Chemistry 11: Equilibrium—Aqueous
A.
Acids and Bases (16.1 to16.11)
1. three theories
a. Arrhenius: acid (H+) reacts with base (OH-) to
form water (neutralization reaction):
H+ + OH-  H2O (not an equilibrium)
b. Brønsted-Lowry: proton donor (HA) transfers
proton (H+) to proton acceptor (:B)
1. HA + :B  :A- + HB+ (equilibrium)
2. free H+ shown as hydronium ion: H3O+
c. Lewis: electron pair acceptor (M+) attaches to an
electron pair donor (:B)
1. M+ + :B  M:B+ (complex ion formation)
2. H3B + :NH3  H3BNH3 (incomplete octet)
2. acids
a. H+ + anion (A-) combine to make (HA)
b. when A- is Cl-, Br-, I-, NO3-, ClO3-, ClO4, SO42c. when A- is any other anion—weak acid/electrolyte
1. % ionization < 5 %
2. organic acids, HC2H3O2 or CH3COOH
3. A- is oxyanion, HXOy (HClO, HBrO2, H2SO3)
a. H is bonded to O
b. weaker O-H bond = stronger acid
1.  electronegativity  (HClO > HBrO)
2.  oxygen  (HClO2 > HClO)
4. nonmetal oxides: CO2 + H2O  H2CO3
d. polyprotic acids (HxA)—first H+ is easiest to
remove from neutral molecule (successive H+ are
harder to remove from anion)
e. acid ionization
1. HA(aq)  H+ + A2. Ka = [H+][A-]/[HA] = [H+]2/[HA]
3. larger Ka = stronger acid
4. polyprotic acids (HxA)
a. H2A  H+ + HA-: Ka1 = [H+]2/[H2A]
b. HA-  H+ + A2-: Ka2 = [H+][A2-]/[HA-]
[HA-]  [H+]  Ka2 = [A2-]
3. bases
a. strong: group 1, Ca2+, Sr2+, Ba2+ hydroxides
b. weak
1. insoluble hydroxides (antacids)
2. ammonia or amines (:NH3, :NH2CH3, etc.)
react with water (hydrolysis reaction)
:NH3(aq) + H2O  NH4+ + :OHKb = [NH4+][OH-]/[NH3] = [OH-]2/[NH3]
(water is not included because it is liquid)
3. anions from weak acids (:A-) = weak
:A- + H2O  HA(aq) + :OHKb = [HA][OH-]/[A-] = [OH-]2/[A-]
4. acid/base properties of aqueous ionic compounds (MX)
strong bases
others
M+
(spectator ions) M+ + H2O  H+ + MOH
X 
strong acids
neutral
acidic
(spectator ions)
acid anions, HX
acidic
acidic
HX-  H+ + X2others
basic
?
X- + H2O  HX + OH5. autoionization of water
a. H2O  H+ + OH1. Kw = [H+][OH-] = 1 x 10-14
2. pure water at 25oC: [H+] = [OH-] = 1 x 10-7 M
b. solutions: acidic: [H+] > [OH-], basic: [H+] < [OH-]
c. pH scale
1. pH = -log [H+], pOH = -log[OH-]
2. pH scale: acidic < 7, neutral = 7, basic > 7
3. pH + pOH = pKW = 14
4. significant figures (right of decimal): 3.45 (2 sf)
Name ____________________________
6.
B.
Brønsted-Lowry reaction: HA + :B  HB+ + :Aa.
HA  H+ + :AKa-HA
+
H + :B  HB+
Ka-HB+
HA + :B  HB+ + :A- K = Ka-HA/Ka-HB+
if K > 1 then HA is a stronger acid than HB+ and
:B is a stronger base than :Ab. conjugate pairs (HA, :A- and :B, HB+)
c. relationship between Ka and Kb for a conjugate pair
1.
HX  H+ + XKa
X- + H2O  HX + OHKb
H2O  H+ + OHKw
2. Ka x Kb = 1 x 10-14 (pKa + pKb = 14)
d. when to use Ka and Kb
K
acid + base  acid
+ base
Ka
HA
+
H2O  H3O+ +
A1/Ka
H3O+ +
B

HB+
+ H2O
 H2O
1/Kb
HA
+
OH
+
AKb
H2O +
B

HB+
+
OHe. amphiprotic
1. proton donor or proton acceptor depending on
other reactant (contains : and H+)
2. HCO3-, HSO3-, H2O, NH3
Acid-Base Equilibrium Problems
1. pure acid/base equilibrium problems
a. determine Ka, given pH or [H+]E and [HA]o

set up "ICE Box" (shaded boxes are given)
[]
HA

H+
+
A[HA]o
0
0
I
C
–[H+]E
+[H+]E
+[H+]E
E
[HA]o – [H+]E
[H+]E
[H+]E
+
+

solve for Ka = [H ]E[A ]E/[HA]E = [H ]E2/([HA]E
determine Kb, given pOH or [OH-]E and [B]o

set up "ICE Box" (shaded boxes are given)
[]
B
+
H 2O

HB+ +
OH[B]o
0
0
I
+[OH-]E
+[OH-]E
C
–[OH-]E
[OH-]E
[OH-]E
E
 [B]o
+

solve for Kb = [OH ]E[HB ]E/[B]E = [OH-]E2/[B]o
b. determine all [ ]E, given [HA]o and Ka

set up "ICE Box" (shaded boxes are given)
[]
HA

H+
+
A[HA]o
0
0
I
–x
+x
+x
C
x
x
E
[HA]o (x < 5%)

solve for x (Ka = [H+]E[A-]E/[HA]E = x2/[HA]o)

solve for [ ]E
determine all [ ]E, for a polyprotic acid, given [H2A]o,
Ka1 and Ka2

set up "ICE Box" (shaded boxes are given)
[]
H2A

H+
+
HA[H2A]o
0
0
I
–x1
+x1
+x1
C
x1
x1
E [H2A]o (x1 < 5%)

solve for x1 (Ka = x12/[H2A]o)

