Overview
This lecture covers covalent bonding, including its principles, examples in common molecules, bond types (single, double, triple), molecular shapes, physical properties, and various giant and simple molecular structures.
Principles of Covalent Bonding
- Covalent bonding involves sharing a pair of electrons between two atoms.
- Each nucleus attracts the shared electron pair, keeping the atoms bonded.
- Covalent bonds typically form between nonmetals aiming for stability via noble gas configuration.
- Atoms may satisfy the octet rule (8 electrons) or duplet rule (2 electrons for hydrogen/helium).
Representing Covalent Bonds
- Dot-and-cross diagrams depict shared electrons and their sources but are symbolic only.
- Structural formulas use lines for bonds; always follow the format requested in exams.
Examples of Covalent Bonding
- H₂: Two hydrogen atoms share electrons, achieving helium's configuration (duplet rule).
- HCl: Hydrogen shares with chlorine; hydrogen attains helium configuration, chlorine attains argon configuration.
- CH₄ (methane): Carbon forms four single bonds with hydrogen to complete its octet.
- NH₃ (ammonia): Nitrogen forms three bonds with hydrogens, retaining one lone pair.
- H₂O (water): Oxygen forms two bonds with hydrogen, leaving two lone pairs.
Types of Covalent Bonds
- Single Bond: Sharing one pair of electrons (e.g., H₂, Cl₂, CH₄).
- Double Bond: Sharing two pairs of electrons (e.g., O₂, CO₂, C₂H₄/ethene).
- Triple Bond: Sharing three pairs of electrons (e.g., N₂).
Molecular Shapes and Polarity
- Electron pairs (bonded and lone pairs) repel, determining molecule shape (e.g., tetrahedral for CH₄, bent for H₂O).
- Water's bent shape and unequal sharing cause polarity (partial charges on atoms).
Complex and Exceptional Molecules
- Some molecules have central atoms with less (e.g., BF₃) or more than an octet (e.g., SF₆, HClO₄).
- Carbon forms four, nitrogen three, oxygen two, and hydrogen/halogens one covalent bond.
Physical Properties of Covalent Compounds
- Simple molecular structures have low melting/boiling points due to weak intermolecular forces.
- Most do not conduct electricity, as there are no free ions or electrons.
- Usually insoluble in water but often soluble in organic solvents.
Giant Covalent Structures
- Diamond: Each carbon bonds to four others in a tetrahedral lattice, very hard, high melting point, does not conduct electricity.
- Graphite: Carbon atoms form hexagonal layers with delocalized electrons, soft, good conductor of electricity, less dense than diamond.
- C₆₀ Fullerene: Spherical molecules of 60 carbons, lower melting point, does not conduct electricity, can dissolve in some solvents.
Key Terms & Definitions
- Covalent bond — a shared pair of electrons between atoms.
- Dot-and-cross diagram — illustration showing electron sharing in covalent bonds.
- Octet rule — atoms tend to have eight electrons in their outer shell for stability.
- Duplet rule — atoms (like H, He) are stable with two electrons in their outer shell.
- Lone pair — non-bonding pair of electrons on an atom.
- Bonded/shared pair — electron pair involved in a covalent bond.
- Polarity — uneven distribution of electron density, causing partial charges.
- Intermolecular force — weak attraction between molecules (not bonds within molecules).
- Giant covalent structure — extensive 3D network of bonded atoms (e.g., diamond, graphite).
- Allotrope — different forms of the same element with distinct structures.
Action Items / Next Steps
- Practice drawing dot-and-cross diagrams for H₂O, NH₃, CH₄, CO₂, O₂, N₂, BF₃, SF₆.
- Memorize typical bonding numbers for C, N, O, H, and halogens.
- Review differences between simple molecular and giant covalent structures.
- Read up on physical properties tied to structure and bonding.