Appendix

Appendix C

A-9

Answers to Practice Problems

D

A P P E N D I X

Answers to Selected Questions and Problems Chapter 1 – Matter and Energy 1.1

1.3 1.5 1.7 1.9 1.11 1.13 1.15 1.17

1.19

1.21 1.23

1.25 1.27 1.29

1.31 1.33 1.35 1.37 1.39 1.41 1.43

1.45

(a) mass; (b) chemical property; (c) mixture; (d) element; (e) energy; (f) physical property; (g) liquid; (h) density; (i) homogeneous mixture; (j) solid (a) 2.95 × 104; (b) 8.2 × 10−5; (c) 6.5 × 108; (d) 1.00 × 10−2 (a) 0.0000186; (b) 10,000,000; (c) 453,000; (d) 0.0061 (a) 6.2 × 103; (b) 3.5 × 107; (c) 2.9 × 10−3; (d) 2.5 × 10−7; (e) 8.20 × 105; (f) 1.6 × 10−6 (a) 3; (b) 2; (c) 4; (d) 2; (e) 3 (a) 1.5; (b) 1.5; (c) 14; (d) 1.20 (a) 2.8; (b) 0.28; (c) 2.8; (d) 0.049 (a) 1.21; (b) 0.204; (c) 1.84; (d) 42.2; (e) 0.00710 (a) 0.036 m; (b) 3.57 × 105 g; (c) 0.07650 L; (d) 8.4670 cm; (e) 5.97 × 10−7 m; (f) 92.7 cm; (g) 7.62 × 104 g; (h) 865 L; (i) 17.5 in; (j) 0.5214 lb; (k) 2.1 qt (a) 0.589 mi; (b) 14.9 lb; (c) 6.617 × 10−2 gal; (d) 2.30 × 102 mL; (e) 4.50 × 103 nm; (f) 0.952 m; (g) 3.12 × 102 kg; (h) 5 × 102 mL; (i) 4.10 ft; (j) 1.19 × 10−3 lb; (k) 6.6 × 10−9 gal (a) 7.38 × 104 ft/min; (b) 1.50 in3; (c) 0.697 lb/in3 (a) homogeneous mixture if the dye is evenly mixed into the water; (b) element; (c) homogeneous mixture; (d) heterogeneous mixture Gasoline, automobile exhaust, oxygen gas, and the iron pipe are matter. Sunlight is energy. Elements are composed of only one type of atom. Compounds are made up of two or more different elements. Metals are lustrous (shiny) and conduct heat and electricity. In addition you can form wires with metals (ductile) and you can make foil out of them by hitting them with a hammer (malleable). (a) titanium; (b) tantalum; (c) thorium; (d) technetium; (e) thallium (a) boron; (b) barium; (c) beryllium; (d) bromine; (e) bismuth (a) nitrogen; (b) iron; (c) manganese; (d) magnesium; (e) aluminum; (f) chlorine (a) Fe; (b) Pb; (c) Ag; (d) Au; (e) Sb Ir is the symbol for the element iridium. Iron’s symbol is Fe. The correct symbol is No. The only pure substance is the salt in the salt shaker (if it is not iodized salt). Hamburger: heterogeneous mixture. Salt: pure substance. Soft drink: heterogeneous mixture (until it goes flat). Ketchup: heterogeneous mixture.

1.47 1.49 1.51

The chemical formula is N2O4. Image D represents a mixture of an element and a compound. The elements are O2, P4, and He. Fe2O3, NaCl, and H2O are compounds.

1.53

Liquid state

1.55 1.57 1.59 1.61 1.63 1.65 1.67 1.69 1.71 1.73 1.75 1.77

1.79 1.81 1.83

gas (a) gas; (b) liquid; (c) solid solid O2(aq) physical properties (a) 0.045 g; (b) 1.6 × 10−3 oz; (c) 9.9 × 10−5 lb (a) 0.10 mg; (b) 1.0 × 102 μg; (c) 1.0 × 10−7 kg (a) 1.2 × 103 mL; (b) 1.2 × 103 cm3; (c) 1.2 × 10−3 m3 1.6 × 102 mL, 0.16 L 1.3 g/cm3, 1.3 g/mL 38.5 mL The molecules in the liquid state are closer together than molecules in the gas state. More matter in the same volume means that the density is higher. oil (least) < plastic < water (greatest) 329 K Scale

1.89

Freezing

Boiling

Difference

Celsius

0°C

100°C

100°C

Kelvin

213.15 K

373.15 K

100 K

32°F

212°F

180°F

Fahrenheit 1.85 1.87

Solid state

no Physical properties are (a) mass, (b) density, and (e) melting point. Chemical properties are (c) flammability, (d) resistance to corrosion, and (f) reactivity with water. Physical changes are (a) boiling acetone, (b) dissolving oxygen gas in water, and (e) screening rocks from sand. Chemical changes are (c) combining hydrogen and oxygen to make water, (d) burning gasoline, and (f) conversion of ozone to oxygen.

H2

A-9

Appendix D

A-10

1.91

Symbolic: Cl2(g) gram is shown.

Answers to Selected Questions and Problems

Cl2(l). A molecular-level dia-

1.121

Condensation

Gas

1.93 1.95

1.119

Liquid

chemical change physical change

1.123 1.125 1.127 1.129 1.131

You could do some library research and find the densities of many different woods. By placing samples of the different woods in water, you can determine if the theory is correct. At high altitude, the air pressure is lower. As a result, when a balloon rises, it expands. Since the mass of air in the balloon has not changed but the volume has increased, the density of the balloon is lower. zinc (a) He; (b) Ne; (c) Ar; (d) Kr; (e) Xe; (f) Rn physical 8 mg 1.6 × 10−4 lb/oz; 0.02 to 0.03 lb