use x1 for the initial concentration of HA- and
H+ in a new "ICE Box" for HA-  H+ + A2determine [ ]E, given [B]o and Kb

set up "ICE Box" (shaded boxes are given)
[]
B
+
H2O

HB+ +
OH[B]o
0
0
I
+x
+x
C
–x
x
x
E
 [B]o

solve for x (Kb = [HB+]E[OH-]E/[B]E = x2/[B]o)

solve for [ ]E
2.
C.
mixed acid/base with conjugate problems
a. two situations
1. incomplete titration of weak acid or weak base
2. conjugate salt is added to weak acid or base
b. system resists change in pH when strong acid or
strong base is added (buffer)
1. addition of acid: A- + H+  HA(aq)
2. addition of base: HA(aq) + OH-  A- + H2O
3. process is reversible, but not equilibrium
c. determine [H+]E, given [HA]o, [A-]o and Ka

set up "ICE Box" (shaded boxes are given)
[]
HA

H+
+
A-]
[HA]
0
[A
I
o
o
–x
+x
+x
C
x
E
 [HA]o
 [A-]o
+

solve for x = [H ]E (Ka = (x)([A ]o)/[HA]o)
determine [OH-]E, given [HA]o, [A-]o and Kb

set up "ICE Box" (shaded boxes are given)
[]
A+
H2O

HA
+
OH[A ]o
[HA]o
0
I
+x
+x
C
–x
x
E
[HA]o
 [A-]o

solve for x = [OH-] (Kb = (x)([HA]o)/[A-]o)
d. determine [A-]o, given [HA]o, equilibrium pH and Ka

set up "ICE Box" (shaded boxes are given)
[]
HA

H+
+
Ax
[HA]o
0
I
C
–[H+]E
+[H+]E
+[H+]E
E
[H+]E
x
 [HA]o

solve for x = [A-]o (Ka = (x)([H+]E/[HA]o)
determine [HA]o, given [A-]o, equilibrium pOH and Kb

set up "ICE Box" (shaded boxes are given)
[]
A+
H2O

HA
+
OHx
[A-]o
0
I
C
+[OH-]E
+[OH-]E
–[OH-]E
E
[OH-]E
x
 [A-]o

solve for x= [HA]o (Kb = (x)([OH-]E/[A-]o)
Acid-Base Titration (17.3)
1. acid or base is added to a fixed amount of base or acid
a. pH is monitored using a pH meter
b. equivalence when moles of H+ = moles OH1. nH+MaVa = nOH-MbVb
2. indicator changes color at the end point
SA + SB
SA + WB
WA + SB
bromthymol blue
methyl red
phenolphthalein
c. buffered solution during incomplete neutralization
2. graphs (0.1 M NaOH added to 25 mL of 0.1 M HCl
(lower) and 25 mL of 0.1 M HAc (upper)
excess NaOH: [OH-] = (nOH- – nH+)/Vtotal
pure F-: [H+] = (Kb[Ac-])½
equivalence
water: [H+] = [OH-]
buffer: [H+] = Ka[HAc]/[Ac-]
pH = pKa
+
pure HAc: [H ] = (Ka[HAc])½
pure HCl: [H+] = 0.1 M
mL of 0.1 M NaOH
D.
E.
Solubility Equilibrium (17.4-17.5)
1. ionic compound (salt) equilibrium with its ions
a. MmXn(s)  m Mn+(aq) + n Xm-(aq)
b. Ksp = [Mn+]m[Xm-]n (Ksp or mass action expression)
1. subscripts become exponents
2. MmXn(s) is not included [ ] doesn't change
2. solubility equilibrium problems
a. determine one [ ]E, given the other [ ]E and Ksp

Write a Ksp expression from formula

fill in Ksp and [ ] (Don't multiple [ ] by subscript)

solve for missing concentration
b. determine solubility (mol/L) “s”, given Ksp

set up "ICE Box" (shaded boxes are given)
[]
MmXn

m Mn+
+
n Xm0
0
I
+m•s
+n•s
C
m•s
n•s
E

solve for s, Ksp = (m•s)m(n•s)n

general solutions
o MX, then Ksp = (s)(s) = s2
o MX2 or M2X, then Ksp = (s)(2s)2 = 4s3
o MX3 or M3X, then Ksp = (s)(3s)3 = 27s4
c. determine solubility “s”, given Ksp and [Xm-]o or [Mn+]o

set up "ICE Box" (shaded boxes are given)
[]
MmXn

m Mn+
+
n Xm0
[Xm-]o
I
C
+m•s
+n•s
m•s
E
 [Xm-]o

solve for s, (Ksp = (m•s)m[Xm-]on)
d. determine Ksp, given solubility (s)

set up "ICE Box" (shaded boxes are given)
[]
MmXn

m Mn+
+
n Xm0
0
I
+m•s
+n•s
C
m•s
n•s
E

solve for Ksp = (m•s)m(n•s)n
e. determine if a precipitate will form, given [ ]o

write Ksp expression, set equal to “Q”

