Halogens: halogens periodic table facts

What are the halogens?

The halogens are fluorine, chlorine, bromine, iodine, astatine and tennessine. The first four are stable nonmetals you can buy, store and study; astatine and tennessine are radioactive and exist only in tiny, short-lived amounts. Anyone comparing sphalerite with zinc blende will find the zinc on the periodic table guide useful for ore types, melting point and density.

Because all the halogens sit one electron short of a full shell, each pulls hard on the outer electron cloud of other atoms. That shared hunger explains why the halogens never occur free in nature and why chemistry teachers treat them as the textbook example of a family with clear, predictable trends.

When one of the halogens grabs that extra electron it becomes a negative ion called a halide: fluoride, chloride, bromide, iodide. Most of what you meet in daily life is a halide rather than the element itself. Table salts, ocean brines, tooth enamel minerals and the silver salts once used in photography are all halides, and the pure elements you see in a laboratory bottle were extracted from them.

Halogens

The elements of column 17; from Greek words meaning "salt" and "to produce".

Halide

The negative ion (X⁻) or a compound containing it, such as NaCl or KI.

Diatomic

Made of two identical atoms bonded together, the natural form of the free halogens (F₂, Cl₂, Br₂, I₂).

Where group 17 sits on the periodic table

You find the halogens in the second column from the right of the periodic table, wedged between the oxygen family (the chalcogens) and the noble gases. Older American tables label the column VIIA and older European ones VIIB; modern IUPAC numbering simply calls it column 17, and some chemists still refer to it as the fluorine group after its lightest member.

Position matters here: one step to the right sit the noble gases with their full shells, which is exactly the configuration halogens are trying to reach. Halogens near the top of the column are each a classic nonmetal, but the character becomes more metallic as you move down, and astatine is predicted to behave partly like a metal.

Why halogens earned the name salt former

The word halogen means "salt producer", built from the Greek hals (salt) and gennan (to produce). Johann Schweigger suggested it in 1811, and Jöns Jacob Berzelius made it stick in 1826 for elements that react with metals to give sea-salt-like solids. Each of the halogens is a salt maker in the most literal sense: chlorine plus sodium gives sodium chloride, the table salt in your kitchen; fluorine plus calcium gives calcium fluoride, the mineral behind fluorite crystals; bromine plus silver gives silver bromide, the light-sensitive salt of film photography; iodine plus potassium gives potassium iodide, the supplement in iodised salt.

The individual names follow a pattern too. Chlorine comes from chloros, Greek for greenish-yellow; bromine from bromos, a stench; iodine from iodes, violet; astatine from astatos, unstable. Fluorine takes its name from the Latin fluere, to flow, and tennessine from the US state of Tennessee.

Fluorine, chlorine, bromine, iodine and astatine: the halogen lineup

The quickest way to see the family resemblance, and the differences, is side by side. The table lists each element's atomic number, standard atomic mass and appearance at room temperature. Exposure limits, health effects and physics of alpha emitters are explained in thorium (Th) with plain language.

As the atomic number climbs, so does the atomic mass, the size of the atom and the strength of the forces holding it together, which is why you move from two gases to a liquid to solids going down the column.

Fluorine

Fluorine is a pale yellow gas and the most reactive element on the periodic table. It attacks glass, ignites many organic materials on contact and even forces the noble gas xenon into compounds. Fluorine never appears free in nature; you find it locked in fluorspar, cryolite and fluorapatite.

Industry handles it in nickel, copper or steel vessels, because those metals grow a thin fluoride skin that stops further attack. Despite the danger of the element, fluorine compounds are some of the most stable materials known, including the non-stick coatings on cookware.

Chlorine

Chlorine is a dense, choking gas with a sharp smell you recognise from swimming pools. It is the most abundant of the halogens in seawater and is mined mainly as halite, the natural rock-salt mineral.

Chlorine is less reactive than fluorine but still reacts with most elements, and its gas is roughly 2.5 times heavier than air, so it pools in low areas after a leak. Chlorine is one of the highest-volume chemicals made by industry, mostly as a building block for other products.

