Reactive nonmetals: reactive nonmetals periodic table facts

Reactive nonmetals on the periodic table: definition and members

A nonmetal is a chemical element that mostly lacks the metallic properties you know from iron or copper: shine, bendability, high density and the ability to conduct electricity easily. Roughly 17 of the 118 known elements are classed this way, against about 90 metals, so nonmetals are a minority on the periodic table even though they dominate the chemistry of air, water and life. Magnets and lasers made with this metal feature in the holmium periodic table, with nuclear and physics notes alongside.

Within that broad family, chemists separate the inert noble gases from the elements that take part readily in chemical reactions. The second set is the reactive group, and on this periodic table it is a family of its own.

What ties the group together is behaviour rather than one column. Each member has four to six electrons in its outer shell (hydrogen has one and needs only one more), so it reacts by gaining electrons or by sharing them in covalent bonds.

That is why every element here forms a huge range of compounds, from simple molecules such as methane to the proteins and DNA in your cells. They show far less metallic character than the metals and only a little more than the neighbouring metalloids.

In chemistry, "reactive" is a relative word. Nitrogen gas barely reacts at room temperature, yet it still counts, because its atoms form strong bonds with hydrogen, oxygen and carbon once the energy barrier is crossed.

The label marks a contrast with the noble gases, not a promise that every member bursts into flame. All descriptions here refer to each element in its stable form at standard temperature and pressure (0 °C and 1 bar) unless noted, which matters because a few of these elements change form with temperature.

Different textbooks draw the family borders in different places. Some classification schemes list only two kinds of nonmetal elements, others list up to seven, and that choice decides where the group 17 elements end up. Astatine, the heaviest halogen, is usually left out of every nonmetal list because the few atoms ever made point to metallic behaviour.

Which elements count: the seven and the eleven

On this periodic table, the reactive nonmetal family holds seven elements:

• Hydrogen (H), element 1, the lightest element and the fuel of stars
• Carbon (C), element 6, the backbone of organic chemistry
• Nitrogen (N), element 7, about four-fifths of the air
• Oxygen (O), element 8, the gas your cells burn food with
• Phosphorus (P), element 15, part of DNA and of the ATP that carries energy in cells
• Sulfur (S), element 16, found in proteins, volcanic gases and gunpowder
• Selenium (Se), element 34, a trace nutrient and a light-sensitive semiconductor

Many other charts, including the scheme used on Wikipedia and in several chemistry courses, use a wider list of eleven by adding the four stable halogens: fluorine, chlorine, bromine and iodine. Both versions are correct within their own system. This site gives group 17 its own family because their chemistry is so uniform, and keeps the seven together because they share a quieter, more varied pattern.

Some older tables call the seven simply the "other nonmetals". Biologists know six of them by the acronym CHNOPS (carbon, hydrogen, nitrogen, oxygen, phosphorus, sulfur), and the literature also calls them the bioelements, because together they make up about 98% of the mass of a living cell.

Where the group sits: groups 1 and 14 to 16

Find the zigzag staircase line that runs from boron down to astatine on the right side of the periodic table, the line that divides metals from nonmetals. The metalloids sit along that line, the noble gases fill the last column, and the reactive group lives in between.

Carbon is in group 14; nitrogen and phosphorus are in group 15; oxygen, sulfur and selenium are in group 16. Hydrogen is the odd one out: it sits above lithium in group 1 because it has one electron, but it behaves far more like a nonmetal than like an alkali metal.

Because the members cluster at the top right, they are small atoms with high nuclear pull on their outer electrons. As you move down a group (oxygen to sulfur to selenium) the atoms grow, the outer electrons are held less tightly, and the elements become a little more metal-like. Selenium's chemistry already shows this shift, and it has a grey form with a metallic lustre, which is why it borders the metalloids.

Types of nonmetals: where the reactive group fits

Most modern tables split the non-metals into three or four types, read from right to left. You meet the noble gases first, then the halogens, then the mixed reactive set described in this guide, and finally the borderline metalloids that some authors count as non-metals and others do not. When a reader checks the highest melting point of any metal, they will find it in the pure tungsten entry along with boiling data.

The boundaries are not sharp, so think of these types as neighbourhoods rather than walled cities. Any classification of elements is a tool for grouping shared properties, and the table below shows how the four types compare with each other and with metals.

