Sulfur

Typical wastewaters

• leather tanning sulfide (HS⁻ and H₂S) from lime and sulfide dehairing liquors, with chromium; BAT-AEL below 1 mg/L for direct and indirect discharge the US pretreatment standard (40 CFR 425.15) allows 24 mg/L daily maximum to sewer
• textile dyeing with sulfur dyes sulfide easily released, below 1 mg/L BAT-AEL for direct and indirect discharge ZDHC also limits sulfite in textile effluent (2, 0.5, 0.2 mg/L by level)
• wool scouring sulfide, mass limited per 1000 kg of wool with phenol and chromium
• mine drainage and industrial brines sulfate from pyrite oxidation, capped by gypsum at 1.5 to 2 g/L after lime treatment; barium reaches lower
• municipal sewage (septic sewers) sulfate reduced to sulfide in the slime layer; H₂S strips to the headspace and is oxidised to sulfuric acid on the sewer crown
• dechlorinated effluent and boiler blowdown sulfite from sulfur dioxide or bisulfite dechlorination and from oxygen scavenging, oxidised to sulfate by dissolved oxygen
• reverse osmosis concentrate sulfate concentrated past gypsum and barite solubility; recovery limited by scaling

1 · Identity

Symbol, number

S, 16

Oxidation states in water

+6 as sulfate SO₄²⁻ (and HSO₄⁻ below pH 2), the stable form in oxic water; +4 as sulfite SO₃²⁻ and bisulfite HSO₃⁻ (dechlorination and oxygen scavenging reagents, unstable in air); -2 as hydrogen sulfide H₂S, bisulfide HS⁻ and metal sulfides in anoxic water; 0 as elemental sulfur from partial oxidation of sulfide; +2 as thiosulfate S₂O₃²⁻ (dechlorination reagent). Organic sulfur in proteins and surfactants.

Note

The element entry covers brimstone, sour gas and the acid industry. This chapter is the redox ladder in water: sulfate down to sulfide in the absence of oxygen, sulfide back to sulfate (by way of sulfuric acid in a sewer crown) when air returns, and the sulfite step that treatment plants buy by the tonne.

2 · Occurrence in water

Natural sources

Dissolution of gypsum, anhydrite and other sulfate minerals and oxidation of pyrite give sulfate; the highest sulfate is in groundwater from natural sources (WHO). Sulfide comes from microbial reduction of sulfate in anaerobic groundwater, sediments and sulfur springs, and from volcanic and geothermal sources.

Anthropogenic sources

Industrial wastes and atmospheric deposition of sulfur dioxide (acid rain); sulfate added by alum and ferric sulfate coagulants and by sulfuric acid dosing (treated water in Ontario carried 22.5 mg/L sulfate against 12.5 mg/L raw); mine drainage from pyrite oxidation; tannery (sulfide dehairing), sulfur dye, pulp and petroleum wastes carry sulfide; dechlorination and boiler blowdown carry sulfite and sulfate.

3 · Speciation

In oxic water sulfur is sulfate, a conservative anion that pairs weakly with calcium and magnesium and precipitates only as gypsum above about 1.5 to 2 g/L or as barium sulfate. Where oxygen runs out and organic carbon is present, sulfate reducing bacteria turn it to sulfide; at pH 7.4 about one third of that sulfide is dissolved H₂S gas and the rest HS⁻, and the odour, toxicity and corrosivity follow the H₂S fraction, which rises as pH falls (pKa₁ about 7.0). Free sulfide is capped by iron: FeS forms wherever ferrous iron is available. Air or chlorine oxidises sulfide back through elemental sulfur to sulfate, and sulfur oxidising bacteria on a sewer crown make sulfuric acid out of it. Sulfite, dosed as sulfur dioxide, bisulfite or metabisulfite, is oxidised to sulfate by chlorine or dissolved oxygen within seconds to minutes.

Solubility

Sulfate is capped by gypsum (CaSO₄.2H₂O), about 2 g/L as the salt in pure water at 25 C, which sets the sulfate ceiling of lime treated acid mine drainage and the scaling limit of reverse osmosis concentrate (Stumm and Morgan chapter 7, from the chapter, not re-read); barium sulfate is far less soluble. Hydrogen sulfide is a soluble gas that strips readily. Metal sulfides (FeS, CuS, ZnS, NiS, PbS, HgS) are among the least soluble solids in water, which is the basis of sulfide precipitation.

