Hybridisation of C₂H₂ (Ethyne/Acetylene)

Ethyne (C₂H₂) is also known as acetylene. It is a hydrocarbon with a linear structure and a triple bond between its two carbon atoms. It's an excellent example of sp hybridisation in organic chemistry.

Let us understand how hybridisation happens in C₂H₂ and how it leads to its bonding and molecular shape.

What is the Hybridisation of C₂H₂?

Ethyne consists of two carbon atoms and two hydrogen atoms. Each carbon atom is bonded to one hydrogen atom and to the other carbon atom. In order to form these bonds and satisfy the octet rule, carbon undergoes sp hybridisation.

Breakdown of C₂H₂ Hybridisation

Ethyne (acetylene) is the first and simplest alkyne, possessing two carbon atoms linked by a triple bond. This triple bond is generally called an acetylenic bond.

Electronic Configuration of Carbon

Ground state vs excited state orbital diagram

The atomic number of carbon is 6. The ground state configuration of carbon is: 1s² 2s² 2p²

Only two unpaired electrons — insufficient to form four bonds.

Excited state configuration: 1s² 2s¹ 2pₓ¹ 2pᵗ¹ 2pᶻ¹

Four unpaired electrons — enough to form four bonds.

Formation of Hybrid Orbitals

sp hybridisation occurs when 1 s orbital and 1 p orbital mix. The result: 2 sp hybrid orbitals per carbon atom, while the remaining two unhybridised p orbitals (2pᵗ and 2pᶻ) stay available for π bonding.

Bond Formation in Ethyne

Each carbon uses one sp orbital to form a σ bond with hydrogen, and one sp orbital to form a σ bond with the other carbon. The two unhybridised p orbitals on each carbon atom overlap sideways to form 2 π bonds, giving rise to a C≡C triple bond.

Result: 1 σ bond and 2 π bonds between the two carbon atoms, plus 2 C–H σ bonds. Hybridisation type: sp. Bond angle: 180°. Geometry: Linear.

Geometry and bonding of ethyne

Property	Details
Molecule	Ethyne (C₂H₂)
Hybridisation	sp
Geometry	Linear
Bond angle	180°
Bonding	3 σ bonds (2 C–H, 1 C–C), 2 π bonds (C≡C)
Unhybridised Orbitals	2 (on each carbon, for π bonding)
Carbon valency satisfied?	Yes, by forming 4 bonds

Formal Charge in C₂H₂

To determine if the Lewis structure of C₂H₂ is stable, we calculate the formal charge on each atom:

Formal charge = Valence electrons - (Lone pair electrons + \frac{1}{2} \times Bonding electrons)

Carbon (C) – each: Valence electrons: 4; Lone pairs: 0; Bonding electrons: 8 (4 from the triple bond with the other C, plus a share from the single bond with H).

$Formal charge = 4 - (0 + \frac{1}{2} \times 8) = 4 - 4 = 0$

Hydrogen (H) – each: Valence electrons: 1; Lone pairs: 0; Bonding electrons: 2 (1 single bond with carbon).

$Formal charge = 1 - (0 + \frac{1}{2} \times 2) = 1 - 1 = 0$

Thus, all atoms in C₂H₂ carry zero formal charge, confirming that the Lewis structure is stable and correct.

Summing Up

Each carbon in C₂H₂ forms 4 bonds: 1 with hydrogen and 3 with the other carbon (1 σ and 2 π). sp hybridisation leads to a linear shape and 180° bond angles. The π bonds are responsible for the triple bond between the carbon atoms.

Frequently Asked Questions

Q1. Why does carbon undergo hybridisation in C₂H₂?
To form 4 covalent bonds, carbon promotes an electron and mixes orbitals to form hybrid orbitals.

Q2. How many σ and π bonds are present in C₂H₂?
There are 3 σ bonds and 2 π bonds in total.

Q3. What is the shape of ethyne?
Linear, due to sp hybridisation.

Q4. Is C₂H₂ polar or non-polar?
Ethyne is non-polar, as the bond dipoles cancel due to the linear geometry.

Q5. What are some uses of alkynes in our lives?
Alkynes are commercially used in the ripening of fruits, as fuel in welding torches, and in the production of many polymers.