Antimony (symbol Sb, atomic number 51) is a lustrous silver-gray metalloid in group 15 (the nitrogen group) of the periodic table It is solid and brittle at normal temperature, often occurring in flaky or crystalline form Natural antimony is usually found combined in minerals, chiefly the sulfide stibnite (Sb₂S₃) rather than in the free metal. Its name derives from Latin antimonium, and the chemical symbol Sb comes from the Latin stibium (the old name for stibnite). Antimony’s common oxidation states are –3 (as in the hydride or Zintl anions), +3 (the dominant state in many compounds), and +5 At standard conditions antimony is a solid (mp ≈903 K; bp on the order of 1650–1900 K) with a density of about 6.7 g·cm⁻³ Its solid form is a poor conductor of heat and electricity (more metallic than a true nonmetal but much less conducting than typical metals).

Atomic Structure and Electron Configuration

Antimony
Atomic number	51
Symbol	Sb
Group	15 (pnictogens)
Boiling point	1587 °C
Electronegativity	2.05 (Pauling)
Electron configuration	[Kr] 4d10 5s2 5p3
Melting point	630.6 °C
Period	5
Main isotopes	121Sb, 123Sb
Oxidation states	−3, +3, +5
Phase STP	Solid
Wikidata	Q1099

Antimony’s atomic structure completes the 5p shell. Its ground-state electron configuration is [Kr] 4d¹⁰ 5s² 5p³ so it has five valence electrons. These are arranged as 5s² 5p³ in the outer shell, giving antimony a half-filled p subshell. Being a heavier pnictogen, antimony has a relatively large atomic radius (empirical atomic radius ≈145 pm and a modest electronegativity (Pauling scale ≈2.05 Its first ionization energy is about 8.64 eV (roughly 835 kJ/mol) reflecting a tendency to lose p-electrons but not as readily as a typical metal. In periodic trends, Sb lies below arsenic and above bismuth, so it is larger and more metallic than As but somewhat smaller than Bi (the lanthanide contraction limits Bi’s increase). Electronegativity and ionization energies decrease down the group: antimony’s electronegativity of ~2.05 is lower than arsenic’s (~2.18), and its first ionization (8.64 eV) is lower than arsenic’s (~9.81 eV). Overall, antimony’s atomic structure — with a partial p-shell and a lone 5s² pair — underpins its amphoteric chemistry (tendency to form covalent bonds but also +3/+5 oxidation states) and its semimetallic properties.

Isotopes and Nuclear Properties

Naturally occurring antimony consists almost entirely of two stable isotopes: ^121Sb (≈57.3%) and ^123Sb (≈42.7%) These isotopes have nuclear spins of 5/2 (^121Sb) and 7/2 (^123Sb), respectively, which make them observable by NMR spectroscopy (though they have large quadrupole moments, so the signals are broad) No other stable isotopes exist, but several radioactive antimony isotopes have moderate half-lives: for example, ^124Sb (t½ ≈ 60.3 days) decays to ^124Te (β⁻), and ^125Sb (t½ ≈ 2.76 years) decays to ^125Te These radioisotopes are mainly of interest in nuclear science or medical tracer production. In fact, ^121Sb can be irradiated to produce ^124I (for diagnostic imaging) and both ^121Sb and ^123Sb can be used to make ^123I (a gamma-imaging isotope) There are no naturally long-lived radioisotopes of antimony of geologic age, so antimony is not used for long-term radiometric dating. Its two stable isotopes do allow mass spectrometric and NMR studies of Sb-containing compounds.

Allotropes and Typical Compounds

Elemental antimony has a limited number of allotropes. The most common form (often called “antimony” or α-Sb) is a silvery-gray crystalline solid with an A7 (rhombohedral) structure (the same structure as gray arsenic) This form is brittle and flakes easily. A second allotrope is a noncrystalline gray ‘amorphous’ form that can be produced, for example, by rapid precipitation or electrolysis. This amorphous Sb is sometimes called “black antimony”. A highly unstable variety, occasionally termed “explosive antimony”, can be made by electrolysis of SbCl₃; it is an amorphous, halogen-rich Sb glass that detonates on friction or impact, reverting explosively to the stable metallic form In practice only the grey crystalline form is of industrial importance.

