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Acid-Base Equilibria and Solubility Equilibria

Acid-Base Equilibria and Solubility Equilibria. Chapter 18. In this chapter we will consider both homogeneous solution equilibria and heterogeneous solution equilibria. Homogeneous solution equilibrium: Heterogeneous solution equilibria: .

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Acid-Base Equilibria and Solubility Equilibria

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  1. Acid-Base Equilibria andSolubility Equilibria Chapter 18

  2. In this chapter we will consider both homogeneous solution equilibria and heterogeneous solution equilibria. Homogeneous solution equilibrium: Heterogeneous solution equilibria:

  3. Consider an equimolar mixture of CH3COOH and CH3COO-Na+ H+(aq) + CH3COO-(aq) CH3COOH (aq) OH-(aq) + CH3COOH (aq) CH3COO-(aq) + H2O (l) • A ______________ solution is a solution of: • A weak acid or a weak base and • The salt of the weak acid or weak base • Both must be present! A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base. Add strong acid Add strong base

  4. Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr (c) Na2CO3/NaHCO3

  5. Maintaining the pH of Blood

  6. Suppose you want to prepare a buffer solution with a specific pH.How would you do that?You’d use the Henderson-Hasselbalch equation!

  7. NaA (s) Na+(aq) + A-(aq) HA (aq) H+(aq) + A-(aq) [H+][A-] Ka = [H+] = [HA] -log [H+] = -log Ka - log [HA] [A-] -log [H+] = -log Ka + log [HA] [A-] pH = pKa + log pH = ___ + log Consider mixture of a weak acid HA and its salt Na+A-. Henderson-Hasselbalch equation Since pKa = -log Ka Or, more generically…

  8. [conjugate base] [acid] We use the Henderson-Hasselbalch equation to prepare a buffer solution with ______________. We often choose a weak acid whose pKa is close to the desired pH and make the [acid]  [conjugate base]. (Why?) pH = pKa + log Because then, log = ____

  9. CH3COO-Na+(s) Na+(aq) + CH3COO-(aq) CH3COOH (aq) H+(aq) + CH3COO-(aq) The ________________________ is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The presence of a common ion __________the ionization of a weak acid or a weak base. Consider a mixture of CH3COO-Na+ (strong electrolyte) and CH3COOH (weak acid).

  10. HCOOH (aq) H+(aq) + HCOO-(aq) Initial (M) Change (M) Equilibrium (M) [HCOO-] pH = pKa + log [HCOOH] What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOO-K+? Mixture of weak acid and conjugate base! 0.30 0.00 0.52 Common ion effect 0.30 – x 0.30 0.52 + x 0.52 pH = HCOOH pKa = _____

  11. NH4+(aq) H+(aq) + NH3(aq) pH = pKa + log [0.30] [NH3] [NH4+] [0.36] Calculate the pH of a 0.30M NH3/0.36M NH4Cl buffer system Ka= 5.6 X 10-10 pKa= _________ pH = ______ + log = ________

  12. 0.050 mol NaOH X 0.020 L = 0.0010 mol OH- L [0.25] 0.025 0.028 [0.28] 0.10 0.10 NH4+ (aq) + OH- (aq) H2O (l) + NH3 (aq) pH = ____ + log [NH4+] = [NH3] = What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the 0.30 M NH3/0.36 M NH4Cl buffer solution? 0.024 start (moles) 0.029 0.001 end (moles) 0.028 0.0 0.025 final volume = 80.0 mL + 20.0 mL = 100 mL = ________

  13. In a ____________, a solution of accurately known concentration is added gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete. ___________ point – the point at which the reaction is complete _______________ – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL… The indicator changes color (__________)

  14. NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq) OH-(aq) + H+(aq) H2O (l) 0.10 M NaOH added to 25 mL of 0.10 M HCl Strong Acid-Strong Base Titrations At equivalence point (pH _______)

  15. CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l) CH3COOH (aq) + OH-(aq) CH3COO-(aq) + H2O (l) CH3COO-(aq) + H2O (l) OH-(aq) + CH3COOH (aq) Weak Acid-Strong Base Titrations At equivalence point (pH _______)

  16. HCl (aq) + NH3(aq) NH4Cl (aq) H+(aq) + NH3(aq) NH4Cl (aq) NH4+(aq) + H2O (l) NH3(aq) + H+(aq) Strong Acid-Weak Base Titrations At equivalence point (pH ______)

  17. HNO2(aq) + OH-(aq) NO2-(aq) + H2O (l) Initial (M) 0.01 Change (M) [NO2-] = = ______ M NO2-(aq) + H2O (l) OH-(aq) + HNO2(aq) 0.200 Equilibrium (M) x2 [OH-][HNO2] = Kb = 0.05-x [NO2-] Exactly 100. mL of 0.100 M HNO2 is titrated with a 0.100 M NaOH solution. What is the pH at the equivalence point ? start (moles) 0.01 0.01 end (moles) 0.0 0.0 0.01 Final volume = _______ mL 0.05 0.00 0.00 -x +x +x 0.05 - x x x pOH = _______ = 2.2 x 10-11 pH = 14 – pOH = ____ 0.05 – x 0.05 x 1.05 x 10-6 = [OH-]

  18. Acid-Base Indicators HIn (aq) H+(aq) + In-(aq) [HIn] [HIn] [HIn]  10  10  1 [In-] [In-] [In-] Color of ________ (HIn) predominates If [HIn]  [In-], combination color (p.669) Color of ______________ (In-) predominates Some Common Acid-Base Indicators

