Sulfur
Sulfur or sulphur (see spelling differences) is a chemical element with symbol S and atomic number 16. It is an abundant, multivalent non-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S8.
Identity
- Atomic mass
- [32.059, 32.076] u
- Category
- Reactive non-metals
- Group
- 16
- Period
- 3
- Block
- p
- Oxidation States
- −2, +2, +4, +6
- Phase (25 °C)
- Solid
- Appearance
- lemon yellow sintered microcrystals
Electronic Structure
- Electron configuration
- 1s2 2s2 2p6 3s2 3p4
- Short configuration
- [Ne] 3s2 3p4
- Electrons per shell
- 2, 8, 6
- Electronegativity (Pauling)
- 2.58
- Electron affinity
- 200 kJ/mol
- 1st ionization energy
- 1.00e+3 kJ/mol
- Ionization energies
- 1000 · 2252 · 3357 · 4556 · 7004… kJ/mol
Physical properties
- Melting point
- 115.2 °C (388 K)
- Boiling point
- 444.6 °C (718 K)
- Density
- 2.07 g/cm³
- Molar heat capacity
- 22.8 J/(mol·K)
Chemical behavior of sulfur
Sulfur has six valence electrons ([Ne] 3s² 3p⁴) and is the non-metal of the two extremes: it is reduced to -2 in sulfides and hydrogen sulfide, and oxidised up to +6 in Sulfuric acid, passing through +2 and +4 in sulfurous. Its medium-high electronegativity (2.58) gives it versatile covalent bonds, including disulfide bridges that fold proteins. Like oxygen, it concatenates: its S₈ rings and chains explain its colours, odours and its volcanic role.
Reactivity
S + O₂ → SO₂
It burns in air with a blue flame giving sulphur dioxide, the gas of volcanoes and acid rain. It is the first industrial step towards sulfuric acid.
H₂ + S → H₂S
With hydrogen it forms hydrogen sulfide, the gas of rotten eggs: flammable and as toxic as cyanide because it blocks cellular respiration. Volcanoes and the intestine emit it.
2 SO₂ + O₂ → 2 SO₃
The dioxide is oxidised to trioxide with vanadium catalyst (contact process): the key step of sulfuric acid. Without this oxidation there would be no modern fertilisers.
Fe + S → FeS
Heated with metals it forms sulfides with incandescence: pyrite (FeS₂, fool's gold) and galena are classic ores. Many sulfides are semiconductors and pigments.
Notable compounds
| Compound | Formula | Use or importance |
|---|---|---|
| Sulfuric acid | H₂SO₄ | Fertilisers, car batteries and refining: the most produced chemical |
| Hydrogen sulfide | H₂S | Volcanic gas, cellular signal and industrial toxic |
| Sulphur dioxide | SO₂ | Wine preservative and precursor of sulfuric acid |
| Pyrite | FeS₂ | Fool's gold and historical sulphur ore |
| Calcium sulfate (gypsum) | CaSO₄·2H₂O | Scales, cement and bricklaying |
Electron shell distribution
View in the atom builderIsotopes
Natural sulfur mixes four stable isotopes: sulfur-32 (³²S, ~95%), sulfur-33 (~0.75%), sulfur-34 (~4.2%) and sulfur-36 (~0.01%), all with sixteen protons. Its proportions reveal the volcanic or biological origin of a sample and reconstruct ancient climate.
Occurrence and production
It is the fifth industrial element by volume: it appears native alongside volcanoes and fumaroles, and as sulfides (Pyrite, galena, blende), galena, blende) and sulfates (gypsum, anhydrite). Today it is obtained above all by desulfurising oil and gas (Claus process): yellow mountains next to each refinery.
History and origin of the name
Known since Antiquity as brimstone (burning stone): Egyptians used it as a fumigant and Homer cites its incense. Antoine Lavoisier consecrated it as an element in 1789. Its smell of hell made it a symbol of the devil in the Bible; its sulfuric acid, in turn, became a measure of a country's industrial power.
