Chemistry: Atoms First 2e · Representative Metals, Metalloids, and Nonmetals

Occurrence, Preparation, and Properties of Sulfur

9 min read

Reference-values note: atomic masses, allotrope transition temperature (95.5 °C), sulfur melting (~115 °C) and boiling (~445 °C) points, and Frasch water temperature (~170 °C) are commonly taught reference values; use the periodic table provided in your course for graded work. Safety note: H₂S is toxic and can deaden the sense of smell; SO₂ irritates the respiratory tract — general ventilation and fume-containment principles only, no lab procedures.

On this page 9 sections

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 (S, atomic number 16) is the yellow, smelly workhorse of the periodic table — a nonmetal that occurs both free and combined, shows more allotropes than any other element, and whose compounds range from the stench of rotten eggs to the world's most important industrial acid. Free sulfur is found as crown-shaped S₈ molecules in volcanic regions and salt-dome deposits, while combined sulfur appears in ores (pyrite, galena, cinnabar) and minerals (gypsum). Sulfur is extracted by the , which melts underground deposits with superheated water and lifts the molten element to the surface with compressed air. Its chemistry rests on oxidation states of −2 (sulfides, H₂S), 0 (elemental S₈), +4 (, sulfites), and +6 (sulfur trioxide, sulfate, sulfuric acid). This topic covers occurrence, preparation, the allotropes, key properties, and the that turns sulfur into sulfuric acid — the most-produced chemical in industry.

Why this matters

Sulfuric acid production is a barometer of industrial output: more H₂SO₄ is made worldwide than any other chemical, and it is used in fertilizers, batteries, petroleum refining, and metal processing. Sulfur also matters for environment and health: burning sulfur-containing coal and oil releases SO₂, the main cause of acid rain; hydrogen sulfide is a toxic, rotten-egg gas that can paralyze the sense of smell; and sulfate aerosols reflect sunlight and cool the atmosphere. In biology, sulfur is essential — the amino acids cysteine and methionine and the disulfide bonds that hold protein structure together all contain it. On exams, sulfur tests allotropy, oxidation-state bookkeeping, and contact-process stoichiometry, plus the acid–base behavior of its oxides.

The college version

Core Concepts

Occurrence: free and combined

Sulfur is about the 16th most abundant element in Earth's crust and, besides carbon, the only common nonmetal found free in large deposits. Native sulfur occurs in volcanic regions and in salt domes (Gulf Coast of the United States), often as crown-shaped S₈ ring molecules in orthorhombic crystals. Combined sulfur is widespread:

Sulfide ores: pyrite (FeS₂, "fool's gold"), galena (PbS), cinnabar (HgS), sphalerite (ZnS).
Sulfate minerals: gypsum (CaSO₄ · 2 H₂O), Epsom salts (MgSO₄ · 7 H₂O).
Organic sulfur: in petroleum, coal, and the amino acids cysteine and methionine.

Preparation: the Frasch process and roasting

Frasch process. Underground sulfur is melted in place by pumping superheated water (about 170 °C) down an outer pipe; molten sulfur (melting point ~115 °C) is forced to the surface through an inner pipe by compressed air. The product is over 99% pure — no chemical purification needed.

From sulfide ores. Roasting sulfides to extract metals also produces sulfur dioxide. Roasting pyrite:

$4 FeS 2 (s) + 11 O 2 (g) Δ ⟶ 2 Fe 2 O 3 (s) + 8 SO 2 (g)$

The SO₂ can be converted back to elemental sulfur (Claus process) or fed into sulfuric acid manufacture.

Allotropes: one element, many structures

Orthorhombic sulfur (S₈ rings): the stable form below 95.5 °C; the pale-yellow crystals found in nature.
Monoclinic sulfur: needle-like crystals of S₈, stable between 95.5 °C and the melting point (~119 °C).
Plastic (amorphous) sulfur: above about 160 °C the S₈ rings open and link into long chains, making the liquid dark and viscous; rapid cooling traps the chains as a rubbery solid that slowly reverts to crystalline sulfur.

Above ~445 °C sulfur vapor contains S₈, S₆, S₄, and S₂ molecules — "molecular sulfur" is a family, not a single species.

