Sigma and pi bonds
Sigma and pi bonds
- Covalent bonds form when orbitals overlap.
- The way they overlap gives a sigma or a pi bond.
- Bond strength links to bond length.
Two kinds of overlap
- a sigma (σ) bond forms by head-on overlap between the two atoms.
- a pi (π) bond forms by the sideways overlap of two p orbitals (above and below).
- single = 1σ; double = 1σ + 1π; triple = 1σ + 2π.
Practice
A pi (π) bond forms by:
A σ bond is head-on overlap; a π bond is the sideways overlap of p orbitals.
Practice
A double bond consists of:
A single bond is 1σ; a double is 1σ + 1π; a triple is 1σ + 2π.
Hybridisation
- Hybridisation mixes orbitals in a shell into new, equal ones:
- sp → 2 orbitals (linear),
- sp² → 3 orbitals (flat, e.g. $\text{C}_2\text{H}_4$),
- sp³ → 4 orbitals (e.g. $\text{CH}_4$).
Practice
Methane (CH₄) uses which hybridisation?
CH₄ is tetrahedral with four equal bonds, from sp³ hybridisation.
Bond energy and length
- bond energy = energy to break one mole of a bond (gas state).
- bond length = distance between the two bonded nuclei.
- A shorter bond is usually stronger: triple > double > single.
Practice
Which bond is the shortest and strongest?
Shorter bonds are stronger: triple > double > single in both strength and shortness.
You've got it
Key idea
- σ = head-on overlap; π = sideways p-orbital overlap
- single = 1σ; double = 1σ + 1π; triple = 1σ + 2π
- hybridisation: sp (linear), sp² (flat), sp³ (tetrahedral)
- shorter bond → stronger (triple > double > single)