Antimony forms two main classes of compounds, corresponding to oxidation states +3 and +5 (and, rarely, –3 as in metal-rich alloys). In the trivalent state Sb³⁺, important compounds include antimony trioxide (Sb₂O₃), antimony trisulfide (Sb₂S₃), and the trihalides (SbCl₃, SbBr₃, SbI₃). For example, Sb₂O₃ is a white solid extensively used as a fire-retardant filler (often with halogenated polymers) and Sb₂S₃ is the mineral stibnite (used historically in cosmetics). Antimony trichloride (SbCl₃, “butter of antimony”) is a covalent liquid, and SbH₃ (stibine) is an unstable, foul-smelling gas analogous to PH₃ and AsH₃ (it decomposes readily to Sb and H₂). Lower oxidation state compounds (e.g. Sb–Sb bonded Zintl phases like Na₃Sb) exist but are mostly of academic interest.

In the pentavalent state Sb⁵⁺, one finds antimony pentoxide (Sb₂O₅), pentachloride (SbCl₅), and pentasulfide (Sb₂S₅). SbCl₅ is a colorless fuming liquid and a very strong Lewis acid (it forms the superacid HSbCl₆ when combined with HCl). Sb₂O₅ occurs only under strongly oxidizing conditions (e.g. in ores altered by weathering) and is also used as a catalyst or in glassmaking. Both Sb(III) and Sb(V) oxides are amphoteric: Sb₂O₃ dissolves in strong base to give antimonite (for example NaSbO₂ in NaOH), and Sb₂O₅ dissolves in alkali to give meta-antimonate (SbO₃⁻). In acids, these oxides form solvated antimony ions; for example, dissolving Sb₂O₃ in strong acid yields Sb³⁺ or mixed oxy-halide complexes. Overall, antimony’s compounds usually display covalent bonding with a tendency for Sb to form three bonds and hold one lone pair (as in SbCl₃) in the +3 state, or six-coordinated octahedral in the +5 state (as in Sb₂O₅). Many Sb compounds share similarities with arsenic and phosphorus analogues but tend to be more metallic in character.

Physical Properties

Solid antimony crystallizes in the rhombohedral A7 structure (a layered arrangement) It has a metallic luster and gray color. At 20 °C its density is about 6.7 g·cm⁻³ Antimony melts at roughly 630 °C (903 K) and boils at on the order of 1600–1650 °C (1873–1923 K), although published boiling values vary (sources range from ~1380 °C to ~1750 °C). Its heat of fusion is about 19.9 kJ/mol and heat of vaporization ~77 kJ/mol. Elemental Sb is a poor thermal conductor compared to metals (thermal conductivity ~18–24 W·m⁻¹·K⁻¹ It is a semimetal or “metalloid”: electrically, it conducts slightly better than an insulator but much worse than a metal (room-temperature resistivity ≈40×10⁻⁶ Ω·m). Magnetically, antimony is diamagnetic (it has no unpaired electrons).

Spectroscopically, antimony and its compounds are notable mostly for flame coloration and emission lines. When burned in a flame, many Sb compounds color the flame a pale green/blue (a distinctive test for Sb). In atomic emission spectra, Sb produces sharp blue-green lines (Sb produces a strong yellow-green line at ~405 nm, and other lines in the blue), but these are rarely used outside analytic chemistry. No prominent public-day emission color is associated, yet since ancient times the green flame of Sb compounds in fireworks and glazes has been noted.

