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Acid-Base Equilibria and Solubility Equilibria Chapter 16 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Presentation on theme: "Acid-Base Equilibria and Solubility Equilibria Chapter 16 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display."— Presentation transcript:

1 Acid-Base Equilibria and Solubility Equilibria Chapter 16 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 Common Ion Effect The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction. AgCl(s) Ag + (aq) + Cl (aq) Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2

3 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 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 H + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) Add strong base OH - (aq) + CH 3 COOH (aq) CH 3 COO - (aq) + H 2 O (l) Consider an equal molar mixture of CH 3 COOH and CH 3 COONa

4 A Buffered Solution... resists change in its pH when either H + or OH are added. 1.0 L of 0.50 M H 3 CCOOH M H 3 CCOONa pH = 4.74 Adding mol solid NaOH raises the pH of the solution to 4.76, a very minor change. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4

5 Which of the following are buffer systems? (a) KF/HF (b) KBr/HBr, (c) Na 2 CO 3 /NaHCO 3 (a) KF is a weak acid and F - is its conjugate base buffer solution (b) HBr is a strong acid not a buffer solution (c) CO 3 2- is a weak base and HCO 3 - is its conjugate acid buffer solution 16.3

6 Henderson-Hasselbalch Equation Useful for calculating pH when the [A ]/[HA] ratios are known. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6

7 What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOONa? Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7

8 Buffered Solution Characteristics Buffers contain relatively large amounts of weak acid and corresponding base. Added H + reacts to completion with the weak base. Added OH reacts to completion with the weak acid. The pH is determined by the ratio of the concentrations of the weak acid and weak base. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8

9 Calculate the pH of the 0.30 M NH 3 /0.36 M NH 4 Cl buffer system. What is the pH after the addition of 20.0 mL of M NaOH to 80.0 mL of the buffer solution? Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9

10 Buffering Capacity... represents the amount of H + or OH the buffer can absorb without a significant change in pH. Buffer solution 1 0.1M HNO 2 /0.1M KNO 2 Buffer solution 2 1.0M HNO 2 /1.0M KNO 2 Both solutions buffer to the same pH Buffer solution 2 has a larger capacity Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10

11 Titrations 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. Equivalence point – the point at which the reaction is complete (stoichiometry) Indicator – substance that changes color at (or near) the equivalence point Endpoint – point where indicator changes color Slowly add base to unknown acid UNTIL The indicator changes color (pink) 4.7

12 Titration (pH) Curve A plot of pH of the solution being analyzed as a function of the amount of titrant added. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12

13 Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) 16.4

14 Weak Acid-Strong Base Titrations CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq) CH 3 COO - (aq) + H 2 O (l) CH 3 COO - (aq) + H 2 O (l) OH - (aq) + CH 3 COOH (aq) At equivalence point (pH > 7): 16.4

15 pH = pKa at halfway point Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Acetic Acid K a =1.8x10 -5 pK a =4.74

16 Strong Acid-Weak Base Titrations HCl (aq) + NH 3 (aq) NH 4 Cl (aq) NH 4 + (aq) + H 2 O (l) NH 3 (aq) + H + (aq) At equivalence point (pH < 7): 16.4 H + (aq) + NH 3 (aq) NH 4 Cl (aq)

17 pH = pKa at halfway point Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Ammonia K b =1.8x10 -5 K a =5.56x pK a =9.25

18 Strong Acid – Strong Base Titration Step 1 -A stoichiometry problem - reaction is assumed to run to completion - then determine remaining species. Step 2 -Determine pH based on concentration of excess reactant (strong acid or base leftover) Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18

19 Consider the titration of 80.0mL of 0.10 M Sr(OH) 2 by 0.4M HCl. Calculate the pH after the following volumes of HCl have been added. a.0.0 mL b.20.0 mL c.30.0 mL d.40.0 mL e.80.0 mL Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19

20 Weak Acid - Strong Base Titration Weak Base – Strong Acid Titration 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 or base equilibrium and calculate pH. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20

21 Consider the titration of 100.0mL of 0.10 M HC 2 H 3 O 2 by 0.2M KOH. Calculate the pH after the following volumes of HCl have been added. a.0.0mL b.20.0 mL c.25.0 mL d.40.0 mL e.50.0 mL f mL Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21

22 Acid-Base Indicator... marks the end point of a titration by changing color. commonly weak acids (HIn) that change color when they dissociate (In - ) Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22

23 Acid-Base Indicators HIn (aq) H + (aq) + In - (aq) 10 [HIn] [In - ] Color of acid (HIn) predominates 10 [HIn] [In - ] Color of conjugate base (In - ) predominates 16.5

24 The titration curve of a strong acid with a strong base. 16.5

25 Which indicator(s) would you use for a titration of HNO 2 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 16.5

26 Solubility Product For solids dissolving to form aqueous solutions. K sp = solubility product constant Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26

27 Solubility Equilibria 16.6 AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ]K sp is the solubility product constant MgF 2 (s) Mg 2+ (aq) + 2F - (aq) K sp = [Mg 2+ ][F - ] 2 Ag 2 CO 3 (s) 2Ag + (aq) + CO (aq) K sp = [Ag + ] 2 [CO ] Ca 3 (PO 4 ) 2 (s) 3Ca 2+ (aq) + 2PO (aq) K sp = [Ca 2+ ] 3 [PO ] 2 Dissolution of an ionic solid in aqueous solution: Q = K sp Saturated solution Q < K sp Unsaturated solution No precipitate Q > K sp Supersaturated solution Precipitate will form

28 Solubility Product Bi 2 S 3 (s) 2Bi 3+ (aq) + 3S 2 (aq) Solubility = s = concentration of Bi 2 S 3 that dissolves, which equals 1/2[Bi 3+ ] and 1/3[S 2 ]. Note:K sp is constant (at a given temperature) s is variable (especially with a common ion present) Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28

29 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. 16.6

30 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 What is the solubility of silver chloride in g/L ? (AgCl K sp = 1.6 x ) What concentration of Ag is required to precipitate ONLY AgBr in a solution that contains both Br - and Cl - at a concentration of 0.02 M? (AgBr K sp = 7.7 x )

31 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) M NaBr? 16.8

32 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) Mg 2+ (aq) + 2OH - (aq) K sp = [Mg 2+ ][OH - ] 2 = 1.2 x K sp = (s)(2s) 2 = 4s 3 4s 3 = 1.2 x s = 1.4 x M [OH - ] = 2s = 2.8 x M pOH = 3.55 pH = At pH less than (Lower [OH - ]) Increase solubility of Mg(OH) 2 At pH greater than 10.45(Raise [OH - ]) Decrease solubility of Mg(OH)

33 Complex Ion Equilibria and Solubility A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions. Co 2+ (aq) + 4Cl - (aq) CoCl 4 (aq) 2- K f = [CoCl 4 ] [Co 2+ ][Cl - ] 4 2- The formation constant or stability constant (K f ) is the equilibrium constant for the complex ion formation. Co(H 2 O) 6 2+ CoCl KfKf stability of complex

34 16.10

35 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 If 2.5g of CuSO 4 are dissolved in 900mL of.30M NH 3, what are the concentrations of Cu 2+, Cu(NH 3 ) 4 2+, and NH 3 at equilibrium? (Cu(NH 3 ) 4 2+ K f =5.0x10 13 ) Calculate the molar solubility of AgCl in 1.0M NH 3. (Ag(NH 3 ) 2 + K f =1.5x10 7 AgCl K sp =1.6x )

36 16.11

37 Qualitative Analysis of Cations 16.11


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