Bromine

Bromine is one of only two elements that are liquid at room temperature (mercury is the other). It is a heavy, red-brown liquid that gives off an orange, foul-smelling vapour, and it burns skin badly on contact.

Bromine is concentrated from salt lakes and underground brines, notably the Dead Sea and wells in Arkansas. Its chemistry sits neatly between chlorine and iodine, which made it the element early chemists used to test whether the column really followed a family pattern.

Iodine

Iodine is a lustrous grey-black solid that sublimes when warmed, turning straight from black crystals into a violet vapor without melting first at normal pressure. It is the least reactive of the stable halogens and the only one that dissolves in a starch solution to give the familiar deep blue-black colour used to test for starch.

Iodine is recovered from Chilean caliche deposits and from brines in Japan and the United States. Your body needs it in tiny, steady amounts, a point the health section returns to.

Astatine and tennessine

Astatine is so radioactive that no one has ever seen a visible sample; any piece large enough to look at would vaporise from its own decay heat. Its longest-lived isotope, astatine-210, has a half-life of only about 8.1 hours, and at any moment the whole Earth's crust holds well under a gram. Astatine was the second synthetic element ever made, after technetium.

Tennessine, element 117, was first produced in 2010 by a Russian-American team and named in 2016. Only a few dozen atoms have ever existed, so its chemistry is predicted rather than measured, and some chemists argue it may not behave like a true halogen at all.

Physical properties of the group 17 elements

The physical properties of the halogens change smoothly down the column, which makes them one of the cleanest illustrations of periodic trends you can find. Colours deepen, densities rise and the state at room temperature shifts from gas to liquid to solid. The numbers below are standard reference values.

Diatomic molecules and intermolecular forces

Free halogens exist as diatomic molecules, two atoms sharing a single covalent bond: F₂, Cl₂, Br₂ and I₂. These halogen molecules are nonpolar, so the only attraction between neighbours is London dispersion forces, the weak, temporary pulls that arise when electron clouds shift for an instant.

Bigger ones have more electrons and larger, more easily distorted clouds, so the attraction grows going down the group. In solid iodine the pairs pack into layers, and the bond length stretches from about 143 pm in F₂ to about 267 pm in I₂.

Melting points and boiling point trends

Because dispersion attraction grows with size, the melting and boiling point climbs steadily from fluorine to astatine. Fluorine boils more than 180 degrees below zero, chlorine just below −34 °C, bromine does not boil until almost 59 °C, and iodine needs about 184 °C.

The atomic radius rises in step, from 71 pm to 133 pm across the stable four. A useful rule of thumb for exams: if you know where an element sits in the column, you can rank its volatility, colour intensity and density without memorising each value.

Halogen reactivity and electronegativity trends

Chemical reactivity runs the opposite way to the physical trends: it falls from fluorine down to iodine. Two numbers explain most of it. Electronegativity, the pull an atom exerts on a shared bond pair, drops from 3.98 for fluorine to 2.66 for iodine, and fluorine is the only element more electronegative than oxygen.

The cause is effective nuclear charge: the nucleus of a small atom sits closer to its outer electrons, so it holds an incoming electron more tightly. Larger atoms in heavier halogens have more shielding inner shells between the nucleus and the outermost electrons, which weakens that pull.

Seven valence electrons and the octet rule

Each halogen atom has an s²p⁵ outer arrangement: two s and five p electrons. The octet rule says atoms are most stable with eight, so each atom needs only one more to match the noble gas configuration of its neighbour to the right. Gaining that one electron releases energy, measured as electron affinity, and halogen atoms have some of the highest values of any element.

Chlorine, not fluorine, has the largest at about 349 kJ/mol, because fluorine's tiny outer shell is so crowded that the extra electron feels some repulsion. That is the first hint that fluorine does not always fit the pattern neatly.