Properties and oxidation states

Elemental sulfur is a yellow, brittle solid, insoluble in water but soluble in carbon disulfide, CS₂. It burns in air with a blue flame to give SO₂. Its chemistry is built on four oxidation states:

−2 (sulfides): H₂S is a weak, toxic acid smelling of rotten eggs; metal sulfides like FeS are common minerals.
0: elemental sulfur, S₈.
+4: SO₂, a pungent gas forming sulfurous acid, H₂SO₃, in water; sulfites contain SO₃^2-.
+6: SO₃ and sulfate, SO₄^2-; the highest and most important state, anchored by sulfuric acid.

Because sulfur spans −2 to +6, its compounds are versatile redox partners: H₂S is a reducing agent, SO₂ can be oxidized or reduced, and concentrated H₂SO₄ is a strong oxidizer as well as a strong acid.

The contact process: from sulfur to sulfuric acid

Sulfuric acid is made in three steps:

$S(s) + O 2 (g) \to SO 2 (g)$

$2 SO 2 (g) + O 2 (g) V 2 O 5, \sim 450 ° C ⟶ 2 SO 3 (g)$

$SO 3 (g) + H 2 O(l) \to H 2 SO 4 (aq)$

The middle step is an equilibrium, so it uses a vanadium(V) oxide catalyst and moderate temperature to balance speed against yield (Le Châtelier's principle). Industrially, SO₃ is first dissolved in concentrated H₂SO₄ to make (H₂S₂O₇), then diluted with water — adding SO₃ directly to water is violently exothermic.

Common Confusions

Do Not Confuse	With	Difference
Sulfide S^2-	Sulfite SO₃^2-	Sulfide is a bare −2 ion (S in −2); sulfite is an oxoanion (S in +4)
Sulfite SO₃^2-	Sulfate SO₄^2-	One more oxygen: sulfite (S+4) vs sulfate (S+6) — a classic spelling/charge trap
Sulfur dioxide SO₂	Sulfur trioxide SO₃	SO₂ is a combustion gas; SO₃ is the +6 oxide made catalytically in the contact process
H₂S smell = reliable warning	H₂S at deadly concentrations	H₂S paralyzes the olfactory nerve: at high levels you may smell nothing
Orthorhombic vs monoclinic sulfur	Different elements	Both are S₈ — only crystal packing differs; temperature decides which is stable
Plastic sulfur	A durable polymer	Amorphous sulfur chains revert to crystalline S₈ over time

Eli explains

The same idea, in plain words

Explain it like I’m 10

Sulfur is the yellow element that smells like rotten eggs and can take many shapes, like a toy that comes as rings, needles, or stretchy rubber. We dig it up by pumping hot water underground to melt it, then blow it to the surface like a bubble through a straw. Its best trick is teaming up with oxygen to make sulfuric acid — the acid in car batteries and a factory workhorse used to make almost everything.

Worked example

Example 1: Contact process — how much acid from how much sulfur?

How many grams of H₂SO₄ can be made from 32.1 g of sulfur, assuming complete conversion? (Molar masses: S = 32.07 g/mol, H₂SO₄ = 98.09 g/mol.)

Step 1 — Note the atom chain: S → SO₂ → SO₃ → H₂SO₄. Every sulfur atom is conserved, so the overall mole ratio is 1:1.

Step 2 — Grams to moles:

$n(S) = 32.1 g \times 1 mol S 32.07 g S = 1.00 mol S$

Step 3 — Apply the 1:1 ratio and convert to grams:

$1.00 mol H 2 SO 4 \times 98.09 g H 2 SO 4 1 mol H 2 SO 4 = 98.1 g H 2 SO 4$

Check: 32 g of sulfur becomes about 98 g of acid — mass increases because oxygen and hydrogen were added. The atom-conservation shortcut saves three separate calculations.

Example 2: Roasting pyrite — sulfur dioxide production

How many grams of SO₂ are released when 480 g of pyrite, FeS₂, is roasted? Balanced equation: 4 FeS₂(s) + 11 O₂(g) → 2 Fe₂O₃(s) + 8 SO₂(g). (Molar masses: FeS₂ = 119.99 g/mol, SO₂ = 64.07 g/mol.)

Step 1 — Grams to moles:

$n(FeS 2) = 480 g \times 1 mol FeS 2 119.99 g FeS 2 = 4.00 mol FeS 2$

Step 2 — Mole ratio (8 mol SO₂ per 4 mol FeS₂):

$4.00 mol FeS 2 \times 8 mol SO 2 4 mol FeS 2 = 8.00 mol SO 2$

Step 3 — Moles to grams:

$8.00 mol SO 2 \times 64.07 g SO 2 1 mol SO 2 = 512.6 g SO 2$

Real-world note: that SO₂ must be captured (scrubbed) or converted to acid; released to air, the same mass contributes directly to acid rain.

