Basics of Acids and Bases

Identifying Acids and Bases

• Acids: Typically have a hydrogen (H) in front of them.
  • Examples: HCl (hydrochloric acid), HF (hydrofluoric acid), HC₂H₃O₂ (acetic acid).
• Bases: Typically have a hydroxide ion (OH⁻).
  • Examples: NaOH, KOH.
• Identify by Charge:
  • Acids: Positively charged (H⁺).
  • Bases: Negatively charged (OH⁻).

Arrhenius Definition

• Acids: Substances that release H⁺ ions in solution.
• Bases: Substances that release hydroxide ions (OH⁻) in solutions.

Bronsted-Lowry Definition

• Acids: Proton donors.
• Bases: Proton acceptors.
• Example: In HCl + H₂O → Cl⁻ + H₃O⁺, HCl is the acid, and H₂O is the base.

Conjugate Acids and Bases

• Conjugate Acid: Formed by adding a H⁺ to a base.
• Conjugate Base: Formed by removing a H⁺ from an acid.
• Example: NH₃ + H₂O → NH₄⁺ + OH⁻
  • NH₃ is the base, NH₄⁺ is the conjugate acid.
  • H₂O is the acid, OH⁻ is the conjugate base.

pH Scale

• 0 - 14 scale, where 7 is neutral.
• Acidic: pH < 7.
• Basic: pH > 7.
• Calculation:
  • pH = -log [H₃O⁺]
  • pOH = -log [OH⁻]
  • pH + pOH = 14 at 25°C.

Strong vs. Weak Acids

• Strong Acids: Ionize completely in solution (e.g., HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄).
• Weak Acids: Partially ionize in solution (e.g., HF, acetic acid).
• Electrolytes: Strong acids/bases are strong electrolytes; weak acids/bases are weak electrolytes.

Strong vs. Weak Bases

• Strong Bases: Soluble ionic compounds that ionize completely (e.g., NaOH, KOH, Ba(OH)₂).
• Weak Bases: Partially ionize (e.g., NH₃, fluoride, acetate).

Mechanisms of Reactions

• Oxide in Water: Produces hydroxide ions.
• Hydride in Water: Produces hydrogen gas and hydroxide ions.

Properties of Acids and Bases

• Acids: Taste sour, turn blue litmus paper red.
• Bases: Taste bitter, feel slippery, turn red litmus paper blue.

Definitions of Acids and Bases

• Lewis Acids: Electron pair acceptors.
• Lewis Bases: Electron pair donors.

Ka and Kb

• Ka: Acid dissociation constant.
• Kb: Base dissociation constant.
• Relationship: Ka × Kb = Kw = 1 × 10⁻¹⁴ at 25°C.

Practice Problems

1. pH Calculation: Given [H₃O⁺], calculate pH, pOH, [OH⁻].
2. Identify Strongest Acids/Bases: Use Ka values to compare.
3. Calculate pKa and pKb: Given Ka, find pKa, pKb, and Kb.

Key Equations

• pH = -log[H₃O⁺]
• pOH = -log[OH⁻]
• pKa = -log(Ka)
• pKb = -log(Kb)
• pH + pOH = 14
• [H₃O⁺] = 10⁻ᵖᴴ
• [OH⁻] = 10⁻ᵖᴼᴴ
• Ka × Kb = Kw

Summary

• Acid-Base Reactions: Acids donate protons, bases accept protons.
• Conjugate Pairs: Acid-base reactions result in the formation of conjugate acids and bases.
• Acid/Base Strength: Determined by their ability to dissociate in water and their corresponding Ka or Kb values.