[41] 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
Moh's hardness

Group 1

Greyish lustrous block with uneven cleaved surface.
Hydrogen in an electrical discharge tube

Hydrogen is a colourless, odourless, and comparatively unreactive diatomic gas with a density of 0.08988 × 10−3 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 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

Several dozen small angular stone like shapes, grey with scattered silver flecks and highlights.
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 MnB, if n > 2. The common oxide of boron (B2O3) is weakly acidic.

Group 14

Greyish lustrous block with uneven cleaved surface.
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 has 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. The common oxide of germanium (GeO2) is amphoteric.

Group 15

Greyish lustrous block with uneven cleaved surface.
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 −210.00 °C. The solid form (density 0.85 g/cm−3; 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 (N2O5) is a strongly acidic oxide. 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.

Greyish lustrous block with uneven cleaved surface.
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[42]

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.

Two dull silver clusters of crystalline shards.
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. The common oxide of arsenic (As2O3) is acidic but weakly amphoteric.

A glistening silver rock-like chunk, with a blue tint, and roughly parallel furrows.
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 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. The common oxide of antimony (Sb2O3) is amphoteric.

Group 16

Greyish lustrous block with uneven cleaved surface.
Liquid oxygen, boiling at room temperature

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".[43] On the other hand, oxygen can react with many inorganic and organic compounds either spontaneously or under the right conditions,[44] (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/cm−3) 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.[45]

Sulfur

Sulfur is a bright-yellow comparatively unreactive 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(?) 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), low electron affinity (95 kJ/mol), and high electronegativity (2.55). It is 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(?), 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

Greyish lustrous block with uneven cleaved surface.
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[46] 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".[47] 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 then 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), 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).

Group 18

Greyish lustrous block with uneven cleaved surface.
Liquified helium

Helium has a density of 0.1785 × 10−3 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.

Greyish lustrous block with uneven cleaved surface.
Neon in an electrical discharge tube

Neon has a density of 0.9002 × 10−3 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.

Greyish lustrous block with uneven cleaved surface.
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.

Greyish lustrous block with uneven cleaved surface.
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.

Greyish lustrous block with uneven cleaved surface.
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
).

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.

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 7p3/2 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.

Cross-cutting relationships

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.[48] 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.[49] Further illustrating the theme is the "extraordinary" similarity between low coordinate phosphorus compounds and unsaturated carbon compounds, and related research into organophosphorus chemistry.[50]

Nitrogen and sufur 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.[49]

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."[51]

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

Main article:
Allotropes of nonmetals
Some allotropes of carbon

Many nonmetals have less stable

ductile mixture of allotropes)[56] have nonmetallic properties. Selenium has several nonmetallic allotropes, all of which are much less electrically conducting than its standard state of grey "metallic" selenium.[57] Iodine is also known in a semiconducting amorphous form.[58] Under sufficiently high pressures, just over half of the nonmetals, starting with phosphorus at 1.7 GPa,[59]
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 per cent 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.[60]

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.[61]

Nonmetals, and metalloids, in their elemental forms are extracted from:

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[63]

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 (

plasma TV display cells), while Rn also sometimes occurs, but then as an unwanted, potentially hazardous indoor pollutant;[64] 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.[65] 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.[66]

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. It 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:

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

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,

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, 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

  1. ^ 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).
  2. ^ Revised Pauling values are used for the metalloids, and reactive nonmetals; Allred-Rochow values for the noble gases
  3. ^ 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
  4. ^ The common oxide is the most stable oxide for that element

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;[67] and standard electrode potentials are from the 1989 compilation by Steven Bratsch.[68]

Citations

  1. ^ Sukys 1999, p. 60
  2. ^ Bettelheim et al. 2016, p. 33
  3. ^ Schulze-Makuch & Irwin 2008, p. 89
  4. ^ Steurer 2007, p. 7
  5. ^ a b 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
  11. ^ Yoder, Suydam & Snavely 1975, p. 58
  12. ^ Kneen, Rogers & Simpson 1972, p. 360
  13. ^ Lee 1996, p. 240
  14. ^ Greenwood & Earnshaw 2002, p. 43
  15. ^ Cressey 2010
  16. ^ Siekierski & Burgess 2002, p. 24–25
  17. ^ Siekierski & Burgess 2002, p. 23
  18. ^ Cox 2004, p. 146
  19. ^ Kneen, Rogers & Simpson 1972, p. 362
  20. ^ Bailar et al. 1989, p. 742
  21. ^ Stein 1983, p. 165
  22. ^ Jolly 1966, p. 20
  23. ^ Clugston & Flemming 2000, pp. 100–1, 104–5, 302
  24. ^ Seaborg 1969, p. 626
  25. ^ Nash 2005
  26. ^ Scerri 2013, pp. 204–8
  27. ^ Challoner 2014, p. 5; Government of Canada 2015; Gargaud et al. 2006, p. 447
  28. ^ Ivanenko et al. 2011, p. 784
  29. ^ Catling 2013, p. 12
  30. ^ Crawford 1968, p. 540
  31. ^ Berkowitz 2012, p. 293
  32. ^ Jørgensen & Mitsch 1983, p. 59
  33. ^ Wulfsberg 1987, p. 159–160
  34. ^ Bettelheim et al. 2016, p. 33—34
  35. ^ Field & Gray 2011, p. 12; see also Myers, Oldham & Tocci 2004, pp. 120–121 who categorize nonmetals as hydrogen; semiconductors "(also known as metalloids)"; less active nonmetals (C, N, O, P, S, Se); halogens; or noble gases
  36. ^ Stein 1969; Pitzer 1975; Schrobilgen 2011
  37. ^ Brasted 1974, p. 814
  38. ^ Sidorov 1960
  39. ^ Cite error: The named reference Rochow 1966, p. 4 was invoked but never defined (see the help page).
  40. ^ Atkins 2006 et al., pp. 8, 122–23
  41. ^ Ritter 2011, p. 10
  42. ^ Wiberg 2001, p. 680
  43. ^ Wiberg 2001, p. 403
  44. ^ Greenwood & Earnshaw 2002, p. 612
  45. ^ Moeller 1952, p. 208
  46. ^ Wiberg 2001, p. 403
  47. ^ Wulfsberg 1987, p. 159
  48. ^ Cronyn 2003
  49. ^ a b Rayner-Canham 2011, p. 126
  50. ^ Dillon, Mathey & Nixon 1998
  51. ^ Kent 2007, p. 104
  52. ^ Cacace, de Petris & Troiani 2002
  53. ^ Koziel 2002, p. 18
  54. ^ Piro et al. 2006
  55. ^ Steudel & Eckert 2003, p. 1
  56. ^ Greenwood & Earnshaw 2002, pp. 659–660
  57. ^ Moss 1952, p. 192; Greenwood & Earnshaw 2002, p. 751
  58. ^ Shanabrook, Lannin & Hisatsune 1981
  59. ^ Yousuf 1998, p. 425
  60. ^ Ostriker & Steinhardt 2001
  61. ^ Nelson 1987, p. 732
  62. ^ Emsley 2001, p. 428
  63. ^ Bolin 2012, p. 2-1
  64. ^ Maroni 1995
  65. ^ King & Caldwell 1954, p. 17; Brady & Senese 2009, p. 69
  66. ^ Nelson 1987, p. 735
  67. ^ Lide 2003
  68. ^ Bratsch 1989

References

Monographs
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