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Understanding Hybridisation of C₂H₂: Ethyne

Understanding Hybridisation of C₂H₂: Ethyne

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. Read on to learn 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 and the other carbon. In order to form these bonds and satisfy the octet rule, carbon undergoes sp hybridisation.

Breakdown of C₂H₂ Hybridisation

Ethyne/acetylene is renowned as the first and simplest form of Alkyne, possessing two carbon atoms linked with triple bonds. The triple bond is generally called an acetylenic bond.

Here is a complete understanding of its hybridisation.

Electronic Configuration of Carbon

The atomic number of carbon is 6.

The ground state of carbon :

1s² 2s² 2p²

Only two unpaired electrons → insufficient to form four bonds

Excited state configuration:

formula

Four unpaired electrons → enough to form four bonds

image

Ground state vs excited state orbital diagram

Formation of Hybrid Orbitals

sp hybridisation occurs when 1 s orbital and 1 p orbital mix.
The result:
formula

Bond Formation in Ethyne

Each carbon uses:

  • 1 sp orbital to form a σ bond with hydrogen
  • 1 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
  • 2 C–H σ bonds
  • Hybridisation type: sp
  • Bond angle: 180°
  • Geometry: Linear
image

 Geometry and bonding of ethyne

Details At A Glance

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 using the formula:

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

Step-by-step for each atom:

Carbon (C) – each

  • Valence electrons: 4
  • Lone pairs: 0
  • Bonding electrons: 8 (4 from one triple bond with another C, 2 from one single bond with H)

Formal charge = 4 – (0 + ½ × 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 + ½ × 2) = 1 – 1 = 0

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

Summing Up

Each carbon in C₂H₂ forms 4 bonds: 1 with hydrogen and 3 with another 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 alkyne in our lives?

Alkyne is commercially used in the ripening of fruits. It helps fuel in welding torches, and in many polymers. It can be used as fuel.

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