CHM 2046C Module 12 Homework Packet Name: __________
Jespersen
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Module 12i:
Acid-Base Equilibria in Aquous
Solutions Jespersen Chapter 16 |
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Study Guide Kotz Sections 18.1-18.3 Buffers/Acid-Base Titrations |
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B1. Preparing a Buffer with a desired pH Section 16.7 |
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C. Derivation of
Henderson-Hasselbalch Equation Section 16.7 |
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Module 12i Total: |
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Study Guide Kotz Sections 18.4-18.7Hetergenous Phase Equilibria |
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Module 12ii: Solubility and
Simultaneous Equilibria Jespersen Chapter 17 |
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Module
12ii Total: |
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Part A. The Common-Ion
and the Common-Ion Effect 1 point
1. What is a
common-ion?
2. What would be the common ion for the following
weak acids and weak bases and in addition to its formula give its name:
(a) Lactic Acid: CH3CH(OH)COOH
(b) Acetic Acid: CH3COOH
(c) Ammonia dissolved
in Water: NH3 in H2O
Page 2
3. What is the
Common-Ion Effect for Weak acids, Weak base, or Salt Solubility?
Discuss Examples in General:
a)
b)
4. Does the pH of the
solution increase, decrease, or stay the same when you:
(a) Add Ammonium
nitrate to a dilute aqueous solution of ammonia?
(b) Add solid potassium acetate to a dilute
aqueous solution of acetic acid?
(c) Add solid sodium
lactate to a dilute aqueous solution of lactic acid?
(d) Add solid
Potassium nitrate to a dilute solution of Potassium hydroxide?
Page 3
Part B. Controlling
pH: Buffer Solutions 1 points
1. What is a buffer
solution?
2. (a) State the Henderson-Hasselbalch
equation for an acid:
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2. (b) State the Henderson-Hasselbalch equation for a base:
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3. According to the
(a)What are they?
The first: The Second
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(b) What are the relative magnitudes of
their contribution to the value of the pH?
4. What are the
two requirements for a buffer solution?
(a)
(b)
5. Explain how
it is that the pH of a buffer solution does not change if the solution is
diluted.
6. Explain what
happens if the concentrations of the buffering components are increased so that
their proportion maintains the same ratio.
Page 4
Part C1. Derivation
of Henderson-Hasselbalch Equations 1 point
1. Given the
ionization of weak acid, HA in water, derive the Henderson-Hasselbalch
equation for a weak acid.
Start with writing the ionization
reaction of the Weak Acid and setup the equilibrium constant expression for Ka:
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2. Given the
ionization of base acid, BOH in water, derive the Henderson-Hasselbalch
equation for a weak base.
Start with writing the ionization
reaction of the weak Base and setup the equilibrium constant expression for Kb:
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Part D
Acid-Base/pH Titrations Curves 3
points 1. How do titration curves differ for
strong acid-strong base, strong acid-weak base; weak acid-strong base, and
weak acid-weak base titrations? Sketch curves for each and describe their pH
equivalence points.
Page
7 Part E.
Acid-Base/pH Titration Points of Interests 3 points
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Page 9:
Part F.
Acid-Base/pH Titration Calculations 3 points
You titrate 25.0 mL of 0.10 M NH3
with 0.10 M HCl. a. What is the
pH of the solution Before acid has been added? b. What is the
pH at the equivalence point? c. What is the pH at the halfway point of the titration? If you can do Part (d) successfully it will be worth a 5 point
bonus: Bonus: d.
Calculate the pH of the solution after adding 5.00, 15.0, 20.0,
22.0, and 30.0 mL of acid. Combine this information with that in parts a., b., and c. and
plot the titration curve.. |
Page 10:
Part G: Writing Ksp
Expressions 1 point
Write the dissociation reactions for the following insoluble salts, then setup the Ksp expressions:
1. Cu(OH)2 (s) + HOH (l) ß à
Ksp = Ksp = 2.2 x 10 -20
2. MgCO3 (s) + HOH (l) ß à
Ksp = Ksp = 6.8 x 10 -6
3. AlPO4 (s) + HOH (l) ß à
Ksp = Ksp = 6.3 x 10 -19
4. Cu2S (s) + HOH (l) ß à
Ksp = Ksp = 2.2 x 10 -48
5. Mg(OH)2 (s) + HOH (l) ß à
Ksp = Ksp = 5.6 x 10 -12
Page 11:
H. Comparing Solubility of Salts 1 point
An insoluble substance is defined as one
that has solubility less than 0.01 moles/liter of water.
