Overview
This lecture covers the fundamental concepts of atoms, the periodic table, chemical bonding, states of matter, mixtures, chemical reactions, energetics, acids and bases, redox reactions, and quantum numbers.
Structure of Atoms
- Atoms are made of a nucleus (protons and neutrons) and electrons in shells.
- The number of protons defines the element.
- Isotopes have different numbers of neutrons; some are unstable and radioactive.
- Atoms with unequal numbers of protons and electrons are ions; cations (+), anions (−).
Periodic Table Organization
- Elements are listed by increasing proton number.
- Columns (groups) share the same number of valence electrons; rows (periods) share the number of shells.
- Metals, non-metals, and semimetals are distinct categories on the periodic table.
- Alkali metals have one valence electron, are shiny and soft.
Chemical Bonds and Molecules
- Valence electrons determine chemical behavior.
- Molecules = two or more bonded atoms; compounds = at least two different elements.
- Lewis Dot Structures represent valence electrons.
- Atoms bond to achieve full outer shells (octet/duet rule).
- Covalent bonds: electrons shared; ionic bonds: electrons transferred; metallic bonds: delocalized electrons.
- Electronegativity measures atom's pull on electrons; increases from bottom left to top right on the table.
- Polar covalent bonds have uneven electron sharing; nonpolar shares evenly.
- Intermolecular Forces: hydrogen bonds (strong dipole), Van der Waals forces (temporary dipoles).
Properties of Matter
- Solids: fixed structure; liquids: move freely but fixed volume; gases: particles move freely and fill volume.
- Temperature = average kinetic energy; entropy = disorder.
- Plasma: ionized gas at high temperatures or voltages.
Mixtures and Pure Substances
- Pure substances: one element or compound.
- Mixtures: homogeneous (solutions), heterogeneous (suspensions), colloids/emulsions (e.g., milk).
Chemical Reactions and Stoichiometry
- Types: synthesis, decomposition, single/double replacement.
- Reactions occur to lower energy and reach stability.
- Stoichiometry: reactants/reactants combine in specific ratios based on conservation of mass.
- Balance equations by trial and error, typically metals first.
Energetics and Thermodynamics
- Activation energy starts reactions; catalysts lower it without being used up.
- Enthalpy: total energy/heat content.
- Exothermic: releases heat; Endothermic: absorbs heat.
- Gibbs Free Energy (ΔG) determines reaction spontaneity (ΔG < 0 is spontaneous).
Acids, Bases, and pH
- Brønsted-Lowry: acids donate protons (H⁺), bases accept them.
- Strong acids dissociate fully; weak acids do not.
- pH = −log[H₃O⁺]; pH + pOH = 14; below 7 is acidic, above is basic.
Redox Reactions
- Redox (reduction-oxidation) involves electron transfer.
- Oxidation number rules: H (+1), O (−2), halogens (−1), elements (0).
- Balance charges and atoms in redox equations.
Quantum Numbers and Electron Configuration
- Four quantum numbers: n (shell), l (subshell), ml (orbital orientation), ms (spin).
- Orbitals: s (2 e⁻), p (6), d (10), f (14).
- Aufbau principle: fill orbitals in a set order for electron configuration.
- Valence electrons determined from electron configuration.
Key Terms & Definitions
- Atom — smallest unit of matter, made of nucleus and electrons.
- Isotope — atoms of an element with different neutrons.
- Ion — charged atom (cation +, anion −).
- Valence electron — electron in the outermost shell.
- Molecule — two or more atoms bonded together.
- Compound — molecule with different elements.
- Stoichiometry — calculation of reactant/product ratios in reactions.
- Enthalpy (ΔH) — total heat content of a system.
- Gibbs Free Energy (ΔG) — energy determining spontaneity of processes.
- pH — scale of acidity based on hydronium ion concentration.
- Quantum number — numbers describing electron properties.
Action Items / Next Steps
- Practice writing and balancing chemical equations.
- Review periodic table groups and electron configurations.
- Complete assigned readings on chemical bonding and quantum mechanics.