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CHEM 163 Chapter 19

CHEM 163 Chapter 19. Spring 2009. Buffers. Solution that resists pH changes Ex. Blood (pH ~ 7.4) Acid must neutralize small amounts of base Base must neutralize small amounts of acid Acid and base must not neutralize each other. Added in as salt (NaCH 3 COO).

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CHEM 163 Chapter 19

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  1. CHEM 163Chapter 19 Spring 2009

  2. Buffers Solution that resists pH changes • Ex. Blood (pH ~ 7.4) • Acid must neutralize small amounts of base • Base must neutralize small amounts of acid • Acid and base must not neutralize each other Added in as salt (NaCH3COO) Use conjugate acid-base pairs! CH3COOH (aq) + H2O (l) CH3COO- (aq) + H3O+ (aq) Common-ion effect Ex: acetate

  3. High concentrations of weak acid/conjugate base • Add H3O+ or OH- • Added amounts are relatively small • Cause only small shifts • React with weak acid or conjugate base HA (aq) + H2O (l) A- (aq) + H3O+ (aq) HA (aq) + OH- (aq) A- (aq) + H2O (l) pH depends on [HA]/[A-] ratio

  4. Making a buffer • Choose the conjugate acid-base pair (pKa ≈ pH) • Calculate the ratio of acid-base concentrations • Determine the buffer concentration • Mix solution; adjust pH Henderson-Hasselbalch equation:

  5. Buffer Properties • Buffer Capacity: • Ability to resist pH change • Unrelated to pH of buffer • Dependent on concentration of weak acid/conj base • Highest when [A-] = [HA] • Buffer Range: • pH range over which buffer is effective • Usually within ±1 pH unit of the pKa of weak acid

  6. Sample Problem Make 200. mL of a pH 3.5 citric acid/sodium citrate buffer with an acid concentration of 0.50 M. We are given solid sodium citrate (294 g/mol) and 5.0 M citric acid. The pKa of citric acid is 3.15.

  7. Measuring pH • Acid-Base Titration Curves: pH v. volume titrant Measuring pH: • pH meter • Acid-base indicators Indicator: • Weak organic acid • HIn different color than In- • Intensely colored (small amount needed) • Changes color over ~ 2 pH units

  8. Titration Curves: Strong acid – Strong base • Low pH (strong acid) • Sudden pH rise (6-8 units) • Slow pH increase  [OH-]added ≈ [H3O+]init Equivalence point: [OH-]added = [H3O+]init pH = 7 End point: when indicator changes color

  9. Calculating pH during titration • Original solution of strong HA • Before the equivalence point • Moles of acid remaining? • Calculate [H3O+] • At the equivalence point: pH = 7 • After the equivalence point • Excess moles of OH- added? • Calculate [OH-] moles base added moles acid total moles acid initial moles acid rxted

  10. Titration Curves:Weak acid – Strong base • Higher initial pH (weak acid, lower Ka) • Buffer region • gradual pH rise • Midpoint: ½ initial HA reacted • Equivalence point: pH > 7.00 • Slow pH increase  [HA] = [A-]  pH = pKa

  11. Calculating pH during titration • Original solution of weak HA • ICE table • Buffer Region • At the equivalence point: • After the equivalence point •  x = [H3O+] • Excess moles of OH- added

  12. Titration Curves:Strong acid – Weak base • Initial pH > 7.00 (weak base) • Buffer region • gradual pH decrease • Equivalence point: pH < 7.00 • Slow pH decrease Less common than strong base-weak acid (fewer appropriate indicators)

  13. Titration Curves:Polyprotic Acids

  14. Salts H2O • soluble NaCl (s) Na+ (aq) + Cl- (aq) • “slightly soluble” • Equilibrium between solid and dissolved ions H2O PbSO4 (s) Pb2+ (aq) + SO42- (aq) Ion-product expression Solubility product Solubility-Product Constant (at saturation) larger Ksp: more dissolution at equil. (saturation) Smaller Ksp: less dissolution at equil. (saturation)

  15. Insoluble Metal Sulfides H2O MnS (s) Mn2+ (aq) + S2- (aq) S2- (aq) + H2O (l) HS- (aq) + OH- (aq) MnS (s) + H2O (l) Mn2+ (aq) + HS- (aq) + OH- (aq)

  16. 3-minute Practice Write Ksp expression for each of the following: Silver bromide in H2O H2O AgBr (s) Ag+ (aq) + Br - (aq) Silver sulfide in H2O Ag2S (s) + H2O (l) 2Ag+ (aq) + HS- (aq) + OH- (aq)

  17. 2-minute Practice Higher Ksp = greater solubility? Yes, for compounds with same total number of ions

  18. What else affects solubility? • Presence of a common ion: Decreases solubility H2O PbSO4 (s) Pb2+ (aq) + SO42- (aq) Add Na2SO4? • pH: ↑ [H3O+] ↑ solubility if compound contains anion of weak acid H2O CaCO3 (s) Ca2+ (aq) + CO32- (aq) CO32- (aq) + H3O+ (aq) H2O (l) + HCO3- (aq)

  19. Homework problems Chap 19: #9, 13, 19, 29, 50, 63, 70, 76, 78, 90 Due Tuesday, 4/28 More lecture notes will be added next week! Stay tuned.

  20. Precipitation Will it occur? • Qsp = Ksp: • Qsp > Ksp: • Qsp < Ksp: • Selective precipitation • Way to separate ions • Form slightly soluble compounds with different Ksp Saturated solution Precipitation occurs Unsaturated solution

  21. Selective Precipitation Mix 0.2 M Zn(NO3)2 and 0.4 M Mn(NO3)2. Precipitate? Add NaOH… Products? Zn(OH)2 and Mn(OH)2 3-minute Practice Ksp Zn(OH)2 = 3.0 x 10-16 KspMn(OH)2 = 1.6 x 10-13 Which product is more soluble? What [OH-] would need to make a saturated solution of the more soluble product? Hint: use Ksp expression!

  22. Complex Ions • Central metal ion + ligands Ionic ligands: Lewis base Lewis acid OH-, CN-, halides Molecular ligands: H2O, NH3 M(H2O)42+ (aq) + 4 NH3(aq) M(NH3)42+ (aq) + 4 H2O(l) Formation constant:

  23. Effects of ligands A slightly soluble compound becomes more soluble when its cation forms a complex ion. AgBr(s) Ag+ (aq) + Br- (aq) Add Na2S2O3: Ag+ (aq) + S2O32- (aq) 2 Ag(S2O3)23-(aq) Amphoteric Hydroxides: • Very slightly soluble in water • More soluble in acidic or basic solutions Al(OH)3 (s) + 3H3O+ Al3+(aq) + 6 H2O (l) Al(H2O)6 (s) + 4 OH- Al(H2O)2(OH)4- (aq) + 4 H2O (l)

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