1. Acid-Base Equilibria and Solubility Equilibria
Chapter 16
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2. 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)
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3. A buffer solution is a solution of: A weak acid or a weak base and
The salt of the weak acid or weak base
Both must be present!
A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base.
Consider an equal molar mixture of CH3COOH and CH3COONa
Add strong acid
H+ (aq) + CH3COO- (aq) CH3COOH (aq)
Add strong base
OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)
16.3

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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.
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5. Which of the following are buffer systems? (a) KF/HF
(b) KBr/HBr, (c) Na2CO3/NaHCO3
(a) KF is a weak acid and F- is its conjugate base
buffer solution
(b) HBr is a strong acid
not a buffer solution
(c) CO32- is a weak base and HCO3- is its conjugate acid
buffer solution
16.3

6. Henderson-Hasselbalch Equation
Useful for calculating pH when the [A]/[HA] ratios are known.
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7. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOONa?
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8. 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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9. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of M NaOH to 80.0 mL of the buffer solution?
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10. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Buffering Capacity
. . . represents the amount of H+ or OH the buffer can absorb without a significant change in pH.
Buffer solution M HNO2/0.1M KNO2
Buffer solution M HNO2/1.0M KNO2
Both solutions buffer to the same pH
Buffer solution 2 has a larger capacity
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11. Titrations
In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.
Equivalence point – the point at which the reaction is complete (stoichiometry)
Indicator – substance that changes color at (or near) the
equivalence point
Endpoint – point where indicator changes color
Slowly add base
to unknown acid
UNTIL
The indicator
changes color
(pink)
4.7

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Titration (pH) Curve
A plot of pH of the solution being analyzed as a function of the amount of titrant added.
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13. Strong Acid-Strong Base Titrations
NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)
OH- (aq) + H+ (aq) H2O (l)
16.4

14. Weak Acid-Strong Base Titrations
CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)
CH3COOH (aq) + OH- (aq) CH3COO- (aq) + H2O (l)
At equivalence point (pH > 7):
CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)
16.4

15. pH = pKa at halfway point
Acetic Acid Ka=1.8x pKa=4.74
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16. Strong Acid-Weak Base Titrations
HCl (aq) + NH3 (aq) NH4Cl (aq)
H+ (aq) + NH3 (aq) NH4Cl (aq)
At equivalence point (pH < 7):
NH4+ (aq) + H2O (l) NH3 (aq) + H+ (aq)
16.4

17. pH = pKa at halfway point
Ammonia Kb=1.8x10-5 Ka=5.56x pKa=9.25
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18. Strong Acid – Strong Base Titration
Step 1 - A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species.
Step 2 - Determine pH based on concentration of excess reactant (strong acid or base leftover)
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19. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Consider the titration of 80.0mL of 0.10 M Sr(OH)2 by 0.4M HCl. Calculate the pH after the following volumes of HCl have been added.
0.0 mL
20.0 mL
30.0 mL
40.0 mL
80.0 mL
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20. Weak Acid - Strong Base Titration Weak Base – Strong Acid 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 or base equilibrium and calculate pH.
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Consider the titration of 100.0mL of 0.10 M HC2H3O2 by 0.2M KOH. Calculate the pH after the following volumes of HCl have been added.
0.0mL
20.0 mL
25.0 mL
40.0 mL
50.0 mL
100.0 mL
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22. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Acid-Base Indicator
. . . marks the end point of a titration by changing color.
commonly weak acids (HIn) that change color when they dissociate (In-)
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23. Acid-Base Indicators HIn (aq) H+ (aq) + In- (aq) [HIn]  10
Color of acid (HIn) predominates
 10
[HIn]
[In-]
Color of conjugate base (In-) predominates
16.5

24. The titration curve of a strong acid with a strong base.
16.5

25. Which indicator(s) would you use for a titration of HNO2 with KOH ?
Weak acid titrated with strong base.
At equivalence point, will have conjugate base of weak acid.
At equivalence point, pH > 7
Use cresol red or phenolphthalein
16.5

26. Solubility Product For solids dissolving to form aqueous solutions.
Ksp = solubility product constant
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27. Solubility Equilibria
AgCl (s) Ag+ (aq) + Cl- (aq)
Ksp = [Ag+][Cl-]
Ksp is the solubility product constant
MgF2 (s) Mg2+ (aq) + 2F- (aq)
Ksp = [Mg2+][F-]2
Ag2CO3 (s) Ag+ (aq) + CO32- (aq)
Ksp = [Ag+]2[CO32-]
Ca3(PO4)2 (s) Ca2+ (aq) + 2PO43- (aq)
Ksp = [Ca2+]3[PO43-]2
Dissolution of an ionic solid in aqueous solution:
Q < Ksp
Unsaturated solution
No precipitate
Q = Ksp
Saturated solution
Q > Ksp
Supersaturated solution
Precipitate will form
16.6

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Solubility Product
Bi2S3(s)  2Bi3+(aq) + 3S2(aq)
“Solubility” = s = concentration of Bi2S3 that dissolves, which equals
1/2[Bi3+] and 1/3[S2].
Note:Ksp is constant (at a given temperature)
s is variable
(especially with a common ion present)
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29. Molar solubility (mol/L) is the number of moles of solute dissolved in 1 L of a saturated solution.
Solubility (g/L) is the number of grams of solute dissolved in 1 L of a saturated solution.
16.6

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What is the solubility of silver chloride in g/L ?
(AgCl Ksp= 1.6 x 10-10)
What concentration of Ag is required to precipitate ONLY AgBr in a solution that contains both Br- and Cl- at a concentration of 0.02 M? (AgBr Ksp= 7.7 x 10-13)
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31. The Common Ion Effect and Solubility
The presence of a common ion decreases the solubility of the salt.
What is the molar solubility of AgBr in (a) pure water and (b) M NaBr?
16.8

32. pH and Solubility Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq)
The presence of a common ion decreases the solubility.
Insoluble bases dissolve in acidic solutions
Insoluble acids dissolve in basic solutions
Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq)
Ksp = [Mg2+][OH-]2 = 1.2 x 10-11
Ksp = (s)(2s)2 = 4s3
(Lower [OH-])
At pH less than 10.45
4s3 = 1.2 x 10-11
Increase solubility of Mg(OH)2
s = 1.4 x 10-4 M
[OH-] = 2s = 2.8 x 10-4 M
At pH greater than 10.45
(Raise [OH-])
pOH = pH = 10.45
Decrease solubility of Mg(OH)2
16.9

33. Complex Ion Equilibria and Solubility
A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions.
Co2+ (aq) + 4Cl- (aq) CoCl4 (aq) 2-
The formation constant or stability constant (Kf) is the equilibrium constant for the complex ion formation.
Kf =
[CoCl4 ]
[Co2+][Cl-]4 2-
Co(H2O)6 2+
CoCl4 2-
stability of complex Kf
16.10

34. 16.10

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If 2.5g of CuSO4 are dissolved in 900mL of .30M NH3, what are the concentrations of Cu2+, Cu(NH3)42+, and NH3 at equilibrium? (Cu(NH3)42+ Kf =5.0x1013 )
Calculate the molar solubility of AgCl in 1.0M NH3. (Ag(NH3)2+ Kf=1.5x AgCl Ksp=1.6x10-10)
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36. 16.11

37. Qualitative Analysis of Cations
16.11