Chapter 17

Sections

17.1, 17.2, 17.4,

17.5(Common Ion Effect)

The Common-Ion Effect

“The extent of ionization of a weak electrolyte is decreased by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte.”

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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)

adding NaCl( ) shifts equilibrium positionaq

See problems 17.13 – 17.18

Problem

Calculate the % Ionization and pH of

(a) a solution that is 0.260 M in formic acid (HCHO2) Ka=1.8x10-4

(b) a solution that is 0.160 M in sodium formate (NaCHO2) and 0.260 M in formic acid (HCHO2)

(c) a solution that is made by combining 125 mL of 0.050 M hydrofluoric acid (Ka=6.8x10-4)with 50.0 mL of 0.10 M sodium fluoride.

Problem

(a) Calculate the percent ionization of 0.085 M lactic acid (HC3H5O3) (Ka = 1.4 × 10−4).

(b) Calculate the percent ionization of 0.085 M lactic acid in a solution containing 0.050 M sodium lactate.

Buffers

Buffers are solutions of a weak conjugate acid–base pair.

They are particularly resistant to pH changes, even when strong acid or base is added.

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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.

Key Points on Buffered Solutions

1. They are weak acids or bases containing a common ion.

2. After addition of strong acid or base, deal with stoichiometry first, then equilibrium.

Henderson-Hasselbalch Equation

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

pH p log( A HA

p log( base acid

a

a

K

K

/ )

/ )

Problem

a. Calculate the pH of a buffer that is 0.225 M in NaHCO3 and 0.290 M in Na2CO3

Calculate problem both ways – equilibrium method

and using Henderson Hasselbalch equation

Carbonic Acid

(H2CO3 )

Ka Values

Ka1 4.5 x 10-7 

Ka2 4.7 x 10-11

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.

pH Range

The pH range is the range of pH values over which a buffer system works effectively.

It is best to choose an acid with a pKa close to the desired pH.

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Making a Buffer SolutionProblem

How many moles of NaBrO should be added to 1.00 L of 5.0 x 10-2 M hypobromous acid (HBrO) to form a buffer solution of pH = 9.14? Assume that no volume change occurs when the NaBrO is added.

See problems 17.21 – 17.26

Buffering Capacity

. . . represents the amount of H+ or OH the buffer can absorb without a significant change in pH.

Buffering Action

See problems 17.24 – 17.28

Buffering Action

See problems 17.27 – 17.30

Solubility Product

For solids dissolving to form aqueous solutions.

Bi2S3(s) 2Bi3+(aq) + 3S2(aq)

Ksp = solubility product constant

and Ksp = [Bi3+]2[S2]3

See problem 17.49

Solubility Product

“Solubility” = x = concentration of Bi2S3 that dissolves, which equals 1/2[Bi3+] and 1/3[S2].

Note:Ksp is constant (at a given temperature)

x is variable (especially with a common ion present)

See problem 17.50

Calculating Ksp

See problems 17.51 – 17.55

Factors Affecting Solubility

The Common-Ion Effect

– If one of the ions in a solution equilibrium is already dissolved in the solution, the equilibrium will shift to the left and the solubility of the salt will decrease:

BaSO4(s) Ba2+(aq) + SO42(aq)

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Factors Affecting SolubilitypH

– If a substance has a basic anion, it will be more soluble in an acidic solution.

– Substances with acidic cations are more soluble in basic solutions.

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Solubility of CaF2 vs. NaF

See problems 17.56 – 17.58

Solubility of CaF2 vs. pH

See problems 17.59 – 17.61

Titration (pH) Curve

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

Equivalence (stoichiometric) point: Enough titrant has been added to react exactly with the solution being analyzed.

15_327

01.0

Vol NaOH added (mL)

50.0

7.0

13.0pH

100.0

Equivalencepoint

Strong Acid – Strong Base Titration

15_328

Vol 1.0 M HCl added

7.0

14.0

50.0 mL

Equivalencepoint

pH

Strong Base – Strong Acid Titration

15_330

Vol NaOH

Strong acid

pH

Weak acid

Weak vs. Strong Acid Titration

15_331

Vol 0.10 M NaOH added (mL)

10 20 30 40 50 60

2.0

4.0

6.0

8.0

10.0

12.0

0

Ka = 10–2

Ka = 10–4

Ka = 10–6

Ka = 10–8

Ka = 10– 10

Strong acid

pH

Weak Acid - Strong Base 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 equilibrium and calculate pH.

Titration Curve pH Calculation

Titration

Acid-Base Indicator

. . . marks the end point of a titration by changing color.

The equivalence point is not necessarily the same as the end point.

15_333

– O

C

O

C O–

O

(Pink base form, In– )

HO

COH

C O–

O

(Colorless acid form, HIn)

OH

Phenolphthalein Indicator

15_3340 1 2 3 4 5 6 7 8 9 10 11 12 13

The pH ranges shown are approximate. Specific transition ranges depend on the indicator solvent chosen.

pH

Crystal Violet

Cresol Red

Thymol Blue

Erythrosin B

2,4-Dinitrophenol

Bromphenol Blue

Methyl Orange

Bromcresol Green

Methyl Red

Eriochrome* Black T

Bromcresol Purple

Alizarin

Bromthymol Blue

Phenol Red

m - Nitrophenol

o-Cresolphthalein

Phenolphthalein

Thymolphthalein

Alizarin Yellow R

* Trademark CIBA GEIGY CORP.

Other Indicators

15_335AB

pH

00

Vol 0.10 M NaOH added (mL)

2

4

6

8

10

12

14

20 40 60 80 100 120

Equivalencepoint pH

00

Vol 0.10 M NaOH added (mL)

2

4

6

8

10

12

14

20 40 60 80 100 120

Equivalencepoint

Phenolphthalein

Methyl red

Phenolphthalein

Methyl red

END

Problem

(a) If the molar solubility of CaF2 at 35°C is 1.24 × 10−3 mol/L, what is Ksp at this temperature?

(b) The Ksp of Ba(IO3)2 at 25°C is 6.0 × 10−10. What is the molar solubility of Ba(IO3)2?

Problem – 17.52

Calculate the solubility of LaF3 (Ksp = 2 x 10-19) in

grams per liter in

(a) pure water

(b) 0.010 M KF solution

(c) 0.050 M LaCl3 solution.

Problem – 17.56

Calculate the molar solubility of Fe(OH)2 (Ksp = 7.9 x 10-16) when buffered at pH

(a) 8.0

(b) 10.0

(c) 12.0.

Equilibria Involving Complex Ions

Complex Ion: A charged species consisting of a metal ion surrounded by ligands (Lewis bases).

Coordination Number: Number of ligands attached to a metal ion. (Most common are 6 and 4.)

Formation (Stability) Constants: The equilibrium constants characterizing the stepwise addition of ligands to metal ions.