Acid-Base Equilibria and Solubility Equilibria
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1 Acid-Base Equilibria and Solubility Equilibria 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. Chapter 16 The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider miture of CH3COONa (strong electrolyte) and CH3COOH (weak acid). CH3COONa (s) Na+ (aq) + CH3COO- (aq) CH3COOH (aq) H+ (aq) + CH3COO- (aq) common ion Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Consider miture of salt NaA and weak acid HA. NaA (s) Na+ (aq) + A- (aq) HA (aq) H+ (aq) + A- (aq) [H+] = Miture of weak acid and conjugate base! HCOOH (aq) Ka -log [H+] = -log Ka + log [conjugate base] [acid] Henderson-Hasselbalch equation -log [H+] = -log Ka - log What is the ph of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? [H+] Ka = Common ion effect pka = -log Ka H+ (aq) + HCOO- (aq) 0.52 [HCOO-] [HCOOH] ph = log [0.52] = 4.01 [0.30] HCOOH pka = 3.77 Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr, (c) Na2CO3/NaHCO3 A buffer solution is a solution of: 1. A weak acid or a weak base and 2. The salt of the weak acid or weak base (a) KF is a weak acid and F- is its conjugate base buffer solution 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. (b) HBr is a strong acid not a buffer solution Consider an equal molar miture of CH3COOH and CH3COONa Add strong acid H+ (aq) + CH3COO- (aq) Add strong base OH- (aq) + CH3COOH (aq) (c) CO32- is a weak base and HCO3- is it conjugate acid buffer solution CH3COOH (aq) 1
2 Calculate the ph of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the ph after the addition of 2 ml of 50 M NaOH to 8 ml of the buffer solution? NH4+ (aq) start (moles) end (moles) [NH3] [NH4+] Maintaining the ph of Blood H+ (aq) + NH3 (aq) pka = 9.25 ph = log NH4+ (aq) + OH- (aq) 28 [0.30] = 9.17 [0.36] 24 H2O (l) + NH3 (aq) 25 final volume = 8 ml + 2 ml = 100 ml [NH4+] = [NH3] = ph = log [0.25] = 9.20 [0.28] Strong Acid-Strong Base Titrations Titrations NaOH (aq) + HCl (aq) In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. OH- (aq) + H+ (aq) H2O (l) + NaCl (aq) H2O (l) 0.10 M NaOH added to 25 ml of 0.10 M HCl Equivalence point the point at which the reaction is complete Indicator substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL The indicator changes color (pink) 4.7 Weak Acid-Strong Base Titrations CH3COOH (aq) + NaOH (aq) Strong Acid-Weak Base Titrations CH3COONa (aq) + H2O (l) CH3COOH (aq) + OH- (aq) HCl (aq) + NH3 (aq) H+ (aq) + NH3 (aq) NH4Cl (aq) NH4Cl (aq) At equivalence point (ph > 7): At equivalence point (ph < 7): NH4+ (aq) + H2O (l) OH- (aq) + CH3COOH (aq) NH3 (aq) + H+ (aq) 2
3 Eactly 100 ml of 0.10 M HNO2 are titrated with a 0.10 M NaOH solution. What is the ph at the equivalence point? start (moles) end (moles) HNO2 (aq) + OH- (aq) NO2- (aq) + H2O (l) HIn (aq) (aq) + In- (aq) [HIn] 10 Color of acid (HIn) predominates [In-] Final volume = 200 ml [NO2-] = = 5 M NO2 (aq) + H2O (l) OH (aq) + HNO2 (aq) Acid-Base Indicators H+ [OH-][HNO2] 2 = Kb = = [NO2-] 5- poh = = [OH-] ph = 14 poh = 8.02 The titration curve of a strong acid with a strong base. [HIn] 10 Color of conjugate base (In-) predominates [In-] Which indicator(s) would you use for a titration of HNO2 with KOH? Weak acid titrated with strong base. At equivalence point, will have conjugate base of weak acid. At equivalence point, ph > 7 Use cresol red or phenolphthalein Solubility Equilibria AgCl (s) Ksp = [Ag+][Cl-] MgF2 (s) Ag2CO3 (s) Ca3(PO4)2 (s) Ag+ (aq) + Cl- (aq) Ksp is the solubility product constant Mg2+ (aq) + 2F- (aq) Ksp = [Mg2+][F-]2 2Ag+ (aq) + CO32- (aq) Ksp = [Ag+]2[CO32-] 3Ca2+ (aq) + 2PO43- (aq) Ksp = [Ca2+]3[PO33-]2 Dissolution of an ionic solid in aqueous solution: Q < Ksp Unsaturated solution Q = Ksp Saturated solution Q > Ksp Supersaturated solution No precipitate Precipitate will form 3
4 Molar solubility (mol/l) is the number of moles of solute dissolved in 1 L of a saturated solution. Solubility (g/l) is the number of grams of solute dissolved in 1 L of a saturated solution. What is the solubility of silver chloride in g/l? AgCl (s) [Ag + ] = M Ag + (aq) + Cl - (aq) 0 0 +s s [Cl - ] = M Solubility of AgCl = mol AgCl 1 L soln +s s K sp = K sp = [Ag + ][Cl - ] K sp = s 2 s = K sp s = g AgCl = mol AgCl -3 g/l If 2.00 ml of M NaOH are added to 1.00 L of M CaCl 2, will a precipitate form? The ions present in solution are Na +, OH -, Ca 2+, Cl -. Only possible precipitate is Ca(OH) 2 (solubility rules). Is Q > K sp for Ca(OH) 2? [Ca 2+ ] 0 = M [OH - ] 0 = M Q = [Ca 2+ ] 0 [OH - 2 ] 0 = 0.10 ( ) 2 = K sp = [Ca 2+ ][OH - ] 2 = Q < K sp No precipitate will form AgBr (s) The Common Ion Effect and Solubility The presence of a common ion decreases the solubility of the salt. What is the molar solubility of AgBr in (a) pure water and (b) 010 M NaBr? K sp = s 2 = K sp s = Ag + (aq) + Br - (aq) NaBr (s) [Br - ] = 010 M AgBr (s) [Ag + ] = s Na + (aq) + Br - (aq) Ag + (aq) + Br - (aq) [Br - ] = s 010 K sp = 010 s s = ph and Solubility The presence of a common ion decreases the solubility. Insoluble bases dissolve in acidic solutions Insoluble acids dissolve in basic solutions Mg(OH) 2 (s) K sp = [Mg 2+ ][OH - ] 2 = K sp = (s)(2s) 2 = 4s 3 4s 3 = s = M [OH - ] = 2s = M poh = 3.55 ph = remove add Mg 2+ (aq) + 2OH - (aq) At ph less than Lower [OH - ] OH - (aq) + H + (aq) H 2 O (l) Increase solubility of Mg(OH) 2 At ph greater than Raise [OH - ] Decrease solubility of Mg(OH) 2 4
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