substitute [ ]o of each ion into the expression

if Q > Ksp, then a precipitate forms

if Q < Ksp, then no precipitate forms
Factors that Affect Solubility (17.5)
1. common ion effect
a. salts are less soluble in solution with common ion
b. Le Chatelier's principle: MX(s)  M+ + X(higher [ ] of a product ion shifts equilibrium left)
2. addition of acid (H+)
a. most anions (except derived from a strong acid) act
as a weak base and form an equilibrium with H+
MX(s) + H+  M+ + HX(aq)
b. Le Chatelier's principle (higher [ ] of reactant (H+)
shifts equilibrium to the right  more soluble)
CaCO3(s) + 2 H+  Ca2+ + CO2(g) + H2O
Mg(OH)2(s) + 2 H+  Mg2+ + 2 H2O
CuS(s) + 2 H+  Cu2+ + H2S(g)
3. addition of ligands (i.e. CN-, Cl-, OH-, NH3)
a. equilibrium exists between cation, ligand and
coordination complex
Cu2+ + 4 CN-  Cu(CN)42Kf = [Cu(CN)42-]/[Cu2+][CN-]4
b. can NH3 dissolve AgCl(s)?
AgCl(s)  Ag+ + ClKsp = 1.8 x 10-10
+
+
Ag + 2 NH3(aq)  Ag(NH3)2
Kf = 1.7 x 107
AgCl(s) + 2 NH3(aq)  Ag(NH3)2+ + ClK = 3.1 x 10-3
answer: AgCl(s) is soluble when [NH3] >> 1 M
c. some metal hydroxides such as Al(OH)3 can act
as an acid (absorb OH-) or base (absorb H+), thus
are soluble in both (amphoterism)
Experiments
1.
Acid-Base Properties (Wear Goggles)—Observe the
properties of acids and bases in a variety of indicators and
the acid/base properties of salts.
Experiment 1—Acid/base indicators: Add 10 drops of 0.1 M
HCl to 4 separate wells. Dip red litmus paper into well #1,
dip blue litmus paper into well #2, add a drop of universal
indicator to well #3 and add a drop of phenolphthalein to
well #4; record color colors of each well. Repeat with 0.1 M
HC2H3O2, NaOH and NH3.
a. Record the colors in the space provided.
red
blue Universal Indicator PhenolAcid/Base
litmus litmus Color
phthalein
pH
HCl
HC2H3O2
NaOH
NH3
Experiment 2—Determine the pH of some salts: Add 10
drops of the salt listed in table b to separate wells. Add 1
drop of universal indicator and record the pH.
b. Complete the table for the four salts tested and
compare the predicted with actual pH range.
Spectator Predicted pH
Well #
Salt
Actual pH
ion(s)
>7 =7 <7
1
Zn(NO3)2
2
NaCl
3
Na3PO4
4
NH4C2H3O2
c. Complete the table by filling in the cation groups and
anion groups that are neutral, acidic or basic.
Acid-Base Property
Cation
Anion
Neutral
Acidic
2.
Buffer Lab (Wear Goggles)—Make buffers using two
techniques and examine their ability to "buffer" a solution.
Experiment 1—Testing water: Add 20 mL of distilled water
to each of two 50 mL beakers labeled A and B. Measure
pH using the pH meter. Add 20 drops of 0.10 M HCl to
beaker A and measure pH. Add 20 drops of 0.10 M NaOH
to beaker B and measure pH. Record the data in table b.
Experiment 2—Partially neutralized weak acid technique:
Add 50. mL of 0.10 M HC2H3O2 to a 150 mL beaker. Add 6
M NaOH, drop by drop, while monitoring the pH with the pH
meter. Stop when the pH = 5.0. Follow the procedure from
experiment 1 using the pH-5 buffer instead of water. Record
the data in table b.
Experiment 3—Add salt of the conjugate acid to a weak
base technique: Add the calculate amount of NH4NO3 (see
part a) to 50 mL of 0.10 M NH3 in a 150 mL beaker.
Measure the pH. Follow the procedure from experiment 1
using the pH-9 buffer instead of water. Record the data in
table b.
a. Calculate the amount of NH4NO3 needed for 50 mL of
a pH 9 buffer made with 0.10 M NH3 (Kb = 1.8 x 10-5).
pOH
[OH-]
[NH4+]
mass
NH4+
b. pH data
Experiment
1
Initial
After HCl
After NaOH
2
c.
3
How effective were the buffered solutions in controlling
changes in pH compared to pure water?
Basic
pH Profile Lab (Wear Goggles)—Monitor the pH while neutralizing acid/base and compare the pH profile to theoretical graphs.
Experiment 1—Strong A/Strong B:
Experiment 2—Weak A/Strong B:
Experiment 3—Weak B/Strong Acid:
Add 25 mL of 0.10 M HCl to 150 mL
Add 25 mL of 0.10 M HC2H3O2 to
Add 25 mL of 0.10 M NH3 to 150 mL
beaker. Measure pH. Add increments
150 mL beaker. Measure pH. Add 10
beaker. Measure pH. Add 10 mL
of 10 mL of 0.10 M NaOH until a total
mL increments of 0.10 M NaOH until
increments of 0.10 M HCl until a total
of 50 mL has been added. Measure
50 mL has been added. Measure the
of 50 mL has been added. Measure
pH after each 10 mL aliquot is added.
pH after each 10 mL aliquot is added.
the pH after each 10 mL aliquot.
a. pH data
c. pH data
e. pH data
mL of NaOH Added
mL of NaOH Added
mL of HCl Added
