Orbital Hybridisation in 3D – sp³, sp², sp, Sigma and Pi Bonds

What it shows

Hybridisation explains why molecules have the shapes they do. The central atom's 2s orbital mixes with one, two or three 2p orbitals to give sp, sp² or sp³ hybrid orbitals pointing 180°, 120° or 109.5° apart; the number of electron domains decides which. In this 3D model you watch the orbitals mix, then see the bonded atoms arrive: end-on overlap along the bond axis forms σ bonds, and sideways overlap of the leftover p orbitals forms π bonds above and below it. Lone pairs in NH₃ and H₂O narrow the angles, while ethene and ethyne show double and triple bonds.

How to use

Choose a molecule and press Next step to move from 1. Atomic orbitals to 2. Hybridisation and 3. σ and π bonds; Replay hybridisation runs the mixing again. Ask students to predict the hybridisation and the number of σ and π bonds before step 3, then check the electron box diagram and the count below. Switch to Ball-and-stick model to compare bond types, and use Bond angles to compare CH₄, NH₃ and H₂O.

Parameters you can change

• Molecule CH₄ – methane, NH₃ – ammonia, H₂O – water, C₂H₄ – ethene, BF₃ – boron trifluoride, C₂H₂ – ethyne, CO₂ – carbon dioxide
• Step shown 1. Atomic orbitals, 2. Hybridisation, 3. σ and π bonds
• Model type Orbital model, Ball-and-stick model
• Show bond angles

Questions to explore

1. Why does carbon in CH₄ not bond with its pure 2p orbitals, and what bond angle would that give?
2. How many σ bonds and how many π bonds are there in ethene and in ethyne?
3. NH₃ and BF₃ both have three bonded atoms, so why are their hybridisation and shape different?