Chemistry 2e · Representative Metals, Metalloids, and Nonmetals

Occurrence, Preparation, and Properties of Sulfur

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

Sulfur is the yellow element of brimstone — found free in volcanic regions, combined in ores, and dissolved in the smelly gas hydrogen sulfide. Its chemistry is built on a simple pattern: sulfur atoms form rings and chains (the S₈ ring is the stable form), and the element adopts oxidation states from -2 (sulfides) through 0 (elemental sulfur) and +4 (SO₂) to +6 (H₂SO₄). This range makes sulfur both the source of some of the foulest-smelling and most toxic compounds in chemistry (H₂S) and the foundation of the most important industrial chemical in the world, sulfuric acid.

Elemental sulfur is extracted either by mining native deposits or, historically, by the , which melted underground sulfur with superheated water and pumped it to the surface. In the laboratory and in nature, sulfur shows dramatic allotropy: rhombic sulfur is the stable form below 95 °C, monoclinic sulfur is stable between 95 °C and the melting point, and plastic sulfur forms when molten sulfur is quenched. When heated, sulfur's physical behavior is itself a classic demo — it melts to a mobile yellow liquid, then turns viscous and dark red as chains form, then thins again near its boiling point.

Why this matters

Sulfuric acid (H₂SO₄) is produced in larger quantities than any other industrial chemical — roughly 250 million tonnes per year worldwide — because it is essential to fertilizer production (converting phosphate rock into usable phosphate), petroleum refining, metal processing, and battery acid (lead–acid car batteries are about 35% H₂SO₄). Sulfur dioxide (SO₂) is both an industrial feedstock and a pollutant: it forms acid rain, which has damaged forests and lakes and corroded buildings, and it is also used (as sulfites) to preserve dried fruit and wine. Hydrogen sulfide is a deadly hazard in oil fields, sewers, and volcanic areas because it paralyzes the sense of smell at dangerous concentrations. In biology, sulfur is part of the amino acids cysteine and methionine, and the disulfide bonds it forms hold protein structure together. Clinically, sulfur compounds appear in drugs from antibiotics (sulfa drugs) to topical acne treatments.

The college version

Core Concepts

Occurrence: native sulfur, sulfides, and sulfates

Sulfur is the 10th most abundant element in the universe and occurs in three forms in nature: (1) native (elemental) sulfur, found around volcanoes and in salt domes (notably along the Gulf of Mexico coast), often formed by the reaction of H₂S with SO₂ from volcanic gases; (2) sulfide ores, such as pyrite (FeS₂, "fool's gold"), galena (PbS), and sphalerite (ZnS), which are the raw material of most metal smelting; and (3) sulfates, such as gypsum (CaSO₄ · 2H₂O) and epsom salts (MgSO₄ · 7H₂O). Hydrogen sulfide, H₂S, is common in natural gas, volcanic gases, and the anaerobic decay of organic matter (the "rotten egg" smell of swamps).

Preparation: mining, the Frasch process, and the Claus process

Native sulfur is mined directly where deposits are shallow. Deep deposits were historically recovered with the Frasch process: three concentric pipes are driven into the deposit; superheated water (~160 °C) melts the sulfur (mp 115 °C), hot compressed air forces the molten sulfur up the middle pipe as a froth, and it is collected and cooled into solid blocks. This process is no longer the major source — most sulfur now comes as a byproduct of petroleum and natural gas refining, where the recovers elemental sulfur from H₂S:

$2H 2 S + 3O 2 \to 2SO 2 + 2H 2 O then 2H 2 S + SO 2 \to 3S + 2H 2 O$

The overall result converts toxic, smelly H₂S into useful solid sulfur — an environmental and economic win.

Allotropes and physical behavior

Sulfur's most stable is rhombic sulfur, built from crown-shaped S₈ rings; it is the pale yellow solid stable below 95.3 °C. Monoclinic sulfur (also S₈ rings, packed differently) is stable from 95.3 °C to the melting point (about 115 °C); heating rhombic sulfur above 95.3 °C converts it to monoclinic. When molten sulfur is heated further, the S₈ rings open and join into long chains, making the liquid dark red and so viscous it will not pour — and if this hot liquid is dumped into cold water, the chains are trapped as rubbery, amorphous plastic sulfur, which slowly reverts to rhombic sulfur. Allotropes of the same element, different properties.