Applications and relevance of sulfur
Uses of the element
It vulcanises rubber (sulfur bridges that make it elastic), acts as a fungicide in vineyards and bleaches sugar. Its thermal baths have relieved skin since Rome. Without sulfur our tyres would be rubber.
Uses of its compounds
Sulfuric acid manufactures fertilisers, car batteries and detergents: its consumption measures a country's industry. SO₂ preserves wine; gypsum raises load-bearing walls; H₂S, in microdoses, regulates blood pressure as a cellular messenger.
Health, safety and environment
Biological role
It is essential: the amino acids cysteine and methionine and B vitamins carry it, and disulfide bridges give shape to hair and proteins. Its lack slows growth; endogenous H₂S, in traces, protects vessels and neurons.
Safety
H₂S kills at 500 ppm by first numbing the sense of smell: victims stop smelling it before collapsing, so detectors are required in refineries and treatment plants. SO₂ irritates eyes and bronchi; sulfuric acid burns and dehydrates tissues instantly.
Environment
SO₂ from power plants and volcanoes causes acid rain that devastated Central European forests in the 80s; conventions drastically reduced it. Sulphate aerosols, moreover, cool the climate by reflecting the sun: big eruptions lower the global temperature.
Frequently asked questions about Sulfur
Why does sulphur smell like rotten eggs?
By hydrogen sulphide (H₂S), which eggs, volcanoes and intestines release by decomposing sulphur-rich proteins. We smell it at parts per billion because our ancestors learned to flee from rotten and dangerous things. Ironically, in microdoses our body makes it as a vascular messenger. Poison is in the dose: the smelly warning is actually always an evolutionary favour.
Why does sulphuric acid measure a country's industry?
Because it participates in fertilisers, batteries, refining, detergents and mining: there is no chemical sector without it and 270 million tonnes are produced each year. Economists use it as a reliable industrial thermometer since the 19th century. Where its consumption rises, industry pulls strongly. It is the most manufactured chemical on the planet: bread, car and mobile phone owe something to sulphuric acid.
Why do volcanoes smell of sulphur?
Because they degas SO₂ and H₂S from magma: the smell of rotten eggs is the unmistakable signature of the Earth breathing. Fumaroles deposit native yellow sulfur that has been mined since Antiquity. That same SO₂ forms aerosols that cool the planet after major eruptions. Smelling sulfur is smelling active geology: hell smells like that because volcanoes are their eternal chimneys.
Does acid rain still exist?
Much less: SO₂ from coal power plants devastated forests and lakes in the 70s and 80s, but conventions, desulphurisation and natural gas reduced it by more than 80% in Europe. Today the problem has shifted to Asia and to ships. The lesson remains: acid rain was cured with policy and technology. What sulfur soiled, regulation cleaned.
What is the 'fool's gold'?
Pyrite (FeS₂), a bright golden sulphide of iron that ruined speculators who believed it to be gold. It is distinguished because it is harder, lighter and leaves a greenish mark. The Romans used it to start fire with a hook. Today it is a sulphur ore and sometimes hides real microscopic gold. Its shine deceived many, but its chemistry is honest: iron and sulphur.
Is wine's sulphur causing headaches?
Probably not: sulphites (SO₂) preserve wine from Rome, preventing it from turning to vinegar, in regulated doses and harmless to almost everyone. Only the most sensitive asthmatics note real discomfort. The headache of red wine comes more from alcohol and amines. 'Sulphite-free' wines spoil much earlier: sulfur is the discreet guardian of each cellar.
What does sulphur do to hair?
It gives it shape and memory: keratin is sewed by disulfide bridges (S–S) and breaking them with heat or thioglycolates allows permanently straightening or curling. The burnt smell when ironing is sulphur released into the air. 'Keratin' treatments restore those broken bridges with new proteins. Each curl and each chemical straightening is pure engineering of sulphur bonds.
Main sources consulted: PubChem, Royal Society of Chemistry, NIST, IUPAC. Last revision: November 2026.