Introduction to Chemistry

Overview

This curriculum map outlines the foundational topics, laboratory objectives, and essential vocabulary for an introductory college-level chemistry course. The structure follows the sequence of chapters and units commonly found in modern chemistry textbooks, providing a comprehensive guide for students preparing for exams and laboratory work.

The Chemical World

Science, Technology, and the Scientific Method

• Science is the systematic study of the natural world through observation and experimentation.

• Technology applies scientific knowledge for practical purposes.

• The scientific method involves making observations, forming hypotheses, conducting experiments, and developing theories or laws.

• Hypothesis: A testable explanation for an observation.

• Scientific Law: A statement that describes consistent natural phenomena (e.g., Law of Conservation of Mass).

• Theory: A well-substantiated explanation of some aspect of the natural world.

Example: The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction.

Laboratory Safety and Equipment

• Identify key safety features: eye wash, safety shower, corrosive and combustible hazards.

• Use appropriate personal protective equipment (PPE) for different experiments.

• Familiarity with laboratory glassware: beaker, graduated cylinder, test tube, Erlenmeyer flask, balance.

• Understand and use Safety Data Sheets (SDS) and the RAMP framework (Recognize, Assess, Minimize, Prepare).

Measurement and Problem Solving

Numbers in Chemistry

• Exact numbers have no uncertainty (e.g., counted objects).

• Measured numbers have uncertainty due to limitations in measurement.

• Precision refers to the reproducibility of measurements; accuracy refers to closeness to the true value.

• Measured numbers always have uncertainty; the last digit is estimated.

• Identify leading zeros (never significant), captive zeros (always significant), and trailing zeros (significant if after a decimal point).

Scientific Notation and Calculations

• Convert numbers to and from scientific notation for clarity and to express uncertainty.

• Perform arithmetic operations with correct significant figures.

Example:

SI Units and Conversions

• The International System (SI) uses base units: kilogram (kg) for mass, meter (m) for length, second (s) for time, liter (L) for volume.

• Use conversion factors to change units.

• Density is defined as mass per unit volume:

Matter and Energy

Classification of Matter

• Matter is anything that has mass and occupies space.

• Atoms are the basic units of matter; molecules are groups of atoms bonded together.

• Pure substances (elements and compounds) vs. mixtures (homogeneous and heterogeneous).

• Physical properties (e.g., melting point) vs. chemical properties (e.g., reactivity).

• Physical changes do not alter composition; chemical changes produce new substances.

States of Matter and Energy

• Three main states: solid, liquid, gas.

• Energy can be kinetic or potential; work is the result of a force acting over a distance.

• Law of Conservation of Energy: Energy cannot be created or destroyed.

• Endothermic reactions absorb energy; exothermic reactions release energy.

• Temperature scales: Celsius (°C), Kelvin (K), Fahrenheit (°F).

• Heat transfer:

Example: Calculate the heat required to raise the temperature of 100 g of water by 10°C using .

Atoms and Elements

Atomic Structure and the Periodic Table

• Atoms consist of protons (positive), neutrons (neutral), and electrons (negative).

• Atomic number (Z): Number of protons; mass number (A): Protons + neutrons.

• Isotopes are atoms of the same element with different numbers of neutrons.

• Ions: Atoms with a net charge due to loss/gain of electrons (cations are positive, anions are negative).

• The periodic table organizes elements by increasing atomic number and similar properties (groups/families).

• Major categories: metals, nonmetals, metalloids.

• Calculate atomic mass as a weighted average of isotopic masses.

Example: Chlorine has two main isotopes: (75.77%) and (24.23%). Calculate the atomic mass.

Molecules and Compounds

Chemical Formulas and Nomenclature

• Chemical formula shows the types and numbers of atoms in a compound.

• Monatomic ions (single atom), polyatomic ions (multiple atoms).

• Ionic compounds (metal + nonmetal), molecular compounds (nonmetals only).

• Empirical formula: Simplest ratio; molecular formula: Actual number of atoms; structural formula: Shows arrangement.

• Naming conventions: Use Roman numerals for transition metals, -ide suffix for anions.