The SO₂ → SO₃ → H₂SO₄ pathway: the contact process

Burning sulfur or roasting sulfide ores produces sulfur dioxide:

$S + O 2 \to SO 2$

SO₂ is a colorless, pungent, toxic gas. The oxidizes it to sulfur trioxide over a vanadium(V) oxide catalyst:

$2SO 2 + O 2 ⇌ 2SO 3$

The reaction is exothermic and decreases gas moles, so moderate temperature (~400–450 °C) and high pressure favor a good yield; the SO₃ is then dissolved in concentrated H₂SO₄ to form oleum, H₂S₂O₇, which is diluted with water to make more sulfuric acid (adding water directly to SO₃ would create a corrosive mist). Sulfuric acid is a strong : it ionizes in two steps, H₂SO₄ → H⁺ + HSO₄^- and HSO₄^- ⇌ H⁺ + SO₄^2-. Concentrated H₂SO₄ is also a powerful — it chars sugar and paper by pulling water out of them — and hot concentrated sulfuric acid acts as an oxidizing agent.

Sulfides, oxidation states, and the smells

Sulfur's oxidation states in common compounds run from -2 (sulfides like Na₂S and H₂S, and metal sulfides) through +4 (SO₂, sulfites) to +6 (SO₃, sulfates, H₂SO₄). Hydrogen sulfide is notorious: it smells like rotten eggs at very low concentrations, but at higher concentrations it deadens the sense of smell, so "if you can smell it, it's probably not yet lethal — but don't trust your nose" is the safety maxim; H₂S is comparable in toxicity to hydrogen cyanide and is heavier than air, so it collects in low-lying, poorly ventilated spaces.

How It Works / Step-by-Step Process

From smelly gas to useful acid: (1) Natural gas and petroleum containing H₂S are processed, and the H₂S is stripped out; (2) in the Claus process, part of the H₂S is burned to SO₂, then the two gases react (2H₂S + SO₂ → 3S + 2H₂O) to deposit molten sulfur; (3) the sulfur is burned to SO₂; (4) SO₂ is oxidized over a vanadium catalyst to SO₃ and absorbed into concentrated sulfuric acid to form oleum; (5) oleum is diluted to the desired acid strength. Every stage involves corrosive, toxic, or hot materials, so the process runs in sealed, ventilated industrial plants — never as a bench-top experiment.

Common Confusions

Do Not Confuse	With	Difference
Rhombic and monoclinic sulfur	Different elements	Both are S₈ ring allotropes of sulfur, stable in different temperature ranges (below/above 95.3 °C).
SO₂ and SO₃	The same oxide	SO₂ is produced by burning sulfur (S at +4); SO₃ is its further oxidation (S at +6) and is the direct precursor to H₂SO₄.
Sulfuric acid being "just" an acid	Concentrated H₂SO₄ as dehydrating/oxidizing agent	Dilute H₂SO₄ is an acid; concentrated acid chars organics (dehydration) and hot concentrated acid oxidizes metals.
Sulfide and sulfate	The same sulfur compound	Sulfide is S at -2 (e.g., FeS₂); sulfate is S at +6 (e.g., CaSO₄). Very different chemistry and hazards.
H₂S smell as a warning	Smell as a reliable safety signal	High H₂S concentrations paralyze the sense of smell; absence of odor does not mean absence of danger.
Sulfur "burns with a blue flame"	Sulfur burning like wood	Sulfur burns with a characteristic blue flame producing SO₂ — the gas, not the flame, is the hazard.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Sulfur is the yellow powder that smells like rotten eggs when it's in the gas hydrogen sulfide. It can take many shapes — like a block of yellow crystals, long stretchy chains, or a dark red goo — depending on how hot it is. Burn it, and it becomes a choking gas that turns into acid rain, but chemists capture that gas and turn it into sulfuric acid, the chemical used to make fertilizer, car batteries, and to refine oil.

