General Chemistry II: Aqueous Equilibria, Thermodynamics, Electrochemistry, and Main Group/Transition Elements

Chapter 14 - Aqueous Equilibria: Acids & Bases

1. What Are Acids and Bases?

Acids and bases are fundamental chemical species that participate in a wide range of chemical reactions, especially in aqueous solutions. Their behavior is described by several definitions and models.

• Arrhenius acid: Produces H+ in water ($\text{HA} \rightarrow \text{H}^+ + \text{A}^-$)

• Arrhenius base: Produces OH- in water ($\text{BOH} \rightarrow \text{B}^+ + \text{OH}^-$)

• Bronsted-Lowry acid: Proton donor

• Bronsted-Lowry base: Proton acceptor

Conjugate pairs: Every acid (HA) has a conjugate base (A-); every base (B) has a conjugate acid (BH+).

2. Strong vs Weak Acids/Bases

Acids and bases are classified by their degree of ionization in water.

• Strong acids/bases: Dissociate 100% (e.g., HCl, HNO3, NaOH)

• Weak acids/bases: Partial dissociation (e.g., CH3COOH, NH3)

Only weak acids/bases establish equilibrium in solution.

3. pH, pOH, and Kw

The pH scale quantifies the acidity or basicity of a solution. The relationship between pH, pOH, and the ion product of water ($K_w$) is essential for calculations.

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

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

• Relationship: $\text{pH} + \text{pOH} = 14$ (at 25°C)

• Water dissociation constant: $K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}$

4. Calculating pH

pH calculations depend on the strength of the acid/base and the concentration.

• Strong acids/bases: $\text{pH} = -\log[\text{H}^+]$

• Weak acids/bases: Use $K_a$ or $K_b$ and ICE tables to solve for [H+]

5. Buffers (Intro)

Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base.

• Made from weak acid + conjugate base

• Buffer capacity: Higher when [acid] ≈ [base]

Chapter 15 - Applications of Aqueous Equilibria

1. Buffers (in Depth)

Buffers are most effective when acid and base are both present in similar concentrations.

• Henderson-Hasselbalch equation: $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{base}]}{[\text{acid}]}\right)$

2. Acid-Base Titrations

Titrations are used to determine the concentration of an acid or base by neutralization.

• Equivalence point: Acid = base (moles)

• Half-equivalence point: $\text{pH} = \text{p}K_a$ for weak acid titration

3. Solubility & Ksp

Solubility product constant ($K_{sp}$) quantifies the equilibrium between a solid and its ions in solution.

• Expression: $K_{sp} = [\text{cation}]^m[\text{anion}]^n$

• Common ion effect: Addition of a common ion decreases solubility

Chapter 16 - Thermodynamics: Entropy & Free Energy

1. Entropy (S)

Entropy is a measure of disorder or randomness in a system.

• Second Law: All spontaneous processes increase the entropy of the universe ($\Delta S_{universe} > 0$)

2. Gibbs Free Energy (G)

Gibbs free energy determines the spontaneity of a process at constant temperature and pressure.

• Equation: $\Delta G = \Delta H - T\Delta S$

• Interpretation: $\Delta G < 0$ is spontaneous; $\Delta G = 0$ is equilibrium; $\Delta G > 0$ is nonspontaneous

3. ΔG and Equilibrium

• Relationship: $\Delta G^\circ = -RT \ln K$

Chapter 17 - Electrochemistry

1. Redox Basics

Redox reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

2. Galvanic (Voltaic) Cells

Galvanic cells use spontaneous redox reactions to generate electricity.

• Key facts: Anode = oxidation; Cathode = reduction; Electrons flow anode → cathode

3. Cell Potential

• Equation: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

• Free energy relation: $\Delta G^\circ = -nFE^\circ_{cell}$

4. Nernst Equation

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

Chapter 19 - The Main-Group Elements

Periodic Trends

Main-group elements show predictable trends in reactivity and properties across the periodic table.

Group	Highlights
Group 1	Alkali metals, very reactive, +1 charge
Group 2	Alkaline earth metals, +2 charge
Group 17	Halogens, -1 charge, very reactive nonmetals
Group 18	Noble gases, mostly unreactive

Chapter 20 - Transition Metals & Coordination Chemistry

1. Transition Metal Properties

Transition metals have variable oxidation states and form colored compounds due to d-orbital electron transitions.

• Variable oxidation states

• Colored compounds

• Magnetic behavior

• Catalysis

2. Coordination Compounds

Coordination compounds consist of a central metal ion bonded to surrounding ligands.

• Ligand: Ion or molecule that binds to metal

• Coordination number: Number of ligands

• Example: [Fe(CN)6]4-

3. Crystal Field Theory (Simplified)

Crystal field theory explains the color and magnetism of transition metal complexes.

• Ligands split d-orbital energies

• Strong field ligands = low spin; weak field ligands = high spin

Final Exam Strategy & Common Mistakes

• Memorize key equations and periodic trends

• Understand concepts, not just facts

• Practice problems by chapter

• Common mistakes: Forgetting to use significant figures, misidentifying anode/cathode, confusing equivalence points

Must-Memorize Equations

• $\text{pH} = -\log[\text{H}^+]$

• $\text{pOH} = -\log[\text{OH}^-]$

• $\text{pH} + \text{pOH} = 14$

• $K_w = [\text{H}^+][\text{OH}^-]$

• $\text{pH} = \text{p}K_a + \log\left(\frac{[\text{base}]}{[\text{acid}]}\right)$

• $K_{sp} = [\text{cation}]^m[\text{anion}]^n$

• $\Delta G = \Delta H - T\Delta S$

• $\Delta G^\circ = -RT \ln K$

• $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

• $\Delta G^\circ = -nFE^\circ_{cell}$

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

Practice Final-Style Questions (Examples)

• Calculate pH of 0.01 M HCl ($\text{pH} = 2$)

• What is the conjugate base of H2SO4? ($\text{HSO}_4^-$)

• At half-equivalence of a weak acid titration, what is pH equal to? ($\text{pH} = \text{p}K_a$)

• What happens to solubility of AgCl if NaCl is added? (Decreases, common ion effect)

• What is the sign of $\Delta S$ for the reaction: ice → liquid water? (Positive, more disorder)

• Where does oxidation occur in a voltaic cell? (Anode)

• Which element has higher electronegativity: N or Cl? (Cl)