Chapter 11: Acids and Bases (Sections 11.1–11.7) – Study Notes

Chapter 11: Acids and Bases

Section 11.1: Characteristics and Nomenclature of Acids and Bases

This section introduces the fundamental properties and naming conventions for acids and bases, essential for understanding their behavior in chemical reactions.

• Properties of Acids:

  • Have a sour taste

  • Produce a stinging sensation

  • Turn litmus paper red

  • Neutralize bases

  • Corrode some metals

• Properties of Bases:

  • Have a bitter (chalky) taste

  • Feel soapy or slippery

  • Turn litmus paper blue

  • Neutralize acids

• Common Acids: Hydrochloric acid (HCl), sulfuric acid (H2SO4), nitric acid (HNO3), acetic acid (CH3COOH), carbonic acid (H2CO3)

• Common Bases: Sodium hydroxide (NaOH), potassium hydroxide (KOH), ammonia (NH3)

Naming Acids

• Rule 1: Acids with hydrogen and a nonmetal: use the prefix hydro- and the suffix -ic acid (e.g., HCl = hydrochloric acid).

• Rule 2: Acids with hydrogen and a polyatomic ion:

  • If the ion ends in -ate, change to -ic acid (e.g., ClO3- = chlorate → HClO3 = chloric acid).

  • If the ion ends in -ite, change to -ous acid (e.g., ClO2- = chlorite → HClO2 = chlorous acid).

• Exception: HCN is named hydrocyanic acid (follows Rule 1 even though CN- is polyatomic).

Naming Bases

• Bases are typically named as ionic compounds containing the polyatomic group hydroxide (OH-).

• Examples: NaOH = sodium hydroxide, KOH = potassium hydroxide, NH3 = ammonia (common name).

Monoprotic, Diprotic, and Triprotic Acids

• Monoprotic acids: Donate one H+ (e.g., HCl, CH3COOH)

• Diprotic acids: Donate two H+ (e.g., H2CO3, H2SO4)

• Triprotic acids: Donate three H+ (e.g., H3PO4)

Additional info: Bases can be classified as monobasic, dibasic, etc., based on the number of OH- ions they provide.

Section 11.1–11.2: Arrhenius and Brønsted–Lowry Definitions

These definitions describe acids and bases based on their behavior in water and their ability to donate or accept protons.

• Arrhenius Acid: Produces H+ ions in water (e.g., HCl → H+ + Cl-).

• Arrhenius Base: Produces OH- ions in water (e.g., NaOH → Na+ + OH-).

• Brønsted–Lowry Acid: Donates a proton (H+).

• Brønsted–Lowry Base: Accepts a proton (H+).

• Hydronium Ion (H3O+): The form in which H+ exists in aqueous solution.

Example: HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)

Section 11.2: Conjugate Acid–Base Pairs

Conjugate acid–base pairs are related by the gain or loss of a single proton (H+).

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

• Example: HF(aq) + H2O(l) ⇌ F-(aq) + H3O+(aq)

  • Pair 1: HF (acid) and F- (conjugate base)

  • Pair 2: H2O (base) and H3O+ (conjugate acid)

Section 11.2: Amphoteric Substances

Amphoteric substances can act as either acids or bases depending on the reaction partner.

• Example: HCO3- can donate H+ (acid) or accept H+ (base).

• Water (H2O): The most common amphoteric substance.

Section 11.3: Strengths of Acids and Bases

The strength of an acid or base depends on its degree of ionization in water.

• Strong Acids: Completely dissociate in water (e.g., HCl, HBr, HI, HNO3, HClO4, H2SO4).

• Weak Acids: Partially dissociate in water (e.g., HF, HCN, H2CO3, CH3COOH).

• Strong Bases: Completely dissociate in water (e.g., NaOH, KOH, Ba(OH)2).

• Weak Bases: Partially ionize in water (e.g., NH3, Mg(OH)2, amines).

Relative Strength and Direction of Acid–Base Reactions

• Strong acids have very weak conjugate bases.

• Reactions favor the formation of weaker acids and bases (from stronger ones).

• Example: H2SO4 + H2O → H3O+ + HSO4- (forward reaction favored).

Table: Relative Strengths of Acids and Bases

Acid	Conjugate Base
HCl	Cl-
HNO3	NO3-
HF	F-
CH3COOH	CH3COO-
H2O	OH-
NH4+	NH3

Additional info: The stronger the acid, the weaker its conjugate base, and vice versa.

Section 11.4: Dissociation Constants for Acids and Bases

Weak acids and bases establish equilibrium in water, described by dissociation constants.

• Acid Dissociation Constant (Ka):

  • For a weak acid: HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)

  • $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• Base Dissociation Constant (Kb):

  • For a weak base: B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)

  • $K_b = \frac{[BH^+][OH^-]}{[B]}$

• Interpretation:

  • Ka or Kb < 1: More reactants (weak acid/base)

  • Ka or Kb > 1: More products (strong acid/base)

pKa and pKb Values

• pKa = -log(Ka)

• Lower pKa = stronger acid

• Lower pKb = stronger base

Section 11.5: Dissociation of Water

Water self-ionizes to a small extent, producing hydronium and hydroxide ions.

• Equation: $2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$

• Water dissociation constant: $K_w = [H_3O^+][OH^-]$

• At 25°C, $K_w = 1.0 \times 10^{-14}$

• In pure water: $[H_3O^+] = [OH^-] = 1.0 \times 10^{-7}$ M

Section 11.6: The pH Scale

The pH scale quantifies the acidity or basicity of a solution using a logarithmic scale.

• $pH = -\log[H_3O^+]$

• pH < 7: Acidic solution

• pH = 7: Neutral solution

• pH > 7: Basic solution

• To find [H3O+] from pH: $[H_3O^+] = 10^{-pH}$

Example: If [H3O+] = 2.5 × 10-3 M, then pH = -log(2.5 × 10-3) ≈ 2.60

Section 11.7: Reactions of Acids and Bases

Acids and bases participate in several important types of chemical reactions.

• Neutralization: Acid + base → salt + water

  • Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

  • Net ionic equation: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$

• Reaction with Metals: Acid + metal → salt + H2(g)

  • Example: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

• Reaction with Carbonates/Bicarbonates: Acid + carbonate/bicarbonate → CO2(g) + salt + water

  • Example: 2HCl(aq) + CaCO3(s) → CO2(g) + CaCl2(aq) + H2O(l)

Application: Antacids neutralize excess stomach acid, often using weakly soluble hydroxides (e.g., Mg(OH)2).