chapter 17 sections 17.1, 17.2, 17.4, 17.5(common ion effect)
TRANSCRIPT
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.