Unit 8: Acids and Bases – General Chemistry Study Notes

Unit 8: Acids and Bases

Overview

This unit explores the chemistry of acids and bases, focusing on their equilibrium behavior, calculations involving pH and pOH, titration analysis, buffer systems, and the molecular structure underlying acid and base strength. The concepts are foundational for understanding chemical equilibria and solution chemistry in general chemistry.

8.1 Introduction to Acids and Bases

pH, pOH, and Water Autoionization

• pH is a measure of the hydronium ion concentration in aqueous solution: $pH = -\log[H_3O^+]$

• pOH is a measure of the hydroxide ion concentration: $pOH = -\log[OH^-]$

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

• The equilibrium constant for water at 25°C is $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

• At 25°C, $pK_w = 14 = pH + pOH$; in pure water, $pH = pOH = 7.0$

• The value of $K_w$ is temperature dependent; at higher temperatures, $K_w$ increases and neutral pH decreases.

Example: In pure water at 0°C, $[H_3O^+] = [OH^-] = 3.376 \times 10^{-8}$ M, so $pH = pOH = 7.47$ and $K_w = 1.14 \times 10^{-15}$.

8.2 pH and pOH of Strong Acids and Bases

Strong Acid and Base Calculations

• Strong acids (e.g., HCl, HBr, HNO3) ionize completely in water: $[H_3O^+] \approx$ initial acid concentration.

• Strong bases (e.g., NaOH, Ba(OH)2) dissociate completely: $[OH^-]$ equals the number of hydroxide ions per formula unit times the base concentration.

• Mixing strong acids and bases results in neutralization; the pH depends on the excess reagent.

Example: Mixing 50 mL of 0.50 M HCl with 50 mL of 0.30 M NaOH leaves excess acid; calculate the new concentration and pH accordingly.

8.3 Weak Acid and Base Equilibria

Equilibrium and Calculations for Weak Acids and Bases

• Weak acids only partially ionize: $HA(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + A^-(aq)$

• Acid dissociation constant: $K_a = \frac{[H_3O^+][A^-]}{[HA]}$; $pK_a = -\log K_a$

• Weak bases only partially ionize: $B(aq) + H_2O(l) \rightleftharpoons HB^+(aq) + OH^-(aq)$

• Base dissociation constant: $K_b = \frac{[HB^+][OH^-]}{[B]}$; $pK_b = -\log K_b$

• Percent ionization: $\% = \frac{[H_3O^+]}{[HA]} \times 100$ for acids

Example: For 0.15 M acetic acid ($K_a = 1.8 \times 10^{-5}$), $pH = 2.78$.

8.4 Acid-Base Reactions and Buffers

Types of Acid-Base Reactions

• Strong acid + strong base: $H^+ + OH^- \rightarrow H_2O$; pH determined by excess reagent.

• Weak acid + strong base: $HA + OH^- \rightleftharpoons A^- + H_2O$; buffer forms if weak acid is in excess; use Henderson-Hasselbalch equation.

• Weak base + strong acid: $B + H_3O^+ \rightleftharpoons HB^+ + H_2O$; buffer forms if weak base is in excess.

• Weak acid + weak base: $HA + B \rightleftharpoons A^- + HB^+$; equilibrium established.

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

8.5 Acid-Base Titrations

Titration Curves and Analysis

• Titration: Laboratory technique to determine the concentration of an unknown acid or base by reacting it with a known titrant.

• Equivalence point: Moles of titrant equal moles of analyte; for strong acid/base titrations, pH is neutral at equivalence.

• Half-equivalence point: For weak acids/bases, $[HA] = [A^-]$ and $pH = pK_a$.

• Polyprotic acids: Multiple equivalence points; each corresponds to the loss of a proton.

8.6 Molecular Structure of Acids and Bases

Structure-Strength Relationships

• Acid and base strength can be inferred from molecular structure.

• Strong acids have very weak, stabilized conjugate bases (by electronegativity, resonance, inductive effects).

• Carboxylic acids (–COOH) are common weak acids.

• Strong bases (group I/II hydroxides) have very weak conjugate acids.

• Nitrogenous bases (amines) are common weak bases.

Example: The stability of conjugate bases increases with resonance and inductive effects (e.g., ClO4− is more stable than ClO3−).

8.7 pH and pKa

Protonation State and Indicators

• The ratio of acid to base forms in solution depends on the relationship between pH and pKa.

• If pH < pKa, the acid form predominates; if pH > pKa, the base form predominates.

• Acid-base indicators change color depending on their protonation state, which is pH-dependent.

Indicator	Color Change	pH Range	pKa
Bromophenol blue	Yellow to blue	3 - 4.5	4
Bromocresol green	Yellow to blue	4 - 5	4.5
Methyl red	Red to yellow	4.5 - 6	5
Bromothymol blue	Yellow to blue	6 - 7.5	7
Phenol red	Yellow to red	7 - 8	7.5
Phenolphthalein	Colorless to pink	8 - 10	9
Alizarin yellow R	Yellow to red	10 - 12	11

8.8 Properties of Buffers

Buffer Action and Preparation

• A buffer contains significant concentrations of both a weak acid and its conjugate base (or weak base and conjugate acid).

• Buffers resist changes in pH when small amounts of strong acid or base are added.

• Preparation methods include mixing a weak acid with its salt, or partially neutralizing a weak acid/base with a strong base/acid.

8.9 Henderson-Hasselbalch Equation

Buffer Calculations

• The Henderson-Hasselbalch equation relates buffer pH to the ratio of conjugate base to acid:

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

• Adding small amounts of acid or base to a buffer does not significantly change the pH.

8.10 Buffer Capacity

Factors Affecting Buffer Capacity

• Buffer capacity increases with the total concentration of buffer components, not their ratio.

• If [acid] > [base], the buffer can neutralize more added base; if [base] > [acid], it can neutralize more added acid.

• Buffer capacity is the amount of acid or base a buffer can neutralize before a significant pH change occurs.

8.11 pH and Solubility

Effect of pH on Salt Solubility

• The solubility of salts containing weak acids or bases is affected by pH.

• For basic salts (e.g., Mg(OH)2), increasing [H+] (lower pH) increases solubility by shifting equilibrium right.

• For acidic salts (e.g., HNO2), increasing [OH−] (higher pH) increases solubility by removing H+.

• Le Châtelier’s principle explains these shifts in equilibrium.

Example: Mg(OH)2 is more soluble in acidic solutions; HNO2 is more soluble in basic solutions.