0
10
20
30
40
50
0
10
20
30
40
50
0
10
20
30
40
50
Graph pH vs. mL of NaOH added.
d.
Graph pH vs. mL of NaOH added.
f.
Graph pH vs. mL of HCl added.
12
12
10
10
10
8
8
8
pH
12
pH
b.
pH
3.
6
6
6
4
4
4
2
2
2
0
20
40
mL of 0.10 M NaOH added
0
20
40
mL of 0.10 M NaOH added
0
20
40
mL of 0.10 M HCl added
g.
Moles of A (first chemical added), moles of B (added in 10 mL aliquots), and total volume. These are the same values for
all titrations.
VA (L)
0.025
0.025
0.025
0.025
0.025
0.025
0.025
VB (L)
0
0.010
0.020
0.025
0.030
0.040
0.050
nA = MAVB
nB = MBVB
Vtot = VA + VB
h. Determine the theoretical pH when 0.10 M NaOH is added to 0.10 M HCl.
titration stage
pure acid
pre equivalence: [HA] > [OH-] equivalence
initial reaction
H+excess + OH-  H+ + H2O
H+ + OH-  H2O
+
calculation
[H ] = excess molesH+  Vtotal
[H+] = [HCl]
pure water
strategy
[H+] = (nA – nB/Vtot
post equivalence: [OH-] > [H+]
H+ + OH-excess  OH- + H2O
[OH-] = excess molesOH-  Vtotal
[OH-] = (nB – nA)/Vtot
[H+]E or [OH-]E
pH = -log[H+]
pH = 14 - pOH = 14 + log[OH-]
pH
Determine the theoretical pH when 0.10 M NaOH is added to 0.10 M HC2H3O2 (Ka = 1.8 x 10-5: Kb = Kw/Ka = 5.6 x 10-10).
HA + OH- 
initial reaction
HAexcess + OH-  HA + A- + H2O
A- + H2O
nHA = nA – nB
moles HA (nHA)
i.
nA- = nB
moles A- (nA-)
equilibrium
reaction
calculation
strategy
HA  H+ + Apure acid
[H+] = (Ka[HA])½
buffer
[H+] = Ka(nHA/nA-)
[]=
A- + H2O 
HA + OHpure base
[OH-] = (Kb[A-])½
same as above for [OH-]
[H+]E or [OH-]E
pH = -log[H+]
pH = 14 - pOH = 14 + log[OH-]
pH
j.
Determine the theoretical pH when 0.10 M HCl is added to 0.10 M NH3 (Kb = 1.8 x 10-5: Ka = Kw/Kb = 5.6 x 10-10).
Bexcess + H+  B + HB+ + H2O B + H+  HB+
initial reaction
nNH3 = nA – nB
moles remaining
nNH4+ = nB
moles formed
[]=
equilibrium
reaction
calculation
strategy
B + H2O  HB+ + OHpure base
[OH-] = (Kb[NH3])½
buffer
[OH-] = Kb(nNH3/nNH4+)
HB+  H+ + B
pure acid
[H+] = (Ka[NH4+])½
same as above except [H+] instead of [OH-]
[H+]E or [OH-]E
pH = 14 - pOH = 14 + log[OH-]
pH
pH = -log[H+]
12
12
12
10
10
10
8
8
8
pH
(3) HCl added to NH3.
pH
Graph the theoretical pH vs. mL of solution added calculated in parts g, h and i below.
(1) NaOH added to HCl.
(2) NaOH added to HC2H3O2.
pH
k.
6
6
6
4
4
4
2
2
2
0
l.
20
40
0
20
40
0
20
40
mL of 0.10 M NaOH added
mL of 0.10 M NaOH added
mL of 0.10 M HCl added
Compare the theoretical pH profiles with the actual pH profiles and complete the following chart.
Best
Comparison between Graphs
Experiment
Indicator
Equivalence pH
Buffer Region
Overall Shape
SA + SB (HCl + NaOH)
WA + SB (HC2H3O2 + NaOH)
WB + SA (NH3 + HCl)
Solubility Product Constant Lab—Determine the [CrO42-] of
saturated Ag2CrO4 solution by spectrophotometry, calculate
Ksp for Ag2CrO4 and compare it to the expected value.
Fill five cuvette tubes with the K2CrO4 standards. Set the
spectrophotometer to 390 nm and measure absorbance.
a. Calculate the CrO42- concentration for each standard
and record its absorbance.
Volume (drops)
[CrO42-]
Absorbance
1.0 x 10-4 M
(mol/L)
Total
K2CrO4
0
80
0
0
b.
20
80
40
80
60
80
80
0
Graph the absorbance vs. [CrO42-].
in the slot opposite your tube). Pour off the liquid, called
the supernatant, while leaving the precipitate in the test
tube (decant). Add 5 mL of distilled water to the test tube.
Stopper, shake for 5 minutes, centrifuge for 5 minutes.
Pipet the clear, pale yellow supernatant to a cuvette (be
careful not to include any precipitate). Measure the
absorbance of the supernatant. Discard the remaining
supernatant in the test tube, add 5 mL distilled water,
stopper, shake, centrifuge and measure the absorbance of
the pipetted supernatant. Repeat until the absorbance
reading is within 0.02 of the previous reading.
d. Record the absorbance, use Beer's law to calculate
[CrO42-], which is the solubility, and then calculate Ksp.
Ksp
Absorbance
[CrO42-]
e.
The actual Ksp for Ag2CrO4 is 1.1 x 10-12. Use this
value and work backward through the calculations to
determine what the absorbance should have been.
f.
Will the following produce an absorbance that is
greater than expected or less? Explain you answer.
(1) Dirt or fingerprints on the cuvette.
0.50
Absorbance
4.
0.40
0.30
0.20
0.10
0
c.
2.5
5.0
7.5
Concentration ( x 10-5 mol/L)
Determine the slope of the line, which equals a in Beer's
law. A (Absorbance) = a (absorptivity) • b (cuvette width—
1 cm) • c (concentration).
