Nonmetal

This article is about the chemical elements that are not metals. For other meanings, see Nonmetal (disambiguation).

In the context of the periodic table, a nonmetal is a chemical element that mostly lacks distinctive metallic properties. They range from colorless gases like hydrogen to shiny crystals like iodine. Physically, they are usually lighter (less dense) than elements that form metals and are often poor conductors of heat and electricity. Chemically, nonmetals have relatively high electronegativity or usually attract electrons in a chemical bond with another element, and their oxides tend to be acidic.

A periodic table extract highlighting nonmetals

17 always/usually considered nonmetals[1][2][3]

6 metalloids, sometimes considered nonmetals

3 status as nonmetal or metal unconfirmed[4]

Seventeen elements are widely recognized as nonmetals. Additionally, some or all of six borderline elements (metalloids) are sometimes counted as nonmetals.

The two lightest nonmetals, hydrogen and helium, together account for about 98% of the mass of the observable universe. Five nonmetallic elements—hydrogen, carbon, nitrogen, oxygen, and silicon—form the bulk of Earth’s atmosphere, biosphere, crust and oceans, although metallic elements are believed to be slightly more than half of the overall composition of the Earth.

Chemical compounds and alloys involving multiple elements including nonmetals are widespread. Industrial uses of nonmetals as the dominant component include in electronics, combustion, lubrication and machining.

Most nonmetallic elements were identified in the 18th and 19th centuries. While a distinction between metals and other minerals had existed since antiquity, a classification of chemical elements as metallic or nonmetallic emerged only in the late 18th century. Since then about twenty properties have been suggested as criteria for distinguishing nonmetals from metals. In contemporary research usage it is common to use a distinction between metal and not-a-metal based upon the electronic structure of the solids; the elements carbon, arsenic and antimony are then semimetals, a subclass of metals. The rest of the nonmetallic elements are insulators, some of which such as silicon and germanium can readily accommodate dopants that change the electrical conductivity leading to semiconducting behavior.

Definition and applicable elements

Unless otherwise noted, this article describes the stable form of an element at standard temperature and pressure (STP).^[a]

While arsenic (here sealed in a container to prevent tarnishing) has a shiny appearance and is a reasonable conductor of heat and electricity, it is soft and brittle and its chemistry is predominately nonmetallic.^[6]

Nonmetallic chemical elements are often broadly defined as those that mostly lack properties commonly associated with metals—namely shininess, pliability, good thermal and electrical conductivity (due to their band structure), and a general capacity to form basic oxides.^[7][8] There is no widely accepted precise definition in terms of these properties;^[9] any list of nonmetals is open to debate and revision.^[1]

Fourteen elements are almost always recognized as nonmetals:^[1][2]

• Hydrogen
• Nitrogen
• Oxygen
• Sulfur

• Fluorine
• Chlorine
• Bromine
• Iodine

• Helium
• Neon
• Argon
• Krypton
• Xenon
• Radon

Three more are commonly classed as nonmetals, but some sources list them as "metalloids",^[3] a term which refers to elements intermediate between metals and nonmetals:^[10]

• Carbon
• Phosphorus
• Selenium

One or more of the six elements most commonly recognized as metalloids are sometimes instead counted as nonmetals:

• Boron
• Silicon
• Germanium
• Arsenic
• Antimony
• Tellurium

About 15–20% of the 118 known elements^[11] are thus classified as nonmetals.^[b]

General properties

Physical

See also § Physical properties by element type

Variety in color and form
of some nonmetallic elements

Boron in its β-rhombohedral phase

Metallic appearance of carbon as graphite

Blue color of liquid oxygen

Pale yellow liquid fluorine in a cryogenic bath

Sulfur as yellow chunks

Liquid bromine at room temperature

Metallic appearance of iodine under white light

Liquefied xenon

Nonmetals vary greatly in appearance, being colorless, colored or shiny. For the colorless nonmetals (hydrogen, nitrogen, oxygen, and the noble gases), no absorption of light happens in the visible part of the spectrum, and all visible light is transmitted.^[14] The colored nonmetals (sulfur, fluorine, chlorine, bromine) absorb some colors (wavelengths) and transmit the complementary or opposite colors. For example, chlorine's "familiar yellow-green colour ... is due to a broad region of absorption in the violet and blue regions of the spectrum".^[15][c] The shininess of boron, graphite (carbon), silicon, black phosphorus, germanium, arsenic, selenium, antimony, tellurium, and iodine^[d] is a result of the electrons reflecting incoming visible light.^[18]

About half of nonmetallic elements are gases under standard temperature and pressure; most of the rest are solids. Bromine, the only liquid, is usually topped by a layer of its reddish-brown fumes. The gaseous and liquid nonmetals have very low densities, melting and boiling points, and are poor conductors of heat and electricity.^[19] The solid nonmetals have low densities and low mechanical strength (being either hard and brittle, or soft and crumbly),^[20] and a wide range of electrical conductivity.^[e]

This diversity stems from variability in crystallographic structures and bonding arrangements. Covalent nonmetals existing as discrete atoms like xenon, or as small molecules, such as oxygen, sulfur, and bromine, have low melting and boiling points; many are gases at room temperature, as they are held together by weak London dispersion forces acting between their atoms or molecules, although the molecules themselves have strong covalent bonds.^[24] In contrast, nonmetals that form extended structures, such as long chains of selenium atoms,^[25] sheets of carbon atoms in graphite,^[26] or three-dimensional lattices of silicon atoms^[27] have higher melting and boiling points, and are all solids. Nonmetals closer to the left or bottom of the periodic table (and so closer to the metals) often have metallic interactions between their molecules, chains, or layers; this occurs in boron,^[28] carbon,^[29] phosphorus,^[30] arsenic,^[31] selenium,^[32] antimony,^[33] tellurium^[34] and iodine.^[35]

Some general physical differences
between elemental metals and nonmetals^[19]

Aspect	Metals	Nonmetals
Appearance and form	Shiny if freshly prepared or fractured; few colored;[36] all but one solid[37]	Shiny, colored or transparent;[38] all but one solid or gaseous[37]
Density	Often higher	Often lower
Plasticity	Mostly malleable and ductile	Often brittle solids
Electrical conductivity[39]	Good	Poor to good
Electronic structure[40]	Metal or semimetalic	Semimetal, semiconductor, or insulator

Covalently bonded nonmetals often share only the electrons required to achieve a noble gas electron configuration.^[41] For example, nitrogen forms diatomic molecules featuring a triple bonds between each atom, both of which thereby attain the configuration of the noble gas neon. In contrast antimony has buckled layers in which each antimony atom is singly bonded with three other nearby atoms.^[42]

Good electrical conductivity occurs when there is metallic bonding,^[43] however the electrons in some nonmetals are not metallic.^[43] Good electrical and thermal conductivity associated with metallic electrons is seen in carbon (as graphite, along its planes), arsenic, and antimony.^[f] Good thermal conductivity occurs in boron, silicon, phosphorus, and germanium;^[21] such conductivity is transmitted though vibrations of the crystalline lattices (phonons of these elements.^[44] Moderate electrical conductivity is observed in the semiconductors^[45] boron, silicon, phosphorus, germanium, selenium, tellurium, and iodine.

