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Fundamentals of Acids and Bases

Mar 12, 2025

Lecture on Basics of Acids and Bases

Identifying Acids and Bases

  • Acids: Typically have a hydrogen in front (e.g., HCl, HF, HC2H3OH).
  • Bases: Typically have a hydroxide ion (e.g., NaOH, KOH).
  • Hydrogen next to a metal (e.g., NaH) indicates a base.
  • Hydrogen attached to a nonmetal indicates an acid.
  • Acids are positively charged; bases are negatively charged.

Definitions

  1. Arrhenius Definition
    • Acids: Release H+ ions in solution (H+ ions become H3O+ in water).
    • Bases: Release OH- ions in solution.
  2. Bronsted-Lowry Definition
    • Acids: Proton donors.
    • Bases: Proton acceptors.

Examples of Acid-Base Reactions

  • HCl in Water:
    • HCl (acid) donates a proton to H2O (base), forming H3O+ (conjugate acid) and Cl- (conjugate base).
  • Ammonia in Water:
    • NH3 gains a hydrogen to become NH4+ (conjugate acid), and water loses hydrogen to become OH- (conjugate base).

Conjugate Acids and Bases

  • Conjugate Acid: Add H+ and increase charge by 1.
  • Conjugate Base: Remove H+ and decrease charge by 1.

pH Scale

  • Scale usually between 0-14.
  • Neutral: pH = 7.
  • Acidic: pH < 7.
  • Basic: pH > 7.
  • Calculating pH: pH = -log[H3O+].
  • pH + pOH = 14 at 25°C.

Strong vs. Weak Acids and Bases

  • Strong Acids: Ionize completely (e.g., HCl, HBr).
  • Weak Acids: Partially ionize (e.g., HF).
  • Strong Bases: Soluble ionic compounds that ionize completely (e.g., NaOH).
  • Weak Bases: Insoluble compounds or those that ionize less than 1% (e.g., NH3).

Properties

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

Conductivity

  • Strong acids/bases ionize completely, acting as strong electrolytes.
  • Weak acids/bases ionize partially, acting as weak electrolytes.

Reactions with Metals

  • Acids react with active metals to produce hydrogen gas (e.g., Zn + HCl).

Lewis Definitions

  • Lewis Acid: Electron pair acceptor.
  • Lewis Base: Electron pair donor.

KA and KB Expressions

  • KA: Acid dissociation constant.
  • KB: Base dissociation constant.
  • Equations:
    • KA = [H3O+][A-]/[HA]
    • KB = [OH-][HB+]/[B]

Temperature Effects

  • KW: Auto-ionization constant of water (1 x 10^-14 at 25°C).
  • KA x KB = KW: Relationship between acid and base dissociation constants.

Practice Problems

  1. Calculating pH given [H3O+].
  2. Calculating pOH and [OH-] given [H3O+].
  3. Calculating [OH-] given [H3O+].
  4. Determining pKA and KA relationships.
  5. Identifying if a solution is strong or weak based on provided information.

Summary

  • Stronger acids have higher KA and lower pKA.
  • Stronger acids produce weaker conjugate bases and vice versa.
  • Relationships between different definitions and properties of acids and bases.

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