Chapter 2 – Atoms, Ions, and the Periodic Table 2.1

2.3 2.5 2.7

1.97 1.99

1.101 1.103 1.105

1.107

1.109

1.111 1.113

1.115 1.117

physical change Anna and Bill would have observed kinetic energy from the movement of the welder and the motion of the sparks. The sparks would have glowed, indicating heat, light, and chemical energy. The molecules in image A have greater kinetic energy because they are moving faster. Any object that would move if allowed has potential energy (e.g., a picture hanging on a wall). The people walking, the wheel chair rolling, and the suitcase being pushed all have kinetic energy. The people, the wall art, and objects on the tables all have potential energy. Many objects have potential and kinetic energy. Consider as an example a car going down the road. A car going up a hill is converting chemical energy in the fuel to mechanical energy to reach the top. At the top of the hill, the car has potential energy. If it rolls down the hill, it gains kinetic energy but loses potential energy. Other energy conversions can be found while driving a car. When the batter swings the bat, potential energy (metabolic energy) is converted into kinetic energy (moving bat). The bat strikes the ball, and the kinetic energy of the bat is transferred to the kinetic energy of the ball. As the ball leaves the bat, it rises against the Earth’s gravity and kinetic energy is converted to potential energy. When the ball starts dropping again, potential energy is converted to kinetic energy. BMI = 21.7 kg/m2; healthy As the water leaves the top of the fountain, it possesses kinetic energy (going up). That energy is converted to potential energy. As the water falls, the potential energy converts back into kinetic energy. It is used to explain and predict scientific results. (a) hypothesis; (b) observation; (c) theory; (d) observation

2.9 2.11 2.13 2.15

(a) neutron; (b) law of conservation of mass; (c) proton; (d) main-group element; (e) relative atomic mass; (f) mass number; (g) isotope; (h) cation; (i) subatomic particle; (j) alkali metal; (k) periodic table Dalton used the laws of conservation of mass (Lavoisier) and definite proportions (Proust). They differ in their atomic masses and chemical properties. Compounds contain discrete numbers of atoms of each element that form them. Because all the atoms of an element have the same relative atomic mass, the mass ratio of the elements in a compound is always the same (law of definite proportions). No. Hydrogen atoms are not conserved. A molecule of hydrogen, H2, should be added to the left panel. Thomson’s cathode-ray experiment electrons The nucleus of helium has two protons and two neutrons. Two electrons can be found outside the nucleus.

proton neutron 2 e–

2.17 2.19

2.21 2.23 2.25 2.27

neutron Carbon has six protons. The relative atomic mass of a carbon atom is 12.01 amu indicating the presence of six neutrons. Protons and neutrons have approximately equal masses, so the nuclear mass is approximately two times the mass of the protons. (a) 1; (b) 8; (c) 47 protons number of protons (b) atomic number

A-11

Appendix D Answers to Selected Questions and Problems

2.29

2.31

The following table displays the atomic, neutron, and mass numbers for the isotopes of hydrogen. 1 1H

2 1H

3 1H

Atomic number

1

1

1

Neutron number

0

1

2

Mass number

1

2

3

(a) 158 O (b) 109 47 Ag 35 (c) 17 Cl

Protons

Neutrons

Electrons

8

7

8

47

62

47

17

18

17

Protons

Neutrons

(a) 56 26 Fe (b) 39 19 K

19

(c) Copper-65 or 65 29 Cu

29

26

2.47

Electrons

(a)

28

30

27

20

(b)

60 + 28 Ni

28

32

27

36

+ (c) 61 28 Ni

28

33

27

(d)

62 + 28 Ni

28

34

27

(e)

64 + 28 Ni

28

36

27

30

Electrons

(a) 23 11Na

11

12

11

(b) 56 25 Mn

25

31

25

8

10

8

9

10

9

They differ in the number of electrons. (a) an anion with a 1- charge is formed; (b) a cation with a 2+ charge is formed (a) Zn2+, cation; (b) P3−, anion

(a) Zn2+

Protons

Electrons

30

28

(b) F-

9

10

+

1

0

(c) H 2.51

2.81 2.83 2.85 2.87 2.89 2.91 2.93 2.95 2.97 2.99 2.101 2.103 2.105 2.107

Protons

Neutrons

Electrons

(a)

37 17 Cl

17

20

18

2.109

(b)

25 2+ 12 Mg

12

13

10

2.111

(c)

13 37N

7

6

10

20

20

18

2+ (d) 40 20 Ca

2.59 2.61

Neutrons

Neutrons

2.49

2.53 2.55 2.57

Protons 58 + 28 Ni

Protons

(c) 188O (d) 199 F

(a) 2 amu; (b) 238 amu The mass of D2 is two times the mass of H2. about 40 amu The numerical values of masses of individual atoms are very small when measured on the gram scale. The size of the atomic mass unit allows us to make easier comparisons and calculations of masses of molecules. The mass number is the sum of the number of protons and neutrons and is always an integer value. In contrast, the mass of an atom is the actual measurement of how much matter is in the atom and is never exactly an integer value (except carbon-12). A mass spectrometer is used to determine the mass of an atom. The mass number of an atom is the sum of the number of protons and neutrons. calcium-40 21.50 amu (a) 58Ni; (b) 64Ni; (c) 58; (d)

seven protons and six neutrons

2.41

2.43 2.45

2.73

2.75 2.77 2.79

(a) 31 H; (b) 94 Be; (c) 31 15 P

2.37

2.39

2.71

(a) Z = 18, N = 18, A = 36; (b) Z = 18, N = 20, A = 38; (c) Z = 18, N = 22, A = 40

2.33

2.35

2.63 2.65 2.67 2.69

potassium, K copper, Cu 7 Li has three protons, three electrons, and four neutrons. 7 + Li has only two electrons, and 6Li has three neutrons. Otherwise they are the same as 7Li. Lithium-6 differs the most in mass. 19 protons and 18 electrons The atomic mass unit is defined as one-twelfth the mass of one carbon-12 atom.

