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Chapter 17

Chapter 17. Common Ion Effect. Drill. Use AP Review Drill # 50-53. Objectives. SWBAT Complete Common Ion Effect calculations. Explain how the Common Ion Effect is a special case of Le Chatelier’s Principle. Common Ion Effect. The Common Ion Effect

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Chapter 17

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  1. Chapter 17 Common Ion Effect

  2. Drill • Use AP Review Drill # 50-53

  3. Objectives • SWBAT • Complete Common Ion Effect calculations. • Explain how the Common Ion Effect is a special case of Le Chatelier’s Principle.

  4. Common Ion Effect • The Common Ion Effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.

  5. Uses of Common Ion Effect • The common ion effect plays an important role in determining: • the pH of a solution • and the solubility of a slightly soluble salt.

  6. The Common Ion Effect • When a salt, with the anion of a weak acid, is added to that acid the addition of the salt reverses the dissociation of the acid. • Acid Dissociation: • CH3COOH CH3COO -1 + H+1 • Salt that is added: • NaCH3COO CH3COO -1 + Na+1

  7. Common Ion Effect • Addition of the salt will increase the concentration of the CH3COO-1 ion. • Le Chatelier’s Principle says that the increased ion concentration will reverse the equilibrium reaction. This reversal decreases production of the products and increases production of the reactants.

  8. Equilibrium Special Case • You have probably realized that the Common Ion Effect is just a special case of Le Chatelier’s Principle.

  9. Common Ion Effect • Lowers the percent dissociation of the acid. (If the equilibrium shifts left, you produce less products.) • The same principle applies to salts with the cation of a weak base. • This is an “ICE BOX” problem with an exception. Now you have an initial concentration of the anion.

  10. Common Ion Effect Example • In this situation, the soluble salt of the weak acid's conjugate base simply provides a source of the conjugate base.

  11. Consider acetic acid in water: • HC2H3O2 (aq) ↔ H+ (aq) + C2H3O2- (aq) • Now, add NaC2H3O2, which dissociates completely, we increase the concentration of C2H3O2-1 (aq). • Le Chatelier's Principle : • the equilibrium will shift to the left (towards the reactants), causing • the [H+] to decrease, and therefore • the pH increases!

  12. Common Ion Effect • The common ion effect is the shift in equilibrium that occurs when an ion ALREADY PRESENT in the equilibrium reaction is added.

  13. EXAMPLE • If 0.100 moles of NaC2H3O2 are added to a 1.00 L of 0.100 M solution of acetic acid, HC2H3O2, what is the resultant pH ?

  14. STEP 1: Identify the major species in solution • HC2H3O2 (weak acid) • Na+1 (neither acidic nor basic which means it is a SPECTATOR ION) • C2H3O2-1 (conjugate base of weak acid) • H2O (very weak acid or base, amphoteric)

  15. STEP 2: Identify the important equilibrium reaction(s). In this case it is a reaction that involves both the weak acid and its conjugate base: HC2H3O2 (aq) ↔ H+ (aq) + C2H3O2-1 (aq)

  16. STEP 3: Make an I.C.E. table to determine the equilibrium concentrations of substances: HC2H3O2 (aq) ↔ H+ (aq) + C2H3O2-1 (aq) http://www.chem.ubc.ca/courseware/pH/section12/index.html

  17. STEP 4: Determine and solve the equilibrium constant expression: • Ka = ([H+][C2H3O2-1] / [HC2H3O2] ) • 1.8 x 10-5 = ([H+][C2H3O2-1] / [HC2H3O2] ) • 1.8 x 10-5 = ( x(0.100 + x) / (0.100 - x) ) • Use the 5% assumption. • 1.8 x 10-5 = ( x(0.100) / (0.100) ) • x = [H+] = 1.8 x 10-5

  18. STEP 5: Calculate the pH! • pH = -log(1.8 x 10-5) = 4.74 • The addition of a common ion can also affect the dissociation of a weak base in a similar manner. • See http://www.chem.ubc.ca/courseware/pH/section12/index.html for more information.

  19. Practice Problems • Try B&L #

  20. Wrap Up • How is the Common Ion Effect a special case of Le Chatelier’s Principle?

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