Solution Manual For Chemistry And Chemical Reactivity, 9th Edition

Name: Solution Manual For Chemistry And Chemical Reactivity, 9th Edition
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Solution Manual For Chemistry And Chemical Reactivity, 9th Edition provides the perfect textbook solutions, giving you the help you need to succeed in your studies.

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iii
Table of Contents
Preface ........................................................................................................................................................................... v
Chapter 1 ....................................................................................................................................................................... 1
Basic Concepts of Chemistry
Let’s Review.................................................................................................................................................................. 9
The Tools of Quantitative Chemistry
Chapter 2 ..................................................................................................................................................................... 20
Atoms, Molecules, and Ions
Chapter 3 ..................................................................................................................................................................... 48
Chemical Reactions
Chapter 4 ..................................................................................................................................................................... 65
Stoichiometry: Quantitative Information about Chemical Reactions
Chapter 5 ..................................................................................................................................................................... 94
Principles of Chemical Reactivity: Energy and Chemical Reactions
Chapter 6 ................................................................................................................................................................... 120
The Structure of Atoms
Chapter 7 ................................................................................................................................................................... 137
The Structure of Atoms and Periodic Trends
Chapter 8 ................................................................................................................................................................... 152
Bonding and Molecular Structure
Chapter 9 ................................................................................................................................................................... 174
Bonding and Molecular Structure: Orbital Hybridization and Molecular Orbitals
Chapter 10 ................................................................................................................................................................. 190
Gases and Their Properties
Chapter 11 ................................................................................................................................................................. 216
Intermolecular Forces and Liquids
Chapter 12 ................................................................................................................................................................. 229
The Chemistry of Solids
Chapter 13 ................................................................................................................................................................. 245
Solutions and Their Behavior
Chapter 14 ................................................................................................................................................................. 270
Chemical Kinetics: The Rates of Chemical Reactions
Chapter 15 ................................................................................................................................................................. 295
Principles of Chemical Reactivity: Equilibria
Chapter 16 ................................................................................................................................................................. 316
The Chemistry of Acids and Bases
Chapter 17 ................................................................................................................................................................. 338
Principles of Reactivity: Other Aspects of Aqueous Equilibria
Chapter 18 ................................................................................................................................................................. 369
Principles of Chemical Reactivity: Entropy and Free Energy

iv
Chapter 19.................................................................................................................................................................. 394
Principles of Chemical Reactivity: Electron Transfer Reactions
Chapter 20.................................................................................................................................................................. 419
Environmental Chemistry
Chapter 21.................................................................................................................................................................. 431
The Chemistry of the Main Group Elements
Chapter 22.................................................................................................................................................................. 457
The Chemistry of the Transition Elements
Chapter 23.................................................................................................................................................................. 470
Carbon: More Than Just Another Element
Chapter 24.................................................................................................................................................................. 494
Biochemistry
Chapter 25.................................................................................................................................................................. 504
Nuclear Chemistry

Chapter 1
1
Chapter 1
Basic Concepts of Chemistry
I NSTRUCTOR ’ S NOTES
This chapter discusses a wide range of basic subjects needed for the study of chemistry. How you use the material
will depend on the preparation level of your students. Essentially all of the topics in this chapter would be found in a
typical high school chemistry class. However, if students took that course two or three years earlier, or if their
performance at that time was inconsistent with their present goal of a professional study for science or engineering,
an extensive review might be necessary. Such a review could take two or three lecture periods. If your class consists
of students who have been screened by a placement exam or other process to reasonably ensure that they have good
abilities in high school chemistry, you may assign most or even all of the chapter for outside reading with only a
single class to review and stress key points. The chapter can stand on its own without direct instruction if your
students learned the topics previously. For an intermediate approach you could assign sections 1.1-1.4 for student
reading then take one lecture for the rest of the chapter.
A pre-quiz selected from Chapter 1 Study Questions could help establish what degree of review is needed for your
students.
SUGGESTED DEMONSTRATIONS
1. Density
 Kolb, K. E.; Kolb, D. K. “Method for Separating or Identifying Plastics,” Journal of Chemical Education
1991, 68, 348.
 Franz, D. A. “Densities and Miscibilities of Liquids and Liquid Mixtures,” Journal of Chemical Education
1991, 68, 594.
 Checkai, G.; Whitsett, J. “Density Demonstration Using Diet Soft Drinks,” Journal of Chemical Education
1986, 63, 515.
 Shakhashiri, B. Z. “Density and Miscibility of Liquids,” Chemical Demonstrations: A Handbook for
Teachers of Chemistry; University of Wisconsin Press, 1989; Vol. 3, pp. 229–233.
2. Illustration of “Physical Change”
 Liquid nitrogen is always a favorite of students. We freeze a banana, a hot dog, a flower, or similar object.
3. Illustration of “Chemical Change”
 The first lecture in the course is often begun with a “bang” by setting off several hydrogen-filled balloons
in a darkened lecture room. The demonstration is described in Chemical Demonstrations: A Handbook for
Teachers of Chemistry; University of Wisconsin Press, 1983; Vol. 1, pp. 106–112.