2.113 2.115 2.117 2.119

2.121

10,810 amu or 1.081 × 104 amu 2500 amu of boron contains more atoms. (a) K; (b) Br; (c) Mn; (d) Mg; (e) Ar; (f) Br, K, Mg, Al, Ar chlorine, Cl titanium, Ti (a) metal; (b) nonmetal; (c) metal; (d) metalloid (a) main group; (b) main group; (c) main group; (d) actinide; (e) transition metal Group VIIA, the halogen family, all occur as diatomic molecules. neon Group VIIIA (18), the noble gases electrons (a) group IA (1); (b) group IIA (2); (c) group VIIA (17); (d) group VIA (16) (a) Na+; (b) O2−; (c) S2−; (d) Cl−; (e) Br− All alkali metal elements (group 1 or IA excluding hydrogen) The mass of oxygen added to form Fe2O3 causes an increase in the mass. The mass ratios of Zn/S are the same for the two samples (within the significant figures given). The mass ratio is approximately 2.0:1.0. Electrons have charge and were readily studied in cathode-ray tubes. nickel-60 19 protons and 20 neutrons The most abundant isotopes of cobalt have masses greater than the masses of the most abundant isotopes of nickel. As a result, the relative atomic mass of cobalt is greater than that of nickel. When there are many isotopes, some of the isotopes can be present in very low abundance. As a result, their masses cannot be determined as accurately and their percentage

A-12

2.123 2.125 2.127

2.129 2.131 2.133

Appendix D Answers to Selected Questions and Problems

contributions are also less well known. Both factors result in a decrease in the precision of the calculated relative atomic mass. 127 carbon 20% boron-10 and 80% boron-11; the relative atomic mass (10.81 amu) is about 80% of the difference between the two isotopes. Br2(l) hydrogen The energy required to force a proton (positive charge) into the nucleus is too great.

3.45 3.47

3.49 3.51 3.53

(a) CaSO4; (b) BaO; (c) (NH4)2SO4; (d) BaCO3; (e) NaClO3 (a) Co2+, cobalt(II) chloride; (b) Pb4+, lead (IV) oxide or plumbic oxide; (c) Cr3+, chromium(III) nitrate; (d) Fe3+, iron(III) sulfate or ferric sulfate (a) CoCl2; (b) Mn(NO3)2; (c) Cr2O3; (d) Cu3(PO4)2 ferrous nitrate and ferric nitrate

Chapter 3 – Chemical Compounds 3.1 3.3 3.5 3.7 3.9 3.11 3.13 3.15 3.17 3.19 3.21 3.23 3.25

Ca2+

Fe2+

K+

Cl–

CaCl2 calcium chloride

FeCl2 iron(II) chloride

KCl potassium chloride

–

CaO calcium oxide

FeO iron(II) oxide

K2O potassium oxide

–

Ca(NO3)2 calcium nitrate

Fe(NO3)2 iron(II) nitrate

KNO3 potassium nitrate

SO32

–

CaSO3 calcium sulfite

FeSO3 iron(II) sulfite

K2SO3 potassium sulfite

OH–

Ca(OH)2 calcium hydroxide

Fe(OH)2 iron(II) hydroxide

KOH potassium hydroxide

ClO3– Ca(ClO3)2 calcium chlorate

Fe(ClO3)2 iron(II) chlorate

KClO3 potassium chlorate

O2

(a) formula unit; (b) strong electrolyte; (c) molecular compound; (d) acid; (e) nonelectrolyte; (f) oxoanion (a) ionic; (b) ionic; (c) molecular (a) molecular; (b) ionic; (c) molecular; (d) ionic (a) molecular; (b) ionic; (c) ionic; (d) molecular The ionic compound LiF would have the highest melting point. (a) Na+, sodium ion; (b) K+, potassium ion; (c) Rb+, rubidium ion (a) Ca2+, calcium ion; (b) N3−, nitride ion; (c) S2−, sulfide ion NO2−, nitrite ion (a) sulfate ion; (b) hydroxide ion; (c) perchlorate ion (a) N3−; (b) NO3−; (c) NO2− (a) CO32−; (b) NH4+; (c) OH−; (d) MnO4− SO3−, sulfite ion IO3−, iodate ion

NO3

3.27

Mn2+

Al3+

NH4+

Cl–

MnCl2 manganese(II) chloride

AlCl3 aluminum chloride

NH4Cl ammonium chloride

–

MnO manganese(II) oxide

Al2O3 aluminum oxide

(NH4)2O ammonium oxide

–

Mn(NO3)2 manganese(II) nitrate

Al(NO3)3 aluminum nitrate

NH4NO3 ammonium nitrate

SO32

–

MnSO3 manganese(II) sulfite

Al2(SO3)3 aluminum sulfite

(NH4)2SO3 ammonium sulfite

OH–

Mn(OH)2 manganese(II) hydroxide

Al(OH)3 aluminum hydroxide

NH4OH ammonium hydroxide

ClO3–

Mn(ClO3)2 manganese(II) chlorate

Al(ClO3)3 aluminum chlorate

NH4ClO3 ammonium chlorate

O2

NO3

3.29 3.31 3.33 3.35 3.37

3.39

3.41 3.43

(a) BaCl2; (b) FeBr3; (c) Ca3(PO4)2; (d) Cr2(SO4)3 (a) K+ and Br−; (b) Ba2+ and Cl −; (c) Mg2+ and PO43−; (d) Co2+ and NO3− (a) Fe2+ will form FeO and Fe3+ will form Fe2O3. (b) Fe2+ will form FeCl2 and Fe3+ will form FeCl3. (a) 2−; (b) 2+ For each “compound” written, the charges do not balance (compounds have no net charge). (a) too many chloride ions, NaCl; (b) not enough potassium ions, K2SO4; (c) There should be three nitrate ions and one aluminum ion, Al(NO3)3. (a) magnesium chloride; (b) aluminum oxide; (c) sodium sulfide; (d) potassium bromide; (e) sodium nitrate; (f) sodium perchlorate MnSO4 and CoCl2 (a) copper(I) oxide or cuprous oxide; (b) chromium(II) chloride; (c) iron(III) phosphate or ferric phosphate; (d) copper(II) sulfide or cupric sulfide

