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.