Many of the nonmetallic elements are hard and brittle,^[20] where dislocations cannot readily move so they tend to undergo brittle fracture rather than deforming.^[46] Some do deform such as white phosphorus (soft as wax, pliable and can be cut with a knife at room temperature),^[47] plastic sulfur,^[48] and selenium which can be drawn into wires from its molten state.^[49] Graphite is a standard solid lubricant where dislocations move very easily in the basal planes.^[50]

Allotropes

Three allotropes of carbon

A transparent electrical insulator

A brownish semiconductor

A blackish semimetal

From left to right, diamond, buckminsterfullerene, and graphite

For a more comprehensive list, see Allotropy § Non-metals, and Single-layer materials.

Over half of the nonmetallic elements exhibit a range of less stable allotropic forms, each with distinct physical properties.^[51] For example, carbon, the most stable form of which is graphite, can manifest as diamond, buckminsterfullerene,^[52] amorphous^[53] and paracrystalline^[54] variations. Allotropes also occur for nitrogen, oxygen, phosphorus, sulfur, selenium and iodine.^[55]

Chemical

See also § Chemical properties by element type

Some general chemistry-based
differences between metals and nonmetals^[19]

Aspect		Metals	Nonmetals
Reactivity[56]		Wide range: very reactive to noble
Oxides	lower	Basic	Acidic; never basic[57]
	higher	Increasingly acidic
Compounds with metals[58]		Alloys	Covalent or Ionic
Ionization energy[59]		Low to high	Moderate to very high
Electronegativity[60]		Low to high	Moderate to very high

Nonmetals have relatively high values of electronegativity, and their oxides are usually acidic. Exceptions may occur if a nonmetal is not very electronegative, or if its oxidation state is low, or both. These non-acidic oxides of nonmetals may be amphoteric (like water, H₂O^[61]) or neutral (like nitrous oxide, N₂O^[62][g]), but never basic.

They tend to gain electrons during chemical reactions, in contrast to metallic elements which tend to donate electrons. This behavior is related to the stability of electron configurations in the noble gases, which have complete outer shells, empirically described by the duet and octet rules of thumb, more correctly explained in terms of valence bond theory.^[65]

The chemical differences between metals and nonmetals stem from variations in how strongly atoms attract and retain electrons. Across a period of the periodic table, the nuclear charge increases as more protons are added to the nucleus.^[66] However, because the number of inner electron shells remains constant, the effective nuclear charge experienced by the outermost electrons also increases, pulling them closer to the nucleus. This leads to a corresponding reduction in atomic radius,^[67] and a greater tendency of these elements to gain electrons during chemical reactions, forming negatively charged ions.^[68] Nonmetals, which occupy the right-hand side of the periodic table, exemplify this behavior.

Nonmetals typically exhibit higher ionization energies, electron affinities, and standard electrode potentials than metals. The higher these values are (including electronegativity) the more nonmetallic the element tends to be.^[69] For example, the chemically very active nonmetals fluorine, chlorine, bromine, and iodine have an average electronegativity of 3.19—a figure^[h] higher than that of any metallic element.

The number of compounds formed by nonmetals is vast.^[70] The first 10 places in a "top 20" table of elements most frequently encountered in 895,501,834 compounds, as listed in the Chemical Abstracts Service register for November 2, 2021, were occupied by nonmetals. Hydrogen, carbon, oxygen, and nitrogen collectively appeared in most (80%) of compounds. Silicon, a metalloid, ranked 11th. The highest-rated metal, with an occurrence frequency of 0.14%, was iron, in 12th place.^[71]

Complications

Adding complexity to the chemistry of the nonmetals are anomalies occurring in the first row of each periodic table block; non-uniform periodic trends; higher oxidation states; multiple bond formation; and property overlaps with metals.

First row anomaly

Condensed periodic table highlighting the first row of each block:  s   p   d  and  f

Period

s-block

1

H 1

He 2

p-block

2

Li 3

Be 4

B 5

C 6

N 7

O 8

F 9

Ne 10

3

Na 11

Mg 12

d-block

Al 13

Si 14

P 15

S 16

Cl 17

Ar 18

4

K 19

Ca 20

Sc-Zn 21-30

Ga 31

Ge 32

As 33

Se 34

Br 35

Kr 36

5

Rb 37

Sr 38

f-block

Y-Cd 39-48

In 49

Sn 50

Sb 51

Te 52

I 53

Xe 54

6

Cs 55

Ba 56

La-Yb 57-70

Lu-Hg 71-80

Tl 81

Pb 82

Bi 83

Po 84

At 85

Rn 86

7

Fr 87

Ra 88

Ac-No 89-102

Lr-Cn 103-112

Nh 113

Fl 114

Mc 115

Lv 116

Ts 117

Og 118

Group

(1)

(2)

(3-12)

(13)

(14)

(15)

(16)

(17)

(18)

The first-row anomaly strength by block is s >> p > d > f.[72][i]

Starting with hydrogen, the first row anomaly primarily arises from the electron configurations of the elements concerned. Hydrogen is notable for its diverse bonding behaviors. It most commonly forms covalent bonds, but it can also lose its single electron in an aqueous solution, leaving behind a bare proton with high polarizing power.^[73] Consequently, this proton can attach itself to the lone electron pair of an oxygen atom in a water molecule, laying the foundation for acid-base chemistry.^[74] Moreover, a hydrogen atom in a molecule can form a second, albeit 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."^[75]

Hydrogen and helium, as well as boron through neon, have small atomic radii. The ionization energies and electronegativities among these elements are higher than the periodic trends would otherwise suggest.