3.55 potassium

iron(III)

strontium

iodide

KI

FeI3

SrI3

oxide

K2O

Fe2O3

SrO

sulfate

K2SO4

Fe2(SO4)3

SrSO4

nitrite

KNO2

Fe(NO2)3

Sr(NO2)2

acetate

KCH3CO2

Fe(CH3CO2)3

Sr(CH3CO2)2

KClO

Fe(ClO)3

Sr(ClO)2

hypochlorite

Appendix D Answers to Selected Questions and Problems

aluminum

cobalt(II)

lead(IV)

3.113

iodide

AlI3

CoI2

PbI4

3.115

oxide

Al2O3

CoO

PbO2

sulfate

Al2(SO4)3

CoSO4

Pb(SO4)2

nitrite

Al(NO2)3

Co(NO2)2

Pb(NO2)4

acetate

Al(CH3CO2)3

Co(CH3CO2)2

Pb(CH3CO2)4

hypochlorite Al(ClO)3 Co(ClO)2 Pb(ClO)4 3.57 AgCl 3.59 NF3, P4O10, C2H4Cl2 3.61 (a) phosphorus pentafluoride; (b) phosphorus trifluoride; (c) carbon monoxide; (d) sulfur dioxide 3.63 (a) SF4; (b) C3O2; (c) ClO2; (d) SO2 3.65 The central image represents H3PO4. 3.67 (a) hydrofluoric acid; (b) nitric acid; (c) phosphorous acid 3.69 (a) HF; (b) H2SO3; (c) HClO4 3.71 hydrogen ions and nitrate ions 3.73 KI, Mg(NO3)2, NH4NO3 3.75 ionic; TiO2, ZnO 3.77 molecular; CO2 and N2O 3.79 potassium sulfide; sodium sulfate, sulfur dioxide 3.81 (a) nitrogen trioxide, molecular; (b) nitrate ion, an ion; (c) potassium nitrate, ionic; (d) sodium nitride, ionic; (e) aluminum chloride, ionic; (f) phosphorus trichloride, molecular, (g) titanium(II) oxide, ionic; (h) magnesium oxide, ionic 3.83 (a) Na2CO3; (b) NaHCO3; (c) H2CO3; (d) HF; (e) SO3; (f) CuSO4; (g) H2SO4; (h) H2S 3.85 HCl(aq) is dissolved in water and ionized; HCl(g) is not ionized; HCl(aq), hydrochloric acid; HCl(g) hydrogen chloride 3.87 (a) No prefixes with ionic compounds. (b) Calcium’s charge should not be stated. (c) Copper’s charge, 2+, should be stated. (d) Prefixes are used for molecular compounds. 3.89 (a) Two potassium ions are needed to balance S2−. (b) Only one Co2+ is needed to balance charge. (c) Nitride is N3−. (d) The prefix tri is associated with iodide, NI3. 3.91 (a) one sodium ion, Na+, and one chloride ion, Cl− (b) one magnesium ion, Mg2+, and two chloride ions, Cl− (c) two sodium ions, Na+, and one sulfate ion, SO42− (d) one calcium ion, Ca2+, and two nitrate ions, NO3− 3.93 (a) electrolyte; (b) electrolyte; (c) electrolyte; (d) nonelectrolyte 3.95 The ions of silver, zinc, and cadmium can each only have one possible charge. They are assumed to have these charges in the compounds they form. 3.97 (a) NO3−; (b) SO32−; (c) NH4+; (d) CO32−; (e) SO42−; (f) NO2−; (g) ClO4− 3.99 (a) magnesium bromide; (b) hydrogen sulfide; (c) hydrosulfuric acid; (d) cobalt(III) chloride; (e) potassium hydroxide; (f) silver bromide 3.101 (a) PbCl2; (b) Mg3(PO4)2; (c) NI3; (d) Fe2O3; (e) Ca3N2; (f) Ba(OH)2; (g) Cl2O5; (h) NH4Cl 3.103 NaHCO3 3.105 Ca(ClO)2 3.107 H2O(l) 3.109 Cu, AgNO3, Ag, Cu(NO3)2 3.111 (a) NH3; (b) HNO3(aq); (c) HNO2(aq)

A-13

Hydrogen atoms attached to oxygen atoms are responsible for the acidic properties. All contain oxygen. Magnesium oxide, MgO, has metal and nonmetal ions which makes it an ionic compound. It is a solid at room temperature. Oxygen and carbon dioxide (O2 and CO2) are both molecular substances and gases at room temperature.

Chapter 4 – Chemical Composition 4.1 4.3 4.5 4.7 4.9 4.11 4.13 4.15

4.17 4.19 4.21 4.23 4.25 4.27 4.29 4.31

4.33 4.35 4.37 4.39 4.41 4.43 4.45 4.47 4.49 4.51 4.53

4.55 4.57 4.59 4.61 4.63

4.65

(a) mole; (b) Avogadro’s number; (c) empirical formula; (d) solute; (e) molarity; (f) concentrated solution 40.0% calcium 60.0% carbon 0.226 g lithium Chemical formulas must have whole-number subscripts. (a) H2S; (b) N2O3; (c) CaCl2 H2SO4, SCl4, C2H4 CO2 A formula unit describes the simplest unit of an ionic or network solid. Sodium chloride, NaCl; chlorine, Cl2; methane, CH4; silicon dioxide, SiO2. 3.0 × 1023 molecules; 3.0 × 1023 N atoms; 9.0 × 1023 H atoms 3.011 × 1023 formula units 1 × 1023 atoms 6.022 × 1023 calcium ions NaCl, 58.44 g/mol; CH4, 16.04 g/mol; Cl2, 70.90 g/mol; SiO2, 60.09 g/mol (a) 472.1 amu; (b) 172.18 amu; (c) 150.90 amu; (d) 120.07 amu (a) 253.8 g/mol; (b) 158.35 g/mol; (c) 56.10 g/mol To measure out a useful number of atoms by counting would not be possible because atoms are too small for us to see and manipulate. 42.394 amu 35.9 g/mol (a) 0.0999 mol; (b) 0.293 mol; (c) 0.127 mol; (d) 0.0831 mol (a) 0.556 mol; (b) 1.388 × 10-3 mol; (c) 733 mol; (d) 5.72 × 10-8 mol Na (a) 428 g; (b) 177 g; (c) 436 g; (d) 2.20 × 102 g 5.0 × 102 g (a) 1.76 mol; (b) 1.06 × 1024 molecules; (c) 1.06 × 1024 N atoms; (d) 5.28 mol of H atoms H2SO4 has the most atoms per mole because it has more atoms per molecule. Na has the least atoms per mole. 1.7 × 1021 molecules (a) 9.421 × 1023 formula units; (b) 1.581 × 1024 formula units; (c) 8.356 × 1024 formula units; (d) 2.696 × 1024 formula units (a) 8.364 × 1025 atoms; (b) 1.541 × 1024 ions; (c) 5.619 × 1024 atoms; (d) 1.451 × 1024 ions 6.8 g SO2 NH3 (82%) no The molecular formula shows the exact numbers of each atom present in a compound. The empirical formula shows the relative amounts of each atom in a compound expressed as small whole numbers. H2O2 (empirical, HO); N2O4 (empirical, NO2)

