Exam 3 Study Guide - Acid Base Equilibria

Chapter 16 - Acid Base Equilibria

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Arrhenius Definition (1880s)

  • Acid: Substance that produces H⁺ ions in water (e.g., HCl).

  • Base: Substance that produces OH⁻ ions in water (e.g., NaOH).

  • Limitations:

    • Restricted to aqueous solutions.

    • Not all acids and bases conform to H⁺ and OH⁻ definitions.

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Bronsted-Lowry Definition (1920s)

  • Acid: Proton (H⁺) donor.

  • Base: Proton (H⁺) acceptor (requires a lone pair).

  • Limitations:

    • Not all acids have a "removable" H⁺.

Strong and Weak Acids/Bases

  • Strong Acids: Completely dissociate in water.

  • Weak Acids: Only partially dissociate in water.

Autoionization of Water

  • Reaction:
    H₂O (l) + H₂O (l) ⇌ H₃O⁺ (aq) + OH⁻ (aq).

  • Equilibrium Expression:

Importance of Kw for Water

  • Kw: Equilibrium constant for water,

    Kw = 1.0 x 10⁻¹⁴ at 25°C.

  • Relationship Between [H⁺] and [OH⁻]:

    • As [H⁺] increases, [OH⁻] decreases, maintaining the constant (Kw).

    • In acidic solutions: [H⁺] > [OH⁻].

    • In basic solutions: [OH⁻] > [H⁺].

pH Scale

  • Definition:
    pH = -log [H⁺] or pH = -log [H₃O⁺].

  • Concentration Relationship:

  • For neutral solution:
    [H⁺] = [OH⁻] = 1.0 x 10⁻⁷ M.

  • For acidic solution:
    [H⁺] > 1.0 x 10⁻⁷ M, [OH⁻] < 1.0 x 10⁻⁷ M.

  • For basic solution:
    [H⁺] < 1.0 x 10⁻⁷ M, [OH⁻] > 1.0 x 10⁻⁷ M.

Identifying Acids and Bases

  • Acids: Typically have H at the beginning of their formula.

    • Examples: HCl, H₂SO₄.

  • Bases: Typically contain OH⁻ in their formula.

    • Examples: NaOH, Ba(OH)₂.

Naming Acids

  • With nonmetal ion: Use "hydro" prefix and end with "ic acid".

    • Example: HCl ➔ Hydrochloric acid.

  • With polyatomic ion: Change the end of the polyatomic ion.

    • Example: H₂SO₄ ➔ Sulfuric acid.

Naming Bases

  • Arrhenius bases are named as hydroxides.

    • Example: NaOH ➔ Sodium hydroxide.

Conjugate Acid-Base Pairs

  • Reversible reactions: ⇌

  • Example:

    • HF ⇌ H⁺ + F⁻

    • H₂O ⇌ OH⁻ + H⁺.

Strengths of Acids & Bases

  • Weak Acids: Partially dissociate in water.

    • Example: Acetic acid.

  • Strong Acids: Fully ionize in solution (100% dissociation).

    • Provide large concentrations of H₃O⁺.

  • Weak Bases: Produce few ions in solution.

  • Strong Bases: Formed from Group 1A(1) & 2A(2) metals, completely dissociate in solution.

Calculating Percent Ionization

  • Formula:
    % Ionization = (Ionized Concentration / Initial Concentration) x 100.

Buffers

  • Definition: A weak acid in equilibrium with its conjugate base or weak base with its conjugate acid.

  • Function:

    • Resist changes in pH when acids or bases are added.

    • Example: HNO₂ ↔ H⁺ + NO₂⁻.

    • Addition of acid or base shifts the equilibrium to maintain pH.

Henderson-Hasselbalch Equation

  • Equation:
    pH = pKa + log([A⁻]/[HA]).

  • Where:

    • HA = weak acid.

    • A⁻ = conjugate base.

Acid-Base Titrations

  • Overview: Analyze unknown concentration of acid/base by measuring the equivalent of a known acid/base.

  • Types:

    • Strong Acid - Strong Base

    • Strong Acid - Weak Base

    • Weak Acid - Strong Base

    • Weak Acid - Weak Base

    • Polyprotic Acid - Strong Base.

Example of Acid-Base Titration Procedure

  1. Write the net ionic equation (e.g., H⁺ + OH⁻ → H₂O).

  2. Calculate moles of HCl and find volume of NaOH required at equivalence point.

  3. Determine pH at equivalence point (often neutral salt formation).

Solubility Equilibria

  • Definition: The equilibrium established when a solid solute is in contact with its saturated solution.

  • Example Reaction: AgCl(s) ⇌ Ag⁺ (aq) + Cl⁻ (aq).

  • Solubility Product Constant (Ksp):
    Ksp = [Ag⁺][Cl⁻]

Calculating Solubility and Ksp

  • Example:

    • For CaSO₄:

    Ksp = [Ca²⁺][SO₄²⁻]

  • Finding Concentrations: Use Ksp and solubility relationships to determine x (concentration of ions).