Why does iron commonly show +2 and +3 oxidation states?

Iron has the electron configuration [Ar] 3d⁶ 4s². Losing the two 4s electrons gives Fe²⁺ ([Ar] 3d⁶). Losing one additional 3d electron gives Fe³⁺ ([Ar] 3d⁵), which has a half-filled d subshell — a particularly stable arrangement. The extra exchange energy of the half-filled 3d⁵ configuration makes Fe³⁺ competitive with Fe²⁺, so both states coexist in chemistry.

What is the difference between ferrous and ferric iron?

Ferrous refers to Fe²⁺ and ferric to Fe³⁺. In aqueous solution Fe²⁺ forms pale green complexes (e.g., [Fe(H₂O)₆]²⁺) while Fe³⁺ forms yellow-brown complexes (e.g., [Fe(H₂O)₆]³⁺). Fe²⁺ is the more reducing and more soluble form; Fe³⁺ readily precipitates as Fe(OH)₃ at neutral pH.

What is the ferrate ion and why is Fe(VI) unusual?

The ferrate ion FeO₄²⁻ contains iron in the +6 oxidation state. It is a powerful oxidizing agent and forms purple salts such as K₂FeO₄. Fe(VI) is unusual because it strips iron of six electrons, pushing well beyond the normal +2/+3 range. Ferrate is studied as an eco-friendly water disinfectant because its reduction product is the non-toxic Fe(OH)₃.

What color is Fe²⁺ vs Fe³⁺ in solution?

Aqueous Fe²⁺ (the hexaaquairon(II) ion) is pale green. Aqueous Fe³⁺ (the hexaaquairon(III) ion) is yellow-brown or amber, which is why rust-contaminated water appears yellow. The color difference arises from different d–d transition energies caused by the different number of d electrons and consequent crystal-field splitting.

Why is Fe³⁺ more stable than Fe²⁺ in air?

Fe²⁺ is readily oxidized to Fe³⁺ by atmospheric oxygen, which is why iron left exposed to damp air rusts (forming Fe₂O₃ and Fe(OH)₃). The half-filled 3d⁵ configuration of Fe³⁺ provides extra exchange energy stability. Additionally, Fe³⁺ precipitates as insoluble hydroxide at near-neutral pH, removing it from solution and driving the equilibrium toward oxidation.

Does iron ever show negative oxidation states?

Yes. In iron carbonyl compounds such as [Fe(CO)₄]²⁻ (tetracarbonylferrate), iron formally carries a −2 oxidation state. These organometallic species exist because the CO ligands back-donate electron density onto iron; they are laboratory curiosities rather than common industrial compounds.

Iron Oxidation States Reference

Name: Iron Oxidation States Reference
Availability: InStock
Author: Nham Vu

Reference for all iron oxidation states: common Fe²⁺ (ferrous) and Fe³⁺ (ferric), plus 0, +1, +4, and +6 (ferrate). Includes d-electron configs, solution colors, and example compounds.

Atomic # 26 Fe Iron

Atomic Mass

55.845 u

Group

8 (VIII B)

Period

Block

d-block

Electronegativity

1.83 (Pauling)

Oxidation States

+2, +3 (common)

As a d-block transition metal, iron can lose varying numbers of electrons from its [Ar] 3d⁶ 4s² ground-state configuration. The most common states are +2 (ferrous) and +3 (ferric). Fe³⁺ benefits from the extra exchange energy of the half-filled 3d⁵ configuration, making it unusually stable among the two. Rarer states include 0 in metal and organometallics, +1, +4 in specialized oxides, and +6 in the purple ferrate ion FeO₄²⁻.

State	Name	d-electrons	Config (after Ar)	Solution Color	Stability
0	Elemental	6 (+ 4s²)	[Ar] 3d⁶ 4s²	N/A (solid metal)	Elemental only
+1	Iron(I)	7	[Ar] 3d⁷	N/A (organometallic)	Rare / organometallic
+2	Ferrous	6	[Ar] 3d⁶	Pale green	Common / stable
+3	Ferric	5	[Ar] 3d⁵	Yellow-brown / amber	Most common
+4	Iron(IV)	4	[Ar] 3d⁴	N/A	Rare / unstable
+6	Ferrate(VI)	2	[Ar] 3d²	Purple (ferrate)	Strong oxidizer

Half-filled d-shell stability
Fe³⁺ has 3d⁵ — one electron in each d orbital. This maximizes exchange energy, conferring extra stability and explaining why Fe²⁺ is easily oxidized to Fe³⁺ in air and why Fe³⁺ often dominates at neutral pH.

Ferrate Fe(VI)
K₂FeO₄ is a purple, highly oxidizing salt. The FeO₄²⁻ ion is isoelectronic with permanganate and chromate. It reduces to Fe(OH)₃, making it useful as a green water treatment agent.

Iron forms compounds across multiple oxidation states. The most common are ferrous (+2) and ferric (+3) compounds, but Fe(0) carbonyls, Fe(IV) oxides, and Fe(VI) ferrates are also known.