While it would normally be expected, on electron configuration consistency grounds, that hydrogen and helium would be placed atop the s-block elements, the significant first row anomaly shown by these two elements justifies alternative placements. Hydrogen is occasionally positioned above fluorine, in group 17, rather than above lithium in group 1. Helium is almost always placed above neon, in group 18, rather than above beryllium in group 2.^[76]

Secondary periodicity

Electronegativity values of the group 16 chalcogen elements showing a W-shaped alternation or secondary periodicity going down the group

An alternation in certain periodic trends, sometimes referred to as secondary periodicity, becomes evident when descending groups 13 to 15, and to a lesser extent, groups 16 and 17.^[77][j] Immediately after the first row of d-block metals, from scandium to zinc, the 3d electrons in the p-block elements—specifically, gallium (a metal), germanium, arsenic, selenium, and bromine—prove less effective at shielding the increasing positive nuclear charge.

The Soviet chemist Shchukarev [ru] gives two more tangible examples:^[79]

"The toxicity of some arsenic compounds, and the absence of this property in analogous compounds of phosphorus [P] and antimony [Sb]; and the ability of selenic acid [H₂SeO₄] to bring metallic gold [Au] into solution, and the absence of this property in sulfuric [H₂SO₄] and [H₂TeO₄] acids."

Higher oxidation states

Roman numerals such as III, V and VIII denote oxidation states

Some nonmetallic elements exhibit oxidation states that deviate from those predicted by the octet rule, which typically results in an oxidation state of –3 in group 15, –2 in group 16, –1 in group 17, and 0 in group 18. Examples include ammonia NH₃, hydrogen sulfide H₂S, hydrogen fluoride HF, and elemental xenon Xe. Meanwhile, the maximum possible oxidation state increases from +5 in group 15, to +8 in group 18. The +5 oxidation state is observable from period 2 onward, in compounds such as nitric acid HN(V)O₃ and phosphorus pentafluoride PCl₅.^[k] Higher oxidation states in later groups emerge from period 3 onwards, as seen in sulfur hexafluoride SF₆, iodine heptafluoride IF₇, and xenon(VIII) tetroxide XeO₄. For heavier nonmetals, their larger atomic radii and lower electronegativity values enable the formation of compounds with higher oxidation numbers, supporting higher bulk coordination numbers.^[80]

Multiple bond formation

Molecular structure of pentazenium, a homopolyatomic cation of nitrogen with the formula N₅⁺ and structure N−N−N−N−N.^[81]

Period 2 nonmetals, particularly carbon, nitrogen, and oxygen, show a propensity to form multiple bonds. The compounds formed by these elements often exhibit unique stoichiometries and structures, as seen in the various nitrogen oxides,^[80] which are not commonly found in elements from later periods.

Property overlaps

While certain elements have traditionally been classified as nonmetals and others as metals, some overlapping of properties occurs. Writing early in the twentieth century, by which time the era of modern chemistry had been well-established^[82] (although not as yet more precise quantum chemistry) Humphrey^[83] observed that:

... these two groups, however, are not marked off perfectly sharply from each other; some nonmetals resemble metals in certain of their properties, and some metals approximate in some ways to the non-metals.

Boron (here in its less stable amorphous form) shares some similarities with metals^[l]

Examples of metal-like properties occurring in nonmetallic elements include:

• Silicon has an electronegativity (1.9) comparable with metals such as cobalt (1.88), copper (1.9), nickel (1.91) and silver (1.93);^[60]
• The electrical conductivity of graphite exceeds that of some metals;^[m]
• Selenium can be drawn into a wire;^[49]
• Radon is the most metallic of the noble gases and begins to show some cationic behavior, which is unusual for a nonmetal;^[87] and
• In extreme conditions, just over half of nonmetallic elements can form homopolyatomic cations.^[n]

Examples of nonmetal-like properties occurring in metals are:

• Tungsten displays some nonmetallic properties, sometimes being brittle, having a high electronegativity, and forming only anions in aqueous solution,^[89] and predominately acidic oxides.^[8][90]
• Gold, the "king of metals" has the highest electrode potential among metals, suggesting a preference for gaining rather than losing electrons. Gold's ionization energy is one of the highest among metals, and its electron affinity and electronegativity are high, with the latter exceeding that of some nonmetals. It forms the Au^– auride anion and exhibits a tendency to bond to itself, behaviors which are unexpected for metals. In aurides (MAu, where M = Li–Cs), gold's behavior is similar to that of a halogen.^[91] The reason for this is that gold has a large enough nuclear potential that the electrons have to be considered with relativistic effects included, which changes some of the properties.^[92]

A relatively recent development involves certain compounds of heavier p-block elements, such as silicon, phosphorus, germanium, arsenic and antimony, exhibiting behaviors typically associated with transition metal complexes. This is linked to a small energy gap between their filled and empty molecular orbitals, which are the regions in a molecule where electrons reside and where they can be available for chemical reactions. In such compounds, this allows for unusual reactivity with small molecules like hydrogen (H₂), ammonia (NH₃), and ethylene (C₂H₄), a characteristic previously observed primarily in transition metal compounds. These reactions may open new avenues in catalytic applications.^[93]

Types

Nonmetal classification schemes vary widely, with some accommodating as few as two subtypes and others up to seven. For example, the periodic table in the Encyclopaedia Britannica recognizes noble gases, halogens, and other nonmetals, and splits the elements commonly recognized as metalloids between "other metals" and "other nonmetals".^[94] On the other hand, seven of twelve color categories on the Royal Society of Chemistry periodic table include nonmetals.^[95][o]

		Group (1, 13−18)				Period
	13	14	15	16	1/17	18	(1−6)
					H	He	1
	B	C	N	O	F	Ne	2
		Si	P	S	Cl	Ar	3
		Ge	As	Se	Br	Kr	4
			Sb	Te	I	Xe	5
						Rn	6

Starting on the right side of the periodic table, three types of nonmetals can be recognized:

   the inert noble gases—helium, neon, argon, krypton, xenon, radon;^[96]

   the reactive halogen nonmetals—fluorine, chlorine, bromine, iodine;^[97] and

   the mixed reactivity "unclassified nonmetals", a set with no widely used collective name—hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, selenium.^[q] The descriptive phrase unclassified nonmetals is used here for convenience.

