General Chemistry Course

Lesson 1 • Dimensional Analysis and Unit Conversions

Teaches the factor-label method for converting between units systematically. This problem-solving strategy applies to stoichiometry, thermodynamics, and beyond.

Lesson 2 • Classification and Properties of Matter

Distinguishes pure substances from mixtures and physical from chemical properties. Provides the conceptual vocabulary used throughout the entire course.

Lesson 3 • SI Units and Measurement Systems

Introduces the International System of Units and metric prefixes for scientific measurement. Accurate unit use is foundational to all quantitative chemistry problems.

Lesson 1 • Periodic Table Organization and Trends

Examines how the periodic table is arranged by atomic number and electron configuration. Periodic trends in atomic radius, ionization energy, and electronegativity are quantified.

Lesson 2 • Historical Development of Atomic Theory

Traces atomic models from Dalton through quantum mechanics, highlighting key experimental evidence. Contextualizes why modern atomic theory is structured as it is.

Lesson 3 • Electron Configuration and Orbitals

Explains quantum numbers, orbital shapes, and rules for filling electrons. Electron configuration directly predicts bonding behavior and periodic trends.

Lesson 1 • Ionic Bonding and Lattice Structures

Describes electron transfer between metals and nonmetals and the resulting lattice energy. Connects ionic bond strength to physical properties such as melting point.

Lesson 2 • Covalent Bonding and Lewis Structures

Covers electron sharing, bond order, and the drawing of Lewis dot structures. Resonance and formal charge are used to evaluate competing structures.

Lesson 3 • Hybridization and Valence Bond Theory

Introduces orbital hybridization (sp, sp2, sp3) to explain observed bond angles. Sigma and pi bonds are distinguished to describe single and multiple bonds.

Lesson 4 • Intermolecular Forces and Physical Properties

Classifies London dispersion, dipole-dipole, and hydrogen bonding forces. Links force strength to boiling point, viscosity, and solubility trends.

Lesson 1 • Naming Covalent Compounds

Uses Greek prefixes to name binary molecular compounds systematically. Distinguishes molecular naming conventions from ionic naming rules.

Lesson 2 • Naming Ionic Compounds

Applies rules for naming binary and polyatomic ionic compounds, including transition metals. Correct naming is essential for reading and writing chemical equations.

Lesson 3 • Hydrates and Organic Compound Basics

Extends naming to hydrated salts and introduces simple organic compound families. Provides a bridge to organic chemistry encountered in later chapters.

Lesson 1 • Balancing Chemical Equations

Applies conservation of mass to balance equations by inspection and algebraic methods. Balanced equations are required for all stoichiometric calculations.

Lesson 2 • Limiting Reagents and Percent Yield

Identifies the limiting reagent and calculates theoretical and percent yield. These concepts are critical for evaluating reaction efficiency in practical settings.

Lesson 3 • Stoichiometric Calculations

Uses mole ratios from balanced equations to convert between masses of reactants and products. Builds the core quantitative skill applied in lab and industrial contexts.

Lesson 4 • The Mole Concept and Molar Mass

Defines Avogadro's number and molar mass as bridges between atomic and macroscopic scales. Mole calculations underpin every quantitative chemistry problem.

Lesson 1 • Solution Concentration and Preparation

Defines molarity, molality, and mass percent and explains how to prepare standard solutions. Accurate concentration calculations are essential for titrations and reaction stoichiometry.

Lesson 2 • Solubility Equilibria and Precipitation

Introduces Ksp to quantify the solubility of sparingly soluble salts. Common-ion effect and selective precipitation are applied to analytical separations.

Lesson 3 • Types of Aqueous Reactions

Classifies precipitation, acid-base, and oxidation-reduction reactions in aqueous solution. Net ionic equations isolate the species actually undergoing chemical change.

Lesson 4 • Acid-Base Equilibria and pH

Applies Brønsted-Lowry definitions and Ka/Kb expressions to calculate pH. Buffer systems and titration curves are analyzed as practical acid-base applications.

Lesson 1 • Enthalpy and Calorimetry

Defines enthalpy change at constant pressure and explains calorimetry measurements. Specific heat capacity is used to calculate heat transfer in coffee-cup and bomb calorimeters.

Lesson 2 • Entropy and the Second Law

Defines entropy as a measure of dispersal and applies the second law to predict spontaneity. Standard entropy values are used to calculate ΔS° for reactions.

Lesson 3 • Gibbs Free Energy and Spontaneity

Combines ΔH and ΔS into Gibbs free energy to determine reaction spontaneity at any temperature. ΔG° is linked to the equilibrium constant K through a fundamental equation.

Lesson 4 • Energy, Heat, and Work in Chemistry

Distinguishes heat from work and defines the system-surroundings framework. The first law of thermodynamics is applied to chemical and physical processes.

Lesson 1 • Oxidation-Reduction and Electrochemical Cells

Assigns oxidation states and balances redox equations using half-reaction methods. Galvanic and electrolytic cells are analyzed using standard reduction potentials.

Lesson 2 • Nuclear Chemistry and Radioactive Decay

Classifies types of radioactive decay and applies first-order kinetics to half-life calculations. Nuclear fission, fusion, and binding energy are introduced as energy sources.

Lesson 3 • Nernst Equation and Electrolysis

Applies the Nernst equation to calculate cell voltage under non-standard conditions. Faraday's laws quantify the mass deposited or dissolved during electrolysis.

Lesson 4 • Activation Energy and Catalysis

Explains the Arrhenius equation and the role of activation energy in reaction rates. Homogeneous and heterogeneous catalysts lower activation energy by alternative mechanisms.