Introductory Chemistry: Core Topics and Study Guide

Ch. 1 Chemistry in Our Lives

Introduction to Chemistry

Chemistry is the study of matter, its properties, the changes it undergoes, and the energy associated with those changes. It plays a vital role in our daily lives, from the food we eat to the medicines we take and the environment we live in.

• Matter: Anything that has mass and occupies space.

• Chemical Change: A process in which substances are transformed into different substances.

• Application: Chemistry is used in healthcare, environmental science, engineering, and many other fields.

Ch. 2 Chemistry and Measurements

Units, Measurement, and Significant Figures

Accurate measurement is fundamental in chemistry. The International System of Units (SI) is used for consistency.

• SI Units: Standard units for measurement (e.g., meter for length, kilogram for mass, second for time).

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

• Scientific Notation: A method to express very large or small numbers, e.g., $1.23 \times 10^4$.

• Conversion Factors: Used to convert between units (e.g., $1 \ \mathrm{kg} = 1000 \ \mathrm{g}$).

Ch. 3 Matter and Energy

Classification and Properties of Matter

Matter can be classified by its physical state and composition. Energy is the capacity to do work or produce heat.

• States of Matter: Solid, liquid, gas.

• Physical vs. Chemical Properties: Physical properties can be observed without changing the substance; chemical properties describe a substance's ability to change into a new substance.

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

Ch. 4 Atoms and Elements

Basic Structure of Matter

All matter is composed of atoms, which are the smallest units of elements. Elements are pure substances that cannot be broken down by chemical means.

• Atom: Consists of protons, neutrons, and electrons.

• Element: Defined by the number of protons in its nucleus (atomic number).

• Periodic Table: Organizes elements by increasing atomic number and similar properties.

Ch. 5 Electronic Structure of Atoms and Periodic Trends

Electron Arrangement and Periodicity

The arrangement of electrons in atoms determines chemical properties and periodic trends.

• Electron Configuration: The distribution of electrons among orbitals.

• Periodic Trends: Patterns such as atomic radius, ionization energy, and electronegativity across periods and groups.

• Example: Atomic radius decreases across a period and increases down a group.

Ch. 6 Ionic and Molecular Compounds

Types of Chemical Bonds

Atoms combine to form compounds through ionic or covalent bonding.

• Ionic Bond: Transfer of electrons from a metal to a nonmetal, forming ions.

• Covalent Bond: Sharing of electrons between nonmetals.

• Example: Sodium chloride (NaCl) is ionic; water (H2O) is molecular (covalent).

Ch. 7 Chemical Quantities

The Mole and Molar Calculations

Chemists use the mole to count particles and relate masses to numbers of atoms or molecules.

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

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

• Example Calculation: To find moles from grams: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Ch. 8 Chemical Reactions

Types and Balancing of Chemical Reactions

Chemical reactions involve the rearrangement of atoms to form new substances. Equations must be balanced to obey the law of conservation of mass.

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

• Balancing Equations: Adjust coefficients to have equal numbers of each atom on both sides.

• Example: $2\mathrm{H}_2 + \mathrm{O}_2 \rightarrow 2\mathrm{H}_2\mathrm{O}$

Ch. 9 Chemical Quantities in Chemical Reactions

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations to determine the amounts of reactants and products.

• Mole Ratio: Derived from coefficients in the balanced equation.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible.

Note: Hess’s Law is omitted in this course.

Ch. 10 Bonding and Properties of Solids and Liquids

Intermolecular Forces and States of Matter

The properties of solids and liquids are determined by the types of bonding and intermolecular forces present.

• Types of Solids: Molecular, ionic, covalent network, metallic.

• Intermolecular Forces: Include hydrogen bonding, dipole-dipole, and London dispersion forces.

• Example: Water has strong hydrogen bonds, leading to high boiling point.

Ch. 11 Gases

Gas Laws and Properties

Gases have unique properties described by several laws relating pressure, volume, temperature, and amount.

• Boyle’s Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Ideal Gas Law: $PV = nRT$

Ch. 12 Solutions

Properties and Concentrations of Solutions

Solutions are homogeneous mixtures of solutes dissolved in solvents. Concentration expresses the amount of solute in a given amount of solution.

• Solubility: Maximum amount of solute that can dissolve in a solvent at a given temperature.

• Concentration Units: Molarity ($M$): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: Preparing a 1.0 M NaCl solution requires dissolving 1 mole of NaCl in 1 liter of water.

Ch. 14 Acids and Bases

Properties and Reactions of Acids and Bases

Acids and bases are important classes of compounds with characteristic properties and reactions.

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

• Base: Substance that accepts a proton or donates OH-.

• pH Scale: Measures acidity or basicity: $\mathrm{pH} = -\log[\mathrm{H}^+]$

• Example: Hydrochloric acid (HCl) is a strong acid; sodium hydroxide (NaOH) is a strong base.

Note: Buffers are omitted in this course.