Compound	Formula	Fe State	Color / Appearance	Notes
Iron(II) Sulfate	FeSO₄	+2	Pale green crystals / solution	Common laboratory reagent; reducing agent; used in water treatment. Mohr's salt is (NH₄)₂Fe(SO₄)₂·6H₂O.
Iron(II) Chloride	FeCl₂	+2	Pale yellow-green solid	Soluble; used as a reducing agent and precursor to other iron compounds.
Iron(II) Oxide	FeO	+2	Black powder	Non-stoichiometric; found in scales on hot-worked steel; reduces to Fe at high temperature.
Iron(II) Hydroxide	Fe(OH)₂	+2	White/green precipitate	Rapidly oxidized in air to Fe(OH)₃ (rust); forms at neutral pH from Fe²⁺ solutions.
Iron(III) Oxide (Rust)	Fe₂O₃	+3	Red-brown solid	Hematite; chief ore of iron and main component of rust; used as a pigment and in magnetic recording.
Iron(III) Chloride	FeCl₃	+3	Brown-black solid / orange soln	Strong Lewis acid; used in PCB etching, water treatment, and as a catalyst.
Iron(III) Hydroxide	Fe(OH)₃	+3	Brown-red gelatinous ppt	Precipitates from Fe³⁺ at pH > ~3; the final product of iron corrosion in water.
Iron(III) Sulfate	Fe₂(SO₄)₃	+3	Yellow solid	Used in water treatment as a coagulant; hydrolyzes in water to give Fe(OH)₃.
Magnetite	Fe₃O₄	+2/+3	Black magnetic solid	Mixed-valence oxide containing both Fe²⁺ and Fe³⁺ in a spinel structure; strongly ferrimagnetic.
Iron Pentacarbonyl	Fe(CO)₅	0	Pale yellow liquid	Organometallic; iron in zero oxidation state; highly toxic; used to make ultra-pure iron powder.
Potassium Ferrate	K₂FeO₄	+6	Purple-red solid	Strong oxidizer; isoelectronic with MnO₄⁻; reduces to Fe(OH)₃, useful for disinfection.

Fe (neutral)

3d orbitals

↑ ↓

6 d-electrons

Noble gas: [Ar] 3d⁶ 4s²

Ground state. 4s fills before 3d per Aufbau; 3d⁶ = 4 paired + 2 unpaired (Hund).

Fe²⁺ (ferrous)

3d orbitals

↑ ↓

6 d-electrons

Noble gas: [Ar] 3d⁶

4s² electrons lost first (higher energy in cation). 4 unpaired e⁻ → paramagnetic.

Fe³⁺ (ferric)

3d orbitals

↑ ↓

5 d-electrons

Noble gas: [Ar] 3d⁵

Half-filled 3d⁵ — one e⁻ per d orbital. Maximum exchange energy; 5 unpaired e⁻.

Why 4s loses before 3d in Fe²⁺? Although 4s fills before 3d in the neutral atom, the increased nuclear charge in cations makes 3d lower in energy than 4s. So both 4s electrons are removed first when forming Fe²⁺. This is why Fe²⁺ has the configuration [Ar] 3d⁶, NOT [Ar] 3d⁴ 4s².

Ionization	Energy (kJ/mol)	Electron removed from
IE₁ → Fe⁺	762.5	4s (highest energy e⁻ in neutral Fe)
IE₂ → Fe²⁺	1561.9	4s (removes second 4s electron)
IE₃ → Fe³⁺	2957	3d (now 4s empty; pulls from 3d⁶ → 3d⁵)
IE₄ → Fe⁴⁺	5290	3d (large jump — disrupts 3d⁵ half-fill)

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Summary

Reference for all iron oxidation states: common Fe²⁺ (ferrous) and Fe³⁺ (ferric), plus 0, +1, +4, and +6 (ferrate). Includes d-electron configs, solution colors, and example compounds.

How it works

Click a tab — Oxidation States, Compounds, Electron Config, or Physical Props — to explore each section.
The Oxidation States panel lists all known states with stability, d-electron count, and solution color.
The Compounds panel shows common iron compounds with their formulas, iron state, and color/appearance.
The Electron Config panel shows the orbital filling diagram for Fe, Fe²⁺, and Fe³⁺ side by side.
The Physical Props panel provides atomic and material data for quick reference.
Click any monospace table cell to copy its value to your clipboard.

Use cases

Students learning d-block chemistry and why transition metals have variable oxidation states.
Chemistry teachers illustrating half-filled d-shell stability with the Fe³⁺ example.
Environmental chemists distinguishing soluble Fe²⁺ from precipitating Fe³⁺ in water analysis.
Biochemists referencing iron redox chemistry in hemoglobin, cytochromes, and iron-sulfur proteins.
Anyone revising for exams covering transition metal chemistry, coordination compounds, or redox reactions.

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