Acid & Base Equilibrium (Chapter 16)

Introduction to Acid-Base Equilibria

Acid-base equilibria involve the reversible reactions of acids and bases in aqueous solution. Understanding these equilibria is essential for predicting pH, calculating concentrations, and analyzing chemical behavior in solution.

• Acids are substances that donate protons (H+), while bases accept protons.

• Acid-base reactions can be described using the Brønsted-Lowry and Lewis definitions.

Completing Acid-Base Reactions

• Identify the acid and base in the reaction.

• Write the products: the conjugate base of the acid and the conjugate acid of the base.

• Example:

Labeling Acids, Bases, Conjugate Acids, and Conjugate Bases

• Acid: Donates H+

• Base: Accepts H+

• Conjugate Acid: Formed when a base gains H+

• Conjugate Base: Formed when an acid loses H+

• Example: In , HCl is the acid, Cl- is the conjugate base, H2O is the base, and H3O+ is the conjugate acid.

Calculating pH and pOH for Strong Acids

• Strong acids dissociate completely in water.

• pH calculation:

• pOH calculation:

• Relationship: at 25°C

• Example: 0.01 M HCl:

Calculating [H3O+] and [OH-] from pH

• Example: pH = 3,

Comparing Acid Strength Using Ka Values

• Acid dissociation constant (Ka): Measures acid strength.

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

• Example: Acetic acid () is weaker than HCl ( very large).

Monoprotic vs. Diprotic Acids

• Monoprotic acid: Donates one proton per molecule (e.g., HCl).

• Diprotic acid: Donates two protons per molecule (e.g., H2SO4).

• Each proton donation requires a separate equilibrium calculation (RICE table).

Calculating pH & pOH of Weak Acids and Bases

• Weak acids/bases only partially ionize.

• Set up an equilibrium (RICE) table to solve for [H3O+] or [OH-].

• For weak acids:

• For weak bases:

Determining Equilibrium Concentrations

• Use the RICE table to find equilibrium concentrations of acid, hydronium, and conjugate base.

• Apply the Ka or Kb expression to solve for unknowns.

Acid & Base Properties of Salts

• Salts can produce acidic, basic, or neutral solutions depending on their ions.

• Example: NH4Cl forms an acidic solution; NaCH3COO forms a basic solution.

Lewis Acids and Bases

• Lewis acid: Electron pair acceptor.

• Lewis base: Electron pair donor.

• Example: (acid) + (base)

Buffers (Chapter 17, Sections 17.1–17.4)

Introduction to Buffers

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

Common Ion Effect

• Occurs when a solution contains two substances that share a common ion.

• Suppresses the ionization of a weak acid or base.

• Example: Adding NaCH3COO to acetic acid solution increases [CH3COO-], shifting equilibrium.

Buffer Solutions

• Made from a weak acid and its conjugate base, or a weak base and its conjugate acid.

• Resist pH changes upon addition of small amounts of strong acid or base.

Effect of Adding a Strong Acid to a Buffer

• Strong acid reacts with the buffer's conjugate base, forming more weak acid.

• pH decreases slightly, but less than in an unbuffered solution.

Effect of Adding a Strong Base to a Buffer

• Strong base reacts with the buffer's weak acid, forming more conjugate base.

• pH increases slightly, but less than in an unbuffered solution.

Henderson-Hasselbalch Equation

• Used to calculate the pH of a buffer solution.

• Where: [A-] = concentration of conjugate base, [HA] = concentration of weak acid.

• Example: For a buffer with 0.10 M acetic acid and 0.10 M sodium acetate, .

Additional info: The RICE table (Reaction, Initial, Change, Equilibrium) is a systematic way to solve equilibrium problems for weak acids and bases. The Henderson-Hasselbalch equation is most accurate when the concentrations of acid and conjugate base are not extremely dilute.