Basic Concepts of Chemistry
2
 Other reactions could be done as well, depending on the facilities available. Possibilities include placing
small pieces of potassium in water, the thermite reaction (Shakhashiri, Volume 1, page 85), or the reaction
of zinc and ammonium nitrate (Shakhashiri, Volume 1, page 51). The latter reaction gives off a large
amount of ZnO dust and other irritating fumes. It is not suitable for a room without good ventilation.
 For fun, and to give some color, as well as talking about our future study of acid-base reactions, we add
aqueous NH 3 to separate flasks containing (i) very dilute acid with phenolphthalein, (ii) Al(NO3) 3, and (iii)
dilute CuSO4 .
 We often have a student contribute a penny to put into concentrated HNO3. The reaction brings up a brief
discussion of oxidation–reduction processes. The NO2 gas generated prompts a discussion of air pollution
problems, as well as the fuels used in the Lunar Lander and in the Space Shuttle. (CAUTION: NO2 is a
very corrosive gas. Use only in a well ventilated room. We do the reaction by putting the penny in a few
milliliters of acid in a 2-L Erlenmeyer flask that is lightly stoppered. This effectively contains the gas.)

Chapter 1
3
SOLUTIONS TO STUDY Q UESTIONS
1.1 (a) C Carbon (c) Cl Chlorine (e) Mg Magnesium
(b) K Potassium (d) P Phosphorus (f) Ni Nickel
1.2 (a) Mn Manganese (c) Na Sodium (e) Xe Xenon
(b) Cu Copper (d) Br Bromine (f) Fe Iron
1.3 (a) barium Ba (d) lead Pb
(b) titanium Ti (e) arsenic As
(c) chromium Cr (f) zinc Zn
1.4 (a) silver Ag (d) tin Sn
(b) aluminum Al (e) technetium Tc
(c) plutonium Pu (f) krypton Kr
1.5 (a) NaCl is a compound; sodium is an element
(b) Sugar is a compound; carbon is an element.
(c) Gold chloride is a compound; gold is an element.
1.6 (a) Pt(NH 3 ) 2Cl2 is a compound; Pt is an element
(b) Copper is an element; copper(II) oxide is a compound
(c) Silicon is an element; sand is a compound
1.7 (a) physical property
(b) chemical property
(c) chemical property
(d) physical property
(e) physical property
(f) physical property
1.8 (a) chemical change
(b) physical change
(c) chemical change
(d) physical change
1.9 (a) Physical properties: color (colorless), physical state (liquid)
Chemical property: reactivity (burns in air)
(b) Physical properties: color (shiny metal, orange), physical state (liquid)
Chemical property: reactivity (aluminum reacts readily with bromine)
1.10 (a) Physical properties: color (white), physical state (solid), density (2.71 g/cm3 )
Chemical properties: reactivity towards acid (reacts to produce gaseous carbon dioxide)
(b) Physical property: color (gray zinc, purple iodine, white compound)
Chemical property: reactivity (zinc and iodine react to give a white compound)

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Basic Concepts of Chemistry
4
1.11. Mechanical energy is used to move the lever, producing electrical and radiant energy
1.12. Mechanical energy, electrical energy, radiant energy
1.13 (a) Kinetic energy
(b) Potential energy
(c) Potential energy
(d) Potential energy
1.14 (a) Potential energy converted to kinetic energy
(b) Kinetic energy converted to potential energy
(c) Potential energy converted to kinetic energy
(d) Kinetic energy converted to potential energy
1.15 (a) Qualitative observations: blue-green color, solid physical state
Quantitative observations: density of 2.65 g/cm3
, mass of 2.5 g
(b) Mass is an extensive property; color, physical state, and density are intensive properties
(c) 2.5 g ·
3
1 cm
2.65 g = 0.94 cm3
1.16 (a) Qualitative observations: shiny golden metallic appearance, crystals in form of perfect cubes
Quantitative observations: length of 0.40 cm on a side, mass of 0.064 g
(b) Mass and length are extensive properties; color, luster, and crystalline form are intensive properties
(c) 0.064 g/(0.40 cm) 3 = 1.0 g/cm3 = 1.0 g/mL
1.17 Observation (c) identifies a chemical property.
1.18 Observations (a) and (b) identify chemical properties.
1.19 Calcium: Ca Fluorine: F
The shape of the fluorite crystals can be described as interwoven cubes.
The overall shape of the crystals indicates that the ions in the solid matrix arrange themselves with
alternating calcium and fluoride ions to produce the crystal appearance.
1.20 (a) Copper, Cu Carbon, C Oxygen: O
(b) Oxygen is a gas, while copper, carbon, and azurite are solids at room temperature. Oxygen is
colorless, while copper has a reddish color and carbon is gray. The gemstone is a bluish color.
1.21 The water can be removed by boiling the solution until all the water is converted to water vapor; the solid
left behind will be NaCl.

