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Appendix C

Answers to Practice Problems



Answers to Selected Questions and Problems

Chapter 1 ­ Matter and Energy

1.1 (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. H2 1.47 1.49 1.51 1.53


Solid state

1.3 1.5 1.7 1.9 1.11 1.13 1.15 1.17

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.

Liquid state


1.21 1.23

1.25 1.27 1.29

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 g; (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 Celsius Kelvin Fahrenheit Freezing 0°C 213.15 K 32°F Boiling 100°C 373.15 K 212°F Difference 100°C 100 K 180°F

1.31 1.33 1.35 1.37 1.39 1.41 1.43

1.85 1.87


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.




Appendix D

Answers to Selected Questions and Problems


Symbolic: Cl2(g) gram is shown.

Cl2(l). A molecular-level dia-






1.93 1.95

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 (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.

2.3 2.5 2.7

1.97 1.99

1.101 1.103 1.105



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

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

Appendix D Answers to Selected Questions and Problems



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

1 1H 2 1H 3 1H

2.63 2.65 2.67 2.69

Atomic number Neutron number Mass number 2.31 2.33 (a) 15 O 8 (b) 109 Ag 47 (c) 35 Cl 17 2.35 2.37 (a) 56 Fe 26 (b) 39 K 19 (c) Copper-65 or 65 Cu 29 2.39 2.41 (a) 23Na 11 (b) 56 Mn 25 (c) 18O 8 (d) 19 F 9 2.43 2.45 2.47 2.49 (a) Zn2+ (b) F(c) H 2.51 (a) (b) (c)

37 17 Cl 25 2+ 12 Mg 13 37N +

1 0 1

1 1 2

1 2 3 2.71

(a) Z = 18, N = 18, A = 36; (b) Z = 18, N = 20, A = 38; (c) Z = 18, N = 22, A = 40 Protons 8 47 17 Neutrons 7 62 18 Electrons 8 47 17 2.75 2.77 2.79 2.73

(a) 3 H; (b) 9 Be; (c) 31 P 1 4 15 Protons 26 19 29 Neutrons 30 20 36

(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) (a) (b) (d)

58 + 28 Ni 60 + 28 Ni

Protons 28 28 28 28 28

Neutrons 30 32 33 34 36

Electrons 27 27 27 27 27

(c) 61 Ni+ 28

62 + 28 Ni 64 + 28 Ni

seven protons and six neutrons Protons 11 25 8 9 Neutrons 12 31 10 10 Electrons 11 25 8 9 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 Electrons 18 10 10 18 2.113 2.115 2.117 2.119 2.109 2.111 20 13 6 20


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 Protons 30 9 1 Protons 17 12 7 20 Electrons 28 10 0 Neutrons

(d) 40 Ca2+ 20 2.53 2.55 2.57

2.59 2.61

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.


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


Appendix D Answers to Selected Questions and Problems

2.123 2.125 2.127

2.129 2.131 2.133

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 Ca2+ Cl­ CaCl2 calcium chloride CaO calcium oxide


Fe2+ FeCl2 iron(II) chloride FeO iron(II) oxide Fe(NO3)2 iron(II) nitrate FeSO3 iron(II) sulfite Fe(OH)2 iron(II) hydroxide Fe(ClO3)2 iron(II) chlorate Al3+ AlCl3 aluminum chloride Al2O3 aluminum oxide Al(NO3)3 aluminum nitrate Al2(SO3)3 aluminum sulfite Al(OH)3 aluminum hydroxide Al(ClO3)3 aluminum chlorate iron(III) FeI3 Fe2O3 Fe2(SO4)3 Fe(NO2)3 Fe(CH3CO2)3 Fe(ClO)3

K+ KCl potassium chloride K2O potassium oxide KNO3 potassium nitrate K2SO3 potassium sulfite KOH potassium hydroxide KClO3 potassium chlorate NH4+ NH4Cl ammonium chloride (NH4)2O ammonium oxide NH4NO3 ammonium nitrate (NH4)2SO3 ammonium sulfite NH4OH ammonium hydroxide NH4ClO3 ammonium chlorate strontium SrI3 SrO SrSO4 Sr(NO2)2 Sr(CH3CO2)2 Sr(ClO)2

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 3.27 (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




Ca(NO3)2 calcium nitrate CaSO3 calcium sulfite Ca(OH)2 calcium hydroxide




ClO3­ Ca(ClO3)2 calcium chlorate Mn2+ Cl­

MnCl2 manganese(II) chloride MnO manganese(II) oxide





Mn(NO3)2 manganese(II) nitrate MnSO3 manganese(II) sulfite Mn(OH)2 manganese(II) hydroxide Mn(ClO3)2 manganese(II) chlorate potassium

SO32 3.29 3.31 3.33 3.35 3.37 (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 iodide oxide sulfate nitrite acetate hypochlorite KI K2O K2SO4 KNO2 KCH3CO2 KClO


3.41 3.43

Appendix D Answers to Selected Questions and Problems


aluminum iodide oxide sulfate nitrite acetate AlI3 Al2O3 Al2(SO4)3 Al(NO2)3 Al(CH3CO2)3

cobalt(II) CoI2 CoO CoSO4 Co(NO2)2 Co(CH3CO2)2

lead(IV) PbI4 PbO2 Pb(SO4)2 Pb(NO2)4 Pb(CH3CO2)4

3.113 3.115

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.

