CHEM 163

Chapter 19

Spring 2009

1

BuffersSolution that resists pH changes

– Ex. Blood (pH ~ 7.4)

• Acid must neutralize small amounts of base• Base must neutralize small amounts of acid• Acid and base must not neutralize each other

2

Use conjugate acid-base pairs!

CH3COOH (aq) + H2O (l)

CH3COO - (aq) + H3O+ (aq)

Added in as salt (NaCH3COO)

Common-ion effectEx: acetate

• High concentrations of weak acid/conjugate base

• Add H3O+ or OH-

– Added amounts are relatively small– Cause only small shifts– React with weak acid or conjugate base

3

HA (aq) + H2O (l) A - (aq) + H3O+ (aq)

[HA]

]O][H[AK 3

-

a

][A

[HA]K]O[H

-a3

pH depends on [HA]/[A-] ratio

HA (aq) + OH- (aq) A - (aq) + H2O (l)

Making a buffer

1. Choose the conjugate acid-base pair (pKa ≈ pH)

2. Calculate the ratio of acid-base concentrations3. Determine the buffer concentration4. Mix solution; adjust pH 4

[acid]

[base]logpKpH aHenderson-Hasselbalch

equation:

][A

[HA]K]O[H

-a3

][A

[HA]logKlog]Olog[H-

-a3

Buffer Properties• Buffer Capacity:

– Ability to resist pH change– Unrelated to pH of buffer– Dependent on concentration of weak acid/conj

base– Highest when [A-] = [HA]

• Buffer Range:– pH range over which buffer is effective

– Usually within ±1 pH unit of the pKa of weak acid

5

Sample Problem

Make 200. mL of a pH 3.5 citric acid/sodium citrate buffer with an acid concentration of 0.50 M.

We are given solid sodium citrate (294 g/mol) and 5.0 M citric acid. The pKa of citric acid is 3.15.

6

Measuring pH• Acid-Base Titration Curves: pH v. volume

titrant

Measuring pH:1. pH meter

2. Acid-base indicators

Indicator:• Weak organic acid• HIn different color than In-

• Intensely colored (small amount needed)• Changes color over ~ 2 pH units

7

Titration Curves: Strong acid – Strong base

• Low pH (strong acid)• Sudden pH rise (6-8

units)• Slow pH increase

8

[OH-]added ≈ [H3O+]init

Equivalence point: [OH-]added = [H3O+]init

pH = 7

End point: when indicator changes

color

Calculating pH during titration• Original solution of strong HA

• Before the equivalence point– Moles of acid remaining?

– Calculate [H3O+]

• At the equivalence point: pH = 7

• After the equivalence point– Excess moles of OH- added?– Calculate [OH-]

9

]OHlog[ 3pH

bbaaa VV MMn moles acid initial moles acid rxted

added basea

a3OH

VV

n

pH

aabbb VV MMn

moles base added

moles acid total

ba

bOHVV

n

pH

Titration Curves:Weak acid – Strong base

• Higher initial pH (weak acid, lower Ka)

• Buffer region– gradual pH rise– Midpoint:

½ initial HA reacted

• Equivalence point: pH > 7.00

• Slow pH increase 10

[HA] = [A-] pH = pKa

Calculating pH during titration• Original solution of weak HA

– ICE table

• Buffer Region

• At the equivalence point:

• After the equivalence point

11

ba

bOHVV

n

pHExcess moles of OH-

added

x = [H3O+] initaK HA

[acid]

[base]logpKpH a

total

acid-

VA

n

AOH3

b

w

K

K

Titration Curves:Strong acid – Weak base

• Initial pH > 7.00 (weak base)• Buffer region

– gradual pH decrease

• Equivalence point: pH < 7.00

• Slow pH decrease

12

Less common than strong base-weak acid(fewer appropriate indicators)

Titration Curves:Polyprotic Acids

13

Salts

• “slightly soluble”• Equilibrium between solid and dissolved ions

PbSO4 (s)

