Sulfur – facts and properties
- Symbol
- S
- Atomic number
- 16
- Atomic mass
- 32.06 u
- Period
- 3
- Group
- 16
- Category
- polyatomic nonmetal
- Electron configuration
- [Ne] 3s² 3p⁴
- Electronegativity
- 2.58
- Atomic radius
- 88 pm
- Valence electrons
- 6
- State at standard conditions
- solid
- Melting point
- 115.21 °C
- Boiling point
- 444.6 °C
- Density
- 2.067 g/cm³
- Discovered
- known since antiquity
Sulfur is a brittle yellow non-metal with atomic number 16. It is one of the few elements found uncombined in nature, so it has been known since antiquity. Its reputation is mixed: burning sulfur gives off a choking smell, hydrogen sulfide reeks of rotten eggs, and older texts link brimstone with fire and hell. Yet sulfur is a cornerstone of industry. It is turned into sulfuric acid, one of the most widely produced chemicals in the world, and sulfur-containing amino acids are found in almost every protein.
Position in the periodic table
Occurrence and extraction
Native sulfur occurs mainly around volcanoes, where sulfurous gases leave yellow crystalline crusts, and in deposits within sedimentary rocks. Far more sulfur is bound up in compounds: in sulfide ores such as pyrite, better known as fool's gold, galena and sphalerite, and in sulfates such as gypsum. Crude oil, natural gas and coal also contain sulfur compounds.
Sulfur used to be dug from mines in Sicily. Later the Frasch process melted it out of underground deposits in Texas and Louisiana with superheated water and forced it to the surface. Today most sulfur is a by-product of cleaning up natural gas and crude oil. The hydrogen sulfide removed from these fuels is partly burned in the Claus process and converted into elemental sulfur, so a pollutant that would otherwise end up in the air as sulfur dioxide becomes a raw material. Nobody can be credited with discovering sulfur. Antoine Lavoisier included it in his list of chemical elements in the late eighteenth century.
Properties and reactions
Sulfur has the electron configuration [Ne] 3s² 3p⁴ and six valence electrons. It sits in period 3 and group 16, directly below oxygen, and has an electronegativity of 2.58. Its most important oxidation states are minus two in sulfides and hydrogen sulfide, plus four in sulfur dioxide and sulfites, and plus six in sulfuric acid and sulfates.
At room temperature sulfur is made of crown-shaped rings of eight atoms, which pack into yellow rhombic crystals. Other allotropes exist too, such as needle-shaped monoclinic sulfur. Sulfur has a density of 2.067 g/cm³, is a poor conductor of heat and electricity and does not dissolve in water. It melts at 115.21 °C to a runny pale yellow liquid. Heated further, the melt turns reddish brown and thick as the rings break open and join into long chains. Pouring the hot liquid into cold water gives rubbery plastic sulfur. Sulfur boils at 444.6 °C.
In air, sulfur burns with a blue flame to form sulfur dioxide, a colourless gas with a sharp, choking smell. It irritates the eyes and lungs, is toxic in higher concentrations and contributes to acid rain. With hydrogen, sulfur forms hydrogen sulfide, a highly toxic gas that smells of rotten eggs. The danger is that at higher concentrations it numbs the sense of smell. Heated with many metals, sulfur forms sulfides, as in the classic school experiment with iron filings.
Uses in everyday life and technology
By far the largest share of sulfur is made into sulfuric acid. In the contact process, sulfur is burned to sulfur dioxide, which is oxidised to sulfur trioxide over a vanadium pentoxide catalyst and then absorbed into sulfuric acid. Most of it goes into fertilisers made from phosphate rock. It is also needed for lead-acid car batteries, dyes, detergents and metal refining.
Vulcanisation means heating natural rubber with sulfur. Sulfur bridges link the long chain molecules together and turn a sticky mass into the tough, elastic rubber used in tyres. Gunpowder is a mixture of saltpetre, charcoal and sulfur, and match heads contain sulfur as well. Winegrowers dust their vines with sulfur against fungal diseases, and wine itself is treated with sulfur dioxide or sulfites to protect it from oxidation and unwanted microbes, which is why wine labels mention sulfites.
Living things cannot do without sulfur. The amino acids cysteine and methionine contain it, and disulfide bridges between cysteine units hold many proteins in shape. Hair, nails and feathers get their strength from keratin, a protein rich in these bridges. The pungent smell of garlic and onions also comes from sulfur compounds.
Molecules containing this element
- Hydrogen sulfide H2S
- Barium sulfide BaS
- Cadmium sulfide CdS
Quick question
Which pair of elements has the same number of valence electrons?
FAQ
Is sulfur poisonous?
Elemental sulfur has very low toxicity. It is even used as a fungicide and in ointments for skin conditions. Some of its compounds are dangerous, though. Sulfur dioxide from burning sulfur irritates the airways, and hydrogen sulfide is highly toxic even at low concentrations. Fine sulfur dust can also form explosive mixtures with air.
What causes the rotten egg smell?
The smell comes from hydrogen sulfide. Bacteria release it when they break down sulfur-containing proteins without oxygen, for example in rotten eggs, in mud at the bottom of ponds or in sewers. Pure sulfur itself is almost odourless.
What is sulfur used for?
Mainly for making sulfuric acid, which goes into fertilisers, batteries and many other chemicals. Sulfur also vulcanises rubber, is found in gunpowder and matches, and protects grapevines from fungi.