Acids and Bases Bootcamp Notes

Acid-Base Theories

  • Arrhenius Theory: Limited to aqueous solutions.
    • Acid: Proton (H+H^+) donor (e.g., HClHCl, HNO3HNO_3, H2SO4H_2SO_4).
    • Base: Hydroxide ion (OH−OH^-) donor (e.g., NaOHNaOH, KOHKOH, Fe(OH)3Fe(OH)_3).
  • Brønsted-Lowry Theory:
    • Acid: Proton (H+H^+) donor.
    • Base: Proton (H+H^+) acceptor (includes NH3NH_3).
  • Lewis Theory:
    • Acid: Electron pair acceptor (e.g., BF3BF_3 with empty p-orbital).
    • Base: Electron pair donor (e.g., NH3NH_3 with lone pair).

Strength and Dissociation Constants

  • Strong Acids and Bases: Completely dissociate in solution; their conjugates are inert ions.
    • Common Strong Acids: HClHCl, HBrHBr, HIHI, H2SO4H_2SO_4, HNO3HNO_3, HClO3HClO_3, HClO4HClO_4.
    • Common Strong Bases: NaOHNaOH, KOHKOH, and hydroxides of Group I and II metals.
  • Weak Acids and Bases: Exist in equilibrium with their conjugate pairs.
  • Autoionization of Water:
    • H2O(l)⇌H+(aq)+OH−(aq)H_2O (l) ⇌ H^+ (aq) + OH^- (aq)
    • Kw=[H+][OH−]=1×10−14K_w = [H^+][OH^-] = 1 \times 10^{-14} at 25ºC.
    • Ka×Kb=Kw=10−14K_a \times K_b = K_w = 10^{-14}.
    • pKa+pKb=pKw=14pK_a + pK_b = pK_w = 14.
  • Dissociation Constant (KK) Scale:
    • K>1K > 1: Strong.
    • 10−14<K<110^{-14} < K < 1: Weak.
    • K<10−14K < 10^{-14}: Inert.
  • Amphoteric Species: Can act as either an acid or a base (e.g., H2OH_2O, HCO3−HCO_3^-).

Quantitative Calculations

  • pH and pOH Definitions:
    • pH=−ln⁡([H+])pH = -\ln([H^+])
    • pOH=−ln⁡([OH−])pOH = -\ln([OH^-])
    • pH+pOH=14pH + pOH = 14
  • Equivalents and Normality:
    • Equivalents: Moles of protons donated or accepted.
    • Normality (NN): N=M⨉equivalents per moleN = M ⨉ \text{equivalents per mole}.
    • Gram-Equivalent Weight (GEWGEW): g eq−1=g mol−1eq mol−1g \, eq^{-1} = \frac{g \, mol^{-1}}{eq \, mol^{-1}}.

Titrations and Buffers

  • Titration Formula: VANA=VBNBV_A N_A = V_B N_B.
  • Equivalence Point: Occurs when all acid/base in the original solution is neutralized.
    • Strong Acid-Strong Base: Equivalence point at pH=7pH = 7.
    • Weak Acid-Strong Base: Equivalence point at pH>7pH > 7 due to conjugate base reaction.
  • Half-Equivalance Point: Enough titrant added to neutralize half the original solute; [HA]=[A−][HA] = [A^-].
  • Buffers: Resist changes in pH; range is typically ±1\pm 1 pH unit around pKapK_a.
  • Henderson-Hasselbalch Equation: pH=pKa+log⁡([conjugate base][conjugate acid])pH = pK_a + \log \left( \frac{[\text{conjugate base}]}{[\text{conjugate acid}]} \right).

Questions & Discussion

  • Question: What does a reaction involving a Brønsted-Lowry acid and base produce?
    • Answer: Salt plus water.
  • Question: Which of the following is a conjugate acid-base pair?
    • Answer: NH4+/NH3NH_4^+/NH_3.
  • Question: All of the following are strong bases EXCEPT: NaOHNaOH, KOHKOH, Ca(OH)2Ca(OH)_2, Al(OH)3Al(OH)_3?
    • Answer: Al(OH)3Al(OH)_3.
  • Question: What is the ionization constant of H2SH_2S?
    • Answer: Much less than 1 (it is a weak acid).
  • Question: Based on ammonium ion pKapK_a values (NH3:9.26NH_3: 9.26, CH3CH2NH2:10.64CH_3CH_2NH_2: 10.64, (CH3CH2)2NH:10.98(CH_3CH_2)_2NH: 10.98, (CH3CH2)3N:10.76(CH_3CH_2)_3N: 10.76), which is the strongest base?
    • Answer: (CH3CH2)2NH(CH_3CH_2)_2NH (highest conjugate acid pKapK_a corresponds to lowest base pKb≈3.02pK_b \approx 3.02).
  • Question: What is the gram-equivalent weight of Arsenic acid (H3AsO4H_3AsO_4) with molecular mass 141.9 g mol−1141.9 \, g \, mol^{-1}?
    • Answer: 47.3 g eq−147.3 \, g \, eq^{-1} (141.9/3141.9 / 3 equivalents).
  • Question: How many equivalence points exist in a titration of phosphoric acid (H3PO4H_3PO_4) with sodium hydroxide?
    • Answer: Three (one for each proton).
  • Question: Which requires more base to neutralize: 49 g49 \, g phosphoric acid (MW 98 g mol−198 \, g \, mol^{-1}) or 36 g36 \, g hydrochloric acid (MW 36 g mol−136 \, g \, mol^{-1}), both in 100 cm3100 \, cm^3 water?
    • Answer: Phosphoric acid because it has greater normality (15 N15 \, N vs. 10 N10 \, N).