Aqueous Solutions and Buffers

Aqueous Solutions

  • Weak acids or bases discussed in Chapter 15:

    • Weak acid: HA
    • Weak base: A-

  • Behavior in Water:

    • Weak Acid: HA + H2O \rightleftharpoons A^- + H3O^+
    • Weak Base: A^- + H_2O \rightleftharpoons HA + OH^-

  • Key Terms to Remember:

    • Acid dissociation constant: K_a
    • Le Châtelier's Principle

Buffers (Chapter 16)

Overview of Chapters:

  • Chapter 16 Topics:
    • 16.2 Buffers
    • 16.3 Titrations
    • 16.4 Solubility Equilibria (Ksp)

Concept of Buffers:

  • Definition: Solutions that resist major changes in pH upon addition of acids or bases.
  • Composition: Contains a significant concentration of a weak acid (HA) and its conjugate base (A-).
  • Efficiency: The higher the concentrations of buffer components, the more efficient the buffer is.

Buffer Equations:

  • Weak Acid Equilibrium: HA + H2O \rightleftharpoons A^- + H3O^+

  • Weak Base Equilibrium: A^- + H_2O \rightleftharpoons HA + OH^-

  • Common Ion Effect: The introduction of a common ion shifts the equilibrium and stabilizes the pH.

Buffer Examples:

  • Common buffers include mixtures such as:
    • Acetic acid (CH3COOH) / acetate (CH3COO-)
    • Hydrofluoric acid (HF) / fluoride (F-)
    • Dihydrogen phosphate (H2PO4-) / hydrogen phosphate (HPO42-)
    • Carbonic acid (H2CO3) / bicarbonate (HCO3-)

Troubleshooting Buffers

Example Problem 1: Calculate pH of Buffer

  1. Components: 0.50 M CH3COOH & 0.50 M CH3COONa; K_a (CH3COOH) = 1.8 imes 10^{-5}
  2. Using Equation:
    • Ka = \frac{[H3O^+][CH_3COO^-]}{[CH3COOH]}
    • This allows for calculation of pH.
    • Final pH = 4.74 (assumes equal concentrations of acid and base, leads to pH = pK_a )

Example Problem 2: After Adding NaOH

  • After adding 0.020 mol of NaOH to 1 L of the buffer:
    • Changes concentrations of HA and A-.
    • Calculation Steps:
    1. Write initial concentrations.
    2. Adjust based on stoichiometry.
    3. Calculate new concentrations and then use the equilibrium expression to find new pH.

Acid-Base Titration

  • General Concept: Quantifies reactions between an acid and a base.
  • Key Definitions:
    • Indicator shows endpoint of reaction through color change.
    • Equivalence Point: Reaction completion (moles acid = moles base).

Types of Titration Curves:

  1. Strong Acid vs Strong Base:
    • Equivalence point pH = 7.
  2. Weak Acid vs Strong Base:
    • Equivalence point pH > 7.
  3. Strong Acid vs Weak Base:
    • Equivalence point pH < 7.

Solubility Equilibria (Chapter 16.8)

  1. Concept of Ksp: Equilibrium constant for a slightly soluble ionic compound in water.

    • K_{sp} = [products] (does not include solids).

  2. Common Ion Effect on Solubility: A common ion decreases the solubility of an ionic compound.

    • Example: PbCrO4(s) \rightleftharpoons Pb^{2+} + CrO4^{2-}

  3. Calculating Solubility: Use the Ksp expression to find molar solubility.

    • Example: For Mg(OH)2 , given K{sp} = 6.3 imes 10^{-10} , use:
      • K_{sp} = [Mg^{2+}][OH^{-}]^2 to derive solubility S.

Environmental Impact: Acid Rain

  • Understanding chemical principles helps explain issues like acid rain, formed from combustion of nitrogen and sulfur oxides.
  • Effects:
    • Toxic heavy metals solubilized, impacting marine life (coral ecosystems, etc.).
  • Acid rain similarly affects structures (e.g., marble) and biological organisms.