AP Chemistry: Chapter 15 Student Notes
Objectives
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15.1:
15.2:
15.3:
15.4:
15.5:
15.6:
15.7:
The Common Ion Effect
What are Buffers
Buffer Capacity
Titrations and pH Curves
Acid-Base Indicators
Solubility Equilbria and Ksp
Ksp and Qualitative Analysis
15.1: The Common Ion Effect
Sometimes equilibrium solutions can have more than one ion—common ions:
NaF(s)  __________ + ____________
This goes into the reaction:
HF(aq) ↔ H+(aq)
+
(dissociation)
F-(aq)
Adding NaF does _________________
Example 1:
What is the percent dissociation of:
a. A 0.10-M solution of acetic acid: CH3COOH?
b. A mixture that contains 0.10-M acetic acid and 0.10-M sodium acetate?
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15.2: What are Buffers
What is a buffer
A buffer ____________ a change in pH
A buffer must have
Example 1
What is the pH of a solution where 50. mL of 0.50M NaC2H3O2 is mixed with 25mL of
0.25M HC2H3O2
A cool Equation: The Henderson-Hasselbach Equation
 [base] 

pH  pKa  log 
 [acid ] 
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Example 2:
Calculate the pH of the following solution: 25 mL of a 0.150M solution of hypochlorous
acid (HOCl) and 32mL of a 0.45M potassium hypochlorite (KOCl)
This can be simplified if you work in mmol throughout the problem
Working in mmol only works if the solution is _______________________
Example 3:
Calculate the pH when 25.0 mL of 0.50 M methylammonium nitrate is mixed with 75 mL
of 0.30 M methylamine.
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Adding Strong Acids and Bases to solutions:
Equivalence Point: When the __________ of ____________ = the ____________ of
___________
How do you know why you have reached the equivalence point?
1.
2.
Diagram for # 2 above
Description of this experiment
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Example 4:[Strong Acid + Salt of Weak Base]
What is the pH when 20mL of 0.25M HCl is reacted with 20 mL of 0.35M sodium nitrite?
Example 5: [Strong Acid + Buffer Solution]
What is the pH when 15 mL of 0.20 M HNO3 is added to a buffer that contains 50.0 mL of
a 0.25 M HCO2H and 0.30 M NaCO2H?
Example 6: [Strong Base + Buffer Solution]
What is the pH when 15 mL of 0.20 M NaOH is added to a buffer that contains 50.0 mL
of a 0.25 M HCO2H and 0.30 M NaCO2H ?
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Example 7: [Strong Base + Weak Base Buffer System]:
What is the pH when 40 mL of 0.25M NaOH is added to a buffer that contains 100 mL of
0.40M Ethalmine (C2H5NH2) and 0.40M ethylammonium chloride (C2H5NH3Cl)?
Example 8: [Strong Base + Weak Acid—ep]
What is the pH when 40 mL of 0.25 M NaOH is mixed with 20 mL of 0.50M propanoic
acid (HC3H5O2)
Example 9: [Strong Base + Weak Acid—beyep]
What is the pH when 42 mL of 0.25 M NaOH is mixed with 20 mL of 0.50M propanoic
acid (HC3H5O2)
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Example 10: [Strong Acid + Weak Base]
What is the pH when 20 mL of 0.20 M methylamine(CH3NH2) is mixed with 10 mL of
0.20 M HNO3.
Example 11: [Strong Acid + Weak Base]
What is the pH when 20 mL of 0.20 M methylamine(CH3NH2) is mixed with 20 mL of
0.20 M HNO3.
Example 11: [Strong Acid + Weak Base]
What is the pH when 20 mL of 0.20 M methylamine(CH3NH2) is mixed with 24 mL of
0.20 M HNO3.
Example 12: [Ratio Problem]
The ratio of NH3 to NH4+ in a buffered solution is 3.2: What is the pH?
