Applications of Aqueous Equilibria

1. Applications of Aqueous Equilibria AP Chemistry Seneca Valley SHS Chapter 15

2. The Common Ion Effect • The solubility of a partially soluble salt is decreased when a common ion is added. • Consider the equilibrium established when acetic acid, HC2H3O2, is added to water. • At equilibrium H+ and C2H3O2- are constantly moving into and out of solution, but the concentrations of ions is constant and equal. • If a common ion is added, e.g. C2H3O2- from NaC2H3O2 (which is a strong electrolyte) then [C2H3O2-] increases and the system is no longer at equilibrium. • So, [H+] must decrease.

3. The Common Ion Effect When a solution contains a salt having an ion common with one in equilibrium, the position of the equilibrium is driven away from the side containing that ion. Compare the pH values for sample problems 5 and 7 to see this!

4. Buffered Solutions • Composition and Action of Buffered Solutions • A buffer resists a change in pH when a small amount of OH- or H+ is added. • A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X-): • The Ka expression is

5. Buffered Solutions • Composition and Action of Buffered Solutions • When OH- is added to the buffer, the OH- reacts with HX to produce X- and water. But, the [HX]/[X-] ratio remains more or less constant, so the pH is not significantly changed. • When H+ is added to the buffer, X- is consumed to produce HX. Once again, the [HX]/[X-] ratio is more or less constant, so the pH does not change significantly.

6. Buffered Solutions Key Points on Buffered Solutions • 1. They are weak acids or bases containing a common ion. • After addition of strong acid or base, deal with stoichiometry first, then equilibrium.

7. Buffered Solutions • Addition of Strong Acids or Bases to Buffers • We break the calculation into two parts: stoichiometric and equilibrium. • The amount of strong acid or base added results in a neutralization reaction: • X- + H3O+ HX + H2O • HX + OH- X- + H2O. • By knowing how must H3O+ or OH- was added (stoichiometry) we know how much HX or X- is formed.

8. Buffered Solutions Addition of Strong Acids or Bases to Buffers

9. Buffered Solutions • Addition of Strong Acids or Bases to Buffers • With the concentrations of HX and X- (note the change in volume of solution) we can calculate the pH from the Henderson-Hasselbalch equation Henderson-Hasselbalch Equation: For a particular buffering system, all solutions that have the same ratio [A-]/[HA] will have the same pH.

10. Buffered Solutions • Buffer Capacity and pH – Henderson-Hasselbach Equation • If Ka is small (i.e., if the equilibrium concentration of undissociated acid is close to the initial concentration), then

11. Buffered Solutions • Buffer Capacity and pH • Buffer capacity is the amount of acid or base neutralized by the buffer before there is a significant change in pH. • Buffer capacity depends on the composition of the buffer. • The greater the amounts of conjugate acid-base pair, the greater the buffer capacity. • The pH of the buffer depends on Ka.

12. Acid-Base Titrations • Strong Acid-Base Titrations • Consider adding a strong base (e.g. NaOH) to a solution of a strong acid (e.g. HCl). • Before any base is added, the pH is controlled by the strong acid solution. Therefore, pH < 7. • When base is added, before the equivalence point, the pH is controlled by the amount of strong acid left in excess. Therefore, pH < 7. • At equivalence point, the amount of base added is stoichiometrically equivalent to the amount of acid originally present. Therefore, the pH is determined by the salt solution. Therefore, pH = 7.

13. Acid-Base Titrations • Strong Acid-Base Titrations • Problems only involve a stoichiometry calculation. • We know the pH at equivalent point is 7.00. • To detect the equivalent point, we use an indicator that changes color somewhere near 7.00. • The equivalence point in a titration is the point at which the acid and base are present in stoichiometric quantities. • The end point in a titration is the observed point. • The difference between equivalence point and end point is called the titration error.

14. Acid-Base Titrations Strong Acid-Base Titrations

15. Acid-Base Titrations • Weak Acid-Strong Base Titrations • Consider the titration of acetic acid, HC2H3O2 and NaOH. • Before any base is added, the solution contains only weak acid. Therefore, pH is given by the equilibrium calculation. • As strong base is added, the strong base consumes a stoichiometric quantity of weak acid: • HC2H3O2(aq) + NaOH(aq)  C2H3O2-(aq) + H2O(l)

16. Acid-Base Titrations Weak Acid-Strong Base Titrations

17. Acid-Base Titrations • Weak Acid-Strong Base Titrations • There is an excess of acetic acid before the equivalence point. • Therefore, we have a mixture of weak acid and its conjugate base. • The pH is given by the buffer calculation. • First the amount of C2H3O2- generated is calculated, as well as the amount of HC2H3O2consumed. (Stoichiometry.) • Then the pH is calculated using equilibrium conditions. (Henderson-Hasselbalch.)

