Hybridisation (Cambridge (CIE) A Level Chemistry): Revision Note
Exam code: 9701
Orbitals & Hybridisation in Covalent Bonding
Bond overlap in covalent bonds
A single covalent bond is formed when two nonmetals combine
Each atom that combines has an atomic orbital containing a single unpaired electron
When a covalent bond is formed, the atomic orbitals overlap to form a combined orbital containing two electrons
This new orbital is called the molecular orbital
The greater the atomic orbital overlap, the stronger the bond
Sigma (σ) bonds are formed by direct overlap of orbitals between the bonding atoms
Pi (π) bonds are formed by the sideways overlap of adjacent above and below the σ bond
σ bonds
Sigma (σ) bonds are formed from the end-on overlap of atomic orbitals
S orbitals overlap this way as well as p orbitals
Forming sigma bonds

The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond
The pair of electrons is found between the nuclei of the two atoms
There is an electrostatic force of attraction between the electrons and nuclei which bonds the atoms to each other
π bonds
Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals
The two lobes that make up the π bond lie above and below the plane of the σ bond
This maximises the overlap of the p orbitals
A single π bond is drawn as two electron clouds one arising from each lobe of the p orbitals
The two clouds of electrons in a π bond represent one bond containing two electrons
Forming pi bonds

Examples of sigma & pi bonds
Hydrogen
The hydrogen atom has only one s orbital
The s orbitals of the two hydrogen atoms will overlap to form a σ bond
Sigma bonding in hydrogen

Ethene
Each carbon atom uses three of its four electrons to form σ bonds
Two σ bonds are formed with the hydrogen atoms
One σ bond is formed with the other carbon atom
The fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a π bond
This means that the C-C is a double bond: one σ and one π bond
Pi bonding in ethene

Sigma and pi bonding in ethene

Ethyne
This molecule contains a triple bond formed from two π bonds (at right angles to each other) and one σ bond
Each carbon atom uses two of its four electrons to form σ bonds
One σ bond is formed with the hydrogen atom
One σ bond is formed with the other carbon atom
Two electrons are used to form two π bonds with the other carbon atom
Sigma and pi bonding in ethyne

Hydrogen cyanide
Hydrogen cyanide contains a triple bond
One σ bond is formed between the H and C atom (overlap of an sp C hybridised orbital with the 1s H orbital)
A second σ bond is formed between the C and N atom (overlap of an sp C hybridised orbital with an sp orbital of N)
The remaining two sets of p orbitals of nitrogen and carbon will overlap to form two π bonds at right angles to each other
Sigma and pi bonding in hydrogen cyanide

Nitrogen
Nitrogen too contains a triple bond
The triple bond is formed from the overlap of the sp orbitals on each N to form a σ bond and the overlap of two sets of p orbitals on the nitrogen atoms to form two π bonds
These π bonds are at right angles to each other
Sigma and pi bonding in nitrogen molecules

Hybridisation
The p atomic orbitals can also overlap end-on to form σ bonds
In order for them to do this, they first need to become modified in order to gain s orbital character
The orbitals are therefore slightly changed in shape to make one of the p orbital lobes bigger
This mixing of atomic orbitals to form covalent bonds is called hybridisation
Mixing one s orbital with three p orbitals is called sp3 hybridisation (each orbital has ¼ s character and ¾ p character)
Mixing one s orbital with two p orbitals is called sp2 hybridisation
Mixing one s orbital with one p orbital forms sp hybridised orbitals
sp hybridisation

sp hybrid orbitals

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