Introductory Chemistry 22: Course Structure, Topics, and Study Guide

Course Overview

Course Description

This course introduces the principles of chemistry, emphasizing problem solving and scientific reasoning. The primary goal is to prepare students for Chemistry 1A and to build foundational knowledge in chemical concepts, laboratory techniques, and scientific thinking.

• Focus: Problem solving, chemical principles, laboratory skills.

• Preparation: Designed for students aiming to continue in Chemistry 1A.

Student Learning Outcomes (SLOs)

• Critical Thinking: Apply scientific reasoning to solve problems related to matter and chemical changes.

• Laboratory Skills: Use scientific technologies and laboratory practices to collect, evaluate, and interpret data.

• Communication: Effectively communicate scientific findings and concepts.

Required Materials

• Textbook: Introductory Chemistry, PCC Custom 5th Edition, Tro.

• Calculator: Scientific calculator (e.g., CASIO fx- 2 5ES PLUS).

• Lab Supplies: Lab goggles, lab notebook, and other specified items.

Course Structure and Policies

Attendance and Participation

Attendance and participation in all assignments are required. This includes discussion posts, turning in assignments on time, and active engagement in labs and lectures.

• Failure to participate may result in being dropped from the course.

Grading

• Exams (3 total): 40%

• Quizzes (lowest score dropped): 15%

• Experiments: 20%

• Problems: 15%

• Participation (POGIL, homework, class worksheets, Ativ chemistry): 5%

• Final Exam: 20%

Approximate Grading Scale

• 90 – 100% = A

• 80 – 89% = B

• 70 – 79% = C

• 60 – 69% = D

• < 60% = F

Assessment Explanations

• Exams: Assess mastery of topics from textbook, lectures, and labs. May include multiple-choice, short answer, and calculation-based questions.

• Quizzes: Test understanding of recent material; may be closed or open book.

• Lab Notebooks: Record experimental details and results; must be completed before lab sessions.

• Experiment Reports: Submit individual lab reports for each experiment.

• Discussions: Participate in online or in-class discussions to reinforce concepts.

• Final Exam: Cumulative, covering all course material.

Late Work Policy

• Late work is penalized: 10% deduction for work submitted within 24 hours, 50% deduction within one week.

• Late pre-lab assignments are not accepted.

Cheating and Academic Integrity

• Cheating includes copying work, using unauthorized resources, and plagiarism.

• Violations result in a zero for the assignment and possible dismissal from the course.

Course Schedule and Topics

Weekly Topics Overview

The following table summarizes the main topics, lab activities, quizzes, and assessments for each week.

Week	Date	Lecture Topics	Lab Activity	Quiz	Experiment	POGIL	P-Set
1	8/25	Measurements and Scientific Notation	Lab safety
1	8/27	Conversions and Density					#1 Due
2	9/3	Dimensional Analysis	Exp 1 - Measurements/Dimensional analysis exercise				#1 Due
3	9/10	Particulate Theory of Matter and Phases; Phase Changes, Separations, and Energy	Exp 2 - Density	Quiz 1	Exp 1		#2 Due
4	9/17	Simple Gas Laws, Ideal Gas Law		Quiz 2	Exp 2		#3 Due
5	9/24	Applications of Gas Laws		Quiz 3	Exp 3		#4 Due
6	10/1	Atomic Theory of Matter; Counting Atoms by Weight	Dispensing volumes using beaker, excel graphing exercise				#5 Due
6	10/3	Describing Composition: Mass %, Empirical and Molecular Formulas, Solutions and Solution Concentrations		Exam 1			#6 Due
7	10/6	EMR, Intro to Quantum Mechanics; Atomic Orbitals and Electron Configurations	Exp 4: % Sugar in beverage (Standard curve)		Exp 4		#7 Due
8	10/13	Periodic Properties; Ionic Compounds: Nomenclature and Properties	Bohr's model activity (in-lab); Nomenclature worksheet				#6 Due
9	10/20	Molecular Compounds: Lewis Structures; VSEPR Theory	Lewis structure activity; Exp 5: Spectroscopy lab (Determination of dye in Gatorade)	Quiz 5	Exp 5		#7 Due
10	10/27	Molecular Polarity; Intermolecular Forces		Exam 2			#8 Due
11	11/3	Chemical Changes and Types of Chemical Reactions; Stoichiometry	Exp 6: Paper chromatography using molecular models to determine polarity worksheet; VSEPR activity		Exp 6		#8 Due
12	11/10	Limiting Reactants, Yields; Gas Stoichiometry, AH	Exp 7: Stoichiometry (CaSO4)	Quiz 7	Exp 7		#9 Due
13	11/17	Solution Stoichiometry; Acids and Bases, Titrations	Exp 8: Acid-base titrations		Exp 8		#10 Due
14	11/24	Acids and Bases, pH; Solubility, Precipitation, Gas-Evolution Reactions	Thanksgiving (Holiday)				#11 Due
15	12/1	Exam 3; Final Exam Review	Exp 9: Quality control of vinegar, lab cleanup		Exp 9
16	12/8	No Class
16	12/10	FINAL EXAM (8am – 10am)

Assessment Topics Table

Assessment	Topics	When
Quiz 1	Sig. Figs., Measurements, Conversions, Density	9/10
Quiz 2	Dimensional Analysis, Phases, Separations, Heat and Energy	9/17
Quiz 3	Gas Laws	9/24
Exam 1	Measurements, Matter, and Phases	10/3
Quiz 4	Atomic Theory, Mole Calculations, Formulas	10/8
Quiz 5	EMR, Bohr Model, Electron Configurations, Spectroscopy	10/22
Exam 2	Atomic Theory, Moles, Quantum Mechanics, Ionic Compounds	10/31
Quiz 6	Lewis Structures, VSEPR, Intermolecular Forces	11/12
Quiz 7	Stoichiometry, Limiting Reactants, Percent Yield	11/26
Exam 3	Molecular Compounds, Lewis Structure, VSEPR, Intermolecular Forces, Chemical Reactions, Stoichiometry, Solutions	12/1
Final	All Chem 22	12/10

Key Chemistry Topics Explained

Measurements and Scientific Notation

Accurate measurement is fundamental in chemistry. Scientific notation is used to express very large or small numbers efficiently.