18. Acid-Base Titrations • Weak Acid-Strong Base Titrations • At the equivalence point, all the acetic acid has been consumed and all the NaOH has been consumed. However, C2H3O2- has been generated. • Therefore, the pH is given by the C2H3O2- solution. • This means pH > 7. • More importantly, pH  7 for a weak acid-strong base titration. • After the equivalence point, the pH is given by the strong base in excess.

19. Acid-Base Titrations • Weak Acid-Strong Base Titrations • The inflection point is not as steep for a weak acid-strong base titration. • The shape of the two curves after equivalence point is the same because pH is determined by the strong base in excess. • Two features of titration curves are affected by the strength of the acid: • the amount of the initial rise in pH, and • the length of the inflection point at equivalence.

20. Acid-Base Titrations Weak Acid-Strong Base Titrations • The weaker the acid, the smaller the equivalence point inflection. • For very weak acids, it is impossible to detect the equivalence point.

21. Acid-Base Titrations • Weak Acid-Strong Base Titrations • Titration of weak bases with strong acids have similar features to weak acid-strong base titrations. • Weak Acid-Strong Base Titrations: A Summary • Step 1 - A stoichiometry problem- reaction is assumed to run to completion - then determine • remaining species. • Step 2 - An equilibrium problem- determine position of weak acid equilibrium and calculate pH.

22. Acid-Base Titrations • Titrations of Polyprotic Acids • In polyprotic acids, each ionizable proton dissociates in steps. • Therefore, in a titration there are n equivalence points corresponding to each ionizable proton. • In the titration of Na2CO3 with HCl there are two equivalence points: • one for the formation of HCO3- • one for the formation of H2CO3.

23. Indicator Color Change • Another way determine the equivalence point of an acid-base titration is through the use of an acid-base indicator. • Careful selection of the indicator will ensure that the end point is close to the equivalence point. • Most common acid-base indicators are complex molecules that are themselves weak acids (HIn). They exhibit one color when the proton is attached and a different color when the proton is absent. • Page 749 contains a list of acid-base indicators. • Henderson-Hasselbach equation is very useful in determining the pH at which an indicator changes color.

24. Solubility Equilibria • Solubility-Product Constant, Ksp • Consider • for which • Ksp is the solubility product. (BaSO4 is ignored because it is a pure solid so its concentration is constant.)

25. Solubility Equilibria • Solubility-Product Constant, Ksp • In general: the solubility product is the molar concentration of ions raised to their stoichiometric powers. • Solubility is the amount (grams) of substance that dissolves to form a saturated solution. • Molar solubility is the number of moles of solute dissolving to form a liter of saturated solution.

26. Solubility Equilibria • Solubility and Ksp • To convert solubility to Ksp • solubility needs to be converted into molar solubility (via molar mass); • molar solubility is converted into the molar concentration of ions at equilibrium (equilibrium calculation), • Ksp is the product of equilibrium concentration of ions.

27. Solubility Equilibria Solubility and Ksp

28. Solubility Equilibria Solubility and Ksp Exercise 15.12 Page 752 Copper (I) bromide has a measured solubility of 2.0x10-4 mol/L at 25C. Calculate its Ksp value.

29. Solubility Equilibria Solubility and Ksp Exercise 15.13 Page 754 Calculate the Ksp value for bismuth sulfide (Bi2S3), which has a solubility of 1.0x10-15mol/L at 25C.

30. Solubility Equilibria Solubility and Ksp Exercise 15.14 Page 755 The Ksp value for copper (II) iodate is 1.4 x 10-7 at 25º C. Calculate its solubility.

31. Solubility Equilibria Solubility and Ksp A potassium chromate solution being added to aqueous silver nitrate, forming silver chromate.

32. Factors That Affect Solubility • Common-Ion Effect • Solubility is decreased when a common ion is added. • This is an application of Le Châtelier’s principle: • as F- (from NaF, say) is added, the equilibrium shifts away from the increase. • Therefore, CaF2(s) is formed and precipitation occurs. • As NaF is added to the system, the solubility of CaF2 decreases.

33. Factors That Affect Solubility Common-Ion Effect

34. Factors That Affect Solubility Common-Ion Effect Exercise 15.15 Page 758 Calculate the solubility of solid CaF2 (Ksp = 4.0x10-11) in a 0.025 M NaF solution.

35. Factors That Affect Solubility • Solubility and pH • Again we apply Le Châtelier’s principle: • If the F- is removed, then the equilibrium shifts towards the decrease and CaF2 dissolves. • F- can be removed by adding a strong acid: • As pH decreases, [H+] increases and solubility increases. • The effect of pH on solubility is dramatic.