To a clean test tube, add 5 mL distilled water, 5 drops of
0.1 M AgNO3 and 15 drops of 0.1 M K2CrO4. Stopper the
tube and shake periodically for 5 minutes. Centrifuge for 5
minutes (be sure a test tube with similar amount of liquid is
(2) The original chromate was not completely washed
from the Ag2CrO4 precipitate.
(3) Some precipitated Ag2CrO4 was included with the
supernatant.
(4) The concentration of K2CrO4 used for the chromate
standards was greater than 1.0 x 10-4 M.
10. Complete the following:
H2O(l) 
Practice Problems
1.
A. Acid and Base
Define acid and base, and write a general acid-base
reaction for each of the three theories.
Theories
Acid
Base
Reaction
[H+] x [OH-] = 1 x 10___  [H+] = [OH-] = 1 x 10___ M
acid solutions [H+] __ [OH-], base solutions [H+] __ [OH-]
pH = _______, pOH = _______and pH + pOH = ___
11. fill in the range of values for each solution.
Solution
Acid
Neutral
Base
pH
Arrhenius
BrønstedLowry
[H+]
Lewis
2.
H+ + H2O(l) 
Complete the chart for a 1 M solution of acid.
Ionization reaction % ionization [HX]
[H+]
100 %
HCl(g) 
HF(g) 
[X-]
8%
HCHO2 
4%
3. Rank the following acids from strongest to weakest:
Acid
HBrO
HClO
HClO2
HIO
Rank
4. Complete the chemical equation and write the Ka
expression for the following weak acids.
Ka expression
Equation
[OH-]
12. Solve for the missing values.
pH
[H+]
4.20
[OH-]
pOH
3.0 x 10-9
13. Given the following:
HC2H3O2(aq)  H+ + C2H3O2HCN(aq)  H+ + CNa. Calculate Kb for C2H3O2- and CN-
Ka = 1.8 x 10-5
Ka = 4.0 x 10-10
C2H3O2CN-
HC2H3O2(aq) 
b.
H2PO4- 
H2CO3(aq) 
HCO3-(aq) 
5.
Ka1 for H2SO3 is 1.3 x 10-2 and Ka2 is 6.3 x 10-8. Write each
dissociation equation and the overall equation. Calculate
K for the reaction H2SO3  2 H+ + SO32-.
Highlight which is stronger.
HC2H3O2 or HCN
C2H3O2- or CNc. Complete the statement about the relative strength of
the acid and base of a conjugate pair.
A strong conjugate acid makes a _______ conjugate base.
d. What is K for the equilibrium below?
HC2H3O2(ag) + CN-(aq)  C2H3O2-(aq) + HCN(aq)
e.
6.
Complete the chemical equation and write the Kb
expression for the following weak bases.
Kb expression
Equation
F-(aq) + H2O 
CH3NH2(aq) + H2O 
7.
Classify the salts as acidic, basic, neutral or can't tell.
NaCl
Cu(NO3)2
KNO2
NH4F
NaClO
8.
BaCl2
Cu(C2H3O2)2
LiF
Write a net ionic equation for each acid-base reaction (SA
= strong acid, SB = strong base, WA = weak acid, WB =
weak base).
SA + SB
HNO3 + NaOH
WA + SB
HF(aq) + Sr(OH)2
SA + WB
H2SO4 + NH3(aq)
SA + WB HCl + Na2S(aq)
9. Explain the observations using chemical equations.
a. Statues made of marble, CaCO3(s), exposed to polluted
(acidic) air lose their definition over time.
b.
Milk of Magnesia is a medication that absorbs excess
stomach acid contains Mg(OH)2(s).
Starting with all reactants and products at 1 M, which
way will the reaction proceed to reach equilibrium?
f.
Complete the statement about the relative strength of
the acids and bases in a Brønsted-Lowry system.
______ acid + ______ base  ______ acid + ______base
14. Kb for NH3 is 1.8 x 10-5. Determine
a. Ka for NH4+.
b.
K for the reactions.
Reaction
NH3(aq) + H2O  NH4+ + OH-
K
NH4+ + OH-  NH3(aq) + H2O
NH3(aq) + H+  NH4+
NH4+  H+ + NH3(aq)
15. Ka for HF is 6.9 x 10-4. Determine
a. Kb for F-:
b.
K for the reactions.
Reaction
HF(aq)  H+ + FH+ + F-  HF(aq)
HF(aq) + OH-  F- + H2O
F- + H2O  HF(aq) + OH-
K
16. Complete the equation, label the acids (A) and bases (B),
and link the conjugate pairs.
NH4+ + H2O 
NO2- + H2O 
HNO2 + H2O 
HNO2 + NH3 
NH3 + H2O 
NH4+ + OH- 
HNO2 + OH- 
NH3 + H3O+ 
21. 0.50 mol of phenol (HOC6H5) in 5.0 L has Ka = 1.6 x 10-10.
a. Determine [H+].
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B. Acid-Base Equilibrium
17. 1.369 g of HClO2 is dissolved in enough water to make
100. mL of solution. The pH is 1.36. Determine
a. Initial concentration of HClO2.
b.
Equilibrium concentration of H+.
E
b.
Determine pH.
22. For H2SO3, Ka1 = 1.3 x 10-2 and Ka2 = 6.3 x 10-8.
a. Calculate [H+] for 0.100 M H2SO3.
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c. Ka.
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b. What are the concentrations of HSO3- and SO32-?
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18. In the first-step ionization of phosphorous acid the acid is
33.3% dissociated in a 0.300 M solution of H3PO3.
Calculate Ka for the first-step ionization of H3PO3.
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23. For phosphoric acid (H3PO4) Ka1 = 7.1 x 10-3, Ka2 = 6.2 x 10-8
