Download chem A exercise package C

130
EXERCISE 1
GRAPHING
Make a full-page graph of the following data: Put the mass on
the vertical axis and the volume on the horizontal axis.
VOLUME (ML)
5.0
10.0
5.0
15.0
20.0
25.0
2O .0
25.0
MASS (G)
4.10
15.50
7.70
11.90
31.10
19.70
15.80
39.90
M/V
QUESTIONS
1. Could you draw one or two lines that would come close to
connecting all the points?
2. For each pair of numbers divide the mass by the volume.
Example: The numerals 5.0 and 4.10 are a pair. You would
divide 5.0 into 4.10. In the column on the right side of
your Data Table enter the results.
3. What regularity exists in your answers for Question 2?
131
EXERCISE 2
DENSITY
Formulas for density problems
M = mass (g)
V = volume (ml)
D = density (g/ml)
D = M/V
V = M/D
M = VxD
1. What is the density of a block of wood that weighs 200
grams and has a volume of 400 ml?
2. What volume of liquid bromine will weigh 2000 g? (Density
of liquid bromine is 3.10 g/ml.)
3. A rectangular block measures 20 cm x 10 cm x 5.0 cm and
has a density of .3 g/ml. What does it weigh?
4. What is the volume of 150 g of nitric acid solution of
density 1.35 g/ml?
5. Concentrated sulfuric acid has a density of 1.860 g/ml.
How much does 1.00 liter of it weigh?
SHOW YOUR WORK.
132
EXERCISE 3
SIGNIFICANT FIGURES CALCULATIONS
DIRECTIONS: On a separate sheet of paper, complete the
following.
A. Determine the number of significant figures in each of the
following measurements.
1.
5.432 g
11.
2500 cm
2.
40.319 g
12.
2500.0 cm
3.
146 mm
13.
1.04 x 1014 g
4.
3.256 cm
14.
3.58 x 10-9 L
5.
0.189 in
15.
46.0979 g
6.
439.6 mL
16.
8365.6 g
7.
2873.0 cc
17.
0.002300 mg
8.
99.9 s
18.
7.500 x 108 oz
9.
0.000235 g
19.
3.92 x 10-4 g
10.
144 cL
20.
1.00000 x 103 lbs.
B. Add the following measurements:
1.
12 cm + 0.031 cm + 7.969 cm
(20. cm)
2.
0.085 cm + 0.062 cm + 0.14 cm
(.29 cm)
3.
3.419 g + 3.912 g + 0.00013 g
(7.331 g)
4.
30.5 g + 16.82 g + 41.07 g + 85.219 g
(173.6 g)
5.
2.3 x 102 g + 4.62 x 102 g + 3.852 x 102 g
133
(1.08 x 103 g)
C. Subtract the following measurements:
1.
41.025 cm – 23.28 cm
(17.74 cm)
2.
289 g – 43.7 g
(245 g)
3.
145.63 mL – 28.9 mL
(116.7 mL)
4.
62.47 g – 39.9 g
(22.6 g)
5.
40.008 mL – 29.0941 mL
(10.914 mL)
D. Multiply the following measurements:
1.
2.89 cm x 4.01 cm
(11.6 cm2)
2.
17.3 cm x 6.2 cm
(110 cm2)
3.
5 cm x 5 cm
(20 cm2)
4.
5.0 cm x 5 cm
(20 cm2)
5.
5.0 cm x 5.0 cm
(25 cm2)
6.
4.8 x 102 m x 2.101 x 103 m
(1.0 x 106 m2)
7.
4.218 cm x 6.5 cm
(27 cm2)
8.
150.0 m x 4.00 m
(6.00 x 102)
9.
282.2 km x 3.0 km
(850 km2)
10.
14 x 10-8 m x 3.25 x 10-6 m
(4.6 x 10-13)
E. Divide the following measurements:
1.
8.071 cm2 / 4.216 cm
(1.914 cm)
2.
109.3758 m2 / 5.813 m
(18.82 m)
3.
24 789.4 km2 / 43.5 km
(5.70 x 102km)
4.
6.058 mm2 / 0.85 mm
(7.1 mm)
5.
4.819 cm2 / 9.852 cm
(.4891 cm)
6.
139.482 m2 / 68.75 m
(2.029 m)
7.
4.23 m2 / 18.941 m
(.223 m)
8.
85.621 km2 / 8.05 km
(10.6 km)
9.
6.93 x 10-5 cm2 / 4.09 x 10-8 cm
(1.69 x 103 cm)
10.
3.005 x 106 m2 / 1.8765 x 104 m
(1.601 x 102 m)
134
EXERCISE 4
SUBATOMIC PARTICLES
ELEMENT
A
B
C
D
E
F
G
H
I
J
K
L
M
N
O
P
R
S
T
ATOMIC
NUMBER
2
ATOMIC
WEIGHT
4
13
ELECTRONS
2
5
PROTONS
2
8
80
76
108
175
31
190
153
16
7
28
59
7
28
28
23
31
12
4
18
93
92
2
8
8
45
98
47
71
NEUTRONS
5
14
28
40
237
146
16
32
135
EXERCISE 5
CHEMICAL FAMILIES
1. A student used the Data Table given below while observing
chemical reactions. N.R. = no reaction and R = a reaction.
Use the data to answer the questions.
Data Table
A.
D. More observations of
Phosphorus:
Observations of Potassium:
and
and
and
and
and
and
and
and
and
and
Argon
Chlorine
Fluorine
Helium
Hydrogen
Krypton
Neon
Nitrogen
Oxygen
Xenon
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(R)
(R)
(R)
(R)
(R)
(R)
(R)
(R)
(R)
(R)
and
and
and
and
and
Hydrogen (N.R.)
Fluorine (N.R.)
Chlorine (N.R.)
Bromine (N.R.)
Iodine
(N.R.)
(R)
(R)
(R)
(R)
(R)
E. More observations of Potassium:
and Bromine (N.R.) (R)
and Iodine
(N.R.) (R)
F. Observations of Chlorine:
B. Observations of Lithium:
and
and
and
and
and
and
Nitrogen
Argon
Helium
Krypton
Neon
Xenon
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(R)
(R)
(R)
(R)
(R)
(R)
and
and
and
and
and
Fluorine
Argon
Helium
Krypton
Neon
Xenon
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(N.R.)
(R)
(R)
(R)
(R)
(R)
G. Observations of some metals
when placed in water
Water and:
Lithium
(N.R.) (R)
Sodium
(N.R.) (R)
Potassium
(N.R.) (R)
Rubidium
(N.R.) (R)
Cesium
(N.R.) (R)
C. Observations of Phosphorus:
and
and
and
and
and
and
Lithium
Sodium
Potassium
Rubidium
Cesium
(R)
(R)
(R)
(R)
(R)
(R)
136
VI. QUESTIONS
1. Look over your first set of observations (Set A). Which elements
do not react with potassium?
2. Look over your second set of observations. Which elements do
not react with lithium?
3. Now check your third set of data. Which elements do not react
with phosphorus?
4. It should now be apparent to you that there is a regularity in the
first three sets of data. The generalization can be made that
helium, neon, argon, krypton, and xenon are all nonreactive
elements. These gases represent a family of elements that are
chemically inert. How many electrons does a neutral atom of
each of these have?
5. How many valence electrons do each of these elements have?
6. Examine your data for the observations made of phosphorous.
Which gases can be said to be reactive? How many electrons do
neutral atoms of these elements have?
7. How many valence electrons do each of these have?
8. Compare these numbers with the numbers of electrons in the
inert gas family. Would there be reason to believe that the gases
that reacted with phosphorus are also a chemical family?
9. Look over the data for chlorine gas in the sixth set of
observations. Is there any regularity? How many electrons do
neutral atoms of these elements have? How does this number
compare with the number in the inert gas family? Can you make
any generalization about these elements?
10. Look over the last set of observations you made. Is there any
evidence to support your answer to Question 6?
137
EXERCISE 6
WHAT MAKES A COMPOUND?
Predict what ions you would get if you dissolved the following ionic
substances.
EXAMPLE:
MgBr2
