Overview
This lecture introduces fundamental chemistry concepts including atomic structure, the periodic table, chemical bonding, states of matter, chemical reactions, and quantum numbers.
Atomic Structure
- All matter consists of atoms, which have a nucleus (protons and neutrons) surrounded by electrons.
- The number of protons determines an element's identity; varying neutrons result in isotopes.
- Electrons occupy shells; those in the outermost are called valence electrons and influence chemical properties.
- Atoms with equal protons and electrons are neutral; more electrons make anions (negative ions), fewer make cations (positive ions).
Periodic Table Organization
- Elements are arranged by increasing proton number; each cell shows name, symbol, protons, and atomic mass.
- Columns (groups) share the same number of valence electrons; rows (periods) have the same number of shells.
- Metals are on the left, nonmetals on the right, and semimetals in between.
- Alkali metals (Group 1, except hydrogen) are soft, shiny, reactive metals with one valence electron.
Chemical Bonds and Molecules
- Atoms form molecules (same or different elements); compounds have at least two different elements.
- Ionic bonds form when electrons transfer between atoms (usually metals to nonmetals); covalent bonds involve sharing electrons.
- Metallic bonds occur in metals with delocalized electrons.
- Bond strength order: Ionic > Covalent > Metallic > Hydrogen > Van der Waals.
- Electronegativity measures atom's electron pull; fluorine is highest.
Intermolecular Forces & Polarity
- Polar molecules have uneven charge distribution; nonpolar share electrons equally.
- Hydrogen bonds (strong IMFs) form between H and F, O, or N.
- Van der Waals forces are weak, momentary attractions in all molecules.
- "Like dissolves like": polar solvents dissolve polar/ionic substances; nonpolar dissolve nonpolar.
- Surfactants (like soap) have both polar and nonpolar parts and form micelles.
States of Matter & Physical Properties
- Solids: fixed structures; liquids: fixed volume, particles move; gases: particles free, fill container.
- Temperature = average kinetic energy; entropy = degree of disorder.
- Higher bond strength increases melting point.
- Plasma is ionized gas found at high temperatures/voltages, emits characteristic light (emission spectrum).
Mixtures and Solutions
- Pure substances are elements or compounds; mixtures combine pure substances.
- Homogeneous mixtures are uniform (solutions); heterogeneous mixtures have visible separation (suspensions).
- Colloids/emulsions (e.g., milk) are intermediate, with evenly distributed but undissolved particles.
Chemical Reactions and Stoichiometry
- Types: synthesis, decomposition, single/double replacement; all seek lower energy/stability.
- Reactions follow conservation of mass; atoms must balance on both sides of equations.
- Stoichiometry uses moles (based on atomic mass) to count particles.
- Physical changes alter form; chemical changes alter substances, often producing gas, odor, or energy.
- Catalysts speed reactions by lowering activation energy but aren't consumed.
Thermodynamics and Equilibrium
- Enthalpy is total heat content; exothermic reactions release heat, endothermic absorb it.
- Gibbs Free Energy predicts reaction spontaneity: ΔG < 0 = spontaneous, ΔG > 0 = non-spontaneous.
- Entropy and temperature influence spontaneity (e.g., melting ice is endothermic but spontaneous above 0°C).
- Equilibrium: forward and reverse reactions at equal rates, concentrations stable.
Acids, Bases, and pH
- Acids donate protons (H⁺); bases accept protons (Brønsted-Lowry theory).
- Amphoteric substances act as acid or base.
- Strong acids/bases dissociate almost completely; weak ones do not.
- pH = -log[H₃O⁺], lower pH is more acidic; pH + pOH = 14.
- Neutralization: strong acid + strong base → salt + water.
Redox Reactions and Oxidation States
- Redox reactions involve electron transfer and changes in oxidation numbers.
- Oxidation: loss of electrons; reduction: gain of electrons.
- Oxidation states follow set rules (H=+1, O=-2, halogens=-1, elements=0).
Quantum Numbers and Electron Configuration
- Four quantum numbers (n, l, ml, ms) describe electron positions and energy.
- s, p, d, f subshells hold 2, 6, 10, and 14 electrons respectively.
- Pauli Exclusion: no two electrons in an atom have the same quantum numbers.
- Aufbau principle: fill subshells in specific order for electron configuration.
Key Terms & Definitions
- Atom — smallest unit of matter with a nucleus and electrons.
- Valence Electrons — electrons in the outermost shell.
- Ion — charged atom (cation = positive, anion = negative).
- Electronegativity — tendency of an atom to attract electrons.
- Isotope — atoms with same protons, different neutrons.
- Ionic/Covalent/Metallic Bond — types of atomic connections via electron transfer or sharing.
- Intermolecular Forces (IMFs) — forces between molecules (hydrogen bonds, Van der Waals).
- Mole — amount of substance (6.022 × 10²³ particles).
- Enthalpy — total heat energy in a system.
- Gibbs Free Energy (ΔG) — determines reaction spontaneity.
- pH — measure of acidity; negative log of hydronium ion concentration.
- Oxidation State — hypothetical charge for atom in molecule.
- Quantum Number — values describing electron's state in atom.
Action Items / Next Steps
- Review periodic table and locate groups, periods, and element properties.
- Practice balancing chemical equations and calculating moles.
- Complete assigned textbook reading on atomic structure, bonding, and basic chemical reactions.