Hydrolysis

H₂S dissociates with pKa₁ near 7.0 at 25 C, so the toxic, odorous, volatile molecule dominates below pH 7 and HS⁻ above it; the second dissociation lies above pH 12, so S²⁻ is negligible in treatment (WHO gives the one third H₂S at pH 7.4 figure). Sulfurous acid: SO₂ hydrates to HSO₃⁻ and SO₃²⁻ with pKa values near 1.9 and 7.2 (Stumm and Morgan chapter 3, from the chapter, not re-read).

Complexation

Sulfate forms weak ion pairs with Ca²⁺, Mg²⁺ and Fe²⁺ and complexes Al³⁺ and Fe³⁺ in acid water; bisulfide complexes copper, mercury and other soft metals, which is why sulfide precipitates can carry dissolved polysulfide metal at very high sulfide.

Precipitates

Gypsum and anhydrite, barite BaSO₄, ettringite in lime treated sulfate water, jarosite and schwertmannite in acid mine drainage; FeS, CuS, ZnS, NiS, PbS in sulfidic water and in sulfide precipitation; elemental sulfur.

4 · Role in treatment

5 · Removal and control

6 · Analytics

Sampling pitfalls

Sulfide is lost from the bottle within minutes by volatilisation and oxidation: fill without headspace and fix with zinc acetate and sodium hydroxide at the tap (Standard Methods 4500-S₂ sample handling, not re-read). Sulfite oxidises in air; titrate on site. Sulfate is stable but a chlorinated or bisulfite dosed sample keeps reacting, so dechlorinated effluent should be sampled after the reaction is complete.

7 · Regulatory limits

Limits change, and many are set locally. Treat these as the published values to start from, not as your compliance target: check the standard in force at your site and the numbers written into your own permit.

8 · Health and environmental effects

Toxicity

Sulfate: no level causing adverse effects identified; laxative at 1000 to 1200 mg/L without diarrhoea, dehydration or weight loss; average daily intake about 500 mg mostly from food (WHO). Hydrogen sulfide: high acute inhalation toxicity, eye irritation at 15 to 30 mg/m₃ in air, but a harmful oral dose from drinking water is not credible because taste and odour give it away (WHO). Sulfite is a food additive; no drinking water value.

Bioaccumulation

Not applicable; sulfur is an essential element taken up as sulfate by plants and algae (the element entry).

Ecotoxicity

US EPA aquatic life criterion for sulfide as hydrogen sulfide: chronic 2.0 µg/L in fresh and salt water (1986), no acute value; sulfate has no criterion but high sulfate depresses osmotic tolerance in freshwater biota (not quantified in the sources read).

Flags

• The H₂S pKa₁ (about 7.0), the sulfurous acid pKa values, the gypsum solubility and the sulfate reduction stoichiometry are cited to Stumm and Morgan chapters from memory, not re-read; WHO supports the pKa with its one third H₂S at pH 7.4 statement.
• The crown corrosion mechanism is cited to Metcalf and Eddy chapter 2 without re-reading.
• The chlorine demand of 8.3 mg Cl₂ per mg H₂S is computed from the stoichiometry written here; the EPA manual states that chlorine oxidises sulfides without printing the dose.
• The chromate reduction by sulfite and the nickel sulfide precipitation are electron and charge balances; the CWW BREF names the techniques without printing equations.
• The sulfate occurrence figures are compilations of the 1980s quoted by WHO in 2004.
• The wool scouring limit is mass based (kg per 1000 kg) and cannot be compared with concentration limits.
• Abu Dhabi values cover two media (marine sulfide 0.1 mg/L; sewer sulfate 1000 and sulfide 1 mg/L).
• Sample preservation for sulfide (zinc acetate) is standard practice cited from memory of Standard Methods 4500-S₂.

Gaps

• No source read gives sulfide or sulfate in municipal wastewater, tannery, textile or pulp effluent as numbers; only the limits are quoted.
• Gypsum and barite scaling limits for reverse osmosis and antiscalant doses were not read.
• Sewer sulfide control doses (oxygen, nitrate, iron, peroxide) were not read; the EPA odour and corrosion design manual could not be opened.
• Detection limits for the Standard Methods sulfate, sulfide and sulfite methods were not read.
• Biogas hydrogen sulfide figures were not sourced.
• Sulfate criteria for aquatic life (state or Canadian) were not read.
• Other GCC discharge standards (Saudi, Oman, Qatar) were not read.

