Lewis Structures, Properties, σ- & π-Bonds

Understanding Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They help us understand how atoms connect and share electrons in a molecule.

For example, the Lewis structure for water (H₂O) shows two hydrogen atoms bonded to an oxygen atom, with two lone pairs on the oxygen:

• Oxygen has 6 valence electrons, and each hydrogen has 1 valence electron.
• In H₂O, oxygen shares one electron with each hydrogen, forming two covalent bonds.

Properties of Ionic, Covalent, and Metallic Bonds

• Ionic Bonds: Formed by the transfer of electrons from one atom to another, creating ions. They typically occur between metals and non-metals. Ionic compounds have high melting and boiling points and conduct electricity when dissolved in water.
• Covalent Bonds: Formed by the sharing of electrons between atoms. They usually occur between non-metals. Covalent compounds can be gases, liquids, or solids and generally have lower melting and boiling points compared to ionic compounds.
• Metallic Bonds: Involve the sharing of free electrons among a lattice of metal atoms. These bonds give metals their characteristic properties, such as conductivity, malleability, and ductility.

Sigma (σ) and Pi (π) Bonds

Sigma and pi bonds are types of covalent bonds that differ in the way atomic orbitals overlap:

• Sigma (σ) Bonds: Formed by the head-on overlap of atomic orbitals. They are the strongest type of covalent bond and are found in all single bonds.
• Pi (π) Bonds: Formed by the side-to-side overlap of atomic orbitals. They are weaker than sigma bonds and are found in double and triple bonds, alongside a sigma bond.

Example: Ethene (C₂H₄)

In ethene, each carbon atom forms three sigma bonds: two with hydrogen atoms and one with another carbon atom. The double bond between the carbon atoms consists of one sigma bond and one pi bond.