The elements in a fourth set are sometimes recognized as nonmetals:

   the generally unreactive^[s] metalloids,^[115] sometimes considered a third category distinct from metals and nonmetals—boron, silicon, germanium, arsenic, antimony, tellurium.

While many of the early workers attempted to classify elements none of their classifications were satisfactory. They were divided into metals and nonmetals, but some were soon found to have properties of both. These were called metalloids. This only added to the confusion by making two indistinct divisions where one existed before.^[116]

Whiteford & Coffin 1939, Essentials of College Chemistry

The boundaries between these types are not sharp.^[t] Carbon, phosphorus, selenium, and iodine border the metalloids and show some metallic character, as does hydrogen.

The greatest discrepancy between authors occurs in metalloid "frontier territory".^[118] Some consider metalloids distinct from both metals and nonmetals, while others classify them as nonmetals.^[119] Some categorize certain metalloids as metals (e.g., arsenic and antimony due to their similarities to heavy metals).^[120][u] Metalloids resemble the elements universally considered "nonmetals" in having relatively low densities, high electronegativity, and similar chemical behavior.^[115][v]

Noble gases

Main article: Noble gas

A small (about 2 cm long) piece of rapidly melting argon ice

Six nonmetals are classified as noble gases: helium, neon, argon, krypton, xenon, and the radioactive radon. In conventional periodic tables they occupy the rightmost column. They are called noble gases due to their exceptionally low chemical reactivity.^[96]

These elements exhibit similar properties, being colorlessness, odorless, and nonflammable. Due to their closed outer electron shells, noble gases possess weak interatomic forces of attraction, leading to exceptionally low melting and boiling points.^[121]

Chemically, the noble gases exhibit relatively high ionization energies, negligible or negative electron affinities, and high to very high electronegativities. The number of compounds formed by noble gases is in the hundreds and continues to expand,^[122] with most of these compounds involving the combination of oxygen or fluorine with either krypton, xenon, or radon.^[123]

Halogen nonmetals

Main article: Halogen

Chemically, the halogen nonmetals have high ionization energies, electron affinities, and electronegativity values, and are relatively strong oxidizing agents.^[124] All four elements tend to form primarily ionic compounds with metals,^[125] in contrast to the remaining nonmetals (except for oxygen) which tend to form primarily covalent compounds with metals.^[w]

Unclassified nonmetals

Selenium conducts electricity around 1,000 times better when light falls on it, a property used in light-sensing applications.^[129]

Hydrogen behaves in some respects like a metallic element and in others like a nonmetal.^[130] Like a metallic element it can, for example, form a solvated cation in aqueous solution;^[131] it can substitute for alkali metals in compounds such as the chlorides (NaCl cf. HCl) and nitrates (KNO₃ cf. HNO₃), and in certain alkali metal complexes^[132][133] as a nonmetal.^[134] It attains this configuration by forming a covalent or ionic bond^[135] or by bonding as an ion to a lone pair of electrons.^[136]

Some or all of these nonmetals share several properties. Being generally less reactive than the halogens,^[137] most of them can occur naturally in the environment.^[138] Collectively, their physical and chemical characteristics can be described as "moderately non-metallic".^[139] When combined with metals, the unclassified nonmetals can form interstitial or refractory compounds.^[140] They also exhibit a tendency to bond to themselves, particularly in solid compounds.^[141] Additionally, diagonal periodic table relationships among these nonmetals mirror similar relationships among the metalloids.^[142]

Abundance, extraction, and uses

Abundance

Approximate composition
(top three components by weight; elements marked *nonmetal  metal 

Universe[143]	75% *hydrogen	23% *helium	1% *oxygen
Atmosphere[144]	78% *nitrogen	21% *oxygen	0.5% *argon
Hydrosphere[145]	86% *oxygen	11% *hydrogen	2% *chlorine
Biomass[146]	63% *oxygen	20% *carbon	10% *hydrogen
Crust[145]	46% *oxygen	27% *silicon	8% aluminium
Earth[147]	32% iron	30% *oxygen	14% magnesium

The abundance of elements in the universe results from nuclear physics processes like nucleosynthesis and radioactive decay.

The volatile noble gas nonmetal elements are less abundant in the atmosphere than expected based upon their overall abundance due to cosmic nucleosynthesis. Mechanisms to explain this difference is an important aspect of planetary science.^[148] The element Xe is unexpectedly depleted, and a possible explanation comes from theoretical models of the high pressures in the Earth's core suggesting that there may be around 10¹³ tons of xenon in the form of stable XeFe₃ and XeNi₃ intermetallic compounds.^[149]

Five nonmetals—hydrogen, carbon, nitrogen, oxygen, and silicon—form the bulk of the directly observable structure of the Earth: about 73% of the crust, 93% of the biomass, 96% of the hydrosphere, and over 99% of the atmosphere, as shown in the accompanying table. Silicon and oxygen form stable tetrahedral structures, known as silicates. Here, "the powerful bond that unites the oxygen and silicon ions is the cement that holds the Earth's crust together."^[150] However, they make up less than 50% of the total weight of the earth.^[147]

In the biomass, the relative abundance of the first four nonmetals (and phosphorus, sulfur, and selenium marginally) is attributed to a combination of relatively small atomic size, and sufficient spare electrons. These two properties enable them to bind to one another and "some other elements, to produce a molecular soup sufficient to build a self-replicating system."^[151]

Extraction

Nine of the 23 nonmetallic elements are gases, or form compounds that are gases, and are extracted from natural gas or liquid air, including hydrogen, nitrogen, oxygen, sulfur, and most of the noble gases. For example, nitrogen and oxygen are extracted from liquid air through fractional distillation^[152] and sulfur from the hydrogen sulfide in natural gas by reacting it with oxygen to yield water and sulfur.^[153] Three nonmetals are extracted from seawater; the rest of the nonmetals – and almost all metals – from mining solid ores.^{[citation needed]}

		Group (1, 13−18)				Period
	13	14	15	16	1/17	18	(1−6)
					H	He	1
	B	C	N	O	F	Ne	2
		Si	P	S	Cl	Ar	3
		Ge	As	Se	Br	Kr	4
			Sb	Te	I	Xe	5
						Rn	6

Nonmetallic elements are extracted from these sources:^[138]

3 from natural gas components

hydrogen (methane), helium, and sulfur (hydrogen sulfide)

6 from liquefied air

nitrogen, oxygen, neon, argon, krypton, and xenon

3 from seawater brine

chlorine, bromine, and iodine

12 from solid ores

boron (borates), carbon (natural graphite), silicon (silica), phosphorus (phosphates), iodine (sodium iodate), radon (uranium ore decay product), fluorine (fluorite); and germanium, arsenic, selenium, antimony, and tellurium (from their sulfides).