Using Ksp values, tell which
salt in each pair is more soluble in water:
1. AgCl (Ksp =
1.8 x 10-10) or AgCN (Ksp
= 6.0 x 10-17)
Answer:
2. Mg(OH)2 (Ksp = 5.6
x 10-12) or Ca(OH)2 (Ksp
= 5.5 x 10-5)
Answer:
3. MgCO3 (Ksp = 6.8
x 10-6) or CaCO3 (Ksp = 3.4 x 10-9)
Answer:
4. FeCO3 (Ksp = 3.1
x 10-11) or Ag2C2O4 (Ksp
= 5.4 x 10-12)
Answer:
Which compound in each pair is more soluble in water than is predicted
by a calculation from Ksp :
5. AgBr (Ksp
= 5.4 x 10-13) or AgCN (Ksp = 6.0 x 10-17)
Answer:
6. PbCl2 (Ksp
= 1.7 x 10-5) or PbCO3 (Ksp
= 7.4 x 10-14)
Answer:
7. PbCl2 (Ksp
= 1.7 x 10-5) or PbF2 (Ksp
= 3.3 x 10-8)
Answer:
Which insoluble
compound in each pair should be more soluble in nitric acid than in pure water:
8. PbCl2 (Ksp
= 1.7 x 10-5) or PbS (Ksp = 9.0 x 10-24)
Answer:
9. AgCl (Ksp =
1.8 x 10-10) or Al(OH)3 (Ksp = not known)
Answer: A
10. AgI (Ksp
= 8.5 x 10-17) or Ag2CO3 (Ksp
= 8.5 x 10-12)
Answer:
Page 13:
I. Estimating Salt Solubility from Ksp 2 points
Given the following insoluble salts, estimate the salt’s molar solubility from the Ksp :
1. Cu(OH)2 (s) + HOH (l) ß à
Ksp = 2.2 x 10 -20
2. MgCO3 (s) + HOH (l) ß à
Ksp = 6.8 x 10 -6
3. Al PO4 (s) + HOH (l) ß à
Ksp = 9.8 x 10 -21
4. Cu2S (s) + HOH (l) ß à
Ksp = 2.2 x 10 -48
Page 14:
J. Determining Ksp
from a Salt’s Solubility 2 points
Given the following insoluble salts, estimate the salt’s molar solubility from the Ksp :
1. Calcium hydroxide, Ca(OH)2 (s) dissolves in water to the extent of 1.04 g/L. What is the value of the Ksp ?
2. You place 1.834 g of Ca(OH)2 (s) in 1.00 L of pure water at 25oC. The pH of the solution is found to be 12.68. What is the value of the Ksp ?
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Part L. Common Ion Effect on Solubility 2 points
What is the solubility, in milligrams per milliliter, of BaCO3, Ksp
=1.2 x 10
-10, in
(a) pure water and
(b) in water containing 10.0
mg/ml of Na2CO3?
(c) What is the pH of the solution? Ka2HCO3 = 4.8 x 10 -11
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Part N: Precipitation
Reactions 2 points
If the concentration of the
strontium ion, Sr 2+, is 2.5 x 10 -4
M does precipitation of SrSO4 occur
when enough of the soluble salt Na2SO4
is added to make the solution,
2.5
x 10 -4 M? Ksp for SrSO4
is 3.4 x 10 -7
Page
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Part
K: Chapter 17 Key Terms 1
point
1. ___________________ – a measure of how much hydronium
or hydroxide can be added to a buffer solution before the buffer can no longer
control the pH
2. .________________ – a material
that is one color below a particular pH and another color above it; typically these
materials are themselves weak acids or bases.
3. ___________________ – a solution that resists a change in pH when
hydroxide or hydronium ions are added.
4. ___________________ - The
limiting of acid (or base) ionization caused by the addition of its conjugate
base (or conjugate acid)
5. ___________________ – the point in a titration at which one reactant
has been exactly consumed by the solution of the other reactant.
6. ___________________ – the equilibrium constant for the formation of a
complex ion.
7. ___________________ – an equation that allows us to calculate the pOH of a buffer:
[conjugate acid] |
pOH = pKb + log ------------------------ |
[base] |
8. ___________________ – the equilibrium constant for the dissociation of
a sparingly soluble salt. For the salt AxBy, it has the
form:
Ksp = [A y+]x [B x-]y
9. ___________________ – the substance added during a titration.
10. __________________ – the addition of a base to an acid (or vice versa)
until the equivalence point is reached