Sigma & Pi Bonds (HL) (DP IB Chemistry): Revision Note

Sigma & pi bonds

Bond overlap in covalent bonds

• A single covalent bond forms when two non-metal atoms each contribute an unpaired electron

• These unpaired electrons occupy atomic orbitals that overlap to form a molecular orbital

  • This molecular orbital contains a shared pair of electrons

  • The shape of the molecular orbital depends on the types of atomic orbitals involved

• The greater the overlap, the stronger the covalent bond

• There are two main types of covalent bond formed by orbital overlap:

  • Sigma (σ) bonds

  • Pi (π) bonds

What is a sigma bond?

• Sigma (σ) bonds are formed by the head-on overlap of atomic orbitals

• The electron density is concentrated along the bond axis

  • The bond axis is an imaginary line between the two nuclei

• Sigma bonds:

  • Are the strongest type of covalent bond

  • Are always present in single covalent bonds

  • Are also found in double and triple bonds

    • In these cases, the sigma bond is accompanied by one or more pi (π) bonds

    • The sigma bond remains the strongest component of the multiple bond

• Sigma bonds can form from:

  • s–s orbital overlap

  • s–p orbital overlap

  • p–p orbital overlap

Sigma bonds from s orbitals - hydrogen

• Each hydrogen atom has a 1s orbital with one unpaired electron

• The two 1s orbitals overlap directly to form a sigma (σ) bond

Sigma bonds from an s and a p orbital - hydrogen fluoride

• The 1s orbital of hydrogen overlaps with 2p orbital of fluorine

• This head-on overlap forms a sigma (σ) bond

Sigma bonds from p orbitals - fluorine

• Each fluorine atom has an unpaired electron in a p orbital

• The p orbitals overlap head-on to form a sigma (σ) bond

What is a pi bond?

• Pi (π) bonds form when adjacent p orbitals overlap sideways (laterally)

• This creates electron density above and below the plane of the sigma (σ) bond

• A single π bond is shown as two electron clouds

  • Each electron cloud comes from one lobe of the overlapping p orbitals

• These clouds together contain two electrons shared between the atoms

• The electron density lies on opposite sides of the bond axis

• π bonds occur only in double and triple bonds

Pi bonds from p orbitals

Examples of sigma & pi bonding in molecules

Methane

• The carbon atom forms four sigma (σ) bonds with hydrogen atoms

• These bonds are formed by the overlap of the hybrid orbitals on carbon with the 1s orbitals of hydrogen

• Each bond involves a head-on overlap, creating a sigma bond

• Methane contains only sigma bonds

Ethene

• Sigma bonds in ethene:

• Each carbon atom in ethene is hybridised

• It forms three sigma (σ) bonds using three hybrid orbitals:

  • Two σ bonds are formed with the hydrogen atoms

  • One σ bond is formed with the other carbon atom

• Pi bonds in ethene:

• The fourth electron from each carbon atom occupies a p orbital

  • This overlaps sideways with another p orbital on the other carbon atom to form a π bond

• This creates a carbon-carbon double bond consisting of:

  • One σ bond

  • One π bond

Ethyne

• Sigma bonds in ethyne:

• Each carbon atom in ethyne is hybridised

• It forms two sigma (σ) bonds using its two hybrid orbitals:

  • One σ bond is formed with a hydrogen atom

  • One σ bond is formed with the other carbon atom

• Pi bonds in ethyne:

• Each carbon atom has two unhybridised p orbitals

  • These overlap sideways with p orbitals on the other carbon atom

  • This forms two perpendicular π bonds

• This creates a carbon-carbon triple bond consisting of:

  • One σ bond

  • Two π bonds

Predicting the type of bonds

• The number and type of sigma (σ) and pi (π) bonds can be deduced by analysing the bonding in a molecule