Uses

Cylinders containing argon gas for use in extinguishing fire without damaging computer server equipment

Nonmetallic elements are present in combination with other elements in almost everything around us, from water to plastics and within metallic alloys. There are some specific uses of the elements themselves, although these are less common; extensive details can be found in the specific pages of the relevant elements. A few examples are:

1. Hydrogen can be used in fuel cells, and is being explored for a possible future low-carbon hydrogen economy.^[154]
2. Carbon has many uses, ranging from decorative applications of diamond jewelry^[155] to diamond in cutting blades^[156] and graphite as a solid lubricant.^[50]
3. Liquid nitrogen is extensively used as a coolant.^[157]
4. Oxygen is a critical component of the air we breath. (While nitrogen is also present, it is less used from the air, mainly by certain bacteria.^[158]) Oxygen gas and liquid is also heavily used for combustion in welding and cutting torches and as a component of rocket fuels.^[159]
5. Silicon is the most widely used semiconductor. While ultra-pure silicon is an insulator, by selectively adding electronic dopants it can be used as a semiconductor where the chemical potential of the electrons can be manipulated, this being exploited in a wide range of electronic devices.^[160]
6. The noble gases have a range of applications, including liquid helium for cryogenic cooling,^[161] and argon to in gaseous fire suppression to -damp fires around sensitive electrical equipment where water cannot be used.^[162]
7. Radon is a potentially hazardous indoor pollutant.^[163]

Taxonomical history

Background

Medieval chemical philosophers focused on metals, rarely investigating nonmetallic minerals.^[164]

French nobleman and chemist Antoine Lavoisier (1743–1794), with a page of the English translation of his 1789 Traité élémentaire de chimie,^[165] listing the elemental gases oxygen, hydrogen and nitrogen (and erroneously including light and caloric); the nonmetallic substances sulfur, phosphorus, and carbon; and the chloride, fluoride and borate ions

Organization of elements by types

See also: Discovery of chemical elements

In the late 1700s, French chemist Antoine Lavoisier published the first modern list of chemical elements in his revolutionary^[166] 1789 Traité élémentaire de chimie. The 33 elements known to Lavoisier were categorized into four distinct groups, including gases, metallic substances, nonmetallic substances that form acids when oxidized,^[167] and earths (heat-resistant oxides).^[168] Lavoisier's work gained widespread recognition and was republished in twenty-three editions across six languages within its first seventeen years, significantly advancing the understanding of chemistry in Europe and America.^[169] Lavoisier's chemistry was "dualistic",: "salts" were combinations of "acid" and "base"; acids where combinations of oxygen and metals while bases where combinations of oxygen and nonmetals. This view prevailed despite increasing evidence that chemicals like chlorine and ammonia contained no oxygen, in large part due the vigious if sometimes misguided defense by the Swedish chemist Berzelius.^[167]: 165

In 1802 the term "metalloids" was introduced for elements with the physical properties of metals but the chemical properties of non-metals.^[170] In 1811 Berzelius used the term "metalloids"^[171] to describe all nonmetallic elements, noting their ability to form negatively charged ions with oxygen in aqueous solutions.^[172][173] Drawing on this, in 1864 the "Manual of Metalloids" divided all elements into either metals or metalloids, with the latter group including elements now called nonmetals.^[174]: 31 Reviews of the book indicated that the term "metalloids" was still endorsed by leading authorities,^[175] but there were reservations about its appropriateness. While Berzelius' terminology gained significant acceptance,^[176] it later faced criticism from some who found it counterintuitive,^[173] misapplied,^[177] or even invalid.^[178] The idea of designating elements like arsenic as metalloids had been considered.^[175] The use of the term "metalloids" persisted in France as textbooks of chemistry appeared in the 1800s. During this period, "metals" as a chemical category were characterized by a single property, their affinity for oxygen, while "metalloids" were organized by comparison of many characteristic analogous to the approach of naturalists.^[179]

Development of types

Bust of Dupasquier (1793–1848) in the Monument aux Grands Hommes de la Martinière [fr] in Lyon, France.

In 1844, Alphonse Dupasquier [fr], a French doctor, pharmacist, and chemist,^[180] established a basic taxonomy of nonmetals to aid in their study. He wrote:^[181]

They will be divided into four groups or sections, as in the following:

Organogens—oxygen, nitrogen, hydrogen, carbon

Sulphuroids—sulfur, selenium, phosphorus

Chloroides—fluorine, chlorine, bromine, iodine

Boroids—boron, silicon.

Dupasquier's quartet parallels the modern nonmetal types. The organogens and sulphuroids are akin to the unclassified nonmetals. The chloroides were later called halogens.^[182] The boroids eventually evolved into the metalloids, with this classification beginning from as early as 1864.^[175] The then unknown noble gases were recognized as a distinct nonmetal group after being discovered in the late 1800s.^[183] This taxonomy was noted as a "natural classification" of the substance considering all aspects rather than an single characteristic like oxygen affinity.^[184] It was a significant departure from other contemporary classifications, since it grouped together oxygen, nitrogen, hydrogen, and carbon.^[185]

In 1828 and 1859, the French chemist Dumas classified nonmetals as (1) hydrogen; (2) fluorine to iodine; (3) oxygen to sulfur; (4) nitrogen to arsenic; and (5) carbon, boron and silicon,^[186] thereby anticipating the vertical groupings of Mendeleev's 1871 periodic table. Dumas' five classes fall into modern groups 1, 17, 16, 15, and 14 to 13 respectively.

Nonmetals as terminology

By as early as 1866, some authors began preferring the term "nonmetal" over "metalloid" to describe nonmetallic elements.^[187] In 1875, Kemshead^[188] observed that elements were categorized into two groups: non-metals (or metalloids) and metals. He noted that the term "non-metal", despite its compound nature, was more precise and had become universally accepted as the nomenclature of choice.