A. The + ion is always written first. So you look on the left side of the
common ion chart and find the positive ion, which is Mg2+ in this
case.
B. The - ion is always written last. So you look on the right side of
the common ion chart and find the negative ion, which is Br- in this
case.
Answer:
Mg2+ and Br-
1. FeCl3
2. NaI
3. Ba(NO3)2
4. (NH4)3PO4
5. Al2(SO4)3
6. MgSO4
7. Na2CO3
8. CaF2
9. PbI2
10. LiClO
138
EXERCISE 7
FORMULA WRITING
DIRECTIONS: On a separate piece of paper, complete the following chart:
Bromide
Br-
Nitrate
NO3-
Hydroxide
OH-
Oxide
O2-
Sulfate
SO42-
Potassium
K+
Silver
Ag+
Ammonium
NH4+
Mercury (I)
Hg22+
Mercury
(II)
Hg2+
Calcium
Ca2+
Copper (II)
Cu2+
Iron (II)
Fe2+
Iron (III)
Fe3+
Antimony(I
II)
Sb3+
139
Phosphate
PO43-
Hydrogen
Carbonate
HCO3-
Chlorate
ClO3-
Acetate
C2H3O2-
EXERCISE 8
FORMULA TO NAME
DIRECTIONS: On a separate sheet of paper name the following compounds
using the stock system.
1.) KCl
2.) NaOH
3.) ZnF2
4.) CuNO2
5.) NaMnO4
6.) KHSO3
7.) Mn (OH)2
8.) CrCl3
9.) (NH4)2SO4
10.) AgCl
11.) PbCrO4
12.) NaOH
13.) NiCl2
14.) Hg2Cl2
15.) Cu(NO3)2
16.) Al2(C2O4)3
17.) Na2Cr2O7
18.) AgClO3
19.) KH2PO4
20.) Cr2S3
21.) NaF
22.) CuNO2
23.) FeBr2
24.) HgCO3
25.) NH4OH
26.) CrCl2
27.) HCl
28.) H2S
29.) Sn(NO3)2
30.) Sn(NO2)4
31.) KOH
32.) Hg(OH)2
33.) KCl
34.) FeCl3
35.) HNO3
36.)
37.)
38.)
39.)
40.)
41.)
42.)
43.)
44.)
45.)
46.)
47.)
48.)
49.)
50.)
51.)
52.)
53.)
54.)
55.)
56.)
57.)
58.)
59.)
60.)
61.)
62.)
63.)
64.)
65.)
66.)
67.)
68.)
69.)
70.)
Cu2O
Al2(SO4)3
N2O 5
HF
Pb(OH)2
NH4NO3
NaHCO3
HgO
Zn(NO2)2
H3PO4
CsOH
Li2O
Ca(OH)2
CaBr2
Fe2O3
H2SO4
FeCO3
SO3
Ba(BrO3)2
Al(OH)3
HClO4
NaC2H3O2
Na2SO3
H2CO3
NH4IO3
LiH
CO
MgBr2
SnBr2
N2O
NH4F
AsCl3
KHCO3
K2O
ZnO
140
EXERCISE 9
NAME TO FORMULA
DIRECTIONS: On a separate sheet of paper, write the correct empirical formula for
each of the following:
1. Sodium nitrate
28. Manganese (VII) oxide
2. Aluminum sulfate
29. Mercury (I) oxide
3. Potassium chromate
30. Iron (II) chloride
4. Lithium hydroxide
31. Copper (II) chloride
5. Ammonium phosphate
32. Tin (II) chloride
6. Magnesium nitrate
33. Titanium (IV) iodide
7. Sodium monohydrogen phosphate
34. Nickel (II) fluoride
8. Sodium bicarbonate
35. Manganese (IV) oxide
9. Potassium bisulfate
36. Lead (IV) oxide
10.
Hydrogen acetate
37. Nickel (II) oxide
11.
Lead (II) chromate
38. Mercury (II) oxide
12.
Copper (II) carbonate
39. Cobalt (III) oxide
13.
Silver nitrate
40. Copper (II) chlorate
14.
Ferrous sulfide
41. Cobalt (III) sulfate
15.
Manganous sulfide
42. Manganese (III) sulfate
16.
Barium sulfite
43. Iron (III) nitrate
17.
Sodium iodide
44. Tin (IV) nitrate
18.
Calcium nitrate
45. Cobalt (II) perchlorate
19.
Zinc fluoride
46. Chromium (III) sulfate
20.
Potassium oxalate
47. Iron (II) hydroxide
21.
Cuprous chloride
48. Copper (II) phosphate
22.
Manganese (III) chloride
49. Mercury (I) nitrite
23.
Iron (III) bromide
50. Lead (II) nitrate
24.
Chromium (III) bromide
51. Mercury (I) sulfate
25.
Tin (IV) chloride
52. Tin (IV) oxide
26.
Chromium (III) oxide
53. Manganese (II) bromide
27.
Lead (II) oxide
141
EXERCISE 10
FORMULAS OF COMPOUNDS
A student performs an experiment similar to our experiment 5.2. His
summary of observations appears below. For convenience he numbered
each of his seven precipitates that formed.
Ions Present In
Each Solution Used
NiCl2
BaCl2
Na2SO4
NaOH
Ba(OH)2
PPT 1
PPT 2
MgSO4
NiCl2
XX
BaCl2
XX
XX
PPT 3
Na2SO4
XX
XX
XX
NaOH
XX
XX
XX
XX
Ba(OH)2
XX
XX
XX
XX
XX
PPT 7
MgSO4
XX
XX
XX
XX
XX
XX
PPT 4
PPT 5
PPT 6
INSTRUCTIONS
1. Fill in the positive and negative ions that are present in each of the
six original solutions in the blanks provided in the Data Table--just
the kinds of ions, not necessarily the numbers (just OH- not 2OHetc.).
2. Now, recalling that unlike charges attract, list the possible formulas
for each of the seven precipitates that formed. Here’s how to do
that. For each of the precipitates, write the two chemicals the
student mixed to get it. (These are called reactants.) Now
exchange the positive ion of each. You need to re-criss-cross the
new formulas to get them right. These are called products. There
should be two. One of each pair of products is the formula of the
precipitate.
3. Go back and eliminate those that can be ruled out by examining the
results of the other combinations tried. Lightly cross out those
formulas. If you do this carefully, you can reason out the correct
formula for each precipitate. In fact, you'll discover that one
precipitate is probably a mixture of two solid substances.
142
EXERCISE 11
ASSIGNING OXIDATION NUMBERS
Directions: On a separate sheet of paper, assign the oxidation number for
each element of the following:
1.
H2SeO3
14.
OF2
2.
Al(NO3)3
15.
MnO2
3.
BaH2
16.
Ca3(PO4)2
4.
Na4SiO4
17.
CuCrO4
5.
NaIO4
18.
Na3AsO3
6.
Na2O2
19.
CaS
7.
HBrO
20.
KHSO3
8.
N2O3
21.
CO3-2
9.
(NH4)2CO3
22.
IO4-
10.
K2Cr2O7
23.
SCN-
11.
Mg(ClO4)2
24.
ClO-
12.
Sb2O3
25.
Hg22+
13.
H3PO4
143
How to Diagram Lewis Structures
1.
Find the total number of valence electrons supplied by all the atoms in
the structure. The number supplied by each A family element is the
same as the group number of the element:
a.
b.
2.
For a negative ion, increase the number by the charge of the ion.
For a positive ion, decrease the number by the charge of the ion.
Determine the number of electrons that would be required to give two
electrons to each hydrogen atom and eight electrons to each of the
other atoms individually:
Num. e- for individual atoms = 2(num. H atoms) + 8(num. other
atoms)
3.
The number obtained in step 2 minus the number obtained in step 1 is
the number of electrons that must be shared in the final structure:
Num. bonding e- = (num. e- for individual atoms) – (total num. e-)
4.
One-half the number of bonding electrons (step 3) is the number of
covalent bonds in the final structure:
Num. e- pair bonds = (num. bonding e-)/2
5.
Write the symbols for the atoms present in the structure, arranging
them in the way that they are found in the structure.
6.
Indicate electron-pair bonds by dashes written between the symbols.
Indicate one bond between each pair of symbols, and then use any
remaining from the total calculated in step 4 to make multiple bonds.