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Chapter 1
5
1.22 The non-uniform appearance of the mixture indicates that samples taken from different regions of the
mixture would be different—a characteristic of a heterogeneous mixture. Recalling that iron is attracted to
a magnetic field while sand is generally not attracted in this way suggests that passing a magnet through
the mixture would separate the sand and iron.
1.23 Physical changes: a, b, d
Chemical changes: c
1.24 Physical changes: a, d
Chemical changes: b, c
1.25 The large colorless block of salt represents the macroscopic view. The spheres represent the microscopic or
particulate view. If one can imagine producing multiple “copies” of the particulate view, the macroscopic
view will result.
1.26 The orange solid and liquid in the bowl and the orange liquid and gas in the round bottom flask represent
the macroscopic views. The spheres represent the particulate view. The particulate view of the solid
displays atoms of bromine tightly packed to produce the solid. The liquid view displays molecules of Br 2
with space separating the individual molecules. The gas view displays molecules of Br2; however, the
molecules are further apart than in the liquid view.
1.27 The plastic (with a much lower density than CCl4) will float. The aluminum (which is more dense than
CCl4 ) will sink.
1.28 Liquids: mercury and water Solid: copper
Of the substances shown, mercury is most dense and water is least dense.
1.29 (a) mixture
(b) mixture
(c) element
(d) compound
1.30 (a) mixture
(b) compound
(c) mixture
(d) mixture

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Basic Concepts of Chemistry
6
1.31 (a) solid iron (b) liquid water (c) water vapor
1.32 (a) water vapor, helium (b) liquid water, solid aluminum (c) brass
1.33 The three liquids will be layered, with hexane floating on top of water, which in turn floats on top of the
perfluorohexane. HDPE will lie at the top of the water layer. PVC will lie on top of the perfluorohexane
layer. Teflon will sink to the bottom of the perfluorohexane layer.
1.34 Melting point. Sugar melts around 160-186 °C while salt melts at 800 °C.
1.35 HDPE will float in liquids with a density higher than 0.97 g/mL, ethylene glycol, water, acetic acid, and
glycerol.
1.36 The sample's density and melting point could be compared with that of silver to prove whether or not the
sample is silver.
1.37 Milk is mostly water. When water freezes its volume increases (its density decreases). When the milk
froze, the increase in volume was so great that it pushed out of the bottle.
1.38 The mass of the object is determined and then the volume is determined by submersion in a known volume
of liquid. The increase in volume would be equal to the volume of the irregularly shaped object. The
density could be calculated by dividing its mass by its volume.
1.39 If too much sugar is excreted, the density of urine will increase. If too much water is excreted, the density
of urine will decrease.
1.40 One could check for an odor, check the boiling or freezing point, or determine the density. If the density is
approximately 1 g/cm3 at room temperature, the liquid could be water. If it boils at about 100 °C and
freezes at about 0 °C, that would be consistent with water. To check for the presence of salt, boil the liquid
away. If a substance remains, it could be salt, but further testing would be required.

