Acids, Bases, and Aqueous Ionic Equilibria – Study Guide

Chapter 16: Acids and Bases

Introduction to Acids and Bases

This chapter explores the fundamental properties of acids and bases, their classification, and the quantitative relationships that govern their behavior in aqueous solutions. Understanding these concepts is essential for predicting chemical reactivity and equilibrium in a variety of chemical systems.

• Brønsted-Lowry Theory: Acids are proton (H+) donors, and bases are proton acceptors.

• Conjugate Acid-Base Pairs: Every acid has a conjugate base, and every base has a conjugate acid, formed by the loss or gain of a proton.

Properties and Identification of Acids and Bases

• Characteristic Properties: Acids typically taste sour, turn litmus paper red, and react with metals. Bases taste bitter, feel slippery, and turn litmus paper blue.

• Strong Acids: HCl, HBr, HI, HNO3, HClO4, H2SO4 (first ionization only).

• Weak Acids: HF, CH3COOH, H3PO4, HNO2, HCN.

• Strong Bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 (contain Group IA or IIA metal cations).

• Weak Bases: NH3, CH3NH2.

Quantitative Aspects of Acid-Base Chemistry

• pKa and pKb: The strength of acids and bases is often expressed as pKa and pKb. Lower pKa indicates a stronger acid; lower pKb indicates a stronger base.

• Relationship between Ka and Kb: For a conjugate acid-base pair: $K_a \times K_b = K_w$ where $K_w$ is the ion-product constant for water ($1.0 \times 10^{-14}$ at 25°C).

• pH and pOH Calculations: $pH = -\log[H^+]$ $pOH = -\log[OH^-]$ $pH + pOH = 14$ (at 25°C)

• Calculating [H+] or [OH-] from pH: $[H^+] = 10^{-pH}$ $[OH^-] = 10^{-pOH}$

• Equilibrium Expressions: For a weak acid HA: $K_a = \frac{[H^+][A^-]}{[HA]}$

Applications and Trends

• Predicting Acid/Base Strength: The direction of acid-base equilibrium favors the formation of the weaker acid and base.

• Polyprotic Acids: These acids can donate more than one proton, each with its own Ka value. The first ionization is always the strongest.

• Salts in Solution: Predict whether a salt will form an acidic, basic, or neutral solution based on the properties of its constituent ions.

Example: Calculating the pH of a Weak Acid Solution

Given 0.10 M acetic acid (CH3COOH), Ka = 1.8 × 10-5:

• Set up the equilibrium: CH3COOH ⇌ H+ + CH3COO-

• Write the Ka expression and solve for [H+]: $K_a = \frac{[H^+][CH_3COO^-]}{[CH_3COOH]}$ Assume x = [H+] at equilibrium, solve for x.

• Calculate pH: $pH = -\log(x)$

Chapter 17: Aqueous Ionic Equilibria (Buffers and Titrations)

Introduction to Buffers and Ionic Equilibria

This chapter focuses on buffer solutions, their preparation, and their behavior during acid-base titrations. Buffers are crucial for maintaining pH stability in chemical and biological systems.

Buffers and the Common Ion Effect

• Buffer Definition: A buffer is a solution that resists changes in pH upon addition of small amounts of acid or base. It typically contains a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Common Ion Effect: Adding a common ion suppresses the ionization of a weak acid or base, shifting the equilibrium according to Le Chatelier's principle.

• Henderson-Hasselbalch Equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$ Used to calculate the pH of buffer solutions.

Buffer Capacity and Selection

• Buffer Capacity: The amount of acid or base a buffer can neutralize before a significant pH change occurs.

• Choosing a Buffer: Select a weak acid/base pair with a pKa close to the desired pH.

Titrations and pH Calculations

• Titration Curves: The shape of the titration curve depends on the strength of the acid and base involved (strong/weak combinations).

• pH at Various Points: Calculate pH before, during, and after the equivalence point using stoichiometry and equilibrium concepts.

• Types of Titrations:

• Strong acid with strong base

• Strong base with strong acid

• Weak acid with strong base

• Weak base with strong acid

Example: Buffer pH Calculation

Given a buffer with 0.20 M acetic acid and 0.10 M sodium acetate (pKa = 4.76):

• Use the Henderson-Hasselbalch equation: $pH = 4.76 + \log\left(\frac{0.10}{0.20}\right) = 4.76 + \log(0.5) = 4.76 - 0.30 = 4.46$

Table: Common Strong and Weak Acids and Bases

Strong Acids	Weak Acids	Strong Bases	Weak Bases
HCl, HBr, HI, HNO3, HClO4, H2SO4 (1st ionization)	HF, CH3COOH, H3PO4, HNO2, HCN	LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2	NH3, CH3NH2