What are the oxidation states of chlorine?

Chlorine has six common oxidation states: -1 (HCl, NaCl), 0 (Cl2), +1 (HOCl, NaOCl), +3 (NaClO2), +5 (NaClO3), and +7 (HClO4, NaClO4). The -1 state is the most stable in ionic and covalent compounds.

Why does chlorine have so many oxidation states?

Chlorine has seven valence electrons (configuration [Ne] 3s2 3p5) and empty 3d orbitals available for bonding. When paired with electronegative oxygen, chlorine can formally adopt positive states (+1 to +7) by shifting electron density to oxygen. This d-orbital availability is why fluorine — which lacks accessible d-orbitals — cannot achieve positive oxidation states.

What is disproportionation of chlorine?

Disproportionation is when an element in a single oxidation state simultaneously oxidizes and reduces itself. Cl2 in water undergoes: Cl2 + H2O → HCl + HOCl. Here Cl2 (0) becomes both Cl- (-1) and Cl+ (+1). This reaction produces the bleaching agent hypochlorous acid (HOCl).

Which chlorine oxidation state is the strongest oxidizer?

In acidic solution the order of oxidizing strength generally runs: Cl2 (+7 in HClO4 is actually the weakest oxoacid oxidant, while +1 in HOCl and Cl2 itself are far stronger under acidic conditions). HClO4 is a strong acid but a poor oxidant at room temperature; HOCl and Cl2 are the practical oxidants in bleaching and disinfection.

What is the difference between NaOCl, NaClO2, NaClO3, and NaClO4?

They differ in the chlorine oxidation state and number of oxygen atoms: NaOCl (hypochlorite, Cl +1, 1 O), NaClO2 (chlorite, Cl +3, 2 O), NaClO3 (chlorate, Cl +5, 3 O), and NaClO4 (perchlorate, Cl +7, 4 O). Each additional oxygen raises the oxidation state by 2 units.

Does chlorine have a +2 or +4 oxidation state?

Not in stable compounds. Even-numbered positive states (+2, +4, +6) are not observed for chlorine because they would require an odd number of unpaired electrons in a way that does not pair with the available oxygen ligands to form a stable octet/expanded structure. Only odd positive states (+1, +3, +5, +7) are stable.

Chlorine Oxidation States

Name: Chlorine Oxidation States
Availability: InStock
Author: Nham Vu

Reference for all chlorine oxidation states (-1, 0, +1, +3, +5, +7) with example compounds, stability notes, and disproportionation context.

Atomic # 17 Cl Chlorine

Atomic Mass

35.45 u

Group

17 (VIIA)

Period

Block

p-block

Electronegativity

3.16 (Pauling)

Oxidation States

-1, 0, +1, +3, +5, +7

Chlorine's seven valence electrons (3s² 3p⁵) and accessible 3d orbitals enable six oxidation states. The -1 state (gaining one electron to reach a full octet) is the most common and stable. Positive states arise when bonded to the more electronegative oxygen.

State	Stability	Example Compound	Notes
-1	Most stable	HCl, NaCl, KCl	Chloride ion — full octet, noble-gas configuration. Found in salts, hydrochloric acid, and biological systems.
0	Elemental	Cl₂	Diatomic gas; strong oxidizing agent; benchmark state for redox calculations.
+1	Moderate	HOCl, NaOCl	Hypochlorous acid / hypochlorite; active bleaching and disinfecting agent. Subject to disproportionation.
+3	Unstable	NaClO₂, ClF₃	Chlorite salts; ClF₃ is a powerful fluorinating agent. Rarely encountered in aqueous chemistry.
+5	Stable salt	NaClO₃, KClO₃	Chlorate; strong oxidizer used in matches and pyrotechnics; KClO3 decomposes to O2 on heating.
+7	Most stable positive	HClO₄, NaClO₄	Perchlorate; highest oxidation state; perchloric acid is one of the strongest known acids.

Key pattern: Only odd positive states exist (+1, +3, +5, +7). Even positive states (+2, +4, +6) are not observed in stable compounds — the electron pairing arithmetic does not yield a closed-shell or half-filled d arrangement that survives in a molecular context.

Chlorine forms a wide range of binary and ternary compounds. The oxoacid series — HOCl, HClO₂, HClO₃, HClO₄ — shows acid strength increasing with oxidation state.

