Acid-Base Chemistry, Buffers, Titrations, and Solubility Equilibria: Study Notes

Acid-Base Chemistry

Introduction

Acid-base chemistry is a fundamental topic in general chemistry, focusing on the properties, reactions, and calculations involving acids, bases, and their solutions. Understanding how to determine pH, the strength of acids and bases, and related equilibrium concepts is essential for mastering this area.

• Solving for pH of Weak Base: Use the base dissociation constant ($K_b$) and equilibrium expressions to calculate the pH of a weak base solution.

• Solving for pH of Strong Base: Strong bases dissociate completely; calculate pOH from concentration, then convert to pH using $pH = 14 - pOH$.

• Solving for pH of Salt: Determine if the salt produces acidic, basic, or neutral solutions by considering the hydrolysis of its ions.

• Qualitatively Assessing pH of Salt: Salts from strong acid and strong base are neutral; from strong acid and weak base are acidic; from weak acid and strong base are basic.

• Strength of Acids and Bases: Strong acids/bases dissociate completely; weak acids/bases only partially dissociate. Common strong acids: HCl, HNO3, H2SO4. Common strong bases: NaOH, KOH.

• Solving for $K_a$ if given $K_b$ (and vice versa): Use the relationship $K_a \times K_b = K_w$, where $K_w = 1.0 \times 10^{-14}$ at 25°C.

• Knowing which to use, $K_a$ or $K_b$: Use $K_a$ for acids and $K_b$ for bases, depending on the species in question.

• $pK_a$, $pK_b$: $pK_a = -\log K_a$; $pK_b = -\log K_b$.

• Polyprotic Acids: Acids that can donate more than one proton (e.g., H2SO4); each dissociation step has its own $K_a$ value.

• pOH: $pOH = -\log [OH^-]$; $pH + pOH = 14$ at 25°C.

Example:

Calculate the pH of a 0.10 M solution of acetic acid ($K_a = 1.8 \times 10^{-5}$):

• Set up the equilibrium: $CH_3COOH \rightleftharpoons CH_3COO^- + H^+$

• Use ICE table and solve for [H+].

• Calculate pH: $pH = -\log[H^+]$

Buffers

Introduction

Buffers are solutions that resist changes in pH upon the addition of small amounts of acid or base. They are essential in many chemical and biological systems.

• Buffer Components: Typically consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Common Ion Effect: The suppression of ionization of a weak acid or base by the presence of a common ion from a strong electrolyte.

• Purpose of a Buffer: To maintain a relatively constant pH in a solution.

• Calculate the pH of a Buffer: Use the Henderson-Hasselbalch equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Buffer Calculations (RICE Table and Henderson-Hasselbalch): RICE tables help track concentrations; Henderson-Hasselbalch simplifies pH calculation for buffers.

• Effect of Adding H+ or OH- to a Buffer: The buffer neutralizes added acid or base, minimizing pH change.

• Buffer Range: The effective pH range is typically $pK_a \pm 1$.

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

• Preparing a Buffer: Mix a weak acid with its salt (conjugate base) or a weak base with its salt (conjugate acid).

Example:

Calculate the pH of a buffer containing 0.20 M acetic acid and 0.20 M sodium acetate ($K_a = 1.8 \times 10^{-5}$):

• $pK_a = -\log(1.8 \times 10^{-5}) = 4.74$

• $pH = 4.74 + \log(\frac{0.20}{0.20}) = 4.74$

Acid-Base Titrations

Introduction

Acid-base titrations are analytical techniques used to determine the concentration of an acid or base by reacting it with a standard solution of known concentration.

• Analyte: The solution of unknown concentration being analyzed.

• Titrant: The solution of known concentration added to react with the analyte.

• Indicators: Substances that change color at (or near) the equivalence point.

• Strong Acid-Strong Base Titration: The equivalence point occurs at pH 7.

• Weak Acid-Strong Base Titration: The equivalence point occurs at pH > 7 due to the formation of a weak base.

• Weak Base-Strong Acid Titration: The equivalence point occurs at pH < 7 due to the formation of a weak acid.

• Polyprotic Acid Titration: Multiple equivalence points corresponding to each ionizable proton.

• Amino Acids as Polyprotic Acids: Amino acids can act as polyprotic acids due to their multiple ionizable groups.

• Four Regions of a Titration Curve:

  1. Initial

  2. Before Equivalence Point (Buffer Region)

  3. Equivalence Point

  4. After Equivalence Point

Example:

Titrating 25.0 mL of 0.10 M acetic acid with 0.10 M NaOH:

• Calculate pH at various points: before titration, halfway to equivalence (buffer region), at equivalence, and after equivalence.

Equilibria of Slightly Soluble Compounds

Introduction

The solubility of ionic compounds in water is governed by equilibrium principles. The solubility product constant ($K_{sp}$) quantifies the extent to which a compound dissolves.

• $K_{sp}$: The solubility product constant; for $AB \rightleftharpoons A^+ + B^-$, $K_{sp} = [A^+][B^-]$.

• Common Ion Effect (Le Châtelier's Principle): The solubility of a salt decreases in the presence of a common ion.

• Effect of pH on Solubility: Solubility of salts containing basic anions increases in acidic solutions.

• When Will the Solution Precipitate? ($Q$ vs $K_{sp}$): If $Q > K_{sp}$, precipitation occurs; if $Q < K_{sp}$, no precipitation; if $Q = K_{sp}$, the solution is saturated.

Example:

Will a precipitate form when 0.010 M AgNO3 is mixed with 0.010 M NaCl? ($K_{sp}$ for AgCl = $1.8 \times 10^{-10}$)

• $Q = [Ag^+][Cl^-] = (0.010)(0.010) = 1.0 \times 10^{-4}$

• Since $Q > K_{sp}$, AgCl will precipitate.