Week 5- Acid Base Equilibrium-A108 (1) (2)

Week 5: Acid-Base Equilibrium

Acid-Base Reactions

• Acids and bases can be defined based on their behavior in water.

• Ammonia (NH₃) acts as a base by accepting protons (H⁺) from water, forming ammonium ions (NH₄⁺).

• Water (H₂O) acts as both an acid (H⁺ donor) and a base (H⁺ acceptor).

• The hydroxide ion (OH⁻) signifies a basic solution.

Learning Goals

• Differentiate between acids, bases, and conjugate acid-base pairs using Arrhenius and Bronsted-Lowry theories.

• Understand the distinction between strong and weak acids/bases.

• Write equations for neutralization reactions.

• Compare strengths of weak acids/bases via equilibrium constants.

• Understand the amphoteric nature of water.

• Write expressions for different equilibrium constants.

• Assess whether a solution is acidic, basic, or neutral based on pH.

• Calculate pH and pOH using

  • [H₃O⁺], [OH⁻], and K_w.

Arrhenius Theory

Arrhenius Acids

• An Arrhenius acid increases [H⁺] in water.

• Example: HCl dissolves, dissociating into H⁺ and Cl⁻ ions.

  • HCl → H⁺ + Cl⁻ (forms hydrochloric acid)

• Acids have:

  • Sour taste

  • Electrolytic behavior

  • Neutralizing capability towards bases

  • pH below 7

Naming of Acids

• Acids with H and a nonmetal ending in -ide use prefix hydro- and end in -ic acid.

  • Example: Cl⁻ → HCl (hydrochloric acid)

• Acids with H and a polyatomic ion change:

  • -ate to -ic acid.

  • -ite to -ous acid.

  • Example: NO₃⁻ → HNO₃ (nitric acid), NO₂⁻ → HNO₂ (nitrous acid)

Arrhenius Bases

Arrhenius Bases

• An Arrhenius base increases [OH⁻] in water.

• Example: NaOH dissociates into Na⁺ and OH⁻ ions:

  • NaOH → Na⁺ + OH⁻

• Bases have:

  • Bitter, chalky taste

  • Soapy/slippery feel

  • Electrolytic behavior

  • pH above 7

Neutralization Reactions

• Neutralization is the reaction between an acid and a base resulting in salt and water:

  • Acid + Base → Salt + Water

  • For example: HBr + LiOH → LiBr + H₂O.

    • Steps:

      1. HBr dissociates: H⁺ + Br⁻

      2. LiOH dissociates: Li⁺ + OH⁻

      3. Products: Li⁺ + Br⁻ → LiBr, H⁺ + OH⁻ → H₂O

Bronsted-Lowry Theory

Definitions

• Bronsted-Lowry Acid: Donates H⁺ (proton).

• Bronsted-Lowry Base: Accepts H⁺ (proton).

• In a Bronsted-Lowry reaction, a H⁺ is transferred from an acid to a base.

• Limitations of Arrhenius theory:

  • Doesn't explain reactions of compounds without H⁺ or OH⁻.

Characteristics of Acids and Bases

Table of Properties

• Acids:

  • Produce H⁺

  • Taste sour

  • Neutralize bases

  • Litmus paper turns red

• Bases:

  • Produce OH⁻

  • Taste bitter

  • Neutralize acids

  • Litmus paper turns blue

Conjugate Acid-Base Pairs

• Bronsted-Lowry reactions form conjugate pairs:

  • Conjugate acid: product from proton acceptance.

  • Conjugate base: product from proton donation.

• Example with NH₃ and H₂O:

  • NH₃/NH₄⁺ and H₂O/OH⁻ are the conjugate pairs.

Strength of Acids and Bases

Strong vs Weak

• Strong acids: Completely dissociate in water; only ions remain in solution.

  • Example: Dissociation of HCl, K_a > 1.

• Weak acids: Partially dissociate; both ions and undissociated species present.

  • Example: Carbonic acid's K_a is 4.5 x 10⁻⁷.

General Characteristics

• Weak Acids: Make up most acids, consist of strong conjugate bases.

• Strong Bases: Hydroxides of Group 1A and 2A metals, fully dissociative, example KOH.

Autoionization of Water

• Water can be an acid or a base (amphoteric).

• Autoionization refers to the spontaneous reaction:

  • 2 H₂O ⇌ H₃O⁺ + OH⁻.

  • At 25°C: K_w = [H₃O⁺][OH⁻] = 1 x 10⁻¹⁴.

pH Scale and Calculations

• pH calculated using:

  • pH = -log[H⁺].

• For neutral solutions: [H₃O⁺] = [OH⁻] = 1 x 10⁻⁷.

• Acidic solutions: pH < 7, Basic: pH > 7, Neutral: pH = 7.

Practice Problems

• Determine concentrations, equilibrium constants, and predict reactions based on acid/base properties.

• Verify knowledge with learning checks on pH, identifying strong/weak acids and bases, and recognizing reactions of acid-base pairs.