Compound	Formula	Cl State	Notes
Hydrogen Chloride	HCl	-1	Strong acid in water; principal industrial acid; Cl is -1 since H is +1.
Sodium Chloride	NaCl	-1	Table salt; ionic; Na+ Cl- — the -1 state at its most stable.
Chlorine Gas	Cl₂	0	Diatomic; homoatomic bond assigns 0 by convention; green-yellow gas.
Hypochlorous Acid	HOCl	+1	Weak acid; active bleaching species in chlorinated water.
Sodium Hypochlorite	NaOCl	+1	Household bleach (3–8% solution); disinfectant and laundry agent.
Chlorous Acid	HClO₂	+3	Unstable weak acid; sodium chlorite (NaClO₂) used in water treatment.
Chlorine Trifluoride	ClF₃	+3	Highly reactive fluorinating agent; T-shaped geometry.
Chloric Acid	HClO₃	+5	Strong acid; chlorates are good oxidizers above ~300 °C.
Potassium Chlorate	KClO₃	+5	Decomposes to KCl + O₂ when heated with MnO₂ catalyst.
Perchloric Acid	HClO₄	+7	One of the strongest known acids; Cl reaches maximum +7 state.
Sodium Perchlorate	NaClO₄	+7	Highly soluble salt; used in rocket propellants and ion-pair reagents.
Chlorine Pentafluoride	ClF₅	+5	Square pyramidal; extremely reactive fluorinating agent.
Chlorine Heptafluoride	ClF₇	+7	Rare; only known Cl +7 binary fluoride; pentagonal bipyramidal.
Dichlorine Heptoxide	Cl₂O₇	+7	Anhydride of perchloric acid; oily explosive liquid.

Ground State Configuration

Core (He)

Core (Ne inner)

Valence s

Valence p — 1 electron short of octet

Notation

Full: 1s² 2s² 2p⁶ 3s² 3p⁵

Noble gas: [Ne] 3s² 3p⁵

How Each Oxidation State Arises

-1 Cl gains 1e⁻ → [Ne] 3s² 3p⁶ (full octet = Ar configuration)

0 Cl₂: shared bond, equal electronegativity → net 0 by convention

+1 Cl donates 1e⁻ to O in HOCl; 3p⁵ loses one paired electron

+3 Cl bonds three O/F atoms, formally loses 3e⁻ from 3s/3p

+5 Five bonds to O/F; 3d orbitals expand the coordination sphere

+7 Maximum: 7 bonds to O/F; all 3s² 3p⁵ electrons formally lost

Why not fluorine? Fluorine (F) also has 7 valence electrons but cannot reach positive oxidation states because it has no available d-orbitals (Period 2, n = 2 stops at p). Chlorine is Period 3, so empty 3d orbitals are energetically accessible for bonding with oxygen — enabling all the oxoacid states.

Disproportionation Reactions

Disproportionation occurs when a single element simultaneously oxidizes and reduces itself. Chlorine's intermediate states (+1, +3, +5) are prone to this in water.

Cl₂ in water (neutral/acid)

Cl₂ (0) + H₂O → HCl (-1) + HOCl (+1)

Produces hypochlorous acid — the active disinfecting species in chlorinated swimming pools and tap water treatment.

Cl₂ in alkali

Cl₂ (0) + 2 NaOH → NaCl (-1) + NaOCl (+1) + H₂O

Basis for industrial bleach production. At high temperature (70 °C+) the chlorate ClO₃⁻ (+5) forms instead of hypochlorite.

HOCl disproportionation (slow, acidic)

3 HOCl (+1) → HClO₃ (+5) + 2 HCl (-1)

Slow at room temperature; accelerated by heat or UV. Explains why hypochlorite bleaches lose potency over time.

KClO₃ thermal decomposition (not disproportionation)

2 KClO₃ (+5) → 2 KCl (-1) + 3 O₂

Redox: Cl is reduced (-1), oxygen is oxidized (0). Catalyzed by MnO₂; standard O₂ lab source.

Oxidizing Strength Comparison (acidic solution)

Oxidant	Cl State	E° (V, vs SHE)	Practical Strength
HClO₄ / Cl⁻	+7 → -1	+1.39 V	Weak oxidant at 25 °C despite highest state
HClO₃ / Cl⁻	+5 → -1	+1.45 V	Moderate; more reactive on heating
Cl₂ / Cl⁻	0 → -1	+1.36 V	Strong, practical industrial oxidant
HOCl / Cl⁻	+1 → -1	+1.49 V	Strongest practical chlorine oxidant in aqueous use

Standard reduction potentials at 25 °C, 1 M acid. Thermodynamics only — kinetics govern actual reaction rates.

Copied!

Summary

Reference for all chlorine oxidation states (-1, 0, +1, +3, +5, +7) with example compounds, stability notes, and disproportionation context.

How it works

Browse the Oxidation States tab to see every chlorine state with stability rating and an example compound.
Switch to the Compounds tab for a full table of common chlorine compounds, their formulas, and assigned oxidation state.
The Electron Config tab shows the orbital diagram and how electron gain or loss produces each state.
The Redox Context tab covers disproportionation reactions and practical oxidizing-agent strength ranking.
Click any formula or value in a table cell to copy it to the clipboard.

Use cases

Students studying halogen chemistry and variable oxidation state patterns.
Chemistry teachers building lesson material on chlorine oxoacids and bleaching.
Lab chemists selecting the right chlorine oxidant for a synthetic procedure.
Water treatment engineers understanding hypochlorite vs. chlorate chemistry.
Anyone preparing for AP or IB chemistry exams covering redox and Group 17.

Frequently Asked Questions

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