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AP Chemistry 11: Equilibrium—Salts and Complex Ions Name _________________________
A.
Ionic Compounds (2.7 to 2.8, 4.2)
1. name and formula
a. cations
1. named the same as metal (Na+ = sodium)
2. polyatomic cations with –ium ending
NH4+ (ammonium), H3O+ (hydronium)
3. charge based on periodic table position
4. assume all transition metals are 2+
a. use Roman numeral for multiple cations
per atom: (Fe2+ -iron(II), Fe3+ -iron(III))
b. important exceptions: Cr3+, Ag+
b. anions
1. –ide ending
a. monatomic: Cl- (chloride)
b. binary: OH- (hydroxide), CN- (cyanide)
2. oxyanion (–ate and –ite ending)
C2H3O2acetate
MnO4permanganate
CO32carbonate
NO3nitrate
CrO42chromate
PO43phosphate
Cr2O72dichromate
SO42sulfate
a. use "bi" prefix when H is added to anion:
HCO3- (bicarbonate). HSO4- (bisulfate)
b. use "ite" suffix when one O is removed:
SO32- (sulfite), NO2- (nitrite)
c. when non-oxygen atom is a halogen:
ClO4ClO3ClO2ClOperchlorate
chlorate
chlorite
hypochorite
c. empirical formula (criss-cross method)
1. charges become opposite's subscript
2. Al3+ and SO42-  Al2(SO4)3
2. predicting solubility
a. solubility guidelines for predicting MX(s) (memorize)
Anions (X)
Cations
(M)
NO3- Cl-, Br-, I- SO42- OH-, S2- Others
Alkali Metal
S
S
S
S
S
NH4+
S
S
S
S
S
Sr2+,Ba2+
S
S
I
S
I
Ag+
S
I
S
I
I
Hg22+, Pb2+
S
I
I
I
I
Others
S
S
S
I
I
b. precipitation reaction (ion exchange)
1. MX(aq) + M'X'(aq)  MX'(s) + M'X(aq)
2. net ionic: M+ + X'-  MX'(s)
B. 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
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. salt are more soluble in solution with H+
b. Le Chatelier's principle: MX(s) + H+  M+ + HX(aq)
(higher [ ] of reactant shifts equilibrium right)
CaCO3(s) + 2 H+  Ca2+ + H2CO3(aq)  CO2(g) + H2O
Mg(OH)2(s) + 2 H+  Mg2+ + 2 H2O
CuS(s) + 2 H+  Cu2+ + H2S(g)
c. salts with Cl-, Br-, I-, SO42- are unaffected by H+
3. formation of complex ions
a. polar molecules or anions are strongly attracted to
transition metal cation forming a complex ion
1. polar molecule or anion is called ligand
a. molecules: H2O, NH3
b. anions: Cl-, CN-, SCN-, OH2. number of ligands (coordination number) is
usually 2, 4 or 6 (often 2 x cation charge)
3. [Cu(NH3)4]SO4—tetraaminecopper(II) sulfate
a. complex ion is placed in brackets
b. counter ion is outside brackets
c. compound is electrically neutral
b. formation constant, Kf
Cu2+ + 4 CN-  Cu(CN)42Kf = [Cu(CN)42-]/[Cu2+][CN-]4
(note: compared to Ksp, Kf is backward and
complex is included)
4. determine equilibrium position
example: Will AgCl dissolve in concentrated NH3?
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 + Cl K = 3.1 x 10-3
answer: K < 1  AgCl is only soluble if [NH3] is high
enough to shift equilibrium toward products.
5. amphoterism: some metal oxides and hydroxides are
soluble in both strong acid and strong base
example Al(OH)3(s)
1. acid: Al(OH)3  Al3+ + 3 OH- Ksp = 2 x 10-31
3 H+ + 3 OH-  3 H2O
Kw-3 = 1 x 1042
+
3+
Al(OH)3 + 3 H  Al + 3 H2O K = 2 x 1011
answer: K > 1  soluble in strong acid (H+)
2. base: Al(OH)3  Al3+ + 3 OH- Ksp = 2 x 10-31
Al3+ + 4 OH-  Al(OH)4Kf = 1 x 1033
Al(OH)3 + OH  Al(OH)4
K = 2 x 102
answer: K > 1  soluble in strong base (OH-)
c.