Note that each H atom is limited to one bond.
7.
The total number of electrons (step 1) minus the number of bonding
electrons (step 3) is the number of unshared electrons:
Num. unshared e- = (total num. e-) – (num. bonding e-)
Complete the electron octet of each atom (other than the H atoms) by
adding dots to represent unshared electrons.
8.
Indicate the formal charges of the atoms where appropriate, and
evaluate the structure.
144
EXERCISE 12
ELECTRON-DOT SYMBOLS (LEWIS STRUCTURES)
DIRECTION: On a separate sheet of paper, construct the electron-dot
symbols for each of the following:
IONIC COMPOUNDS
1.
NaI
5.
BaO
2.
CaF2
6.
K3P
3.
AlCl3
7.
RbBr
4.
Na2S
8.
Li3N
MOLECULES AND POLYATOMIC IONS
9.
HClO3
26.
OF2
10. ClO3-
27.
H2S
11. ClO2-
28.
PCl3
12. ClO-
29.
SiO2
13. NH3
30.
CO2
14. NH4+
31.
HClO4
15. NH2-
32.
N2
16. HNO3
33.
H2SO4
17. NO3-
34.
PH3
18. HNO2
35.
HCN
19. NO2-
36.
HClO
20. H2O
37.
C2H4
21. H3O+
38.
C2H2
22. OH23. CHI3
24. CH3OH
145
EXERCISE 13 INSTRUCTION
Many formulas for substances cannot be explained in terms of ionic bonding.
Consider the substance Cl2O. Both the chlorine and the oxygen atom need
more electrons for a stable electron population. A model proposed that
would allow both atoms to gain electrons is shown in the diagram on this
page. The model proposes that atoms overlap and that each atom put one
electron into this overlapping region or into an electron "pool." By doing this,
each atom appears to gain an electron within its original boundary. For every
overlapping region an atom appears to gain one electron. Two overlapping
regions, such as for oxygen, will result in the gain of two electrons. This
process of overlapping atoms is called covalent bonding. The substance that
results from covalent bonding is called a covalent substance. The process of
overlapping atoms will keep occurring for a particular atom until it has a
stable number of electrons in its original boundary. In general, covalent
substances form when both atoms want to gain electrons. Usually both
atoms will be found toward the right side of the periodic chart.
An important difference between ionic and covalent bonding is that in
covalent bonding ions are not formed. The group of atoms that forms as a
result of covalent bonding is electrically neutral. As a whole there is the
same number of electrons as protons. The group of atoms is referred to as a
molecule.
Formation of Cl2O
146
EXERCISE 13
COVALENT BONDING
WORKSHEET ON COVALENT BONDING
1. Make a diagram similar to those on formation of Cl2O and Br2 for the
bonding arrangement in the following molecules:
A. Br2O
B. NH3
C. CH4
D. PF3
Show on your diagram:
1. The number of electrons involved or not involved with bonding
2. Each covalent bond
147
EXERCISE 14 INFORMATION
148
EXERCISE 14
PREDICTING FORMULAS OF COMPOUNDS
Using the Chart "Some Compounds of Third Row Elements" predict the
formulas of the compounds produced by combining the following elements:
1. Calcium and Chlorine
2. Boron and Oxygen
3. Nitrogen and Hydrogen
4. Hydrogen and Oxygen
5. Fluorine and Oxygen
6. Selenium and Fluorine
7. Carbon and Bromine
8. Magnesium and Tellurium
9. Gallium and Sulfur
10. Sodium and Iodine
11. Barium and Selenium
12. Germanium and Oxygen
13. Indium and Iodine
14. Arsenic and Hydrogen
15. Antimony and Selenium
149
EXERCISE 15
FORMULA MASS
DIRECTIONS: On a sheet of paper, determine the formula mass of each of
the following:
1.
H3PO4
16.
K4Fe(CN)6
2.
AlCl3
17.
Nd2O3
3.
Dy(OH)3
18.
Sb(NO3)3
4.
K2C4H4O6
19.
K3PO4
5.
H2SO4
20.
Ga2(SO4)3
6.
N2O5
21.
zinc acetate
7.
HNO3
22.
copper (I) sulfate
8.
NiSO4
23.
carbon dioxide
9.
Sn(OH)2
24.
calcium bicarbonate
10. (NH4)3PO4
25.
carbonic acid
11. Fe(C2H3O2)3
26.
aluminum nitrate
12. SO2
27.
ammonium sulfate
13. AgNO3
28.
barium chloride
14. NaIO4
29.
iron (II) phosphate
15. Pr(OH)3
30.
strontium sulfite
150
EXERCISE 16
MOLE THEORY
Directions: On a separate sheet of paper, complete the following table using
mole theory.
SUBSTANCE
1. BaBr2
GRAM
FORMULA
MASS
QUANTITY
IN GRAMS
NUMBER
OF MOLES
35.0 g
0.120
42.0 g
XX
5. H2S
0.223
12.0X1023
6. Bi2O3
XX
5.00 g
8. Cu(OH)2
0.100
XX
3.00X1022
9. FeCl2
10. B2H6
26.7
11. CO2
1.80
XX
XX
6.00X1022
12. LiH
13. Na2SO4
XX
6.00X1024
3. C2H6
7. H2O
VOLUME IN
LITERS AT
STP
XX
2. AuCl3
4. PbCO3
NUMBER OF
MOLECULES
36.9 g
XX
XX
14. FeCl2
13.2
XX
7.94X1024
15. C6H6
151
EXERCISE 17
EMPIRICAL FORMULAS
Directions: On a separate sheet of paper, calculate the empirical formula for
each of the following:
1.
88.8 % copper and 11.2 % oxygen
2.
40.0 % carbon; 6.67 % hydrogen and 53.33 % oxygen
3.
92.3 % carbon; and 7.7 % hydrogen
4.
70.0 % iron and 30.0 % oxygen
5.
5.88 % hydrogen and 94.12 % oxygen
6.
79.90 % copper and 20.10 % oxygen
7.
56.4 % potassium; 8.7 % carbon and 34.9 % oxygen
8.
10.04 % carbon; 0.84 % hydrogen and 89.12 % chlorine
9.
42.50 % chromium and 57.50 % chlorine
10.
15.8 % carbon and 84.2 % sulfur
11.
30.43 % nitrogen and 69.57 % oxygen
12.
82.4 % nitrogen and 17.60 % hydrogen
13.
12.5 % hydrogen; 37.5 % carbon and 50.0 % oxygen
14.
75.0 % carbon and 25.0 % hydrogen
15.
29.40 % calcium; 23.56 % sulfur and 47.04 % oxygen
152
EXERCISE 18
TYPICAL MOLE PROBLEMS
1.
Determine the number of moles of each kind of atom in each of the
following:
5 moles C3H5Br
7 moles Ni(OH)2
1 mole CO2
2 moles (NH4)2SO4
2.
Find the molecular weight of these:
H2O
F2
NaNO3
KH2PO4
3.
Find the number of moles of each compound:
36g H2O
4.
38g F2
Given Ba(NO3)2 + NaCl
 BaCl2 + NaNO3
How many moles of NaCl react with:
1 mole Ba(NO3)2
5 moles Ba(NO3)2
5.
How many moles of copper atoms are in 2 moles of CuSO4?
6.
How many moles of molecules would there be in 10.g of
H2O
CaCO3
7. If you had 5 moles of CH3CCl3 molecules, how many grams would they
weigh?
8. How many moles of particles are in 300. x 1023 particles?
9. Which would contain more molecules : 20. g of NH4Cl or 20. g of H20?
10. How many moles of oxygen atoms are in 10 moles of Fe3(PO4)2
molecules?
153
EXERCISE 19
BALANCING EQUATIONS
Directions: On a separate sheet of paper, balance the following equations by
inspection:
1.
Na
+
Br2
2.
Al
3.
P + O2 
4.
N2O4 
5.
HClO3
6.
Al2O3 +
7.
C12H22O11
8.
NH4NO3 
9.
H3PO3