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)

Chapter 4 ­ Chemical Composition

4.1 4.3 4.5 4.7 4.9 4.11 4.13 4.15 (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)

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



Appendix D Answers to Selected Questions and Problems

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

(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

5.9 5.11

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

5.13 5.15 5.17 5.19 5.21

5.23 5.25

5.27 5.29 5.31 5.33

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)

Chapter 5 ­ Chemical Reactions and Equations

5.1 (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.


5.3 5.5 5.7



(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


5.75 5.77 5.79


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) CuSO45H2O(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


5.105 5.107 5.109


5.113 5.115

5.117 5.119

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

5.121 5.123 5.125

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

K+ I­ O C

Solution of KI

Solution of CO

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 6.13 6.15 6.17 (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 (a) 6 mol HNO3 3mol H 2 2 mol Al ; (b) ; (c) 3mol H 2 2 mol Al 2 mol Al

(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. (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

6.19 6.21 6.23 6.25 6.27 6.29 6.31 6.33 6.35

(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) (a)

(b) oxygen (c) hydrogen 6.37 Initially mixed How much reacts Composition of mixture 2C2H2(g) + 6 molecules 4 molecules 2 molecules 7O2(g) 18 molecules 14 molecules 4 molecules 8 molecules 12 molecules 4CO2(g) + 6H2O(g) 0 molecules 0 molecules

Appendix D Answers to Selected Questions and Problems


6.39 Initially mixed How much reacts Composition of mixture 6.41 6.43 6.45 6.47

2C2H10(g) + 13O2(g) 3.10 mol 2.00 mol 1.10 mol 13.0 mol 13.0 mol 0.0 mol

8CO2(g) 0.00 mol


10H2O(g) 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 How much reacts Composition of mixture 0.260 g 0.260 g 0.000 g 1.00 g 0.968 g 0.03 g 0.761 g 0.467 g 4CO2(g) + 6H2O(g) 0.00 g 0.00 g

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

(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)

6.115 6.117 6.119 6.121


Appendix D Answers to Selected Questions and Problems

6.123 6.125

54.8°C 0.450 J/(g °C); chromium


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 (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 7.43 7.45

7.3 7.5

2p 1s 1s 1s

3p 2s 2s 2s 2p 2p 2p 3s 3s 3s 3p 3p 3p

(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


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

7.47 7.49 7.51 7.53 7.55 7.57 7.59 7.61


7.65 7.67 7.69 7.71 7.73 7.75




7.83 7.85

(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



7.89 7.91


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<S<O Sodium's valence electron is in a higher principle energy level. This means it is not as close to the nucleus and is therefore easier to remove. The fluorine has more protons in its nucleus than oxygen. As a result, the electrons of fluorine are more strongly attracted to the nucleus; the ionization energy is higher. First ionization, IE1, Mg(g) Mg+(g) + 1e- + Second ionization, IE2, Mg (g) Mg2+(g) + 1e- Magnesium because the third electron is removed from the core electrons. (a) Fluorine has a very high ionization energy. (b) Elements do not lose core electrons when forming ions (under normal conditions). (c) high ionization energy because Ne has a core configuration O<S<P The valence electrons of chorine and sulfur are both in principal energy level 3, but chlorine has more protons in its nucleus to attract the electrons. (a) Mg; (b) P3- Both ions are isoelectronic with argon, but potassium ion is larger because it has fewer protons in its nucleus. For hydrogen the energies of the sublevels are all the same. (a) [Xe]6s24f 145d106p3; (b) [Rn]; (c) [Rn]7s2 12 filled p orbitals This means that there are 10 filled d orbitals. Both are related to the attraction of electrons for the nucleus. As attraction gets higher, electrons are more tightly held. This results in greater ionization energy and smaller radius. The radius of the sodium ion is smaller and the chloride ion greater than for their respective elements. Neon's red color indicates lower energy, lower frequency, and longer wavelength. Argon's blue color indicates higher energy, higher frequency, and shorter wavelength.

8.23 8.25 8.27

(a) H-F ; (b) I-Cl ; (c) B-H







(a) H-H < C-H < O-H < F-H (b) O-Cl < C-Cl < F-Cl < HCl The Lewis symbol can be used to determine the number of valence electrons and the number of electrons gained or lost to reach an octet configuration. (a) C ; (b) I ; (c) Se ; (d) Sr ; (e) Cs ; (f) Ar


8.31 8.33

(a) Cl ­ ; (b) Sc3+; (c) S (a) Li+ (b) Cl


; (d) Ba2+; (e) B3+



Ba2+ S



(c) Ba2+ 8.35



8.41 8.43 8.45

The second electron removed from potassium, K, would be taken from the core. These electrons are very strongly held by the atom and are not easily removed. When fluorine gains an electron and sodium loses one electron, both have satisfied the octet rule. The ionic bond formed during the process is very stable and the compound is electrically neutral. A crystal lattice is the repeating pattern observed in all crystal structures. An ionic crystal is the structure that forms when ions form crystals by minimizing the repulsive energies of like-charged ions and maximizing the contact of oppositely charged ions. Six chloride ions surround each sodium ion. The ratio of the cations and anions is different, so different ionic crystals are formed. O O F F (a) 8 valence electrons around each atom in O2 and F2; (b) The bond in O2 is a double bond. The bond in F2 is a single bond. Each hydrogen needs only one electron to satisfy its valence shell. It does not have another electron to form another bond. (a) 1; (b) 3; (c) 1; (d) none (a) boron (b) oxygen (c) nitrogen (a) H C N (b) H C H H C N C H C H H