Sources

Identity

Name and symbol

Sulfur, S

Atomic number

16 protons

Position

group 16 · period 3 · p-block · polyatomic nonmetal

CAS number

7704-34-9

Atomic structure

Atomic mass

32.066 u

Electron configuration

1s² 2s² 2p⁶ 3s² 3p⁴
[Ne] 3s²³p⁴

Electrons per shell

2, 8, 6

Valence electrons

6 outer shell

Physical properties

State at room temperature

Solid

Melting point

388.36 K (115.21 °C)

Boiling point

717.75 K (444.6 °C)

Density

2.067 g/cm3

Appearance

Lemon yellow sintered microcrystals

Thermal conductivity

0.205 W/(m·K)

Electrical resistivity

2×10¹⁵ Ω·m at 20 °C

Electrical conductivity

5.00e-16 S/m

Crystal structure

orthorhombic

Molar heat capacity

22.75 J/(mol·K)

Chemical properties

Oxidation states

+6, +4, -2

Electronegativity

2.58 (Pauling Scale)

Ionisation energy

10.36 eV
1st 999.6, 2nd 2,252, 3rd 3,357 kJ/mol

Electron affinity

2.077 eV

Atomic radius

empirical 105, covalent 105, van der Waals 180 pm

Ionic radius

S²⁻ 184; S⁴⁺ 37; S⁶⁺ 29 pm

Reactivity

A chalcogen with oxidation states from -2 to +6 that combines with nearly every element except the noble gases, oxidising less electronegative partners to sulfides and being oxidised itself by oxygen and the halogens; many of its reactions need heat.

with water

Insoluble in water and does not react with it under ordinary conditions beyond an extremely slow hydrolysis to hydrogen sulfide and sulfuric acid.

with oxygen, air

Stable in air at room temperature; burns with a blue flame to sulfur dioxide, and the trioxide forms only at 400 to 600 C over a catalyst: $\ce{S + O2 -> SO2}$

with acids

Not attacked by dilute non-oxidising acids; concentrated nitric acid oxidises it to sulfate with evolution of nitrogen dioxide: $\ce{S + 6 HNO3 -> H2SO4 + 6 NO2 + 2 H2O}$

with halogens

Reacts vigorously in liquid fluorine to give sulfur(IV) and (VI) fluorides (SF4, SF6); chlorine gives the low-valent chlorides S2Cl2 and SCl2, bromine the dibromide, iodine no stable binary compound: $\ce{S + 3 F2 -> SF6}$

Typical compounds

• H₂SO₄ sulfuric acid the most important manufactured chemical
• SO₂ sulfur dioxide from burning sulfur or sulfide ores; air pollutant
• H₂S hydrogen sulfide rotten-egg gas of natural gas; deadly at high concentration
• FeS₂ iron pyrite fool's gold, the commonest sulfide mineral
• CaSO₄.2H₂O gypsum calcium sulfate dihydrate; plaster
• CS₂ carbon disulfide volatile solvent; rayon and cellophane manufacture

Occurrence, production and use

Crustal abundance

3.50×102 milligrams per kilogram

Oceanic abundance

9.05×102 milligrams per liter

Occurrence and sources

Sulfur is found in meteorites. R.W. Wood suggests that the dark area near the crater Aristarchus is a sulfur deposit.

Sulfur occurs native in the vicinity of volcanos and hot springs. It is widely distributed in nature as iron pyrites, galena, sphalerite, cinnabar, stibnite, gypsum, epsom salts, celestite, barite, etc.

• dissolved in seawater about 905 mg/L; crustal estimate 350 mg/kg (Jefferson Lab figures via PubChem)
• sulfur in natural gas, petroleum and tar sands sour gas fields and refineries worldwide; recovered at petroleum refineries, gas plants and coking plants; Louisiana and Texas gave 52 percent of US output in 2024
• metal sulfides: pyrite, galena, sphalerite, cinnabar, stibnite ore deposits; by-product sulfuric acid from nonferrous smelters was about 8 percent of US sulfur in 2024
• native sulfur; gypsum and anhydrite (CaSO4) volcanic areas and evaporites; gypsum is mined at about 100 million tonnes a year for cement and plaster

Extraction, production

Recovery of elemental sulfur from sour natural gas and crude oil

$\ce{4 H2S (g) + 2 SO2 (g) -> 3 S2 (g) + 4 H2O (g)}$

Desulfurisation of fuels is now almost the whole supply; formerly the sulfur was destroyed chemically, and large amounts are recovered from Alberta gas fields (RSC, LANL). US elemental sulfur production was 8.2 million tonnes in 2024 (USGS, printed pp. 172 to 173). No equation is printed in the sources.

