General Chemistry I: Comprehensive Study Guide

Chapter 1: Introduction to Chemistry

Classifying Matter: Pure Substances and Mixtures

Chemistry studies matter, which is anything that has mass and occupies space. Matter can be classified as pure substances or mixtures, and further as elements, compounds, homogeneous mixtures, or heterogeneous mixtures.

• Pure Substances: Consist of only one type of substance. Can be elements (single type of atom) or compounds (two or more elements chemically joined).

• Mixtures: Combinations of two or more substances. Can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

Example: Gold is a pure substance (element), coffee is a homogeneous mixture, and chocolate chip cookie is a heterogeneous mixture.

Physical and Chemical Properties

• Physical Properties: Characteristics that can be observed without changing the substance's identity (e.g., color, melting point).

• Chemical Properties: Describe a substance's ability to undergo chemical changes (e.g., flammability).

Physical and Chemical Changes

• Physical Change: Alters the form or appearance but not the chemical identity (e.g., melting ice).

• Chemical Change: Results in a new substance with different properties (e.g., rusting iron).

Measurement and Significant Figures

• Significant figures reflect the precision of a measurement.

• Rules:

  • All nonzero digits are significant.

  • Zeros between significant digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

Density

• Density is the ratio of mass to volume.

$\text{Density} = \frac{\text{mass}}{\text{volume}}$

Example: A brick stamped 14 karat gold has a mass of 110g and volume of 95 mL. Density = 1.3 g/mL.

Chapter 2: Matter and Its Components

Subatomic Particles

Atoms are composed of protons, neutrons, and electrons.

• Protons: Positive charge, found in nucleus.

• Neutrons: No charge, found in nucleus.

• Electrons: Negative charge, found outside nucleus.

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in an atom; defines the element.

• Mass Number (A): Total number of protons and neutrons.

$\text{Mass Number} = \text{Number of Protons} + \text{Number of Neutrons}$

Isotopes

• Atoms of the same element with different numbers of neutrons.

• Isotopes have the same atomic number but different mass numbers.

Nuclear Radiation and Radioactivity

• Unstable nuclei emit radiation (alpha, beta, gamma) to become more stable.

• Alpha Particle (α): Helium nucleus ($^4_2\text{He}$).

• Beta Particle (β): High-energy electron ($^0_{-1}\text{e}$).

• Gamma Ray (γ): High-energy electromagnetic radiation.

Half-Life

• The time required for half the atoms in a radioactive sample to decay.

$\text{Number of half-lives} = \frac{\text{total time}}{\text{half-life}}$

Chapter 3: Stoichiometry

Electron Arrangement and Periodic Table

• Electrons are arranged in shells and subshells (2n2 rule).

• Valence electrons are in the outermost shell and determine chemical properties.

• Groups (columns) indicate the number of valence electrons for main group elements.

• Periods (rows) indicate the number of energy levels.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

• Electronegativity: Decreases down a group, increases across a period.

Ion Formation

• Atoms gain or lose electrons to achieve noble gas configuration.

• Metals form cations (positive ions), nonmetals form anions (negative ions).

Writing Formulas and Naming Compounds

• Use the periodic table to determine charges and write correct formulas.

• Name ionic and covalent compounds according to standard rules.

Chapter 4: Solutions and Aqueous Reactions

Solutions and Solubility

• Solute: Substance being dissolved.

• Solvent: Substance doing the dissolving (usually present in greater amount).

• Solubility depends on temperature, pressure, and nature of solute/solvent.

Types of Mixtures

• Suspensions: Particles settle out.

• Colloids: Particles do not settle out (e.g., milk).

• Solutions: Homogeneous mixtures, particles do not settle out (e.g., salt water).

Concentration Units

• Molarity (M): Moles of solute per liter of solution.

$\text{Molarity} = \frac{\text{moles of solute}}{\text{liters of solution}}$

Electrolytes

• Strong Electrolytes: Completely ionize in solution (e.g., NaCl, HCl).

• Weak Electrolytes: Partially ionize (e.g., acetic acid).

• Nonelectrolytes: Do not ionize (e.g., sugar).

Chapter 5: Thermal Changes in Chemical Reactions

Thermodynamics and Reaction Types

• Exothermic Reaction: Releases heat (ΔH < 0).

• Endothermic Reaction: Absorbs heat (ΔH > 0).

• Enthalpy (ΔH): Heat content of a system.

• Entropy (ΔS): Measure of disorder.

• Free Energy (ΔG): Determines spontaneity of a reaction.

$\Delta G = \Delta H - T\Delta S$

Chapter 6: Introduction to Quantum Mechanics

Electron Configuration

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Pauli exclusion principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's rule: Electrons fill degenerate orbitals singly before pairing.

Chapter 7: Electron Configuration and Periodic Trends

Properties of Gases

• Gases have indefinite shape and volume, expand to fill container.

• Low density, easily compressed, diffuse rapidly.

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T)

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

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V)

• Ideal Gas Law: $PV = nRT$

Where R = 0.0821 L·atm/(mol·K), T in Kelvin.

Intermolecular Forces

• London Dispersion Forces: Present in all molecules, strongest in large nonpolar molecules.

• Dipole-Dipole Forces: Between polar molecules.

• Hydrogen Bonding: Strongest, occurs when H is bonded to N, O, or F.

Chapter 8: Chemical Bonding and Structure of Molecules

Ionic and Covalent Bonds

• Ionic Bonds: Transfer of electrons from metal to nonmetal.

• Covalent Bonds: Sharing of electrons between nonmetals.

Chapter 9: Molecular Shapes and Molecular Orbital Theory

VSEPR Theory

• Predicts molecular shapes based on electron pair repulsion.

Chapter 10: Gases and Their Properties

See Chapter 7 for gas laws and properties.

Chapter 11: Liquids, Solids, and Intermolecular Forces

See Chapter 7 for intermolecular forces and properties of liquids and solids.

Chapter 13: Solutions and Colligative Properties

Colligative Properties

• Depend on the number of solute particles, not their identity (e.g., boiling point elevation, freezing point depression).

Chapter 14: Reaction Kinetics

Reaction Rates

• Rate depends on concentration, temperature, surface area, and presence of a catalyst.

Chapter 15: Principles of Chemical Equilibrium

Equilibrium Constant

• At equilibrium, the rate of forward and reverse reactions are equal.

• Equilibrium constant expression (K):

  • For reaction: aA + bB ⇌ cC + dD

$K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

Concentrations of solids and liquids do not appear in the equilibrium expression.

Chapter 16: Acid-Base Equilibria

Acids and Bases

• Acid: Donates H+ ions in solution.

• Base: Accepts H+ ions in solution.

• Strong acids/bases dissociate completely; weak acids/bases only partially.

pH and pOH

• pH measures acidity; pOH measures basicity.

$\text{pH} = -\log[\text{H}^+]$

$\text{pOH} = -\log[\text{OH}^-]$

$\text{pH} + \text{pOH} = 14$

• pH < 7: Acidic; pH = 7: Neutral; pH > 7: Basic.

BONUS: Mathematical Operations and Functions

• Scientific notation is used for very large or small numbers.

• Conversion between units is essential (e.g., metric prefixes, dimensional analysis).

BONUS: Lab Techniques and Procedures

• Distinguish between chemical and physical properties and changes.

• Identify types of mixtures and methods of separation (filtration, distillation, etc.).