Unclassified nonmetals: the elements of life

Wikipedia and several research papers call the seven the unclassified nonmetals, a deliberately neutral name, because no single collective label has won out. Their chemistry is best described as moderately nonmetallic: less aggressive than fluorine or chlorine, so most of them occur naturally in the environment, sometimes even as free elements (native sulfur, diamond, the nitrogen and oxygen of the air).

Several shared habits set them apart. They like to bond to themselves, which gives carbon its chains and rings, sulfur its S₈ crowns and phosphorus its P₄ tetrahedra. With transition metals, small atoms such as hydrogen, carbon and nitrogen slip into the gaps of the metal lattice to form interstitial compounds such as steel's iron carbide, and tungsten carbide and titanium nitride are hard refractory compounds used on cutting tools.

Hydrogen even mimics a metal at times: it swaps places with sodium in chlorides (NaCl compared with HCl) and with potassium in nitrates (KNO₃ compared with HNO₃). Selenium is the most metal-like of the set; its electrical conductivity jumps roughly a thousandfold in light, the effect behind early photocopiers and light-sensing cells.

Halogen nonmetals and the salt-formers of group 17

The halogen nonmetals are fluorine, chlorine, bromine and iodine, the stable members of group 17. Each atom has seven valence electrons, one short of a full octet, so it grabs one electron from almost anything and becomes an ion with a 1− charge.

These ions are the halides (fluoride, chloride, bromide, iodide) found in salts and seawater. That single missing electron explains why the reactive halogen nonmetals are listed with the reactive group on so many charts.

All four exist as diatomic molecules (F₂, Cl₂, Br₂, I₂), and together they cover all three everyday states: fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. They are strong oxidizing agents, and with metals they form mainly ionic salts such as sodium chloride and potassium iodide.

Reactivity falls as you move down the group, because each larger atom holds the incoming electron less tightly: fluorine tops the list of every element, and iodine is the gentlest of the four. Read the dedicated halogens page for the full story of halogen chemistry; here they matter as the benchmark the seven are measured against.

Noble gases and metalloids on either side

To the right, the noble gases already have full outer shells, so they are relatively unreactive and form only a few hundred compounds, almost all of xenon or krypton with fluorine or oxygen. In bonding chemistry, a noble gas is the target configuration the reactive group is chasing whenever it bonds.

To the left, the metalloids boron, silicon, germanium, arsenic, antimony and tellurium look shiny but behave chemically much like weak nonmetals. Carbon, phosphorus and selenium border the metalloids so closely that some sources list them there instead, which is one reason element counts differ between charts.

How the nonmetal classification took shape

Antoine Lavoisier's 1789 textbook of modern chemistry, with its list of 33 substances, already placed oxygen, hydrogen and nitrogen among the gases and sulfur, phosphorus and carbon among the "nonmetallic substances" that form acids when burned. In 1811 the Swedish chemist Berzelius called every such element a "metalloid", a word that stuck in France for decades.

In 1844 the French pharmacist Alphonse Dupasquier split the nonmetals into four sets: organogens (oxygen, nitrogen, hydrogen, carbon), sulphuroids (sulfur, selenium, phosphorus), chloroides (fluorine to iodine) and boroids (boron, silicon). Put his first two sets together and you have almost exactly the seven-member reactive family used today; his chloroides became the halogens and his boroids grew into the metalloids.

Mendeleev's periodic table of 1869 to 1871 then gave these elements fixed columns, and by the 1870s "non-metal" had replaced "metalloid" as the accepted term, and the noble gases joined as a separate set after their discovery in the 1890s. Once x-ray crystallography and quantum chemistry arrived in the twentieth century, chemists could explain these old groupings by how atoms and electrons are arranged, which is also why the most abundant element in the universe, hydrogen, still sits awkwardly in every classification scheme and every periodic table design.

Physical properties of the reactive group of nonmetals

The physical properties of this family vary more than in any metal family. Three are gases, four are solids, and the solids range from transparent diamond to waxy white phosphorus.

What they share is how their properties contrast with metals: low density, weak or brittle solids, and generally low conductivity. In chemistry classes this contrast is usually the first test for sorting elements into metals and nonmetals.

