Introductory Chemistry 22: Course Syllabus 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 a foundation in chemical concepts, laboratory skills, and critical thinking.

• Focus: Principles of chemistry, scientific method, laboratory techniques.

• Preparation: Designed for students aiming to continue in advanced chemistry courses.

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 in written and oral formats.

Course Structure and Requirements

Required Materials

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

• Calculator: Scientific calculator (e.g., CASIO fx-52ES PLUS).

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

Attendance and Participation

• Attendance is mandatory for lectures and labs.

• Participation includes discussion posts, assignments, and in-class activities.

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

Grading Policy

• Exams (3 total): 40%

• Quizzes: 15%

• Participation: 5%

• Problem Sets: 15%

• Lab Reports: 20%

• Final Exam: Comprehensive, 20%

Approximate Grading Scale

• 90 – 100% = A

• 80 – 89% = B

• 70 – 79% = C

• 60 – 69% = D

• < 60% = F

Major Topics and Weekly Schedule

Topic Outline

• Measurements and Scientific Notation

• Conversions and Density

• Dimensional Analysis

• Particulate Theory of Matter, Phases, Separations, and Energy

• Gas Laws (Ideal Gas Law, Charles' Law, etc.)

• Atomic Theory of Matter

• Counting Atoms by Weight

• Describing Composition: Mass %, Empirical and Molecular Formulas, Solutions

• EMR, Quantum Mechanics, Atomic Orbitals

• Periodic Properties

• Ionic Compounds: Nomenclature and Properties

• Molecular Compounds: Lewis Structures, VSEPR Theory

• Molecular Polarity, Intermolecular Forces

• Chemical Changes and Types of Chemical Reactions

• Stoichiometry, Limiting Reactants, Percent Yield

• Acids and Bases, pH

• Solubility, Precipitation, Gas-Evolution Reactions

Assessment Schedule

Assessment	Topics	Date
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 Exam	All Chem 22	12/10

Key Topic Summaries

Measurements and Scientific Notation

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

• Measurement: Determining the quantity of a substance using standard units (e.g., grams, liters).

• Scientific Notation: A method to write numbers as a product of a coefficient and a power of ten, e.g., $6.02 \times 10^{23}$.

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

• Example: The mass of a proton is $1.67 \times 10^{-24}$ grams.

Conversions and Density

Unit conversions and density calculations are essential for quantitative chemical analysis.

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

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

• Example: Water has a density of $1.00 \text{ g/mL}$ at room temperature.

Dimensional Analysis

Dimensional analysis is a systematic approach to problem solving that uses conversion factors to move from one unit to another.

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

• Method: Multiply the given value by conversion factors to cancel units.

• Example: Converting 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, Phases, Separations, and Energy

Chemistry studies matter at the particulate level, including its phases and energy changes.

• Phases of Matter: Solid, liquid, gas.

• Separation Techniques: Filtration, distillation, chromatography.

• Energy: The capacity to do work or produce heat; includes kinetic and potential energy.

• Example: Distillation separates mixtures based on differences in boiling points.

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 by Weight

Atomic theory explains the structure of matter. Counting atoms by weight uses the mole concept.

• Atom: The smallest unit of an element retaining its properties.

• Mole: $1 \text{ mole} = 6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar Mass: Mass of one mole of a substance, in grams per mole.

• Example: $\text{Moles} = \frac{\text{Mass}}{\text{Molar Mass}}$

Describing Composition: Mass %, Empirical and Molecular Formulas, Solutions

Chemists describe the composition of substances using formulas and concentration units.

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

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

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

• Solution Concentration: $M = \frac{\text{moles solute}}{\text{liters solution}}$

EMR, Quantum Mechanics, Atomic Orbitals

Electromagnetic radiation (EMR) and quantum mechanics explain atomic structure and electron behavior.

• EMR: Energy transmitted through space as waves (e.g., light).

• Quantum Mechanics: Describes the behavior of electrons in atoms.

• Atomic Orbitals: Regions in space where electrons are likely to be found.

• Example: The s, p, d, and f orbitals have distinct shapes and energy levels.

Periodic Properties and Ionic Compounds

The periodic table organizes elements by properties; ionic compounds form from the transfer of electrons.

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

• Ionic Compounds: Formed by the transfer of electrons between metals and nonmetals.

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

Molecular Compounds: Lewis Structures, VSEPR Theory

Molecular compounds are described by Lewis structures and VSEPR theory, which predict molecular shapes.

• Lewis Structure: Diagram showing valence electrons and bonds in a molecule.

