1.8 sp2 Hybrid Orbitals and the Structure of Ethylene

sp2 Hybridization and Bonding in Ethylene

sp2 Hybrid Orbitals in Ethylene

In ethylene ($C_2H_4$), each carbon atom undergoes sp2 hybridization. This means one 2s orbital and two 2p orbitals mix together to form three equivalent sp2 hybrid orbitals. These three orbitals arrange themselves in a trigonal planar geometry, pointing away from each other at 120° angles to minimize electron repulsion.

That accounts for three of carbon's four orbitals. The leftover unhybridized 2p orbital sits perpendicular to the plane of the sp2 orbitals. This p orbital is what makes the pi bond possible.

Each carbon in ethylene uses its sp2 orbitals to form three sigma bonds:

• Two C–H bonds: overlap of a carbon sp2 orbital with a hydrogen 1s orbital
• One C–C sigma bond: head-on overlap of one sp2 orbital from each carbon, along the internuclear axis

The molecular geometry of ethylene is planar. All six atoms (two carbons and four hydrogens) lie in the same plane.

Sigma and Pi Bond Formation

The carbon-carbon double bond in ethylene consists of one sigma (σ) bond and one pi (π) bond. These form through different types of orbital overlap.

Sigma bond formation:

1. One sp2 orbital from each carbon points directly at the other carbon.
2. These orbitals overlap head-on along the internuclear axis, creating a strong, cylindrically symmetric bond.

Pi bond formation:

1. The unhybridized 2p orbital on each carbon is perpendicular to the molecular plane.
2. These two p orbitals overlap laterally (side-by-side), above and below the plane of the molecule.
3. This creates a pi bond, which is weaker than the sigma bond because the overlap is less direct.

Together, the sigma and pi bonds produce a carbon-carbon double bond that is both shorter and stronger than a single bond:

Notice that the double bond is not simply twice as strong as the single bond. The sigma bond contributes roughly 347 kJ/mol, so the pi bond accounts for only about 264 kJ/mol. That weaker pi bond is the one that breaks most easily during reactions.

Ethylene vs. Ethane Structure

Comparing ethylene and ethane highlights how hybridization shapes molecular properties.

• Hybridization: Carbon in ethylene is sp2 (trigonal planar, 120° bond angles). Carbon in ethane is sp3 (tetrahedral, 109.5° bond angles).
• Geometry: Ethylene is flat; all atoms lie in one plane. Ethane has a three-dimensional, tetrahedral arrangement around each carbon.
• Bonding: Ethylene has a C=C double bond (one σ + one π). Ethane has a C–C single bond (one σ only).
• Bond length and strength: The double bond in ethylene (1.34 Å, 611 kJ/mol) is shorter and stronger than the single bond in ethane (1.54 Å, 347 kJ/mol).

Reactivity is where this really matters. The pi bond in ethylene is relatively electron-rich and exposed above and below the molecular plane, making it a target for electrophiles. This is why ethylene readily undergoes addition reactions like hydrogenation (adding $H_2$) and halogenation (adding $Br_2$ or $Cl_2$). Ethane, with only a sigma bond, is far less reactive and typically requires more extreme conditions (like free-radical initiation) to react.

Valence Bond Theory and Molecular Structure

Valence bond theory provides the framework for understanding why ethylene looks and behaves the way it does. The key idea is that hybridization determines geometry:

• sp2 hybridization produces a trigonal planar arrangement with 120° bond angles.
• sp3 hybridization produces a tetrahedral arrangement with 109.5° bond angles.

In ethylene, the sp2 hybridization of each carbon dictates the planar structure, and the leftover p orbitals dictate where the pi bond forms. So hybridization doesn't just predict shape; it also predicts what types of bonds are available and, by extension, how the molecule will react.