Pb2+ (aq) + SO42- (aq)

H2O

NaCl (s) Na+ (aq) + Cl- (aq) H2O• soluble

spQ ]][SO[Pb 24

2

Solubility productIon-product

expressionspK ]][SO[Pb 2

42

(at saturation)Solubility-Product

Constant

larger Ksp: more dissolution at equil. (saturation)

Smaller Ksp: less dissolution at equil. (saturation)

Insoluble Metal Sulfides

MnS (s) Mn2+ (aq) + S2- (aq) H2O

S2- (aq) + H2O (l) HS- (aq) + OH- (aq) MnS (s) Mn2+

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

(aq)

spK ]][OH][HS[Mn -2

3-minute Practice

Write Ksp expression for each of the following:

Silver bromide in H2O

Silver sulfide in H2O

AgBr (s) Ag+ (aq) + Br - (aq) H2O

spK ]][Br[Ag

+ H2O (l)

Ag2S (s) 2Ag+ (aq) +

HS- (aq) + OH- (aq)

spK ]][OH[HS][Ag 2

Higher Ksp = greater solubility?

Yes, for compounds with same total number of ions

Compound

Ksp Solubility

Ca(OH)2 6.5 x 10-6 1.2 x 10-2

PbSO4 1.6 x 10-8 1.3 x 10-4

MgCO3 3.5 x 10-8 1.9 x 10-4

BaF2 3.2 x 10-11 7.2 x 10-3

2-minute Practice

What else affects solubility?• Presence of a common ion:

PbSO4 (s) Pb2+ (aq) + SO42- (aq)

H2O

Add Na2SO4?

Decreases solubility

• pH:

CaCO3 (s) Ca2+ (aq) + CO32- (aq)

H2O

↑ [H3O+] ↑ solubility

if compound contains anion of weak acid

CO32- (aq) + H3O+

(aq) H2O (l) + HCO3

- (aq)

Homework problems

Chap 19: #9, 13, 19, 29, 50, 63, 70, 76, 78, 90

Due Tuesday, 4/28

More lecture notes will be added next week! Stay tuned.

19

Precipitation

Will it occur?• Qsp = Ksp:

• Qsp > Ksp:

• Qsp < Ksp:

• Selective precipitation– Way to separate ions– Form slightly soluble compounds with different

Ksp

Saturated solutionPrecipitation occursUnsaturated solution

Selective Precipitation

21

Mix 0.2 M Zn(NO3)2 and 0.4 M Mn(NO3)2.

Precipitate?

Add NaOH… Zn(OH)2 and Mn(OH)2

Ksp Zn(OH)2 = 3.0 x 10-16 Ksp Mn(OH)2 = 1.6 x 10-13

3-minute Practice

Which product is more soluble?

What [OH-] would need to make a saturated solution of the more soluble product? Hint: use Ksp expression!

Products?

Complex Ions• Central metal ion + ligands

Ionic ligands:OH-, CN-, halides

Molecular ligands:H2O, NH3

Lewis acid Lewis base

M(H2O)42+ (aq) + 4 NH3

(aq)M(NH3)4

2+ (aq) + 4 H2O (l)

Formation constant:

fK 43

242

243

]][)([

])([

NHOHM

NHM

Effects of ligands

A slightly soluble compound becomes more soluble when its cation forms a complex ion.

AgBr (s) Ag+ (aq) + Br-

(aq)Add Na2S2O3:Ag+ (aq) + S2O3

2- (aq)

Ag(S2O3)23-

(aq)2

Amphoteric Hydroxides:• Very slightly soluble in water• More soluble in acidic or basic solutions Al(OH)3

(s)Al3+ (aq) + 6 H2O

(l) +

3H3O+Al(H2O)6 (s)+ 4 OH- Al(H2O)2(OH)4

- (aq)+ 4 H2O (l)