and Ka3 = 4.5 x 10-13. Determine for 0.100 M H3PO4.
a. [H+] and [H2PO4-]
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19.
HC2O4-  H+ + C2O427.35% of the HC2O4- is dissociated in a 0.0100 M solution.
Calculate Ka for the second-step dissociation of oxalic acid.
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b. [HPO42-]
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20. Acetic Acid (HC2H3O2) has Ka = 1.8 x
a. The [H+] in 0.100 M solution.
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10-5.
Determine
C
E
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c. [PO43-]
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b.
The percent ionization.
% = [H+]/[HA]o x 100 =
c.
Was the assumption that the equilibrium concentration
of HC2H3O2 is 0.100 M valid? Explain your answer.
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24. 0.10 M methylamine, CH3NH2, has Kb = 5.0 x 10-4.
Determine [OH-]E.
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29. A liter of benzoic acid (HBen) and sodium benzoate (Ben-)
buffer is prepared using 1.00 mol of benzoic acid and 0.50
mol of sodium benzoate. The Ka for HBen is 6.3 x 10-5.
a. 0.100 mol of H+ is added. Calculate
I
mol HBen
C
mol Ben-
E
[H+]E
pH
25. 6.6 g of hydroxylamine, HONH2, are in a 1-liter solution.
a. Determine the initial concentration of HONH2.
b.
0.100 mol of OH- is added. Calculate
mol HBen
mol Ben-
b. Determine [OH-]E given Kb = 9.1 x 10-9
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pH
C
30. A buffer is prepared by adding 2.00 mol of HC2H3O2 and
2.00 mol of NaC2H3O2 to enough water to make 5.00 L of
solution. Acetic acid has a Ka = 1.8 x 10-5. Calculate
a. the pH of the buffer.
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26. 4.00 g of NaF are in 0.500 L solution.
a. Determine the initial concentration of NaF.
b.
b.
Determine Kb for
F-
given that Ka for HF is 6.7 x
10-4.
c. Determine the equilibrium concentration of OH-.
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the pH when 0.50 mol of H+ is added.
mol HC2H3O2
mol C2H3O2[H+]E
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pH
E
c.
the pH when 0.50 mol of OH- is added.
mol HC2H3O2
d.
Determine the pH of the solution.
27. 2.70 g of HCN (Ka = 4.0 x 10-10) and 2.45 g of NaCN are
added to water to make 1.00 L of solution
a, What are the initial concentrations of HCN and CN-?
b. What is the pH of the solution?
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pH
C. Acid-Base Titration
31. Determine the pH after the following aliquots of 0.10 M
NaOH are added to 20 mL of 0.1 M HF (Ka = 6.7 x 10-4).
mL NaOH
pH
C
0 mL
E
10 mL
20 mL
28. Calculate [H+] for a solution that is 0.10 M in both H2CO3
and NaHCO3. The Ka for H2CO3 is 4.2 x 10-7.
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30 mL
D. Solubility Equilibrium
32. A saturated solution of Ba3(PO4)2 (Ksp = 6 x 10-39) has a
[Ba2+] = 5 x 10-4 M. Calculate [PO43-].
33. A saturated solution of PbSO4 (Ksp = 1.8 x 10-8) has a
[SO42-] = 2 x 10-4 M. Calculate [Pb2+].
34. A solution contains [Ba2+] = 0.0040 M and [Pb2+] = 0.0060 M.
What concentration of F- will just precipitate one of the ions?
Which one will precipitate first?
(BaF2 Ksp = 1.8 x 10-7, PbF2 Ksp = 7.1 x 10-7)
35. Consider Ag3PO4 (Ksp = 1.0 x 10-16).
a. What is the solubility in pure water?
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41. Will the precipitate, Al(OH)3 (Ksp = 2 x 10-31), form when
200. mL of 1 x 10-6 M of Al(NO3)3 is mixed with 300. mL of
5 x 10-6 M of Ba(OH)2? Determine
[Al3+]
[OH-]
Q
ppt?
42. Will the precipitate, Ca3(PO4)2 (Ksp = 1 x 10-33), form when
250 mL of 0.40 M of ammonium phosphate is mixed with
450 mL of 0.125 M of calcium chloride?
[Ca2+]
[PO43-]
Q
b. What is the solubility in 0.0010 M Na3PO4?
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36. Consider Ag2CrO4 (Ksp = 1 x 10-12).
a. What is the solubility in pure water?
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ppt?
43. The initial concentrations are [Mg2+] = [Sr2+] = 0.02 M and
[CO32-] = 2 x 10-7 M. Will a precipitate form, and if so what
is it? (MgCO3 Ksp = 7 x 10-6, SrCO3 Ksp = 6 x 10-10)
SrCO3
MgCO3
C. Factors that Affect Solubility
44. What does the difference between the two answers from
questions (8.a) and (8.b) illustrate?
45. Complete the complex ion synthesis reactions and write a
Kf expression.