8.49 8.51

Chapter 8 ­ Chemical Bonding

8.1 (a) single bond; (b) alkane; (c) covalent bonding; (d) ionic crystal; (e) octet rule; (f) polar covalent bond; (g) alkyne; (h) electronegativity; (i) triple bond; (j) crystal lattice; (k) Lewis formula; (l) chemical bond A chemical bond is an attractive force between atoms or ions in a substance. Metals form ionic bonds with nonmetals. HF, NCl3, CF4 (a) ionic; (b) covalent; (c) covalent; (d) ionic A and D CH4, SF4, HCl (a) high; (b) high; (c) low; (d) high; (3) low The electronegativity value decreases. bonds between two identical atoms (a) N = Br < Cl < O < F (highest electronegativity); (b) H < C < N < O < F 8.53

8.3 8.5 8.7 8.9 8.11 8.13 8.15 8.17 8.19 8.21

(c) H C (d) H H (e) C





Appendix D Answers to Selected Questions and Problems


(a) O




8.75 8.77 8.79


(b) O


(a) alcohol; (b) alkane; (c) ether; (d) alkene; (e) ketone (a) alcohol; (b) alkyne; (c) aldehyde; (d) ether (a) alkane; (b) amine; (c) ester H H H O H H H The geometric arrangement of electron pairs and atoms is predicted by finding the geometry that gives the greatest distance between the electron groups and atoms (Table 8.5). The shape is determined by the arrangement of the bonded atoms (Table 8.6). Shape is determined by first knowing the number of bonded atoms and unshared electron pairs (Table 8.5). Once that is determined, the shape is given by the arrangement of atoms (Table 8.6). (a) H C C C C H



(c) O (d) O (e) 8.57 8.59


+ ­





8.61 8.63

(a) different; (b) same; (c) same; (d) same; (e) same Resonance descriptions are required when a molecule or ion can be represented by two or more reasonable Lewis structures that differ only in the positions of the bonding and lone pair electrons. The positions of the atomic nuclei do not change. Resonance is exhibited in (c) and (d). (a) resonance structures for NO2- N O O (b) O O S


(b) (c) (d) 8.89 (a) linear; (b) tetrahedral; (c) tetrahedral; (d) trigonal planar; (e) tetrahedral; (f) tetrahedral; (g) trigonal planar; (h) tetrahedral (a) linear, 180°; (b) trigonal pyramidal, 109.5°; (c) bent, 109.5° (tetrahedral); (d) bent, 120° (trigonal planar); (e) bent, 109.5° (tetrahedral); (f) tetrahedral, 109.5°; (g) trigonal planar, 120°; (h) bent, 109.5° (tetrahedral) (a) 109.5°; (b) 109.5°; (c) not defined; (d) 180°; (e) 120°; (f) 120°; (g) 109.5° (a) BF3; (b) NH3; (c) SCl2 (a) BCl3; (b) H2O; (c) BeCl2; (d) NH4+ ClO3- image B Each nitrogen has one unshared electron pair. The bond angles on the carbon are 109.5°. The bond angles on the nitrogen are 120°. A bond is polar if there is a difference in electronegativity between the two atoms that are bonded. A molecule is polar if the polarity of the bonds to a central atom do not cancel out. The chlorines are all attracting the electrons of the carbon with equal force. Because the forces are in opposite direction and the molecule is geometrically symmetric, the polarity of the bonds cancels in the molecule. (a) polar; (b) polar; (c) polar; (d) polar (a) SO2 has an unshared pair of electrons on the central atom and is not symmetrical, so the bond polarities do not cancel. As a result SO2 is polar. CO2 is linear and contains polar bonds but is symmetrical. The bond polarities in CO2 cancel, so it is a nonpolar molecule. (b) SO2 has an unshared pair of electrons on the central atom and is not symmetrical, so the bond polarities do not cancel. As a result SO2 is polar. SO3 is trigonal planar and contains polar bonds but is symmetrical. The bond polaritites in








(c) O C O (d) O O

8.93 O C O O O






8.95 8.97 8.99 8.101 8.103 8.105 8.107




8.65 8.67 8.69 8.71 8.73

Hydrogen's valence shell can only hold two electrons. (a) octet obeyed; (b) expanded octet; (c) octet obeyed; (d) octet obeyed (a) S, expanded octet; (b) B, incomplete octet; (c) Xe, expanded octet; (d) Cl, odd electron species alkanes with only single bonds, alkenes with double bonds, and alkynes with triple bonds H H H C C C C C C H H


8.111 8.113


Appendix D

Answers to Selected Questions and Problems


8.115 8.117 8.119 8.121 8.123 8.125 8.127

SO3 cancel, so it is a nonpolar molecule. (c) SeCl2 has two unshared pairs of electrons on the central atom and is not symmetrical, so the bond polarities do not cancel. As a result SeCl2 is polar. BeCl2 is linear and contains polar bonds but is symmetrical. The bond polarities in BeCl2 cancel, so it is a nonpolar molecule. (d) Both are tetrahedral, but in the case of CH3I, the bond polarities do not cancel and the molecule is polar. The bonds in CH4 are all equivalent, so the bond polarities cancel. image B The molecule CCl2F2 should be soluble in water. image B (a) Br (b) Pb (c) S (d) Ca (e) Be (f) Xe F > Cl > Br > I (a) covalent; (b) ionic; (c) covalent; (d) covalent; (e) ionic (a) H N N H (b) S C F S