10.
Sb2S3
+
11.
(NH4)2Cr2O7 
12.
C4H10
13.
TiCl4 + H2O
14.
H2SO3
15.
Fe
16.
CaO
17.
Sc2O3
+
18.
Pb
+
H 2O
+
O2

Pb(OH)2
19.
Ag
+
H 2S
+
O2

Ag2S
20.
P4
+
S8 
21.
NaF
22.
Al
23.
N2 +
24.
Zn
25.
C12H26
+
S
Al2S3

P2O5
NO2
ClO2 +

Cl2
+
PH3
+
H2S
Cr2O3
+
H2O + N2
CO2
+
H2O

 TiO2 + HCl
SO2 + H2O
Fe3O4 +

SO3


HNO3
O2
H2
Ca3(PO4)2

Sc2(SO4)3

+
H 2O
P4S10
FeCl3 
+
SbCl3 +

P2O5
H2
+ CO
H2O
H3PO4 +
O2
HCl
AlCl3
H2O
H2O
+

+ O2
+

+
C
N2 O
HCl
+
H2O
C+

+
+
NaBr

Na3FeF6 +
AlCl3 +
NaCl
H2
NH3


Zn(NO3)2
+
CO2 +
H2O
154
H2
EXERCISE 20
COMPLETE, BALANCE AND TYPE OF REACTION
Directions: On a separate sheet of paper, complete; balance and indicate the
type of reaction for the following:
1.
iron + copper (II) sulfate 
2.
potassium chloride + mercury (I) sulfate 
3.
sulfur + oxygen 
4.
sulfuric acid + cadmium 
5.
zinc + silver sulfate 
6.
aluminum sulfite + sodium hydroxide 
7.
iron (III) oxide + hydrogen 
8.
sodium sulfate + barium bromide 
9.
mercury (II) nitrate + potassium sulfate 
10.
nitrogen + hydrogen 
11.
silver nitrate + sodium carbonate 
12.
mercury + oxygen 
13.
sodium hydroxide + barium bromide 
14.
silicon + sulfur 
15.
antimony (III) sulfide + iron 
16.
potassium chloride + aluminum chloride 
17.
copper + oxygen 
18.
calcium sulfide + hydrochloric acid 
19.
antimony + sulfur 
20.
iron (III) oxide + aluminum 
21.
mercury (II) oxide 
22.
iron (III) bromide + barium hydroxide 
23.
magnesium + lead (II) sulfate 
24.
aluminum + nitrogen 
25.
hydrosulfuric acid + iron (II) chloride 
155
EXERCISE 21
COMPLETE AND BALANCE
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
SODIUM CARBONATE 
LITHIUM + NITRIC ACID 
CARBON + SULFUR 
SILVER ACETATE + POTASSIUM CHLORATE 
ACETIC ACID + POTASSIUM CARBONATE 
OXYGEN + TITANIUM 
LEAD (II) PHOSPHATE + SODIUM BROMIDE 
C7H16 BURNS 
CESIUM OXALATE + GOLD (III) NITRITE 
IRON (III) CHLORIDE 
CALCIUM + SILICON 
ALUMINUM PHOSPHATE + MOLYBDENUM (VI) ACETATE 
HYDROGEN SULFIDE + AMMONIUM HYDROXIDE 
MANGANESE + SILVER 
SELENIUM + CHLORINE 
PHOSPHORUS BURNS 
POTASSIUM CARBONATE + HYDROCHLORIC ACID 
VANADIUM (V) OXIDE 
COBALT (III) SULFIDE + CARBONIC ACID 
POTASSIUM CHLORATE 
HYDROGEN PERCHLORATE + AMMONIUM HYDROXIDE 
CARBON + IODINE 
COPPER (II) HYDROGEN SULFATE + HYDROGEN HYDROXIDE 
C21H44 BURNS 
MANGANESE (II) CARBONATE + SULFURIC ACID 
OXYGEN + FLUORINE 
POTASSIUM THIOCYANATE + IRON (III) CHLORIDE 
TUNGSTEN (VI) OXIDE + SODIUM 
LEAD + COBALT (III) NITRITE 
BARIUM HYDROXIDE + HYDROGEN SULFITE 
CALCIUM + WATER 
CARBON DIOXIDE + NICKEL (III) OXIDE  (ONE PRODUCT)
SULFUR BURNS 
PHOSPHORUS (V) BROMIDE 
MAGNESIUM PHOSPHATE + COPPER (I) ACETATE 
AMMONIUM CARBONATE  (THREE PRODUCTS)
ZINC + CARBONIC ACID 
CARBON DIOXIDE + WATER  (ONE PRODUCT)
BORON + FLUORINE 
MERCURY (I) NITRATE + STRONTIUM IODIDE 
MANGANESE OXIDE 
SULFURIC ACID + CESIUM 
SULFURIC ACID  (TWO PRODUCTS)
MANGANESE (II) HYDROXIDE + HYDROCHLORIC ACID 
SELENIUM + SULFUR 
156
EXERCISE 22
CHEMICAL EQUATIONS AND MOLES
DIRECTIONS: On a separate sheet of paper, do the following: (Note: Equations are
unbalanced.)
1. How many moles of cadmium nitrate will react with 6.678 moles of sulfuric acid?
Cd(NO3)2 + H2SO4  CdSO4 + HNO3 (6.678 mole)
2. How many atoms of iron will react with an excess of sulfur to form 2.0 moles of iron (II)
product? Fe + S  FeS (1.2 x 1024 atoms)
3. How many moles of silver nitrate are needed to react with 0.34654 mole of sodium
(0.34654 mole)
bromide? AgNO3 + NaBr  AgBr + NaNO3
4. Aluminum reacts with oxygen to form the oxide. If we want to produce 45,670 mol of
the oxide how much of each is needed?
(91.340 mol Al)
(68.505 mol O2)
Al + O2  Al2O3
5. How many moles of mercury are formed from the decomposition of 6 908 moles of
mercury (II) oxide? HgO  Hg + O2 (6 908 mol)
6. How many liters of nitrogen at STP are needed to react with 9.000 7g of aluminum?
(3.73 L)
Al + N2  AlN
7. How many grams of sodium hydroxide are needed to react with 2.22kg of barium
bromide? NaOH + BaBr2  NaBr + Ba(OH)2 (598g)
8. How many molecules of hydrogen are needed to react with 55.55g iron (III) oxide? H2 +
(6.29x1023 mcls)
Fe2O3  H2O + Fe
9. How many grams of zinc are needed to react with .000 006 785 grams of silver sulfate?
(1.42x10-6g)
Zn + Ag2SO4  ZnSO4 + Ag
10. How many grams of silicon are needed to react with 23.90g of sulfur?
Si + S  SiS2 (10.5g)
11. How much silver phosphate is produced if 10.0g of silver acetate is reacted with an
excess of sodium phosphate? AgC2H3O2 + Na3PO4  Ag3PO4 + NaC2H3O2 (4.05g)
12. What mass of sodium hydroxide is needed to completely react with 25.0g of sulfuric
(20.4g)
acid?
NaOH + H2SO4  Na2SO4 + H2O
13. Sulfur (IV) oxide reacts with water to form sulfurous acid, how much sulfurous acid is
produced from 2.75g of sulfur (IV) oxide and an excess of water?
SO2 + H2O  H2SO3 (3.52g)
14. 739g of bromine reacts with potassium iodide in excess. How much of each product is
formed in the reaction? Br2 + KI  KBr + I2 (1 170g I2) (1.10x103g KBr)
15. 27.0g of silver oxide decomposes into silver and oxygen. How many liters of oxygen are
(1.30L O2)
produced? Ag2O  Ag + O2
157
EXERCISE 23 INFORMATION
There are two factors to consider when expressing the energy change of a substance:
1. The temperature change of the substance must be considered.
2. The amount of substance that changes temperature must be considered.
Both these factors are combined in a unit of energy called the calorie. A calorie is a unit of
energy. It is defined as the amount of energy it takes to heat one gram of water 1 °C. This
definition can be expressed by three equations:
1. No. of calories = (Change in temperature) x (no. of grams of water)
2. t = _____no. of calories________
no. of grams of water
3. No. of grams of water = no. of calories
t
Typical problems involving the calorie:
1. 400 grams of water are heated from 250C to 400C.
required?
Solution:
Use equation 1.
How many calories are
The number of grams of water = 400
The change in temperature (t)
= 400 - 250
= 150
Thus:
Number of calories = (400) x (15) = 6,000 calories
2. 200 calories of energy are added to 50 grams of water. What temperature change
occurs in the water?
Solution:
Use equation 2. The number of grams of water = 50
The number of calories
= 200
Thus:
t = 200/50 = 4oC
