Acid-Base Equilibria, Solubility, and Electrochemistry: Exam 3 Study Guide

Chapter 16: Acid-Base Equilibria

Classifications of Acids and Bases

Acids and bases can be classified according to their ability to donate or accept protons, and their strength in aqueous solution.

• Arrhenius Acids/Bases: Acids produce H+ in water; bases produce OH-.

• Brønsted–Lowry Acids/Bases: Acids are proton donors; bases are proton acceptors.

• Lewis Acids/Bases: Acids accept electron pairs; bases donate electron pairs.

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

Example: HCl is a strong acid; CH3COOH is a weak acid.

Conjugate Acid–Base Pairs

Every acid-base reaction involves two conjugate acid–base pairs.

• Conjugate Acid: The species formed when a base gains a proton.

• Conjugate Base: The species formed when an acid loses a proton.

Example: NH3 (base) + H2O (acid) ⇌ NH4+ (conjugate acid) + OH- (conjugate base)

Autoionization of Water

Water can act as both an acid and a base, undergoing autoionization:

• $2\mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{H}_3\mathrm{O}^+(aq) + \mathrm{OH}^-(aq)$

• The equilibrium constant is called the ion-product constant for water ($K_w$):

• $K_w = [\mathrm{H}_3\mathrm{O}^+][\mathrm{OH}^-] = 1.0 \times 10^{-14}$ at 25°C

Weak Acids and Weak Bases

Weak acids and bases only partially ionize in solution, establishing equilibrium.

• Acid Dissociation Constant ($K_a$): Measures acid strength.

• $K_a = \frac{[\mathrm{H}_3\mathrm{O}^+][\mathrm{A}^-]}{[\mathrm{HA}]}$

• Base Dissociation Constant ($K_b$): Measures base strength.

• $K_b = \frac{[\mathrm{BH}^+][\mathrm{OH}^-]}{[\mathrm{B}]}$

Example: For acetic acid, $K_a = 1.8 \times 10^{-5}$

Relationship Between $K_a$ and $K_b$

The strengths of conjugate acid–base pairs are related:

• $K_a \times K_b = K_w$

• For a conjugate acid–base pair, as $K_a$ increases, $K_b$ decreases.

Acid–Base Properties of Salt Solutions

Salts can affect the pH of a solution depending on the acid–base properties of their constituent ions.

• Salts from strong acid and strong base: neutral solution.

• Salts from strong base and weak acid: basic solution.

• Salts from strong acid and weak base: acidic solution.

Acid–Base Behavior and Chemical Structure

The strength of acids and bases is influenced by molecular structure, including bond polarity, bond strength, and resonance stabilization.

• For binary acids, acid strength increases with increasing bond polarity and decreasing bond strength.

• For oxoacids, acid strength increases with the number of oxygen atoms attached to the central atom.

Chapter 17: Buffers, Titrations, and Solubility Equilibria

The Common-Ion Effect and Buffers

The common-ion effect occurs when a solution contains two substances that share a common ion, suppressing the ionization of a weak acid or base.

• Buffer: A solution that resists changes in pH when small amounts of acid or base are added.

• Buffers are made from a weak acid and its conjugate base, or a weak base and its conjugate acid.

Preparation of Buffers

Buffers are prepared by mixing a weak acid with its salt (conjugate base) or a weak base with its salt (conjugate acid).

• Example: Acetic acid and sodium acetate.

Henderson–Hasselbalch Equation

The Henderson–Hasselbalch equation relates the pH of a buffer to the concentrations of acid and base:

• $\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\mathrm{A}^-]}{[\mathrm{HA}]} \right)$

Acid–Base Titrations and Titration Curves

Titrations are used to determine the concentration of an acid or base by reacting it with a standard solution. Titration curves plot pH versus volume of titrant added.

• Strong Acid–Strong Base: Sharp pH change at equivalence point (pH = 7).

