General Chemistry: Solutions, Acids & Bases, and Atomic Structure Study Guide

Chapter 4: Solutions and Reactions

4.1–4.3: Solution Concentration and Dilution

Understanding solution concentration is fundamental in chemistry, as it allows chemists to quantify the amount of solute in a given volume of solvent. Molarity (M) is the most common unit of concentration.

• Molarity (M): The number of moles of solute per liter of solution.

• Calculation: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution: The process of reducing the concentration of a solution by adding more solvent. The relationship is: $M_1V_1 = M_2V_2$ where $M_1$ and $V_1$ are the initial molarity and volume, and $M_2$ and $V_2$ are the final molarity and volume.

• Applications: Preparing solutions of desired concentrations for reactions or analysis.

4.4–4.5: Electrolytes and Types of Reactions

Substances in aqueous solution can conduct electricity depending on their ability to dissociate into ions.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, HCl).

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

• Nonelectrolytes: Do not dissociate (e.g., sugar in water).

• Types of Reactions: Precipitation, acid-base neutralization, and oxidation-reduction (redox) reactions.

4.6–4.7: Writing Chemical Equations

Chemical reactions in aqueous solution are represented by molecular, total ionic, and net ionic equations.

• Molecular Equation: Shows all reactants and products as compounds.

• Total Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that actually change during the reaction.

4.8: Solubility Guidelines

Solubility rules help predict whether an ionic compound will dissolve in water. These rules are essential for determining the products of precipitation reactions.

Ion	Solubility in Water	Common Exceptions
Li+, Na+, K+, Rb+, Cs+, NH4+	Soluble	None
NO3-, C2H3O2-, ClO4-	Soluble	None
Cl-, Br-, I-	Soluble	Ag+, Hg22+, Pb2+
SO42-	Soluble	Sr2+, Ba2+, Pb2+, Ag+, Ca2+
OH-	Insoluble	Group 1A, Ca2+, Sr2+, Ba2+
CO32-, PO43-	Insoluble	Group 1A, NH4+

4.9: Precipitation Reactions

Precipitation reactions occur when two soluble salts are mixed and an insoluble product (precipitate) forms. Use solubility rules to predict the outcome and write net ionic equations.

4.10: Naming and Formulas of Oxyanions

Oxyanions are polyatomic ions containing oxygen. Their names and formulas follow specific patterns.

Acid Name	Formula	Anion Name	Anion Formula
Perchloric acid	HClO4	Perchlorate	ClO4-
Chloric acid	HClO3	Chlorate	ClO3-
Chlorous acid	HClO2	Chlorite	ClO2-
Hypochlorous acid	HClO	Hypochlorite	ClO-

4.11–4.12: Acid-Base Reactions

Acid-base neutralization involves the reaction of an acid and a base to form water and a salt. Write molecular, total ionic, and net ionic equations for these reactions.

4.13–4.18: Stoichiometry and Redox Reactions

• Stoichiometry: Use mole and volume relationships to solve for unknowns in reactions.

• Oxidation Numbers: Assign to atoms to identify redox changes.

• Redox Reactions: Involve transfer of electrons; identify oxidizing and reducing agents.

• Activity Series: Predicts if a redox reaction will occur.

• Titration: Used to determine the concentration of an oxidizing or reducing agent.

Chapter 5: Atomic Structure and Quantum Theory

5.1–5.3: Electromagnetic Radiation

Light exhibits both wave and particle properties. Key characteristics include wavelength, frequency, and amplitude.

• Wavelength (λ): Distance between two consecutive peaks.

• Frequency (ν): Number of cycles per second (Hz).

• Relationship: $c = \lambda \nu$ where $c$ is the speed of light.

• Energy of a photon: $E = h\nu$ where $h$ is Planck's constant.

5.4–5.5: Photoelectric Effect and Atomic Spectra

The photoelectric effect demonstrates the particle nature of light. Atomic spectra provide evidence for quantized energy levels.

• Photoelectric Effect: Electrons are ejected from a metal when light of sufficient frequency shines on it.

• Continuous vs. Line Spectrum: Continuous spectrum contains all wavelengths; line spectrum contains only specific wavelengths.

5.6–5.8: Bohr Model and Quantum Numbers

The Bohr model explains electron transitions in atoms. Quantum numbers describe the properties of atomic orbitals and electrons.

• de Broglie Equation: $\lambda = \frac{h}{mv}$

• Quantum Numbers: Principal (n), Angular momentum (l), Magnetic (ml), Spin (ms).

5.9–5.17: Electron Configuration and Periodic Trends

Electron configurations describe the arrangement of electrons in atoms. The periodic table helps predict atomic properties.

• Aufbau Principle: Electrons fill orbitals from lowest to highest energy.

• Hund's Rule: Every orbital in a subshell is singly occupied before any is doubly occupied.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Periodic Trends: Atomic radius, ionization energy, electron affinity, and electronegativity vary predictably across the table.

Chapter 6: Periodic Properties and Ionic Compounds

6.1–6.3: Electron Configurations and Unpaired Electrons

Electron configurations for main group and transition metal ions are essential for understanding chemical behavior.

• Ground State: The lowest energy arrangement of electrons.

• Unpaired Electrons: Important for magnetism and reactivity.

6.4–6.7: Atomic and Ionic Size

Atomic and ionic radii change across periods and down groups due to effective nuclear charge and electron shielding.

• Isoelectronic Ions: Ions with the same number of electrons but different nuclear charges.

• Trends: Atomic radius decreases across a period, increases down a group.

6.8–6.9: Electron Affinity

Electron affinity is the energy change when an atom gains an electron. Trends can be explained by atomic structure.

6.10–6.12: Ionic Compounds and Lattice Energy

The octet rule helps predict ion charges and formulas. Lattice energy is the energy released when ions form a solid lattice.

• Born-Haber Cycle: A thermochemical cycle used to calculate lattice energy.

• Lattice Energy: Increases with higher charge and smaller ionic radius.