3. 1000 calories of energy are added to a water sample. The temperature of the
water changes from 400C to 650C. How much did the water sample weigh?
Solution:
Use equation 3. The number of calories
Change in temperature t
Weight of water sample = 1000/25 = 40 g
158
= 1000
= 250
EXERCISE 23 INFORMATION, CONTINUED
In the experiment you made an important discovery: A substance doesn't have to change
temperature in order for it to gain or lose heat. You found that when moth ball crystals froze
and changed from a liquid to a solid, there was no change in temperature. In fact, you could
use the procedure followed in this experiment to determine the freezing point of many
substances. Phase change energy is the name given to the energy necessary to change a liquid
to a solid or a solid to a liquid. A molecule does not change its molecular formula when it
undergoes a phase change. If a substance changes from a solid to a liquid, heat must be
added. This is called endothermic change. If a substance changes from a liquid to a solid, heat
must be released. This is called exothermic change.
How would you report this exothermic or endothermic energy change? In other words, what
units would you use to report this energy?
There are two basic ways to express the energy change of a substance:
1. One is calories per gram.
2. The other is calories per mole.
The two basic formulas are:
calories per gram =
no. of calories_____
no. of grams of substance
calories per mole = (calories per gram) x (molecular weight)
An example: When 40 grams of a substance are frozen, this causes 220 grams of water to rise
in temperature from 250C to 330C. Find the number of calories per mole if the
molecular weight is 120 g.
Solution:
To determine the calories per mole we must find the following:
1. No. of calories
2. No. of calories per gram
3. No. of calories per mole
Step 1. No. of calories
= (220) x (330-250)
= (220) x (8) = 1760 calories
Step 2. No. of calories per gram = 1760 calories
40 grams
Step 3. No. of calories per mole
= 44 calories
gram
= (44 cal. per gram) x (120 g/mol)
= 5,280 cal/mol
159
EXERCISE 23-CALORIES
Show your work on each problem.
1. How many calories would be required to heat 220 grams of water
from 200C to 350C?
2. 30 grams of a substance are burned, and the heat produced is used
to warm 2000 grams of water from 220C to 30C.
A. How many calories are released by the substance?
B. How many calories would be released if only 1.0 g of the
substance were burned?
C. If this substance has a molecular weight of 60 grams per mole,
how many calories are released by one mole?
3. Two grams of candle are burned. Each gram of candle releases
7000 cal when burned. If all the heat from the burning candle goes
into a sample of water, what is the weight of the water? The water
temperature increases 250C.
4. If 400 grams of water at 250C are heated by adding 4000 calories to
it,
A. What is t?
B. What is the final temperature of the water?
5. Find the laboratory data in your notebook on Experiment 11.1. Use
the data to answer these questions:
A. Determine the temperature change of a hot-water sample.
B. Determine the weight of the hot-water sample.
C. How many calories of energy were lost by the hot water?
D. Determine the temperature change of the cold-water sample.
E. Determine the weight of the cold-water sample.
F. How many calories of energy were gained by the cold-water
sample?
G. Should there be a similarity in the answers for Questions C and
F?
Explain.
H. What might account for the difference in the answers for
Questions C and F?
160
EXERCISE 24
CALORIMETRY
DIRECTIONS: On a separate sheet of paper, complete the following.
1.) How many calories are needed to raise 145 g of water from 40.0OC to
65.0OC? (3620 cal)
2.) What temperature change is produced when 898 cal. are absorbed by
134 g of water? (6.70OC)
3.) How many grams of water can be heated from 34.5OC to 89.6OC using
3566 calories? (64.7g)
4.) How many kilocalories are required to change the temperature of 321 g
of water 45.6OC? (14.6 kcal)
5.) How many grams of water can be heated 65.0OC by 4.56 kcal? (70.2g)
6.) What is the final temperature after 90.0 calories are absorbed by 11.0 g
of water at 25.0OC? (33.2OC)
7.) How many calories are released when 344 g of water-cools from 70.0OC
to 30.0OC? (13 800 cal)
8.) What temperature change is produced when 666 g of water gives off
9.88 kcal? (14.8OC)
9.) What is the final temperature when 245 calories are given off by 45.5 g
of water at 46.0OC? (40.6OC)
10.) How many calories are needed if the temperature of 34.6 g of water
changes 3.67OC? (127 cal)
161
EXERCISE 25
ENTHALPY
DIRECTIONS: On a separate sheet of paper do the following:
1.) Calculate the change in enthalpy for 9279 g of oxygen combining with
hydrogen. Product is a liquid.
( -166 000 kJ)
2.) Calculate the change in enthalpy for the burning of 3 940 kg of methane
CH4. All substances are gases. (-1.98X108kJ)
3.) Calculate the change in enthalpy for the formation of 0.000 297 5g of SO3
from sulfur and oxygen. (-0.00164 kJ)
4.) Calculate the change in enthalpy for the formation of 28.54 g of sliver
chloride from sliver and chlorine. (-25.2 kJ)
5.) Calculate the change in enthalpy for the formation of CuI from 354 g of
copper and an excess of iodine. (-378 kJ)
Use the following heats of formation to solve the above
H2O (l) = - 286 kJ/mol
CO2 = - 393.5 kJ/mol
H2O(g) = - 242 kJ/mol
CH4 = -74.8 kJ/mol
SO3 = - 441 kJ/mol
AgCl = -127 kJ/mol
CuI = - 67.8 kJ/mol
162
EXERCISE 26
GAS LAWS
DIRECTIONS: On a separate sheet of paper, answer the following questions.
1.)
842 L of a gas is at 132oC, what is its volume at 152oC? (884 L)
2.)
If 496 ml of a gas is at 3.21 atm, what is its pressure at 694 mL?
(2.29 atm.)
3.)
If 62.7 mL of gas is at 307 K, what is its temp at 37.8 mL? (185 K)
4.)
If a gas is at 444 mm Hg and 0.277 L, what is its volume at 684 mm
Hg? (0.180 L)
5.)
If 697 L of a gas is at 2.11 atm and 419 K, what is its temp. at 3.11
atm. and 546 L? (484K)
6.)
If 0.0677 L of a gas is at 432 PSI and 267 K, what is the pressure of
the gas at 342 K and .148 L? (253 PSI)
7.)
What is the volume of a gas at STP if 4 944 ft3 is at 760 mm and
86.0OC?
(3 760 ft3)
8.)
89 246 cm3 of a gas is at 0.0178 atm, what is the pressure of the gas
at 422 cm3? (3.76 atm.)