States of matter, color and density

Hydrogen, nitrogen and oxygen are gases at room temperature because their small diatomic molecules attract each other only through weak dispersion forces, even though the bonds inside each molecule are strong. Carbon, phosphorus, sulfur and selenium are solids, because their atoms link into extended networks, rings or chains that need far more heat to break up. That is why their melting and boiling points span from about −259 °C for hydrogen to roughly 3,600 °C for carbon.

Colour tells the same story. Hydrogen and nitrogen are colorless in every state; liquid oxygen turns pale blue. Sulfur is lemon yellow, red phosphorus is dark red, and selenium appears red or grey depending on its form.

Among the halogens, fluorine is pale yellow, chlorine yellow-green, bromine reddish-brown, and solid iodine grey-black with a sheen, but it sublimes on warming into a violet vapour. Every one of the seven has a low density: selenium, the heaviest, is about 4.8 g/cm³, well under iron's 7.9 g/cm³. The solids are brittle rather than malleable, so a lump of sulfur shatters where a lump of lead would dent.

Electrical conductivity and the band gap

In solid-state chemistry terms, metals conduct because their outer electrons move freely through overlapping energy bands. Most nonmetals have a band gap: an energy jump that electrons must clear before they can carry current. Wide gaps make hydrogen, nitrogen, oxygen and sulfur insulators, and each behaves as an electrical insulator whether solid, liquid or gas; sulfur's electrical conductivity is about 10⁻¹⁸ S/cm, among the lowest of any element.

Two members break the rule in useful ways. Graphite has no band gap along its layers, so it conducts electricity better than some metals and is used in electrodes and pencil "leads". Selenium and black phosphorus are semiconductors with narrow gaps, which is why selenium was used in early solar cells and black phosphorus is being studied as a semiconductor for thin electronic films.

Carbon spans the whole range by itself: graphite is a conductor, diamond is a wide-gap insulator, and graphene-based forms are tuned into semiconductor devices. Thermal conductivity follows a similar spread: diamond moves heat faster than copper and most other metals because its rigid lattice carries vibrations so efficiently, yet the gaseous members are poor at it.

Allotropes: one element, several forms

More than half of the nonmetallic elements exist as allotropes, different structures of the same element with very different properties. Carbon is the famous case in materials chemistry: graphite (soft, grey, a solid lubricant), diamond (the hardest natural material and an insulator) and fullerene cages such as C₆₀. Phosphorus comes as white phosphorus (waxy, molecular, glows and ignites in air), red phosphorus (polymeric, used on match strikers) and black phosphorus (layered, semiconducting).

Oxygen exists as O₂ and as ozone, O₃; sulfur has dozens of ring and chain forms; selenium has red molecular and grey chain-like forms. When you compare chemistry data for these elements, always check which form the number describes.

Chemical properties and reactivity of non-metals

The chemical properties of the group all come back to one tendency that runs through nonmetal chemistry: in chemical reactions they gain or share enough electrons to reach a noble-gas arrangement. How strongly each element pulls on electrons sets how vigorously it reacts and what kind of bonds it forms, and those properties change smoothly with position on the periodic table.

Valence electrons and the octet rule

The octet rule says that atoms are most stable when their electron configuration ends with eight electrons in their outermost shell (two for hydrogen, the duet rule). Carbon has four valence electrons, nitrogen and phosphorus five, oxygen, sulfur and selenium six. To reach eight, they either gain electrons from a metal, forming anions, or share pairs with other nonmetal atoms:

$$\mathrm{O} + 2e^- \rightarrow \mathrm{O}^{2-} \qquad \mathrm{N} + 3e^- \rightarrow \mathrm{N}^{3-} \qquad \mathrm{H} + e^- \rightarrow \mathrm{H}^-$$

Ions such as oxide (O²⁻), sulfide (S²⁻) and nitride (N³⁻) appear throughout mineral chemistry and in everyday salts. Why do these atoms pull so hard?

Across a period, each step adds one more of the protons in the nucleus and raises the nuclear charge, while the number of inner electron shells stays the same, so the outer electrons are drawn closer. Smaller atoms with stronger pull sit at the top right of the table, exactly where this family lives.