C.
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 10 minutes. Centrifuge for 3
minutes (be sure a test tube with similar amount of liquid is
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 0.1 M NaNO3 to the test tube.
Repeat the stopper, shake, centrifuge, decant and wash
cycle. Centrifuge. Pipet the clear, pale yellow supernatant
to a cuvette (be careful not to include any precipitate).
Measure the absorbance of the solution.
d. Record the absorbance, use Beer's law to calculate
[CrO42-], which is the solubility, and then calculate Ksp.
Absorbance
[CrO42-]
Ksp
Experiments
1.
Solubility Lab—Mix cations and anions together to
determine solubility and compare results to solubility table.
a. Fill in the formula of any expected precipitate.
NO3SO42S2OHPO43I+
Na
NH4+
Ba2+
Ag+
Pb2+
Al3+
Add 1 drop of each cation (nitrate salt) along a row.Add 1
drop of each anion (sodium salt) along a column. Write the
color for any precipitate that forms.
b. Fill in the color of any precipitate that forms.
NO3SO42S2OHPO43I+
Na
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.
Pb2+
Will the following produce an experimental result that
is greater than expected or less? Explain you answer.
(1) Dirt or fingerprints on the cuvette.
Al3+
c. Are there any precipitates that you expected to form that
didn't or that didn't form when you expected them to?
(2) The original chromate was not completely washed
from the Ag2CrO4 precipitate.
NH4+
Ba2+
Ag+
Solubility Product Constant Lab—Determine the [CrO42-] of
saturated AgCrO4 solution by spectrophotometry, calculate
Ksp for AgCrO4 and compare it to the expected value.
Fill three cuvette tubes with the K2CrO4 standards. Set the
spectrophotometer to 375 nm and measure absorbance.
a. Calculate the CrO42- concentration for each standard
and record its absorbance.
Volume (mL)
[CrO42-]
Absorbance
1.0 x 10-3 M
(mol/L)
Total
K2CrO4
0
20
0
0
b.
1.0
20.
2.0
20.
3.0
20.
Graph the absorbance vs. [CrO42-].
0.50
Absorbance
2.
0.40
0.30
0.20
0.10
0
c.
0.50
1.0
1.5
Concentration ( x 10-4 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).
(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 10-3 M.
(5) The concentration of K2CrO4 used for the
chromate standards was less than 1.0 x 10-3 M.
3.
Qualitative Analysis Lab—Perform a series of procedures
to isolate and test for the presence of one or more of the
"group 1" cations (Ag+, Pb2+ and Hg22+) in an unknown.
Step 1: Add 2 drops of 6 M HCl to 20 drops of the known
solution in a small test tube. Place in a boiling water bath
(250 mL beaker half filled with tap water) for 2 minutes and
stir occasionally with a glass rod. Centrifuge the hot
solution and quickly decant. Save the supernatant for step
2 and the precipitate for step 3.
Step 2: Add 1 drop of 6 M HC2H3O2 and 2 drops of 0.1 M
K2CrO4 to the supernatant from step 1. A yellow
precipitate confirms the presence of Pb2+.
Step 3: Add 10 drops of 6 M NH3 to the precipitate from
step 1 and stir thoroughly. Centrifuge and decant. Save
the supernatant for step 4. A gray or black precipitate
confirms the presence of Hg22+.
Step 4: Add 6 M HNO3 to the supernatant from step 3 until
it is acidic toward litmus paper (dip the stirring rod into the
supernatant then place it on blue litmus paper—turns red
in acid). A white precipitate confirms the presence of Ag+.
Repeat steps 1-4 with the unknown and record your results.
a. What do Ag+, Pb2+ and Hg2+ have in common?
b.
How is Pb2+ separated from AgCl(s) and Hg2Cl2(s)?
c.
How is the presence of Pb2+ confirmed?
d.
Ag+ separated from Hg2Cl2(s)?
e.
How is the presence of Hg22+ confirmed?
f.
How is the presence of Ag+ confirmed?
g. Which ions are present in your unknown?
Pb2+
Ag+
Hg22+
Practice Problems
1.
2.
A. Ionic Compounds
Fill in the charge of each cation and anion, and then
complete the formula and name of the ionic compound.