(a) Both structures are tetrahedral. The CCl4 structure is nonpolar due to symmetry. However the different types of atoms (with different electronegativities) bonded to the carbon in CH2Cl2 causes this molecule to be polar. (b) Both structures are tetrahedral and polar but the CH3F molecule is more polar than the CH3Br molecule because the C-F bond is more polar than the C-Br bond. (c) Both structures are trigonal pyramidal and polar but the NF3 molecule is more polar than the NH3 molecule because the N-F bonds are more polar than the N-H bonds. (d) Both structures are bent and polar but the H2O molecule is more polar than the OF2 molecule because the O-H bonds are more polar than the O-F bonds.

Chapter 9 ­ The Gaseous State

9.1 (a) effusion; (b) Boyle's law; (c) combined gas law; (d) ideal gas; (e) pressure; (f) Dalton's law of partial pressures; (g) molar volume; (h) ideal gas constant, R Y axis units and label are missing, x axis units are missing, the points should not be connected by lines (a straight best-fit line that comes closest to all points should be used), gridlines should be uniformly labeled (for example, 10, 20, 30, 40, 50, etc.), the x axis maximum should be 60 instead of 120, a graph title is missing. There is a trend, but it is not linear.

Boiling Point vs. Number of Carbon Atoms 175 125 Boiling point (°C)


(c) F As F (d) O C O (e) C O 8.129 (a) O Cl O O

­ ­




(b) N O O


75 25 ­25 ­75 ­125 ­175 0 1 2 3 4 5 6 7 8 9 10 11 Number of carbon atoms

(c) N C O








(d) H C O O


(e) F 8.131



9.7 F

Pressure and volume are inversely related.

Inches Hg vs. Inches Gas 100 90 80 70 60 50 40 30 20 10 0 0 10 20 30 40 Inches gas 50 60


(a) tetrahedral; (b) trigonal planar; (c) bent; (d) trigonal pyramidal; (e) tetrahedral; (f) bent; (g) tetrahedral; (h) trigonal planar; (i) bent; (j) bent (a) BeCl2 is nonpolar because the molecule is symmetrical. OCl2 is polar. The molecule has polar bonds and is bent (not symmetrical). (b) PH3 is polar. The bonds are polar and the molecule is not symmetrical. BH3 is nonpolar because the molecule is symmetrical. (c) BCl3 is nonpolar because the molecule is symmetrical. AsCl3 is polar because the bonds are polar and the molecule is not symmetrical. (d) SiH4 is nonpolar because the molecule is symmetrical. NH3 is polar because the bonds are polar and the molecule is not symmetrical.

Inches Hg


Appendix D Answers to Selected Questions and Problems

9.9 9.11 9.13 9.15 9.17

9.19 9.21

The vapor pressure is 21 torr and 50 torr at 22°C and 38°C, respectively. (a) x = 0.4; (b) x = 0.206; (c) x = ±4; (d) x = 2 (a) 100°C; (b) 26.7°C; (c); 0°C; -40.0°C (a) 24.4; (b) 0.116; (c) 17.2; (d) 4.70 When compared to other states of matter, gases have low densities and are very compressible. They also take the shape of any container they are put in. The density of warm air is lower than the density of cold air.


9.49 9.51 9.53 9.55

9.23 9.25

Gas pressure is the amount of force exerted by the gas divided by the area over which the force is exerted. Absolute pressure is measured with a device called a barometer. A tire gauge measures the pressure above atmospheric pressure.

9.57 9.59 9.61 9.63 9.65 9.67


9.69 9.71 9.73


9.31 9.33 9.35

(a) 0.980 atm; (b) 935 torr; (c) 0.119 atm; (d) 643 Pa; (e) 1.01 × 103 mm Hg; (f) 563 torr; (g) 1.06 × 105 Pa; (h) 293 mm Hg 1.024 × 105 Pa As pressure increases, the volume decreases at constant temperature. The particles collide more frequently with the container walls causing an increase in pressure. Since the number of moles has not changed, we expect to find twice as many molecules in the same space. The velocity of the molecules should not change.

9.75 9.77 9.79 9.81 9.83 9.85 9.87 9.89 9.91 9.93 9.95

9.97 9.99 9.101 9.103 9.37 9.39 9.41 9.43 9.45 (a) 3.60 L; (b) 26.7 mL; (c) 0.392 mL (a) 161 torr; (b) 0.100 torr; (c) 17.9 atm 182 atm 0.577 L Charles's law states that if the pressure is constant, the volume and temperature (in kelvin) are directly proportional to each other (volume increases when temperature increases; volume decreases when temperature decreases).