Structure, quantum mechanics and band structure

The early terminologies were empirical categorizations based upon observables. As the 20th century started there were significant changes in understanding. The first was due to methods, mainly x-ray crystallography, for determining how atoms are arranged in the various materials. As early as 1784 René Just Haüy discovered that every face of a crystal could be described by simple stacking patterns of blocks of the same shape and size (law of decrements).^[189] Haüy's study led to the idea that crystals are a regular three-dimensional array (a Bravais lattice) of atoms and molecules, with a single unit cell repeated indefinitely, along with other developments in the early days of physical crystallography. After Max von Laue demonstrated in 1912 that x-rays diffract,^[190] fairly quickly William Lawrence Bragg and his father William Henry Bragg were able to solve previously unknown structures.^{[191][192][193]} Building on this, it became clear that most of the simple elemental metals had close packed structures. With this determined the concept of dislocations originally developed by Vito Volterra in 1907^[194] became accepted, for instance being used to explain the ductility of metals by G. I. Taylor in 1934.^[195]

The second was the advent of quantum mechanics. In 1924 Louis de Broglie in his PhD thesis Recherches sur la théorie des quanta^[196] introduced his theory of electron waves. This rapidly became part of what was called by Erwin Schrödinger undulatory mechanics,^[197] now called the Schrödinger equation, wave mechanics or more commonly in contemporary usage quantum mechanics. While it was not so easy to solve the mathematics in the early days, fairly rapidly ideas such as the chemical bond terminology of Linus Pauling^[198] as well as electronic band structure concepts were developed.^[160]

Filling of the electronic states in various types of materials at equilibrium. Here, height is energy while width is the density of available states for a certain energy in the material listed. The shade follows the Fermi–Dirac distribution (black: all states filled, white: no state filled). In metals and semimetals the Fermi level E_F lies inside at least one band.

In insulators and semiconductors the Fermi level is inside a band gap; however, in semiconductors the bands are near enough to the Fermi level to be thermally populated with electrons or holes. "intrin." indicates intrinsic semiconductors.

edit

From this the concept of nonmetals as "not-a-metal" originates. The original approach to describe metals and nonmetals was a band-structure with delocalized electrons (i.e. spread out in space). A nonmetal has a gap in the energy levels of the electrons at the Fermi level.^{[160]: Chpt 8 & 19} In contrast, a metal would have at least one partially occupied band at the Fermi level;^[160] in a semiconductor or insulator there are no delocalized states at the Fermi level, see for instance Ashcroft and Mermin.^[160] (A semimetal is similar to a metal, with a slightly more complex band structure.) These definitions are equivalent to stating that metals conduct electricity at absolute zero, as suggested by Nevill Francis Mott,^[199]: 257 and the equivalent definition at other temperatures is also commonly used as in textbooks such as Chemistry of the Non-Metals by Ralf Steudel^[200] and work on metal–insulator transitions.^[201][202]

Originally^[203][204] this band structure interpretation was based upon a single-electron approach with the Fermi level in the band gap as illustrated in the Figure, not including a complete picture of the many-body problem where both exchange and correlation terms matter, as well as relativistic effects such as spin-orbit coupling. For instance, the passivity of gold is typically associated with relativistic terms.^[205] A key addition by Mott and Rudolf Peierls was that these could not be ignored.^[206] For instance, nickel oxide would be a metal if a single-electron approach was used, but in fact has quite a large band gap.^[207] As of 2024 it is more common to use an approach based upon density functional theory where the many-body terms are included.^[208][209] Although accurate calculations remain a challenge, reasonable results are now available in many cases.^[210][211]

It is common to nuance the early definition of Alan Herries Wilson and Mott which was for a zero temperature. As discussed by Peter Edwards and colleagues,^[212] as well as Fumiko Yonezawa,^{[199]: 257–261}it is important to consider the temperatures at which both metals and nonmetals are used. Yonezawa provides a general definition for both general temperatures and conditions (such as standard temperature and pressure):^[199]: 260

When a material conducts and at the same time the temperature coefficient of the electric conductivity of that material is not positive under a certain environmental condition, the material is metallic under that environmental condition. A material which does not satisfy these requirements is not metallic under that environmental condition.

The precise meaning of semiconductor needs a little care. In terms of the temperature dependence of their conductivity they are all classified as insulators; the pure forms are intrinsic semiconductors. When they are doped their conductivity continues to increase with temperature,^[160] and can become substantial; hence the ambiguities with an empirical categorisation using conductivity described earlier. Indeed, some elements that are normally considered as insulators have been exploited as semiconductors. For instance diamond, which has the largest band gap of the elements that are solids under normal conditions,^[213] has a number of semiconductor applications.^[214][215]

Band structure definitions of metals and nonmetals are widely used in current research into materials, and apply both to single elements such as insulating boron^[216] as well as compounds such as strontium titanate.^[217] The characteristics associated with metals and nonmetals in early work such as their appearance and mechanical properties are now understood to be consequences of how the atoms and electrons are arranged.

Comparison of selected properties

The two tables in this section list some of the properties of five types of elements (noble gases, halogen nonmetals, unclassified nonmetals, metalloids and, for comparison, metals) based on their most stable forms at standard temperature and pressure. The dashed lines around the columns for metalloids signify that the treatment of these elements as a distinct type can vary depending on the author, or classification scheme in use.

Physical properties by element type

See also § Physical

Physical properties are listed in loose order of ease of their determination.