Cation
Anion
Formula
Name
Fe(II)2+
O2Ca2+
H-
Al3+
Br-
Ag+
S2-
Zn2+
I-
Pb(IV)4+
PO43-
Cu(II)2+
SO42-
Cr3+
NO3-
Co(III)3+
Cl-
Sn(IV)4+
MnO4-
Hg2(I)2+
CO32-
Na+
Cr2O72-
Mg2+
C2H3O2-
Ni2+
CrO42-
CO32K+
Write formulas for the cation, anion and compound named.
Name
Cation
Anion
Formula
Ammonium carbonate
3.
Write the formula for the oxyanions of iodine.
IO4IO3IO2-
4.
For each pair of salts, cross out the spectator ions and
then write a net ionic equation for the precipitation reaction.
sodium carbonate +
barium nitrate
mercury(I) sulfate +
ammonium chloride
magnesium nitrate +
sodium hydroxide
lead(II) nitrate +
potassium bromide
B. Solubility Equilibrium
5. A saturated solution of Ba3(PO4)2 (Ksp = 6 x 10-39) has a
[Ba2+] = 5 x 10-4 M. Calculate [PO43-].
6.
A saturated solution of PbSO4 (Ksp = 1.8 x 10-8) has a
[SO42-] = 2 x 10-4 M. Calculate [Pb2+].
7.
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)
8.
Consider Ag3PO4 (Ksp = 1.0 x 10-16).
a. What is the solubility in pure water?
[]
I
C
E
b. What is the solubility in 0.0010 M Na3PO4?
[]
I
C
E
Consider Ag2CrO4 (Ksp = 1 x 10-12).
a. What is the solubility in pure water?
[]
I
C
E
9.
Aluminum acetate
Silver sulfate
Lead(IV) chlorite
Copper(II) nitrite
IO-

b.
How many grams are dissolved in water to make one
liter of solution?
c.
How many grams dissolve in 0.100 M K2CrO4 to
make 100 mL of solution?
Sodium cyanide
Copper(I) fluoride
Manganese(IV) oxide
Iron(III) sulfite
Magnesium dichromate
Mercury(II) hydroxide
Potassium phosphate
Sodium bicarbonate
Tin(II) oxalate
[]
I
C
E
10. The solubility of AgCl is 1.3 x 10-5 M. What is Ksp?
11. The solubility of LaF3 is 9.3 x 10-6 M. What is Ksp?
12. 500. mL of a saturated solution contains 0.0651 g of MgF2
at 25oC.
a. What is the solubility in mol/L?
b.
What is Ksp?
13. The solubility of BaC2O4 is 22 mg/L.
a. What is the solubility in mol/L?
b.
What is Ksp?
19. Correct the formulas of the complex compounds.
Name
Formula
Corrected
hexammineiron(II)
[Fe(NH3)4](SO4)2
sulfate
potassium
[Ni(CN)4]K4
tetracyanonickelate(II)
20. Complete the complex ion synthesis reactions and write a
Kf expression.
Ag+ + 2 NH3 
Cu2+ + 4 Cl- 
Al3+ + 4 OH- 
Fe3+ + SCN- 
21. Explain the observations using chemical equations.
a. AgCl is soluble in concentrated ammonia solution.
b.
Ammonia added to a solution of Cu(NO3)2 turns the
solution from light blue to dark blue.
22. Given the following equilibriums with their constants.
BaF2(s)  Ba2+ + 2 FKsp = 1.8 x 10-7
14. 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-]
Cr3+ + 4 OH-  Cr(OH)4-
Kf = 8.0 x 1029
Cr(OH)3(s) 
Ksp = 1.6 x 10-30
Cr3+
+3
OH-
Cu2+ + 4 NH3(aq)  Cu(NH3)42+
Kf = 5.0 x 1012
Cu(OH)2(s) 
Ksp = 4.8 x 10-20
Cu2+
+2
OH-
HF(aq)  H+ + F-
Ka = 6.8 x 10-4
H2O(l) 
Kw = 1.0 x 10-14
H+
+
OH-
H2S(aq)  2 H+ + S2-
Q
ppt?