b.
How many grams are dissolved in water to make one
liter of solution?
c. What is the solubility in 0.100 M K2CrO4?
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37. The solubility of AgCl is 1.3 x 10-5 M. What is Ksp?
38. The solubility of LaF3 is 9.3 x 10-6 M. What is Ksp?
39. 500. mL of a saturated solution contains 0.0651 g of MgF2
at 25oC.
a. What is the solubility in mol/L?
b.
Al3+ + 4 OH- 
Cu2+ + 4 Cl- 
Fe3+ + SCN- 
46. Given the following equilibriums with their constants.
BaF2(s)  Ba2+ + 2 FKsp = 1.8 x 10-7
3+
Cr + 4 OH  Cr(OH)4
Kf = 8.0 x 1029
Cr(OH)3(s)  Cr3+ + 3 OHKsp = 1.6 x 10-30
2+
2+
Cu + 4 NH3(aq)  Cu(NH3)4
Kf = 5.0 x 1012
Cu(OH)2(s)  Cu2+ + 2 OHKsp = 4.8 x 10-20
HF(aq)  H+ + FKa = 6.8 x 10-4
+
H2O(l)  H + OH
Kw = 1.0 x 10-14
H2S(aq)  2 H+ + S2K = 1.0 x 10-20
2+
2MnS(s)  Mn + S
Ksp = 2.5 x 10-13
Calculate K for the equilibriums below.
a. MnS(s) + 2 H+  Mn2+ + H2S(aq)
b.
Cu(OH)2(s) + 4 NH3(aq)  Cu(NH3)42+ + 2 OH-.
c.
BaF2(s) + 2 H+  Ba2+ + 2 HF(l)
d.
Cr(OH)3(s) + 3 H+  Cr3+ + 3 H2O(l)
What is Ksp?
40. The solubility of BaC2O4 is 22 mg/L.
a. What is the solubility in mol/L?
b.
Ag+ + 2 NH3 
What is Ksp?
e.
Cr(OH)3(s) + OH-  Cr(OH)4-
f.
What property of Cr(OH)3 is illustrated by the answers
from parts (d) and (e)?
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
1. Which is the net ionic equation for the neutralization reaction
between barium hydroxide and ammonium chloride?
(A) H+ + OH-  H2O
(C) OH- + NH4+  NH4OH
+
+
(B) H + NH3  NH4
(D) OH- + NH4+  NH3 + H2O
2.
3.
4.
Which 1-M solution is the most basic?
(A) NaNO3 (B) Na2CO3 (C) NaCI
11. When phenolphthalein is used as the indicator in a titration
of HCl with NaOH, it undergoes a rapid color change from
clear to red at the end point of the titration because
(A) phenolphthalein is a very strong acid that is capable of
rapid dissociation
(B) the solution being titrated undergoes a large pH
change near the equivalence point of the titration
(C) phenolphthalein undergoes an irreversible reaction in
basic solution
(D) OH- is the catalyst for phenolphthalein decomposition
(D) NaHSO4
Which is generally true as the number of oxygen atoms
increases in a series of acids, such as HXO, HXO2, HXO3?
(A) The acid strength decreases only if X is a nonmetal.
(B) The acid strength decreases only if X is a metal.
(C) The acid strength decreases.
(D) The acid strength increases.
HSO4– + H2O  H3O+ + SO42–
In the equilibrium represented above, the species that act
as bases include which of the following?
I. HSO4II. H2O
III. SO42(A) II only (B) II only (C) III only (D) II and III
5.
HC2H3O2 + CN-  HCN + C2H3O2–
The equilibrium constant, K = 3.7 x 104. Which can be
concluded from this information?
(A) CN- is a stronger base than C2H3O2-.
(B) HCN is a stronger acid than HC2H3O2.
(C) The conjugate base of CN- is C2H3O2-.
(D) K will increase with an increase in temperature.
6.
Equal volumes of 0.10 M H3PO4 and 0.20 M KOH are
mixed. After equilibrium is established, the type of ion in
solution in largest concentration, other than the K+ ion, is
(A) H2PO4- (B) HPO42- (C) PO43(D) OH-
7.
Which of the following species is in the greatest
concentration in a 0.100 M solution of H2SO4 in water?
(A) H2SO4 (B) H3O+
(C) HSO4– (D) SO42–
8.
Which of the following reactions does NOT proceed
significantly to the right in aqueous solutions?
(A) H3O+ + OH–  2 H2O
(B) HCN + OH–  H2O + CN–
(C) H2SO4 + H2O  H3O+ + HSO4–
(D) H2O + HSO4–  H2SO4 + OH–
9.
10. A molecule or an ion is classified as a Lewis acid if it
(A) accepts a proton from water
(B) accepts a pair of electrons to form a bond
(C) donates a pair of electrons to form a bond
(D) donates a proton to water
Which of the following is NOT amphiprotic?
(A) HCO3- (B) H2PO4- (C) NH4+
(D) H2O
12. The pH of 0.1 M ammonia is approximately
(A) 1
(B) 7
(C) 11
(D) 14
13. What is NH4+ in the reaction: 2 NH3  NH4+ + NH2-?
(A) a catalyst
(B) both an acid and a base
(C) the conjugate acid of NH3
(D) the reducing agent
14. At 25°C, an aqueous solutions with a pH of 8 has a [OH-] of
(A) 10-14 M (B) 10-8 M (C) 10-6 M (D) 1 M
15. In the titration of a weak acid of unknown concentration
with a standard solution of a strong base, a pH meter was
used to follow the progress of the titration. Which of the
following is true for this experiment?
(A) The pH is 7 at the equivalence point.
(B) The pH at equivalence depends on the indicator used.
(C) The graph of pH versus volume of base added rises
gradually at first and then much more rapidly.
(D) The graph of pH versus volume of base added shows
no sharp rise.
16. How can 100. mL of NaOH solution with a pH of 13 be
converted to a NaOH solution with a pH of 12?
(A) Add 10 mL of distilled water to the 100 mL of NaOH.
(B) Add 100 mL of distilled water to the 100 mL of NaOH.
(C) Add 900 mL of distilled water to the 100 mL of NaOH.
(D) Add 100. mL of 0.10 M HCI to the 100 mL of NaOH.
17. A 0.2 M solution of a weak monoprotic acid, HA, has a pH
of 3. The ionization constant of this acid is
(A) 5 x 10-7 (B) 2 x 10-7 (C) 5 x 10-6 (D) 5 x 10-3
Questions 18-19 Oxalic acid, H2C2O4, is a diprotic acid with
K1 = 5 x 10-2 and K2 = 5 x 10-5.
18. Which equals the equilibrium constant for the reaction:
H2C2O4 + 2 H2O  2 H3O+ + C2O42–?
(A) 5 x 10-2 (B) 5 x 10-5 (C) 25 x 10-7 (D) 5 x 10-7
19. Which species is in highest concentration in 0.1 M H2C2O4?
(A) H2C2O4 (B) H3O+
(C) HC2O4- (D) C2O42-
27. The most nearly neutral solution
28. A buffer at a pH > 8
20.
Acid
Acid Dissociation Constant, Ka
H3PO4
7 x 10-3
–
H2PO4
8 x 10-8
HPO42–
5 x 10-13
On the basis of the information above, a buffer with a
pH = 9 can best be made by using
(A) H3PO4 + H2PO4–
(B) H2PO4- + PO42–
2–
(C) H2PO4 + HPO4
(D) HPO42- + PO43–
29. A buffer at a pH < 6
30.
