What are the common oxidation states of sulfur?

Sulfur most commonly shows −2 (sulfide: H₂S, Na₂S), 0 (elemental: S₈), +2 (S₂Cl₂, some thiosulfates), +4 (SO₂, H₂SO₃, SO₃²⁻), and +6 (SO₃, H₂SO₄, SO₄²⁻). Intermediate states such as +1, +3, and +5 appear in polysulfides and radical species but are less common.

What is the electron configuration of sulfur in its oxidation states?

Neutral sulfur is [Ne] 3s² 3p⁴. In the −2 state the 3p shell is complete ([Ne] 3s² 3p⁶). In 0 it retains [Ne] 3s² 3p⁴. As sulfur is oxidized toward +6, electrons are formally removed from the 3p and then 3s subshells, reaching [Ne] at the +6 limit.

Why can sulfur form so many oxidation states?

Sulfur has accessible 3d orbitals (or, more precisely, low-energy empty orbitals in the valence shell) that allow it to expand its octet and accommodate more than four bonds. This distinguishes it from oxygen, which is locked at −2 or −1 and cannot form +4 or +6 species.

What is the most stable oxidation state of sulfur?

Under standard aerobic conditions, +6 as sulfate (SO₄²⁻) is the thermodynamic sink. The −2 state (sulfide) is stable under anoxic (reducing) conditions such as deep sediments and volcanic environments.

What is a disproportionation reaction of sulfur?

In hot concentrated NaOH, elemental sulfur (0) disproportionates: 3 S + 6 NaOH → 2 Na₂S (−2) + Na₂SO₃ (+4) + 3 H₂O. Thiosulfate (S₂O₃²⁻) also disproportionates in acid: S₂O₃²⁻ + 2H⁺ → S (0) + SO₂ (+4) + H₂O.

What oxidation state is sulfur in H₂SO₄?

Sulfur is +6 in sulfuric acid (H₂SO₄). Oxygen is −2 (four oxygen atoms × −2 = −8) and each hydrogen is +1 (two × +1 = +2), so sulfur must be +6 to give a net charge of 0.

Sulfur Oxidation States

Name: Sulfur Oxidation States
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Author: Nham Vu

Reference for sulfur's common oxidation states (−2, 0, +2, +4, +6), example compounds, and redox/disproportionation context.

Atomic # 16 S Sulfur

Atomic Mass

32.06 u

Group

16 (VIA)

Period

Block

p-block

Electronegativity

2.58 (Pauling)

Oxidation States

−2, 0, +2, +4, +6

Sulfur spans one of the widest oxidation state ranges of any non-metal — from −2 in sulfide minerals to +6 in sulfate. Each state is stabilized by a distinct set of compounds found across industrial chemistry, geochemistry, and biology.

State	Common Species	Electron Config	Example Compound	Context
−2	S²⁻	[Ne] 3s² 3p⁶	H₂S	Volcanic gas, metal sulfide minerals, biological thiols
0	S₈	[Ne] 3s² 3p⁴	S₈	Elemental sulfur; rhombic/monoclinic allotropes
+2	S₂Cl₂	[Ne] 3s² 3p²	S₂Cl₂	Disulfur dichloride; thiosulfate ion edge case
+4	SO₂	[Ne] 3s² 3p⁰ 3d⁰	SO₂	Combustion product, acid rain precursor, food preservative
+6	SO₄²⁻	[Ne]	H₂SO₄	Sulfuric acid; most produced industrial chemical worldwide

Electron configuration (neutral S): [Ne] 3s² 3p⁴
Sulfur has 6 valence electrons and can accept 2 more (→ −2) or lose up to 6 (→ +6) using its available 3p and 3d orbitals. This breadth is unique among period 2–3 non-metals.

S²⁻

Sulfide

Oxidation State

−2

Key Ion/Species

S²⁻ (sulfide)

Electron Config

[Ne] 3s² 3p⁶

Valence Electrons

8 (full octet)

Geometry

Tetrahedral in H₂S

Character

Reducing agent; nucleophile

The −2 state is the most reduced form of sulfur. The sulfide ion (S²⁻) has a complete 3p⁶ subshell, making it isoelectronic with argon. H₂S is a colorless, toxic gas with a distinctive rotten-egg odor; it is produced naturally in volcanic emissions, deep-sea hydrothermal vents, and anaerobic decomposition. Metal sulfides (pyrite FeS₂, galena PbS, cinnabar HgS) are economically important ores. In biology, cysteine and methionine carry the −2 state via thiol (−SH) groups.

