Understanding Hybridisation of SO₂: Sulfur Dioxide

Sulfur Dioxide (SO₂) is also known as sulfurous anhydride. It is a molecule with a bent shape and is part of the main pollutants in the atmosphere. It is the toxic component in acid rains. It’s a good example of sp² hybridisation in inorganic chemistry.

Let us understand how hybridisation happens in SO₂. Read on to learn how it leads to its bonding and molecular shape.

What is the Hybridisation of SO₂?

Sulfur Dioxide consists of one sulfur atom and two oxygen atoms. Each of the oxygen atoms is bonded to the sulfur atom. In all, the molecule also contains a lone pair of electrons. In order to form these bonds and satisfy the octet rule, sulfur undergoes sp² hybridisation.

Using the Hybridisation Formula

We can determine the hybridisation of sulphur dioxide using the simple formula:

Step-by-step calculation:

Valence electrons of central atom (S): 6
Number of sigma bonds with central atom: 2
Negative charge: 0
Positive charge: 0

Interpretation:

A hybridisation number of 4 corresponds to sp² hybridisation (one orbital remains for π bonding).
This explains the bent molecular geometry of SO₂, with a bond angle of approximately 119°, due to one lone pair on sulfur repelling the bonding pairs.

Breakdown of SO₂ Hybridisation

The sulfur dioxide’s hybridisation has a resonant structure as it has a delocalised π system over the two S–O bonds, making both bonds equivalent. SO₂ also contains a lone pair of electrons.

Here is a complete understanding of its hybridisation.

Electronic Configuration of Sulfur

The atomic number of sulfur is 16.

The ground state of sulfur :

1s² 2s² 2p⁶ 3s² 3p⁴

Only two unpaired electrons → insufficient to form two required double bonds and also satisfy the octet rule

Excited state configuration:

1s² 2s² 2p⁶ 3s¹ 3p³ 3d¹

Four unpaired electrons → enough to form two bonds with O which are strong, and along with that also undergo resonance

Formation of Hybrid Orbitals

sp² hybridisation occurs when 1 s orbital and 2 p orbitals mix.
The result:
→ 3 sp² hybrid orbitals on the sulfur atom
→ The remaining 1 unhybridised p orbital stays available for π bonding

Bond Formation in Sulfur Dioxide

The sulfur atom in SO₂ undergoes sp² hybridisation, resulting in three sp² hybrid orbitals and one unhybridised p orbital.

Two sp² orbitals form σ bonds with the two oxygen atoms.
One sp² orbital holds the lone pair of electrons on sulfur.
The unhybridised p orbital on sulfur overlaps sideways with a p orbital of oxygen to form a π bond. Due to resonance, this π bond is delocalised over both S–O bonds, making them equivalent.

Result:

2 S–O σ bonds
1 delocalised π bond over the two S–O bonds

Hybridisation type: sp²
Bond angle: ~119° (caused by repulsion between lone pair and bonding pairs)
Geometry: Bent

Details At A Glance

Property	Details
Molecule	Sulfur Dioxide (SO₂)
Hybridisation	sp²
Geometry	Bent
Bond angle	~119°
Bonding	2 σ bonds (S–O), π bond (delocalised over both single bonds)
Unhybridised Orbitals	1 (on sulfur for π bonding)
Sulfur valency satisfied?	Yes, by forming 2 bonds and having 1 lone pair of electrons

Formal Charge in SO₂

To determine if the Lewis structure of SO₂ is stable, we calculate the formal charge on each atom using the formula:

Formal charge = Valence electrons – (Lone pair electrons + ½ × Bonding electrons)

Step-by-step for each atom:

Sulfur (S) – central atom

Valence electrons: 6
Lone pairs: 1
Bonding electrons: 8
(from double bond with both O (presence of resonance))

Formal charge = 6 – (2 + ½×8) = 6 – (2 + 4) = 6 – (6) = 0

Oxygen (O) – each

Valence electrons: 6
Lone pairs: 2
Bonding electrons: 4
(1 double bond with sulfur)

Formal charge = 6 – (4 + ½×4) = 6 – (4 + 2) = 6 – (6) = 0

Thus, all atoms in SO₂ carry zero formal charge, confirming that the Lewis structure is stable and correct.

Summing Up

The sulfur in SO₂ forms 2 bonds with oxygen and shares a π bond (delocalised) with the same atoms. sp² hybridisation leads to a bent shape and ~119° bond angles. The change in bond angle is because of the repulsion between lone pair and bonding pairs.