Overview
This lecture covers the structure of atoms, chemical bonding, the periodic table, types of matter, chemical reactions, energetics, acids and bases, redox reactions, and atomic orbitals.
Atomic Structure & Elements
- All matter is made of atoms; atoms have a nucleus (protons and neutrons) and electrons.
- The number of protons determines the element; different neutron numbers create isotopes.
- Outermost electrons are "valence electrons" and drive chemical behavior.
- Atoms with equal protons and electrons are neutral; more electrons = anion, fewer = cation.
The Periodic Table
- Elements are arranged by increasing protons; rows = periods (same number of electron shells), columns = groups (same valence electrons).
- Group 1 (alkali metals) are soft, shiny, reactive metals with one valence electron.
- Transition metals’ electron patterns are complex.
- Metals are left of the periodic line, nonmetals right, semimetals on the line.
Molecules, Compounds, and Bond Types
- Two or more atoms bonded = molecule; different elements = compound.
- Molecular formulas count atoms, but isomers (same formula, different structure) exist.
- Lewis-Dot Structures show valence electrons and bonds.
- Atoms bond to achieve a full outer shell (usually 8 electrons).
Chemical Bonds & Forces
- Covalent bonds: electrons shared between atoms.
- Electronegativity: atom's pull on shared electrons; increases up and right on the periodic table.
- Ionic bonds: electron transfer between atoms with high electronegativity difference (>1.7).
- Metallic bonds: electrons delocalized among metal atoms.
- Polar covalent bonds: unequal sharing (difference 0.5–1.7); nonpolar: nearly equal sharing (<0.5).
- Hydrogen bonds: between hydrogen and highly electronegative atoms (F, O, N).
- Van der Waals forces: temporary dipoles from moving electrons.
- Intermolecular forces affect properties like solubility and melting points.
States of Matter & Physical Properties
- Solids: fixed structure, low entropy; liquids: particles move freely, fixed volume; gases: fill space, high entropy.
- Temperature = average kinetic energy; entropy = disorder.
- Strong bonds yield high melting points.
Mixtures & Solutions
- Pure substances: single element or compound; mixtures: blend of substances.
- Homogeneous mixtures (solutions) are uniform; heterogeneous mixtures have distinct regions.
- Colloids (emulsions) have medium-sized particles, like milk.
Chemical Reactions & Stoichiometry
- Types: synthesis, decomposition, single/double replacement.
- Reactions seek lower energy, proceed in specific ratios (stoichiometry).
- Mass is conserved; equations must be balanced for each element.
- The mole is the amount containing Avogadro's number (count) of particles.
Chemical and Physical Changes
- Physical changes alter appearance, not substance; chemical changes alter substance (e.g., bubbles, odor, energy release).
- Reactions need activation energy; catalysts lower this energy and are not consumed.
Energetics: Enthalpy, Entropy, and Spontaneity
- Enthalpy (ΔH): heat content; negative ΔH = exothermic, positive = endothermic.
- Gibbs Free Energy (ΔG) = drives spontaneity; ΔG < 0 is spontaneous.
- Entropy increase can make endothermic reactions spontaneous at high temperature.
Equilibrium, Acids, Bases & pH
- Chemical equilibrium: forward and reverse reactions at equal rates, concentrations constant.
- Brønsted-Lowry acid: proton donor; base: proton acceptor.
- Strong acids/bases dissociate fully; weak, partially.
- pH = –log[hydronium]; pH 7 is neutral, <7 acidic, >7 basic.
- pH + pOH = 14.
Redox Reactions & Oxidation Numbers
- Redox: electron transfer; oxidation = loss, reduction = gain.
- Oxidation numbers follow fixed rules (e.g., O usually –2, H +1).
- Balancing redox equations may involve adding water or ions.
Atomic Orbitals & Quantum Numbers
- Electrons described by four quantum numbers: n (shell), l (subshell), ml (orbital), ms (spin).
- Subshells: s (2e–), p (6e–), d (10e–), f (14e–).
- Aufbau principle: fill lower energy subshells before higher.
- Only two electrons (opposite spins) per orbital due to Pauli exclusion.
Key Terms & Definitions
- Atom — the basic unit of matter, made of a nucleus and electrons.
- Isotope — atoms of the same element with different neutrons.
- Valence electrons — electrons in the outermost shell.
- Ions — charged atoms; cations (+), anions (–).
- Electronegativity — tendency to attract electrons in a bond.
- Mole — amount of substance containing Avogadro’s number of particles.
- Enthalpy (ΔH) — heat content of a system.
- Entropy (ΔS) — measure of disorder in a system.
- Gibbs Free Energy (ΔG) — determines spontaneity of a reaction.
- pH — measure of hydronium ion concentration.
- Redox reaction — reaction involving electron transfer.
- Quantum number — number describing electron properties/orbitals.
Action Items / Next Steps
- Practice drawing Lewis-Dot structures for simple molecules.
- Review periodic table group and period trends.
- Complete assigned homework on balancing chemical reactions.
- Read textbook section on electron configuration and quantum numbers.