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15.4: Titrations and pH Curves
Strong Acid with Strong Base
Strong Base Strong Acid
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Weak Acid-Strong Base
Strong Acid-Weak Acid with Strong Base Comparison
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Ka and Titration Curves
Weak Base with Strong Acid
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A Special Point: when the pH = pKa
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15.5: Acid-Base Indicators
Equivalence Point is determined by
1. Colorometrically:
2. Using a pH meter:
How do you pick an indicator—for colorometric analysis?
1.
2.
Determine the pH at the equivalence point either from a calculation or using a
pH meter
Choose the appropriate indicator from the chart below:
Indicator
Gentian violet (Methyl violet)
Low pH color
Transition pH
range
High pH color
yellow
0.0–2.0
blue-violet
Leucomalachite green (first transition) yellow
0.0–2.0
green
Leucomalachite green (second
transition)
green
11.6–14
colorless
Thymol blue (first transition)
red
1.2–2.8
yellow
Thymol blue (second transition)
yellow
8.0–9.6
blue
Methyl yellow
red
2.9–4.0
yellow
Bromophenol blue
yellow
3.0–4.6
purple
Congo red
blue-violet
3.0–5.0
red
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Methyl orange
red
3.1–4.4
orange
Bromocresol green
yellow
3.8–5.4
blue-green
Methyl red
red
4.4–6.2
yellow
Methyl red / Bromocresol green
red
4.5–5.2
green
Azolitmin
red
4.5–8.3
blue
Bromocresol purple
yellow
5.2–6.8
purple
Bromothymol blue
yellow
6.0–7.6
blue
Phenol red
yellow
6.8–8.4
purple
Neutral red
red
6.8–8.0
yellow
Naphtholphthalein
colorless to
reddish
7.3–8.7
greenish to
blue
Cresol Red
yellow
7.2–8.8
reddish-purple
Phenolphthalein
colorless
8.3–10.0
fuchsia
Thymolphthalein
colorless
9.3–10.5
blue
Alizarine Yellow R
yellow
10.2–12.0
red
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Which indicator would you pick?
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How do indicators work?
All indicators are ____________________
Their “partners” have color and change when the equilibrium is shifted.
HIn ↔ H+ + InBlue
green
Or some such derivation
Each indicator has a Ka and the color change occurs at approximately the –log of
the Ka.
Example 1:
The equivalence point of a titration is 4.3: Which of the following indicators would be a
good choice? Explain your choice.
Indicator
Thymol Blue
Eriochrome Black T
Alizarin
m-Nitrophenol
Thymolphthalien
Alizarin Yellow R
Ka
2.3 x 10-2
5.4 x 10-5
6.6 x 10-6
8.3 x 10-8
2.5 x 10-9
4.3 x 10-11
15.6: Solubility Equilbria and Ksp
Ksp: the study of the solubility’s of ____________ ionic ________________
Writing Solubility Product Expressions
CaF2
Mg3(PO4)2
BaSO4
Ksp values are _______________
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Ksp Values
Bromides
Carbonates
Hydroxides
Iodides
Oxalates
Phosphates
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PbBr2
AgBr
BaCO3
CaCO3
CoCO3
CuCO3
FeCO3
PbCO3
MgCO3
MnCO3
NiCO3
Ag2CO3
ZnCO3
AgOH
Al(OH)3
Ca(OH)2
Cr(OH)3
Co(OH)2
Cu(OH)2
Fe(OH)2
Fe(OH)3
Pb(OH)2
Mg(OH)2
Mn(OH)2
Ni(OH)2
Zn(OH)2
PbI2
AgI
BaC2O4
CaC2O4
MgC2O4
AlP04
Ba3(P04)2