The velocity of the gas molecules increases as temperature increases. The volume of the container will increase as the temperature increases. (a) 5.41 L; (b) 671 mL; (c) 75.5 L (a) 273°C; (b) -265°C; (c) -18°C 0.169 mL Nothing happens to the particles if the temperature, pressure, and volume are constant. If the tank is sealed so that no gas molecules are allowed to escape, then nothing happens to the pressure. If the tank is opened, then the molecules of gas in the tank will leave the tank. Particles leave the tank to maintain an equilibrium pressure with the air outside the tank. (a) 418 torr; (b) 673 torr; (c) 4.44 atm (a) -91.3°C; (b) 750°C; (c) 437°C 7.71 atm (a) 1.2 L; (b) -195.3°C; (c) 9.10 × 102 torr 1.61 L Guy-Lussac's law states that volumes of gas react in simple whole-number ratios when the volumes of reactants and products are measured at the same temperature and pressure. 22.414 L/mol (a) 0.385 mol, 6.17 g; (b) 1.562 × 10-2 mol, 2.050 g; (c) 2.15 mol, 60.1 g The number of particles is the same. The balloon containing argon has the greater mass, and therefore greater density since the volumes are the same. (a) 0.34 mol, 7.6 L; (b) 1.5 mol, 34 L; (c) 2.70 mol, 60.5 L 2380 balloons 11.0 L Any gas whose behavior is described by the five postulates of kinetic molecular theory is an ideal gas. (a) 0.34 mol, 0.69 L; (b) 1.5 mol, 3.1 L; (c) 2.70 mol, 5.51 L (a) 0.245 mol, 3.92 g; (b) 4.33 × 10-3 mol, 8.74 × 10-3 g; (c) 0.288 mol, 8.06 g The balloon filled with CO2 sinks because the density of CO2 is greater than the density of air. (a) 0.7600 g/L; (b) 1.250 g/L; (c) 1.964 g/L (a) 0.673 g/L; (b) 1.11 g/L; (c) 1.74 g/L There are 0.208 moles of each gas. (a) 0.42 g; (b) 3.3 g; (c) 13 g Dalton's law of partial pressures states that the total pressure in a container is the sum of the pressures exerted by each of the individual gases in a mixture of gases. 708 torr (a) 2.78 mol; (b) 2.08 mol; (c) 2.15 atm; (d) 1.60 atm (a) 40.01 g/mol; (b) 30.03 g/mol; (c) 45.99 g/mol; (d) 20.3 g/mol; (d) 16.0 g/mol The five postulates of kinetic molecular theory are 1) Gases are composed of small particles widely separated. 2) Particles behave independently of each other. 3) Gas particles move in straight lines. 4) Gas pressure results from force exerted by the molecules in the container. The force is the sum of the forces exerted by each molecule as it bounces off the walls of the container. 5) The average kinetic energy of the gas particles depends only on the absolute temperature.

Appendix D Answers to Selected Questions and Problems




9.109 9.111 9.113 9.115 9.117 9.119 9.121 9.123 9.125 9.127

As the temperature is increased, the kinetic energy of the particles increases. Recall that kinetic energy is the energy of motion. Increasing temperature increases molecular velocity. The particles strike the walls of the container with greater force and, therefore, the pressure is higher. Pressure depends on the force of collisions and the frequency of collisions in a given area. The density of gas particles is inversely proportional to the volume. If the volume decreases, the density increases. Since the density of gas particles increases with decreasing volume, the frequency of collisions in a given area has increased. This means that the pressure has increased. (lowest velocity) CO2 < CH4 < He < H2 (highest velocity) Molecules with lower masses have higher average velocities and as a result diffuse faster. helium 24 L H2 48.0 L CO2 and 76.0 L O2 1.68 × 103 g The balloon expands or contracts so that the external pressure and the internal pressure are the same. The density of a gas decreases if it is heated and the volume is allowed to expand. The air in the hot-air balloon is less dense than the surrounding air, so the balloon floats. 1100 K What will happen if . . . Temperature increases and pressure remains constant? Macroscopic View Balloon gets bigger Microscopic View Molecules move faster Molecules collide harder Collision frequency increases Molecules are further apart Molecules move slower Less energetic collisions Collision frequency decreases Molecules are closer together Molecules move slower Less energetic collisions Lower collision frequency

Temperature decreases and pressure remains constant?

Balloon gets smaller

10,000 feet (temperature decreases and pressure decreases) 9.129 9.131 9.133 9.135 9.137 9.139 9.141 9.143

Balloon probably gets a little bigger

1.75 L 1.21 L 27.7 L (a) 5.00 L; (b) 3.33 L; (c) 1.67 L 228 K (a) If the pressure and temperature are constant, increasing the moles of gas increases the volume occupied by the gas. (b) pressure and temperature 131 g/mol 0.718 L

Chapter 10 ­ The Liquid and Solid States

10.1 10.3 Particle spacing Intermolecular attraction Kinetic energy 10.5 10.7 10.9 (a) vapor pressure; (b) melting point; (c) equilibrium; (d) evaporation; (e) induced dipole; (f) alloy; (g) boiling point; (h) sublimation; (i) molecular solid; (j) intermolecular force; (k) London dispersion force; (l) melting point; (m) crystal Liquid Dense or closely spaced Moderate Low Gas Molecules far apart Weak High


A liquid is fluid because the molecules can move past each other. Attractive forces hold the particles in solids together in a rigid, three-dimensional array. The substance is in the solid state. If the substance were a liquid or gas, it would not maintain its shape. In addition, if the substance was not a solid, the images on the surface would not maintain their shape. The image is a molecular-level representation of the liquid state. The liquid state and solid state both have dense particle spacing. Basically, the molecules are always touching each other. However, in the image, the molecules do not show any clear pattern. This is a characteristic of the liquid state. The solid, liquid, and gas states are represented in circles 1, 3, and 2, respectively. The phase changes are deposition, vaporization, and melting for phase changes A, B, and C, respectively.


