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The Common Ion Effect

The Common Ion Effect. Common Ions Not all solutions contain only one compound Common Ions will shift Eq. The Common Ion Effect. HC 2 H 3 O 2 + H 2 O  H 3 O + + C 2 H 3 O 2 - Suppose we add NaC 2 H 3 O 2 , which way will the reaction shift?. pH and Common Ions.

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The Common Ion Effect

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  1. The Common Ion Effect Common Ions • Not all solutions contain only one compound • Common Ions will shift Eq.

  2. The Common Ion Effect HC2H3O2 + H2O  H3O+ + C2H3O2- Suppose we add NaC2H3O2, which way will the reaction shift?

  3. pH and Common Ions • What will happen to the pH of a soln of benzoic acid, HC7H5O2, if you add potassium benzoate, KC7H5O2? • What will happen to the pH of a soln of NH3 if you add NH4NO3? • What will happen to the pH of a soln of HCl if you add NaCl? • What will happen to the pH of a soln of HCl if you add CH3NH2?

  4. pH and Common Ions • What is the pH of a solution containing 0.30 mol of acetic acid and 0.20 mol of sodium acetate in 1.00 L of solution? (Ans: pH = 4.6) • What is the pH of a soln containing 0.16 M HNO2 (Ka = 4.5 X 10-4) and 0.10 M KNO2? (Ans: 3.1)

  5. pH and Common Ions • What is the fluoride concentration and pH of a soln containing 0.10 mole of HCl and 0.20 mol of HF in 1.0 L of soln? (Ans: 1.4 X 10-3 M, pH = 1.00) • Calculate the CHO2- ion concentration in a solution that is 0.050 M HCHO2 and 0.10 M HNO3. (Ans: 9.0 X 10-5, pH = 1.00)

  6. Buffers 1. A solution that resists a change in pH 2. Buffer capacity – the amount of acid or base needed to overwhelm the buffer. 3. Composed of a weak acid and its conjugate base: Ex: HC2H3O2/NaC2H3O2 HC2H3O2 + H2O  H3O+ + C2H3O2-

  7. Buffers 4. Examples • Blood H2CO3, HCO3- • Sea Water • Pool water • Buffered aspirin

  8. Buffers 1. Adding a base (NaOH) HC2H3O2 + OH- H2O + C2H3O2- Less HC2H3O2 More C2H3O2-

  9. Buffers 2. Adding an acid (HCl) C2H3O2- + H+ HC2H3O2 Less C2H3O2- More HC2H3O2

  10. Writing Titration Reactions Examples: HCl + NaOH  HC2H3O2 + NaOH  H2SO4 + KOH  Ca(OH)2 + HNO3  NH3 + HCl  CH3NH2 + HBr 

  11. Writing Titration Reactions Student Practice HC6H7O6 + NaOH  H2SO4 + NaOH  HClO4 + KOH  C5H5N + HCl  H2SO4 + NH3 C6H5NH2 + HNO3 C2H5NH2 + HBr 

  12. How much more 0.012 M NaOH does it take to neutralize a solution of HCl that is pH=2 compared to one that is pH 4? • Is 0.012 M NaOH an appropriate concentration to use for both pH values? Suggest any changes you might want to make.

  13. Strong Acid/Strong Base Titrations 1. A reaction to determine the molarity of a solution • Uses an indicator or a pH meter to find the eq. point • pH at eq. is always 7 (neutral salt is formed) • Graph if titrating w/ strong base • Graph if titrating w/ strong acid

  14. Strong Acid/Strong Base Titrations 1. How many mL of 0.0978 M NaOH are needed to titrate 15.00 mL of 0.0831 M HBr? (Ans: 12.8 mL) 2. How many mL of 0.0978 M NaOH are needed to completely titrate 13.82 mL of 0.104 M H2SO4? (Ans: 29.4 mL)

  15. Strong Acid/Strong Base Titrations 3. What is the pH of 50.0 mL of 0.100 M HCl when the following volumes of 0.100 M NaOH have been added: a) 0 mL b) 25.0 mL c) 50.0 mL d) 51.0 mL

  16. a) Before addition (No NaOH yet) HCl + H2O  Cl- + H3O+ pH = 1

  17. b) Before Equivalence For HCl Total volume = 75.0 mL M1V1 = M2V2 (0.100 M)(50.0 mL) = x(75.0 mL) x = 0.0666 M

  18. b) Before Equivalence HCl + NaOH  NaCl + H2O I 0.0666 0.0333 C-0.0333 -0.0333 + 0.0333 E 0.0333 0 0.0333 Ans: pH = 1.5

  19. c) At Equivalence HCl + NaOH  NaCl + H2O I 0.05 0.05 C-0.05 -0.05 +0.05 E 0 0 0.05 Ans: pH = 7

  20. d) After Equivalence HCl + NaOH  NaCl + H2O I 0.0495 0.0505 C-0.0495 -0.0495 +0.0495 E 0 0.001 0.0495 Ans: pH = 11

  21. Strong Acid/Strong Base Titrations 4. Calculate the pH when the following volumes of 0.100 M HNO3 have been added to 25.00 mL of 0.100 M KOH. • 0 ml (pH =13) • 24.90 mL (pH = 10.30) • 25.0 mL (pH = 7) • 25.10 mL (pH = 3.70)

  22. Weak Acid/Weak Base Titrations 1. pH at eq. Is NOT 7 if a weak acid or weak base is involved. 2. pH depends on what is “leftover” after the rxn occurs 3. Graph

  23. Weak Acid/Weak Base Titrations • HC2H3O2 is titrated with KOH. Is the pH at equivalence acidic, basic or neutral? • NH3 is titrated with HCl. Is the pH at equivalence acidic, basic or neutral?

