AP chem
Text: chapters 14, 15
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•
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K a, K b, K w
Strong acids, weak acids, polyprotic acids
Titrations
PROBLEMS
1. A 0.1-molar solution of acetic acid
-5
(Ka = 1.8 x 10 ) has a pH of about
a. 1
b. 3
c. 7
d. 10
e. 14
2.
3.
Ka, the acid dissociation constant,
-4
for an acid is 9 x 10 at room
temperature. At this temperature,
what is the approximate percent
dissociation of the acid in a 1.0 M
solution?
a. 0.03%
b. 0.09%
c. 3%
d. 5%
e. 9%
a.
b.
c.
d.
e.
5.
6.
7.
Using the information below, choose
the best answer for preparing a pH =
8 buffer.
Acid
Ka value
-3
H3PO4
7.2 x 10
-8
H2PO4
6.3 x 10
2-13
HPO4
4.2 x 10
a. K2HPO4 + KH2PO4
b. H3PO4
c. K2HPO4 + K3PO4
d. K3PO4
e. K2HPO4 + H3PO4
For 4-6, use the following:
-4
Ka = 1.8 x 10 for HCOOH
-4
Kb = 4.4 x 10 for CH3NH2
-2
Ka1 = 3 x 10 for H3PO2
-7
Ka2 = 1.7 x 10 for H2PO2
4.
Review - 11
(Acids and bases)
a solution with pH = 7
a solution with pH < 7, which is
not a buffer solution
a solution with pH < 7, which is
a buffer
a solution with a pH > 7, which
is not a buffer solution
a solution with pH > 7, which is
a buffer
A solution with an initial KCOOH
concentration of 1 M, and an initial
K2HPO2 concentration of 1 M is…
A solution with an initial H3PO2
concentration of 1 M, and an initial
KH2PO2 concentration of 1 M.
Name: ____________________
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Buffers
Salts
Acid base equilibria
A solution with an initial CH3NH2
concentration of 1 M, and an initial
CH3NH3Cl concentration of 1 M.
A solution of a weak base is titrated
with a solution of a standard strong
acid. The progress of the titration is
followed with a pH meter. Which of
the followin observations would
occur?
a. The pH of the solution gradually
decreases throughout the
experiment.
b. Initially the pH of the solution
drops slowly, and then it drops
much more rapidly.
c. At the equivalence point the pH
is 7.
d. After the equivalence point, the
pH becomes constant because
this is the buffer region.
e. The pOH at the equivalence
point equals the pKb of the
base.
d.
e.
The solution is acidic because
of hydrolysis of the NO2 ion.
The solution is basic because of
hydrolysis of the NO2 ion.
For 12-15, use the following:
a.
b.
c.
d.
e.
H2C2O4 (oxalic acid) and
KHC2O4 (potassium hydrogen
oxalate)
KNO3 (potassium nitrate) and
HNO3 (nitric acid)
NH3 (ammonia) and NH4NO3
(ammonium nitrate)
C2H5NH2 (ethylamine) and KOH
(potassium hydroxide)
CH3NH2 (methylamine) and
HC2H3O2 (acetic acid)
12. The most acidic solution (lowest
pH).
13. The solution with a pH nearest 7.
14. A buffer with a pH > 7.
8.
9.
What is the ionization constant, Ka,
for a weak monoprotic acid if a 0.30molar solution has a pH of 4.0?
-10
a. 9.7 x 10
-2
b. 4.7 x 10
-6
c. 1.7 x 10
-4
d. 3.0 x 10
-8
e. 3.3 x 10
Phenol, C6H5OH, has a Ka = 1.0 x
-10
10 . What is the pH of a 0.010 M
solution of phenol?
a. Between 3 and 7
b. 10
c. 2
d. Between 7 and 10
e. 7
10. You are given equimolar solutions of
each of the following. Which has the
lowest pH?
a. NH4Cl
b. NaCl
c. K3PO4
d. Na2CO3
e. KNO3
11. When sodium nitrite is dissolved in
water
a. The solution is acidic because
of hydrolysis of the sodium ion.
b. The solution is neutral.
c. The solution is basic because of
hydrolysis of the sodium ion.
15. A buffer with a pH < 7.
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16. Determine the OH (aq)
concentration in 1.0 M aniline
(C6H5NH2) solution. (Kb for aniline is
-10
4.0 x 10 .)
-5
a. 2.0 x 10 M
-10
b. 4.0 x 10 M
-6
c. 3.0 x 10 M
-7
d. 5.0 x 10 M
0
e. 1.0 x 10 M
-
17. For IO3 + HC2H3O2 ⇌ HIO3 +
C2H3O2 , the above equation has an
equilibrium constant less than 1.
What are the relative strengths of
the acids and bases?
Acids
Bases
a. HIO3 < HC2H3O2 IO3 < C2H3O2
b. HIO3 < HC2H3O2 IO3 > C2H3O2
c. HIO3 > HC2H3O2 IO3 > C2H3O2
d. HIO3 > HC2H3O2 IO3 < C2H3O2
e. HIO3 = HC2H3O2 IO3 = C2H3O2
+
2+
18. For: ZnS(s) + 2H (aq) ⇌ Zn (aq) +
H2S(aq), what is the equilibrium
constant for the above reaction?