Ca3(P04)2
CrP04
Pb3(P04)2
Ag3P04
Zn3(P04)2
6.3 x 10-6
3.3 x 10-13
8.1 x 10-9
3.8 x 10-9
8.0 x 10-13
2.5 x 10-10
3.5 x 10-11
1.5 x 10-13
4.0 x 10-5
1.8 x 10-11
6.6 x 10-9
8.1 x 10-12
1.5 x 10-11
2.0 x 10-8
1.9 x 10-33
7.9 x 10-6
6.7 x 10-31
2.5 x 10-16
1.6 x 10-19
7.9 x 10-15
6.3 x 10-38
2.8 x 10-16
1.5 x 10-11
4.6 x 10-14
2.8 x 10-16
4.5 x 10-17
8.7 x 10-9
1.5 x 10-16
1.1 x 10-7
2.3 x 10-9
8.6 x 10-5
1.3 x 10-20
1.3 x 10-29
1.0 x 10-25
2.4 x 10-23
3.0 x 10-44
1.3 x 10-20
9.1 x 10-33
Chlorides
Chromates
Cyanides
Fluorides
Sulfates
Sulfides
Sulfites
PbCl2
AgCl
BaCrO4
CaCrO4
PbCrO4
Ag2CrO4
Ni(CN)2
AgCN
Zn(CN)2
BaF2
CaF2
PbF2
MgF2
BaS04
CaS04
PbS04
Ag2S04
CaS
CoS
CuS
FeS
Fe2S3
PbS
MnS
NiS
Ag2S
ZnS
BaS03
CaS03
Ag2S03
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1.7 x 10-5
1.8 x 10-10
2.0 x 10-10
7.1 x 10-4
1.8 x 10-14
9.0 x 10-12
3.0 x 10-23
1.2 x 10-16
8.0 x 10-12
1.7 x 10-6
3.9 x 10-11
3.7 x 10-8
6.4 x 10-9
1.1 x 10-10
2.4 x 10-5
1.8 x 10-8
1.7 x 10-5
8 x 10-6
5.9 x 10-21
7.9 x 10-37
4.9 x 10-18
1.4 x 10-88
3.2 x 10-28
5.1 x 10-15
3.0 x 10-21
1.0 x 10-49
2.0 x 10-25
8.0 x 10-7
1.3 x 10-8
1.5 x 10-14
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Example 1:
The solubility of copper I bromide is 2.0 x 10-4M. What is the value of Ksp
Example 2:
What is the molar solubility of Silver Sulfide?
Example 3:
What is the molar solubility of Bismuth III Sulifde. The Ksp = 1.1 x 10-73
Example 4:
What is the molar solubility of Iron III hydroxide
Silly Hydroxides
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Will a ppt form:
Compare the Ksp to the Qsp
Qsp> Ksp ________________
Qsp < Ksp ________________
Example 5: Will a ppt form
50.0 mL of 0.00025 M Na3PO4 is mixed with 50.0 mL of 0.0025 M BaCl2. Will a ppt
form? Show all calculations to support your answer.
Example 6: Competing ppt
Sodium chloride is added to a 50 mL beaker that contains a mixture or 0.00015 M lead II
nitrate and 0.00035 M silver nitrate. What ppt will form first. Show all work.
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Example 7: Competing ppt and how much remains.
At 25˚C the solubility product constant, Ksp, for strontium sulfate, SrSO4, is 7.610-7.
The solubility product constant for strontium fluoride, SrF2, is 7.910-10.
(a) What is the molar solubility of SrSO4 in pure water at 25˚C?
(b) What is the molar solubility of SrF2 in pure water at 25˚C?
(c) An aqueous solution of Sr(NO3)2 is added slowly to 1.0 litre of a well-stirred
solution containing 0.020 mole F- and 0.10 mole SO42- at 25˚C. (You may assume
that the added Sr(NO3)2 solution does not materially affect the total volume of the
system.)
1. Which salt precipitates first?
2.
What is the concentration of strontium ion, Sr2+, in the solution when the first
precipitate begins to form?
(d)
As more Sr(NO3)2 is added to the mixture in (c) a second precipitate begins to
form. At that stage, what percent of the anion of
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AP Chemistry Chapter 15 Practice Exam
1. A buffer solution contains 0.40 mole of formic acid, HCOOH, and 0.60 mole of
sodium formate, HCOONa, in 1.00 liter of solution. The ionization constant, Ka, of
formic acid is 1.8x10-4.
