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Understanding Ka and Kb Calculations for Weak Acid and Base Reactions

This resource explores the calculations of acid dissociation constant (Ka) and base dissociation constant (Kb) for weak acids and bases. It explains the chemical reactions involved, including the formation of hydronium ions (H3O+) and hydroxide ions (OH-) in solution. It highlights the relationship between Ka, Kb, and the ion product of water (Kw) and provides practical examples for determining the acidity or basicity of solutions by calculating H+ and OH- concentrations. The information is crucial for understanding equilibrium in acid-base chemistry.

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Understanding Ka and Kb Calculations for Weak Acid and Base Reactions

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  1. Ka and Kb Calculations

  2. For Weak Acid Reactions: HA + H2O  H3O+ + A- Ka = [H3O+][A-] Ka < 1 [HA]

  3. For Weak Base Reactions: B + H2O  HB+ + OH- Kb= [HB+][OH-] Kb < 1 [B]

  4. Notice that Ka and Kb expressions look very similar. The difference is that a base produces the hydroxide ion in solution, while the acid produces the hydronium ion in solution.

  5. Another note on this point: H+ and H3O+ are both equivalent terms here. Often water is left completely out of the equation since it does not appear in the equilibrium. This has become an accepted practice. (*However, water is very important in causing the acid to dissociate.)

  6. H2O(l) + H2O(l) <=> H3O+(aq) + OH-(aq) Keq[H2O]2 = Kw = [H3O+][OH-] Kw = 1.0 x 10-14 Kw = Ka x Kb (Another useful equation) Knowing this value allows us to calculate the OH- and H+ concentration for various situations.

  7. [OH-] = [H+] : solution is neutral (in pure water, each of these is 1.0 x 10-7) [OH-] > [H+] : solution is basic [OH-] < [H+] : solution is acidic

  8. Kw = [H+] [OH-] = 1.0 X 10-14 M Calculate [H+] or [OH-] as required for each of the following solutions at 25°C, and state whether the solution is neutral, acidic, or basic. a. 1.0 X 10-5M OH- b. 1.0 X 10-7M OH- c. 10.0 M H+

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