Key Compounds

Compound	Formula	Notes
Hydrogen Sulfide	H₂S	Toxic gas; weak acid (pKa 6.98); used in H₂S cracking and analytical precipitation.
Sodium Sulfide	Na₂S	Yellow flakes; highly alkaline; used in pulp and paper (kraft process).
Iron(II) Sulfide	FeS	Black solid; formed in anaerobic sediments; precursor to pyrite FeS₂.
Lead(II) Sulfide	PbS	Black mineral galena; main lead ore; infrared detector material.

S₈

Elemental

Oxidation State

Common Allotrope

S₈ (crown ring)

Electron Config

[Ne] 3s² 3p⁴

Physical Form

Yellow crystalline solid

Melting Point

113 °C (rhombic)

Character

Can be oxidized or reduced

Elemental sulfur is the reference state (oxidation state = 0). At room temperature the stable allotrope is rhombic sulfur — puckered S₈ crown rings. Sulfur is a disproportionation midpoint: it can act as an oxidizing agent (accepting electrons toward −2) or a reducing agent (donating electrons toward +4 or +6). Vast quantities are recovered as a by-product of petroleum hydrodesulfurization. The Claus process converts H₂S to elemental sulfur for storage and further processing.

Key Forms and Uses

Form / Use	Formula / Name	Notes
Rhombic Sulfur	α-S₈	Stable below 96 °C; used in fertilizers, rubber vulcanization.
Monoclinic Sulfur	β-S₈	Stable 96–119 °C; needle-like crystals; converts to rhombic on cooling.
Plastic (amorphous) S	Sₙ	Formed by quenching molten sulfur; elastic, dark rubber-like; unstable.
Flowers of Sulfur	S (fine)	Powdered sulfur from condensation; used in agriculture as fungicide.

S(+2)

Uncommon

Oxidation State

Example Species

S₂Cl₂, S in S₂O₃²⁻ (avg.)

Electron Config

[Ne] 3s² 3p²

Stability

Limited; unstable in water

Occurrence

Rare; intermediates and halides

Character

Mild oxidizing agent

The +2 state is the least common of sulfur's main oxidation states. It appears clearly in disulfur dichloride (S₂Cl₂) and sulfenyl chloride (RSCl) derivatives. In thiosulfate (S₂O₃²⁻) the formal average S state is +2, though the two sulfur atoms are inequivalent — one is formally +5 and the other −1 in a bonding model that considers the S–S bond. The +2 state does not have a simple aqueous oxo-ion; it is mainly encountered as covalent halide compounds and short-lived intermediates in sulfur combustion.

Key Compounds

Compound	Formula	Notes
Disulfur Dichloride	S₂Cl₂	Fuming yellow liquid; reacts violently with water; used in rubber vulcanization.
Thiosulfate Ion	S₂O₃²⁻	Average S oxidation state +2; widely used in photography (fixer) and iodometric titrations.
Sulfenyl Chloride	CH₃SCl	Organic S(+2) derivative; used as electrophilic sulfur source in synthesis.
Disulfur Monoxide	S₂O	Transient gas-phase species; detected in volcanic and combustion environments.

SO₂

Sulfite / SO₂

Oxidation State

Key Species

SO₂, SO₃²⁻, H₂SO₃

Electron Config

[Ne] 3s² (formally)

Geometry (SO₂)

Bent, 119°

Aqueous Acid

Sulfurous acid (H₂SO₃)

Character

Reducing or mild oxidizing agent

Sulfur dioxide (SO₂) is the dominant +4 species. It is produced by burning sulfur or sulfide ores and is the primary precursor to acid rain when oxidized to SO₃ in the atmosphere. SO₂ dissolves in water to give sulfurous acid (H₂SO₃) and the sulfite ion (SO₃²⁻). It is used industrially as a preservative (E220 in food), a bleaching agent in paper manufacture, and as an intermediate in sulfuric acid production. The +4 state can disproportionate: in hot concentrated sulfuric acid, SO₂ is both oxidized and reduced.

Key Compounds

Compound	Formula	Notes
Sulfur Dioxide	SO₂	Colorless gas; pungent; acid rain precursor; industrial bleach/preservative.
Sulfurous Acid	H₂SO₃	Weak diprotic acid; exists only in aqueous solution; pKa1 1.81.
Sodium Sulfite	Na₂SO₃	White solid; reducing/antioxidant agent; used in photography and winemaking.
Sulfuryl Difluoride	SF₄	Toxic gas; S(+4); Lewis acid; used in organic fluorination reactions.