  19. The titration curve of a strong acid with a strong base.

  20. Which indicator(s) would you use for the titration of HNO2 versus KOH ? Weak acid titrated with strong base. At equivalence point, will have conjugate base of weak acid. At equivalence point, pH > 7 Use _________ _____ or _________________ Some Common Acid-Base Indicators

  21. Solubility Equilibria AgCl (s) Ag+(aq) + Cl-(aq) MgF2(s) Mg2+(aq) + 2F-(aq) Ag2CO3(s) 2Ag+(aq) + CO32-(aq) Ca3(PO4)2(s) 3Ca2+(aq) + 2PO43-(aq) Ksp= [Ag+][Cl-] Ksp is the ________________________ Ksp= [Mg2+][F-]2 Ksp= [Ag+]2[CO32-] Ksp= [Ca2+]3[PO33-]2 Dissolution of an ionic solid in aqueous solution: _________________ ____________ solution Q < Ksp Q = Ksp ____________ solution _________________ Q > Ksp ____________ solution

  22. Solubility Products of Some Slightly Soluble Ionic Compounds at 25ºC

  23. _________ ___________(mol/L) is the number ofmoles of solute dissolved in 1 L of a saturated solution. _______________(g/L)is the number ofgrams of solute dissolved in 1 L of a saturated solution.

  24. AgCl (s) Ag+(aq) + Cl-(aq) Initial (M) Change (M) Equilibrium (M)  s = Ksp mol AgCl 1 L soln x g AgCl 1 mol AgCl What is the solubility of silver chloride in g/L ? Ksp= 1.6 x 10-10 0.00 0.00 Ksp= [Ag+][Cl-] +s +s Ksp= s2 s s s = _________ [Cl-] = _________ M [Ag+] = _________ M = ______ g/L Solubility of AgCl =

  25. Relationship between Ksp and Molar Solubility (s)

  26. Solubility of Common Ionic Compounds in water at 250C

  27. 2 Q = [Ca2+]0[OH-]0 If 2.00 mL of 0.200 M NaOH are added to 1.00 L of 0.100 M CaCl2, will a precipitate form? The ions present in solution are Na+, OH-, Ca2+, Cl-. Only possible precipitate is Ca(OH)2 (solubility rules p.95). Is Q > Kspfor Ca(OH)2? [OH-]0 = 4.0 x 10-4M [Ca2+]0 = 0.100 M = 0.10 x (4.0 x 10-4)2 = 1.6 x 10-8 Ksp = [Ca2+][OH-]2 = 8.0 x 10-6 Q < Ksp No precipitate will form

  28. The _____________________________ is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The addition of a common ion _______________the solubility of a salt.

  29. AgBr (s) Ag+(aq) + Br-(aq) • What is the molar solubility of AgBr in • pure water and • (b) 0.0010 M NaBr? Ksp = 7.7 x 10-13 s2 = Ksp s = 8.8 x 10-7 = [Ag+] = [AgBr]

  30. AgBr (s) Ag+(aq) + Br-(aq) NaBr (s) Na+ (aq) + Br-(aq) • What is the molar solubility of AgBr in • pure water and • (b) 0.0010 M NaBr? [Br-] = 0.0010 M [Ag+] = s [Br-] = 0.0010 + s 0.0010 Ksp = 0.0010 x s = 7.7 x 10-13 s = 7.7 x 10-10 = [Ag+] = [AgBr] So adding 0.0010 M NaBr to the solution lowers the solubility of AgBr from 8.8 x 10-7 to 7.7 x 10-10

  31. OH-(aq) + H+(aq) H2O (l) Mg(OH)2(s) Mg2+(aq) + 2OH-(aq) pH and Solubility • The presence of a common ion _________solubility. • Insoluble bases dissolve in ____________ solutions • Insoluble acids dissolve in ____________ solutions What happens if you add H+ to Mg(OH)2 (aq)? Ksp= [Mg2+][OH-]2 = 1.2 x 10-11 Ksp = (s)(2s)2 = 4s3 At pH < 10.45 4s3 = 1.2 x 10-11 Lower [OH-] s = 1.4 x 10-4M [OH-] = 2s = 2.8 x 10-4M Increase solubility of Mg(OH)2 pOH = 3.55 pH = 10.45

  32. Mg(OH)2(s) Mg2+(aq) + 2OH-(aq) pH and Solubility • The presence of a common ion decreasessolubility. • Insoluble bases dissolve in acidic solutions • Insoluble acids dissolve in basic solutions What happens if you add OH- to Mg(OH)2 (aq)? Ksp= [Mg2+][OH-]2 = 1.2 x 10-11 Ksp = (s)(2s)2 = 4s3 4s3 = 1.2 x 10-11 At pH > 10.45 s = 1.4 x 10-4M Raise [OH-] [OH-] = 2s = 2.8 x 10-4M Decrease solubility of Mg(OH)2 pOH = 3.55 pH = 10.45

  33. 2- Co2+(aq) + 4Cl-(aq) CoCl4(aq) 2- stability of complex Kf 2+ Co(H2O)6 2- CoCl4 [CoCl4 ] Kf = [Co2+][Cl-]4 Complex Ion Equilibria and Solubility A _________ ______ is an ion containing a central metal cation bonded to one or more molecules or ions. The ________ ________ or ________ ________ (Kf) is the equilibrium constant for the complex ion formation.

  34. Formation Constants of Selected Complex Ions in Water at 25ºC

  35. Separation of Cations into Groups According to Their Precipitation Reactions with Various Reagents

  36. Qualitative Analysis of Cations And flame tests for cations

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