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Chapter 1
7
1.41 (a) Solid potassium reacts with liquid water to produce gaseous hydrogen and a homogeneous mixture of
potassium hydroxide in liquid water.
(b) The reaction is a chemical change.
(c) Potassium and water are reactants, hydrogen and potassium hydroxide are products.
(d) Among the qualitative observations are (i) the reaction is violent; and (ii) heat and light (a purple
flame) are produced.
1.42
Least dense liquid: water
Most dense liquid: mercury
Medium density: carbon tetrachloride
1.43 Any balloon filled with a gas having a density less than 1.12 g/L will float in air. Helium and neon balloons
will float.
1.44 A copper-colored metal could be copper, but it may also be an alloy of copper, for example, brass or
bronze. Testing the material’s density and melting temperature would be one way to find out if it is copper.
1.45 Physical change
1.46 (a) Mass CHCl3 = 10.0 mL · 1.492 g/mL = 14.9 g
Mass CHBr 3 = 5.0 mL · 2.890 g/mL = 14 g
Mass of solution = 14.9 g + 14 g = 29 g
Density = mass of solution/volume of solution = 29 g/15.0 mL = 1.9 g/mL
(b) Mass of solution = 2.07 g/mL · 20.0 mL = 41.4 g
Also, mass of solution = mass CHCl3 + mass CHBr 3
Let X = volume CHCl3 and (20.0 mL − X) = volume CHBr 3
Then, mass of solution = (1.492 g/mL)X + (2.890 g/mL)(20.0 mL – X) = 41.4 g
X = 11.7 mL CHCl3
(20.0 – X) = 8.3 mL CHBr 3
To obtain 20.0 mL of solution with d = 2.07 g/cm3 , mix 11.7 mL CHCl3 and 8.3 mL CHBr 3 .

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Let’s Review
9
Let’s Review
The Tools of Quantitative Chemistry
I NSTRUCTOR ’ S NOTES
This section reviews a wide range of mathematical topics that will be needed by the student in the course, so
mastery of this material should be stressed. The need for accurate measurement of physical quantities and the
correct recording of this information can be characterized as necessary for reasons of safety using medical (drug
dosage) and engineering examples (material composition and properties). Mistakes in this area can mean the
difference between life and death.
A pre-quiz selected from the Study Questions at the end of this section could help establish what degree of review is
needed for your students.
SUGGESTED DEMONSTRATIONS
1. Units of Measurement
 Add a few drop of bromcresol green to a 2-L flask before the lecture. On filling with water during lecture,
the water becomes blue. When the flask is almost full, it is topped off with dilute HCl, and the solution
turns yellow. Next, the water is poured into an ordinary 1-L flask or beaker that already contains some
dilute base, and the solution turns blue again. This is then poured into a graduated cylinder containing
some phenolphthalein. We use this sequence to comment on the relative accuracies of the different types of
glassware. We take some containers such as soda cans so students can connect metric units with familiar
objects.
 A weighed piece of fruit or some other solid gives some meaning to mass expressed in grams.
 Earley, C. W. “A Simple Demonstration for Introducing the Metric System to Introductory Chemistry
Classes,” Journal of Chemical Education 1999, 76, 1215.
2. Demonstrations Using Significant Figures
 Kirksey, H. G. “Significant Figures: A Classroom Demonstration,” Journal of Chemical Education 1992,
69, 497.
 Abel, K. B.; Hemmerlin, W. M. “Significant Figures,” Journal of Chemical Education 1990, 67, 213.

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The Tools of Quantitative Chemistry
10
SOLUTIONS TO STUDY Q UESTIONS
LR.1 (25 °C + 273.15 °C) 1 K
1 C = 298 K
LR.2 (5.5  10 3 °C + 273.15 °C) 1 K
1 C = 5.8  10 3 K
LR.3 (a) 289 K (b) 97 °C (c) 310 K
LR.4 (a) –196 °C (b) 336 K (c) 1177 °C
LR.5 42.195 km · 1000 m
1 km = 4.2195  10 4 m
42.195 km · 0.62137 miles
1 km = 26.219 miles
LR.6 19 cm · 10 mm
1 cm = 190 mm
19 cm · 1 m
100 cm = 0.19 m
LR.7 2.5 cm  2.1 cm = 5.3 cm2
5.3 cm2 ·
2
1 m
100 cm
 
 
  = 5.3  10 –4 m2
LR.8
2
2 211.8 cm
Area = r = = 109 cm
2
 
   
 
2
2 –2 21 m
109 cm = 1.09 10 m
100 cm
 
  
 
LR.9 250. mL ·
3
1 cm
1 mL = 250. cm3
250. mL · 3
1 L
10 mL = 0.250 L
250. mL ·
3
1 cm
1 mL ·
3
1 m
100 cm
 
 
  = 2.50  10 –4 m3
250. mL · 3
1 L
10 mL ·
3
1 dm
1 L = 0.250 dm3

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Let’s Review
11
LR.10 1.5 L ·
3
10 mL
1 L = 1.5  10 3 mL
1.5 L ·
3
10 mL
1 L ·
3
1 cm
1 mL = 1.5  10 3 cm3
1.5  10 3 cm3 ·
3
1 dm
10 cm
 