9.)
A sample of gas occupies 75.0 ml at 725 mm Hg and 18OC. Calculate
its volume at 801 mm Hg and 298 K? (69.5 mL)
10.) A sample of gas occupies 75.0 mL at 1.15 atm. and 25OC. At what
temp. will it occupy 4.00 L at 1.00 atm?: (13 800 K)
11.) A sample of gas occupies 8.25 L at STP. What will its volume become
at 735 mm Hg and 20O0C? (9.16 L)
12.) A sample of gas occupies 0.575 L at 1.50 atm. and 125OC. What
pressure will it have if its volume changes to 0.300 L and its
temperature to 200OC? (3.42 atm.)
13.
A sample of gas occupies 2.30 L at 825 mmHg and 70oC. What is its
volume at STP? (1.99 L)
163
EXERCISE 27
MORE GAS LAW CALCULATIONS
1. Calculate the volume of 1.00 mole of CO2 at STP.
2. At sea level the average lung capacity for humans is about 4.0 liters. At
37o C (body temperature) how many grams of oxygen gas could your lungs
hold?
3. If 40.0 grams of methane, CH4, is confined to 2500 mL at 200oC, what
pressure does it exert?
4. Calculate the relative rates of diffusion of methane and sulfur dioxide.
5. What volume tank would be needed to contain 3 900 g of helium at 3.45
atm. and 45.0 o C?
6. What temp. would be on a 34.0 L container with 856 g of oxygen at 45
atm?
7. What is the diffusion rate of carbon dioxide to Radon?
8. What pressure is exerted on a 35 678 mL container, if it has 3 899 g of
chlorine at 34.9 o C?
9. What is the diffusion rate of Argon to Krypton?
10. If 5.03 x 1024 atoms of neon are in a 21.9 L tank at 765K, What is the
pressure?
11. If gas A has a pressure of 235mm and gas B has a pressure of 567
mm and gas C has a pressure of 242 mm, what is the total pressure on
the tank?
164
EXERCISE 28
SOLUTION CHEMISTRY
DIRECTIONS: On a separate sheet of paper, do the following:
A. Calculate the molarity of the following solution.
1. 316g of MgBr2 in 859 mL of solution. (2.00M)
2. 4.67g of Ca(C2H3O2)2 in 465 mL of solution. ( .0635M)
3. 4.009g of LiClO3 in 297 mL of solution. ( .149M)
B. Calculate the molality of the following solutions:
1. 199g of NiBr2 in 599g of water. (1.52m)
2. 67.0g of RbI in 2 007g of water. ( .157m)
3. .00987g of AgClO2 in .333kg of water. ( .000169m)
C. Calculate the mass of solute needed to make the following solutions:
1. 1 206mL of .0464M Sc(NO3)3 (12.9g)
2. 134mL of 1.345M VBr3 (52.4g)
3. 5.166L of .397M IrCl4 (685g)
D. Calculate the volume of solution that can be made with the given
information.
1. 90.0g of NH4Cl to make 0.750M solution. (2.24L)
2. 356g of NaNO3 to make 0.00467M solution. (897L)
3. .0493g of AgNO3 to make 0.00000111M solution. (261L)
165
EXERCISE 29
RATE OF REACTION
DIRECTIONS: On a separate sheet of paper, do the
following:
1. Write the rate equation for the following:
a. 2Q + 3Z -- 4Y
b. H2 + Cl2 -- 2HCl
c. 2 C18H38 + 55 O2 -- 36 CO2 + 38 H2O
2. List the five factors that affect the rate of a reaction and
explain in terms of kinetic theory.
3. Differentiate between a heterogeneous and a homogeneous
catalyst.
4. How does a catalyst affect reaction rate?
5. Diagram and label a potential energy curve for a catalyzed and
uncatalyzed exothermic reaction.
166
EXERCISE 30
REACTION RATES
1. What is the rate expression for the following reaction? What is the value
of the rate constant for the reaction?
H2O2 + 2HI  2H2O + I2
Trial
1
2
3
[H2O2]
0.1M
0.1M
0.2 M
[HI]
0.1M
0.2M
0.1 M
Rate ((mol/L)/s)
0.0076
0.0152
0.0152
2. Assume that NO(g) and H2(g) react according to the rate expression: rate
= k[NO]2[H2] . How does the rate change if:
a. the concentration of H2 is doubled?
b. The volume of the enclosing vessel is suddenly halved?
c. The temperature is decreased?
3. For the reaction: H2(g) + I2(g)  2HI(g) , the following data were
obtained:
Experime
nt
Initial [H2]
Initial [I2]
1
2
3
1.0M
1.0M
2.0M
1.0M
2.0M
2.0M
Initial rate of
formation of HI
((mol/L)/s)
0.20
0.40
0.80
a. Write the rate expression for this reaction.
b. What would be the initial rate of formation of the hydrogen iodide if
the initial concentrations of the hydrogen gas and the iodine gas were
each 0.50M?
167
EXERCISE 30 CONTINUED
4. The reaction CH3COCH3 + I2  CH3COCH2I + HI is run under carefully
controlled conditions in the presence of an excess of acid. Write the rate
expression for the reaction using the following data:
Initial
concentration of
CH3COCH3
0.100M
0.0500M
0.0500M
Initial concentration of Initial rate
I2
((mol/L)/s)
1.16x10-7
5.79x10-8
5.78x10-8
0.100M
0.100M
0.500M
5. NO will react with bromine according to the reaction 2NO + Br2  2NOBr.
Experimentally, the following data were obtained:
Trial
Initial [NO]
Initial [Br2]
1
2
3
1.0M
1.0 M
2.0 M
1.0 M
2.0 M
2.0 M
Initial rate of
formation of
NOBr
((mol/L)/s)
0.80
1.60
6.40
a. Write the rate expression for this reaction.
b. Calculate the rate constant, k.
c. What would be the initial rate of formation of NOBr if the initial
concentrations of NO and Br2 were 0.60M and 0.25M, respectively?
d. If the initial concentration of NO were 1.4M, what initial concentration
of Br2 would produce NOBr at an initial rate of 0.75(mol/L)/s?
e. Note that concentrations and rates are all specified “initial”. How
would rate differ if measured later in the reaction? Why would it
differ?
168
EXERCISE 31
LeCHATELIER
For these questions predict what changes will take place resulting from the
changes in concentration, pressure, or heat. Le Chatelier's Principle will be
helpful in making these predictions.
Le Chatelier's Principle: If a system at equilibrium is changed, the system
will partially counteract that change.
Predict the result of these changes on the following equilibrium
system(s):
1. Cu2+ (light blue) + 4NH3 (colorless)  Cu(NH3)42+ (dark blue)
a. Add Cu2+ solution
b. Add NH3 solution
c. Remove NH3
d. Add Cu(NH3)42+
2. 2CO + 02  2C02 + heat
a. Add CO2
b. Remove O2
c. Add CO
d. Add heat
e. Remove heat
3. N2 + 3H2  2NH3 + heat
a. Remove NH3
b. Add H2
c. Remove N2
d. Add heat
e. Remove heat (cool)
4. Cr2O72- (orange) + OH- (colorless)  2Cr2O72- (yellow) + H+ (colorless)
a. Add H+
b. Remove OHc. Add Ba2+ (reacts with CrO42- but not with Cr2O72-)
d. Remove Cr2O72-
169
EXERCISE 32
EQUILIBRIUM
DIRECTIONS: On a separate sheet of paper, do the following.
1.) Write the expressions for the equilibrium constant for each of the following
reversible reactions.
a. H2O + CO
b. N2 + 3H2


