Chapter 17: Acids and Bases – Comprehensive Study Notes

Acids and Bases: Definitions and Concepts

Acid-Base Definitions

Acids and bases can be defined in several ways, each with increasing generality. Understanding these definitions is fundamental to classifying substances and predicting their behavior in chemical reactions.

• Arrhenius Definition:

  • Acid: Produces H+ ions in aqueous solution. Example: HCl(aq) → H+(aq) + Cl–(aq)

  • Base: Produces OH– ions in aqueous solution. Example: NaOH(aq) → Na+(aq) + OH–(aq)

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor

  • Base: Proton (H+) acceptor

• Lewis Definition:

  • Acid: Electron-pair acceptor

  • Base: Electron-pair donor

Note: All Arrhenius acids are also Brønsted-Lowry acids, and all Brønsted-Lowry acids are also Lewis acids.

Acid and Base Strength

Strength and Dissociation in Water

The strength of an acid or base refers to its degree of ionization or dissociation in water.

• Strong acids: Dissociate completely in water (irreversible reaction). Examples: HCl, HBr, HI, HNO3, H2SO4, HClO4

• Weak acids: Dissociate partially (reversible, equilibrium reaction). Examples: HF, CH3COOH, HCN, H2CO3

• Strong bases: Dissociate completely. Examples: NaOH, KOH, Ca(OH)2, Ba(OH)2

• Weak bases: React partially with water. Examples: NH3, amines, anions of weak acids

Relative Strengths and Conjugate Pairs

Acid and base strength is inversely related to the strength of their conjugate partners.

• Acid strength order (strongest to weakest): HI > HBr > HCl > HNO3 > H2SO4 > HF > CH3COOH > H2CO3 > HCN > H2O

• Key rule: The stronger the acid, the weaker its conjugate base (and vice versa).

• Example: Cl– (conjugate base of HCl) is essentially neutral; CH3COO– (conjugate base of acetic acid) is a weak base.

Acid Dissociation Constant (Ka) and pKa

The acid dissociation constant quantifies the extent of acid dissociation in water.

• General reaction: HA(aq) + H2O(l) ↔ H3O+(aq) + A–(aq)

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• Larger Ka → stronger acid; lower pKa → stronger acid

• $pK_a = -\log(K_a)$

Acid	Ka	pKa
HF	6.8 × 10–4	3.17
CH3COOH	1.8 × 10–5	4.74
HCN	4.9 × 10–10	9.31
H2O	1.0 × 10–14	14.00

Water: Autoionization and Amphiprotism

Autoionization of Water and Kw

Water can ionize to form hydronium and hydroxide ions, even in pure samples.

• Reaction: H2O(l) + H2O(l) ↔ H3O+(aq) + OH–(aq)

• $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ (at 25 °C)

• In pure water: [H3O+] = [OH–] = 1.0 × 10–7 M → pH = 7 (neutral)

• pH = –log[H3O+]; pOH = –log[OH–]; pH + pOH = 14 (at 25 °C)

Amphiprotism of Water

Water can act as either an acid or a base, depending on the reaction partner.

• As a base: HCl + H2O → Cl– + H3O+

• As an acid: NH3 + H2O → NH4+ + OH–

• Other amphiprotic species: HCO3–, H2PO4–, HPO42–

Conjugate Acid-Base Pairs and Equilibrium

Conjugate Acid-Base Pairs

Each Brønsted-Lowry acid-base reaction involves two conjugate pairs, differing by one proton.

• Acid loses H+ → conjugate base; base gains H+ → conjugate acid

• Example: CH3COOH + H2O ↔ H3O+ + CH3COO–

• Pairs: CH3COOH / CH3COO– and H3O+ / H2O

Equilibrium Favorability

Acid-base reactions favor the formation of the weaker acid and base (more stable, lower energy products).

• Equilibrium lies toward the side with weaker acid/base (Kc >> 1)

• Example: HCl + NaOH → H2O + NaCl; Kc ~ 1014

Quantitative Aspects of Acid-Base Chemistry

Percent Dissociation of Weak Acids

The percent of acid molecules that ionize increases as the acid is diluted.

• $\%\ \text{dissociation} = \left( \frac{[H_3O^+]_{eq}}{[HA]_{initial}} \right) \times 100\%$

• As [HA]initial decreases, % dissociation increases (Le Chatelier's principle)

• Example: 1.0 M acetic acid → 0.42% dissociated; 0.01 M → 4.2% dissociated

Polyprotic Acids

Polyprotic acids donate protons in sequential steps, each with its own Ka value.