Appendix D Answers to Selected Questions and Problems



The vapor pressure of a liquid increases as the temperature is increased. A liquid boils when its vapor pressure is equal to the pressure above the liquid. The temperature at which this occurs is the boiling point. The normal boiling point is the temperature at which the vapor pressure of the liquid is equal to 1 atm. In order to convert a substance from the liquid to the gas state, the energy of the molecules must increase enough to allow them to overcome their attractions for each other and move about independently. Because the molecules must absorb energy to go into the gas state, the process is endothermic (i.e., absorbs energy from the surroundings).



10.61 10.63


The vapor pressures of pure liquids depend on the strength of the intermolecular forces between the particles of the liquids. When a molecule is polar, it means that some portion of the molecule appears permanently slightly negative to other molecules and another portion is slightly positive. This comes from the unequal sharing of electrons (Chapter 8). The oppositely signed charges on different molecules attract each other (dipole-dipole interaction), so the molecules stick together. The instantaneous and induced dipoles in nonpolar molecules are not permanent. As molar mass increases the importance of London dispersion forces increases. Dispersion Force (a) C6H6 (b) NH3 (c) CS2 (d) CHCl3 x x x x Dispersion Force (a) Kr (b) CO (c) CH4 (d) NH3 x x x x x x x x DipoleDipole Hydrogen Bonding x x DipoleDipole Hydrogen Bonding

10.19 10.21




10.29 10.31 10.33 10.35 10.37 10.39 10.41 10.43 10.45 10.47 10.49

10.51 10.53 10.55

solid state to the gas state, sublimation; endothermic All molecules and atoms have some degree of attraction for each other. As a gas is cooled, the molecules are losing kinetic energy. Eventually the attractive forces are greater than the kinetic energy of the molecules and the molecules begin to coalesce (stick together) to form the liquid state. Evaporation of water is endothermic. The energy water needs to evaporate comes from the surrounding air, cooling it. The cooled air can be used to cool a house (if it is circulated by a fan). Water boils when it has been heated to the point that its vapor pressure equals the atmospheric pressure. At high elevation, the atmospheric pressure can be significantly lower than 1 atm (about 0.6 to 0.7 atm at 12,000 ft). Since the water boils at a lower temperature, the pasta cooks slower. As a liquid is heated the molecules of the liquid acquire more energy. The percent of molecules on the liquid surface with enough energy to go into the gas phase increases. This causes an increase in the vapor pressure. Gallium is a liquid at 100°C. At 15°C gallium is a solid. BC 237 kJ 41.1 kJ 1.69 × 103 kJ 124 kJ An attractive force between individual particles (molecules or atoms) of a substance is called an intermolecular force. gas state Ar NO, NF3, and CH3Cl A bond is an attractive force between two atoms in a molecule. A hydrogen bond is not a true bond because it usually takes place between two atoms on different molecules. In comparison to covalent bonds, the strength of attraction is much weaker and the electrons are not shared. Only B can hydrogen-bond. (a) and (b) (b)


10.67 10.69 10.71

10.73 10.75 10.77 10.79


10.83 10.85

(lowest boiling point) CH3OH < CH3CH2OH < CH3CH2CH2OH (highest boiling point) NO2 Hydrogen, H2, is larger than (but not more massive than) helium. The size of the hydrogen molecule allows the electrons to be moved around more easily and therefore the dispersion force is greater. More energy is required to vaporize liquid hydrogen than helium. HCl is polar, while F2 is nonpolar. (a) CH3OH; (b) CH3OH; (c) NH3 (lowest) He < N2 < H2O < I2 (highest) Both molecules are polar, but water has very strong hydrogen bonding while acetone has dipole-dipole interactions. As a result, water has a higher boiling point. On a molecular level, the particles of a substance in the liquid state are higher in energy than for the same substance in the solid state and the molecules have freedom to move apart from each other. In addition, the liquid state does not have a long range structure like the crystal lattice of many solids. On a macroscopic level, liquids flow and can take the shapes of their containers. Liquids tend to be less dense than solids. Viscosity is the resistance of a liquid to flow. The more viscous a liquid, the slower it will flow when it is poured. All molecules of a liquid experience attractive forces to other molecules in the liquid. Molecules in the middle experience forces that are equal in all directions. On the surface, the molecules only experience attractive forces