c. 4NH3 + 5O2
CO2 + H2
2NH3

4NO + 6H2O
2.) Determine the equilibrium constant for the following reaction if the concentration
of [N2O4] = .00150 and [NO2] = .571
N2O4

2NO2
3.) For the reaction 2SO2 + O2  2SO3, the concentrations of the sulfur oxides are
[SO2] =2.00 and [SO3] = 10.0. What is the concentration of oxygen when the Keq
= 800.0 for the reaction?
4.) For the gaseous reactions below, indicate what happens to the equilibrium
position (shift to left or right) when the indicated stress on the system occurs.
a. remove NH3
b. decrease pressure
N2 + 3H2
c. decrease temp.
d. add a catalyst
CO2 + H2 + heat
e. increase [SO2]
f. increase temp.
2SO2 + O2
g. increase pressure
h. increase [CO]
CO2(g) + C(s) + heat
i. decrease pressure
j. remove N2O4
N2O4(g) + heat
k. increase [H2]
l. increase pressure
H2(g) + Cl2(g)
 2HCl(g) + heat
m. decrease [O2]
n. add a catalyst
N2 + O2 + heat

o. increase [Cl2]
p. decrease pressure


2NH3 + heat

CO + H2O
2SO3 + heat


2CO(g)
2NO2(g)
2NO
PCl3(s) + Cl2(g)  PCl5(s) + heat
170
EXERCISE 33
SOLUBILITY EQUILIBRIA
Directions: On a separate sheet of paper, do the following:
1.
2.
Write the solubility product equilibrium constant for the following:
a.
lead (II) sulfate
b.
silver chloride
c.
calcium phosphate
In a saturated solution of silver oxalate, the concentration of the silver
ion is .000 340 M. Calculate the value for Ksp.
3.
What is the concentration of the silver ion in a saturated solution of
silver chloride? Ksp = 1.80 x 10-10
4.
In a saturated solution of silver chromate, the concentration of silver
ion is 4.00 x 10-6 M and the chromate ion concentration is .0990 M.
What is the value for the Ksp?
5.
The Ksp for lead (II) chloride is 1.60 x 10-5. If the concentration of the
lead ion is .0178 M, what is the concentration of chloride ion?
6.
If the Ksp for AgBr is 4.90 x 10-13, and the concentration of silver and
bromide ions are each 9.1 x 10-5 M in a solution of silver bromide, will
a precipitate form?
7.
The Ksp for HgS is 3x 10-53. Explain what this means in terms of
solubility.
171
EXERCISE 34
ACIDS AND BASES
1. Which of the following are acids, bases, both, or neither?
a. Fb. H2C204
c. S032d. HC4H406e. H2S
f. HC032g. K+
2. For the pairs of reactants given below:
A. Show with a chemical equation what products could form from the
equilibrium mixture.
B. Label which reactant is the acid and which is the base.
C. Predict whether the equilibrium mixture favors the reactants or the
products.
CH3COOH + HS-