Worked example

Worked Example 1 — Sulfur dioxide from burning sulfur. How many grams of SO₂ form when 16.0 g of sulfur burns completely? The reaction is S + O₂ → SO₂. Molar masses: S = 32.07 g/mol; SO₂ = 32.07 + 2(16.00) = 64.07 g/mol. Conversion path first, then substitute:

$g SO 2 = 16.0 g S \times 1 mol S 32.07 g S \times 1 mol SO 2 1 mol S \times 64.07 g SO 2 1 mol SO 2 = 32.0 g SO 2$

Unit check: g S → mol S → mol SO₂ → g SO₂. Burning 16.0 g of sulfur yields 32.0 g of SO₂ — one mole of sulfur makes one mole of the dioxide, and the mass doubles because each S gains two O atoms.

Worked Example 2 — Sulfuric acid from sulfur. In the contact process, one atom of sulfur ultimately produces one molecule of H₂SO₄ (S → SO₂ → SO₃ → H₂SO₄). What mass of H₂SO₄ can be made from 32.07 g of sulfur? Molar mass of H₂SO₄ = 2(1.008) + 32.07 + 4(16.00) = 98.09 g/mol.

$g H 2 SO 4 = 32.07 g S \times 1 mol S 32.07 g S \times 1 mol H 2 SO 4 1 mol S \times 98.09 g H 2 SO 4 1 mol H 2 SO 4 = 98.1 g H 2 SO 4$

So each gram of sulfur yields about 3.06 g of sulfuric acid (98.09/32.07) — the mass gain comes from the oxygen and water added along the route, which is why acid plants are built next to sulfur sources and air.

Key takeaways

Sulfur occurs native (volcanoes, salt domes), in sulfide ores (FeS₂, PbS, ZnS), and as sulfates (CaSO₄ · 2H₂O).
Frasch process: superheated water + compressed air melt and lift underground sulfur; modern sulfur is mostly a refinery byproduct recovered by the Claus process from H₂S.
Allotropes: rhombic (stable < 95.3 °C), monoclinic (95.3 °C to mp ~115 °C), plastic (quenched molten sulfur, amorphous).
Combustion: S + O₂ → SO₂; contact process: 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, ~400–450 °C) → H₂SO₄.
H₂SO₄ is the world's most-produced industrial chemical; it is a strong diprotic acid and a dehydrating agent when concentrated.
Oxidation states: -2 (sulfides, H₂S), 0 (elemental), +4 (SO₂), +6 (SO₃, H₂SO₄).
H₂S is toxic and deadens smell at high concentrations — never rely on odor for safety.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

Name the three natural forms in which sulfur occurs and give one example mineral for each.
Show answer
Native sulfur (volcanic deposits); sulfide ores (pyrite FeS₂, galena PbS, sphalerite ZnS); sulfates (gypsum CaSO₄ · 2H₂O). H₂S is also common in gas/volcanic emissions.
What is the Frasch process, and why is most sulfur now obtained differently?
Show answer
It melts underground sulfur with superheated water and lifts it with compressed air. Most sulfur now comes as a byproduct of petroleum/natural gas refining via the Claus process.
Why does molten sulfur become dark red and very viscous when heated further?
Show answer
The S₈ rings open and join into long polymer chains, which entangle and raise viscosity; the chains also absorb light differently, giving the dark red color.
Write the two-step Claus process reaction for converting H₂S to elemental sulfur.
Show answer
2H₂S + 3O₂ → 2SO₂ + 2H₂O, then 2H₂S + SO₂ → 3S + 2H₂O.
What mass of SO₂ is produced from 64.1 g of sulfur? (Molar masses: S 32.07, SO₂ 64.07 g/mol.)
Show answer
64.1 g S × (1 mol/32.07 g) × (64.07 g SO₂/1 mol) = 128 g SO₂.
What oxidation states does sulfur show in H₂S, S₈, SO₂, and H₂SO₄?
Show answer
H₂S: -2; S₈: 0; SO₂: +4; H₂SO₄: +6.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

allotrope: Different structural form of the same element
S₈ ring: The crown-shaped eight-atom ring that is sulfur's stable molecular unit
Frasch process: Historic method of melting underground sulfur with superheated water and lifting it with compressed air
Claus process: Converting H₂S from refining into elemental sulfur
contact process: Catalytic oxidation of SO₂ to SO₃, then hydration to H₂SO₄
dehydrating agent: Substance that removes water from other materials
diprotic acid: Acid that can donate two protons
sulfide: Compound of sulfur at oxidation state -2

Sources & references

openstax.org — Chemistry 2e

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