Overview
This lecture covers atomic structure, the periodic table, chemical bonding, states of matter, chemical reactions, quantum numbers, and key chemistry concepts essential for understanding molecular behavior.
Atomic Structure and Elements
- Everything is made of atoms, which have a nucleus (protons and neutrons) and electrons.
- The number of protons determines the element; different neutron counts create isotopes.
- Atoms with equal protons and electrons are neutral; more or less electrons create ions (anions or cations).
The Periodic Table
- Elements are organized by increasing proton number; groups have the same number of valence electrons and similar chemical properties.
- Periods (rows) indicate the number of electron shells.
- The table is divided into metals, nonmetals, and semimetals.
- Each element’s cell lists name, symbol, proton count, and atomic mass.
Chemical Bonds and Compounds
- Atoms bond to achieve full outer shells (octet rule).
- Covalent bonds share electrons; ionic bonds transfer electrons (metal + nonmetal); metallic bonds involve delocalized electrons.
- Bond type depends on electronegativity difference: ionic (>1.7), polar covalent (0.5–1.7), nonpolar covalent (<0.5).
- Intermolecular forces (IMFs) include hydrogen bonds and Van der Waals forces.
- Isomers have the same formula but different structures.
- Lewis-dot structures represent valence electrons and bonds.
States of Matter and Solutions
- Solid: fixed structure; Liquid: particles move within fixed volume; Gas: particles move freely and fill space.
- Plasma: ionized gas at high temperature or voltage.
- Mixtures: homogeneous (solutions) or heterogeneous (suspensions); colloids are intermediate (e.g., milk).
Chemical Reactions and Stoichiometry
- Types: synthesis, decomposition, single/double replacement; all aim for lower energy/more stability.
- Stoichiometry ensures reactants combine in specific ratios; law of conservation of mass.
- 1 mole = atomic mass in grams; use for balancing reaction equations.
- Physical changes alter state/appearance; chemical changes create new substances.
Thermodynamics and Equilibrium
- Activation energy is needed for reactions; catalysts lower this energy without being consumed.
- Enthalpy: heat content; exothermic releases heat, endothermic absorbs heat.
- Gibbs free energy predicts spontaneity (ΔG < 0 exergonic/spontaneous).
- Equilibrium is when forward and reverse reactions occur at the same rate.
Acids, Bases, and Redox
- Acids donate protons; bases accept protons (Brønsted-Lowry).
- pH = -log[H₃O⁺]; pH < 7 acidic, pH = 7 neutral, pH > 7 basic; pH + pOH = 14.
- Amphoteric substances act as acids or bases.
- Redox reactions involve electron transfer and changes in oxidation numbers.
Quantum Mechanics and Electron Configuration
- Electrons are organized by quantum numbers: n (shell), l (subshell), ml (orbital), ms (spin).
- Subshell types: s, p, d, f; shells fill in order (Aufbau principle).
- Each orbital holds 2 electrons (opposite spin); configuration predicts chemical behavior.
Key Terms & Definitions
- Atom — Smallest unit of matter, with a nucleus and electrons.
- Isotope — Atoms of the same element different in neutron number.
- Ion — Charged atom (anion = negative, cation = positive).
- Molecule — Two or more atoms bonded together.
- Compound — Molecule with at least two different elements.
- Valence Electrons — Electrons in the outermost shell.
- Stoichiometry — The calculation of reactant/product amounts in reactions.
- Enthalpy — Heat content of a system.
- Gibbs Free Energy (ΔG) — Determines reaction spontaneity.
- pH — Scale for hydronium ion concentration.
- Redox Reaction — Chemical reaction involving electron transfer.
- Quantum Numbers — Describe electron position and energy in an atom.
Action Items / Next Steps
- Practice drawing Lewis structures for common molecules.
- Review balancing chemical equations using stoichiometry.
- Learn to assign oxidation numbers and predict redox changes.
- Memorize periodic table groups and their properties.