Basics of Acids and Bases

Identifying Acids and Bases

• Acids:
  • Typically have a hydrogen (H) in front.
  • Examples: HCl (Hydrochloric acid), HF (Hydrofluoric acid), HC2H3O2 (Acetic acid).
  • Tend to be positively charged.
• Bases:
  • Typically have a hydroxide ion (OH-) attached.
  • Examples: NaOH, KOH.
  • Tend to be negatively charged.
• Special Cases:
  • Hydrogen next to a metal (e.g., sodium hydride) indicates a base.
  • Acids release H+ ions in solution; bases release OH- ions.

Definitions

Arrhenius Definition

• Acids: Release H+ ions in solution.
• Bases: Release OH- ions in solution.

Bronsted-Lowry Definition

• Acids: Proton donors.
• Bases: Proton acceptors.

Acid-Base Reactions

• Example: HCl in water
  • HCl donates a proton to water (base) forming Cl- and H3O+.
• Example: Ammonia in water
  • NH3 accepts a proton from water, forming NH4+ and OH-.

Conjugate Acids and Bases

• Conjugate Acid: Formed by adding H+ to a base.
• Conjugate Base: Formed by removing H+ from an acid.
• Example:
  • Conjugate acid of water: H3O+.
  • Conjugate base of water: OH-.

pH and pOH

• pH Scale: Typically 0-14, but can go beyond.
  • Neutral at 7.
  • Acidic if < 7.
  • Basic if > 7.
• Calculating pH: -log[H3O+].
• Calculating pOH: -log[OH-].
• Relationship: pH + pOH = 14 at 25°C.

Strong vs. Weak Acids and Bases

• Strong Acids: Ionize completely (e.g., HCl, HBr).
• Weak Acids: Ionize partially.
• Strong Bases: Soluble ionic compounds (e.g., NaOH).
• Weak Bases: Partially ionize (e.g., NH3).

Chemical Reactions

• Strong Acid Reaction: Single arrow, complete ionization.
• Weak Acid Reaction: Double arrow, partial ionization.

Additional Concepts

• Electrical Conductivity: Strong acids and bases conduct more due to complete ionization.
• Active Metals: React with acids to produce hydrogen gas.
• Lewis Definitions:
  • Lewis Acid: Electron pair acceptor.
  • Lewis Base: Electron pair donor.

Amphoteric Substances

• Act as both acid and base (e.g., water).

Equilibrium Constants

• Ka (Acid dissociation constant): Strength increases with higher Ka.
• Kb (Base dissociation constant): Calculated similar to Ka.
• Kw: Autoionization constant of water.
  • Kw = 1 x 10^-14 at 25°C.

Practice Problems

• Examples involving calculation of pH, pOH, concentration calculations, and equilibrium expressions.

Key Concepts Review

• Higher Ka = Stronger Acid.
• Higher Kb = Stronger Base.
• Stronger acid = Weaker conjugate base, and vice versa.
• Acid strength inversely related to pKa.

Example Reactions

• HCl dissociating in water.
• Ammonia as a Lewis base reacting with Lewis acids like AlCl3.

Common Misconceptions

• Strong acids and bases are strong electrolytes and conduct electricity well.
• Acids turn blue litmus paper red; bases turn red litmus paper blue.