Why is +3 the most stable oxidation state for antimony?

Antimony has the electron configuration [Kr] 4d¹⁰ 5s² 5p³. In the +3 state it uses only the three 5p electrons for bonding, leaving the 5s² pair inert. This "inert pair effect" — where relativistic contraction makes the 5s² pair energetically reluctant to bond — makes Sb(III) the thermodynamically preferred state under most conditions.

What compounds contain antimony in the +5 state?

Antimony pentoxide (Sb₂O₅), antimony pentafluoride (SbF₅), antimony pentachloride (SbCl₅), and potassium hexafluoroantimonate (KSbF₆) all feature Sb(V). SbF₅ is one of the strongest known Lewis acids and reacts with fluoride donors to form the extremely strong superacid fluoroantimonic acid (HSbF₆).

What is the −3 oxidation state of antimony?

Antimony achieves −3 in metal stibides — ionic compounds where antimony accepts three electrons from electropositive metals, forming the stibide ion Sb³⁻. Sodium stibide (Na₃Sb) is a classic example. Stibine (SbH₃) is the −3 hydride analogue, analogous to ammonia and phosphine, though far less stable.

How does the inert pair effect explain antimony chemistry compared to nitrogen and phosphorus?

In nitrogen (Period 2) and phosphorus (Period 3), both the ns² and np³ electrons readily participate in bonding, so +5 is easily achieved. As you descend Group 15 to arsenic, antimony, and bismuth, relativistic effects stabilize the ns² pair, making it less available for bonding. This is why Sb(III) is more stable than Sb(V), and why bismuth Bi(III) is overwhelmingly dominant with virtually no stable Bi(V) chemistry.

What are the most important industrial uses of antimony compounds?

Antimony trioxide (Sb₂O₃) is the most commercially important compound, used primarily as a flame-retardant synergist with halogenated polymers in plastics, textiles, and electronics. Antimony is also alloyed with lead to harden battery grids. Antimony sulfide (Sb₂S₃) has historical use in pyrotechnics and is found in the mineral stibnite (the primary ore of antimony).

Yes. Antimony and its compounds are toxic to humans, with biological effects resembling arsenic poisoning. Stibine (SbH₃) is a highly toxic gas, similar in hazard profile to arsine (AsH₃). Chronic antimony exposure targets the respiratory system, liver, and skin. Occupational exposure limits are strictly regulated in most jurisdictions.

Antimony Oxidation States

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

Reference for antimony (Sb, Z=51) oxidation states (+5, +3, 0, −3), stability, electron configuration, and example compounds.

Atomic # 51 Sb Antimony

Atomic Mass

121.760 u

Group

15 (VA)

Period

Block

p-block

Electronegativity

2.05 (Pauling)

Oxidation States

+5, +3 (dominant), 0, −3

Antimony exhibits four oxidation states: +5, +3, 0, and −3, with +3 the most stable and most common. Its ground-state configuration is [Kr] 4d¹⁰ 5s² 5p³. The three 5p electrons are available for bonding in the +3 state, leaving the 5s² pair inert — the classic inert pair effect. Reaching +5 requires enough oxidizing power to engage both 5s electrons, while −3 is achieved when antimony acts as a formal anion in metal stibides.

Oxidation State	Stability	Notes
+5	Moderate — strong oxidant needed	Found in Sb₂O₅, SbF₅, SbCl₅, and KSbF₆. SbCl₅ is a strong Lewis acid. SbF₅ is one of the strongest Lewis acids known, forming superacids with HF. Requires both 5s² electrons to participate in bonding — energetically costly due to the inert pair effect.
+3	Stable — dominant	Most common state. Includes Sb₂O₃, SbCl₃, Sb₂S₃, and antimony tartrate. The 5s² lone pair remains non-bonding (stereochemically active). Dominant in aqueous chemistry and in flame-retardant applications as antimony trioxide.
0	Elemental only	Metallic and amorphous allotropes. The stable metallic form has a layered rhombohedral structure. Used in lead-acid battery alloys and semiconductor applications (e.g., InSb infrared detectors).
−3	Stibide / hydride only	Present in metal stibides such as Na₃Sb and in stibine (SbH₃). Stibine is a toxic gas (bp −17 °C), thermally less stable than phosphine (PH₃) due to the weaker Sb–H bond. The Sb³⁻ ion is strongly reducing.

