Course Schedule Overview

Introduction

This document outlines the week-by-week schedule for a General Chemistry I college course, including lecture topics, quizzes, exams, and project due dates. The schedule closely follows the standard curriculum for an introductory chemistry sequence, covering foundational topics essential for further study in chemistry and related sciences.

Chapter Topics and Sequence

Chapter 1: Matter & Measurement

This chapter introduces the fundamental concepts of chemistry, including the nature of matter, types of substances, and the importance of measurement in scientific inquiry.

• Matter: Anything that has mass and occupies space. Classified as elements, compounds, or mixtures.

• Measurement: The process of obtaining the magnitude of a quantity relative to an agreed standard. Includes units, significant figures, and uncertainty.

• Example: Measuring the mass of a sample using a balance and reporting the value with the correct number of significant figures.

Chapter 2: Atoms, Molecules & Ions

This chapter explores the basic building blocks of matter, atomic theory, and the formation of molecules and ions.

• Atoms: The smallest unit of an element that retains its chemical properties.

• Molecules: Groups of two or more atoms bonded together.

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a charge.

• Example: Sodium (Na) loses an electron to form Na+, while chlorine (Cl) gains an electron to form Cl-.

Chapter 3: Chemical Reactions & Stoichiometry

This chapter covers the types of chemical reactions, balancing equations, and quantitative relationships in reactions (stoichiometry).

• Chemical Reaction: A process in which substances are transformed into different substances.

• Stoichiometry: The calculation of reactants and products in chemical reactions using balanced equations.

• Key Equation: $ \mathrm{aA} + \mathrm{bB} \rightarrow \mathrm{cC} + \mathrm{dD} $

• Example: Calculating the mass of water produced from a given mass of hydrogen gas and oxygen gas.

Chapter 4: Reactions in Aqueous Solution

This chapter examines reactions that occur in water, including precipitation, acid-base, and redox reactions.

• Aqueous Solution: A solution in which water is the solvent.

• Types of Reactions: Precipitation, acid-base neutralization, and oxidation-reduction (redox).

• Example: Mixing solutions of silver nitrate and sodium chloride to form a precipitate of silver chloride.

Chapter 5: Thermochemistry

This chapter introduces the study of energy changes in chemical reactions, focusing on heat transfer and the concept of enthalpy.

• Thermochemistry: The study of the energy and heat associated with chemical reactions.

• Enthalpy (H): The heat content of a system at constant pressure.

• Key Equation: $ \Delta H = H_{\text{products}} - H_{\text{reactants}} $

• Example: Calculating the heat released when methane is combusted in oxygen.

Chapter 6: Electronic Structure of Atoms

This chapter explores the arrangement of electrons in atoms, quantum theory, and the periodic trends that arise from electronic structure.

• Quantum Theory: Describes the behavior of electrons in atoms using quantum numbers and orbitals.

• Electron Configuration: The distribution of electrons among the orbitals of an atom.

• Example: Writing the electron configuration for oxygen: 1s2 2s2 2p4.

Chapter 7: Periodic Properties of the Elements

This chapter discusses the periodic table, trends in atomic and ionic size, ionization energy, and electron affinity.

• Periodic Law: Properties of elements are periodic functions of their atomic numbers.

• Trends: Atomic radius decreases across a period, increases down a group; ionization energy shows the opposite trend.

• Example: Comparing the atomic radii of sodium and chlorine.

Chapter 8: Basic Concepts of Chemical Bonding

This chapter covers the types of chemical bonds (ionic, covalent, metallic), bond polarity, and the concept of electronegativity.

• Chemical Bond: The force holding two atoms together in a molecule or compound.

• Types: Ionic (transfer of electrons), covalent (sharing of electrons), metallic (delocalized electrons).

• Example: Sodium chloride (NaCl) is an ionic compound; water (H2O) is covalently bonded.

Chapter 9: Molecular Geometry and Bonding Theories

This chapter introduces the shapes of molecules, VSEPR theory, and bonding models such as hybridization and molecular orbital theory.

• VSEPR Theory: Predicts the geometry of molecules based on electron pair repulsion.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Example: Methane (CH4) has a tetrahedral geometry due to sp3 hybridization.

Chapter 10: Gases

This chapter examines the properties of gases, gas laws, and the kinetic molecular theory.

• Gas Laws: Mathematical relationships between pressure, volume, temperature, and amount of gas.

• Key Equations:

  • Boyle's Law: $ P_1V_1 = P_2V_2 $

  • Charles's Law: $ \frac{V_1}{T_1} = \frac{V_2}{T_2} $

  • Ideal Gas Law: $ PV = nRT $

• Example: Calculating the volume of a gas at different temperatures and pressures.

Chapter 11: Liquids and Intermolecular Forces

This chapter explores the properties of liquids and the forces that hold molecules together, such as hydrogen bonding, dipole-dipole, and London dispersion forces.

• Intermolecular Forces: Forces of attraction between molecules, influencing boiling and melting points.

• Types: Hydrogen bonding (strongest), dipole-dipole, London dispersion (weakest).

• Example: Water's high boiling point is due to hydrogen bonding.

Assessment Structure

• Quizzes: Regular quizzes after each chapter to reinforce understanding.

• Exams: Four major tests, each covering multiple chapters as outlined in the schedule.

• Projects: Assigned projects to apply concepts in practical or research contexts.

Summary Table: Chapter Coverage by Exam

Additional Info

• Some chapters (e.g., Ch. 12-24) are not covered in this schedule but are part of the broader General Chemistry curriculum.

• "FLIPPED" indicates a flipped classroom approach, where students review material before class and engage in active learning during class time.