Introductory General Chemistry (Lecture) – Syllabus and Study Guide

Course Overview

This course, Introductory General Chemistry (Lecture), covers the fundamental principles of general chemistry, including metric measurements, matter and energy, atomic structure, chemical nomenclature, chemical bonding, chemical reactions, stoichiometry, gas laws, molecular theory, properties of liquids, solids, solutions, acids, and bases.

Course Structure and Logistics

Class Meetings

• Days/Times: Mondays and Wednesdays, 10:00–11:15 AM

• Location: Room P-203

• Instructor: Dr. Beverly Symmann, Ph.D.

• Contact: Email and office hours via Zoom (see syllabus for details)

Required Materials

• Textbook: Nivaldo Tro, Introductory Chemistry, 7th Edition, Pearson, 2024

• Calculator: Non-graphing, scientific calculator

• Safety: Glasses or goggles (for lab)

Course Topics and Weekly Schedule

The course is organized into weekly topics, each focusing on a core area of general chemistry. Below is a summary of the main topics and their sequence:

Week	Topic	Chapters
1	Introduction, Matter & Energy	CH01, CH02
2	Molecules & Compounds	CH03
3	Chemical Composition	CH04
4	Chemical Reactions	CH05
5	Electrons in Atoms & The Periodic Table	CH06, CH07
6	Chemical Bonding	CH08
7	Gases	CH11
8	Liquids, Solids & Intermolecular Forces	CH12
9	Solutions	CH13
10	Acids & Bases	CH14
11	Oxidation & Reduction	CH16
12	Radioactivity & Nuclear Chemistry	CH17

Student Performance Objectives (Exit Skills)

Key Learning Goals

• Explain and relate chemical principles to the observed behavior of matter and energy in the laboratory.

• Interpret natural chemical phenomena in terms of physical causes and in the context of scientific theories.

• Demonstrate safe and effective laboratory techniques and procedures.

• Describe and apply concepts of atomic structure, valence, the periodic table, bonding, and reactivity.

• Classify and solve problems on stoichiometry, acids and bases, gas laws, Lewis structures, oxidation-reduction, and solubility.

• Work simple problems on stoichiometry and chemical formulas.

Student Learning Outcomes (SLOs)

• Use chemical concepts to solve quantitative problems and evaluate the results of calculations for reasonableness.

• Clearly explain qualitative chemical concepts and trends.

• Describe, explain, and model chemical and physical processes at the molecular level.

• Perform laboratory techniques correctly using appropriate safety procedures.

• Evaluate experimental values from laboratory data and interpret the results.

Assessment and Grading

• Homework: 10%

• Lab Reports: 20%

• Quizzes: 15%

• Final Exam: 20%

Grades are assigned based on a weighted average. Homework is assigned for each chapter, and lab reports are required for each experiment. Quizzes and exams are scheduled throughout the semester.

Key Chemistry Concepts Covered

Metric Measurements and Matter

• Metric System: Standard units for measuring mass, length, and volume (e.g., grams, meters, liters).

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

Atomic Structure

• Atoms: The basic unit of matter, composed of protons, neutrons, and electrons.

• Atomic Number ($Z$): Number of protons in the nucleus.

• Mass Number ($A$): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Chemical Nomenclature

• Ionic Compounds: Named using the cation and anion names (e.g., NaCl is sodium chloride).

• Covalent Compounds: Use prefixes to indicate the number of atoms (e.g., CO2 is carbon dioxide).

Chemical Bonding

• Ionic Bonds: Formed by the transfer of electrons from metals to nonmetals.

• Covalent Bonds: Formed by the sharing of electrons between nonmetals.

• Lewis Structures: Diagrams that show the bonding between atoms and the arrangement of valence electrons.

Chemical Reactions and Stoichiometry

• Chemical Equation: Represents a chemical reaction using symbols and formulas.

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

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

Example Equation:

$\ce{2H_2 + O_2 -> 2H_2O}$

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant temperature)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant pressure)

• Ideal Gas Law: $PV = nRT$

Acids, Bases, and Solutions

• Acids: Substances that donate protons ($H^+$) in solution.

• Bases: Substances that accept protons or donate hydroxide ions ($OH^-$).

• pH Scale: Measures the acidity or basicity of a solution; $pH = -\log[H^+]$

Oxidation and Reduction

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Redox Reactions: Chemical reactions involving the transfer of electrons.

Radioactivity and Nuclear Chemistry

• Radioactivity: The spontaneous emission of particles or energy from unstable atomic nuclei.

• Types of Radiation: Alpha ($\alpha$), beta ($\beta$), and gamma ($\gamma$) radiation.

Laboratory and Safety

• Proper use of safety glasses/goggles is required at all times in the laboratory.

• Lab reports must be submitted for each experiment, either handwritten or typed.

• Late lab reports and homework are penalized 5% per day late.

Academic Integrity and Policies

• All work must be your own; plagiarism or copying will result in a zero.

• During quizzes and exams, no phones or smart watches are allowed.

• Attendance is important for success; missing more than three classes may result in being dropped from the course.

Support Services

• Student Accessibility Services (SAS): Accommodations for students with documented disabilities.

• Counseling Services: Academic and personal counseling available.

• Learning Center: Free tutoring for chemistry and other subjects.