Chemical Reactivity and Trends

Antimony’s chemistry is intermediate between metals and nonmetals, with amphoteric behavior. In air at ordinary temperatures, the metal tarnishes slowly to form a thin oxide film (mostly Sb₂O₃). When heated strongly in oxygen it oxidizes to Sb₂O₃ or to Sb₂O₅ under very oxidizing conditions. At red-heat it glows but does not readily ignite. Antimony does not react with cold water and is only slowly attacked by hot acids. Concentrated nitric acid or hot concentrated sulfuric acid oxidize it to H₃SbO₃ or H₅SbO₆ (soluble antimonic acids) and release inert gas (NO, SO₂). Aqua regia dissolves Sb completely, yielding soluble SbCl₃. Antimony reacts vigorously with halogens: for example, Cl₂ and Br₂ give SbCl₃/SbCl₅ and SbBr₃, respectively. Sb metal also dissolves in warm alkali under strong oxidizers, producing antimonate (SbO₄³⁻) species. In simple terms, Sb³⁺ tends to form covalent-cationic salts (e.g. SbCl₃) and oxides/hydroxides, while Sb⁵⁺ forms more ionic oxoanions (SbO₄³⁻). The hydride SbH₃ (stibine) is easily decomposed by heat or catalysts, so it is best known as a laboratory curiosity.

Chemically, antimony is more reactive than bismuth but less so than arsenic or phosphorus. It does not, for example, form a stable oxide as easily as phosphorus (no Sb analogue of P₂O₅ except Sb₂O₅ under forcing conditions). In water, elemental Sb remains largely inert (drinking water has no free Sb³⁺ except small traces). In organic chemistry it shows little reactivity apart from formation of alkyl antimonides (e.g. R₃Sb, often prepared from SbCl₃). Antimony corrosion is generally mild: environments that attack iron or copper often leave Sb untouched, and it often passivates by forming an oxide layer. The reactivity trends reflect its place mid-way down the pnictogens: it shows no “supermicidal” behavior but can form stable higher oxides and colored complexes (e.g. complex antimonates) more readily than heavier Bi.

Occurrence and Production

Antimony is a relatively rare element, about one-fifth as abundant in Earth’s crust as arsenic – roughly 1 gram per ton of crust It is not found free in nature; instead, major ores are sulfides and sulfosalts. The primary mineral is stibnite (Sb₂S₃), which historically funded most Sb production. Other antimony minerals include valentinite and senarmontite (both Sb₂O₃), kermesite (Sb₂S₂O), and various complex sulfosalts. Significant ore deposits exist in China, Russia, Tajikistan, Bolivia, and South Africa; China dominates global supply by a wide margin In fact, China has been the world’s largest producer (often roughly half of annual output) and possesses extensive reserves (currently ~20–25 years of supply at current demand Smaller quantities of Sb ore have been mined historically in Europe (e.g. Italy and Cornwall, UK) and North America, but most industrial antimony today originates from Asia.

Commercial production typically begins with roasting or oxidizing Sb₂S₃ ore to form Sb₂O₃ and SO₂. For example, 2 Sb₂S₃ + 9 O₂ → 2 Sb₂O₃ + 6 SO₂. The resulting Sb₂O₃ (a white sesquioxide) is then reduced to Sb metal. The classic method (described by Basil Valentine in the 17th century was to mix Sb₂O₃ with carbon (e.g. charcoal) and heat strongly: Sb₂O₃ + 3 C → 2 Sb + 3 CO₂. An alternative route (Andreas Libavius, 1615 reduces stibnite directly with iron: 2 Sb₂S₃ + 3 Fe → 3 FeS₂ + Sb. Modern producers often use either smelting or a hydrometallurgical extraction depending on ore type. China’s Xikuangshan deposit in Hunan province (the world’s largest antimony mine) historically followed these routes. In recycled sources, antimony can also be recovered from scrap (e.g. lead-acid battery lead is often 2–10% Sb), but such recycling is minor compared to primary mining. Global annual production in recent years has been on the order of 150,000–200,000 metric tons (China providing most of that)

Applications and Technology

Antimony’s chief industrial use is in flame-retardant formulations. Antimony trioxide (Sb₂O₃), often combined with halogenated flame retardants, greatly enhances fire resistance in plastics, textiles, and paints For example, polyvinyl chloride (PVC) and polypropylene used in electronics often contain Sb₂O₃; when exposed to fire, the Sb compound promotes char formation and suppresses smoke. Antimony compounds also appear in other safety applications: they are used in many sound-deadening foams, wire coatings, and building materials to slow combustion.