• Strong Acid–Weak Base or Weak Acid–Strong Base: Equivalence point pH is not 7; buffer region present.

pH Calculations During Titration

• Before equivalence: Use stoichiometry and buffer equations.

• At equivalence: Only salt present; calculate pH from hydrolysis.

• After equivalence: Excess titrant determines pH.

Solubility Equilibria and Heterogeneous Equilibrium

Solubility equilibria involve the dissolution and precipitation of sparingly soluble salts.

• Solubility Product Constant ($K_{sp}$): $K_{sp} = [\text{cation}]^m [\text{anion}]^n$

• Heterogeneous equilibrium: Solid and aqueous phases coexist.

Applications of $K_{sp}$ and Factors Affecting Solubility

• Common-Ion Effect: Addition of a common ion decreases solubility.

• Effect of pH: Solubility of salts containing basic anions increases in acidic solution.

• Complex-Ion Formation: Increases solubility by forming soluble complexes.

• Temperature: Solubility generally increases with temperature for most solids.

Amphoteric Substances

Amphoteric substances can act as either acids or bases (e.g., Al(OH)3).

Precipitation, Separation of Ions, and Selective Precipitation

Precipitation occurs when the product of ion concentrations exceeds $K_{sp}$.

• Reaction Quotient ($Q$): Used to predict precipitation.

• If $Q > K_{sp}$, precipitation occurs; if $Q < K_{sp}$, no precipitation.

• Selective Precipitation: Used to separate ions by adding reagents that precipitate specific ions first.

Qualitative Analysis for Metallic Elements

Qualitative analysis involves identifying metal ions based on their chemical reactions and solubility properties.

• Grouping and identification based on selective precipitation and complex formation.

• Qualitative analysis focuses on presence/absence; quantitative analysis measures amount.

Chapter 20: Electrochemistry

Oxidation States and Redox Reactions

Redox (reduction–oxidation) reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons; increase in oxidation state.

• Reduction: Gain of electrons; decrease in oxidation state.

• Assign oxidation states to track electron transfer.

Balancing Redox Equations

Redox equations are balanced by separating into half-reactions and ensuring mass and charge balance.

• Balance atoms other than O and H.

• Balance O by adding H2O; balance H by adding H+ (or OH- in basic solution).

• Balance charge by adding electrons.

Voltaic Cells and Standard Cell Potentials

Voltaic (galvanic) cells use spontaneous redox reactions to generate electrical energy.

• Anode: Site of oxidation; electrons flow from anode to cathode.

• Cathode: Site of reduction.

• Cell potential ($E_{cell}$): $E_{cell} = E_{cathode} - E_{anode}$

• Standard cell potentials are measured under standard conditions (1 M, 1 atm, 25°C).

Free Energy and Redox

The relationship between cell potential and free energy is given by:

• $\Delta G = -nFE_{cell}$

• Where $n$ = number of moles of electrons, $F$ = Faraday's constant (96,485 C/mol).

Nonstandard Conditions (Nernst Equation)

The Nernst equation allows calculation of cell potential under nonstandard conditions:

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

• Where $Q$ is the reaction quotient.

Batteries

Batteries are practical applications of voltaic cells. Types include:

• Lead/Acid Battery: Used in cars; rechargeable.

• Alkaline Battery: Common household battery; non-rechargeable.

• Nickel/Cadmium (NiCd): Rechargeable; used in portable electronics.

• Lithium Ion: High energy density; used in modern electronics.

Corrosion

Corrosion is the deterioration of metals by redox reactions with the environment (e.g., rusting of iron).

• Prevention methods include coatings, sacrificial anodes, and alloying.

Battery Type	Electrodes	Electrolyte	Rechargeable?	Main Use
Lead/Acid	Pb, PbO2	H2SO4	Yes	Automobiles
Alkaline	Zn, MnO2	KOH	No	Household devices
Nickel/Cadmium	Cd, NiO(OH)	KOH	Yes	Portable electronics
Lithium Ion	LiC6, LiCoO2	Li+ salt in organic solvent	Yes	Modern electronics