Contact process for sulfuric acid: sulfur burning or sulfide roasting, catalytic oxidation, absorption

SO₂ + 1/2 $\ce{O2 <=> SO3}$

Sulfur dioxide from combustion of elemental sulfur or roasting of metal sulfides is converted on a catalyst in a gas-phase equilibrium, delta H0 = -99 kJ/mol; conversion rate is the key performance measure (AAF BREF, PDF p179). The trioxide is then absorbed, $\ce{SO3 + H2O -> H2SO4 (SIC BREF, PDF}$ p178). The acid is the most important manufactured chemical and its largest use is phosphoric acid.

Frasch mining (historical)

Superheated water melted the sulfur in salt domes for pumping.

Uses

Most of the sulfur that is produced is used in the manufacture of sulfuric acid (H2SO4). Large amounts of sulfuric acid, nearly 40 million tons, are used each year to make fertilizers, lead-acid batteries, and in many industrial processes. Smaller amounts of sulfur are used to vulcanize natural rubbers, as an insecticide (the Greek poet Homer mentioned "pest-averting sulphur" nearly 2,800 years ago!), in the manufacture of gunpowder and as a dying agent.

In addition to sulfuric acid, sulfur forms other interesting compounds. Hydrogen sulfide (H2S) is a gas that smells like rotten eggs. Sulfur dioxide (SO2), formed by burning sulfur in air, is used as a bleaching agent, solvent, disinfectant and as a refrigerant. When combined with water (H2O), sulfur dioxide forms sulfurous acid (H2SO3), a weak acid that is a major component of acid rain.

Sulfur is a component of black gunpowder, and is used in the vulcanization of natural rubber and a fungicide. It is also used extensively in making phosphatic fertilizers. A tremendous tonnage is used to produce sulfuric acid, the most important manufactured chemical.

It is used to make sulfite paper and other papers, to fumigate, and to bleach dried fruits. The element is a good insulator.

Sulfur is essential to life. It is a minor constituent of fats, body fluids, and skeletal minerals.

• Sulfuric acid and phosphate fertilisers: phosphoric acid, $\ce{Ca3(PO4)2 + 3 H2SO4 -> 2 H3PO4 + 3 CaSO4 (AAF BREF, PDF}$ p242); hydrogen fluoride, $\ce{CaF2 + H2SO4 -> 2 HF + CaSO4 (PDF}$ p286); ammonium sulfate fertiliser about 90 percent of US sulfur is consumed as sulfuric acid; phosphate fertiliser demand sets sulfur demand, with new projects in Africa and west Asia
• Mining and metals: by-product sulfuric acid from nonferrous smelters; sulfuric acid for high-pressure acid leaching of nickel for batteries; sulfur recovered from oil and gas processing
• Rubber and chemicals: vulcanisation of rubber; carbon disulfide, mercaptans as gas odorants, pesticides and herbicides; surfactants and detergents as sulfate derivatives
• Food and beverage: sulfites as preservatives; burning sulfur to preserve wine is ancient; bleaching and fumigating dried fruits
• Pulp and paper: sulfite pulping and sulfite bleaching of paper
• Agriculture and construction: fungicide; gypsum for plaster

Safety, toxicity

Carbon disulfide, hydrogen sulfide, and sulfur dioxide should be handled carefully. Hydrogen sulfide in small concentrations can be metabolized, but in higher concentrations it quickly can cause death by respiratory paralysis.

It quickly deadens the sense of smell. Sulfur dioxide is a dangerous component in atmospheric air pollution.

GHS classification, signal word Danger

• H315 Causes skin irritation Skin corrosion/irritation
• H228 Flammable solid Flammable solids
• H370 Causes damage to organs Specific target organ toxicity, single exposure
• H373 May causes damage to organs through prolonged or repeated exposure Specific target organ toxicity, repeated exposure

Discovery and name

Discovered by

Ancient china

Discovered

before 2000 BCE

First isolated

not in sources

Named by

not in sources

Origin of the name

from the Latin sulpur

Sulfur is pale yellow, odorless, brittle solid, which is insoluble in water but soluble in carbon disulfide. In every state, whether gas, liquid or solid, elemental sulfur occurs in more than one allotropic form or modification; these present a confusing multitude of forms whose relations are not yet fully understood.

In 1975, University of Pennsylvania scientists reported synthesis of polymeric sulfur nitride, which has the properties of a metal, although it contains no metal atoms. The material has unusual optical and electrical properties.

High-purity sulfur is commercially available in purities of 99.999+%.

Amorphous or "plastic" sulfur is obtained by fast cooling of the crystalline form. X-ray studies indicate that amorphous sulfur may have a helical structure with eight atoms per spiral. Crystalline sulfur seems to be made of rings, each containing eight sulfur atoms, which fit together to give a normal X-ray pattern.