Chapter 17: Acids and Bases – Study Guide

Chapter 17: Acids and Bases

17.1 Heartburn

Heartburn is a common condition caused by excess stomach acid irritating the esophagus. Understanding acids and bases helps explain the chemistry behind antacids and their effectiveness in neutralizing stomach acid.

• Heartburn: Occurs when gastric acid flows back into the esophagus, causing discomfort.

• Antacids: Substances that neutralize excess stomach acid, often containing weak bases like magnesium hydroxide or calcium carbonate.

• Chemical Reaction Example: Neutralization of hydrochloric acid by calcium carbonate: $\mathrm{2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)}$

17.2 The Nature of Acids and Bases

Acids and bases are fundamental chemical species with distinct properties and behaviors in aqueous solutions.

• Acids: Substances that taste sour, turn blue litmus paper red, and react with metals to produce hydrogen gas.

• Bases: Substances that taste bitter, feel slippery, turn red litmus paper blue, and react with acids to form water and salts.

• Examples: Hydrochloric acid (HCl), sulfuric acid (H2SO4), sodium hydroxide (NaOH), ammonia (NH3).

17.3 Definitions of Acids and Bases

There are several models for defining acids and bases, each broadening the scope of acid-base chemistry.

• Arrhenius Definition:

  • Acid: Produces H+ ions in aqueous solution.

  • Base: Produces OH− ions in aqueous solution.

  • Example: HCl (acid), NaOH (base).

• Brønsted-Lowry Definition:

  • Acid: Proton (H+) donor.

  • Base: Proton (H+) acceptor.

  • Example: NH3 + H2O → NH4+ + OH−

• Conjugate Acid-Base Pairs:

  • When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

  • Example: In the reaction above, NH3 is the base, NH4+ is its conjugate acid; H2O is the acid, OH− is its conjugate base.

17.4 Acid Strength and Acid Ionization Constant (Ka)

Acids vary in their ability to donate protons, which is quantified by their ionization constant.

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

• Weak Acids: Partially dissociate in water (e.g., acetic acid, HF).

• Acid Ionization Constant (Ka): Measures the extent of acid dissociation: $K_a = \frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

• Relationship: Larger Ka = stronger acid; smaller Ka = weaker acid.

17.5 Autoionization of Water and pH

Water can ionize to a small extent, producing hydronium and hydroxide ions, which is the basis for the pH scale.

• Autoionization of Water: $2\mathrm{H_2O(l)} \rightleftharpoons \mathrm{H_3O^+(aq)} + \mathrm{OH^-(aq)}$

• Ion Product Constant for Water (Kw): $K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] = 1.0 \times 10^{-14}$ (at 25°C)

• pH and pOH:

  • $\mathrm{pH} = -\log_{10}[\mathrm{H_3O^+}]$

  • $\mathrm{pOH} = -\log_{10}[\mathrm{OH^-}]$

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

• Relationship: Lower pH = higher acidity; higher pH = more basic.

17.6 Finding the [H3O+] and the pH of Strong and Weak Acid Solutions

Calculating the concentration of hydronium ions and pH depends on whether the acid is strong or weak.

• Strong Acids: [H3O+] equals the initial acid concentration.

• Weak Acids: Use an ICE table and Ka to solve for [H3O+].

• Relationship between Ka and Acid Strength: Higher Ka means stronger acid.

• Relationship between pKa and Acid Strength: Lower pKa means stronger acid. $\mathrm{p}K_a = -\log_{10} K_a$

• Example Calculation: For 0.10 M HCl (strong acid): [H3O+] = 0.10 M; pH = 1.0

17.7 Base Solutions

Bases can be strong or weak, and their strength is quantified by the base ionization constant.

• Base Ionization Constant (Kb): $K_b = \frac{[\mathrm{BH^+}][\mathrm{OH^-}]}{[\mathrm{B}]}$

• Relationship: Larger Kb = stronger base; smaller Kb = weaker base.

• Relationship between pKb and Base Strength: Lower pKb = stronger base. $\mathrm{p}K_b = -\log_{10} K_b$

• Calculating [OH−]: For strong bases, [OH−] equals the initial base concentration.

• Example: 0.10 M NaOH (strong base): [OH−] = 0.10 M; pOH = 1.0; pH = 13.0

17.8 The Acid-Base Properties of Ions and Salts

Ions from salts can affect the pH of a solution depending on their origin from strong or weak acids and bases.

• Predicting Salt Solution pH:

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

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

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

• Relationship of Ka, Kb, and Kw: $K_a \times K_b = K_w$

• Relationship of pKa, pKb, and pKw: $\mathrm{p}K_a + \mathrm{p}K_b = \mathrm{p}K_w = 14$ (at 25°C)

• Example: Ammonium chloride (NH4Cl) forms an acidic solution because NH4+ is a weak acid.

17.9 Polyprotic Acids

Polyprotic acids can donate more than one proton, with each ionization step having its own Ka value.

• Polyprotic Acid: An acid with more than one ionizable proton (e.g., H2SO4, H3PO4).

• Ionization Steps: Each proton is lost in a separate step, with decreasing Ka values: Ka1 > Ka2 > Ka3

• Example: Sulfuric acid (H2SO4): First ionization is strong; second is weak.

17.10 Acid Strength and Molecular Structure

The strength of an acid is influenced by its molecular structure, including factors such as bond strength, electronegativity, and resonance stabilization.

• Bond Strength: Weaker H–A bonds make it easier to donate H+.

• Electronegativity: Higher electronegativity of A stabilizes the negative charge, increasing acid strength.

• Resonance: Delocalization of charge via resonance stabilizes the conjugate base, increasing acid strength.

• Inductive Effect: Electronegative atoms near the acidic proton withdraw electron density, stabilizing the conjugate base.

• Formal Charge: The location and magnitude of formal charge affect acid strength.

• Example: Acetic acid is stronger than ethanol due to resonance stabilization of its conjugate base.

17.11 The Lewis Acid-Base Model

The Lewis model expands the definition of acids and bases to include electron pair transfer.

• Lewis Acid: Electron pair acceptor.

• Lewis Base: Electron pair donor.

• Comparison:

  • Arrhenius: H+ or OH− in water.

  • Brønsted-Lowry: Proton transfer.

  • Lewis: Electron pair transfer.

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

Additional info: Practice problems are recommended for each section to reinforce understanding. For calculations involving weak acids and bases, use ICE tables and quadratic equations as needed. For polyprotic acids, the first ionization usually dominates the pH calculation.