Electronegativity, ionization energy and electron affinity

In chemistry, electronegativity measures how strongly an atom attracts the shared electrons in a bond. On the Pauling scale the seven run from 2.19 (phosphorus) to 3.44 (oxygen), all above the metal average of roughly 1.5. Their mean is easy to check:

$$\bar{\chi}_{\text{reactive}} = \frac{2.20 + 2.55 + 3.04 + 3.44 + 2.19 + 2.58 + 2.55}{7} = \frac{18.55}{7} \approx 2.65$$

The four halogens average 3.19, which matches their stronger appetite for electrons. Electronegativity also predicts bond type: a large difference, such as oxygen (3.44) against sodium (0.93), gives an ionic bond, while two similar values, such as carbon and hydrogen, give a covalent bond. The same periodic trends appear in two related measures.

Ionization energy, the energy needed to remove an electron, is high (about 1,000 to 1,400 kJ/mol for most of the seven), so these elements rarely form positive ions. Electron affinity, the energy released when an atom accepts an electron, is large for oxygen and sulfur but close to zero for nitrogen, a known exception caused by its stable half-filled 2p subshell. All three values fall as atomic radius grows down a group, which is why selenium is less reactive than sulfur, and sulfur less than oxygen.

Ranking the most reactive nonmetals

Ranked across every nonmetal on the periodic table, fluorine comes first, followed by oxygen and chlorine. Fluorine is so highly reactive that it attacks glass and even forms compounds with xenon. Within the seven elements, the order is roughly:

1. Oxygen: high reactivity with nearly every element, the driver of rusting and burning
2. White phosphorus: catches fire in air at about 30 °C, so it is stored submerged
3. Sulfur: moderately reactive, burns with a blue flame and combines with most metals when heated
4. Hydrogen: stable when cold but explosive with oxygen once ignited
5. Selenium: similar to sulfur but slower, showing the decrease in reactivity down group 16
6. Carbon: graphite and diamond react only when hot
7. Nitrogen: the N≡N triple bond (about 945 kJ/mol) makes the gas the least reactive member

In practical chemistry the ranking depends on form as well as element: black phosphorus is far more stable than the white allotrope, and atomic nitrogen is more reactive than oxygen gas, even though N₂ gas is almost inert.

Covalent compounds and acidic oxides

When two nonmetals meet, neither wants to give up electrons, so they share them in covalent bonds. That is why the group forms millions of covalent compounds: water (H₂O), ammonia (NH₃), carbon dioxide (CO₂), hydrogen sulfide (H₂S) and the whole of organic chemistry. With reactive metals, oxygen and the halogens tend to form ionic compounds instead, and the other members form salt-like, covalent or even metallic compounds depending on the partner.

A classic test in school chemistry is the oxide. Metal oxides are basic, but nonmetal oxides are usually acidic: CO₂ makes carbonic acid, SO₃ makes sulfuric acid, and NO₂ leads to nitric acid. These acidic oxides drive the chemistry of acid rain.

Dissolved, their acids release oxyanions such as sulfate, nitrate and phosphate. A few oxides are neutral (CO, N₂O) or amphoteric like H₂O, but none is basic. Hydrogen fluoride, from the halogen side, is the reminder that the halogens build acids too, without needing oxygen at all.

Where the reactive nonmetal family breaks the pattern

Textbook trends are tidy, but the seven members of this reactive family of nonmetals (H, C, N, O, P, S, Se) bend them in predictable places. The light members break the rules in one direction and the heavy members in another, and knowing those exceptions makes the chemistry of the group much easier to follow.

First-row anomaly and multiple bonds

The lightest member of each group behaves differently from the rest, an effect called the first-row anomaly. Hydrogen, carbon, nitrogen and oxygen are tiny atoms with no d orbitals and unusually high electronegativity for their position. Hydrogen can lose its single electron entirely in solution, leaving a bare proton that drives acid–base chemistry, and it forms hydrogen bonds that raise water's boiling point and hold DNA strands together.

Small size also lets carbon, nitrogen and oxygen form multiple bonds: C=C, C≡C, C=O, N≡N and O=O. Their heavier relatives in period 3 and below rarely do, so oxygen is a gas of O₂ molecules while sulfur forms rings of single bonds, and CO₂ is a gas while SiO₂ is solid sand. In compounds of nitrogen and oxygen, electron density is shared across several atoms at once, which gives the nitrogen oxides their many stoichiometries.