• Measurement: Quantitative description of physical properties (e.g., mass, volume).

• Scientific Notation: Numbers written as $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Example: $0.00056 = 5.6 \times 10^{-4}$

Conversions and Density

Unit conversions and density calculations are essential for laboratory work and chemical analysis.

• Unit Conversion: Changing from one unit to another using conversion factors.

• Density: Ratio of mass to volume, $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Example: If a sample has a mass of 10 g and a volume of 2 mL, its density is $5 \text{ g/mL}$.

Dimensional Analysis

Dimensional analysis is a method for converting units and solving problems using conversion factors.

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit.

• Method: Multiply by conversion factors to cancel units and obtain the desired unit.

• Example: Convert 5.0 cm to meters: $5.0 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.050 \text{ m}$

Particulate Theory of Matter and Phases

Matter is composed of particles (atoms, molecules, ions) and exists in different phases: solid, liquid, and gas.

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume, indefinite shape; particles can move past each other.

• Gas: Indefinite shape and volume; particles are far apart and move freely.

• Phase Changes: Transitions between solid, liquid, and gas (e.g., melting, boiling).

Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

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

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

• Ideal Gas Law: $PV = nRT$

• Example: Calculate the volume of 1 mole of gas at STP: $V = \frac{nRT}{P}$

Atomic Theory and Counting Atoms

Atomic theory explains the structure of atoms and how to count atoms using moles and Avogadro's number.

• Atom: Smallest unit of an element, composed of protons, neutrons, and electrons.

• Mole: $6.022 \times 10^{23}$ particles (Avogadro's number).

• Example: 1 mole of carbon contains $6.022 \times 10^{23}$ atoms.

Describing Composition: Mass %, Empirical and Molecular Formulas

Chemists describe the composition of substances using mass percent, empirical, and molecular formulas.

• Mass Percent: $\text{Mass %} = \frac{\text{mass of element}}{\text{total mass}} \times 100$

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

Quantum Mechanics and Electron Configurations

Quantum mechanics describes the behavior of electrons in atoms, including their arrangement in orbitals.

• Electron Configuration: Distribution of electrons among atomic orbitals.

• Example: Carbon: $1s^2 2s^2 2p^2$

Periodic Properties and Ionic Compounds

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Periodic Trends: Atomic radius, ionization energy, electronegativity.

• Ionic Compounds: Formed by transfer of electrons; consist of cations and anions.

• Nomenclature: Naming rules for ionic compounds (e.g., NaCl: sodium chloride).

Molecular Compounds, Lewis Structures, and VSEPR Theory

Molecular compounds are formed by sharing electrons. Lewis structures and VSEPR theory predict molecular shapes.

• Lewis Structure: Diagram showing bonding and lone pairs of electrons.

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

• Example: Water ($H_2O$) has a bent shape due to two lone pairs on oxygen.

Molecular Polarity and Intermolecular Forces

Molecular polarity depends on the distribution of electrons and affects intermolecular forces.

• Polar Molecule: Unequal sharing of electrons; has a dipole moment.

• Intermolecular Forces: Forces between molecules (e.g., hydrogen bonding, dipole-dipole, London dispersion).

• Example: Water is polar and exhibits hydrogen bonding.

Chemical Changes and Types of Chemical Reactions

Chemical reactions involve the transformation of substances. Types include synthesis, decomposition, single replacement, and double replacement.

• Stoichiometry: Quantitative relationships in chemical reactions.

• Balanced Equation: $aA + bB \rightarrow cC + dD$

• Limiting Reactant: Reactant that determines the amount of product formed.

Solution Stoichiometry, Acids and Bases, Titrations

Solution stoichiometry involves calculations with solutions. Acids and bases are defined by their ability to donate or accept protons.

• Acid: Proton donor; Base: Proton acceptor.

• pH: $\text{pH} = -\log[H^+]$

• Titration: Technique to determine concentration by reacting with a standard solution.

Solubility, Precipitation, Gas-Evolution Reactions

Solubility rules predict whether a compound will dissolve. Precipitation and gas-evolution reactions are common in aqueous chemistry.

• Solubility: Ability of a substance to dissolve in a solvent.

• Precipitation Reaction: Formation of an insoluble product (precipitate).

• Gas-Evolution Reaction: Produces a gas as a product.

Supplemental Instruction and Support

Supplemental Instruction (SI)

SI sessions provide additional support and review of key concepts. Attendance can earn extra credit.

• Weekly sessions outside of class hours.

• Review and practice of challenging topics.

Important Dates

• Drop Deadline: September 7, 2025

• Withdrawal Deadline: November 14, 2025

• Final Exam: December 10, 2025 (8am – 10am)

Summary Table: Key Chemistry Concepts

Concept	Definition	Example
Density	Mass per unit volume	$\text{Density} = \frac{10\text{ g}}{2\text{ mL}} = 5\text{ g/mL}$
Ideal Gas Law	Relates pressure, volume, temperature, and moles	$PV = nRT$
Empirical Formula	Simplest ratio of elements	CH2O for glucose
pH	Measure of acidity	$\text{pH} = -\log[H^+]$

Additional info: Some details (e.g., specific lab activities, SI session structure) were inferred from standard introductory chemistry syllabi and the provided schedule.