University Chemistry Chapter 12: Acid-Base Equilibria and Solubility Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 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. The presence of a common ion suppresses the ionization of a weak acid or a weak base. Consider a mixed solution of CH 3 COONa (strong electrolyte) and CH 3 COOH (weak acid). CH 3 COONa (s) Na + (aq) + CH 3 COO - (aq) CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) common ion The Common Ion Effect

3 Consider a mixture of salt NaA and weak acid HA. HA (aq) H + (aq) + A - (aq) NaA (s) Na + (aq) + A - (aq) K a = [H + ][A - ] [HA] [H + ] = K a [HA] [A - ] -log [H + ] = -log K a - log [HA] [A - ] -log [H + ] = -log K a + log [A - ] [HA] Henderson-Hasselbalch equation

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7 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

8 Buffered SolutionUnbuffered Solution Bromophenol blue: blue-purple >pH 4.6 and yellow< pH 3.0. HCl added

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13 HCl H + + Cl - HCl + CH 3 COO - CH 3 COOH + Cl -

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16 Maintaining the pH of Blood CA: carbonic anhydrase, HHb: hemoglobin molecule HHbO 2 : Oxyhemoglobin pH 7.40 pH 7.25

17 Titrations In a titration a solution of accurately known concentration (standard solution) 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 acid has completely reacted with or been neutralized by the base. Indicator – substance that changes color at (or near) the equivalence point. Slowly add base to unknown acid UNTIL The indicator changes color (pink) End point – the point where the indicator changes color.

18 pH meter – alternate method to detect the equivalence point

19 Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) pH profile for the titration of a strong acid with a strong base

20 Consider the addition of a M NaOH solution (from a buret) to an Erlenmeyer flask containing 25.0 mL of M HCl. You can calculate the pH in the following cases:

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22 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) pH Profile for the Titration of a Weak Acid with a Strong Base. At equivalence point (pH > 7)

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28 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) H + (aq) + NH 3 (aq) NH 4 + (aq) pH Profile for the Titration of a Strong Acid with a Weak Base.

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32 Acid-Base Indicators HIn (aq) H + (aq) + In - (aq)  10 [HIn] [In - ] Color of acid (HIn) predominates  0.1 [HIn] [In - ] Color of conjugate base (In - ) predominates An indicator is a substance (usually a weak organic acid or base) that has distinctly different colors in its nonionized and ionized forms. The end point of a titration occurs when the indicator changes color. Phenolph- thalein H 3 In + H 2 In In 2− In(OH) 3− Structure + pH = pK a + log 10 [In - ] [HIn] Color changes around pKa of the indicator.

33 The titration curve of a strong acid with a strong base

34 pH Many indicators are plant pigments.

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39 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb NO K + + 2I - PbI 2 (s) + 2K + + 2NO 3 - K + and NO 3 - are spectator ions Pb(NO 3 ) 2 (aq) + 2KI (aq) PbI 2 (s) + 2KNO 3 (aq) precipitate Pb I - PbI 2 (s) PbI 2

40 Pb I - PbI 2 (s)

41 Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.

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44 Solubility Equilibria 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 Solubility product can be approximated as the product of the molar concentrations of the constituent ions, each raised to the power of its stoichiometric coefficient in the equilibrium equation. Experimental concentrations may differ from ideal, depending on concentration and reactions involving any of the ions.

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46 Possible relationships between Q and K sp Q = K sp Saturated solution Q < K sp Unsaturated solutionNo precipitate Q > K sp Supersaturated solution Precipitate will form Equilibrium Ion product (Q) Q has the same form as the K sp except that the concentrations of ions are not equilibrium concentrations.

47 Molar solubility (mol L -1 ) is the number of moles of solute dissolved in 1 L of a saturated solution. Solubility (g L -1 ) is the number of grams of solute dissolved in 1 L of a saturated solution. Relating solubility and the K sp

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53 Predicting Precipitation Reactions

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55 Separation of Ions by Fractional Precipitation

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57 The Solubility of a Substance Is Affected by a Number of Factors Common Ion Effect: Common ion Consider a solution containing both AgCl and AgNO 3 The solubility of AgCl is less in a solution of AgNO 3 than in water.

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60 pH: The solubilities of many substances are affected by hydronium and/or hydroxide ion concentration [OH - ] is a function of pH. Increasing pH will decrease the solubility. Milk of magnesia, which contains Mg(OH) 2, is used to treat acid indigestion.

61 [F - ] is a function of pH. Decreasing pH will increase the solubility. The solubilities of salts containing anions that do not hydrolyze are unaffected by pH. Examples of such anions are Cl -, Br -, and I -.

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66 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 4 2- KfKf stability of complex Complex Ion Equilibria Co 2+ is hydrated in solution as Co(H 2 O) 6.

67 Addition of concentrated NH 3 (aq) to CuSO 4 (aq) blue solution light blue precipitate dark blue complex ion Increasing [NH 3 ]

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74 NaCl(aq) added to AgNO 3 (aq)NH 3 (aq) added to AgCl(s) AgCl(s)Ag(NH 3 ) 2 + (aq)

75 Amphoteric hydroxides can react with both acids and bases. Lewis base Lewis acid Soluble in acid Soluble in base

76 Qualitative Analysis of Groups 1 – 5 Cations Qualitative analysis: the determination of the types of ions present in a solution.

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78 Flame Test for Cations lithium sodium potassiumcopper