The successive acid dissociation
-8
constants for H2S are 9.5 x 10 (Ka1)
-19
and 1 x 10 (Ka2). Ksp for ZnS
-24
equals 1.6 x 10 .
-24
-8
a. 1.6 x 10 / 9.5 x 10
-79
-24
b. 1 x 10 / 1.6 x 10
© S. O’Malley 2010
c.
d.
e.
-27
-24
9.5 x 10 / 1.6 x 10
-8
-24
9.5 x 10 / 1.6 x 10
-24
-27
1.6 x 10 / 9.5 x 10
19. In the equilibrium below, which
species behave as bases?
+
2H2PO4 + H2O ⇌ H3O + HPO4
2i. HPO4
ii. H2PO4
iii. H2O
a. i only
b. i and ii
c. ii and iii
d. i and iii
e. iii only
+
20. H2C3H2O4 + 2H2O ⇌ 2H3O +
2C3H2O4 . As shown above, malonic
acid is a diprotic acid. The
successive equilibrium constants are
-3
-6
1.5 x 10 (Ka1) and 2.0 x 10 (Ka2).
What is the equilibrium constant for
the above reaction?
-14
a. 1.0 x 10
-6
b. 2.0 x 10
-12
c. 4.0 x 10
-9
d. 3.0 x 10
-3
e. 1.5 x 10
-
21. HC3H5O2(aq) + HCOO (aq) ⇌
HCOOH(aq) + C3H5O2 (aq). The
equilibrium constant, K, for the
-2
above equilibrium is 7.2 x 10 . This
value implies which of the following?
a. A solution with equimolar
amounts of HC3H5O2(aq) and
HCOO (aq) is neutral.
b. C3H5O2 (aq) is a stronger base
than HCOO (aq).
c. HC3H5O2(aq) is a stronger acid
than HCOOH(aq).
d. HCOO (aq) is a stronger base
than C3H5O2 (aq).
e. The value of the equilibrium
does not depend on the
temperature.
Free-response
-8
22. Hypochlorous acid, HOCl (Ka = 3.2 x 10 ), is a weak acid in water. (2005 AP form B free response section)
http://www.collegeboard.com/prod_downloads/ap/students/chemistry/ap05_sg_chemistry_form_b.pdf
a. Write a chemical equation showing how HOCl behaves as an acid in water.
b. Calculate the pH of a 0.175 M solution of HOCl.
c. Write the net ionic equation for the reaction between HOCl(aq) and the strong base NaOH(aq).
d. In an experiment, 20.00 mL of 0.175 M HOCl(aq) is placed in a flask and titrated with 6.55 mL of 0.435 M
NaOH(aq).
i. Calculate the number of moles of NaOH(aq) added.
+
ii. Calculate [H3O ] in the flask after the NaOH(aq) has been added.
iii. Calculate [OH ] in the flask after the NaOH(aq) has been added.
23. Lactic acid, HC3H5O3, reacts with water to produce an acidic solution. Shown below are the complete Lewis
structures of the reactants. (2007 AP form B free response section)
http://www.collegeboard.com/prod_downloads/ap/students/chemistry/ap07_sg_chemistry_form_b.pdf
a.
b.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
B
A
C
D
C
E
B
E
A
A
E
B
E
C
A
A
D
E
D
D
B
c.
Complete the equation by drawing the complete Lewis structures of the reaction products.
Choosing from the chemicals and equipment listed below, describe how to prepare 100.00 mL of a 1.00 M
-1
aqueous solution of NH4Cl (molar mass = 53.5 g mol ). Include specific amounts of equipment where
appropriate.
NH4Cl(s)
50 mL buret
100 mL graduated cylinder
100 mL pipet
Distilled water
100 mL beaker 100 mL volumetric flask
balance
Two buffer solutions, each containing acetic acid and sodium acetate, are prepared. A student adds 0.10 mol of
HCl to 1.0 L of each of these buffer solutions and to 1.0 L of distilled water. The table below shows the pH
measurements made before and after the 0.10 mol of HCl is added.
i.
ii.
iii.
22.
23.
Write the balanced net-ionic equation for the reaction that takes place when the HCl is added to
buffer 1 or buffer 2.
Explain why the pH of buffer 1 is different from the pH of buffer 2 after 0.10 mol of HCl is added.
Explain why the pH of buffer 1 is the same as the pH of buffer 2 before 0.10 mol of HCl is added.
(a) HOCl + H2O ⇌ H3O+ + OCl-, (b) 4.13, (c) HOCl + OH- → H2O + OCl-, (d) (i) 0.00285 mol, (ii) 7.3 x 10-9 M, (iii) 1.4 x 10-6 M
(see the link in the problem above for a solution to this problem)
Problems taken from “5 Steps to a 5”, McGraw Hill