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Acid-Base Chemistry

Acid-Base Chemistry. Review Unit 6 Chapters 4, 15, 16. Properties of Acids. Acids Taste sour Turn litmus red, phth colorless React with Metals to make H 2 and salt (redox) Carbonates to make CO 2 and H 2 O and salt Base to make H 2 O and salt (neutralization, K=10 14 ).

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Acid-Base Chemistry

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  1. Acid-Base Chemistry Review Unit 6 Chapters 4, 15, 16

  2. Properties of Acids • Acids • Taste sour • Turn litmus red, phth colorless • React with • Metals to make H2 and salt (redox) • Carbonates to make CO2 and H2O and salt • Base to make H2O and salt (neutralization, K=1014)

  3. Properties of Bases • Taste bitter • Feel slippery • Turn litmus blue, phth pink • React with • Fats, oils, and waxes to make soap and water • Acids to make salt and H2O

  4. Models of Acid-Base Behavior • Arrhenius • Acids dissolve in water to form H+/H3O+ ions • Bases dissolve in water to form OH- ions • Built on the conjugate species of water • Brönsted-Lowry • Acids are proton (H+) donors • Bases are proton acceptors • Lewis • Acids are electron-pair acceptors • Bases are electron-pair donors • Model of choice for coordination chemists

  5. Strong vs. Weak Acids • Acids ionize HX + H2O <---> X- + H3O+ • Ka=[H3O+] [X-]/[HX] • Ka= hydronium x conjugate/original • Strong acids ionize completely, irreversibly, no K • HCl, HBr, HI, HNO3, H2SO4, HClO4 • Weak acids ionize < 5%, reversibly, use K • Reversibility means the species produced (X-) can reaccept a proton, making it a conjugate base

  6. Strong vs. Weak Bases • Bases hydrolyze (split water) • B + H2O <---> HB+ + OH- • Kb = [OH-] [HB+]/[B] • Kb = hydroxide x conjugate/original • Strong bases are the soluble hydroxides • Group I and heavy Group II hydroxides • Weak bases hydrolyze < 5% • Reversibility means the species produced (HB+) can redonate a proton

  7. pH • Def’n: pH = -log [H3O+] • [H3O+] = 10-pH • [H3O+] [OH-] = 1.0 x 10-14 • H2O(l) + H2O(l) <===> H3O+(aq) + OH-(aq) • Kw = [H3O+] [OH-] = 1.0 x 10-14 • pH + pOH = 14

  8. Key Equilibria • Autoionization of water • Occurs in any aqueous solution; Kw applies in all • 2 H2O <===> H3O+ + OH- Kw = 1.0 x 10-14 • Ionization of an acid • HX(aq) + H2O(l) <===> X-(aq) + H3O+(aq) • Ka = [H3O+] [X-]/[HX] • Hydrolysis of a base • B(aq) + H2O(l) <===> HB+(aq) + OH-(l) • Kb = [OH-] [HB+]/[B]

  9. Uses for Ka and Kb (simple sol’ns) • Measure pH ---> Find Ka or Kb • Given Ka ---> Find pH • Given Kb ---> Find pOH • All can be solved with The Grid, but • Grid always gives • Ka = x2/([original] - x) where x = [H3O+] or • Kb = x2/([original] - x) where x = [OH-]

  10. Buffers • Sol’ns that are resistant to changes in pH • Roughly equimolar mix of acid and conjugate • Conjugates don’t neutralize each other • Conjugates can neutralize all other acids and bases • Types of buffer problems • Design a buffer with a certain pH • Find pH of a buffer with given composition • Find pH of a buffer after something is added • All are solved with either Ka or Kb

  11. Buffer Form of the Ka Equation • Ka = [H3O+] [conj base]/[acid] applies to buffers just as for simple solutions • Differences • Due to common ion effect, there is no dissociation • So [acid]original = [acid]equilibrium (same for base) • We rearrange Ka/b equation in “buffer form” • [H3O+] = Ka (acid/base) [OH-] = Kb (base/acid) • Can use original [ ]’s = equilibrium [ ]’s • Can use M’s or moles in “the ratio”

  12. Making a Buffer • Mix HF(aq) with NaF(s)or • Partially neutralize HF • Neutralizing HF produces F- • HF + OH- ---> H2O + F- • An equimolar buffer has a “ratio” of 1, so its pH = pKa of the weak acid chosen • Fine-tune pH by tweaking the “ratio”

  13. pH Curves • Measure pH as base is titrated into an acid • Equivalence point is when moles H3O+ = moles OH- • Big change in pH at the equivalence point • Biggest for strong acid/strong base titration • Smallest for weak acid/weak base titration • Choose indicator with pKa ~ pH at equiv point • pH at equivalence point • = 7 for strong acid/strong base (no conjugates left) • > 7 for weak acid/strong base (conj base left) • < 7 for strong acid/weak base (conj acid left) • Use equivalence point to calculate [ ] of UK • Use half-equivalence point to find pKa

  14. Common Weak Species • Acids • Organic acid (RCOOH), HCO3-, Zn+2, NH4+ • Bases • NH3, amines (RNH2), conjugates

  15. Polyprotic Acids • First ionization is more extensive than second • Ka1 > Ka2 • pH curve has two eq. points

  16. Miscellaneous • Anhydrides • Metal oxides are base anhydrides • CaO + H2O ---> Ca+2 + 2 OH- • Nonmetal oxides are acid anhydrides • SO2 + H2O ---> H2SO3 • Amphiprotic • HCO3- can either gain or lose a proton • Amphoteric • Al species can act as acids or bases (in a Lewis way)

  17. Types of Problems • Simple solution • Only one ingredient in water • Ka = x2/([original] - x) where x = [H3O+] • Buffer • Acid and it’s conjugate • [H3O+] = Ka (acid/base) or use Kb • Titration • Acid and a non-conjugate base • Two steps • Neutralization goes to completion • What remains is • Simple solution if you go to the equivalence point • Buffer is you only partially neutralize

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