A-14

4.67 4.69 4.71 4.73 4.75 4.77 4.79 4.81 4.83 4.85

4.87 4.89

4.91 4.93 4.95 4.97 4.99 4.101 4.103 4.105 4.107 4.109 4.111 4.113 4.115 4.117 4.119 4.121 4.123 4.125 4.127 4.129

Appendix D Answers to Selected Questions and Problems

(a) P2O5; (b) same as molecular; (c) same as molecular; (d) C3H5O2 (a) C3H2Cl; (b) same as molecular; (c) same as molecular both NO2 and N2O4 (a) Fe3O4; (b) C6H5NO2 C5H6O C7H5N3O6 percent composition by mass and the molar mass C3H6O3 C6H12O6 (a) 50.05% S, 49.95% O; (b) 47.27% Cu, 52.73% Cl; (c) 42.07% Na, 18.89% P, 39.04% O; (d) 16.39% Mg, 18.89% N, 64.72% O Chalcocite (Cu2S) and cuprite (CuO) are both relatively high in copper content (79.85% and 79.88%, respectively). A solution is a homogeneous mixture of two or more substances. Some common solutions: clear drinks (coffee and tea), window cleaner, soapy water, tap water, air, brass (a homogeneous mixture of copper and zinc). Water is the solvent because it is present in the largest amount. Calcium chloride, CaCl2, is the solute. Concentrated means high in solute concentration, and dilute means low in solute concentration. Concentration describes the quantity of solute in a given amount of solvent or solution. solution A (a) 2.03 M; (b) 0.540 M; (c) 3.20 M; (d) 3.21 M 0.0186 mol Na2SO4; 0.0372 mol Na+; 0.0186 mol SO42(a) 0.375 moles, 28.0 g; (b) 0.512 mol, 72.8 g (a) 1.00 L; (b) 0.0833 L 40.0 mL 69.2 mL (a) 0.01814 M; (b) 0.2974 M; (c) 0.04020 M 9.76 g 39.95 amu, 6.634 × 10-23 g (a) 0.904 mol, 5.44 × 1023 atoms; (b) 0.741 mol, 4.46 × 1023 atoms; (c) 0.169 mol, 1.02 × 1023 atoms 7.493 × 10-3 mol C6H11OBr 2.458 × 1023 oxygen atoms 1.31 × 1025 molecules CO2 (a) C8H8O3; (b) C8H8O3 Al2O3

Chapter 5 – Chemical Reactions and Equations 5.1

5.3 5.5 5.7

(a) single-displacement reaction; (b) anhydrous; (c) molecular equation; (d) decomposition reaction; (e) balanced equation; (f) reactant; (g) spectator ion; (h) combustion; (i) precipitate The reactants are aluminum and oxygen gas. The product is aluminum oxide. Image A represents the reactants, and image C represents the products. The numbers of hydrogen atoms do not match. One hydrogen molecule should be added to the reactant image.

5.9

The product image should show five XeF2 molecules and one unreacted Xe atom.

5.11

5.13 5.15 5.17 5.19 5.21

5.23 5.25

5.27 5.29 5.31

Three major signs are visible: (1) a brown gas is formed, (2) bubbles, (3) color change. It is not a chemical change because no new substance is formed. New substances are formed, so a chemical reaction has taken place. No new chemical substances have formed, so no chemical reaction has taken place. A chemical equation is a chemist’s way of showing what happens during a chemical reaction. It identifies the formulas for the reactants and products and demonstrates how mass is conserved during the reaction. (a) chemical reaction; (b) physical change; (c) chemical reaction Balancing a chemical equation demonstrates how mass is conserved during a reaction. This makes the equation a quantitative tool for determining the amount of reactant used and product produced. (a) NaH(s) + H2O(l) H2(g) + NaOH(aq) (b) 2Al(s) + 3Cl2(g) 2AlCl3(s) N2(g) + 3Cl2(g) 2NCl3(g) Image B: 2Mg(s) + O2(g) 2MgO(s)

5.33

5.35

5.37

5.39

(a) 2Al(s) + 3Cl2(g) 2AlCl3(s) (b) Pb(NO3)2(aq) + K2CrO4(aq) PbCrO4(s) + 2KNO3(aq) (c) 2Li(s) + 2H2O(l) 2LiOH(aq) + H2(g) (d) 2C6H14(g) + 19O2(g) 12CO2(g) + 14H2O(l) (a) CuCl2(aq) + 2AgNO3(aq) Cu(NO3)2(aq) + 2AgCl(s) (b) S8(s) + 8O2(g) 8SO2(g) (c) C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