15. 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+]
K = 1.0 x 10-20
MnS(s) 
+
Ksp = 2.5 x 10-13
Calculate K for the equilibriums below.
a. MnS(s) + 2 H+  Mn2+ + H2S(aq)
Mn2+
S2-
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)
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)?
[PO43-]
Q
ppt?
16. 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
17. What does the difference between the two answers from
questions (8.a) and (8.b) illustrate?
18. Explain the observations using chemical equations.
a. Statues made of marble (CaCO3) that are displayed in
polluted (acidic air) cities lose their definition over time.
b.
Milk of Magnesia is a medication that absorbs excess
stomach acid contains Mg(OH)2.
Practice Multiple Choice
Briefly explain why the answer is correct in the space provided.
1. What is the correct formula for hypochlorite?
(A) ClO(B) ClO2(C) ClO3(D) ClO4-
2.
What is the correct formula for copper(II) sulfate?
(A) CuS
(B) Cu2S
(C) CuSO3 (D) CuSO4
3.
Which of the following is the correct name for the
compound with formula Ca3P2?
(A) Calcium phosphite
(B) Tricalcium diphosphorus
(C) Calcium phosphate (D) Calcium phosphide
4.
5.
6.
Which is the net ionic equation for the precipitation reaction
between Lead(II) nitrate and sodium phosphate?
(A) Na+ + NO3-  NaNO3(s)
(B) 3 Pb(NO3)2(aq) + 2 Na3PO4(aq) 
6 NaNO3(s) + Pb3(PO4)2(aq)
(C) 3 Pb2+ + 2 PO43-  Pb3(PO4)2(s)
(D) 3 Pb(NO3)2(aq) + 2 Na3PO4(aq) 
Pb3(PO4)2(s) + 6 NaNO3(aq)
0.20-L of 0.20 M K2CO3 is added to 0.30-L of 0.40 M
Ba(NO3)2. Barium carbonate precipitates. The
concentration of barium ion, Ba2+, remaining in solution is
(A) 0.15 M (B) 0.16 M (C) 0.20 M (D) 0.24 M
Na3PO4(aq) + 3 AgNO3(aq)  Ag3PO4(s) + 3 NaNO3(aq)
100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0 M
AgNO3. Which is the order of increasing concentration of
the ions remaining in solution?
(A) [PO43-] < [NO3-] < [Na+]
(B) [PO43-] < [Na+] < [NO3-]
(C) [NO3-] < [PO43-] < [Na+]
(D) [Na+] < [NO3-] < [PO43-]
7.
An aqueous solution contains Pb2+, Fe2+, and Cu2+ ions.
Which will separate Pb2+ from the other ions at 25oC?
(A) Adding dilute Na2S solution
(B) Adding dilute HCI solution
(C) Adding dilute NaOH solution
(D) Adding dilute NH3 solution
8.
Which is LEAST soluble in water?
(A) (NH4)2SO4
(B) KMnO4
(C) BaCO3
(D) Zn(NO3)2
9.
Which compound is NOT appreciably soluble in water but
is soluble in dilute hydrochloric acid?
(A) Mg(OH)2
(B) K2CO3
(C) CuSO4
(D) (NH4)2SO4
Questions 10-12 refer to the following solid compounds.
(A) PbSO4 (B) CuSO4 (C) KMnO4 (D) KCl
10. Is purple in aqueous solution
13. Which occurs when NH3 is mixed with 0.1 M Cu(NO3)2?
(A) A dark red precipitate forms and settles out.
(B) Separate layers of immiscible liquids form.
(C) The color turns from light blue to dark blue.
(D) Bubbles of ammonia gas form.
14. The formula for hexaamineiron(II) sulfate is
(A) [Fe(NH3)6]SO4
(B) SO4[Fe(CN)6]
(C) [Fe(NH3)4]SO4
(D) [Fe(NH3)6]2(SO4)3
15. Cl-, OH- and SO42- are added to three samples of a colorless,
0.10 M solution, respectively. Which cation could be present
in the solution if no change is observed in any sample?
(A) Ni2+
(B) Ag+
(C) Ba2+
(D) Na+
Questions 16-20 Answer the following questions that relate to
the solubility of Al(OH)3 (Ksp is 2 x 10-31 at 25oC).
