General Chemistry I & II: Course Syllabus and Topic Overview

General Chemistry I & II: Course Syllabus and Topic Overview

This syllabus outlines the sequence of topics, key concepts, and reading assignments for a two-semester General Chemistry course. The course covers foundational principles, chemical reactions, thermodynamics, equilibrium, electrochemistry, nuclear chemistry, and biochemistry.

Course Structure

• Lecture Schedule: 44 lectures, including review sessions and exams.

• Reading Assignments: Corresponding textbook sections for each topic.

• Major Exams: Four in-class exams and a standardized final exam.

Main Topics and Subtopics

1. Introduction to Organic Chemistry

• Hydrocarbons: Structure, isomerism, and line-angle formulas.

• Nomenclature: Naming conventions for alkanes, alkenes, alkynes, and functional groups.

• Structural Isomers: Compounds with the same molecular formula but different connectivity.

• Functional Groups: Key reactive groups in organic molecules (e.g., alcohols, carboxylic acids).

2. Solutions and Solubility

• Solution Formation: Factors affecting solubility, including temperature and pressure.

• Solubility Product (Ksp): Equilibrium constant for sparingly soluble salts.

• Colligative Properties: Properties that depend on solute particle number (e.g., boiling point elevation, freezing point depression).

• Raoult’s Law: $P_A = X_A P_A^*$, where $P_A$ is the vapor pressure of component A, $X_A$ is its mole fraction, and $P_A^*$ is the vapor pressure of pure A.

3. Chemical Kinetics

• Rate Laws: Mathematical relationships between reactant concentrations and reaction rate.

• Integrated Rate Laws: Equations describing concentration changes over time for zero, first, and second order reactions.

• Half-life: Time required for half of a reactant to be consumed.

• Arrhenius Equation: $k = Ae^{-E_a/RT}$, where $k$ is the rate constant, $A$ is the frequency factor, $E_a$ is activation energy, $R$ is the gas constant, and $T$ is temperature in Kelvin.

4. Chemical Equilibrium

• Equilibrium Constant (K): Ratio of product and reactant concentrations at equilibrium.

• Le Châtelier’s Principle: Predicts the effect of changes in concentration, temperature, or pressure on equilibrium position.

• Reaction Quotient (Q): Used to determine the direction a reaction will proceed to reach equilibrium.

• ICE Tables: Tool for calculating equilibrium concentrations (Initial, Change, Equilibrium).

5. Acids, Bases, and Buffers

• Acid-Base Theories: Arrhenius, Brønsted-Lowry, and Lewis definitions.

• pH and pOH: $pH = -\log[H^+]$, $pOH = -\log[OH^-]$

• Buffer Solutions: Resist changes in pH upon addition of small amounts of acid or base.

• Henderson-Hasselbalch Equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Titrations: Analytical technique to determine concentration of an acid or base.

6. Thermodynamics

• First Law of Thermodynamics: Energy conservation in chemical systems.

• Enthalpy (ΔH): Heat content at constant pressure.

• Entropy (ΔS): Measure of disorder or randomness.

• Gibbs Free Energy (ΔG): $\Delta G = \Delta H - T\Delta S$; predicts spontaneity of reactions.

7. Electrochemistry

• Redox Reactions: Electron transfer processes; oxidation and reduction.

• Balancing Redox Equations: Using half-reactions in acidic or basic solution.

• Galvanic (Voltaic) Cells: Electrochemical cells that generate electrical energy from spontaneous redox reactions.

• Standard Electrode Potentials (E0): Measured under standard conditions; used to calculate cell potential.

• Nernst Equation: $E = E^0 - \frac{0.0592}{n} \log Q$ (at 25°C), where $n$ is the number of electrons transferred.

8. Nuclear Chemistry

• Types of Radioactive Decay: Alpha, beta, and gamma decay.

• Half-life Calculations: Time required for half of a radioactive sample to decay.

• Nuclear Fission and Fusion: Splitting and combining of atomic nuclei, respectively.

• Applications: Nuclear power, medical imaging, and food irradiation.

9. Biochemistry Introduction

• Proteins and Amino Acids: Structure and function of biological macromolecules.

• Fats and Oils: Structure and biological roles.

• Sugars: Carbohydrate chemistry basics.

10. Green Chemistry

• Principles of Green Chemistry: Designing chemical products and processes to reduce or eliminate hazardous substances.

Sample Table: Course Schedule Overview

Lecture	Date	Topics	Reading Section
1	8/18	Hydrocarbons, structural isomers, line-angle formulas	22.1 – 22.4
7	9/1	Solubility rules, net ionic equations, chemical kinetics	14.7 – 15.1
15	9/28	Acid, base, pH, calculations of strong acids and strong bases	16.7 – 16.8
29	10/2	Redox reactions, oxidation numbers	5.9, 20.1
36	11/1	Radioactivity detection, half-lives	21.5 – 21.6
41	11/12	Green Chemistry

Additional info: This syllabus provides a comprehensive overview of the topics typically covered in a two-semester General Chemistry sequence, including both foundational and applied aspects of chemistry. Students are encouraged to use the reading assignments to supplement lecture material and prepare for exams.