Property	Element type
	Metals	Metalloids	Unc. nonmetals	Halogen nonmetals	Noble gases
General physical appearance	lustrous[19]	lustrous[218]	◇ lustrous: carbon, phosphorus, selenium[219] ◇ colored: sulfur[220] ◇ colorless: hydrogen, nitrogen, oxygen[221]	◇ lustrous: iodine[3] ◇ colored: fluorine, chlorine, bromine[222]	colorless[223]
Form and density[224]	solid (Hg liquid)	solid	solid or gas	solid or gas (bromine liquid)	gas
	often high density such as iron, lead, tungsten	low to moderately high density	low density	low density	low density
	some light metals including beryllium, magnesium, aluminium	all lighter than iron	hydrogen, nitrogen lighter than air[225]		helium, neon lighter than air[226]
Plasticity	mostly malleable and ductile[19]	often brittle[218]	phosphorus, sulfur, selenium, brittle[x]	iodine brittle[230]	not applicable
Electrical conductivity	good[y]	◇ moderate: boron, silicon, germanium, tellurium ◇ good: arsenic, antimony[z]	◇ poor: hydrogen, nitrogen, oxygen, sulfur ◇ moderate: phosphorus, selenium ◇ good: carbon[aa]	◇ poor: fluorine, chlorine, bromine ◇ moderate: I[ab]	poor[ac]
Electronic structure[40]	metal (beryllium, strontium, α-tin, ytterbium, bismuth are semimetals)	semimetal (arsenic, antimony) or semiconductor	◇ semimetal: carbon ◇ semiconductor: phosphorus ◇ insulator: hydrogen, nitrogen, oxygen, sulfur	semiconductor (I) or insulator	insulator

Chemical properties by element type

See also § Chemical

Chemical properties are listed from general characteristics to more specific details.

Property	Element type
	Metals	Metalloids	Unc. nonmetals	Halogen nonmetals	Noble gases
General chemical behavior	◇ strong to weakly metallic[234] ◇ noble metals are relatively inert[235]	weakly nonmetallic[ad]	moderately nonmetallic[236]	strongly nonmetallic[237]	◇ inert to nonmetallic[238] ◇ radon shows some cationic behavior[239]
Oxides	basic; some amphoteric or acidic[8]	amphoteric or weakly acidic[240][ae]	acidic[af] or neutral[ag]	acidic[ah]	metastable XeO3 is acidic;[247] stable XeO4 strongly so[248]
	few glass formers[ai]	all glass formers[250]	some glass formers[aj]	no glass formers reported	no glass formers reported
	ionic, polymeric, layer, chain, and molecular structures[252]	polymeric in structure[253]	◇ mostly molecular[253] ◇ carbon, phosphorus, sulfur, selenium have 1+ polymeric forms	◇ mostly molecular ◇ iodine has a polymeric form, I2O5[254]	◇ mostly molecular ◇ XeO2 is polymeric[255]
Compounds with metals	alloys[19] or intermetallic compounds[256]	tend to form alloys or intermetallic compounds[257]	◇ salt-like to covalent or metallic: hydrogen†, carbon, nitrogen, phosphorus, sulfur, selenium[10] ◇ mainly ionic: oxygen[258]	mainly ionic[125]	simple compounds at STP not known[ak]
Ionization energy (kJ mol−1)[59] ‡	low to high	moderate	moderate to high	high	high to very high
	376 to 1,007	762 to 947	941 to 1,402	1,008 to 1,681	1,037 to 2,372
	average 643	average 833	average 1,152	average 1,270	average 1,589
Electronegativity (Pauling)[al][60] ‡	low to high	moderate	moderate to high	high	high (radon) to very high
	0.7 to 2.54	1.9 to 2.18	2.19 to 3.44	2.66 to 3.98	ca. 2.43 to 4.7
	average 1.5	average 2.05	average 2.65	average 3.19	average 3.3

† Hydrogen can also form alloy-like hydrides^[133]
‡ The labels low, moderate, high, and very high are arbitrarily based on the value spans listed in the table

See also

• CHON (carbon, hydrogen, oxygen, nitrogen)
• List of nonmetal monographs
• Metallization pressure
• Nonmetal (astrophysics)
• Period 1 elements (hydrogen & helium)
• Properties of nonmetals (and metalloids) by group

