Acid Base Definitions and Influence of Acid Structure

Acid Base Definitions

  • Weak bases: Bases that do not completely ionize in H₂O.
    • Example: NH<em>3(aq)+H</em>2O(l)⇌NH4+(aq)+OH−(aq)NH<em>3(aq) + H</em>2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)
      • KbK_b is the base dissociation constant.
  • Monoprotic Acids: Acids that have only one acidic proton.
    • Example:
      HNO<em>3(aq)→H+(aq)+NO</em>3−(aq)HNO<em>3(aq) \rightarrow H^+(aq) + NO</em>3^-(aq)
  • Polyprotic Acids: Acids with more than one acidic proton.
    • Example: H<em>2SO</em>4(aq)⇌H+(aq)+HSO<em>4−(aq)H<em>2SO</em>4(aq) \rightleftharpoons H^+(aq) + HSO<em>4^-(aq)HSO</em>4−(aq)⇌H+(aq)+SO42−(aq)HSO</em>4^-(aq) \rightleftharpoons H^+(aq) + SO_4^{2-}(aq)
      • KaK_a is the acid dissociation constant.

Acid-Base Equilibrium and Constants

  • General acid dissociation: HA(aq)+H<em>2O(l)⇌H</em>3O+(aq)+A−(aq)HA(aq) + H<em>2O(l) \rightleftharpoons H</em>3O^+(aq) + A^-(aq)Ka=[H+][A−][HA]K_a = \frac{[H^+][A^-]}{[HA]}
    • pK<em>a=−log⁡[K</em>a]pK<em>a = -\log[K</em>a]: A lower pKapK_a indicates a stronger acid.
    • [H3O+]=[H+][H_3O^+] = [H^+]
  • General base dissociation: B(aq)+H<em>2O(l)⇌BH+(aq)+OH−(aq)B(aq) + H<em>2O(l) \rightleftharpoons BH^+(aq) + OH^-(aq)K</em>b=[BH+][OH−][B]K</em>b = \frac{[BH^+][OH^-]}{[B]}
    • pK<em>b=−log⁡[K</em>b]pK<em>b = -\log[K</em>b]
  • Autoionization of water: H<em>2O(l)+H</em>2O(l)⇌H3O+(aq)+OH−(aq)H<em>2O(l) + H</em>2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)
    • Ion product of water: K<em>w=[H</em>3O+][OH−]K<em>w = [H</em>3O^+][OH^-]
      • At 25°C, Kw=1.0×10−14K_w = 1.0 \times 10^{-14}
  • pH and pOH:
    • pH=−log⁡[H3O+]pH = -\log[H_3O^+]
    • pOH=−log⁡[OH−]pOH = -\log[OH^-]
  • Relationship between K<em>aK<em>a, K</em>bK</em>b, pK<em>apK<em>a, pK</em>bpK</em>b, and KwK_w:
    • pK<em>w=pK</em>a+pKb=14pK<em>w = pK</em>a + pK_b = 14 for an acid/base conjugate pair
    • pH+pOH=pKw=14pH + pOH = pK_w = 14

Influence of Acid Structure on Acidity and Other Aspects

  • Solute Effects on pH
    • Salts can affect the pH of a solution depending on the acidity/basicity of their ions.
    • Example 1: NaCl(s)→Na+(aq)+Cl−(aq)NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)
      • Na+Na^+: Cation from a strong base - No effect on pH.
      • Cl−Cl^-: Anion from a strong acid - No effect on pH.
      • The solution remains neutral.
    • Example 2: NH<em>4Cl(aq)→NH</em>4+(aq)+Cl−(aq)NH<em>4Cl(aq) \rightarrow NH</em>4^+(aq) + Cl^-(aq)
      • Adding NH4ClNH_4Cl will lower the pH.
      • NH<em>4+(aq)⇌NH</em>3(aq)+H+(aq)NH<em>4^+(aq) \rightleftharpoons NH</em>3(aq) + H^+(aq)
      • NH<em>4+NH<em>4^+ is the conjugate acid of a weak base (NH</em>3NH</em>3).
      • Cl−Cl^-: Anion from a strong acid - No effect.
    • Example 3: NaOCl(aq)→Na+(aq)+OCl−(aq)NaOCl(aq) \rightarrow Na^+(aq) + OCl^-(aq)
      • Adding NaOClNaOCl will increase the pH.
      • Na+Na^+: Cation of a strong base - No effect.
      • OCl−(aq)+H2O(l)⇌HOCl(aq)+OH−(aq)OCl^-(aq) + H_2O(l) \rightleftharpoons HOCl(aq) + OH^-(aq)
      • OCl−OCl^- is the conjugate base of a weak acid (HOClHOCl).
  • Oxyacids: Oxygen-containing acids.
    • Acidic hydrogen attached to an oxygen bonded to another element.
    • Acidity increases with the electronegativity of the central element and the number of oxygen atoms.
    • Example: HNO<em>2HNO<em>2 vs. HNO</em>3HNO</em>3
      • HNO<em>2HNO<em>2 (weak acid) vs. HNO</em>3HNO</em>3 (strong acid)
      • Resonance structures contribute to the stability of the conjugate base.
      • More resonance structures lead to greater stability.