 
  = 1.5 dm3
LR.11 2.52 kg ·
3
10 g
1 kg = 2.52  10 3 g
LR.12 2.265 g · 3
1 kg
10 g = 2.265  10 –3 kg
2.265 g ·
3
10 mg
1 g = 2.265  10 3 mg
LR.13 500. mL ·
3
1 cm
1 mL · 3
1.11 g
1 cm = 555 g
LR.14
3
1 cm
2.365 g 10.5 g
 = 0.225 cm3
LR.15 density = 2.361 g
(2.35 cm 1.34 cm 0.105 cm)  = 7.14 g/cm3
The metal is (c) Zinc
LR.16 600 g H2O ·
3
1 cm
0.995 g = 600 cm3
600 g Pb ·
3
1 cm
11.35 g = 50 cm3
Given equal masses the less dense substance (water) will have a greater volume.
LR.17 1200 Cal · 3
10 cal
1 Cal · 4.184 J
1 cal = 5.0  10 6 J
LR.18 1670 kJ · 3
10 J
1 kJ · 1 cal
4.184 J · 3
1 Cal
10 cal = 399 Cal
LR.19 170 kcal · 3
10 cal
1 kcal · 4.184 J
1 cal · 3
1 kJ
10 J = 710 kJ
The food product with 170 kcal per serving has a greater energy content per serving.

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The Tools of Quantitative Chemistry
12
LR.20 130 Calories × 4.184 J/cal × 1000 cal/Calorie = 5.4 × 105 J = 5.4 × 10 2 kJ. The juice provides more total
energy.
5.4 × 10 2 kJ/335 mL = 1.6 kJ/mL 630 kJ/295 mL = 2.1 kJ/mL. The juice provides more energy per
milliliter.
LR.21 (a) Method A average = 2.2 + 2.3 + 2.7 + 2.4
4 = 2.4 g/cm3
Method B average = 2.703 + 2.701 + 2.705 + 5.881
4 = 3.480 g/cm3
For method B the reading of 5.811 can be excluded because it is more than twice as large as all other
readings. Using only the first three readings, average = 2.703 g/cm3
.
(b) Method A: Percent error = 2.4 – 2.702 100 %
2.702  = 10 %
Method B: Percent error = 2.703 – 2.702 100 %
2.702  = 0.04 % (omitting data point)
(c) Method A: Standard deviation =
2 2 2 2
0.2 + 0.1 + 0.3 + 0
3 = 0.2 g/cm3
Method B: Standard deviation =
2 2 2
0 + 0.002 + 0.002
2 = 0.002 g/cm3 (omitting data point)
(d) Before excluding a data point for B, method A is more accurate and more precise. After excluding the
fourth data point for B, this method gives a more accurate and more precise result.
LR.22 (a) Student A: Average = 135 °C Percent error = 135 – 135 100 %
135  = 0 %
Student B: Average = 138 °C Percent error = 138 – 135 100 %
135  = 2 %
(b) Student B is more precise; Student A is more accurate
LR.23 (a) 5.4  10 –2 g (2 significant figures) (c) 7.92  10 –4 g (3 significant figures)
(b) 5.462  10 3 g (4 significant figures) (d) 1.6  10 3 mL (2 significant figures)
LR.24 (a) 1623 (4 significant figures) (c) 0.0632 (3 significant figures)
(b) 0.000257 (3 significant figures) (d) 3404 (4 significant figures)
LR.25 (a) 9.44  10 –3 (c) 11.9
(b) 5.694  10 3 (d) 0.122
LR.26 (a) 2.44  10 8 (c) 0.133
(b) 4.85  10 –2 (d) 0.0286

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Let’s Review
13
LR.27
y = 0.1637x + 0.0958
0
1
2
3
4
5
6
7
0 20 40
Number of kernels
slope = 0.1637 g/kernel The slope represents the average mass of a kernel.
Using the straight-line equation, mass = (0.1637 g/kernel)(20 kernels) + 0.0958 g = 3.370 g
20.88 g = (0.1637 g/kernel)(number of kernels) + 0.0958 g
number of kernels = 127
LR.28 (a) 0.21
(b) 5.6
(c) slope = 5.6 – 4.0
0.30 – 0.21 = 18; intercept = 0.20
(d) y = 18x + 0.20 = (18)(1.0) + 0.20 = 18
LR.29 (a) slope = 20.00 – 4.00
0.00 – 4.00 = –4.00, intercept = 20.00, y = –4.00x + 20.00
(b) y = –4.00x + 20.00 = (–4.00)(6.0) + 20.00 = –4.00
LR.30 (a)
y = 4E-05x - 0.0416
0
1
2
3
4
5
6
7
0.E+00 1.E+05 2.E+05
1/speed
(b) y = 4.0  10 –5
x – 0.0416
y-intercept = –0.0416
LR.31 C = 0.0823
LR.32 n = 1.63
LR.33 T = 295

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