Chapter 15: Applications of Aqueous Equilibria

1
Chapter 15 Applications of Aqueous Equilibria

• Buffers
• Common Ion Effect
• Henderson-Hasselbalch Equation
• Buffer Capacity
• Acid-Base Titrations Titration Curves
• Strong Acid-Strong Base Titrations
• Weak Acid-Strong Base Titrations
• Strong Acid-Weak Base Titrations
• Acid-Base Indicators
• Solubility Equilibria
• Calculating Solubility
• Common Ion Effect
• Selective Precipitation
• Qualitative Analysis
• Complex Ion Equilibria

2
Example 1

• Find the pH of a solution that is 0.20M
• KCNO and 0.10 M HCNO.
• Ka for HCNO is 3.5 x 10-4.
• Answer pH 3.76

3
Example 2

• The ratio of HCO3 to H2CO3 in
• human blood is 201. Find the pH of an
• aqueous solution with this composition.
• Answer pH 7.67

4
Example 3

• Find the pH of a buffer that is made of
• 0.20M KCNO and 0.10 M HCNO. (Note
• This is the same as Example 1, but here
• we will double check our answer using
• the Henderson-Hasselbalch equation.)
• Ka for HCNO is 3.5 x 10-4.
• Answer pH 3.76

5
Example 4

• What pH change would result from the
• addition of 5.0 mL of 0.10M HCl to 50.0
• mL of a buffer containing 0.10 M NH3
• and 0.10 M NH4Cl?
• How much would the pH of 50.0 mL of
• pure water change if the same amount
• of acid was added to it?
• Answers DpH-.09 DpH-4.96

6
Example 5

• How many grams of Na2CO3 should be
• added to 1.5 L of 0.20 M NaHCO3 to
• make a buffer of pH 10.00?
• Ka2 of H2CO3 5.6 x 10-11
• Answer 18 g of Na2CO3

7
Example 6

• What is the pH of a buffer made by
• adding 5.0 mL of 0.20 M NaOH to 25.0
• mL of 0.10 M HC2H3O2?
• Answer pH 4.56

8
Example 7

• Calculate the pH change that occurs if
• 1.0 mL of 0.10 M HCl is added to 50.0
• mL of a buffer containing
• 0.30 M pyruvic acid (HC3H3O3) and 0.30 M
  potassium pyruvate? Ka for pyruvic acid is 1.4
  x 10-4.
• 0.0030 M pyruvic acid and 0.0030 M potassium
  pyruvate?
• Answers a. DpH0.00 b. DpH-0.70

9
Example 8

• A 75.0 mL sample of 0.200 M HBr is
• titrated with 0.100 M KOH to a
• phenolphthalein endpoint. How much
• KOH solution is needed to reach the
• equivalence point? What is the pH of
• the solution at the equivalence point?

10
Strong Acid-Strong Base Titration Curves
Figures 15.1 15.2
11
Acid-Base Indicators
12
Example 9

• When a 50.0 mL sample of 0.250 M
• nitrous acid is titrated with 0.100 M
• NaOH, what volume of NaOH is needed
• to reach the equivalence point? What is
• the pH at the equivalence point? What
• is the pH at the halfway point?
• Ka for nitrous acid is 4.0 x 10-4.

13
Weak Acid-Strong Base Titration Curves
14
Example 10

• Find the volume of 0.100 M HCl needed
• to reach the equivalence point in the
• titration of 25.0 mL of 0.100 M NH3.
• Also find the pH of the solution in the
• flask
• Initially
• After 10.0 mL of HCl have been added
• At the halfway point
• At the equivalence point
• After 35.0 mL of HCl have been added
• Kb for NH3 is 1.8 x 10-5

15
Example 11

• A 100.0 mL sample of a weak,
• monoprotic acid with a concentration
• of 0.200 M is titrated with 0.100 M
• NaOH. After 10.0 mL of NaOH have
• been added, the pH is 5.79. What is
• Ka for this acid?

16
Example 12

• Silver bromide has a solubility of 0.133
• mg per 1.00 L of water. Find Ksp for
• silver bromide.

17
Example 13

• Mercury (I) chloride has a Ksp of
• 1.3 x 10-18. Find its solubility in units of
• mole/L and g/L.

18
Example 14

• Which of the following ionic compounds
• is more soluble in water? (i.e. Which
• will dissolve more moles per liter?)
• CaSO4 or CaCO3
• CaSO4 or Ca(OH)2
• Ksp for CaSO4 6.1 x 10-5
• Ksp for CaCO3 8.7 x 10-9
• Ksp for Ca(OH)2 1.3 x 10-6

19
Example 15

• Calcium oxalate has a solubility of
• 6.1x10-3 g/L in water. Find its solubility
• in 0.20 M CaCl2.
• Ksp for CaC2O4 2.3x10-9

20
Example 16

• One type of kidney stones is made of
• calcium phospate. If Ca2 in urine is
• 0.080 g/L, what is the minimum
• molarity of phosphate that will cause
• kidney stones to form?
• Ksp for calcium phosphate 1.3x10-32

21
Example 17

• A 65.0 mL sample of 0.010 M Pb(NO3)2
• was added to a beaker containing 40.0
• mL of 0.035 M KCl. Will a precipitate
• form?

22
Example 18

• What percentage of Ca2 ions remain in
• solution after CaCO3 precipitates when
• 25.0 mL of 0.10 M CaCl2 is added to
• 25.0 mL of 0.10 M Na2CO3?
• Ksp for CaCO3 is 8.7 x 10-9.

23
Example 19

• When 1.0 M AgNO3 is slowly added to a
• solution containing 0.015M Cl- and
• 0.015M CrO42-, what percent of Cl-
• remains in solution when the Ag2CrO4
• begins to precipitate? (i.e. What is the
• maximum separation of Cl- from CrO42-
• that can be achieved?)

24
Example 20

• A solution contains 0.10 M Cd2 and 0.10 M
• Ni2. What concentration of S2- will precipitate
• a maximum amount of one cation without
• precipitating the other?
• Ksp NiS 3.0 x 10-21
• Ksp CdS 1.0 x 10-28

25
Qualitative Analysis Scheme
26
Example 21

• How much Zn2 ion remains in solution
• in a mixture that is 0.010 M Zn(NO3)2
• and 0.10 M NH3?
• Kf Zn(NH3)42 2.9 x 109

27
Example 22

• Calculate the solubility of AgI in
• 0.10 M KCN Kf Ag(CN)2- 5.6 x 1018
• water
• Ksp for AgI 1.5 x 10-16

28
Example 23

• Will nickel (II) hydroxide precipitate from
• in a solution that is 0.0020 M NiSO4,
• 0.010 M NaOH, and 0.10 M NH3?
• Kf for Ni(NH3)62 5.6 x 108
• Ksp for Ni(OH)2 2.0 x 10-15

29
(No Transcript)