NH4+ + F-

HSO4- + NH3

HNO2 + S2-

NH3 + HPO42-

HPO42- + S2-

C6H5COO- + HClO

H2CO3 + OH-

172
EXERCISE 35
MOLARITY CALCULATIONS
1. How many moles of NaOH are dissolved in 5OO ml of .10 M
solution?
2. If 4 moles of substance are dissolved in 500 ml of solution,
what in the molarity?
3. If you had .3 moles of a substance and you needed to make a
solution with a concentration of .6 M, what volume of solution
could you make?
4. If you had 15 g of acetic acid (CH3COOH) and added this
amount to enough water to make 2.5 liters of solution, what
would be the molarity of the solution?
5. How many liters of a .5 M NaCl solution could you make,
using 29 grams of NaCl?
173
EXERCISE 36
ACID-BASE CHEMISTRY
Directions: On a separate sheet of paper, do the following:
1.
Calculate the Ka for each of the following weak electrolytes.
a.
A .00100 M acetic acid solution with
[H3O+]= 1.27 x 10-4 M.
b.
(1.85 X 10-5)
A .0070 M NH4OH solution with
[OH-]= 3.46 x 10-4 M.
c.
(1.80 x 10-5)
A .100 M HCN solution with [H3O+] = 7.85 x 10-6M
(6.16 x 10-10)
2.
Calculate the Ka for the following weak electrolytes:
a.
.100M HF
b.
.0500M HC2H3O2 1.8 % ionized
c.
.200M HAsO2
3.
Calculate the pH and pOH of each of the following:
a.
.0889M HCl
(1.051 and 12.949)
b.
.333M Ba(OH)2
(13.823 and .177)
c.
35.6 g of HNO3 in 355 mL of solution.
(-.202 and 14.202)
d.
45.6 g of HC2H3O2 in 794 mL of solution.
8.50 % ionized
.00550 % ionized
(7.90 x 10-4)
(1.80 x 10-5)
(6.05 x 10-10)
Ka= 1.80 x 10-5
(2.38 and 11.62)
e.
.345M H2X
(.161 and 13.839)
f.
1.09M C6H5NH2
Kb= 4.20 x 10-10
174
(9.330 and 4.670)
exercise 36 continued
4.
5.
Complete the following tables:
[H+]
pH
a.
3.4 x 10-8
_____
_____
b.
1.44 x 10-14
_____
_____
c.
_____
5.667
_____
d.
_____
_____
1.229
e.
_____
-.087
_____
f.
_____
_____
10.876
pOH
What is the pH of a buffer containing .125M HCN and .0117M
NaCN?
Ka= 4.00 x 10-10
6.
(8.369)
Using the buffer system in problem #4, what concentration of
NaCN is needed to set up a buffer with a pH of 7.000 if we start with
0.1000 HCN? (4.00X10-4)
175
EXERCISE 37
WRITING REDOX REACTIONS
Write the equation for the cell reaction that will occur spontaneously in
a cell that contains the two electrodes given.
1.
Zn2+/Zn and I2/I-
2.
Br2/Br- and Pb2+/Pb
3.
Ca2+/Ca and Cu2+/Cu
4.
Mg2+/Mg and Cu2+/Cu
5.
K+/K and Pb2+/Pb
6.
Al3+/Al and Br2/Br-
Write the two half reactions and the final balanced equation for each of
the following, and state whether or not the products or the reactants
are favored:
7. Zn + Ag+ 
8. F2 + Fe 
9. MnO4- + H2O + Fe2+ 
10. AgCl + H2O 
11. Zn + OH- + Ag2O2 + H+ 
12. Mg2+ + Hg 
13. Br2 + Ce3+ 
14. H3O+ + Pb 
15. Cr2O72- + H3O+ + Cd 
176
EXERCISE 38
REDOX BALANCING
Directions: On a separate sheet of paper, balance the following
equations using the bracket method:
Cu(NO3)2 + NO + H2O
1.
Cu + HNO3
2.
Fe(NO3)2 + HNO3
3.
Zn + HNO3
4.
Sb + H2SO4
5.
H2S + H2SO3
6.
Ag + HClO3 + HCl
7.
HNO3 + H2SO4 + Hg
8.
CO + I2O5
9.
Zn + HNO3
10.
FeSO4 + KMnO4 + H2SO4  Fe2(SO4)3 + MnSO4 + K2SO4 + H2O
11.
K2Cr2O7 + KBr + H2SO4 
12.
Cu + H2SO4
13.
Tl2O3
+



Zn(NO3)2 + NO2 + H2O
Sb2(SO4)3 + SO2 + H2O
S + H 2O

AgCl + H2O

Hg2SO4 + NO + H2O

CO2 + I2

HCl
Fe(NO3)3 + NO + H2O




Zn(NO3)2 + NH4NO3 + H2O
Br2 + Cr2 (SO4)3 + K2SO4 + H2O
CuSO4 + SO2 + H2O
TlCl +
Cl2
177
+
H2O
EXERCISE 38, PAGE 2
14. As2O3 + HNO3 + H2O

15.
Pt + HCl + HNO3
H2PtCl6 + NO2 + H2O
16.
H2SO3 + I2 + H2O
17.
SbCl5 + KI
18.
CdS + I2 + HCl
19.
AuCl3 + KI
20.
Ti2(SO4)3 + Fe2(SO4)3

21.
BrO3- + Br- + H+
Br2 + H2O
22.
Cl2 + OH-
23.
Fe + Cr2O72- + H+
24.
Ce4+ + C2O42-
25.
I- + Cr2O72-
H2SO4 + HI

SbCl3 + KCl + I2

CdCl2 + HI + S

AuCl + KCl + I2



H3AsO4 + NO2

Ti(SO4)2 + FeSO4
Cl- + ClO3- + H2O


Cr3+ + Fe3+ + H2O
Ce3+ + CO2
+ H+ 
Cr3+ + H2O + I2
178
EXERCISE 39
NUCLEAR CHEMISTRY
Balancing nuclear equations is somewhat more complicated than
balancing chemical equations because atoms do not have to be
conserved in a nuclear reaction. There are two rules to observe when
balancing nuclear reactions:
1. Charges on the product and reactant sides of the equation must
be equal.
2. Total mass on the product and reactant sides must be equal. This
rule ignores the small defect in mass.
EXAMPLE
Conservation of charge
92U + On = 56Ba + 36Kr + 30n
Conservation of mass
235
U + 1n = 141Ba + 92Kr + 31n
The complete reaction
235
92U
+
1
On
141

56Ba
+
92
36Kr
+3
1
0n
For the reactions below fill in the missing number
(e = electron and n = neutron).
1.
238
2.
9
92U
4Be
+
3.
7
+
4.
239
5.
3Li
92U
226
1
+
???
1

88Ra
0n
2He
1H

239

???
=

4
1
0n
2He
93Np
???
92U
+
???Rn
+
12
???C
+ ___?___
???
-1e
+
4
(electron)
2He
Find the symbol for the element by using the atomic number.
179
EXERCISE 40
NUCLEAR EQUATIONS AND
HALF-LIFE
DIRECTIONS: On a sheet of paper, do the following:
1.
a.
Complete and balance the following nuclear equations.
14
14
6C
7N

218
b.
c.
d.
e.
f.
g.
h.
84Po
+
4
2He

22
11Na
k.
86Rn + ___?___
-1e
+
l.
___?___
m.
140

__?__ 
230
88Ra
j.
10Ne + ___?___
0

140
59Pr
+ ___?___
222

234
90Th
i.
22

226
88Ra
___?___
58Ce + ___?___
24
12Mg +
4

2He
n.
0
-1e
234
91Pa
234
92U

9
5B
1
1H

8
1

13
3
2He

16
8O
+
7
2He
6C
+
0n
___?___
+
___?___
+
___?___
7N
+
___?___
2He
+
___?___
12

1H 
___?___
3
2.
The half-life of iodine-131
is 8 days. Calculate the amount
remaining after 8 days, 16 days,
24 days and 32 days if we start
with 21.8 grams of the
radioisotope.
+ ___?___
3.
What does the half-life
indicate about an unstable
isotope?
180
181