Oxidation state limits and metal-like behaviour

The octet rule predicts −3 for group 15 and −2 for group 16, as in NH₃ and H₂S, but the heavier members also reach high positive states. Phosphorus reaches +5 in PCl₅ and sulfur reaches +6 in SF₆ and sulfuric acid, while nitrogen's maximum oxidation state is +5 and oxygen almost never exceeds +2. Down group 16 a zigzag known as secondary periodicity appears, so selenium chemistry does not always sit neatly between sulfur and tellurium.

Property overlaps blur the metal border too. Graphite conducts better than manganese, selenium can be drawn into a wire like ductile metals, and under extreme pressure hydrogen is predicted to turn metallic. Several of these elements also build long chain molecules: carbon makes the plastics and other polymers we rely on, and sulfur forms a rubbery chain-like plastic form when molten sulfur is cooled fast.

Nonmetal elements compared in one property table

Use the table below to compare the seven elements side by side. Values are standard reference data for the most stable form at 25 °C and 1 atm (white phosphorus for melting and boiling points, graphite for carbon). The standard state column tells you whether you would hold a gas or a solid at room conditions.

Three patterns stand out. The atomic number jumps from 8 to 15 and then to 34, and atomic weight climbs from about 1 for hydrogen to about 79 for selenium, so the group is spread across four periods of the periodic table rather than packed into one.

The gases are all period 1 or 2 elements, and every period 3 and 4 member is a solid. And the electron configuration column shows why the chemistry repeats: oxygen, sulfur and selenium all end in s²p⁴, two electrons short of a noble gas.

Abundance, extraction and uses of nonmetallic elements

For a small corner of the periodic table, these nonmetallic elements carry an enormous share of the world's mass. Four members of this family of nonmetals, hydrogen, carbon, nitrogen and oxygen, make up most of the planet's air, oceans and living things, and silicate minerals build much of its rock.

Abundance on Earth and in living things

Hydrogen is about 75% of the ordinary matter in the universe by mass. On Earth the picture changes: the atmosphere is about 78% nitrogen and 21% oxygen, the hydrosphere is 86% oxygen and 11% hydrogen by mass, and oxygen alone makes up about 46% of Earth's crust, locked in silicate rock.

The biomass of living things is about 63% oxygen, 20% carbon and 10% hydrogen. Your own body follows the same pattern: roughly 65% oxygen, 18% carbon, 10% hydrogen, 3% nitrogen, 1% phosphorus and 0.25% sulfur, with selenium present only in micrograms as part of antioxidant enzymes.

Extraction from air, gas and rock

Industrial chemistry separates nitrogen and oxygen by cooling liquid air and boiling off each gas at its own temperature, a process called fractional distillation. Hydrogen mostly comes from natural gas by steam reforming of methane, and sulfur is recovered from the hydrogen sulfide in sour natural gas.

Carbon comes from coal, graphite mines and diamond deposits, phosphorus from phosphate rock heated with coke and sand, and selenium as a by-product of copper refining. Most of it comes from the Claus process, which partly burns that sour gas and lets the rest react with the dioxide it produces.

Everyday uses of each element

Hydrogen

Ammonia and fertiliser production, oil refining, and fuel cells that turn hydrogen and oxygen into electricity with water as the only exhaust.

Carbon

Steel, graphite electrodes and pencils, graphite as a solid lubricant, diamond cutting blades and jewellery, and activated charcoal filters.

Nitrogen

Fertilisers, an inert blanket for food packaging, and liquid nitrogen as a coolant at −196 °C for freezing samples.

Oxygen

Medical breathing gas, steelmaking, combustion in welding torches, and liquid oxygen as an oxidiser in rocket fuels.

Phosphorus

Phosphate fertilisers, safety-match strikers made from red phosphorus, and detergents.

Sulfur

Sulfuric acid, the most produced industrial chemical, plus vulcanised rubber, fungicides and gunpowder.

Selenium

Glass decolouring and red tinting, pigments, and photoconductive and light-sensitive devices.

Whether you are reading a chemistry syllabus or a food label, the same seven elements keep turning up, and their spot on the periodic table explains most of their properties. Understand how they share electrons, and most of everyday chemistry starts to make sense.