2. pH and pOH

3. Ionization of water Experiments have shown that pure water ionizes very slightly: 2H 2 O  H 3 O + + OH - Measurements show that: [H 3 O + ] = [OH - ]=1 x 10 -7 M Pure water contains equal concentrations of H 3 O + + OH -, so it is neutral.

4. pH pH is a measure of the concentration of hydronium ions in a solution. pH = -log [H 3 O + ] or pH = -log [H + ]

5. Sig. Figs. for Logarithms The rule is that the number of decimal places in the log is equal to the number of significant figures in the original number. Example: [H+] = 1.0 x 10 -9 M (2 significant figures) pH = -log(1.0 x 10 -9 ) = 9.00 (2 decimal places)

6. Example: What is the pH of a solution where [H 3 O + ] = 1 x 10 -7 M? pH = -log [H 3 O + ] pH = -log(1 x 10 -7 ) pH = 7.0

7. Example: What is the pH of a solution where [H 3 O + ] = 1 x 10 -5 M? pH = -log [H 3 O + ] pH = -log(1 x 10 -5 ) pH = 5.0 When acid is added to water, the [H 3 O + ] increases, and the pH decreases.

8. Example: What is the pH of a solution where [H 3 O + ] = 1 x 10 -10 M? pH = -log [H 3 O + ] pH = -log(1 x 10 -10 ) pH = 10.0 When base is added to water, the [H 3 O + ] decreases, and the pH increases.

9. The pH Scale AcidNeutral Base 0 7 14

10. pOH pOH is a measure of the concentration of hydroxide ions in a solution. pOH = -log [OH - ]

11. Example: What is the pOH of a solution where [OH - ] = 1 x 10 -5 M? pOH = -log [OH - ] pOH = -log(1 x 10 -5 ) pOH = 5.0

12. How are pH and pOH related? At every pH, the following relationships hold true: [H + ] [OH - ] = 1 x 10 -14 M pH + pOH = 14

13. Example 1: What is the pH of a solution where [H + ] = 3.4 x 10 -5 M? pH = -log [H + ] pH = -log(3.4 x 10 -5 M) pH = 4.47

14. Example 2: The pH of a solution is measured to be 8.86. What is the [H + ] in this solution? pH = -log [H + ] 8.86 = -log [H + ] -8.86 = log [H + ] [H + ] = antilog (-8.86) [H + ] = 10 -8.86 [H + ] = 1.4 x 10 -9 M

15. Example 3: What is the pH of a solution where [H + ] = 5.4 x 10 -6 M? pH = -log [H + ] pH = -log(5.4 x 10 -6 ) pH = 5.27

16. Example 4: What is the [OH - ] and pOH for the solution in example #3? [H 3 O + ][OH - ]= 1 x 10 -14 (5.4 x 10 -6 )[OH - ] = 1 x 10 -14 [OH - ] = 1.9 x 10 -9 M pH + pOH = 14 pOH = 14 – 5.27 = 8.73

17. Buffered Solutions A solution of a weak acid and a common ion is called a buffered solution.

18. Consider the following buffered solution… HAc  H + + Ac - H 2 O  H + + OH - Add additional acid…(H + ) The H + will combine with the Ac - producing HAc. There is an excess of Ac - from the common ion salt. HAc  H + + Ac -

19. Now, add additional base (OH - ) The OH - will combine with the H + to produce water… H 2 O  H + + OH - The H+ comes from the HAc HAc  H + + Ac -

20. Thus, the solution maintains it’s pH in spite of added acid or base.