Oxidizing agents and reduction potential

A substance that takes electrons from another is an oxidizer, and the halogens are among the strongest oxidizing agents in chemistry. Their standard reduction potentials rank the strength directly:

The practical consequence is the displacement reaction: a halogen higher in the column pushes a lower one out of its salt. Bubble chlorine through a bromide solution and the solution turns orange as free bromine appears:

$$\mathrm{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2}$$

The reverse never happens on its own, which is why iodine cannot pull chlorine out of common salt.

Bond energy and the fluorine anomaly

You might expect the X–X bond to weaken steadily down the group, but fluorine breaks the trend. Its bond is weaker than chlorine's or bromine's:

The two fluorine atoms are so small that their lone pairs crowd each other and push apart. That weak F–F link is easy to break, while the bonds fluorine then forms with carbon, hydrogen or metals are among the strongest known. Together these two facts make fluorine both violently reactive and the maker of extremely durable products.

Oxidation state patterns

All halogens show an oxidation state of −1 in its halides. Fluorine almost never shows anything else, because nothing is electronegative enough to strip charge from it.

Chlorine, bromine and iodine can also share electrons with more electronegative partners such as oxygen, so they reach positive oxidation states of +1, +3, +5 and +7. The chlorine oxyacids show the full ladder:

• +1 in hypochlorous acid, HClO, the active agent in chlorine pool treatments
• +3 in chlorous acid, HClO₂
• +5 in chloric acid, HClO₃
• +7 in perchloric acid, HClO₄, used in rocket propellant salts

Iodine reaches +7 most comfortably, in periodates, because its large atom has room for many neighbours around it.

Halogen compounds: halides, acids and interhalogens

Halogens bond in two broad ways. With reactive metals they take charge outright and form ionic bonds, giving hard crystalline salts with high melting temperatures.

With nonmetals they share electrons through covalent bonds, giving substances that are often gases or volatile liquids. Most of the useful compounds of these elements fall into three families: halides of hydrogen, halides of metals and compounds made only of halogens.

Hydrogen halides

Each element in the column combines with hydrogen to give a binary compound HX. The general reaction is:

$$\mathrm{H_2 + X_2 \rightarrow 2HX}$$

How violently it goes follows the reactivity trend. Hydrogen and fluorine explode even in the dark and cold; hydrogen and chlorine explode when lit by sunlight or a spark; bromine needs heat; iodine reacts only partially and reaches an equilibrium. The four hydrogen halides are hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr) and hydrogen iodide (HI).

All of them dissolve in water and release H⁺ ions. Hydrochloric, hydrobromic and hydroiodic acid are strong acids that ionise completely, while hydrofluoric acid is a weak acid because the H–F bond is so strong. Do not let the label fool you, though: hydrofluoric acid is one of the most dangerous laboratory chemicals, because it soaks through skin and strips calcium from bone.

Hydrogen fluoride also stands apart by forming hydrogen bonds, so it boils at about 20 °C while hydrogen chloride boils near −85 °C. Hydrochloric acid is the same acid your stomach produces to digest food.

Metal halides

Sodium burns in chlorine with a bright orange flame to give the archetype of a metal halide:

$$\mathrm{2Na + Cl_2 \rightarrow 2NaCl}$$

Iron shows how the halogen's strength sets the product. Fluorine, chlorine and bromine all take iron to its +3 state, and hot iron wire glows in chlorine to give black-brown iron(III) chloride, used to etch copper circuit boards.

Iodine is too weak an oxidizer to do the same, so iron and iodine stop at iron(II) iodide. Silver halides show the trend in colour: silver chloride is white, silver bromide cream and silver iodide yellow, a contrast you can use in the lab to tell which halide a solution contains.

Interhalogen compounds and the superhalogen

Two different halogens can bond to each other, giving interhalogen compounds of the form XY_n, where n is 1, 3, 5 or 7 and the larger halogen sits in the middle. Examples include ClF, BrF₃, IF₅ and IF₇, the only stable compound with seven fluorine atoms around one centre.

They are generally more reactive than the parent elements, and chlorine trifluoride is notorious for setting sand, asbestos and concrete alight. A related idea is the superhalogen: a cluster such as AlCl₄ or BF₄ whose electron affinity is higher than that of any single halogen atom, which lets it act as an extra-strong electron acceptor in modern materials and battery research.