SO₄²⁻

Sulfate

Oxidation State

Key Species

SO₄²⁻, H₂SO₄, SO₃

Electron Config

[Ne] (formally 3d-expanded)

Geometry (SO₄²⁻)

Tetrahedral

Aqueous Acid

Sulfuric acid (H₂SO₄), strong diprotic

Character

Thermodynamic sink; strong oxidant when conc.

The +6 state is the highest and most oxidized form of sulfur. Sulfuric acid (H₂SO₄) is the world's highest-volume manufactured chemical (~270 million tonnes/year) and is produced via the contact process: S is burned to SO₂ (+4), which is oxidized over a V₂O₅ catalyst to SO₃ (+6), then absorbed into oleum and diluted. Sulfate (SO₄²⁻) is the thermodynamically stable sulfur species in aerobic aqueous environments. Sulfur hexafluoride (SF₆) also carries the +6 state and is an extremely stable, non-toxic gas used as an electrical insulator.

Key Compounds

Compound	Formula	Notes
Sulfuric Acid	H₂SO₄	World's most-produced industrial chemical; strong diprotic acid; contact process product.
Sulfur Trioxide	SO₃	Reacts violently with water; intermediate in H₂SO₄ production.
Sulfur Hexafluoride	SF₆	Colorless, odorless, non-toxic gas; extreme stability; used in high-voltage switchgear.
Calcium Sulfate	CaSO₄	Anhydrite/gypsum; plaster of Paris on hydration; used in construction and soil amendment.

Standard reduction potentials (E°) are given for acidic aqueous solution (pH 0, 25 °C) vs. SHE. More positive E° means the oxidized form is a stronger oxidant.

SO₄²⁻

Oxidation State +6

SO₄²⁻ + 4H⁺ + 2e⁻ → H₂SO₃ + H₂O

E°

+0.17 V

SO₂

Oxidation State +4

SO₂ + 4H⁺ + 4e⁻ → S + 2H₂O

E°

+0.50 V

Oxidation State 0

S + 2H⁺ + 2e⁻ → H₂S

E°

+0.14 V

H₂S

Oxidation State −2

2H⁺ + 2e⁻ → H₂ (ref.)

E°

~0.00 V

Selected Half-Reactions

Half-Reaction	E° (V vs SHE)	Context
SO₄²⁻ + 4H⁺ + 2e⁻ → H₂SO₃ + H₂O	+0.17	Sulfate reduction; relevant in anaerobic sulfate-reducing bacteria
SO₄²⁻ + 8H⁺ + 6e⁻ → S + 4H₂O	+0.36	Full sulfate-to-sulfur reduction; used in geochemical modeling
SO₄²⁻ + 10H⁺ + 8e⁻ → H₂S + 4H₂O	+0.30	Sulfate to sulfide; key in anoxic sediment chemistry
SO₂ + 4H⁺ + 4e⁻ → S + 2H₂O	+0.50	SO₂ used as reducing bleach; pulp and paper
S + 2H⁺ + 2e⁻ → H₂S	+0.14	Elemental sulfur to sulfide; Claus process reverse
2SO₃²⁻ + 3H₂O + 4e⁻ → S₂O₃²⁻ + 6OH⁻	−0.58	Sulfite to thiosulfate; alkaline solutions

Disproportionation of elemental sulfur:
In hot concentrated NaOH: 3 S + 6 NaOH → 2 Na₂S + Na₂SO₃ + 3 H₂O
S(0) is simultaneously reduced to −2 and oxidized to +4. Thiosulfate also disproportionates in acid: S₂O₃²⁻ + 2H⁺ → S↓ + SO₂ + H₂O

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Summary

Reference for sulfur's common oxidation states (−2, 0, +2, +4, +6), example compounds, and redox/disproportionation context.

How it works

Select an oxidation state tab (−2, 0, +2, +4, or +6) to view its properties.
The properties panel shows electron configuration, key aqueous or gaseous species, and stability notes.
The compounds table lists common examples and their industrial or biological relevance.
The Redox & Disproportionation tab shows half-reactions and standard potentials.
The Overview tab provides a side-by-side comparison of all five common states.
Click any formula cell in a table to copy its text.

Use cases

Students revising p-block oxidation state trends for chemistry exams.
Teachers explaining the industrial sulfuric acid (contact) process.
Geochemists modeling the sulfur cycle and sulfide mineral formation.
Environmental chemists tracking SO₂ and SO₃ in atmospheric chemistry.
Lab chemists balancing redox equations involving sulfur species.
Engineers working with hydrogen sulfide removal in natural gas processing.

Frequently Asked Questions

Last updated: 2026-07-08 · Reviewed by Nham Vu