Covalent Bonding

Single Covalent Bond: Formed when two atoms share exactly one pair of electrons (2 electrons total). It is represented in structural formulas by a single line (e.g., $H-H$).

Double Covalent Bond: Occurs when two atoms share two pairs of electrons (4 electrons total) to satisfy their valency. This is common in molecules like Oxygen ($O=O$) and is stronger and shorter than a single bond.

Triple Covalent Bond: Involves the sharing of three pairs of electrons (6 electrons total) between two atoms. These bonds, such as in Nitrogen ($N \equiv N$), are exceptionally strong and require significant energy to break.

Visualizing Valence: Dot and cross diagrams are used to track the origin of electrons in a bond. One atom's electrons are represented by dots ($\cdot$), while the other's are represented by crosses ($imes$).

Outer Shell Focus: These diagrams typically only show the electrons in the highest energy level (valence shell), as these are the only ones involved in chemical bonding.

Lone Pairs: Electrons in the outer shell that are not involved in bonding are called lone pairs. These must be included in diagrams as they influence the molecule's shape and reactivity.

Electron Deficiency: Some molecules contain central atoms that are stable with fewer than eight electrons. For example, Boron often forms compounds where it is surrounded by only six valence electrons.

Expanded Octet: Elements in Period 3 and below have access to $d$-orbitals, allowing them to accommodate more than eight electrons in their valence shell. This is seen in molecules where a central atom like Phosphorus or Sulfur bonds with five or six other atoms.

Odd-Electron Molecules: Rare cases exist where a molecule has an odd number of total valence electrons, meaning at least one atom cannot have a complete octet (e.g., Nitrogen Monoxide).

Check the Valency: Always verify the group number of an element to determine its initial valence electrons. For instance, Group 14 elements like Carbon need four more electrons, meaning they will typically form four covalent bonds.

Account for All Electrons: A common mistake is forgetting to draw lone pairs in dot-and-cross diagrams. Ensure the total number of dots and crosses matches the sum of valence electrons from all participating atoms.

Identify Exceptions: If the central atom is from Period 3 (like $S$ or $P$), be prepared for an expanded octet. Do not force these atoms to follow the octet rule if the molecular formula suggests more bonds are formed.

Verify Stability: After drawing a structure, count the electrons around each atom. Except for known exceptions, every atom should be surrounded by 8 electrons (or 2 for Hydrogen).