Overview
This lecture covers key concepts and types of chemical bonding—ionic, covalent, and metallic—as well as molecular shapes, intermolecular forces, and solubility, tailored to the Edexcel specification.
Ionic Bonding
- Ionic bonds form between oppositely charged ions due to electrostatic attraction.
- Ions are formed by electron loss (positive ion) or gain (negative ion) to achieve full outer shells.
- Group 1 forms +1, Group 2 +2, Group 6 –2, Group 7 –1 ions; polyatomic ions include OH–, NO3–, NH4+, SO42–, and CO32–.
- "Swap and drop" method is used to determine ionic compound formulas.
- Ionic compounds have giant lattice structures, high melting points, are brittle, soluble in water, and conduct electricity when molten or in solution.
- Higher charge and smaller ionic radius increase bond strength and melting point (charge density).
- Isoelectronic ions: same electron number, ionic radius decreases with increasing proton number.
Evidence for Charged Particles
- Electrolysis of colored salts (e.g., copper chromate) demonstrates movement of ions toward electrodes, proving ions are charged.
Covalent Bonding
- Covalent bonds involve sharing electrons to achieve full outer shells.
- Single, double, and triple covalent bonds differ by the number of shared electrons.
- Dative (coordinate) bonds: both bonding electrons come from one atom.
- Bond enthalpy increases as bond length decreases; triple bonds are strongest/shortest.
Molecular Shape and VSEPR
- Shape depends on number of bonding and lone electron pairs (VSEPR theory).
- Lone pairs repel more, reducing bond angles.
- Key geometries: linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal bipyramidal (90°,120°), octahedral (90°).
- Lone pairs modify shapes: trigonal pyramidal (107°), bent (104.5°), see-saw, T-shaped, square planar, etc.
Giant Covalent Structures
- Diamond: each C bonded to 4 others, very hard, high melting point, does not conduct electricity.
- Graphite: each C bonded to 3 others, layers slide, conducts electricity via delocalized electrons.
- Graphene: single graphite layer, strong, lightweight, conducts electricity.
- Silicon dioxide has a diamond-like giant covalent structure.
Metallic Bonding
- Metal atoms form positive ions in a lattice with delocalized electrons ("sea of electrons").
- Metals conduct heat and electricity, are malleable and ductile, and have high melting points.
Electronegativity and Polarity
- Electronegativity: an atom's ability to attract electrons in a covalent bond; increases up and right on the periodic table (excluding noble gases).
- Large differences in electronegativity yield ionic bonds; small/no difference means covalent.
- Polar covalent bonds have unequal electron sharing, causing dipoles (δ+ and δ–).
Intermolecular Forces
- London forces: temporary dipole-induced dipole, present in all molecules, weakest.
- Permanent dipole-dipole: between molecules with permanent dipoles, stronger than London.
- Hydrogen bonding: occurs when H is bonded to N, O, or F—strongest intermolecular force.
Solubility
- Polar substances dissolve in polar solvents; nonpolar dissolve best in nonpolar solvents.
- Ionic compounds dissolve in water via hydration; exceptions exist (e.g., Al2O3).
- Alcohols dissolve in water due to hydrogen bonding; larger hydrocarbon chains decrease solubility.
- Haloalkanes do not dissolve in water due to weak dipoles.
Summary Table
- Giant covalent: solid, high melting, not conductive (except graphite/graphene), insoluble.
- Simple molecular: low melting/boiling point, usually gases/liquids, non-conductive, solubility varies.
- Giant ionic: solid, high melting, conductive when molten/solution, soluble in water.
- Metallic: solid, high melting, conductive, not soluble.
Key Terms & Definitions
- Ion — an atom or molecule with a net electric charge.
- Electronegativity — ability of an atom to attract electrons in a bond.
- Ionic bond — electrostatic attraction between oppositely charged ions.
- Covalent bond — sharing of electron pairs between atoms.
- Dative/coordinate bond — covalent bond where both electrons originate from one atom.
- Lone pair — a pair of valence electrons not involved in bonding.
- VSEPR — Valence Shell Electron Pair Repulsion theory for molecular shapes.
- London force — weak intermolecular attraction due to temporary dipoles.
- Hydrogen bond — strong attraction involving H and N, O, or F.
- Hydration — process where ions are surrounded by water molecules during dissolution.
- Charge density — ratio of an ion’s charge to its volume, affects bond strength.
Action Items / Next Steps
- Memorize molecular shapes, angles, and VSEPR rules.
- Learn key ion formulas and charges.
- Review hydrogen bonding criteria and examples.
- Review the provided summary table for comparative properties.
- Practice applying the swap and drop method for ionic formulas.