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pH and pOH

pH and pOH. 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. pH.

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pH and pOH

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  1. pH and pOH

  2. Ionization of water • Experiments have shown that pure water ionizes very slightly: • 2H2O  H3O+ + OH- • Measurements show that: [H3O+] = [OH-]=1 x 10-7 M • Pure water contains equal concentrations of H3O+ + OH-, so it is neutral.

  3. pH • pH is a measure of the concentration of hydronium ions in a solution. • pH = -log [H3O+] or • pH = -log [H+]

  4. 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)

  5. Example: What is the pH of a solution where [H3O+] = 1 x 10-7 M? • pH = -log [H3O+] • pH = -log(1 x 10-7) • pH = 7.0

  6. Example: What is the pH of a solution where [H3O+] = 1 x 10-5 M? • pH = -log [H3O+] • pH = -log(1 x 10-5) • pH = 5.0 • When acid is added to water, the [H3O+] increases, and the pH decreases.

  7. Example: What is the pH of a solution where [H3O+] = 1 x 10-10 M? • pH = -log [H3O+] • pH = -log(1 x 10-10) • pH = 10.0 • When base is added to water, the [H3O+] decreases, and the pH increases.

  8. The pH Scale 0 7 14 Acid Neutral Base

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

  10. 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

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

  12. 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

  13. 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

  14. 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

  15. Example 4: What is the [OH-] and pOH for the solution in example #3? • [H3O+][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

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

  17. Consider the following buffered solution… HAc  H+ + Ac- H2O  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-

  18. Now, add additional base (OH-) The OH- will combine with the H+ to produce water… H2O  H+ + OH- The H+ comes from the HAc HAc  H+ + Ac-

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

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