In metallurgy and alloys, small amounts of antimony impart hardness and strength. Lead–antimony alloys (typically 2–5% Sb) are common in lead-acid battery plates and in bullet/shot casting; Sb improves mechanical strength and ease of casting. (Older “type metal” used for printing contained ~12% Sb to produce a light, hard alloy.) Antimony is also an ingredient in various solder and bearing alloys (e.g. Babbitt metal).

In electronics, antimony is important in specialized semiconductors. III–V antimonide compounds (such as indium antimonide InSb, gallium antimonide GaSb, and aluminum antimonide AlSb) cover bandgaps from ~1.6 eV down to 0.17 eV This makes them very useful for infrared detectors, LEDs, and diode lasers. For example, InSb is used in infrared thermal cameras and high-speed transistors, while GaSb and related alloys are used in diode lasers and photo-detectors around 1–2 µm (common in fiber communications and LiDAR) Antimony is also used as a dopant in silicon and germanium; as a heavy group-15 dopant it creates deep n-type levels useful in some device designs.

Antimony compounds have niche uses in glass, ceramics, and catalysts. Antimony pentoxide (Sb₂O₅) is widely used as a catalyst in synthesizing polyethylene terephthalate (PET) plastic – it promotes the condensation reaction that forms the polymer. Some specialty glasses (used in optics or infrared windows) are doped with small amounts of Sb₂O₃ to modify refractive index or to scavenge oxygen. Certain types of enamel paints and ceramics also use antimony compounds as opacifiers or pigments (e.g. antimony yellow, a lead–antimony oxide pigment). In the past, antimony compounds were used in medicine (notably antimony potassium tartrate as “tartar emetic” to induce vomiting or treat parasitic diseases), but these uses have largely ceased due to toxicity In modern medicine, a related Sb(III) drug (sodium stibogluconate) is still used against leishmaniasis in some regions though safer treatments are preferred where available.

Biology, Environment, and Safety

Antimony has no known biological role in humans or animals and is considered a toxic heavy element. Its toxicity is chemically comparable to arsenic: trivalent antimony compounds (Sb³⁺) in particular interfere with cellular enzymes and can damage multiple organ systems High-dose exposure (as in industrial accidents) can cause vomiting, abdominal pain, diarrhea, and even heart or neurological effects. Chronic exposure (workers in mining or processing) is associated primarily with lung irritation and lung disease, and also cardiovascular problems. The International Agency for Research on Cancer (IARC) has classified trivalent antimony and its compounds as “probably carcinogenic” to humans (Group 2A) based on animal studies of lung cancer In contrast, pentavalent antimony (Sb⁵⁺) is considered not classifiable (Group 3) for carcinogenicity

Environmental anthropogenic sources of Sb include mining/waste, flame-retardant-containing consumer goods, and lead-acid battery recycling. Near shooting ranges and metal foundries, elevated Sb levels are found in dust and soil In water, Sb(III) and Sb(V) species can leach from polyethylene terephthalate (PET) bottles or glass, but tap-water concentrations are normally very low. Because Sb(III) and Sb(V) oxidize and precipitate easily, it tends to bind to soils and sediments rather than circulate freely. Regulatory limits reflect its toxicity: for example, the WHO guideline for drinking water recommends a tolerable daily intake of about 6 μg Sb per kg body-weight (roughly 20 ng/mL in water) and U.S. occupational exposure limits for Sb (as metal or compounds) are an 8-hour time-weighted average of 0.5 mg/m³

Safe handling of antimony (especially dust or soluble compounds) requires ventilation and respiratory protection. Inhalation of Sb₂O₃ dust or SbCl₃ fumes can irritate lungs. Chronic exposure should be minimized: employers monitor Sb in air and provide protective clothing and training. Ingestion is usually avoided (food containers now use cobalt or zinc oxides instead of Sb₂O₃ to avoid leaching). Because antimony compounds can accumulate in the body, biological monitoring (urine/ blood Sb) may be required in industrial settings. Overall, normal industrial hygiene prevents serious health effects, but antimony’s potential as a carcinogen and toxin means it is treated with caution in environmental and workplace standards.

