General Chemistry Study Notes: Aqueous Equilibria, Thermodynamics, Electrochemistry, Main Group & Transition Elements

Chapter 14: Aqueous Equilibria—Acids and Bases

Key Definitions and Acid-Base Concepts

This chapter introduces the fundamental concepts of acids and bases, their equilibria in aqueous solutions, and calculations involving pH and buffer systems.

• Arrhenius acid: Generates H+ ions in water.

• Arrhenius base: Generates OH- ions in water.

• Brønsted-Lowry acid: Proton donor.

• Brønsted-Lowry base: Proton acceptor.

• Conjugate acid-base pairs: Differ by one proton (H+).

Equilibrium constant for acids (Ka): Measures acid strength. For a generic acid HA:

$K_a = \frac{[H^+][A^-]}{[HA]}$

pH and pOH: pH = -log[H+], pOH = -log[OH-], and pKa = -log Ka.

Relationship: $pH + pOH = 14$ at 25°C.

Calculating pH and Equilibrium Concentrations

• For strong acids/bases, assume complete dissociation.

• For weak acids/bases, use ICE tables and Ka or Kb to solve for equilibrium concentrations.

• Example: Calculate pH for 0.10 M acetic acid (Ka = 1.8 × 10-5).

Acid-Base Equilibria and Buffers

• Buffer: Solution that resists changes in pH upon addition of small amounts of acid or base. Composed of a weak acid and its conjugate base (or vice versa).

• Henderson-Hasselbalch equation:

$pH = pK_a + \log \left(\frac{[A^-]}{[HA]}\right)$

Chapter 15: Applications of Aqueous Equilibria

Buffer Solutions and Titrations

This chapter explores practical uses of aqueous equilibria, including buffer systems and titration curves.

• Buffer capacity: Amount of acid/base a buffer can neutralize before pH changes significantly.

• Titration: Gradual addition of one solution to another to determine concentration.

• Acid-base titration curve: Plots pH vs. volume of titrant added; equivalence point is where moles of acid = moles of base.

Solubility and Precipitation

• Solubility product (Ksp): Equilibrium constant for dissolution of a sparingly soluble salt.

• Example: For AgCl: $K_{sp} = [Ag^+][Cl^-]$

• Common ion effect: Solubility of a salt decreases in the presence of a common ion.

Chapter 16: Thermodynamics—Entropy, Free Energy, and Equilibrium

Energy, Entropy, and Spontaneity

This chapter covers the laws of thermodynamics, entropy, and the criteria for spontaneous processes.

• First law: Energy is conserved.

• Second law: Spontaneous processes increase the entropy of the universe.

• Entropy (S): Measure of disorder.

• Gibbs free energy (G): Determines spontaneity: $\Delta G = \Delta H - T\Delta S$

• Spontaneous process: $\Delta G < 0$

Equilibrium and Thermodynamic Relationships

• Relationship between free energy and equilibrium constant:

$\Delta G^\circ = -RT \ln K$

• Summary of formulas: $\Delta S_{univ} = \Delta S_{sys} + \Delta S_{surr}$; $\Delta G = \Delta H - T\Delta S$; $\Delta G^\circ = -RT \ln K$

Chapter 17: Electrochemistry

Redox Reactions and Electrochemical Cells

This chapter introduces oxidation-reduction reactions and the construction of electrochemical cells.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Galvanic (voltaic) cell: Converts chemical energy to electrical energy; spontaneous reaction.

• Cell potential (Ecell): Measured in volts; calculated from standard reduction potentials.

Cell potential equation:

$E_{cell} = E_{cathode} - E_{anode}$

• Nernst equation: Relates cell potential to concentrations:

$E = E^\circ - \frac{0.0592}{n} \log Q$

Chapter 19: The Main-Group Elements

Properties and Trends of Main-Group Elements

This chapter surveys the chemistry of Groups 1-2 and 13-18, focusing on periodic trends and reactivity.

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

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

• Group 1 (alkali metals): Highly reactive, form +1 ions.

• Group 2 (alkaline earth metals): Less reactive than Group 1, form +2 ions.

• Group 17 (halogens): Very reactive nonmetals, form -1 ions.

• Group 18 (noble gases): Inert, very low reactivity.

Chapter 20: Transition Elements and Coordination Chemistry

Transition Metals and Coordination Compounds

This chapter discusses the unique properties of transition metals and the structure of coordination complexes.

• Transition metals: Groups 3-12; partially filled d orbitals.

• Variable oxidation states: Many transition metals can adopt multiple oxidation states.

• Coordination compound: Contains a central metal ion bonded to ligands.

• Ligand: Ion or molecule that donates electron pairs to the metal.

• Crystal field theory: Explains color and magnetism of complexes.

Property	Main-Group Elements	Transition Elements
Oxidation States	Usually one	Multiple possible
Color	Usually colorless	Often colored
Magnetism	Usually diamagnetic	Often paramagnetic

Example: [Fe(CN)6]4- is a coordination complex with Fe2+ and six cyanide ligands.

Additional info: Crystal field splitting and ligand field strength determine the color and magnetic properties of transition metal complexes.