Covalent Bonds (DP IB Chemistry: HL)

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

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

Hydrogen fluoride has sigma bonds between s and p orbitals

Fluorine has sigma bonds between 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

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 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

π orbitals are formed by the end-on overlap of p orbitals

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

Direct overlap of the 1s orbitals of the hydrogen atoms results in the formation of a σ bond