Overview
This lecture covers atomic structure, the periodic table, types of bonds, properties of matter, chemical reactions, and fundamental quantum concepts in chemistry.
Atomic Structure & Elements
- Atoms are made of a core (protons and neutrons) and electrons in shells.
- The number of protons determines the element.
- Isotopes are atoms of the same element with different numbers of neutrons.
- Atoms with unequal protons and electrons are ions (cations: positive, anions: negative).
The Periodic Table
- Elements arranged by increasing number of protons (atomic number).
- Groups (columns) share the same number of valence electrons and similar chemical behavior.
- Periods (rows) share the same number of electron shells.
- Metals are on the left, non-metals on the right, and semimetals form a dividing line.
Chemical Bonds & Molecules
- Molecules: two or more atoms bonded; compounds: molecules with different elements.
- Covalent bonds share electrons; ionic bonds transfer electrons, forming crystal lattices (salts).
- Metallic bonds involve delocalized electrons, giving metals their properties.
- Bond types ordered by strength: Ionic > Covalent > Metallic > Hydrogen > Van der Waals.
- Polar covalent bonds create electric dipoles (e.g., in water).
- Hydrogen bonds and Van der Waals forces are types of intermolecular forces (IMFs).
States of Matter & Solutions
- Main states: solid, liquid, gas (and plasma at high temperatures or voltages).
- Temperature is average kinetic energy; entropy is disorder.
- Strong bonds lead to higher melting points.
- Mixtures: homogeneous (solutions), heterogeneous (suspensions), or colloids (emulsions).
Chemical Reactions & Stoichiometry
- Types of reactions: synthesis, decomposition, single and double replacement.
- Chemical equations must be balanced according to conservation of mass.
- Stoichiometry uses moles (amount of substance) based on atomic/molecular mass.
Energy & Equilibrium in Reactions
- Physical changes affect appearance; chemical changes alter substances.
- Activation energy is required for reactions; catalysts lower this barrier.
- Enthalpy (ΔH): heat content; exothermic reactions release heat, endothermic absorb it.
- Gibbs Free Energy (ΔG) combines enthalpy and entropy to determine reaction spontaneity.
- Chemical equilibrium: forward and reverse reactions occur at equal rates.
Acids, Bases, & Redox Reactions
- Acids donate protons (H+), bases accept protons (Brønsted-Lowry).
- pH measures hydronium ion concentration; lower pH = more acidic.
- pH + pOH = 14; neutralization forms water and salts.
- Redox reactions involve electron transfer; oxidation number rules help track changes.
Quantum Structure of Atoms
- Electrons occupy shells (n), subshells (l), orbitals (ml), and have spins (ms).
- Subshells: s (2 e-), p (6 e-), d (10 e-), f (14 e-); filling order follows the Aufbau principle.
- Electron configuration describes electron arrangement; noble gas shorthand used for simplicity.
- Valence electrons defined by electrons outside the last full noble gas shell.
Key Terms & Definitions
- Atom — Smallest unit of matter, made of protons, neutrons, electrons.
- Valence Electrons — Electrons in the outermost shell, important for bonding.
- Ion — Atom with a net charge from unequal protons and electrons.
- Isotope — Atoms with same protons, different neutrons.
- Mole — Amount of substance containing Avogadro's number of particles.
- Enthalpy (ΔH) — Heat content of a system.
- Gibbs Free Energy (ΔG) — Energy determining spontaneity of reactions.
- Acid/Base (Brønsted-Lowry) — Acid: proton donor; Base: proton acceptor.
- Redox Reaction — Reaction involving transfer of electrons.
- Orbitals/Subshells — Regions where electrons are likely found; s, p, d, f types.
Action Items / Next Steps
- Practice writing and balancing chemical equations.
- Memorize periodic table group properties and bond types.
- Review calculation of pH and basic rules for assigning oxidation numbers.
- Read next lesson on quantum mechanics (as indicated).