Appendix D Answers to Selected Questions and Problems




10.91 10.93 10.95 10.97 10.99 10.101 10.103

10.105 10.107

10.109 10.111 10.113 10.115 10.117 10.119 10.121 10.123 10.125 10.127 10.129

to the sides and "downward" (toward the center of the liquid). Surface tension causes liquids to bead up on solid surfaces and is also why liquids tend to form spheres. In addition, surface tension is related to capillary action and the formation of a meniscus (curved surface) of a liquid in a container. There is an imbalance of forces on the molecules at the surface of a liquid. As the surface area of the liquid is minimized, the imbalance of energy is also minimized. Substances assume the shape with the least surface area for the given volume. As a result, liquids have a tendency to form spherical shapes to minimize this energy difference. When ice forms on bodies of water such as lakes and rivers, it stays at the surface and provides an insulating layer. Energy loss through evaporation is minimized and less water freezes than would if the ice did not form at the top. Ice would continually form and sink. Many bodies of water would freeze completely making life in them impossible during the coldest months. Crystals are defined as solids that, on a molecular or atomic level, have a regular repeating pattern of atoms. See Figures 10.31 and 10.37. Crystal structures help us understand the macroscopic properties of the solid state (i.e., conductivity, hardness, brittleness). Close packing is a common structure of metals. Amorphous solids have a tendency to flow because the molecules are not locked into a crystal lattice. Solids are classified by the strength of attraction between the subunits of the solid. This allows us to broadly classify the characteristics of the solid substances. ionic solid There are two reasons. First, the attractive forces between metal atoms are relatively weak. This allows them to be separated from each other. Secondly, the electrons in metal atoms are delocalized among all their neighboring atoms. Atoms can be moved with relatively little effect on the charge distribution. In ionic solids, cations and anions cannot be separated easily, so particles do not move apart easily. This causes ionic solids to be brittle. Both ionic and metallic solids form crystal lattices. Ionic crystals resemble metal crystals with smaller ions fitting into holes between larger ions. Molecular solids are held together either by dipole-dipole interactions, hydrogen bonding, or London dispersion forces. Ionic solids are held together by interionic attractions (i.e., forces of attraction between oppositely charged ions). Because ionic interactions are much stronger than intermolecular forces, the ionic solids tend to melt at much higher temperatures. (a) covalent bonds; (b) London dispersion forces; (c) London dispersion forces, dipole-dipole interactions, and hydrogen bonding; (d) covalent bonds (a) ionic; (b) molecular; (c) metallic; (d) network; (e) molecular (b) and (e) are molecular solids molecular solid SiF4 Below its freezing point, ice is lost by sublimation. CH4 (-162°C), C2H6 (-88°C), C3H8 (-42°C), C4H10 (0°C); London dispersion forces increase with molecular size. A glass plate (SiO2) is very polar. Polar substances are attracted to polar substances. Water is attracted to glass, so it spreads out, but mercury is repelled by the glass and tends to bead up. (c) is an ionic solid (a) Cl2; (b) CH3SH; (c) NF3 Substance (solid type) NaCl (ionic) SiC (network) SiCl4 (covalent) Fe (metallic) Electrical Conductivity conductive in molten state not conductive not conductive Hardness hard hardest softest Melting Point high highest lowest high Vapor Pressure low lowest highest low

always conductive soft*

*Metals are soft relative to network or ionic solids. 10.131 melting; freezing; sublimation; deposition; vaporization; condensation 10.133


Appendix D Answers to Selected Questions and Problems

Chapter 11 ­ Solutions

11.1 (a) saturated solution; (b) aqueous solution; (c) molarity; (d) solvent; (e) entropy; (f) solubility; (g) parts per million; (h) miscible; (i) osmosis; (j) colligative property; (k) percent by volume; (l) mass/volume percent Solute (a) sodium chloride (and other salts) (b) carbon (c) oxygen (O2) 11.5 Solvent water iron water



11.41 11.43

11.7 11.9 11.11 11.13

Strong electrolytes are substances that fully ionize when placed in solution. Strong acids, strong bases, and soluble ionic compounds are all strong electrolytes. nonelectrolyte (a) H+(aq) and Br- (aq); (b) NH4+(aq) and Cl-(aq); (c) CH3CH2CH2CH2OH(aq) A HO (a) Ca(OH)2(s) 2 Ca2+(aq) + 2OH-(aq) (b) N2(g)


11.45 11.47 11.49 11.51 11.53 11.55 11.57 11.59 11.61 11.63 11.65 11.67 11.69






11.21 11.23

11.25 11.27 11.29


11.33 11.35


CH3OH(aq) (c) CH3OH(l) 2 In the solvent, hydrogen-bonding and dipole-dipole interactions are broken. In the solute, ionic bonds are broken. Ion-dipole interactions are formed in the solution. In both the solute and the solvent, London dispersion forces must be overcome. The new forces are also London dispersion forces. The attraction between the products (dissolved ions) and water molecules is weaker than the forces of attraction in the pure substances. Energy and entropy drive the formation of solutions. The intermolecular attractions between water molecules and oil molecules are not the same (water is polar and cooking oil is nonpolar). Since they are not "like," they are immiscible. Both molecules are polar and can engage in hydrogen bonding. (a) insoluble; (b) soluble; (c) soluble Ion-ion interactions are very strong. There is not enough energy released or entropy gained in the formation of ion-solvent interactions to make the dissolution process favorable. The ions in NaCl are attracted by ionic bonding. The most significant attractive force between water molecules is hydrogen bonding. Although disrupting the attractive forces in this solute and solvent requires considerable energy, formation of the ion-dipole interactions between solute and solvent particles releases enough energy to allow the ionic compound to dissolve. The solubility of gases such as O2 decreases as temperature increases. The solubility of oxygen and nitrogen gases should be lower at high altitude. The solubility of these gases is directly proportional to their pressures above the solution and these pressures are lower at high altitude. The insoluble gases that cause the bends can be redissolved into the blood by putting the divers in pressurized chambers. This increases the solubility of the gas in the blood. Decreasing the chamber pressure slowly (decompression) allows the gases (mostly excess nitrogen) to leave through the lungs.

