lecture recording on 12 March 2025 at 09.49.58 AM

Water Autoionization

• Water can react with itself through the process of autoionization.

• This process generates both hydronium ions (H₃O⁺) and hydroxide ions (OH⁻).

• Water acts as an amphoteric substance, meaning it can function as both an acid (donating H⁺) and a base (accepting H⁺).

Strong Acids

• Identify strong acids to determine hydronium and hydroxide concentrations:

  • Common strong acids include Hydrochloric acid (HCl), Hydrobromic acid (HBr), Hydroiodic acid (HI), Nitric acid (HNO₃), Sulfuric acid (H₂SO₄), Perchloric acid (HClO₄), and Chloric acid (HClO₃).

• Strong acids have a very high acid dissociation constant (Ka) that is much greater than 1, indicating almost complete ionization in water.

• The ionization reaction for HCl in water:

  • HCl + H₂O → H₃O⁺ + Cl⁻

• At equilibrium, since we assume complete ionization, the concentration of hydronium ions equals the initial concentration of HCl.

  • Example: For a 0.01 M HCl solution, [H₃O⁺] = 0.01 M.

pH Scale

• The pH scale measures the concentration of hydronium ions in a solution:

  • Calculated as pH = -log[H₃O⁺].

• A sample calculation for 0.01 M HCl:

  • pH = -log(0.01) = 2.00.

• pH is a logarithmic scale; pH < 7 indicates an acidic solution, pH = 7 indicates neutrality, and pH > 7 indicates a basic solution.

• Significant figures in pH: The digits after the decimal point are significant while digits before are not.

Relationships Between pH and Ion Concentrations

• To convert from pH back to [H₃O⁺] when given pH:

  • [H₃O⁺] = 10^(-pH).

  • Example: pH 5.25 → [H₃O⁺] = 10^(-5.25) = 5.6 × 10⁻⁶ M.

• Use the ion-product constant of water (Kw) to relate hydronium and hydroxide concentrations:

  • Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ M² at 25°C.

• Neutral solutions: [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M, pH = 7.

Acids and Hydroxide Concentrations

• In acidic solutions (pH < 7), [H₃O⁺] > [OH⁻].

• In basic solutions (pH > 7), [OH⁻] > [H₃O⁺].

• The relationship of the hydronium ion concentration must always be greater than the hydroxide concentration in acidic solutions and vice versa.

Strong Bases

• Identifying strong bases:

  • Common strong bases include Sodium hydroxide (NaOH), Potassium hydroxide (KOH), and other soluble metal hydroxides.

• Strong bases also completely ionize:

  • Example: KOH → K⁺ + OH⁻.

• For a solution of strong base, the hydroxide concentration equals the original base concentration.

Weak Acids

• Weak acids partially ionize in solution.

• Example: For acetic acid (CH₃COOH) with a Ka of 1.8 × 10⁻⁵,

  • Establish a reaction: CH₃COOH ⇌ H₃O⁺ + CH₃COO⁻.

• Set up an ICE table (Initial, Change, Equilibrium) to solve for concentrations at equilibrium.

• Ka expression: Ka = [H₃O⁺][CH₃COO⁻]/[CH₃COOH].

• If Ka is small and the concentration of the acid is significantly higher (e.g., >100 times Ka), an approximation can simplify calculations without needing a quadratic equation:

  • Assume [CH₃COOH] at equilibrium is approximately equal to its initial concentration.

Conclusion

• Utilizing both the pH scale and the autoionization of water allows for effective calculations of ion concentrations in various solutions.

• Practice various scenarios to increase comfort with these concepts.