Discovery history of the halogen elements

Chlorine came first. In 1774 Carl Wilhelm Scheele heated hydrochloric acid with manganese dioxide and collected a yellow-green gas, but he believed it was an oxide.

Humphry Davy proved it was an element in 1810 and named it. Hydrochloric acid itself had been known for centuries, and alchemists used it in aqua regia, the nitric-hydrochloric mixture that dissolves gold.

Iodine followed in 1811, when Bernard Courtois added too much sulfuric acid to seaweed ash and saw purple fumes condense into dark crystals. Bromine was discovered independently by Antoine Balard in 1825 and Carl Löwig in 1826, both from salt waters.

Fluorine was the hardest: its compounds were known from the 1500s, but several chemists were poisoned or blinded trying to free it. Henri Moissan finally isolated it in 1886 by electrolysis of potassium bifluoride dissolved in liquid HF, cooled to −50 °C, and won the 1906 Nobel Prize for the work.

Astatine was predicted by Mendeleev as "eka-iodine" and made in 1940 by Dale Corson, Kenneth MacKenzie and Emilio Segrè, who bombarded bismuth with alpha particles. Tennessine completed the column in 2010 when a team at Dubna, working with Oak Ridge and other US laboratories, fired calcium-48 into berkelium-249.

Producing group VIIA elements from brine, minerals and seaweed

Because the free elements are too reactive to survive in nature, all halogens you buy have been extracted from its salts. The method depends on how hard the halide clings to its electron.

• Fluorine: made only by electrolysis of potassium fluoride dissolved in liquid HF. No chemical oxidizer is strong enough. The HF itself comes from treating fluorspar with sulfuric acid.
• Chlorine: made by electrolysis of brine (sodium chloride solution) in the chlor-alkali process, which yields sodium hydroxide and hydrogen at the same time.
• Bromine: freed from bromide-rich brine by bubbling chlorine through it, then blowing the bromine out with air or steam.
• Iodine: extracted from iodate in Chilean nitrate deposits or from oilfield brines; historically it came from burning seaweed.

To see the bromine step in numbers, take one cubic metre of seawater carrying about 67 grams of bromide ions. That is 67 ÷ 79.90 ≈ 0.839 mol of Br⁻, and since each Cl₂ frees two bromide ions, you need about 0.419 mol of chlorine, or roughly 29.7 g of Cl₂, to release the full 67 g of bromine. In practice producers start from salt lakes far richer than open seawater, because pumping thousands of tonnes of water for a few grams per cubic metre costs more energy than the bromine is worth.

Natural abundance in Earth's crust and seawater

Chlorine dominates the oceans, while fluorine is the most common of the halogens in rock. Typical reference values are:

Those figures are in parts per million by mass for rock. The contrast tells a story: chloride and bromide salts dissolve easily and have been washing into the ocean for billions of years, which is why salt water is so rich in them, while fluoride stays bound in minerals such as fluorite and apatite.

Isotopes and radioactive forms

Fluorine and iodine each have just one stable isotope (fluorine-19 and iodine-127). Chlorine has two, chlorine-35 at about 76 per cent and chlorine-37 at about 24 per cent, which is why its atomic mass is 35.45 rather than a whole number. Bromine's two stable isotopes, bromine-79 and bromine-81, occur in almost equal amounts.

Several radioactive isotopes do real work. Fluorine-18, with a half-life near 110 minutes, labels the glucose tracer used in positron emission tomography (PET) scans. Iodine-131 treats an overactive thyroid and its cancers, because the gland absorbs iodine and concentrates the dose right where it is needed.

Chlorine-36 dates groundwater up to about a million years old, and isotopes of iodine released by nuclear tests now trace how waste spreads through the environment. Astatine has no stable isotope at all.

How halogens are used: disinfectants, lighting, plastics and medicine

The same electron hunger that makes these elements dangerous also makes them useful. Their compounds kill microbes, carry light, build durable materials and deliver drugs. Some uses, such as Freon refrigerants, have been phased out after their side effects became clear.

