1. Aqueous Equilibria Chapter 17 Additional Aspects of Aqueous Equilibria You love Chemistry You might have doubts, but deep, deep, deep down you know there is a little love for the central science!

2. Aqueous Equilibria THE COMMON ION EFFECT How do they effect dissociation?

3. Aqueous Equilibria Compare: Calculate the pH of a 0.25 M propionic acid solution (K a =1.3 x 10 -5 ) Calculate the pH of a 0.25 M propionic acid solution that also has 0.10 M sodium propionate add.

4. Aqueous Equilibria Common Ion Effect How does LeChatelier support the previous calculations?

5. Aqueous Equilibria Common Ion Effect Summarize the Common Ion Effect:

6. Aqueous Equilibria Buffers: Solutions of a weak conjugate acid-base pair. They are particularly resistant to pH changes, even when strong acid or base is added.

7. Aqueous Equilibria How do Buffers Resist pH Changes? Consider a buffer composed of equal concentrations of nitrous acid and nitrite ion. 1) Does this meet the criteria for a buffer? 2) What would happen if a volume of HCl was added to the buffer? 3) What would happen if a volume of NaOH was added to the buffer

8. Aqueous Equilibria How do Buffers Resist pH Changes? The pH of a buffer will change somewhat, but not significantly.  The Balance between the conjugate acid/base pair is disrupted  Either the conjugate acid or the conjugate base will be present in a higher concentration after the addition  This will cause a minor change to the pH

9. Aqueous Equilibria Calculating the pH of a Buffer Solution In order to calculate the pH of a buffer solution you will need to use the Henderson-Hasselbalch Equation: pH = pK a + log [base] [acid]

10. Aqueous Equilibria Henderson–Hasselbalch Equation What is the pH of a buffer that is 0.12 M in benzoic acid and 0.20 M in sodium benzoate? K a for benzoic acid is 6.3  10 −5.

11. Aqueous Equilibria Another Application of the Henderson-Hasselbalch Equation How many moles of NH 4 Cl must be added to 2.0 L of 0.10 M NH 3 to form a buffer whose pH is 9.00? (Assume that the addition of NH 4 Cl does not change the volume of the solution.) Kb for NH 3 is 1.8 x 10 -5.

12. Aqueous Equilibria pH Range and Buffer Capacity The pH range is the range of pH values over which a buffer system works effectively. It is best to choose an acid with a pK a close to the desired pH. Buffer Capacity is the amount of acid or base a buffer can neutralize before significant pH changes  The higher the molarity or volume of the conjugate pairs, the greater the capacity of the buffer.

13. Aqueous Equilibria Calculating pH Changes in Buffers A buffer is made by adding 0.300 mol HC 2 H 3 O 2 and 0.300 mol NaC 2 H 3 O 2 to enough water to make 1.00 L of solution. Calculate the original pH of the buffer and the pH after 0.020 mol of NaOH is added. K a of acetic acid is 1.8 x 10 -5

14. Aqueous Equilibria Homework Ch. 17: 15, 17, 21, 23, 25

15. Aqueous Equilibria Solubility Product Constant Using Equilibrium to Determine the dissociation of a solid in solution.

16. Aqueous Equilibria Solubility Products Consider the equilibrium that exists in a saturated solution of BaSO 4 in water: Write the Equilibrium Expression for this reaction. BaSO 4 (s) Ba 2+ (aq) + SO 4 2− (aq)

17. Aqueous Equilibria Solubility Products The equilibrium constant expression for this equilibrium is K sp = [Ba 2+ ] [SO 4 2− ] where the equilibrium constant, K sp, is called the solubility product. The solubility product defines the dissociation of the solid in solution

18. Aqueous Equilibria Practice Problem Write separate expressions for the solubility product constant for CaF 2 and Silver Sulfate

19. Aqueous Equilibria Practice Problem Determine the concentration of each ion in a saturated solution of zinc hydroxide. Zinc hydroxide has a K sp =3.0 x 10 -16

20. Aqueous Equilibria Solubility Products K sp is not the same as solubility. Solubility is generally expressed as the mass of solute dissolved in 1 L (g/L) or 100 mL (g/mL) of solution, or in mol/L (M).

