1. Applications of Aqueous Equilibria
Chemistry: McMurry and Fay, 6th Edition
Chapter 15: Applications of Aqueous Equilibria
4/27/2017 2:44:03 AM
Chapter 15
Applications of Aqueous Equilibria
Copyright © 2011 Pearson Prentice Hall, Inc.

2. Chemistry: McMurry and Fay, 6th Edition
Chapter 15: Applications of Aqueous Equilibria
4/27/2017 2:44:03 AM
The Common-Ion Effect
Common-Ion Effect: The shift in the position of an equilibrium on addition of a substance that provides an ion in common with one of the ions already involved in the equilibrium
CH3CO2H(aq) + H2O(l)
H3O+(aq) + CH3CO2–(aq)
What are the ions present in the solution?
Copyright © 2011 Pearson Prentice Hall, Inc.

3. The Common-Ion Effect CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2–(aq)
Chemistry: McMurry and Fay, 6th Edition
Chapter 15: Applications of Aqueous Equilibria
The Common-Ion Effect
4/27/2017 2:44:03 AM
Le Châtelier’s Principle
CH3CO2H(aq) + H2O(l)
H3O+(aq) + CH3CO2–(aq)
The addition of acetate ion to a solution of acetic acid suppresses the dissociation of the acid. The equilibrium shifts to the left.
Copyright © 2011 Pearson Prentice Hall, Inc.

4. Chapter 15: Applications of Aqueous Equilibria
4/27/2017
15.3 Buffer Solutions
Buffer Solution: A solution which contains a weak acid and its conjugate base and resists drastic changes in pH.
Weak acid +
Conjugate base
CH3CO2H + CH3CO21-
HF + F1-
NH41+ + NH3
H2PO41- + HPO42-
For
Example:
Acidic/basic salts with the weak base/acid.
Copyright © 2008 Pearson Prentice Hall, Inc.

5. Chemistry: McMurry and Fay, 6th Edition
Chapter 15: Applications of Aqueous Equilibria
4/27/2017 2:44:03 AM
Example
Calculate the pH of a solution that is prepared by dissolved 0.10 mol of acetic acid and 0.10 mol sodium acetate in enough water to make L of solution. Is this a buffer solution?
Ka = 1.8 x 10-5
25 °C p
This assumes the complete dissociation of sodium acetate as shown previously.
Copyright © 2011 Pearson Prentice Hall, Inc.

6. Example
Calculate the pH of 0.15M HF and 0.25NaF mixture. Is this a buffer solution?

7. Example Calculate the pH of 100.0 mL DI water
Calculate the new pH after adding 1.0 mL of 0.10M HCl to the above water solution.

8. Buffer Solutions H3O1+(aq) + A1-(aq) HA(aq) + H2O(l) Conjugate base
(M+A-)
Weak acid
Add a small amount of base (-OH) to a buffer solution
Acid component of solution neutralizes the added base
Addition of OH1- to a buffer:
H2O(l) + A1-(aq)
HA(aq) + OH1-(aq)
100%

9. Buffer Solutions H3O1+(aq) + A1-(aq) HA(aq) + H2O(l) Conjugate base
(M+A-)
Weak acid
Add a small amount of acid (H3O+) to a buffer solution
Base component of solution neutralizes the added acid
Addition of H3O1+ to a buffer:
H2O(l) HA(aq)
A1-(aq) + H3O1+(aq)
100%

10. Buffer Solutions
The addition of –OH or H3O+ to a buffer solution will change the pH of the solution, but not as drastically as the addition of –OH or H3O+ to a non-buffered solution

11. Buffer Solutions

12. Example
pH of human blood (pH = 7.4) controlled by conjugated acid-base pairs (H2CO3/HCO3-). Write an equation for this buffer mixture then neutralization equation for the following effects
With addition of HCl
With addition of NaOH

13. Example
50.0 mL of M HCl was added to a .100L buffer consisting of moles of sodium acetate and moles of acetic acid. What is the pH of the buffer before and after the addition of the acid? Ka of acetic acid is 1.7 x Assume the volume is constant

14. Example
Calculate the pH of 0.100L of a buffer solution that is 0.25M in HF and 0.50 M in NaF.
What is the change in pH on addition of moles KOH
Calculate the pH after addition of moles HBr
Assume the volume remains constant
Ka = 3.5 x 10-4

15. Example
Calculate the pH of the buffer that results from mixing 60.0mL of M HCHO2 and 15.0 mL of 0.500M NaCHO2
Ka = 1.7 x 10-4
Calculate the pH after addition of 10.0 mL of MHBr. Assume volume is additive

16. Buffer Capacity
A measure of amount of acid or base that the solution can absorb without a significant change in pH.
Depends on how many moles of weak acid and conjugated base are present.
For an equal volume of solution: the more concentrated the solution, the greater buffer capacity
For an equal concentration: the greater the volume, the greater the buffer capacity

17. Example
The following pictures represent solutions that contain a weak acid HA and/ or its sodium salt NaA. (Na+ ions and solvent water molecules have been omitted for clarity
Which of the solutions are buffer solution?
Which solution has the greatest buffer capacity?

18. Example
What is the maximum amount of acid that can be added to a buffer made by the mixing of 0.35 moles of sodium hydrogen carbonate with 0.50 moles of sodium carbonate? How much base can be added before the pH will begin to show a significant change?

19. 15.4 The Henderson-Hasselbalch Equation
Chapter 15: Applications of Aqueous Equilibria
4/27/2017
15.4 The Henderson-Hasselbalch Equation
Weak acid
Conjugate base
H3O1+(aq) + Base(aq)
Acid(aq) + H2O(l)
[H3O1+][Base]
[Acid]
[Acid]
[Base]
[H3O1+] = Ka
Ka =
[Base]
[Acid]
pH = pKa + log
Copyright © 2008 Pearson Prentice Hall, Inc.

20. Examples
Calculate the pH of a buffer solution that is 0.50 M in benzoic acid (HC7H5O2) and M in sodium benzoate (NaC7H5O2). Ka = 6.5 x 10-5

21. Example
How would you prepare a NaHCO3-Na2CO3 buffer solution that has pH = Ka2 = 4.7 x 10-11

22. Example
You prepare a buffer solution of .323 M NH3 and (NH4)2SO4. What molarity of (NH4)2SO4 is necessary to have a pH of 8.6? (pKb NH3= 4.74)