Appendix D Answers to Selected Questions and Problems

5.41 5.43

5.45 5.47 5.49 5.51

5.53 5.55 5.57 5.59

5.61 5.63 5.65

5.67 5.69 5.71

5.73

5.75 5.77 5.79

5.81

Cu(NO3)2(aq) + 2Ag(s) Cu(s) + 2AgNO3(aq) decomposition: one reactant, more than one product combination: more than one reactant, one product single-displacement: element and compound as reactants, different element and compound as products double-displacement: two compounds as reactants, two compounds as products (a) combination; (b) single-displacement; (c) decomposition double-displacement (a) combination; (b) single-displacement (a) CaCl2(aq) + Na2SO4(aq) CaSO4(s) + 2NaCl(aq): double-displacement (b) Ba(s) + 2HCl(aq) BaCl2(aq) + H2(g): singledisplacement (c) N2(g) + 3H2(g) 2NH3(g): combination (d) FeO(s) + CO(g) Fe(s) + CO2(g): not classified (e) CaO(s) + H2O(l) Ca(OH)2(aq): combination (f) Na2CrO4(aq) + Pb(NO3)2(aq) PbCrO4(s) + 2NaNO3(aq): double-displacement (g) 2KI(aq) + Cl2(g) 2KCl(aq) + I2(aq): single-displacement (h) 2NaHCO3(s) Na2CO3(s) + CO2(g) + H2O(g): decomposition NiCO3(s) NiO(s) + CO2(g) (a) CaCO3(s) CaO(s) + CO2(g) (b) CuSO4⋅5H2O(s) CuSO4(s) + 5H2O(g) 2Mg(s) + O2(g) 2MgO(s) (a) 3Ca(s) + N2(g) Ca3N2(s) (b) 2K(s) + Br2(l) 2KBr(s) (c) 4Al(s) +3O2(g) 2Al2O3(s) (a) Zn(s) + 2AgNO3(aq) 2Ag(s) + Zn(NO3)2(aq) (b) 2Na(s) + FeCl2(s) 2NaCl(s) + Fe(s) Zn(s) + SnCl2(aq) ZnCl2(aq) + Sn(s) (a) Ca reacts with water: Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) Ca reacts with HCl: Ca(s) + 2HCl(aq) CaCl2(aq) + H2(g) (b) Fe reacts with HCl: Fe(s) + 2HCl(aq) FeCl2(aq) + H2(g) (c) no reaction (a) yes; (b) no; (c) no; (d) yes (a) soluble; (b) soluble; (c) insoluble; (d) insoluble (a) K2CO3(aq) + BaCl2(aq) BaCO3(s) + 2KCl(aq) (b) CaS(s) + Hg(NO3)2(aq) Ca(NO3)2(aq) + HgS(s) (c) Pb(NO3)2(aq) + K2SO4(aq) PbSO4(s) + 2KNO3(aq) (a) CaCO3(s) + H2SO4(aq) CaSO4(s) + CO2(g) + H2O(l) (b) SnCl2(aq) + 2AgNO3(aq) Sn(NO3)2(aq) + 2AgCl(s) lead(II) sulfate CaCl2(aq) + K2CO3(aq) CaCO3(s) + 2KCl(aq) (a) precipitation of HgS(s) (b) formation of the insoluble gas H2S(g) (c) formation of the stable molecular compound H2O(l) and precipitate BaSO4(s) (a) H2S(aq) + Cu(OH)2(s) CuS(s) + 2H2O(l) (b) no reaction (c) KHSO4(aq) + KOH(aq) K2SO4(aq) + H2O(l)

5.83 5.85 5.87 5.89 5.91 5.93 5.95 5.97

5.99 5.101

5.103

5.105 5.107 5.109

5.111

5.113 5.115

5.117 5.119

A-15

O2(g) (a) Cs2O(s); (b) PbO(s) or PbO2(s); (c) Al2O3(s); (d) H2O(g); (e) CO(g) or CO2(g) (a) CO(g) or CO2(g) and H2O(g); (b) CO2(g); (c) Al2O3(s); (d) CO(g) or CO2(g) and H2O(g) An electrolyte produces ions when dissolved in water. A nonelectrolyte does not produce ions in water. (a) and (b) are electrolytes; (c) nonelectrolyte nonelectrolyte (a) no ions; (b) no ions; (c) Na+(aq) and Cl-(aq) Molecular: all substances written as complete chemical formulas or atoms Ionic: all soluble salts, strong acids and bases are written as ions Net ionic: all spectator ions are removed from the ionic equation ions that do not participate in a reaction (a) 2NaCl(aq) + Ag2SO4(s) Na2SO4(aq) + 2AgCl(s) 2Cl-(aq) + Ag2SO4(s) 2AgCl(s) + SO42-(aq) (b) Cu(OH)2(s) + 2HCl(aq) CuCl2(aq) + 2H2O(l) Cu(OH)2(s) + 2H+(aq) Cu2+(aq) + 2H2O(l) (c) BaCl2(aq) + Ag2SO4(s) BaSO4(s) + 2AgCl(s) Ba2+(aq) + 2Cl-(aq) + Ag2SO4(s) BaSO4(s) + 2AgCl(s) (a) calcium carbonate, CaCO3 (b) Na2CO3(aq) + CaBr2(aq) CaCO3(s) + 2NaBr(aq) (c) 2Na+(aq) + CO32-(aq) + Ca2+(aq) + 2Br-(aq) CaCO3(s) + 2Na+(aq) + 2Br-(aq) (d) Na+(aq) and Br-(aq) (e) CO32-(aq) + Ca2+(aq) CaCO3(s) Ag+(aq) + Cl-(aq) AgCl(s) 2Na(s) + 2H2O(l) 2Na+(aq) + 2OH-(aq) + H2(g) (a) Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s) Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s) (b) FeO(aq) + 2HCl(aq) FeCl2(aq) + H2O(l) O2-(aq) + 2H+(aq) H2O(l) (a) Sr(NO3)2(aq) + H2SO4(aq) SrSO4(s) + HNO3(aq) Sr2+(aq) + SO42-(aq) SrSO4(s) (b) no reaction (c) CuSO4(aq) + BaS(aq) CuS(s) + BaSO4(s) + Cu2 (aq) + SO42-(aq) + Ba2+(aq) + S2-(aq) CuS(s) + BaSO4(s) (d) NaHCO3(aq) + CH3CO2H(aq) CO2(g) + H2O(l) + NaCH3CO2(aq) HCO3-(aq) + CH3CO2H(aq) CO2(g) + H2O(l) + CH3CO2-(aq) combination (a) ZnSO4(aq) + Ba(NO3)2(aq) BaSO4(s) + Zn(NO3)2(aq) (b) 3Ca(NO3)2(aq) + 2K3PO4(aq) Ca3(PO4)2(s) + 6KNO3(aq) (c) ZnSO4(aq) + BaCl2(aq) BaSO4(s) + ZnCl2(aq) (d) 2KOH(aq) + MgCl2(aq) 2KCl(aq) + Mg(OH)2(s) (e) CuSO4(aq) + BaS(aq) CuS(s) + BaSO4(s) 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) (b)