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

• Example: Water (H2O) has a bent shape due to two lone pairs on oxygen.

Molecular Polarity and Intermolecular Forces

Polarity and intermolecular forces determine physical properties of substances.

• Polarity: Unequal sharing of electrons creates dipoles in molecules.

• Intermolecular Forces: Forces between molecules, including hydrogen bonding, dipole-dipole, and London dispersion.

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

Chemical Changes and Types of Chemical Reactions

Chemical reactions involve the transformation of substances through breaking and forming bonds.

• Types: Synthesis, decomposition, single replacement, double replacement, combustion.

• Balancing Equations: Ensures conservation of mass; same number of atoms on both sides.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Stoichiometry, Limiting Reactants, Percent Yield

Stoichiometry calculates quantities in chemical reactions; limiting reactants and percent yield assess reaction efficiency.

• Stoichiometry: Quantitative relationship between reactants and products.

• Limiting Reactant: The reactant that is completely consumed first.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Acids and Bases, pH

Acids and bases are defined by their ability to donate or accept protons; pH measures solution acidity.

• Acid: Substance that donates a proton (H+).

• Base: Substance that accepts a proton.

• pH: $\text{pH} = -\log[H^+]$

• Example: A solution with $[H^+] = 1 \times 10^{-7}$ M has pH 7.

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) from soluble reactants.

• Gas-Evolution Reaction: Produces a gas as one of the products.

• Example: Mixing Na2SO4 and BaCl2 forms insoluble BaSO4.

Laboratory Component

Lab Safety and Reports

• Lab Safety: Follow all safety protocols; wear goggles and appropriate attire.

• Lab Reports: Submit individual reports for each experiment; late reports receive reduced credit.

• Lab Notebook: Record all data and observations; required for each lab session.

Academic Integrity and Policies

Cheating and Plagiarism

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

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

Late Work Policy

• Late submissions receive a 10% penalty per day, up to 50% maximum reduction.

• Late pre-lab assignments are not accepted.

Dropping and Withdrawal

• Consistent participation is required to remain enrolled.

• Withdrawal deadlines: September 7 (no record), November 14 (W grade).

Supplemental Instruction

• SI sessions are available for extra support and may offer extra credit.

• Sessions are held outside of class hours and cover key topics.

Summary Table: Weekly Topics and Activities

Week	Lecture Topics	Lab Activity	Quiz	Experiment	Problem Set Due
1	Measurements, Scientific Notation	Lab Safety			#1 Due
2	Conversions, Density, Dimensional Analysis	Exp 1: Measurements/Dimensional Analysis		1	#2 Due
3	Particulate Theory, Phases, Separations, Energy	Exp 2: Density	Quiz 1	2	#3 Due
4	Gas Laws	Exp 3: Charles' Law	Quiz 2	3	#4 Due
5	Atomic Theory, Counting Atoms	Dispensing Volumes, Excel Graphing	Quiz 3	4	#5 Due
6	Describing Composition, Mass %, Empirical/Molecular Formulas, Solutions	Excel	Exam 1		#6 Due
7	EMR, Quantum Mechanics, Atomic Orbitals	Exp 4: % Sugar in Beverages	Quiz 4	5	#7 Due
8	Periodic Properties, Ionic Compounds	Bohr's Model Activity, Nomenclature Worksheet	Quiz 5	6	#8 Due
9	Molecular Compounds, Lewis Structures, VSEPR Theory	Lewis Structure Activity, Spectroscopy Lab		7	#9 Due
10	Molecular Polarity, Intermolecular Forces	Lewis Polarity Worksheet	Exam 2		#10 Due
11	Chemical Changes, Types of Chemical Reactions, Stoichiometry	Exp 6: Paper Chromatography, VSEPR Activity		8	#11 Due
12	Limiting Reactants, Yields, Gas Stoichiometry	Exp 7: Stoichiometry (CaSO4)	Quiz 6	9	#12 Due
13	Solution Stoichiometry, Acids and Bases, Titrations	Exp 8: Acid-Base Titrations	Quiz 7	10	#13 Due
14	Acids and Bases, pH, Solubility, Precipitation, Gas-Evolution Reactions	Thanksgiving (Holiday)		11	#14 Due
15	Final Exam Review	Exp 9: Quality Control of Vinegar, Lab Cleanup	Exam 3		#15 Due
16	No Class
17	Final Exam

Additional info:

• Some details inferred from standard chemistry curriculum and syllabus structure.

• Specific experiment titles and worksheet activities are based on common introductory chemistry lab practices.