4.14.25 -- Acids/Bases Lecture Notes

Strong Acids and Bases

• Strong Acids:

  • Hydrochloric Acid (HCl)
  • Hydrobromic Acid (HBr)
  • Hydroiodic Acid (HI)
  • Perchloric Acid (HClO₄)
  • Nitric Acid (HNO₃)
  • Sulfuric Acid (H₂SO₄)

• Strong Bases:

  • Potassium Hydroxide (KOH)
  • Sodium Hydroxide (NaOH)
  • Magnesium Hydroxide (Mg(OH)₂)
  • Calcium Hydroxide (Ca(OH)₂)
  • Barium Hydroxide (Ba(OH)₂)

• Note: Familiarize yourself with examples of strong acids and bases for college chemistry. Some lists might differ in length or content, but knowing a foundational group is most important.

Auto-Ionization of Water

• Definition: Auto-ionization is the process where water generates ions by itself.

  • Example: 2 H₂O ⇌ H₃O⁺ + OH⁻

• Types of Solutions:

  • Neutral Solutions: [H₂O] = [H₃O⁺] = [OH⁻]
  • Acidic Solutions:
  • More H₃O⁺ than OH⁻
  • Example: Acetic Acid (CH₃COOH)
  • Basic Solutions:
  • More OH⁻ than H₃O⁺

Naming Acids

• Binary Acids (hydrogen + one other element):
  • Named with the prefix "hydro-" and root of the anion with the suffix "-ic".
  • Example: HCl (Hydrochloric Acid)
• Oxyacids (contain oxygen):
  • Named without "hydro"; the suffix depends on the anion:
  • Anion suffix "-ate" → acid suffix "-ic" (H₂SO₄: Sulfuric Acid)
  • Anion suffix "-ite" → acid suffix "-ous" (H₂SO₃: Sulfurous Acid)

Dissociation of Acids

• Dissociation in water separates acids into ions:
  • For example: HCl → H⁺ + Cl⁻
• Acetic Acid: Mostly remains undissociated in water, a characteristic of weak acids.
• Strong Acids: Completely dissociate, while weak acids do not.

Understanding pH Scale

• pH Definition: Measure of acidity/basicity based on H⁺ concentration.
  • 0-7: Acidic solutions
  • 7: Neutral (pure water)
  • 7-14: Basic solutions
• Logarithmic Nature: Every unit change in pH represents a tenfold change in H⁺ concentration.
  • Example: pH 6 is ten times more acidic than pH 7.

Definitions of Acids and Bases

1. Arrhenius Definition:

   • Acids produce H⁺ in aqueous solutions;
   • Bases produce OH⁻ in aqueous solutions.

2. Bronsted-Lowry Definition:

   • Acids donate protons (H⁺);
   • Bases accept protons.

3. Lewis Definition:

   • Acids accept electron pairs;
   • Bases donate electron pairs.

Conjugate Acid-Base Pairs

• Conjugate Bases are what acids become after donating a proton.
• Conjugate Acids are what bases become after accepting a proton.
  • Example: H₂SO₄ (strong acid) → HSO₄⁻ (conjugate base)
  • Example: NH₃ (base) + H₂O → NH₄⁺ (conjugate acid)

Neutralization Reactions

• Reaction between an acid and a base yielding water and a salt:
  • Example: HNO₃ (Nitric Acid) + NaOH (Sodium Hydroxide) → H₂O + NaNO₃ (Sodium Nitrate)
• Net Ionic Equation:
  • Strong acids and bases are fully dissociated in water, while weak acids remain as is.

Key Points for Exam Preparation

• Familiarize with strong acids and bases.
• Understand the naming conventions for acids.
• Practice molecular and net ionic equations for neutralization reactions.
• Be clear on definitions of acids and bases in different contexts (Arrhenius, Bronsted-Lowry, Lewis).