16. The [Al3+] at 25oC when [OH-] = 1 x 10-4 M is
(A) (2 x 10-31)/(1 x 10-12) (B) (2 x 10-31)/3(1 x 10-4)
(C) [(2 x 10-31)/(1 x 10-4)]3 (D) (1 x 10-12)/(2 x 10-31)
17. The solubility at 25oC is
(A) 27(2 x 10-31)4
(C) 27(2 x 10-31)¼
(B) 4(2 x 10-31)3
(D) [(2 x 10-31)/27]¼
18. Ksp at 50oC when the solubility is 1.0 x 10-6 M is
(A) 5.3 x 10-19
(B) 2.7 x 10-19
-23
(C) 2.7 x 10
(D) 2.0 x 10-14
19. Which is true of a mixture where the initial concentrations
of [OH-] and [Al3+] are 1 x 10-8 M at 25oC?
(A) Q = 1 x 10-8  Al(OH)3 will precipitate
(B) Q = 27 x 10-8  Al(OH)3 will precipitate
(C) Q = 1 x 10-32 
3 will NOT precipitate
(D) Q = 27 x 10-32 
3 will NOT precipitate
20. What is the equilibrium constant for the following reaction:
Al(OH)3 + OH-  Al(OH)4-? (Kf = 1.0 x 1033 for Al(OH)4-)
(A) 2.0 x 10-2
(B) 2.0 x 102
4
(C) 2.0 x 10
(D) 8.0 x 106
21.
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
(C) 2 x 10-6
(D) 5 x 105
22. 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
11. Is blue and very soluble in water
12. Is white and insoluble in water
23. What is the solubility of Ag2CrO4 (Ksp = 8 x 10-12)?
(A) 8 x 10-12 M
(B) 2 x 10-12 M
-12
½
(C) (4 x 10 ) M
(D) (2 x 10-12)⅓ M
24. 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
25. 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+
e.
3.
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.
How many grams of Ca(OH)2 can you dissolve in 100
mL of 1.0 x 10-3 M NaOH?
f.
At a different temperature, the solubility of Ca(OH)2 is
0.15 g/L. What is the Ksp?
g.
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.
h.
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)
i.
Calculate K for: Ca(OH)2(s) + 2 H+  Ca2+ + 2 H2O.
(H2O  H+ + OH- Kw = 1.0 x 10-14)
j.
Is Ca(OH)2 soluble in HF? Justify your answer.
(HF(aq)  H+ + F- Ka = 6.8 x 10-4)
Practice Free Response
1.
For each reaction, in part (1) write a net ionic balanced
equation and in part (2) answer the question about the
reaction. Assume that solutions are aqueous unless
otherwise indicated.
a. A solution of sodium hydroxide is added to a solution
of lead(II) nitrate.
(1) Balanced equation:
(2) If 1.0 L volumes of 1.0 M solutions of sodium
hydroxide and lead(II) nitrate are mixed together,
how many moles of product(s) will be produced?
b.
Excess nitric acid is added to solid calcium carbonate.
(1) Balanced equation:
(2) Briefly explain why statues made of marble
(calcium carbonate) displayed outdoors in urban
areas are deteriorating.
2.
The identity of an unknown solid is to be determined. The
compound is one of the seven salts in the following table.
NaCl
BaCl2• 2 H2O
CaCO3
CuSO4• 5 H2O
Al(NO3)3• 9 H2O
BaSO4
Ni(NO3)2• 6 H2O
Use the results of the following observations or laboratory
tests to explain how each compound may be eliminated or
confirmed. The tests are done in sequence from (a)
through (e). Explain your reasoning.
a. The unknown solid is white. Which two compounds
can be eliminated because they have color?
b.
The unknown is soluble in water. Which other two
compounds can be eliminated because they are
insoluble in water.
c.
The unknown forms a white precipitate when AgNO3(aq)
is added. Which other compound can be eliminated
because it is soluble in AgNO3(aq)?
d.
The unknown loses mass when carefully heated.
Which other compound can be eliminated because it
does not lose mass when carefully heated.
Describe a test that can be used to confirm the identity
of the remaining unknown compound. Limit your
confirmation test to a reaction between an aqueous
solution of the unknown compound and an aqueous
solution of one of the other soluble salts listed in the
table. Describe the expected results of the test;
include the formula of the product.