1. Ionic Equilibria: Part II Buffers and Titration Curves16
Ionic Equilibria: Part II
Buffers and Titration Curves

2. The Common Ion Effect and Buffer Solutions
Example 16-1: Calculate the concentration of H+and the pH of a solution that is 0.15 M in acetic acid and 0.15 M in sodium acetate.
This is another equilibrium problem with a starting concentration for both the acid and anion.

3. The Common Ion Effect and Buffer Solutions
Substitute the quantities determined in the previous relationship into the ionization expression.

4. The Common Ion Effect and Buffer Solutions
Apply the simplifying assumption to both the numerator and denominator.

5. Weak Bases plus Salts of Weak Bases
Example 16-2: Calculate the concentration of OH- and the pH of the solution that is 0.15 M in aqueous ammonia, NH3, and 0.30 M in ammonium nitrate, NH4NO3.

6. Weak Bases plus Salts of Weak Bases
Substitute the quantities determined in the previous relationship into the ionization expression for ammonia.

7. Buffering Action
Example 16-3: If mole of gaseous HCl is added to 1.00 liter of a buffer solution that is M in aqueous ammonia and M in ammonium chloride, how much does the pH change? Assume no volume change due to addition of the HCl.
Calculate the pH of the original buffer solution.

8. Buffering Action

9. Buffering Action
Next, calculate the concentration of all species after the addition of the gaseous HCl.
The HCl will react with some of the ammonia and change the concentrations of the species.
This is another limiting reactant problem.

10. Buffering Action

11. Buffering Action
Using the concentrations of the salt and base and the Henderson-Hassselbach equation, the pH can be calculated.

12. Buffering Action

13. Buffering Action
Finally, calculate the change in pH.

14. Buffering Action
Example 16-4: If mole of NaOH is added to 1.00 liter of solution that is M in aqueous ammonia and M in ammonium chloride, how much does the pH change? Assume no volume change due to addition of the solid NaOH.
You do it!

15. Buffering Action
pH of the original buffer solution is 8.95, from above.
First, calculate the concentration of all species after the addition of NaoH.
NaOH will react with some of the ammonium chloride.
The limiting reactant is the NaOH.

16. Buffering Action

17. Buffering Action
Calculate the pH using the concentrations of the salt and base and the Henderson-Hasselbach equation.

18. Buffering Action
Calculate the change in pH.

19. Preparation of Buffer Solutions
Example 16-5: Calculate the concentration of H+ and the pH of the solution prepared by mixing 200 mL of M acetic acid and 100 mL of M sodium hydroxide solutions.
Determine the amounts of acetic acid and sodium hydroxide prior to the acid-base reaction.

20. Preparation of Buffer Solutions
Sodium hydroxide and acetic acid react in a 1:1 mole ratio.

21. Preparation of Buffer Solutions
After the two solutions are mixed, the total volume of the solution is 300 mL (100 mL of NaOH mL of acetic acid).
The concentrations of the acid and base are:

22. Preparation of Buffer Solutions
Substitution of these values into the ionization constant expression (or the Henderson-Hasselbach equation) permits calculation of the pH.

23. Preparation of Buffer Solutions
Example 16-6:Calculate the number of moles of solid ammonium chloride, NH4Cl, that must be used to prepare 1.00 L of a buffer solution that is 0.10 M in aqueous ammonia, and that has a pH of 9.15.
Because pH = 9.15, the pOH can be determined.

24. Preparation of Buffer Solutions
The appropriate equilibria representations are:

25. Preparation of Buffer Solutions
Substitute into the ionization constant expression (or Henderson-Hasselbach equation) for aqueous ammonia

26. Preparation of Buffer Solutions

27. Synthesis Question
Bufferin is a commercially prepared medicine that is literally a buffered aspirin. How could you buffer aspirin? Hint - what is aspirin?

28. Synthesis Question
Aspirin is acetyl salicylic acid. So to buffer it all that would have to be added is the salt of acetyl salicylic acid.