(a) Calculate the pH of this solution.
(b) If 100. milliliters of this buffer solution is diluted to a volume of 1.00 liter with pure
water, the pH does not change. Discuss why the pH remains constant on dilution.
(c) A 5.00 milliliter sample of 1.00 molar HCl is added to 100. milliliters of the original
buffer solution. Calculate the [H3O+] of the resulting solution.
(d) A 800.-milliliter sample of 2.00-molar formic acid is mixed with 200. milliliters of
4.80-molar NaOH. Calculate the [H3O+] of the resulting solution.
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2. In water, hydrazoic acid, HN3, is a weak acid that has an equilibrium constant, Ka,
equal to 2.8x10-5 at 25ºC. A 0.300 liter sample of a 0.050 molar solution of the acid is
prepared.
(a) Write the expression for the equilibrium constant, Ka, for hydrazoic acid.
(b) Calculate the pH of this solution at 25ºC.
(c) To 0.150 liter of this solution, 0.80 gram of sodium azide, NaN3, is added. The salt
dissolved completely. Calculate the pH of the resulting solution at 25ºC if the
volume of the solution remains unchanged.
(d) To the remaining 0.150 litre of the original solution, 0.075 liter of 0.100 molar
NaOH solution is added. Calculate the [OH-] for the resulting solution at 25ºC.
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3. The equations and constants for the dissociation of three different acids are given
below.
HCO3- <=> H+ + CO32Ka = 4.2 x 10-7
H2PO4- <=> H+ + HPO42Ka = 6.2 x 10-8
HSO4- <=> H+ + SO42Ka = 1.3 x 10-2
(a) From the systems above, identify the conjugate pair that is best for preparing a buffer
with a pH of 7.2. Explain your choice.
(b) Explain briefly how you would prepare the buffer solution described in (a) with the
conjugate pair you have chosen.
(c) If the concentrations of both the acid and the conjugate base you have chosen were
doubled, how would the pH be affected? Explain how the capacity of the buffer is
affected by this change in concentrations of acid and base.
(d) Explain briefly how you could prepare the buffer solution in (a) if you had available
the solid salt of the only one member of the conjugate pair and solution of a strong
acid and a strong base.
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4. The solubility of iron(II) hydroxide, Fe(OH)2, is 1.43x10-3-gram per litre at 25oC.
(a) Write a balanced equation for the solubility equilibrium.
(b) Write the expression for the solubility product constant, Ksp, and calculate its value.
(c) Calculate the pH of a saturated solution of Fe(OH)2 at 25oC.
(d) A 50.0-millilitre sample of 3.00x10-3 molar FeSO4 solution is added to 50.0millilitres of 4.00x10-6-molar NaOH solution. Does a precipitate of Fe(OH)2 form?
Explain and show calculations to support your answer.
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5. Solve the following problem related to the solubility equilibria of some metal
hydroxides in aqueous solution.
(a) The solubility of Cu(OH)2(s) is 1.72x10-6-gram per 100. milliliters of solution at
25°C.
(i) Write the balanced chemical equation for the dissociation of Cu(OH)2(s) in
aqueous solution.
(ii) Calculate the solubility (in moles per liter) of Cu(OH)2 at 25oC.
(iii) Calculate the value of the solubility-product constant, Ksp, for Cu(OH)2 at 25oC.
(b) The value of the solubility-product constant, Ksp, for Zn(OH)2 is 7.7x10-17 at 25oC.
(i) Calculate the solubility (in moles per liter) of Zn(OH)2 at 25oC in a solution
with a pH of 9.35.
(ii) At 25oC, 50.0-milliliters of 0.100-molar Zn(NO3)2 is mixed with 50.0-milliliters
of 0.300-molar NaOH. Calculate the molar concentration of Zn2+(aq) in the
resulting solution once equilibrium has been established. Assume that volumes
are additive.
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