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

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., .

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

• Example: The mass of a proton is 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, .

• Example: Water has a density of 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: .

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: (at constant temperature)

• Charles' Law: (at constant pressure)

• Ideal Gas Law:

• Example: Calculate the volume of 1 mole of gas at STP:

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: particles (Avogadro's number).

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

• Example:

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

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

• Mass Percent:

• 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:

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:

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:

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:

• Example: A solution with 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

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.