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

Acid-Base Equilibria andSolubility Equilibria

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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

adding NaCl( ) shifts equilibrium positionaq

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.

16.3

Add strong acid

H+ (aq) + CH3COO- (aq) CH3COOH (aq)

Add strong base

OH- (aq) + CH3COOH (aq) CH3COO- (aq) + H2O (l)

Consider an equal molar mixture of CH3COOH and CH3COONa

A Buffered Solution

•. . . resists change in its pH when either H+ or OH are added.

1.0 L of 0.50 M H3CCOOH

+ 0.50 M H3CCOONa

•pH = 4.74•Adding 0.010 mol solid NaOH raises the pH of the solution to 4.76, a very minor change.


4

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

(a) HF is a weak acid and F- is its conjugate basebuffer solution

(b) HBr is a strong acidnot a buffer solution

(c) CO32- is a weak base and HCO3

- is its conjugate acidbuffer solution

16.3

Henderson-Hasselbalch Equation

Useful for calculating pH when the [A]/[HA] ratios are known.


6

Acid

BasepK

HA

ApKpH aa loglog

What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOONa?


7

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.


8

Calculate the pH of the 0.30 M NH3/0.36 M NH4Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution?


9

Buffering Capacity• . . . represents the amount of H+ or OH

the buffer can absorb without a significant change in pH.

• Buffer solution 1 0.1M HNO2/0.1M KNO2

• Buffer solution 2 1.0M HNO2/1.0M KNO2

• Both solutions buffer to the same pH• Buffer solution 2 has a larger capacity


10

TitrationsIn 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 baseto unknown acid

UNTIL

The indicatorchanges color

(pink) 4.7

Titration (pH) Curve

• A plot of pH of the solution being analyzed as a function of the amount of titrant added.


12

Strong Acid-Strong Base Titrations

NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)

OH- (aq) + H+ (aq) H2O (l)

16.4

Weak Acid-Strong Base Titrations

CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)

CH3COOH (aq) + OH- (aq) CH3COO- (aq) + H2O (l)

CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)

At equivalence point (pH > 7):

16.4

pH = pKa at halfway point


15

Acetic Acid Ka=1.8x10-5 pKa=4.74

Strong Acid-Weak Base Titrations

HCl (aq) + NH3 (aq) NH4Cl (aq)

NH4+ (aq) + H2O (l) NH3 (aq) + H+ (aq)

At equivalence point (pH < 7):

16.4

H+ (aq) + NH3 (aq) NH4Cl (aq)

pH = pKa at halfway point


17

Ammonia Kb=1.8x10-5 Ka=5.56x10-10 pKa=9.25

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)


18

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 mLb. 20.0 mLc. 30.0 mLd. 40.0 mLe. 80.0 mL


19

Sr(OH)2 + 2HCl SrCl2 + 2H2O

Moles OH-

Moles H+ Moles of Excess

Total Volume (mL)

pH Calculations

0.016 0 0.016 mol OH- 80

0.016 0.008 0.008 mol OH- 100

0.016 0.012 0.004 mol OH- 110

0.016 0.016 0 mol (equivalence

point)

120

0.016 0.032 0.016 mol H+ 160

20

Weak Acid - Strong Base TitrationWeak 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.


21

Consider the titration of 100.0mL of 0.10 M HC2H3O2 by 0.2M KOH. Calculate the pH after the following volumes of KOH

have been added.

a. 0.0mLb. 20.0 mLc. 25.0 mLd. 40.0 mLe. 50.0 mLf. 100.0 mL


22

KOH + HC2H3O2 KC2H3O2 + H2OKOH + HC2H3O2 KC2H3O2 + H2O

Before 0.004 0.01 0 NA

Change -0.004 -0.004 +0.004

After 0 0.006 0.004


23


Before 0.005 0.01 0 NA

Change -0.005 -0.005 +0.005

After 0 0.005 0.005


24


Before 0.008 0.01 0 NA

Change -0.008 -0.008 +0.008

After 0 0.002 0.008


25


Before 0.01 0.01 0 NA

Change -0.01 -0.01 +0.01

After 0 0 0.01


26


Before 0.02 0.01 0 NA

Change -0.01 -0.01 +0.01

After 0.01 0 0.01


27

Acid-Base Indicator

•. . . marks the end point of a titration by changing color. •commonly weak acids (HIn) that change color when they dissociate (In-)


28

OHInOHHIn 32

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

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

16.5

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

16.5

Solubility Product

•For solids dissolving to form aqueous solutions.

•Ksp = solubility product constant


32

Solubility Equilibria

16.6

AgCl (s) Ag+ (aq) + Cl- (aq)

Ksp = [Ag+][Cl-] Ksp is the solubility product constant

MgF2 (s) Mg2+ (aq) + 2F- (aq) Ksp = [Mg2+][F-]2

Ag2CO3 (s) 2Ag+ (aq) + CO32- (aq) Ksp = [Ag+]2[CO3

2-]

Ca3(PO4)2 (s) 3Ca2+ (aq) + 2PO43- (aq) Ksp = [Ca2+]3[PO4

3-]2

Dissolution of an ionic solid in aqueous solution:

Q = Ksp Saturated solution

Q < Ksp Unsaturated solution No precipitate

Q > Ksp Supersaturated solution Precipitate will form

Solubility Product

•Bi2S3(s) 2Bi3+(aq) + 3S2(aq)

•“Solubility” = s = concentration of Bi2S3 that dissolves, which equals

1/2[Bi3+] and 1/3[S2].

• Note:Ksp is constant (at a given temperature)

•s is variable (especially with a common ion present)


34

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


36

What is the solubility of silver chloride in g/L ? (AgCl Ksp= 1.6 x 10-10)


37

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 Ksp= 7.7 x 10-13)

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) 0.0010 M NaBr?

16.8

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) Mg2+ (aq) + 2OH- (aq)

Ksp = [Mg2+][OH-]2 = 1.2 x 10-11

Ksp = (s)(2s)2 = 4s3

4s3 = 1.2 x 10-11

s = 1.4 x 10-4 M

[OH-] = 2s = 2.8 x 10-4 M

pOH = 3.55 pH = 10.45

At pH less than 10.45 (Lower [OH-])

Increase solubility of Mg(OH)2

At pH greater than 10.45 (Raise [OH-])

Decrease solubility of Mg(OH)2

16.9

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