CHEM 10123-015 General Chemistry II: Course Schedule and Topic Overview (Spring 2026)

Course Overview

This course schedule outlines the sequence of topics, readings, and assessments for CHEM 10123-015, a second-semester general chemistry course. The schedule is organized by class date, with each session focusing on key concepts from the textbook Chemistry, 8th edition by Robinson, McMurry, & Fay, supplemented by additional lecture material. The course covers advanced solution chemistry, chemical equilibrium, acid-base chemistry, thermodynamics, electrochemistry, kinetics, and an introduction to organic and coordination chemistry.

Major Topics and Subtopics

1. Solutions and Their Properties

Understanding the formation and properties of solutions is fundamental in chemistry, as many reactions occur in solution.

• Solution Formation, Entropy, and Free Energy (ΔG): Explores why solutions form, focusing on the roles of entropy (randomness) and Gibbs free energy ($\Delta G$). Key Equation: $\Delta G = \Delta H - T\Delta S$

• Concentration Units: Includes molarity, mole fraction, and mass percent. Example: Calculating mole fraction for a component in a solution.

• Solubility and Colligative Properties: Examines how solutes affect boiling point, freezing point, and vapor pressure. Key Equation (Boiling Point Elevation): $\Delta T_b = iK_b m$ Key Equation (Freezing Point Depression): $\Delta T_f = iK_f m$ van’t Hoff Factor (i): Number of particles a solute produces in solution.

2. Chemical Equilibrium

Chemical equilibrium describes the state where the rates of the forward and reverse reactions are equal, and concentrations remain constant.

• Equilibrium Constant (K) and Reaction Quotient (Q): $K$ expresses the ratio of product to reactant concentrations at equilibrium. $Q$ is calculated the same way but for any point in the reaction. Key Equation: $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Le Châtelier’s Principle: Predicts how a system at equilibrium responds to disturbances (changes in concentration, pressure, temperature).

• Equilibrium Calculations & RICE Charts: RICE (Reaction, Initial, Change, Equilibrium) tables help organize data for equilibrium problems.

3. Acids, Bases, and Aqueous Equilibria

This section covers acid-base theories, calculations, and the behavior of acids and bases in solution.

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

• pH and pOH Calculations: Key Equation: $\mathrm{pH} = -\log[H^+]$

• Strong vs. Weak Acids/Bases: Strong acids/bases dissociate completely; weak acids/bases only partially.

• Conjugate Acid-Base Pairs: Species differing by one proton.

• Ka, pKa, and pH Relationships: Key Equation: $K_a = \frac{[H^+][A^-]}{[HA]}$ $\mathrm{p}K_a = -\log K_a$

• Polyprotic Acids: Acids that can donate more than one proton.

• Salt Solutions and Hydrolysis: Not all salts form neutral solutions; hydrolysis can make them acidic or basic.

• Lewis Acids and Bases: Lewis acids accept electron pairs; Lewis bases donate them.

4. Buffers and Titrations

Buffers resist changes in pH, and titrations are used to determine concentrations of acids or bases.

• Buffer Solutions: Contain a weak acid and its conjugate base (or vice versa). Key Equation (Henderson-Hasselbalch): $\mathrm{pH} = \mathrm{p}K_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Buffer Capacity: The amount of acid or base a buffer can neutralize before pH changes significantly.

• Common-Ion Effect: The suppression of ionization of a weak acid/base by adding a common ion.

• Titration Curves: Graphs showing pH changes during titration; equivalence point is where moles of acid = moles of base.

• Indicators: Substances that change color at a specific pH range, used to detect the equivalence point.

5. Solubility Equilibria

Solubility equilibria involve the dissolution and precipitation of ionic compounds.

• Solubility Product Constant (Ksp): Key Equation: $K_{sp} = [A^+]^m [B^-]^n$

• Molar Solubility: The number of moles of solute that dissolve per liter of solution.

• Precipitation and Selective Precipitation: Predicting when a precipitate will form based on $Q$ and $K_{sp}$.