Disinfectants and water treatment

Chlorine is the world's most widely used chemical for water treatment. A small residual dose keeps drinking water free of cholera and typhoid bacteria all the way to the tap; the US EPA caps that residual at 4 mg/L.

Sodium hypochlorite is the active ingredient in bleach, iodine tinctures and povidone-iodine scrub skin before surgery, and bromine compounds sanitise hot tubs where chlorine breaks down too fast. Fluoride is added to drinking water at about 0.7 mg/L in the United States to reduce tooth decay.

The halogen lamp and lighting

A halogen lamp is a tungsten filament bulb filled with a little iodine or bromine gas. Tungsten that evaporates from the hot filament combines with the halogen atoms near the cooler glass, drifts back, and is redeposited on the filament. This regenerative cycle lets the lamp run hotter and brighter than an ordinary incandescent bulb without blackening its quartz envelope, though LEDs have now replaced most of them.

Refrigerants, PVC and flame retardants

Chlorine goes into polyvinyl chloride (PVC), the plastic used for pipes, cables and window frames, and into solvents such as chloroform. Fluorine gives Teflon its non-stick surface and gave the chlorofluorocarbons (CFCs) sold as Freon their stability.

Organobromides are the most important class of flame retardants, added to device casings, foams and textiles to slow ignition. Halogenated pesticides such as DDT were once used worldwide before their persistence led to bans.

Drug components and medicine

About a fifth of modern drugs contain fluorine, because a carbon–fluorine bond blocks the body from breaking a molecule down too quickly. Familiar examples include fluoxetine and several statins. Chlorine appears in many antibiotics and antihistamines, iodine-based contrast agents make blood vessels visible on X-rays and CT scans, and potassium iodide tablets protect the gland after a nuclear accident by saturating it with stable iodine.

Halogen toxicity, health and environmental concerns

All four stable elements are toxic as free elements. Fluorine and chlorine gases destroy lung tissue, bromine causes deep skin burns, and toxic iodine vapour irritates the eyes and airways. US workplace limits reflect the danger: 0.1 parts per million for fluorine and a 1 ppm ceiling for chlorine.

Chlorine gas was released as a weapon at Ypres in 1915 because it is heavier than air and sinks into trenches. Hydrogen fluoride is corrosive, highly toxic and can be lethal after a skin splash covering only a few per cent of the body.

Concentrated hydrogen bromide and hydrogen iodide are similarly harsh irritants. The toxicity of a halogen compound, however, depends far more on the molecule than on the element: chloride in salt is essential, chlorine gas is deadly.

Biological role of iodine, fluoride and chloride

Chloride is the main negative ion in your blood and stomach acid, and you need several grams a day from salt. Iodine is a trace nutrient: adults need about 150 micrograms daily to make the hormones thyroxine and triiodothyronine, and deficiency causes goitre and developmental problems.

Fluoride strengthens tooth enamel by turning some of it into fluorapatite, though too much causes mottling. Bromide has a small role in building tissue collagen, and astatine has none, being far too radioactive.

Halogen-free electronics and dioxins

Electronics makers have been moving to halogen-free circuit boards, solder fluxes and cable jackets. The worry is not normal use but disposal: when brominated or chlorinated plastics burn in uncontrolled fires or poor incinerators, they can form dioxins and furans, persistent pollutants linked to cancer. Modern high-temperature incinerators largely prevent this, but standards such as IEC 61249-2-21 still set limits on chlorine and bromine content in boards so that waste stays safer wherever it ends up.

Ozone layer damage and PFAS

The stability that made CFCs attractive let them drift into the stratosphere, where ultraviolet light frees chlorine atoms that destroy ozone in a chain reaction; a single chlorine atom can break down thousands of ozone units. The 1987 Montreal Protocol phased them out, and the ozone layer is now slowly recovering.

A newer concern involves PFAS, the "forever chemicals" built on chains of carbon–fluorine bonds. They resist breakdown in the environment and build up in people and wildlife, and in 2024 the EPA set the first enforceable US drinking water limits for several PFAS at 4 parts per trillion.