Notes

1. The most stable forms are: diatomic hydrogen H₂; β-rhombohedral boron; graphite for carbon; diatomic nitrogen N₂; diatomic oxygen O₂; tetrahedral silicon; black phosphorus; orthorhombic sulfur S₈; α-germanium; gray arsenic; gray selenium; gray antimony; gray tellurium; and diatomic iodine I₂. All other nonmetallic elements have only one stable form at STP.^[5]
2. At higher temperatures and pressures the numbers of nonmetals can be called into question. For example, when germanium melts it changes from a semiconductor to a metallic conductor with an electrical conductivity similar to that of liquid mercury.^[12] At a high enough pressure, sodium (a metal) becomes a non-conducting insulator.^[13]
3. The absorbed light may be converted to heat or re-emitted in all directions so that the emission spectrum is thousands of times weaker than the incident light radiation.^[16]
4. Solid iodine has a silvery metallic appearance under white light at room temperature. At ordinary and higher temperatures it sublimes from the solid phase directly into a violet-colored vapor.^[17]
5. The solid nonmetals have electrical conductivity values ranging from 10⁻¹⁸ S•cm⁻¹ for sulfur^[21] to 3 × 10⁴ in graphite^[22] or 3.9 × 10⁴ for arsenic;^[23] cf. 0.69 × 10⁴ for manganese to 63 × 10⁴ for silver, both metals.^[21] The conductivity of graphite and arsenic (both semimetals) exceed that of manganese.
6. Thermal conductivity values for metals range from 6.3 W m⁻¹ K⁻¹ for neptunium to 429 for silver; cf. antimony 24.3, arsenic 50, and carbon 2000.^[21] Electrical conductivity values of metals range from 0.69 S•cm⁻¹ × 10⁴ for manganese to 63 × 10⁴ for silver; cf. carbon 3 × 10⁴,^[22] arsenic 3.9 × 10⁴ and antimony 2.3 × 10⁴.^[21]
7. While CO and NO are commonly referred to as being neutral, CO is a slightly acidic oxide, reacting with bases to produce formates (CO + OH⁻ → HCOO⁻);^[63] and in water, NO reacts with oxygen to form nitrous acid HNO₂ (4NO + O₂ + 2H₂O → 4HNO₂).^[64]
8. Electronegativity values of fluorine to iodine are: 3.98 + 3.16 + 2.96 + 2.66 = 12.76/4 3.19.
9. Helium is shown above beryllium for electron configuration consistency purposes; as a noble gas it is usually placed above neon, in group 18.
10. The net result is an even-odd difference between periods (except in the s-block): elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods (except the first) differ in the opposite direction. Many properties in the p-block then show a zigzag rather than a smooth trend along the group. For example, phosphorus and antimony in odd periods of group 15 readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth in even periods prefer to stay at +3.^[78]
11. Oxidation states do not reflect the actual net charge of atoms in molecules or ions, they represents the valence which refers more to how many bonds there are. For instance carbon typically has a valence of +4, but that only means that it forms three bonds. Electronegative elements such as fluorine are conventionally associated with negative valence, while electropositive ones have positive valence.
12. Greenwood^[84] commented that: "The extent to which metallic elements mimic boron (in having fewer electrons than orbitals available for bonding) has been a fruitful cohering concept in the development of metalloborane chemistry ... Indeed, metals have been referred to as "honorary boron atoms" or even as "flexiboron atoms". The converse of this relationship is clearly also valid."
13. For example, the conductivity of graphite is 3 × 10⁴ S•cm^−1.[85] whereas that of manganese is 6.9 × 10³ S•cm⁻¹.^[86]
14. A homopolyatomic cation consists of two or more atoms of the same element bonded together and carrying a positive charge, for example, N₅⁺, O₂⁺ and Cl₄⁺. This is unusual behavior for nonmetals since cation formation is normally associated with metals, and nonmetals are normally associated with anion formation. Homopolyatomic cations are further known for carbon, phosphorus, antimony, sulfur, selenium, tellurium, bromine, iodine and xenon.^[88]
15. Of the twelve categories in the Royal Society periodic table, five only show up with the metal filter, three only with the nonmetal filter, and four with both filters. Interestingly, the six elements marked as metalloids (boron, silicon, germanium, arsenic, antimony, and tellurium) show under both filters. Six other elements (113–118: nihonium, flerovium, moscovium, livermorium, tennessine, and oganesson), whose status is unknown, also show up under both filters but are not included in any of the twelve color categories.
16. The quote marks are not found in the source; they are used here to make it clear that the source employs the word non-metals as a formal term for the subset of chemical elements in question, rather than applying to nonmetals generally.
17. Varying configurations of these nonmetals have been referred to as, for example, basic nonmetals,^[98] bioelements,^[99] central nonmetals,^[100] CHNOPS,^[101] essential elements,^[102] "non-metals",^[103][p] orphan nonmetals,^[104] or redox nonmetals.^[105]
18. Arsenic is stable in dry air. Extended exposure in moist air results in the formation of a black surface coating. "Arsenic is not readily attacked by water, alkaline solutions or non-oxidizing acids".^[110] It can occasionally be found in nature in an uncombined form.^[111] It has a positive standard reduction potential (As → As³⁺ + 3e = +0.30 V), corresponding to a classification of semi-noble metal.^[112]
19. "Crystalline boron is relatively inert."^[106] Silicon "is generally highly unreactive."^[107] "Germanium is a relatively inert semimetal."^[108] "Pure arsenic is also relatively inert."^[109][r] "Metallic antimony is … inert at room temperature."^[113] "Compared to S and Se, Te has relatively low chemical reactivity."^[114]
20. Boundary fuzziness and overlaps often occur in classification schemes.^[117]
21. Jones takes a philosophical or pragmatic view to these questions. He writes: "Though classification is an essential feature of all branches of science, there are always hard cases at the boundaries. The boundary of a class is rarely sharp ... Scientists should not lose sleep over the hard cases. As long as a classification system is beneficial to economy of description, to structuring knowledge and to our understanding, and hard cases constitute a small minority, then keep it. If the system becomes less than useful, then scrap it and replace it with a system based on different shared characteristics."^[117]
22. For a related comparison of the properties of metals, metalloids, and nonmetals, see Rudakiya & Patel (2021), p. 36.
23. Metal oxides are usually somewhat ionic, depending upon the metal element electropositivity.^[126] On the other hand, oxides of metals with high oxidation states are often either polymeric or covalent.^[127] A polymeric oxide has a linked structure composed of multiple repeating units.^[128]
24. All four have less stable non-brittle forms: carbon as exfoliated (expanded) graphite,^[227][228] and as carbon nanotube wire;^[229] phosphorus as white phosphorus (soft as wax, pliable and can be cut with a knife, at room temperature);^[47] sulfur as plastic sulfur;^[48] and selenium as selenium wires.^[49]
25. Metals have electrical conductivity values of from 6.9×10³ S•cm⁻¹ for manganese to 6.3×10⁵ for silver.^[231]
26. Metalloids have electrical conductivity values of from 1.5×10⁻⁶ S•cm⁻¹ for boron to 3.9×10⁴ for arsenic.^[232]
27. Unclassified nonmetals have electrical conductivity values of from ca. 1×10⁻¹⁸ S•cm⁻¹ for the elemental gases to 3×10⁴ in graphite.^[85]
28. Halogen nonmetals have electrical conductivity values of from ca. 1×10⁻¹⁸ S•cm⁻¹ for F and Cl to 1.7×10⁻⁸ S•cm⁻¹ for iodine.^[85][233]
29. Elemental gases have electrical conductivity values of ca. 1×10⁻¹⁸ S•cm⁻¹.^[85]
30. Metalloids always give "compounds less acidic in character than the corresponding compounds of the [typical] nonmetals."^[218]
31. Arsenic trioxide reacts with sulfur trioxide, forming arsenic "sulfate" As₂(SO₄)₃.^[241] This substance is covalent in nature rather than ionic;^[242] it is also given as As₂O₃·3SO₃.^[243]
32. NO
    ₂, N
    ₂O
    ₅, SO
    ₃, SeO
    ₃ are strongly acidic.^[244]
33. H₂O, CO, NO, N₂O are neutral oxides; CO and N₂O are "formally the anhydrides of formic and hyponitrous acid, respectively viz. CO + H₂O → H₂CO₂ (HCOOH, formic acid); N₂O + H₂O → H₂N₂O₂ (hyponitrous acid)."^[245]
34. ClO
    ₂, Cl
    ₂O
    ₇, I
    ₂O
    ₅ are strongly acidic.^[246]
35. Metals that form glasses are: vanadium; molybdenum, tungsten; alumnium, indium, thallium; tin, lead; and bismuth.^[249]
36. Unclassified nonmetals that form glasses are phosphorus, sulfur, selenium;^[249] CO₂ forms a glass at 40 GPa.^[251]
37. Disodium helide (Na₂He) is a compound of helium and sodium that is stable at high pressures above 113 GPa. Argon forms an alloy with nickel, at 140 GPa and close to 1,500 K, however at this pressure argon is no longer a noble gas.^[259]
38. Values for the noble gases are from Rahm, Zeng and Hoffmann.^[260]