• Complex Ions and Formation Constant (Kf): Complex ions increase solubility of some salts. Key Equation: $K_f = \frac{[Complex]}{[Metal][Ligand]^n}$

6. Chemical Kinetics

Kinetics studies the rates of chemical reactions and the factors affecting them.

• Reaction Rate: Change in concentration of a reactant or product per unit time. Key Equation: $\text{Rate} = -\frac{1}{a}\frac{d[A]}{dt}$

• Rate Laws: Express the relationship between rate and concentration. General Form: $\text{Rate} = k[A]^m[B]^n$

• Integrated Rate Laws: Used to determine concentrations over time for zero, first, and second order reactions.

• Half-Life ($t_{1/2}$): Time required for half the reactant to be consumed. First Order: $t_{1/2} = \frac{0.693}{k}$

• Arrhenius Equation: Describes temperature dependence of rate constant. $k = A e^{-E_a/RT}$

• Reaction Mechanisms and Catalysis: Steps by which a reaction occurs; catalysts lower activation energy.

7. Thermochemistry and Thermodynamics

Thermochemistry focuses on heat changes in reactions; thermodynamics studies energy, entropy, and spontaneity.

• Heat, Work, and Internal Energy: $\Delta E = q + w$

• Enthalpy ($\Delta H$): Heat change at constant pressure.

• Hess’s Law: The total enthalpy change is the sum of individual steps.

• Entropy ($S$): Measure of disorder; increases with randomness.

• Gibbs Free Energy ($\Delta G$): Determines spontaneity. $\Delta G = \Delta H - T\Delta S$

• Relationship between $\Delta G$ and $K$: $\Delta G^\circ = -RT \ln K$

8. Electrochemistry

Electrochemistry deals with redox reactions and their applications in cells and batteries.

• Redox Reactions: Involve transfer of electrons; oxidation is loss, reduction is gain.

• Balancing Redox Equations: Use half-reactions for balancing in acidic or basic solution.

• Voltaic (Galvanic) Cells: Convert chemical energy to electrical energy; use spontaneous redox reactions.

• Cell Potential ($E^\circ_{cell}$): Calculated from standard reduction potentials. $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

• Nernst Equation: Calculates cell potential under nonstandard conditions. $E = E^\circ - \frac{0.0592}{n} \log Q$

• Electrolysis and Electroplating: Use electrical energy to drive nonspontaneous reactions.

• Batteries and Corrosion: Practical applications of electrochemistry.

9. Organic and Coordination Chemistry

Introduction to the structure, nomenclature, and properties of organic and coordination compounds.

• Hydrocarbons: Compounds containing only carbon and hydrogen; includes alkanes, alkenes, and alkynes.

• Isomerism: Structural and geometric isomers; optical isomers (enantiomers) and chirality.

• Functional Groups: Specific groups of atoms that determine chemical properties (e.g., alcohols, ketones, carboxylic acids).

• Nomenclature: Systematic naming of organic compounds.

• Coordination Compounds: Complexes formed between metal ions and ligands; nomenclature and properties.

• Crystal Field Theory: Explains color and magnetism of coordination complexes; crystal field splitting energy ($\Delta$), weak-field and strong-field ligands.

10. Molecular Structure and Bonding

Understanding the three-dimensional structure of molecules is essential for predicting reactivity and properties.

• VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.

• Valence Bond Theory (VBT): Describes bonding using atomic orbitals and hybridization.

• Drawing 3D Structures: Practice with organic and inorganic molecules.

Assessment and Practice

• Quizzes: Regular quizzes assess understanding of recent material.

• Exams: Three midterm exams and a cumulative final exam.

• Practice Sets: Problem sets and extra practice worksheets are provided for each major topic.

• Aktiv Chemistry Practice: Online practice sets organized by module.

Additional Info

• Some lecture content is not in the textbook; students are encouraged to attend class and review slides.

• Suggested textbook problems and additional resources are posted on the course D2L site.

• Students are encouraged to review mistakes and seek connections between concepts and mathematical approaches.