Shapes of organic molecules; σ and π bonds — common mistakes

When counting bonds, always draw the full displayed formula first. It's easy to forget the C-H bonds in a skeletal or structural formula. A common mistake is to count a double bond as two $\sigma$ bonds; remember it's always one $\sigma$ and one $\pi$.

A π bond relies on the sideways overlap of p-orbitals. For this to be effective, the atoms must be held close together at a fixed distance. This is the job of the much stronger σ bond, which forms the primary framework of the molecule.

No. A C=C bond consists of one σ bond and one π bond. The π bond is weaker than the σ bond. Therefore, the bond enthalpy of a C=C bond (e.g., ~614 kJ mol⁻¹) is greater than a C-C bond (e.g., ~348 kJ mol⁻¹), but less than double.

Hybridisation is a more advanced model to explain these shapes. A tetrahedral carbon (like in ethane, 109.5°) is described as sp³ hybridised. A trigonal planar carbon (like in ethene, 120°) is sp² hybridised. A linear carbon (like in ethyne, 180°) is sp hybridised. You will explore this in more detail at A2 Level.