1. Sections 17.1, 17.2, 17.4, 17.5(Common Ion Effect)
Chapter 17
Sections
17.1, 17.2, 17.4,
17.5(Common Ion Effect)

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The Common-Ion Effect
“The extent of ionization of a weak electrolyte is decreased by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte.”
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3. AgCl(s)  Ag+(aq) + Cl(aq)
Common Ion Effect
The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction.
AgCl(s)  Ag+(aq) + Cl(aq)
See problems – 17.18

4. Problem Calculate the % Ionization and pH of
a solution that is M in formic acid (HCHO2) Ka=1.8x10-4
a solution that is M in sodium formate (NaCHO2) and M in formic acid (HCHO2)
a solution that is made by combining 125 mL of M hydrofluoric acid (Ka=6.8x10-4)with 50.0 mL of 0.10 M sodium fluoride.

5. Problem
Calculate the percent ionization of M lactic acid (HC3H5O3) (Ka = 1.4 × 10−4).
Calculate the percent ionization of M lactic acid in a solution containing M sodium lactate.

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Buffers
Buffers are solutions of a weak conjugate acid–base pair.
They are particularly resistant to pH changes, even when strong acid or base is added.
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7. A Buffered Solution
. . . resists change in its pH when either H+ or OH are added.
1.0 L of 0.50 M H3CCOOH
M H3CCOONa
pH = 4.74
Adding mol solid NaOH raises the pH of the solution to 4.76, a very minor change.

9. Key Points on Buffered Solutions
1. They are weak acids or bases containing a common ion.
2. After addition of strong acid or base, deal with stoichiometry first, then equilibrium.

10. Henderson-Hasselbalch Equation
Useful for calculating pH when the [A]/[HA] ratios are known.

11. Problem
Carbonic Acid (H2CO3 )
Ka Values
Ka1
4.5 x 10-7 
Ka2
4.7 x 10-11
Calculate the pH of a buffer that is M in NaHCO3 and M in Na2CO3
Calculate problem both ways – equilibrium method and using Henderson Hasselbalch equation

12. Buffered Solution Characteristics
Buffers contain relatively large amounts of weak acid and corresponding base.
Added H+ reacts to completion with the weak base.
Added OH reacts to completion with the weak acid.
The pH is determined by the ratio of the concentrations of the weak acid and weak base.

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pH Range
The pH range is the range of pH values over which a buffer system works effectively.
It is best to choose an acid with a pKa close to the desired pH.
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14. Making a Buffer Solution Problem
How many moles of NaBrO should be added to 1.00 L of 5.0 x 10-2 M hypobromous acid (HBrO) to form a buffer solution of pH = 9.14? Assume that no volume change occurs when the NaBrO is added.
See problems – 17.26

15. Buffering Capacity
. . . represents the amount of H+ or OH the buffer can absorb without a significant change in pH.

16. Buffering Action
See problems – 17.28

17. Buffering Action
See problems – 17.30

18. Solubility Product For solids dissolving to form aqueous solutions.
Bi2S3(s)  2Bi3+(aq) + 3S2(aq)
Ksp = solubility product constant
and
Ksp = [Bi3+]2[S2]3
See problem 17.49

19. Solubility Product
“Solubility” = x = concentration of Bi2S3 that dissolves, which equals 1/2[Bi3+] and 1/3[S2].
Note: Ksp is constant (at a given temperature)
x is variable (especially with a common
ion present)
See problem 17.50

20. Calculating Ksp
See problems – 17.55

21. Factors Affecting Solubility
The Common-Ion Effect
If one of the ions in a solution equilibrium is already dissolved in the solution, the equilibrium will shift to the left and the solubility of the salt will decrease:
BaSO4(s)
Ba2+(aq) + SO42(aq)
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22. Factors Affecting SolubilitypH
If a substance has a basic anion, it will be more soluble in an acidic solution.
Substances with acidic cations are more soluble in basic solutions.
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23. Solubility of CaF2 vs. NaF
See problems – 17.58

24. Solubility of CaF2 vs. pH
See problems – 17.61

25. Titration (pH) Curve
A plot of pH of the solution being analyzed as a function of the amount of titrant added.
Equivalence (stoichiometric) point: Enough titrant has been added to react exactly with the solution being analyzed.

27. Strong Acid – Strong Base Titration

28. Strong Base – Strong Acid Titration

29. Weak vs. Strong Acid Titration

31. Weak Acid - Strong Base Titration
Step 1 - A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species.
Step 2 - An equilibrium problem - determine position of weak acid equilibrium and calculate pH.

32. Titration Curve pH Calculation

33. Titration

34. Acid-Base Indicator
. . . marks the end point of a titration by changing color.
The equivalence point is not necessarily the same as the end point.

35. Phenolphthalein Indicator

36. Other Indicators

40. END

41. Problem
If the molar solubility of CaF2 at 35°C is × 10−3 mol/L, what is Ksp at this temperature?
The Ksp of Ba(IO3)2 at 25°C is 6.0 × 10−10. What is the molar solubility of Ba(IO3)2?

42. Problem – 17.52 pure water 0.010 M KF solution 0.050 M LaCl3 solution.
Calculate the solubility of LaF3 (Ksp = 2 x 10-19) in grams per liter in
pure water
0.010 M KF solution
0.050 M LaCl3 solution.

43. Problem – 17.56
Calculate the molar solubility of Fe(OH) (Ksp = 7.9 x 10-16) when buffered at pH
8.0
10.0
12.0.

44. Equilibria Involving Complex Ions
Complex Ion: A charged species consisting of a metal ion surrounded by ligands (Lewis bases).
Coordination Number: Number of ligands attached to a metal ion. (Most common are 6 and 4.)
Formation (Stability) Constants: The equilibrium constants characterizing the stepwise addition of ligands to metal ions.