• Example: Phosphoric acid (H3PO4):

  • First dissociation: H3PO4 ↔ H+ + H2PO4–; Ka1 = 7.5 × 10–3

  • Second: H2PO4– ↔ H+ + HPO42–; Ka2 = 6.2 × 10–8

  • Third: HPO42– ↔ H+ + PO43–; Ka3 = 4.8 × 10–13

• Ka1 >> Ka2 >> Ka3; first step provides nearly all [H3O+]

• H2SO4 is a special case: first ionization is complete (strong), second is weak (Ka2 = 1.2 × 10–2)

Factors Affecting Acid Strength

Electronegativity, Bond Polarity, and Bond Strength

Several structural factors influence acid strength:

• Binary acids (H–X):

  • Down a group: Bond strength decreases, acid strength increases (HI > HBr > HCl > HF)

  • Across a period: Electronegativity increases, bond polarity increases, acid strength increases (HF > H2O > NH3)

• Oxyacids: More oxygen atoms increase acid strength (HClO4 > HClO3 > HClO2 > HClO)

• Resonance: Delocalization of charge in the conjugate base stabilizes it, increasing acid strength

• Inductive effects: Electronegative substituents near the acidic H stabilize the conjugate base, increasing acid strength

Bases: Kb and pKb

Base Dissociation Constant (Kb) and pKb

The base dissociation constant quantifies the extent of base ionization in water.

• General reaction: B(aq) + H2O(l) ↔ BH+(aq) + OH–(aq)

• $K_b = \frac{[BH^+][OH^-]}{[B]}$

• Larger Kb → stronger base; lower pKb → stronger base

• $pK_b = -\log(K_b)$

• Example: NH3: Kb = 1.8 × 10–5, pKb = 4.74; methylamine (CH3NH2): Kb = 4.4 × 10–4

Relationship Between Ka and Kb

• For any conjugate acid-base pair at 25 °C:

  • $K_a \times K_b = K_w = 1.0 \times 10^{-14}$

  • $pK_a + pK_b = 14$

• Stronger acid (larger Ka) → weaker conjugate base (smaller Kb)

• To find Kb of A– from Ka of HA:

  • $K_b = \frac{K_w}{K_a}$

• Example: Ka(CH3COOH) = 1.8 × 10–5 → Kb(CH3COO–) = 5.6 × 10–10

Buffer Solutions and the Henderson-Hasselbalch Equation

Relative Concentrations and pH

The Henderson-Hasselbalch equation relates pH to the ratio of conjugate base to acid in solution.

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

• If [HA] > [A–]: pH < pKa (acidic)

• If [HA] = [A–]: pH = pKa (half-equivalence point)

• If [A–] > [HA]: pH > pKa (basic relative to pKa)

• Example: [CH3COO–] = 0.10 M, [CH3COOH] = 0.010 M → pH = 4.74 + 1.0 = 5.74

Buffer solutions are mixtures of a weak acid and its conjugate base (or weak base and its conjugate acid) that resist changes in pH.

Weak Bases in Water

Ammonia, Amines, and Weak Acid Anions

These species act as weak bases by accepting protons from water, generating OH– ions.

• Ammonia: NH3 + H2O ↔ NH4+ + OH– (Kb = 1.8 × 10–5)

• Amines: RNH2 + H2O ↔ RNH3+ + OH– (generally stronger than NH3)

• Weak acid anions: e.g., CH3COO– + H2O ↔ CH3COOH + OH–

• Rule: The anion of a weak acid is a weak base; the anion of a strong acid is essentially neutral in water.

Acidic, Basic, and Neutral Salt Solutions

Salt Hydrolysis and Solution pH

The pH of a salt solution depends on the acid and base from which the salt is derived.

Acid + Base Combination	Example	Ion Behavior in Water	pH
Strong acid + Strong base	NaCl	Na+ & Cl– do not hydrolyze	Neutral (7)
Strong acid + Weak base	NH4Cl	NH4+ donates H+ to H2O (weak acid)	Acidic (< 7)
Weak acid + Strong base	NaCH3COO	CH3COO– accepts H+ from H2O (weak base)	Basic (> 7)
Weak acid + Weak base	NH4CH3COO	Compare Ka(cation) vs. Kb(anion): Ka > Kb → acidic; Kb > Ka → basic	Depends on Ka vs. Kb

Lewis Acids and Bases

Lewis Acid-Base Theory

The Lewis definition is the most general, focusing on electron pairs rather than protons.

• Lewis acid: Electron-pair acceptor (must have an empty or available orbital)

• Lewis base: Electron-pair donor (must have a lone pair)

• Examples:

  • BF3 + :NH3 → F3B–NH3 (BF3 = Lewis acid; NH3 = Lewis base)

  • Cu2+ + 4 :NH3 → [Cu(NH3)4]2+ (metal cation = Lewis acid)

  • H+ + :OH– → H2O (H+ is both Brønsted and Lewis acid)

Model	Acid Definition	Base Definition	Scope
Arrhenius	Produces H+	Produces OH–	Aqueous only
Brønsted-Lowry	H+ donor	H+ acceptor	Any solvent
Lewis	e– pair acceptor	e– pair donor	Most general

Additional info: These notes synthesize and expand upon the provided study guide, adding definitions, examples, and context for clarity and completeness.