Hybridization of C2H4 – Ethene (Ethylene)

Ethene, commonly known as ethylene, serves as the simplest member of the alkene family. This small molecule, with its seemingly straightforward structure of two carbon atoms double-bonded to each other and four hydrogen atoms, plays a pivotal role both in nature and industry.

Understanding ethene’s hybridization not only provides insights into its physical properties and chemical behavior but also sets the stage for exploring the broader class of unsaturated hydrocarbons.

What is the Hybridization of Ethene?

The hybridization of ethene (C₂H₄) involves sp² hybridization for each of its carbon atoms. Here’s a quick overview:

• Carbon Ground State: Each carbon starts with an electron configuration of 1s² 2s² 2p².
• Promotion: One 2s electron is promoted to a 2p orbital, making it 1s² 2s¹ 2p³.
• Hybridization: Each carbon then combines one 2s orbital with two 2p orbitals (let’s say 2px and 2py, assuming the molecule lies in the xy plane) to form three sp² hybrid orbitals. The remaining 2p orbital (2pz) is unhybridized.
• Bond Formation:
  • σ (Sigma) Bonds: Each carbon uses its three sp² orbitals to form three sigma bonds: one with the other carbon atom and two with hydrogen atoms. This results in four C-H σ bonds and one C-C σ bond in ethene.
  • π (Pi) Bond: The unhybridized p orbitals on each carbon overlap laterally to form a π bond, which exists above and below the plane of the σ bond framework between the carbon atoms.

Hybridization Process

1. Carbon Ground State Configuration:
   • Carbon in its ground state has an electron configuration of 1s² 2s² 2p².
2. Excitation:
   • Before bonding, one of the 2s electrons is promoted to the 2p orbital, resulting in 1s² 2s¹ 2p³. Now, carbon has four unpaired electrons available for bonding.
3. Hybridization:
   • In ethene, each carbon atom undergoes sp² hybridization. This involves combining one s orbital and two p orbitals to form three sp² hybrid orbitals. The remaining p orbital (let’s say pz if we consider the hybridized orbitals in the x-y plane) remains unhybridized.
4. Formation of Orbitals:
   • Three sp² hybrid orbitals: These are oriented in a trigonal planar arrangement with an angle of 120 degrees between them. Each of these orbitals will form a σ (sigma) bond.
   • One p orbital: This remains perpendicular to the plane of the sp² orbitals and will later form a π (pi) bond.

Bond Formation in Ethene:

• σ Bonds:
  • Each carbon atom uses its three sp² hybrid orbitals to form:
    • One σ bond with another carbon atom.
    • Two σ bonds with two hydrogen atoms.
  • So, in total, there are five σ bonds in ethene: one C-C σ bond and four C-H σ bonds.
• π Bond:
  • The unhybridized p orbitals on each carbon overlap side-by-side, forming one π bond. This π bond is located above and below the plane of the σ bond framework between the two carbon atoms.

Molecular Geometry and Properties:

• Geometry: The molecule has a planar geometry due to the sp² hybridization, with all atoms lying in the same plane.
• Bond Lengths and Strength:
  • The C-C σ bond in ethene is shorter and stronger than in ethane (C₂H₆) due to the increased s-character in sp² orbitals compared to sp³ in ethane.
  • The π bond, being weaker than the σ bond, makes the double bond in ethene less than twice as strong as a single bond, but it does make the molecule more reactive.
• Rotational Barrier:
  • The presence of the π bond restricts rotation around the C=C bond, which is why isomers (cis/trans) can exist for substituted ethenes.

Significance in Chemistry:

• Reactivity: The π bond in ethene makes it more reactive than ethane. It can easily undergo addition reactions where the π bond breaks, and new σ bonds form.
• Polymerization: Ethene is a crucial industrial chemical, primarily used in the production of polyethylene, where the double bond enables polymerization.