AP Chemistry: Chapter 17 Student Notes
Objectives
• 17.1: The Common Ion Effect
• 17.2: What are Buffers
• 17.3: Titrations and pH Curves
• 16.4: Acid-Base Indicators
• 17.4: Solubility Equilbria and Ksp
• 17.5: Ksp and Qualitative Analysis
17.1: The Common Ion Effect
Sometimes equilibrium solutions can have more than one ion—common ions:
NaF(s) __________ + ____________ (dissociation)
This goes into the reaction:
HF(aq) ↔H+(aq) + F-(aq)
Adding NaF does _________________
Example 1:
What is the pH 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?
17.2: Buffers
What is a buffer?
A buffer ____________ a change in pH
A buffer must have:
1)
2)
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 ]
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.
17.3 Titrations and pH Curves
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
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 ?
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)
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?
17.4: Titrations and pH Curves
Strong Acid with Strong Base
Strong Base Strong Acid
Weak Acid-Strong Base
Strong Acid-Weak Acid with Strong Base Comparison
Ka and Titration Curves
17.5: Acid-Base Indicators
Weak Base with Strong Acid
Equivalence Point is determined by
1. Colorometrically:
2. Using a pH meter:
How do you pick an indicator—for colorometric analysis?
1. Determine the pH at the equivalence point either from a calculation or using a
pH meter
2. Choose the appropriate indicator from the chart below:
Which indicator would you pick?
How do indicators work? All indicators are ____________________
Their “partners” have color and change when the equilibrium is shifted.
HIn  H+ + InEach 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 agood choice? Explain
your choice.
Indicator Ka
Thymol Blue 2.3 x 10-2
Eriochrome Black T 5.4 x 10-5
Alizarin 6.6 x 10-6
m-Nitrophenol 8.3 x 10-8
Thymolphthalien 2.5 x 10-9
Alizarin Yellow R 4.3 x 10-11
17.6: 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 _______________
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
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.
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 the second form exists?