Chapter 4 – Reactions in Aqueous Solution

4.1 – General Properties of Aqueous Solutions

An aqueous solution is a solution in which water is the solvent. Understanding the behavior of substances in water is fundamental to chemistry.

• Electrolyte: A substance that dissolves in water to produce ions, thus conducting electricity. Strong electrolytes dissociate completely (e.g., NaCl), while weak electrolytes only partially dissociate (e.g., acetic acid).

• Nonelectrolyte: A substance that dissolves in water but does not produce ions (e.g., sugar).

• Example: NaCl is a strong electrolyte; glucose is a nonelectrolyte.

4.2 – Precipitation Reactions

Precipitation reactions occur when two solutions of ionic compounds are mixed and an insoluble product (precipitate) forms.

• Use a solubility chart to determine if a compound is soluble (aq) or insoluble (s) in water.

• Precipitate: An insoluble solid formed in a reaction, usually denoted as (s).

• Ionic equation: Shows all ions present in the reaction.

• Net ionic equation: Shows only the ions and molecules directly involved in the reaction.

• Spectator ions: Ions that do not participate in the reaction.

Example: Mixing AgNO3 and NaCl forms AgCl (s) as a precipitate.

4.3 – Acids, Bases, and Neutralization Reactions

Acids and bases are fundamental to aqueous chemistry. Their strength and reactions are important for understanding solution behavior.

• Acid: Produces H+ ions in solution; often has a sour taste and reacts with metals.

• Base: Produces OH- ions in solution; often has a bitter taste and feels slippery.

• Strong acids/bases: Completely dissociate in water. Weak acids/bases: Partially dissociate.

• Neutralization reaction: Acid reacts with base to form water and a salt.

• Example: HCl + NaOH → NaCl + H2O

4.4 – Oxidation-Reduction Reactions

Redox reactions involve the transfer of electrons between substances. Understanding oxidation numbers is key to identifying these reactions.

• Oxidation: Loss of electrons (OIL: Oxidation Is Loss).

• Reduction: Gain of electrons (RIG: Reduction Is Gain).

• Oxidation number: A value assigned to an element in a compound based on electron distribution.

• Example: In NaCl, Na has an oxidation number of +1, Cl is -1.

4.5 – Concentrations of Solutions

Concentration describes the amount of solute in a given volume of solution. Molarity is the most common unit.

• Molarity (M):

• Conversion: 1 L = 1000 mL

• Dilution equation:

• Example: If you dilute 50 mL of 2 M NaCl to 100 mL, the new concentration is 1 M.

4.6 – Solution Stoichiometry and Chemical Analysis

Stoichiometry in solutions involves using concentrations and volumes to calculate reactant and product quantities.

• Use molarity and volume to find moles:

• Apply balanced equations to relate moles of reactants and products.

• Titration: A technique to determine concentration by reacting a known volume with a standard solution.

Chapter 6 – Electronic Structure of Atoms

6.1-6.4 – Nature/Behavior of Light, Quantized Energy, and the Bohr Model

Light and energy are quantized, meaning they exist in discrete packets. The Bohr model explains electron transitions in atoms.

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

• Wavelength (λ): Distance between wave peaks (meters).

• Relationship: where is the speed of light ( m/s).

• Energy of a photon: where is Planck's constant ( J·s).

• Bohr model: Electrons occupy quantized energy levels; transitions emit or absorb photons.

• Quantized energy: Only specific energy values are allowed.

• Electron transitions:

• Emission: Electron drops to a lower energy level, releasing energy.

• Excitation: Electron absorbs energy and moves to a higher level.

6.5/6.6 – Quantum Mechanics, Atomic Orbitals/Representations of Orbitals

Quantum numbers describe the properties of atomic orbitals and electrons.

• Principal quantum number (n): Indicates energy level; n = 1, 2, 3, ...

• Angular momentum quantum number (l): Defines orbital shape; l = 0 (s), 1 (p), 2 (d), 3 (f)

• Magnetic quantum number (ml): Specifies orbital orientation; ml = -l to +l

• Example: For n = 3, l can be 0, 1, or 2; ml ranges from -2 to +2 for d orbitals.

6.7-6.9 – Many-Electron Atoms, Orbital Diagrams, and Electron Configurations

Electron configurations and orbital diagrams show how electrons are arranged in atoms.

• Spin quantum number (ms): Indicates electron spin; ms = +1/2 or -1/2

• 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.

• Electron configuration: Notation showing distribution of electrons among orbitals (e.g., 1s2 2s2 2p6).

• Condensed configuration: Uses noble gas core to simplify notation (e.g., [Ne] 3s2 3p1).

• Orbital capacity: s (2 electrons), p (6), d (10)

Additional info: Figures 6.27 and 6.28 referenced in the notes are not provided, but the periodic table (image_3) is useful for electron configuration practice.