Inert Pair Effect
Moving down Group 15 from N → P → As → Sb → Bi, the ns² electron pair becomes progressively less likely to bond. Relativistic contraction of the 5s orbital makes the 5s² pair in antimony more stable and less chemically available. This raises the energy needed to reach +5, making +3 the preferred state — a trend that culminates in bismuth, where Bi(III) is essentially the only accessible oxidation state.

Most common antimony compounds feature Sb in the +3 oxidation state. The most commercially significant is antimony trioxide (Sb₂O₃), which is added to polymers and textiles as a flame-retardant synergist with halogenated compounds. The +5 state is accessed with strong oxidants, while −3 appears in specialized intermetallic and hydride chemistry.

Compound	Formula	Sb State	Notes
Antimony Trioxide	Sb₂O₃	+3	White solid; the most important commercial antimony compound. Flame-retardant synergist in plastics, textiles, and rubber when combined with halogenated compounds.
Antimony Trichloride	SbCl₃	+3	Colorless hygroscopic solid; mild Lewis acid. Used as a catalyst in organic synthesis and as a Baeyer–Villiger oxidation catalyst precursor. Hydrolyzes readily in water.
Antimony Sulfide	Sb₂S₃	+3	Orange-red solid; the mineral stibnite, the principal ore of antimony. Historical use in pyrotechnics and cosmetics (kohl). Found in volcanic environments.
Potassium Antimony Tartrate	K(SbO)C₄H₄O₆	+3	Known as tartar emetic. Historical use in medicine as an emetic and anti-parasitic. Still used in veterinary medicine and as a mordant in textile dyeing.
Antimony Pentoxide	Sb₂O₅	+5	Yellow solid; used as a flame retardant and as an ion-exchange material. Less common than the trioxide; requires oxidizing conditions to form.
Antimony Pentafluoride	SbF₅	+5	One of the strongest Lewis acids known. Reacts with HF to produce fluoroantimonic acid (HSbF₆), one of the strongest known superacids. Extremely reactive.
Antimony Pentachloride	SbCl₅	+5	Fuming liquid; strong Lewis acid and chlorinating agent. Reacts with organic compounds and used as a catalyst. Hydrolyzes violently in water to give Sb₂O₅ and HCl.
Potassium Hexafluoroantimonate	KSbF₆	+5	Air-stable salt of the SbF₆⁻ anion. Used in electrochemistry and as a non-coordinating anion source in synthesis.
Sodium Stibide	Na₃Sb	−3	Intermetallic compound; formal Sb³⁻ anion. Strong reducing agent. Reacts vigorously with water. Representative metal stibide.
Stibine	SbH₃	−3	Colorless toxic gas (bp −17 °C); the antimony hydride analogue of ammonia. Less thermally stable than PH₃. Produced by reduction of Sb³⁺ with strong reducing agents.

Ground State Configuration

Inert pair — remains non-bonding in Sb(III)

Three bonding electrons — used in +3 and +5 states

Notation

Full: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p³

Noble gas: [Kr] 4d¹⁰ 5s² 5p³

Sb³⁺: [Kr] 4d¹⁰ 5s² (inert pair remains)

Sb⁵⁺: [Kr] 4d¹⁰

Ionization Steps

[Kr] 4d¹⁰ 5s² 5p³

Neutral atom — three 5p and two 5s valence electrons

Sb³⁺

[Kr] 4d¹⁰ 5s²

Lose three 5p electrons → most stable ion; 5s² pair is inert (lone pair)

Sb⁵⁺

[Kr] 4d¹⁰

Lose two further 5s electrons → requires strong oxidants; very Lewis acidic

Sb³⁻

[Kr] 4d¹⁰ 5s² 5p⁶

Gain three electrons → stibide ion in Na₃Sb; strongly reducing

Why the 5s² pair stays put in Sb(III)
In Sb³⁺, the 5s² pair is retained as a lone pair that sits in a large, asymmetric orbital — a stereochemically active lone pair that determines the geometry of Sb(III) compounds (pyramidal, analogous to ammonia). Removing this pair to reach Sb⁵⁺ costs far more energy than removing the 5p³ electrons, which is why +5 requires fluorine or concentrated nitric acid to achieve.

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Summary