MnS(s) + 2 H+  Mn2+ + H2S(g)
Ksp, for MnS in 5 x 10-15 and K for the dissociation of H2S is
1 x 10-20. What is the equilibrium constant for the reaction?
(A) 5 x 10-15
(B) 5 x 10-8
-6
(C) 2 x 10
(D) 5 x 105
Questions 21-22 The graph shows the titration curve when 100.
mL of 0.0250 M acetic acid is titrated with 0.100 M NaOH.
31. The solubility of CuI is 2 x 10-6 M. What is the solubility
product constant, Ksp, for CuI?
(A) 1.4 x 10-3
(B) 2 x 10-6
(C) 4 x 10-12
(D) 2 x 10-12
32. What is the solubility of Ag2CrO4 (Ksp = 8 x 10-12)?
(A) 8 x 10-12 M
(B) 2 x 10-12 M
(C) (4 x 10-12)½ M
(D) (2 x 10-12)⅓ M
21. Which of the following indicators is the best for this titration?
Indicator
pH Range of Color Change
(A) Methyl orange
3.2 - 4.4
(B) Methyl red
4.8 - 6.0
(C) Bromthymol blue
6.1 - 7.6
(D) Phenolphthalein
8.2 - 10.0
22. What part of the curve corresponds to the optimum buffer
action for the acetic acid/acetate ion pair?
(A) Point V (B) Point X (C) Point Z (D) Along WY
Questions 23-24 Ka for HCN is 5.0 x 10-10.
23. What is the H+ concentration in 0.05 M HCN?
(A) 2.5 x 10-11 M
(B) 2.5 x 10-10 M
-10
(C) 5.0 x 10 M
(D) 5.0 x 10-6 M
33. How many moles of NaF must be dissolved in 1 L of a
saturated solution of PbF2 at 25oC to reduce the [Pb2+] to
1 x 10-6 M? (Ksp PbF2 at 25oC = 4 x 10-8)
(A) 0.02
(B) 0.04
(C) 0.1
(D) 0.2
34. A solution containing cations was treated with 0.1 M HCl.
The white precipitate formed was filtered and washed with
hot water. A few drops of 0.1 M K2CrO4 were added to the
hot water filtrate and a bright yellow precipitate was
produced. The white precipitate remaining on the filter
paper was readily soluble in ammonia solution. What two
ions could have been present in the unknown?
(A) Ag+ and Hg22+
(B) Ag+ and Pb2+
(C) Ba2+ and Ag+
(D) Ba2+ and Hg22+
Practice Free Response
24. What is the pH of a 0.02 M solution of HCN?
(A) Between 7 and 10
(B) 7
(C) Between 4 and 7
(D) 4
1.
b.
25. If the acid dissociation constant Ka for an acid HA is
8 x 10-4 at 25oC, what percent of the acid is dissociated in
a 0.5 M solution of HA at 25oC?
(A) 0.08% (B) 0.2%
(C) 1%
(D) 4%
2.
Questions 26-29 Assume all solutions are 1 M.
(A) NH3 and NH4CI
(B) H3PO4 and NaH2PO4
(C) HCI and NaCI
(D) NH3 and HC2H3O2 (acetic acid)
26. The solution with a pH = 0 and is not a buffer
Excess nitric acid is added to solid calcium carbonate.
a. Write a net ionic balanced equation:
Briefly explain why statues made of marble (calcium
carbonate) displayed outdoors in urban areas are
deteriorating.
25 mL of 0.40 M HF (Ka = 7.2 x 10-4) reacts with 15 mL of
0.40 M NaOH according to the reaction.
HF(aq) + OH-(aq)  H2O(l) + F-(aq)
a. Calculate the moles of HF remaining.
b.
Calculate the initial concentration of HF.
c.
Calculate the initial concentration of F-.
d. Calculate the equilibrium concentration of H+.
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e. Calculate the pH.
3.
Propanoic acid, HC3H5O2, ionizes in water:
HC3H5O2(aq)  C3H5O2-(aq) + H+(aq) Ka = 1.34 x 10-5
a. Calculate the pH of 0.265 M propanoic acid.
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b. 0.496 g of NaC3H5O2 is added to a 50.0 mL of 0.265 M
propanoic acid. Calculate
(1) the initial concentration of C3H5O2-.
5.
(2) the equilibrium concentrations of H+.
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Methanoic acid forms when methanoate reacts with water:
CHO2-(aq) + H2O(l)  HCHO2(aq) + OH-(aq)
c. Given that [OH-]E is 4.18 x 10-6 M in a 0.309 M solution
of sodium methanoate, calculate
(1) Kb for CHO2-.
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(2) Ka for HCHO2.
d.
4.
f.
A solution contains [Ca2+] = 4.0 x 10-3 M and [Mg2+] =
1.0 x 10-4 M. What concentration of OH- will just
precipitate one of the ions? Which one will precipitate
first? Justify your answer. (Mg(OH)2 Ksp = 6.0 x 10-12)
g.
Calculate K for: Ca(OH)2(s) + 2 H+  Ca2+ + 2 H2O.
(H2O  H+ + OH- Kw = 1.0 x 10-14)
h.
Is Ca(OH)2 soluble in HF? Justify your answer.
(HF(aq)  H+ + F- Ka = 6.8 x 10-4)
i.
At a different temperature, the solubility of Ca(OH)2 is
0.15 g/L. What is the Ksp at the different temperature?
AgBr is a cream-colored salt with Ksp, = 5.0 x 10−13 at 298 K.
a. What is [Ag+] in 50 mL of a saturated AgBr(aq) at 298K?
b.
50 mL water is added to the solution in part (b), which
contains solid AgBr at the bottom. Equilibrium is
reestablished and solid AgBr remains. Is the value of
[Ag+] greater than, less than, or equal to the value you
calculated in part (b)? Justify your answer.
c.
Calculate the minimum volume of water needed to
completely dissolve 5.0 g of AgBr(s) at 298 K.
d.
10.0 mL of 1.5 x 10−4 M AgNO3 is mixed with 2.0 mL
of 5.0 x 10−4 M NaBr. Will a precipitate form? Justify
your answer with calculations.
e.
The color of another salt of silver, AgI(s), is yellow. A
student adds a solution of NaI to a test tube
containing a small amount of solid, cream-colored
AgBr. After stirring the contents of the test tube, the
student observes that the solid in the test tube
changes color from cream to yellow.
(1) Write the chemical equation for the reaction that
occurred in the test tube.
Which acid is stronger, propanoic acid or methanoic
acid? Justify your answer.
Answer the questions that relate to Ca(OH)2 Ksp is 1.8 x 10-11.
a. A saturated solution has a [OH-] = 2.0 x 10-4 M.
Calculate [Ca2+].
b.
Calculate the solubility.
c.
Calculate the mass of Ca(OH)2 that can be dissolved
in water to make 1.50 L of solution?
d.
Calculate the solubility of Ca(OH)2 in 1.0 x 10-3 M NaOH.
e.
Will Ca(OH)2 precipitate when 250 mL of 0.0040 M of
sodium hydroxide is mixed with 450 mL of 0.0125 M of
calcium chloride? Justify your answer.
(2) Which salt has the greater value of Ksp : AgBr or
AgI? Justify your answer.