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).