21. Aqueous Equilibria K sp from Solubility Data Solid silver chromate is added to pure water at 25°C. Some of the solid remains undissolved at the bottom of the flask. Analysis of the equilibrated solution shows that its silver ion concentration is 1.3  10 –4 M. Assuming that Ag 2 CrO 4 dissociates completely in water and that there are no other important equilibria involving the Ag + or CrO 4 2– ions in the solution, calculate K sp for this compound.

22. Aqueous Equilibria Solubility from K sp The K sp for CaF 2 is 3.9  10 –11 at 25°C. Assuming that CaF 2 dissociates completely upon dissolving and that there are no other important equilibria affecting its solubility, calculate the solubility of CaF 2 in grams per liter.

23. Aqueous Equilibria Refresher Problems Getting our mind back into Chemistry Mode.

24. Aqueous Equilibria Equilibrium Expressions Write equilibrium expressions for the following reactions: 1) Ni(OH) 2 (s)  Ni 2+ (aq) + 2 OH - (aq) 2) 2NOBr (g)  2 NO (g) + Br 2 (g) 3) HClO 3 (ag)  H + (aq) + ClO 3 - (aq) 4) NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH - (aq)

25. Aqueous Equilibria pH What is the pH of a solution that consists of 250 ml of 0.75 M hydrofluoric acid and 2.1 grams of sodium fluoride? (K a =6.8 x 10 -4 )

26. Aqueous Equilibria Solubility What’s the molar concentration of each ion in an equilibrated solution of lead (II) fluoride. K sp of lead (II) fluoride= 3.6 x 10 -8 Can we figure out the pH of this solution?

27. Aqueous Equilibria Buffer Solutions What is the pH of a 500.0 mL buffer solution that consists of 1.25 M acetic acid and 1.00 M sodium acetate if 15.0 mL of 0.750 M nitric acid is added to it? K a of acetic acid is 1.80 x 10 -5

28. Aqueous Equilibria HOMEWORK Read and Take Notes on section 17.3 Acid- Base Titrations

29. Aqueous Equilibria Factors that Affect Solubility

30. Aqueous Equilibria Factors that Affect Solubility From your understanding of equilibrium and how it relates to acids/bases, what would be some factors that either increase or decrease the dissociation of a solid in solution?

31. Aqueous Equilibria Factors Affecting Solubility 1) The Common-Ion Effect  If one of the ions in a solution equilibrium is already dissolved in the solution, the equilibrium will shift to the left and the solubility of the salt will decrease. BaSO 4 (s) Ba 2+ (aq) + SO 4 2− (aq)

32. Aqueous Equilibria Factors Affecting Solubility 2) pH  If a substance has a basic anion, it will be more soluble in an acidic solution.  Substances with acidic cations are more soluble in basic solutions.

33. Aqueous Equilibria Factors Affecting Solubility pH  Explain how the solubility of Mg(OH) 2 would be affected by the presence of an acid.

34. Aqueous Equilibria Practice Problem Which of the following substances will be more soluble in acidic solution than in basic solution: (a)Ni(OH) 2 (s) (b) CaCO 3 (s) (c) AgCl(s) (d) BaF 2 (s)

35. Aqueous Equilibria Evaluating a Solution Will more solid dissolve or will a precipitate form?

36. Aqueous Equilibria Will a Precipitate Form? In a solution,  If Q = K sp, the system is at equilibrium and the solution is saturated.  If Q < K sp, more solid will dissolve until Q = K sp.  If Q > K sp, the salt will precipitate until Q = K sp.

37. Aqueous Equilibria Practice Problem Will a precipitate form when 0.10 L of 8.0  10 –3 M Pb(NO 3 ) 2 is added to 0.40 L of 5.0  10 –3 M Na 2 SO 4 ? (K sp for PbSO 4 =6.3 x 10 -9 )