1.7 sp3 Hybrid Orbitals and the Structure of Ethane

Structure and Bonding in Ethane

Ethane is the simplest molecule that contains a carbon-carbon bond, making it the perfect starting point for understanding how sp3 hybridization works in practice. By studying ethane, you can see how hybrid orbitals explain molecular geometry, bond strength, and the ability of single bonds to rotate.

Structure of the Ethane Molecule

Ethane ($CH_3CH_3$) consists of two methyl groups connected by a single C-C bond, giving it a total of two carbons and six hydrogens.

Each carbon atom in ethane undergoes sp3 hybridization. This means one 2s orbital and three 2p orbitals on carbon mix together to produce four equivalent sp3 hybrid orbitals. These four orbitals point toward the corners of a tetrahedron, which minimizes electron-pair repulsion (VSEPR) and produces bond angles of approximately 109.5°.

Here's how the bonds form:

• The C-C bond results from end-on (head-to-head) overlap of one sp3 orbital from each carbon. This creates a sigma ($\sigma$) bond.
• Each of the remaining three sp3 orbitals on each carbon overlaps with a hydrogen 1s orbital, forming six C-H sigma ($\sigma$) bonds total.

Every bond in ethane is a sigma bond. There are no pi bonds here.

Carbon-Carbon Bonding with sp3 Orbitals

The sp3 hybridization process is what allows carbon to form four bonds instead of the two you might expect from its ground-state electron configuration ($1s^2\, 2s^2\, 2p^2$). By mixing the 2s and all three 2p orbitals, carbon gets four equivalent hybrid orbitals, each holding one electron and ready to bond.

The C-C sigma bond in ethane forms through end-on overlap of two sp3 hybrid orbitals, one from each carbon. End-on overlap is the most effective type of orbital overlap, which is why sigma bonds are strong and directional.

This same bonding pattern extends to all saturated hydrocarbons (alkanes). Any time two sp3-hybridized carbons share a single bond, the mechanism is the same: end-on sp3–sp3 overlap producing a sigma bond.

C-H vs. C-C Bonds in Ethane

All bonds in ethane are sigma bonds, but the C-H and C-C bonds differ in length and strength:

The C-H bond is both shorter and stronger than the C-C bond. Two factors explain this:

• Atomic size: Hydrogen is much smaller than carbon, so the bonding electrons sit closer to both nuclei, resulting in a shorter bond with stronger attraction.
• Electronegativity difference: Carbon (2.5) and hydrogen (2.2) have a small electronegativity difference, which creates a slight bond polarity. The C-C bond, by contrast, is completely nonpolar since both atoms are identical. This polarity contributes modestly to C-H bond strength.

Hybridization and Bond Rotation

Carbon's ground-state electron configuration ($1s^2\, 2s^2\, 2p^2$) shows only two unpaired electrons in the 2p orbitals. To form four bonds, carbon promotes one 2s electron into the empty 2p orbital and then hybridizes all four orbitals into equivalent sp3 hybrids. Each sp3 orbital has 25% s character and 75% p character.

One important consequence of sigma bonding: free rotation around the C-C bond. Because a sigma bond has cylindrical symmetry along the bond axis, rotating one methyl group relative to the other doesn't break or weaken the orbital overlap. This rotation gives ethane access to different conformations (staggered, eclipsed, and everything in between), which you'll explore further when studying conformational analysis.