Overview
This lecture explains the three main types of chemical bonds: ionic (electrovalent), covalent, and coordinate bonds, their formation, properties, and related examples, including electron dot structures and key distinctions.
Types of Chemical Bonding
- Chemical bond is a force that connects atoms, elements, or molecules.
- The three main types of bonds: Ionic (Electrovalent), Covalent, and Coordinate.
Ionic (Electrovalent) Bond
- Forms between a metal (loses electrons) and a non-metal (gains electrons).
- Involves complete transfer of electrons from metal to non-metal.
- Example: Na (281) loses one electron, becomes Na+, Cl (287) gains one, becomes Cl–, forming NaCl.
- Ionic bond results from electrostatic attraction between cations (+) and anions (–).
- The formation involves oxidation (loss of electrons) for metals and reduction (gain of electrons) for non-metals.
- Bond formation is a redox (oxidation-reduction) reaction.
- Properties:
- Strong electrostatic forces.
- Solid at room temperature due to strong attraction.
- High melting and boiling points.
- Conduct electricity in molten or aqueous state, not in solid.
- Soluble in water (polar solvent), insoluble in organic solvents.
Covalent Bond
- Forms between two non-metals via mutual sharing of electrons.
- No ions are produced; the bond involves molecules only.
- Types:
- Single bond: One pair shared (e.g., H₂).
- Double bond: Two pairs shared (e.g., O₂).
- Triple bond: Three pairs shared (e.g., N₂).
- Structures can be represented using Lewis electron dot structures.
- Examples: CH₄, CCl₄, H₂O, NH₃, with respective dot structures.
- Properties:
- Weak Van der Waals (intermolecular) forces.
- Often gases, liquids, or soft solids.
- Low melting and boiling points.
- Generally poor conductors of electricity.
- Usually insoluble in water (unless polar).
Polar and Non-polar Covalent Compounds
- Non-polar: Equal sharing, no charges, typically insoluble in water (e.g., H₂, O₂, CCl₄).
- Polar: Unequal sharing due to electronegativity differences (e.g., HCl, NH₃, H₂O); display partial charges (δ+ and δ–), can dissolve in water, and may conduct electricity.
- Geometry can affect polarity (e.g., CCl₄ non-polar despite polar bonds; CHCl₃ is polar).
Coordinate (Dative) Bond
- Formed when both electrons in a shared pair come from the same atom (the donor) to an acceptor atom.
- Example: Hydronium ion (H₃O⁺) and ammonium ion (NH₄⁺), where O or N donates a lone pair to H⁺.
- Arrow in the structure points from donor to acceptor.
- Coordinate bonds exist alongside covalent bonds within the same species.
Key Terms & Definitions
- Ionic Bond (Electrovalent Bond) — Complete transfer of electrons from metal to non-metal, forming cations and anions.
- Covalent Bond — Mutual sharing of one or more pairs of electrons between two non-metals.
- Coordinate Bond — Shared pair of electrons both supplied by one atom to another atom.
- Redox Reaction — Simultaneous occurrence of oxidation and reduction during bond formation.
- Electronegativity — Tendency of an atom to attract shared electrons in a bond.
- Lone Pair — A pair of valence electrons not involved in bond formation.
- Dielectric Property — Water's ability to reduce electrostatic force between ions.
- Polar Solvent — Solvent that dissolves ionic and polar compounds due to its charge separation.
Action Items / Next Steps
- Review electron dot structures for examples (NaCl, MgCl₂, MgO, CH₄, CCl₄, H₂O, NH₃, H₃O⁺, NH₄⁺).
- Study the difference between ionic, covalent, and coordinate bonds.
- Understand the role of polarity and electronegativity.
- Complete readings on redox reactions, electrolysis, and properties of acids/bases as referenced.