Covalent Bonds

I. The Purpose of Covalent Bonding 

By sharing electrons, atoms of a molecule or lattice can acquire more stable electron configurations. 

Example: Two hydrogen atoms and one oxygen atom join to form H₂O. Each hydrogen atom has one electron, but it needs one more to fill its outer shell. The oxygen atom has six electrons in its outer shell and needs two more to be stable. To solve this, the oxygen atom shares one of its electrons with each hydrogen atom, while each hydrogen atom shares its electron with oxygen. This sharing of electrons establishes covalent bonds and makes all the atoms more stable. 

Note: Each bond contains two electrons. Single (1), double (2), and triple (3) bonds exist in covalent bonding. 

II. The Visualization of Covalent Bonding 

Introduction to Orbital Overlap:

Recall orbitals from electron configuration. In covalent bonds, orbitals from different atoms overlap to create a region of shared electron density between them. 

• A direct overlap of orbitals right in between of two atoms is called a σ-bond. All covalent bonds include σ-bonds. 

• An indirect, perpendicular overlap of orbitals above and below the space between two atoms is called a π-bond. The second and third covalent bonds between certain atoms consist of π-bonds. 

Bonding with Atomic Orbitals:

This section refers to the use of native electron configurations of atoms to form covalent bonds. 

Essentially, unpaired orbitals of atoms overlap to create bonds. Two electrons per bond, one originating from each orbital. If more sharing is needed to achieve stable configurations, then more bonds (double or triple) will form.