Example 3: Percent sulfur in pyrite

What is the mass percent of sulfur in FeS₂? (Molar masses: Fe = 55.85, S = 32.07 g/mol.)

Step 1 — Formula for molar mass:

$M(FeS 2) = 55.85 + 2(32.07) = 119.99 g/mol$

Step 2 — Percent sulfur:

$% S = 2(32.07) g S 119.99 g FeS 2 \times 100 % = 53.5 %$

More than half the mass of "fool's gold" is actually sulfur — which is why roasting it produces so much SO₂.

Key takeaways

Native sulfur is S₈ crown rings; allotropes: orthorhombic (stable < 95.5 °C), monoclinic (95.5–119 °C), plastic/amorphous (quenched polymer chains).
Frasch process: superheated water (~170 °C) melts underground S; compressed air lifts it; product >99% pure.
Roasting sulfides produces SO₂: 4 FeS₂ + 11 O₂ → 2 Fe₂O₃ + 8 SO₂.
Oxidation states: −2 (H₂S, sulfides), 0 (S₈), +4 (SO₂, SO₃^2-), +6 (SO₃, SO₄^2-, H₂SO₄).
Contact process: S → SO₂ → SO₃ ^V₂O₅⟶ H₂SO₄; oleum (H₂S₂O₇) intermediates the SO₃ + H₂O step.
H₂SO₄ is the most-produced industrial chemical; concentrated acid is a strong acid, oxidizer, and dehydrating agent.
SO₂ from burning sulfur fuels causes acid rain; H₂S is toxic and paralyzes the sense of smell.
Sulfur is essential in biology: cysteine, methionine, and protein disulfide bonds.

Check yourself

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

Describe how the Frasch process works and why no chemical purification is needed.
Show answer
Superheated water (~170 °C) is pumped underground to melt the sulfur (mp ~115 °C); compressed air then forces the molten sulfur up through an inner pipe. Because the deposit is already nearly pure, the melt comes out >99% pure without chemical treatment.
What are the three main allotropes of solid sulfur, and what happens to molten sulfur above ~160 °C?
Show answer
Orthorhombic (stable below 95.5 °C), monoclinic (95.5–119 °C), and plastic/amorphous (chains formed above ~160 °C, then rapid cooling). Above ~160 °C the S₈ rings open into long polymer chains, making the liquid dark and viscous.
List sulfur's four important oxidation states with one example compound each.
Show answer
−2: H₂S or FeS; 0: S₈; +4: SO₂ or SO₃^2-; +6: SO₃, SO₄^2-, or H₂SO₄.
How many grams of H₂SO₄ can be made from 64.1 g of sulfur (S = 32.07, H₂SO₄ = 98.09 g/mol)?
Show answer
64.1 g ÷ 32.07 g/mol = 2.00 mol S; 1:1 atom conservation → 2.00 mol H₂SO₄ = 2.00 × 98.09 = 196.2 g H₂SO₄.
Write the three steps of the contact process and name the catalyst.
Show answer
(1) S + O₂ → SO₂; (2) 2 SO₂ + O₂ ⇌ 2 SO₃ over V₂O₅ at ~450 °C; (3) SO₃ + H₂O → H₂SO₄ (via oleum industrially).

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Allotrope: Different structural form of the same element
Frasch process: Mining molten sulfur with superheated water + compressed air
Sulfide: Compound with S^2- (S in −2)
Sulfur dioxide: SO₂, pungent gas with S in +4
Sulfate: SO₄^2- ion (S in +6)
Contact process: Catalytic route SO₂ + 12O₂ → SO₃ → H₂SO₄
Oleum: H₂S₂O₇, SO₃ dissolved in H₂SO₄

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

Occurrence: free and combined

Preparation: the Frasch process and roasting

Allotropes: one element, many structures

Properties and oxidation states

The contact process: from sulfur to sulfuric acid

Common Confusions

The same idea, in plain words

Worked example

Example 1: Contact process — how much acid from how much sulfur?

Example 2: Roasting pyrite — sulfur dioxide production

Example 3: Percent sulfur in pyrite

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references