1. Chapter 15 Applications of Aqueous Equilibria

2. The Common Ion Effect A common ion is an ion that is produced by multiple species in solution (other than H) It comes from having a salt dissolved in the same solution as an acid. If the acid is a weak acid, additional ions of the conjugate base may affect the pH of the solution.

3. Consider the Following… HF   H + + F - and NaF  Na + + F - What happens to the equilibrium expression when the NaF dissociates (because it is ionic and that’s what ionic compounds do?) Equilibrium Shifts To Left

4. Check This Out Calculate [H + ] and percent dissociation of HF in a solution of 1.0 M HF (K a = 7.2 x 10 -4 ) and 1.0 M NaF. Compare those values to a 1.0 M HF solution with no NaF…

5. BUFFERS!! Solutions that resist pH change when small amounts of acid or base are added. HA + H 2 O H 3 O + + A - Based on the Common Ion Effect Two Types Weak Acid + Its Salt Weak Base + Its Salt

6. Buffers Try Me A buffered solution contains 1.00 M acetic acid (K a = 1.80 x 10 -5 ) and 1.00 M sodium acetate. Calculate the pH of the solution.

7. Buffy Buff Buffers Calculate the change in pH that occurs when 0.03 mole of solid NaOH is added to 1.0 Liter of the buffered solution from the problem we just finished.

8. Henderson Hasselbalch Useful for calculating the pH of solutions when the ratios [HA]/[A-] is known. pH = pKa + log ([A-]/[HA])

9. Note: When a solution contains equal concentrations of an acid and its conjugate base, [H + ] = K a the pH of the buffer is equal to the log of the K a

10. What if I graphed it all? buffered unbuffered

11. Try This Buffer Problem Calculate the pH of a solution containing 0.75 M lactic acid (K a = 1.4 x 10 -4 ) and 0.25 M sodium lactate. Lactic acid (HC 3 H 5 O 3 ) is a common constituent of biologic systems. For example, it is found in milk and is present in human muscle tissue during exertion. A)Use Henderson Hasselbach OR B) Use an I.C.E. Chart

12. Solubility Equilibria Ion dissociation is an equilibrium process You can write an equilibrium expression for it. Ksp = solubility product constant Solubility product ≠ solubility

13. Calculating Ksp Copper (I) bromide has a measured solubility of 2.0 x 10 -4 mol/L at 25 o C. Calculate its K sp value.

14. Calculating Solubility The Ksp value for copper (II) iodate is 1.4 x 10 -7 at 25 o C. Calculate its solubility at this temperature.

16. Comparing Solubilities # of ions produced matters: If the number is the same AgClAgOHAgI If the number is different CuSAg 2 SBi 2 S 3 chart

17. pH and Solubility Example of Mg(OH) 2 solubility Common ion effect causes milk of magnesia to dissolve in stomach fluids but it is relatively insoluble in non-acidic solutions because of the presence of other OH- ions. If X- is an effective base, (HX is a weak acid) The salt MX will be more soluble in acid. Precipitation occurs as a result of decreased solubility due to the Common Ion Effect

18. Precipitation Reactions To calculate the likelihood of the formation of precipitate formation upon mixing two solutions, we use Only this time we call it the ion product. Q > Ksp Precipitation! Q < Ksp No Precipitation!

19. Try It Out A solution is prepared by adding 750.0 mL of 4.00 x 10 -3 M Ce(NO 3 ) 3 to 300.0 mL of 2.00 x 10 -2 M KIO 3. Will Ce(IO 3 ) 3 precipitate from this solution (K sp = 1.9 x 10 - 10 )

20. Precipitation Try ME A solution is prepared by mixing 150.0 mL of 1.00 x 10 -2 M Mg(NO 3 ) 2 and 250.0 mL of 1.00 x 10 -1 M NaF. Calculate the concentrations of Mg 2+ and F - at equilibrium with solid MgF 2 (Ksp = 6.4 x 10 -9 )

21. Titrations We’ve been doing tons of these. Lets look a little closer. Titration curves are graphical representations of the pH changes the solution experiences as it is titrated. Strength of the acid/base, and the molarity of each affect the features of the curves.

22. Titration Curves The pH at the equivalence point is determined by the strength of the acid and base. The starting pH and ending pH are determined by the pH (concentration) of the acid and the base. Pay attention to the direction of the titration curve! Are you adding base? Are you adding acid?

23. pH at the Equivalence Point For strong acid/strong base, pH = 7 at the equivalence point. When a weak acid or base involved, you need to calculate the pH. –Use stoichiometry and an ICE chart –The pH at the volume half way to the equivalence point is equal to the pKa!

24. Try Me A chemist has synthesized a monoprotic weak acid and wants to determine the K a value. To do so, the chemist dissolves 2.00mmol of the solid acid into 100.0 mL water and titrates the resulting solution with 0.0500M NaOH. After 20.0mL NaOH has been added, the pH is 6.00. What is the K a value for the acid?

25. Try Me Again You titrate an unknown polyprotic acid with 0.75M NaOH. Using the titration curve, determine concentration of the acid, and the Ka values for each hydrogen in the acid.