Overview
This lecture covers the structure, naming, and properties of ionic and molecular compounds, focusing on electron transfer, charge balance, naming rules, and 3D molecular geometry.
Ionic Bonds and Ions
- Ionic bonds involve the complete transfer of electrons from one atom (typically s-block metal) to another (typically p-block nonmetal).
- S-block elements (columns 1 & 2) usually form positive ions (cations); p-block elements (columns 5–7) form negative ions (anions).
- Atoms become ions when they lose or gain electrons to achieve a stable electron configuration (usually like a noble gas).
Determining Charges and Chemical Formulas
- Group 1 metals form +1, group 2 metals form +2, Aluminum (group 3) forms +3, group 5 nonmetals form –3, group 6 form –2, and group 7 halogens form –1 ions.
- Chemical formulas for ionic compounds must balance positive and negative charges to have an overall charge of zero (e.g., NaCl, MgCl₂).
- The lowest common denominator method ensures charge balance when combining ions (e.g., Mg₃P₂).
Naming Ionic Compounds
- Name the cation (metal) first with its element name; name the anion (nonmetal) with the suffix “-ide”.
- Transition metals (d-block) with variable charges specify the charge in Roman numerals (e.g., copper(II) chloride for CuCl₂).
- Polyatomic ions are named as listed in standard tables; their charges must be accounted for in formulas (e.g., Mg(NO₃)₂).
Polyatomic Ions
- Polyatomic ions are groups of atoms with a net charge (e.g., SO₄²⁻, NO₃⁻).
- The names and charges for polyatomic ions are provided in reference tables; memorize common ones as needed.
Electronegativity and Periodic Trends
- Electronegativity is the tendency of an atom to attract electrons; it increases up a group and across a period (fluorine is the most electronegative).
- S-block metals have low electronegativity and lose electrons easily; p-block nonmetals have high electronegativity and gain electrons.
Intermolecular Forces
- Intermolecular forces occur between molecules: dipole-dipole (polar molecules), hydrogen bonding (H bonded to N, O, or F), and dispersion forces (nonpolar molecules).
- Hydrogen bonding is the strongest, followed by dipole-dipole, then dispersion (weakest).
Molecular Compounds and Naming
- Molecular (covalent) compounds involve sharing electrons, usually between two p-block nonmetals.
- Use prefixes to indicate the number of each atom: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa- (e.g., carbon dioxide: CO₂).
Lewis Structures and Molecular Geometry
- Draw the least electronegative atom in the center; total valence electron count = sum of valence electrons from all atoms.
- Distribute electrons to satisfy the octet rule (8 electrons around C, N, O, F; exceptions: H—2, Be—4, B/Al—6).
- Formal charge = (group number) – (non-bonding electrons + ½ bonding electrons); aim for structures with formal charges close to zero.
- The AXE method (A = central atom, X = bonded atoms, E = lone pairs) is used for determining 3D shapes (trigonal pyramidal, linear, etc.).
Key Terms & Definitions
- Ionic Bond — Transfer of electrons from metal to nonmetal.
- Covalent Bond — Sharing of electrons between nonmetals.
- Cation — Positively charged ion.
- Anion — Negatively charged ion.
- Polyatomic Ion — A charged group of covalently bonded atoms.
- Electronegativity — The ability of an atom to attract electrons.
- Formal Charge — Difference between valence electrons and those assigned in a structure.
- Octet Rule — Atoms tend to have 8 electrons in their valence shell.
- Lewis Structure — Representation showing bonds and lone pairs in a molecule.
- Intermolecular Forces — Attractions between molecules (dipole-dipole, hydrogen bond, dispersion).
- AXE Method — System to predict molecular shape (A = central, X = surrounding, E = lone pairs).
Action Items / Next Steps
- Practice naming and writing formulas for ionic and molecular compounds.
- Reference polyatomic ion tables for homework and exams.
- Use MoleView or similar tools to visualize molecular shapes in 3D.
- Complete assigned textbook problems on Lewis structures, formal charge, and molecular geometry.