Example: is sodium chloride (ionic); is carbon dioxide (molecular).

Chemical Composition

Mole Concept and Calculations

• Mole: Amount of substance containing entities (Avogadro's number).

• Conversions: grams ↔ moles ↔ number of particles.

• Molar mass: Mass of one mole of a substance (g/mol).

• Percent composition: Mass percent of each element in a compound.

• Determine empirical formula from experimental data; calculate molecular formula using molar mass.

Example: Calculate the number of moles in 18 g of water (): mol.

Chemical Reactions

Types and Evidence of Chemical Reactions

• Evidence: color change, gas formation, precipitate, temperature change.

• Write and balance chemical equations.

• Use solubility rules to predict precipitation reactions.

• Write molecular, complete ionic, and net ionic equations.

• Types: acid-base, gas-evolution, redox, combustion.

Quantities in Chemical Reactions

Stoichiometry and Redox

• Use balanced equations to relate quantities of reactants and products.

• Oxidation-reduction (redox) reactions involve electron transfer.

• Key terms: oxidation state, oxidizing agent, reducing agent, galvanic cell, cathode, anode.

Chemical Bonding and Molecular Shape

VSEPR Theory and Polarity

• VSEPR theory predicts molecular shapes based on electron group repulsion.

• Sigma (σ) and pi (π) bonds describe types of covalent bonds.

• Polarity depends on bond dipoles and molecular geometry.

• Intermolecular forces (IMFs): dipole-dipole, London dispersion, hydrogen bonding.

Phases of Matter and Intermolecular Forces

Phase Changes and IMFs

• Phases: solid, liquid, gas; transitions require energy changes.

• London forces (dispersion), dipole-dipole, hydrogen bonds affect physical properties.

• Nonmolecular solids (network covalent, metallic) have high melting points.

Gases and the Ideal Gas Law

Gas Laws and Calculations

• Pressure is force per unit area; measured in atm, torr, Pa.

• Relationships: Boyle's Law (P-V), Charles's Law (V-T), Avogadro's Law (V-n).

• Ideal Gas Law:

• Calculate molar mass and density of gases using the ideal gas law.

Solutions

Properties and Calculations

• Solution: Homogeneous mixture of solute and solvent.

• Solubility depends on temperature, pressure, and intermolecular forces.

• Molarity (M):

• Colligative properties: vapor pressure lowering, boiling point elevation, freezing point depression.

Chemical Kinetics and Equilibrium

Reaction Rates and Equilibrium

Reaction mechanism: Sequence of steps in a reaction.

• Activation energy (Ea): Minimum energy required for a reaction.

• Rate law:

• Chemical equilibrium: Forward and reverse reactions occur at equal rates.

• Equilibrium constant (Keq) expresses the ratio of product to reactant concentrations at equilibrium.

• Le Chatelier's Principle: A system at equilibrium responds to disturbances to restore equilibrium.

Acids, Bases, and Electrolytes

Acid-Base Chemistry

• Acids donate protons (H+); bases accept protons.

• Strong vs. weak electrolytes: degree of dissociation in water.

• pH:

• Conjugate acid-base pairs differ by one proton.

• Titration is used to determine concentration of an acid or base.

Nuclear Chemistry

Radioactivity and Nuclear Reactions

• Types of radiation: alpha (α), beta (β), gamma (γ).

• Fission (splitting nuclei), fusion (combining nuclei), decay reactions.

• Penetrating power: gamma > beta > alpha.

• Half-life: Time for half of a radioactive sample to decay.

• Applications: nuclear power, medical imaging, carbon dating.

Key Laboratory Skills

• Proper measurement of mass, volume, and temperature.

• Use of laboratory glassware and safety equipment.

• Data analysis: graphing, interpreting experimental results, and applying calculations to laboratory data.

Sample Table: SI Base Units and Prefixes

Additional info: This guide is based on the curriculum map and vocabulary from "Introductory Chemistry" by Tro, covering all major topics in a standard introductory chemistry course. Laboratory objectives are integrated to reinforce conceptual understanding and practical skills.