Recording-2025-03-04T16:52:21.792Z

Module 9: Acids and Bases

Overview

• Focus on acids and bases as specialized solutions.

• Builds on prior knowledge of solutions, solutes, solvents, etc.

• Definitions of acids and bases will be provided, emphasizing multiple perspectives.

Definitions

Arrhenius Definition

• Acid: A compound that releases hydrogen ions (H⁺) in solution.

  • Example: HCl (hydrochloric acid) dissociates to H⁺ and Cl⁻ in water.

• Base: A compound that releases hydroxide ions (OH⁻) in solution.

  • Example: NaOH (sodium hydroxide) dissociates in water to release OH⁻.

Properties of Acids and Bases

• Acids tend to:

  • Be sour in taste.

  • React with metals to release hydrogen gas.

  • Cause color changes in indicators (e.g., litmus paper).

• Bases tend to:

  • Have a slippery feel.

  • Taste bitter.

  • Also cause color changes in indicators.

Significance of Hydrogen Ions and Hydroxide Ions

• Hydrogen ions are often referred to as protons (due to hydrogen's atomic structure: atomic number 1).

• Reactions between acids and bases can result in neutralization, producing water (H₂O).

Bronsted-Lowry Definition

• Bronsted-Lowry Acid: Any compound that donates protons (H⁺) in solution.

• Bronsted-Lowry Base: Any compound that accepts protons (H⁺) in solution.

• This definition expands the types of compounds classified as bases.

• Example Reaction: In a reaction with formic acid (HNO₂) and ammonia (NH₃), NH₃ acts as a base by accepting H⁺ to become NH₄⁺ (ammonium).

Conjugate Acids and Bases

• Conjugate Acid: The acid formed after a base gains a proton.

• Conjugate Base: The base formed after an acid donates a proton.

• Example: In the reaction of HNO₂ and water, if water accepts a proton, it becomes hydronium (H₃O⁺), making it a conjugate acid.

Equilibrium and Le Chatelier's Principle

• Chemical reactions can reach equilibrium, where the forward and reverse reactions occur at the same rate.

• Le Chatelier's Principle: If an external factor changes the concentration of a substance in equilibrium, the reaction will shift in a direction to counter that change.

Water Dissociation and pH Scale

• Water can dissociate into hydronium (H₃O⁺) and hydroxide (OH⁻) ions:

  • 2 H₂O ⇌ H₃O⁺ + OH⁻

• Concentration of hydronium in pure water is always equal to that of hydroxide: [H₃O⁺] = [OH⁻] = 1 x 10⁻⁷ M

• pH is defined as:

  • pH = -log[H₃O⁺]

• In pure water, pH = 7. For any solution:

  • pH < 7 indicates acidic solution.

  • pH > 7 indicates basic solution.

Calculating pH

• To calculate pH from hydronium concentration:

  1. Use the formula: pH = -log[H₃O⁺].

  2. If given hydroxide concentration (OH⁻), use the water dissociation constant (Kw = 1 x 10⁻¹⁴) to find [H₃O⁺]:

     • [H₃O⁺] = Kw / [OH⁻]

  3. Calculate pH as above after determining [H₃O⁺].

Example Questions

• If given [H₃O⁺] = 3.4 x 10⁻⁴ M, calculate: pH = -log(3.4 x 10⁻⁴) = 3.47.

• If given [OH⁻] = 6.7 x 10⁻⁵ M:

  1. Find [H₃O⁺]: [H₃O⁺] = 1 x 10⁻¹⁴ / [OH⁻] = 1.49 x 10⁻¹⁰ M.

  2. pH = -log(1.49 x 10⁻¹⁰) ≈ 9.83.

Summary

• Definitions of acids and bases are crucial for understanding their behavior in reactions.

• The discussion of conjugate pairs and the equilibrium states helps to grasp complex acid-base interactions.

• The pH scale is fundamentally tied to water's dissociation and the balance of hydrogen and hydroxide ions.