1. Chapter 15
Acid–Base Equilibria

2. 15.1 Solutions of Acids or Bases Containing a Common Ion
Buffered Solutions
Buffering Capacity
15.4 Titrations and pH Curves
Acid–Base Indicators
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3. An application of Le Châtelier’s principle.
Common Ion Effect
Shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction.
An application of Le Châtelier’s principle.
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4. HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)
Example
HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)
Addition of NaCN will shift the equilibrium to the left because of the addition of CN-, which is already involved in the equilibrium reaction.
A solution of HCN and NaCN is less acidic than a solution of HCN alone.
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5. Key Points about Buffered Solutions
Buffered Solution – resists a change in pH.
They are weak acids or bases containing a common ion.
After addition of strong acid or base, deal with stoichiometry first, then the equilibrium.
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6. Henderson–Hasselbalch Equation
For a particular buffering system (conjugate acid–base pair), all solutions that have the same ratio [A–] / [HA] will have the same pH.
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7. Exercise
What is the pH of a buffer solution that is 0.45 M acetic acid (HC2H3O2) and 0.85 M sodium acetate (NaC2H3O2)? The Ka for acetic acid is 1.8 × 10–5.
pH = 5.02
pH = –logKa + log([C2H3O2–] / [HC2H3O2]) = –log(1.8 × 10–5) + log(0.85 M / 0.45 M) = 5.02
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9. Determined by the magnitudes of [HA] and [A–].
The amount of protons or hydroxide ions the buffer can absorb without a significant change in pH.
Determined by the magnitudes of [HA] and [A–].
A buffer with large capacity contains large concentrations of the buffering components.
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10. Buffer capacity measures how well a solution resists changes in pH when acid or base is added. The greater the buffer capacity, the less the pH changes. 
The amount of H+ or OH- that buffered solution can absorb without a significant change in pH

11. 2) Magnitudes of [HA] and [A-]
 the capacity of a buffered soln.
Ex : soln A : 5.00 M HOAc M NaOAc
soln B : 0.05 M HOAc M NaOAc
pH change when 0.01 mol of HCl(g) is added

12. 3) [A-] / [HA] ratio  the pH of a buffered soln.

13. The pH Curve for the Titration of 50. 0 mL of 0. 200 M HNO3 with 0
The pH Curve for the Titration of 50.0 mL of M HNO3 with M NaOH
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14. The pH Curve for the Titration of 100. 0 mL of 0. 50 M NaOH with 1
The pH Curve for the Titration of mL of 0.50 M NaOH with 1.0 M HCI
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15. Concept Check
Consider a solution made by mixing 0.10 mol of HCN (Ka = 6.2 x 10–10) with mol NaOH in 1.0 L of aqueous solution.
What are the major species immediately upon mixing (that is, before a reaction)?
HCN, Na+, OH–, H2O
Major Species: HCN, Na+, OH-, H2O.
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16. K = 1.8 x 109 Let’s Think About It… HC2H3O2(aq) + OH–  C2H3O2–(aq)
Before
0.20 mol mol
Change
–0.030 mol –0.030 mol
mol
After
0.17 mol
0.030 mol
K = 1.8 x 109
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17. Steps Toward Solving for pH
HC2H3O2(aq) +
H2O
H3O+
+ C2H3O2-(aq)
Initial
0.170 M ~0
0.030 M
Change –x +x
Equilibrium
0.170 – x x x
Ka = 1.8 x 10–5
pH = 3.99
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18. Exercise
Calculate the pH of a mL solution of M acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5.
pH = 2.87
pH = 2.87
Major Species: HC2H3O2, H2O
Dominant reaction: HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2-(aq)
The mL is not necessary to know for this problem.
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19. Concept Check
Calculate the pH of a solution made by mixing mL of a M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, and 50.0 mL of a 0.10 M NaOH solution.
pH = 4.74
pH = 4.74
Major Species: HC2H3O2, Na+, OH-, H2O
Dominant reaction: HC2H3O2(aq) + OH-(aq)  H2O + C2H3O2-(aq)
In general, the best acid (HC2H3O2 in this case) will tend to react with the best base (OH- will always be the best base, if present). Discuss how to calculate the K value for this reaction (note -- it is not a Ka or Kb expression). In this case, K = Ka/Kw = 1.8 x 109. We can assume the reaction goes to completion. In this problem, OH- is limiting.
After the reaction takes place, the major species are: HC2H3O2, C2H3O2-, H2O
Primary reaction is: HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2-(aq)
When students solve this, make sure they include the initial concentration of the conjugate base (they tend to forget and call it zero).
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20. Concept Check
Calculate the pH of a solution at the equivalence point when mL of a M solution of acetic acid (HC2H3O2), which has a Ka value of 1.8 x 10–5, is titrated with a 0.10 M NaOH solution.
pH = 8.72
pH = 8.72
Major Species: HC2H3O2, Na+, OH-, H2O
Dominant reaction: HC2H3O2(aq) + OH-(aq)  H2O + C2H3O2-(aq)
In general, the best acid (HC2H3O2 in this case) will tend to react with the best base (OH- will always be the best base, if present).
At the equivalence point, there are the same number of moles of HC2H3O2(aq) + OH-(aq) to exactly react due to the 1:1 ratio in the balanced equation. This turns out to be mol of both the acid and base in this case. This means that mL of NaOH had to be added.
After the reaction takes place, the major species are: C2H3O2-, H2O, Na+
Primary reaction is: C2H3O2-(aq) + H2O  OH-(aq) + HC2H3O2(aq)
Now use an ICE chart to determine the equilibrium concentrations and finally the pH. Take note that the initial concentration of C2H3O2-(aq) is M (0.010 mol / L).
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21. The pH Curve for the Titration of 50. 0 mL of 0. 100 M HC2H3O2 with 0
The pH Curve for the Titration of 50.0 mL of M HC2H3O2 with M NaOH
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22. The pH Curves for the Titrations of 50. 0-mL Samples of 0
The pH Curves for the Titrations of 50.0-mL Samples of 0.10 M Acids with Various Ka Values with 0.10 M NaOH
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23. The pH Curve for the Titration of 100. 0mL of 0. 050 M NH3 with 0
The pH Curve for the Titration of 100.0mL of M NH3 with 0.10 M HCl
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24. Kjeldahl Nitrogen Analysis : developed in 1883, the analysis remains one of the most widely used methods for determining nitrogen in organic substances such as protein, cereal, and flour.
P.228

25. Melamine contaminated milk

26. Marks the end point of a titration by changing color.
The equivalence point is not necessarily the same as the end point (but they are ideally as close as possible).
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27. The Acid and Base Forms of the Indicator Phenolphthalein
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28. The Methyl Orange Indicator is Yellow in Basic Solution and Red in Acidic Solution
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29. HIn(aq)  H+(aq) + In-(aq) pKHIn
1) Usually a weak organic acid or base that has distinctly different colors in its nonionized & ionized forms.
HIn(aq)  H+(aq) + In-(aq) pKHIn
nonionized ionized
form form

30. 1) 2) 3)

31. Useful pH Ranges for Several Common Indicators
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