Appendix D Answers to Selected Questions and Problems

A-16

5.121 5.123

Aqueous solutions of HCl and NaOH, HNO3 and KOH, HCl and KOH will give the desired net ionic equation. No reaction occurs. K+

5.125

I–

O C

Solution of CO

Solution of KI

5.127 5.129

Any metal higher in activity can be used. (a) Al; (b) Zn; (c) Mg; (d) Sn (a) Potassium chromate precipitates iron(III) chromate leaving Al3+ in solution. (b) Sodium sulfate precipitates barium sulfate leaving Mg2+ in solution. (c) Silver perchlorate precipitates silver chloride leaving NO3- in solution. (d) Barium sulfate is insoluble in water; MgSO4 will dissolve in water.

Chapter 6 – Quantities in Chemical Reactions 6.1 6.3 6.5 6.7 6.9 6.11

(a) stoichiometry; (b) heat; (c) endothermic reaction; (d) specific heat; (e) actual yield (a) C6H12O6 + 6O2 6CO2 + 6H2O; (b) 72 CO2 molecules; (c) 5 C6H12O6 molecules The coefficients give the relative ratios of products used and reactants produced. They can represent the actual numbers of molecules or moles of each substance. (a) 1 Mg2+(aq) and 2 NO3−(aq) ions (b) 1 mole of Mg2+(aq) and 2 moles of NO3−(aq) ions The best representation is (d); however, (c) could also be chosen because it has the reactants and products in the correct proportions. In reaction (a), the reactants are not diatomic. In reaction (b), mass is not conserved (not balanced). 15mol O2 (a) (b) 7.5 mol O2; (c) 2.9 mol C6H6 2 mol C6 H 6 6 mol HNO3 3mol H 2 2 mol Al ; (b) ; (c) 3mol H 2 2 mol Al 2 mol Al

6.13

(a)

6.15

(a) not conserved; (b) conserved; (c) conserved; (d) conserved; (e) moles of molecules may be conserved in some cases. The others are conserved in all cases.

6.17

(a)

6.19 6.21 6.23 6.25 6.27 6.29 6.31 6.33

(b) 3.50 mol C, 3.50 mol CO, 1.50 mol N2, 2.50 mol H2O (c) 21.9 mol C, 21.9 mol CO, 9.38 mol N2, 15.6 mol H2O (a) 249.70 g/mol; (b) 159.62 g/mol; (c) 0.639 g CuSO4 (a) 2.05 g; (b) 5.89 g; (c) 7.94 g (a) 21.0 g CO; (b) 38.0 g I2 (a) 1.23 g Cl2; (b) 0.0460 g H2O; (c) 2.13 g O2; (d) 3.06 g Cl2 1.0 g CaCO3 the fuel (a) 12 sandwiches; (b) bread is limiting; (c) 3 pieces of turkey (a) F2; (b) N2; (c) N2 and F2 (both react completely)

6.35

(a)

3mol N 2 5mol H 2 O 7 mol C 7 mol CO , , , 2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3

(b) oxygen (c) hydrogen 6.37

2C2H2(g) +

7O2(g)

Initially mixed

6 molecules

18 molecules

How much reacts

4 molecules

14 molecules

Composition of mixture

2 molecules

4 molecules

4CO2(g) + 6H2O(g) 0 molecules

0 molecules

8 molecules

12 molecules

Appendix D Answers to Selected Questions and Problems

6.39

6.41 6.43 6.45 6.47

6.49 6.51 6.53 6.55 6.57 6.59 6.61 6.63 6.65 6.67 6.69 6.71 6.73 6.75 6.77 6.79 6.81 6.83 6.85 6.87 6.89 6.91 6.93 6.95 6.97 6.99 6.101 6.103 6.105 6.107 6.109 6.111 6.113

6.115 6.117 6.119 6.121

2C2H10(g) + 13O2(g) Initially mixed

3.10 mol

13.0 mol

How much reacts

2.00 mol

13.0 mol

Composition of mixture

1.10 mol

0.0 mol

8CO2(g)

+

A-17

10H2O(g)

0.00 mol

0.00 mol

8.00 mol

10.0 mol

(a) P4; (b) both are consumed completely; (c) O2 (a) F2; (b) 12.5 g NF3 limiting reactant, NaHCO3; 4.2 g CO2 produced 2C2H6(g) + 7O2(g) Initially mixed

