Basics of Covalent Bonding

Study Notes

Covalent bonds form when atoms share electron pairs to achieve a stable electron configuration.
This bonding typically occurs between non-metal atoms.
Atoms strive to achieve a stable electron configuration, resembling noble gases with filled valence electron shells.
The shared electron pair creates a bond that holds the atoms together.
The number of shared electron pairs determines the bond type:
- Single Bond: one shared pair
- Double Bond: two shared pairs
- Triple Bond: three shared pairs

Bond Strength: Triple bonds are stronger than double bonds, which are stronger than single bonds.
Bond Length: Triple bonds are shorter than double bonds, which are shorter than single bonds. Shorter bonds are typically stronger.
Molecular Shape: The Valence Shell Electron Pair Repulsion (VSEPR) theory explains molecular shapes.
Electron pairs arrange themselves to minimize repulsion, leading to specific molecular geometries.
- Linear: Two bonding pairs, no lone pairs (e.g., carbon dioxide (CO₂)).
- Trigonal Planar: Three bonding pairs, no lone pairs (e.g., boron trifluoride (BF₃)).
- Tetrahedral: Four bonding pairs, no lone pairs (e.g., methane (CH₄)).
- Bent: Two bonding pairs and one or two lone pairs (e.g., water (H₂O)).

Nonpolar Covalent Bonds: Identical non-metal atoms share electrons equally. Minimal or zero electronegativity difference. (e.g., diatomic hydrogen (H₂)).
Polar Covalent Bonds: Different non-metal atoms share electrons unequally due to electronegativity differences. (e.g., water (H₂O)).
Coordinate Covalent Bonds: One atom donates both electrons in a shared pair. (e.g., ammonium ion (NH₄⁺)).

Organic Molecules: Contain carbon and hydrogen atoms bonded covalently.
- Hydrocarbons: Methane (CH₄), ethane (C₂H₆), ethene (C₂H₄)
- Functional Groups: Hydroxyl (-OH), carboxyl (-COOH)
Inorganic Molecules:
- Water (H₂O): Polar molecule with covalent bonds between hydrogen and oxygen.
- Carbon Dioxide (CO₂): Linear molecule with double covalent bonds between carbon and oxygen atoms.

Pharmaceuticals: Covalent bonding is essential in drug design. Understanding how drugs bind to receptors leads to new medications. (e.g., Aspirin (acetylsalicylic acid)).
Materials Science: Covalent bonds are fundamental in polymers (e.g., plastics and fibers). (e.g., Polyethylene).
Environmental Chemistry: Covalent bonds are involved in water purification processes (e.g., activated carbon in water filters).

Resonance Structures: Multiple Lewis structures can describe a molecule's bonding due to electron delocalization. (e.g., nitrate ion (NO₃⁻)).
Hybridization: Atomic orbitals mix to form new hybrid orbitals used to create covalent bonds. (e.g., methane (CH₄) carbon undergoes sp³ hybridization).