Hybridisation in Organic Molecules (CIE AS Chemistry)

• Each carbon atom has four electrons in its outer shell (electronic configuration: 1s²2s²2p²)
• Carbon atoms share these four electrons in four covalent bonds with other atoms to achieve a full outer shell configuration
• These electrons are found in orbitals within the respective atoms
• When forming a covalent bond, the orbitals overlap in such a way to form two types of bonds
  • Sigma bonds (σ)
  • Pi bonds (π)

Hybridisation: sp³

• The electron pair in a σ bond is found in a region of space between the nuclei of the two atoms that are sharing the electrons
• The electrostatic attraction between the electrons (negatively charged) and the two nuclei (positively charged) holds the two atoms together
• Carbon atoms that form four σ bonds are said to be sp³ hybridised
• The four pairs of electrons around each carbon repel each other forcing the molecule to adopt a configuration in which the bonding pairs of electrons are as far away from each other as possible
  • The molecule adopts a tetrahedral arrangement with bond angles of 109.5 ^o

 Bonding in ethane

The diagram shows a molecule of ethane in which each carbon atom forms four σ bonds to adopt a tetrahedral configuration and minimise the repulsion between the bonding pairs of electrons

Hybridisation: sp²

• When carbon atoms use only three of their electron pairs to form a σ bond, they are said to be sp² hybridised
  • Each carbon atom will have a p orbital with contains one spare electron

• When the p orbitals of two carbon atoms overlap with each other, a π bond is formed (the π bond contains two electrons)
• The two orbitals that form the π bond lie above and below the plane of the two carbon atoms to maximise bond overlap
• The three bonding pairs of electrons are in the plane of the molecule and repel each other
• The molecule adopts a planar arrangement with bond angles of 120 ^o

 Bonding in ethene

The overlap of the two p orbitals results in the formation of a π bond in ethene (sp² hybridised molecule) in which the bonding pair of electrons repel each other to force the molecule into a planar configuration with bond angles of 120^o

Hybridisation: sp

• Carbon atoms can also use only one of their electron pair to form a σ bond, in which case the carbon atoms are said to be sp hybridised
  • Each carbon atom will have two p orbitals with one spare electron each

• When the four p orbitals of the carbon atoms overlap with each other, two π bonds are formed (each π bond contains two electrons)
• The two orbitals that form the π bond lie above and below the plane of the carbon atoms
• The two orbitals of the other π bond lie in front and behind the plane of the atoms
  • This maximises the overlap of the four p orbitals

• The molecule adopts a linear arrangement with bond angles 180 ^o

Bonding in ethyne

The overlap of the p orbitals results in the formation of two π bonds in ethyne (sp hybridised molecule) which adopts a linear arrangement with bond angles of 180^o

σ bonds

• Sigma bonds are formed from the end-on overlap of atomic orbitals
  • S orbitals overlap this way as well as p orbitals

Forming sigma bonds

Sigma bonds can be formed from the end-on overlap of s or p orbitals

• 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
  • The electrostatic attraction between the electrons and nuclei bonds the atoms to each other

The arrangement of the σ bond in sp³, sp² and sp hybridised carbon atoms

 

The σ orbitals are formed from the end-on overlap of the atomic orbitals resulting in symmetrical electron density on the atoms

π bonds

• Pi (π) bonds are formed from the sideways overlap of p orbitals
• The two lobes that make up the π bond lie above and below the plane of the atoms
  • This maximises the overlap of the p orbitals

 Forming pi bonds 

π bonds can be formed from the end-on overlap of p orbitals

• In triple bonds, there is an additional overlap of p orbital
• The two lobes of the π bond lie in front of and behind the plane of the atoms in the molecule
  • This maximises the overlap of the p orbitals

The arrangement of the π bond in sp³, sp² and sp hybridised carbon atoms

 

The π bonds are formed from the sideway overlap of the atomic orbitals