0.260 g

1.00 g

How much reacts

0.260 g

0.968 g

Composition of mixture

0.000 g

0.03 g

4CO2(g) + 6H2O(g) 0.00 g

0.00 g

0.761 g

0.467 g

(a) NaOH(aq) + HNO3(aq) H2O(l) + NaNO3(aq) (b) HNO3 (c) basic actual yield An actual yield greater than the theoretical yield can be caused by contamination. The solid may have not been dry, causing the apparent mass of product to be high. (a) 7.44 g; (b) 8.95 g; (c) 83.1% 90.5% 0.47 mol I2 84% The potential energy is converted to kinetic energy (including heat), some of which is transferred to the ground as heat. The potential energy of the hydrogen and oxygen is used to provide electrical energy which runs electric motors and produces kinetic energy. The reaction is endothermic because the decrease in temperature of the surroundings indicates energy is absorbed by the reaction. As the liquid evaporates, the molecules going into the gas state must absorb some energy from their surrounding. As a result, the surroundings (water and your skin) lose energy and feel colder. The products are lower in potential energy. 2.20 × 103 J 3.47 × 104 cal 0.209 Cal 114 Cal lead 5470 J or 1310 cal −210 J A calorimeter should have good insulation and a way to accurately and precisely measure the temperature. (a) The rock must have lost heat. (b) qrock = −1.3 × 103 J (a) lower; (b) gain; (c) 2.8 kJ 63.5°C Because a reaction is a process, you can only measure the effect it has on the surroundings. The chemical reaction is the system, and the water and calorimeter are the surroundings. (a) exothermic; (b) -525 kJ/mol (a)147 kJ; (b) -24.5 kJ/g (a) qnut = −2.6 × 104 J so 2.6 × 104 J is released; (b) 6.3 × 103 cal or 6.3 Cal; (c) 3.2 Cal/g −256 kJ (a) 3.00 mol CO2, 1.50 mol N2, 0.250 mol O2, 2.50 mol H2O (b) 7.50 mol CO2, 3.75 mol N2, 0.625 mol O2, 6.25 mol H2O (a) 30.6 AgNO3; (b) 25.8 g 9330 g O2 (a) 2K(s) + Cl2(g) 2KCl(s) (b) Cl2 (c) There should be some gray solid and a white solid (KCl) inside the container. (d) 11 g (e) 4.5 g K 1mol O2 ; (b) 7.9 g O2 and 8.9 g H2O; (c) 4.5 g H2O 2 mol H 2 31.8°C 0.450 J/(g °C); chromium 2.80 × 103 kJ/mol, 6.70 × 102 kcal/mol (a)

A-18

6.123 6.125

Appendix D Answers to Selected Questions and Problems

54.8°C 0.450 J/(g °C); chromium

7.41

Their primary differences are size and energy. The 3p orbital is larger and higher in energy.

Chapter 7 – Electron Structure of the Atom 7.1

7.3

(a) electromagnetic radiation; (b) frequency; (c) ionization energy; (d) Hund’s rule; (e) electron configuration; (f) core electron; (g) orbital; (h) continuous spectrum; (i) isoelectronic Infrared, microwave, and radio frequency

7.5

7.43 7.45

(a) 1; (b) 3; (c) 5; (d) 7; (e) 1; (f) 3 Si B P

Lower frequency ( ) with higher frequency ( ) superimposed

7.7 7.9 7.11 7.13 7.15 7.17 7.19 7.21 7.23 7.25

7.27 7.29 7.31 7.33

7.35 7.37

7.39

blue, yellow, orange, red As wavelength increases, the frequency decreases. As wavelengths decrease, frequencies increase. infrared Gamma photons have the highest energy and highest frequency. Radio frequency waves are the longest. 4.00 × 1015 Hz, ultraviolet 4.27 × 10−19 J, visible light (blue) white light no No. If they did, line spectra would not exist. The energy of the electron is quantized—it can only have certain values. To move to a higher orbit, the electron would have to absorb energy. To go to a lower orbit, it would have to release energy (e.g., a photon). A single photon is released. The n = 6 to n = 3 transition gives the highest-energy photon. The n = 5 to n = 3 transition gives the lowest-energy photon and therefore the longest wavelength. The four lines are a result of four different transitions in the hydrogen atom. These transitions are n = 6 to n = 2 violet (highest energy) n = 5 to n = 2 blue n = 4 to n = 2 green n = 3 to n = 2 red (lowest energy) 4.58 × 10-19 J Bohr’s orbits required that the orbits be a fixed distance from the nucleus with the electron following a specific pathway. Modern orbitals describe the region of space surrounding the nucleus where we are most likely to find the electron. B

3p

2p

7.47 7.49 7.51 7.53 7.55 7.57 7.59 7.61

7.63

7.65 7.67 7.69 7.71 7.73 7.75

7.77

7.79

7.81

7.83 7.85

1s

2s

2p

3s

3p

1s

2s

2p

3s

3p

1s

2s

2p

3s

3p

(a) 1s22s22p63s23p2; (b) 1s22s1; (c) 1s22s22p63s2 All orbitals can hold 2 electrons (including a 2p orbital). 8 electrons 3 elements in group VIIA (17) transition metals silicon (a) 1s22s22p63s1 (b) 1s22s22p63s23p64s23d5 (c) 1s22s22p63s23p64s23d104p4 Bromine should be 1s22s22p63s23p64s23d104p5. The 4p orbital given in the problem has one too many electrons and the 10 electrons in the d orbital are in the third principle energy level. (a) chlorine; (b) cobalt; (c) cesium (a) [Ne]3s1; (b) [Ar]4s23d5; (c) [Ar]4s23d104p4 bromine (1s22s22p63s23p64s23d104p5) Valence electrons are the electrons in the highest principle energy level. No, the d-orbital electrons are always one energy level lower than the valence electrons. As you go left to right on the periodic table, one electron is added each time the group number increases. The group number is the number of electrons in the s and p orbitals of the highest energy level. Electrons in the d orbitals are not included since they are always one energy level below the highest energy level. (a) third valence level; 3 valence electrons (b) third valence level; 6 valence electrons (c) fourth valence level; 5 valence electrons Cations always have fewer electrons than the element from which they are formed. They are similar in that they possess the same core of electrons. (a) 1s22s22p6; [Ne]; F − (b) 1s22s22p6; [Ne]; N3− (c) 1s22s22p63s23p63d10; [Ar]3d10; Zn2+ O2−, N3−, Na+ Valence electrons farther from the attraction of the nucleus are easier to remove. This means that potassium (4s valence electrons) has a lower ionization energy

Appendix D Answers to Selected Questions and Problems

7.87

7.89 7.91

7.93

7.95 7.97 7.99

7.101 7.103

7.105 7.107 7.109 7.111 7.113 7.115 7.117

7.119 7.121

than sodium (3s valence electrons). Since less energy is expended to remove potassium’s electron, more energy is left to drive the reaction. The first ionization energy of calcium is higher than that of potassium. Since less energy is expended to remove potassium’s electron, more energy is left to drive the reaction. In addition, calcium must lose a second electron before it becomes stable. Even more energy is expended to remove this electron. P