History and Etymology

Antimony has a rich history dating to antiquity. Its compounds (especially the sulfide, stibnite) were known and used by ancient civilizations. Egyptian make-up (“kohl”) famously used ground Sb₂S₃ to darken eyelids (records trace back to ~4000 BCE) The Greek physician Dioscorides (~1st century CE) mentioned metallic antimony and recommended it for treating skin ailments Roman writers such as Pliny the Elder also noted medicinal uses of antimony minerals, calling them stibi or stimmi

Pure antimony metal was first isolated in the Middle Ages. Geological alchemists roasting stibnite and reducing the oxide produced a brittle silvery metal. The 17th-century alchemist Basil Valentine (author of Triumphal Chariot of Antimony, 1604) described obtaining metallic Sb by roasting the sulfide and reducing with charcoal German chemist Andreas Libavius later (1615) described a more efficient process: reducing stibnite directly with iron filings to get free Sb Over time, chemists realized this element was distinct: in 1818, German chemist Friedrich Wöhler showed Sb is different from arsenic.

The name antimony comes from medieval Latin antimonium, whose origin is uncertain; it may be a corruption of Greek terms or an Arabic miner name. An etymology often cited (though debated) is Greek anti- + monos (“not alone”), supposedly referring to the fact that antimony is never found in a pure native form In any case, the ancient term stibium (from Greek stubo, “mark” or “paint”) gave rise to the symbol Sb. Some old alchemical terms were “stibium” or “regulus of antimony,” reflecting confusion and superstition (antimony was thought by alchemists to cure poison).

Industrial and chemical milestones include the 1777 recognition of antimony’s distinct element status, the 19th-century development of Sb-based alloys (type metal, pewter), and the 20th-century expansion of Sb₂O₃ fire retardants. In the early 20th century, antimony alloys were added to lead-acid batteries and cable sheathing. During World War II, antimony bromide and oxide found use in pilotless drone fuels and flamethrower agents. In recent decades, safety regulations have been driven by antimony’s toxicity (recent classification as a probable carcinogen) and by its critical role: modern high-speed electronics and Li-ion batteries have created new demand. Etymologically and historically, antimony exemplifies an “ancient element” that found roles from cosmetics to semiconductors, with a legacy of alchemy and modern industry interwoven.

Data Table

Property	Value (and Units)
Symbol	Sb
Atomic Number (Z)	51
Element Category	Metalloid (pnictogen, Group 15, p-block)
Standard Atomic Weight	121.76 (1)
Phase at STP	Solid (gray metallic)
Crystal Structure	Rhombohedral (A7 structure)
Density	~6.7 g·cm⁻³ at 20 °C
Melting Point	~904 K (631 °C)
Boiling Point	~1900–2000 K (~1620–1720 °C)
Covalent Radius (empirical)	~139 pm
Electronegativity (Pauling scale)	2.05
First Ionization Energy	8.64 eV (≈834 kJ/mol)
Electron Configuration	[Kr] 4d¹⁰ 5s² 5p³
Oxidation States	–3, +3, +5 (others possible but rare)
Stable Isotopes (natural abundances)	^121Sb (57.3%, spin 5/2) and ^123Sb (42.7%, spin 7/2)
Notable Compounds	Sb₂O₃ (white oxide), Sb₂O₅ (orange oxide), SbCl₃, SbCl₅, Sb₂S₃ (stibnite), SbH₃ (stibine)

(1) Supported by IUPAC-certified atomic mass. (Masses of individual isotopes: ^121Sb ≈120.9038, ^123Sb ≈122.9042.)