  24. Weak Acid/Weak Base Titrations 4. Two step problems a) Stoichiometric step – same as strong acid/strong base b) Equilibrium Step – uses the “leftovers”

  25. Weak Acid/Weak Base Titrations 1. Calculate the pH when the following volumes of 0.0500 M KOH are added to 50.0 mL of 0.0250 M Benzoic acid: a) 20.0 mL b) 25.0 mL c) 30.0 mL

  26. a) Before equivalence Stoichiometric step KOH + HC7H5O2 H2O+ KC7H5O2 I 0.0143 0.0179 C 0.0143 0.0143 0.0143 E 0 0.0036 0.0143

  27. a) Before equivalence Equilibrium Step HC7H5O2 + H2O  H3O+ + C7H5O2- I 0.0036 0 0.0143 C -x x x E 0.0036-x x 0.0143 + x Ka = 6.3 X 10-5 = x(0.0143) pH = 4.80 0.0036

  28. b) At Equivalence Stoichiometric step KOH + HC7H5O2 H2O+ KC7H5O2 I 0.0167 0.0167 C 0.0167 0.0167 0.0167 E 0 0 0.0167

  29. b) At Equivalence Equilibrium Step C7H5O2- + H2O  OH- + HC7H5O2 I 0.0167 C -x x x E 0.0167-x x x Kb = 1.54 X 10-10 = x2/0.0167 pH = 8.2

  30. c) After Equivalence Stoichiometric step KOH + HC7H5O2 H2O+ KC7H5O2 I 0.0188 0.0156 C0.0156 0.0156 0.0156 E 0.0032 0 0.0156 pH = 11.5

  31. Strong Acid/Weak Base Titrations 2. Calculate the pH when the following volumes of 0.0300 M HCl are added to 30.0 mL of 0.0200 M NH3: a) 15.0 mL (pH = 8.8) b) 20.0 mL (pH = 5.6) c) 21.0 mL (pH = 3.2)

  32. a) Before equivalence Stoichiometric Step NH3 + HCl  NH4+ + Cl- I 0.0133 0.0100 C 0.0100 0.0100 0.0100 0.01 E 0.0033 0 0.0100 0.01

  33. a) Before equivalence Equilibrium Step NH3 + H2O  NH4+ + OH- I 0.0033 0.0100 C -x x x E Kb = 1.8 X 10-5 = [x(0.01)]/0.0033 pH = 8.8

  34. b) At Equivalence Stoichiometric Step NH3 + HCl  NH4+ + Cl- I 0.0120 0.0120 C 0.0120 0.0120 0.0120 0.012 E 0 0 0.0120 0.012

  35. b) At Equivalence Equilibrium Step NH4+ + H2O NH3 + H3O+ I 0.0120 C -x x x E Ka= 5.56 X 10-10 = x2/0.0120 pH = 5.6

  36. c) After Equivalence Stoichiometric Step NH3 + HCl  NH4+ + Cl- I 0.0118 0.0124 C 0.0118 0.0118 0.0118 0.0118 E 0 0.0006 0.0118 0.0118 pH = 3.22

  37. Weak Acid/Weak Base Titrations 3. What is the pH at eq. If you titrate 0.100 M HC2H3O2 with 0.0800 M NaOH? (pH = 8.7) 4. What is the pH at eq. if you titrate 0.0934 M CH3NH2 with 0.103 M HCl? (pH = 6.1)

  38. H3PO3 + NaOH  H3PO3 + OH- H2PO3- + H2O H2PO3- + OH-  HPO32- + H2O HPO32- + OH-  PO33- + H2O

  39. Polyprotic Titrations

  40. CH3NH2 + H2O  CH3NH3+ + OH- (add CH3NH3+) b) HCN + H2O  CN- + H3O+ (add HCl) c) HF + H2O  F- + H3O+ (add NaF) d) NH3 + H2O  NH4+ + OH- (add NaOH) e) NH4++ H2O  NH3 + H3O+ (add NH3)

  41. Estimate the pH at equivalence if you titrate HC3H5O2 with NaOH? Estimate the pH at equivalence if you titrate CH3NH2 with HNO3.

  42. a) 3.62 b) 5.01 c) 3.07 • a) 4.1 % b) 1.9% 22. a) 10.33 b) 10.19 24 a) 9.30 26. 24 grams 40. a) 40.7 mL b) 25.3 mL c) 40.2 mL 42. a) 12.10 b) 11.28 c) 7.00 d) 2.73 e) 2.15 • a) 10.87 b) 9.67 c) 9.16 d) 7.7 e) 5.56 f) 3.4 46. a) 7.00 b) 8.25 c) 9.59

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