11.71 11.73 11.75 11.77 11.79 11.81



11.87 11.89 11.91 11.93 11.95

17.0 g KCl On a macroscopic level, you can't dissolve more solid in a saturated solution, but you can dissolve more in an unsaturated solution. On a molecular level, the ions in a saturated solution are in equilibrium with the solid. You could add a little solid at a time until no more solid dissolves and you see traces of undissolved solid in the solution. The solution will be saturated and 0.85 g Ca(OH)2(s) will be left undissolved. 14.3% 105 g KI 5.423% Percent by volume is often used for mixtures of liquids. This is done because liquids are usually measured by volume. 30.4% (a) 2.2 ppb; (b) Since the maximum contamination level is 5 ppb, the water is safe to drink. 0.015 mol 2.41 m 3.0 M; 3.0 m 9.6 × 10-4%; 9.6 ppm; 9600 ppb (a) 9.0 ppb; yes, the concentration is below the EPA standard; (b) 9.0 × 10-6 mg/mL; (c) 9.0 × 10-4 mg; (d) 4.3 × 10-9 mol (a) 0.030 L KI; (b) 6.9 g PbI2 0.30 mol Na2CO3 57.14 mL 0.8812 M H2SO4 (a) 10.00 mL; (b) 52.50 mL; (c) 37.5 mL A colligative property is a physical property of a solvent that varies with the number of solute particles dissolved and not on the identity of those solute particles. Osmosis occurs when two solutions of different solute particle concentrations are separated by a membrane that allows the solvent but not the solute molecules to pass (semipermeable membrane). The volume of the more concentrated solution increases because solvent molecules migrate through the membrane to the more concentrated solution. Reverse osmosis occurs when solvent is forced to travel in a direction opposite of the direction it would travel during osmosis. Water will flow out of the blood cell, and the cell will collapse or shrink. (a) decrease; (b) increase; (c) decrease The boiling point of the water is elevated slightly because of the presence of the sugar. (a) 1.3°C; (b) 101.3°C The sodium chloride solution will have the higher boiling point because NaCl produces two solute particles when it dissolves in water.

Appendix D Answers to Selected Questions and Problems




11.99 11.101

11.103 11.105

11.107 11.109 11.111


11.115 11.117 11.119

Solution B has the highest concentration of dissolved particles and therefore has the lowest freezing point. Glucose is molecular, and images B and C show molecular compounds. Sodium chloride produces two ions of slightly different size as shown in image A. (a) 2.0 mol; (b) 5.0 mol; (c) 4.0 mol The osmotic pressure of MgCl2 will be greater than that of NaNO3 because each formula unit of MgCl2 produces three ions when it dissolves while NaNO3 produces only two ions. Vitamin D is nonpolar. The breaking of the crystal lattice and disruption of the intermolecular attractions between water molecules require energy. The formation of ion-dipole interactions between water and the ions releases energy. 182 g 8.90 M Volume changes with temperature. If the concentration units are based only on mass or number of moles, then they will not change with temperature. Percent by mass (a) and molality (d) do not change with temperature. The density of a dilute solution is approximately the density of the solvent. In aqueous solutions, for example, one liter of water has a mass of one kilogram. So the values for molarity (moles/liter) and molality (moles/kilogram) will be very similar. (a) 10.5 g BaSO4; (b) 2.33 g BaSO4; (c) 2.33 g BaSO4 2.0 × 103 g solution B


Ea EReactants

EProducts Reaction progress


12.23 12.25


12.29 12.31 12.33 12.35 12.37

Chapter 12 ­ Reaction Rates and Chemical Equilibria

12.1 (a) activation energy, Ea; (b) intermediate; (c) homogeneous equilibrium; (d) Le Chatelier's principle; (e) collision theory; (f) equilibrium constant expression increasing reactant concentrations, increasing temperature, or providing a catalyst grind the calcium carbonate to a fine powder; increase the temperature; increase the concentration of HCl catalyst No. If the reactants do not collide with sufficient energy or with the proper orientation, then the activated complex cannot be formed. (a) improper orientation; (b) proper orientation; (c) improper orientation activated complex minimum energy required to break bonds in reactants before reactants can be converted into products For similar reactions, the higher the activation energy, the slower the reaction.

12.3 12.5 12.7 12.9

12.39 12.41 12.43 12.45 12.47

12.11 12.13 12.15 12.17

12.49 12.51 12.53


Increasing temperature increases the kinetic energy and velocity of the molecules reacting. The fraction of colliding molecules with energy greater than the activation energy increases when temperature is increased. increasing temperature At room temperature, molecules of propane and oxygen do not collide with sufficient energy to cause a reaction. However, by providing a spark (heat), the kinetic energy of the molecules is increased. Combustion reactions are exothermic, so the reaction is sustained by the heat produced by the reaction. Roasts are generally thicker and more massive than steaks. For the same amount of heat, the steak gets hotter faster and therefore cooks faster. A catalyst provides a lower-energy route (pathway) to the same products. Catalytic converters use palladium and platinum catalysts (Figure 12.11). active site Catalysts are only temporarily changed during the course of a reaction (e.g., during the binding of the substrate). Chlorine atoms in the upper atmosphere act as catalysts for the decomposition of ozone. Since they act catalytically, they are not lost until they diffuse into outer space (which is a slow process) or are removed by some other reaction. intermediate (a) NO3 is an intermediate; (b) N2O5 + NO 3NO2 (a) H+ and Br- are catalysts and Br2 is an intermediate. (b) 2H2O2 2H2O + O2 When a reaction reaches equilibrium, the concentrations of reactants and products stop changing. A chemical equilibrium is a state of a reaction where the rate of the forward reaction is equal to the rate of the reverse reaction. The system has reached equilibrium in images D and E. (a) yes; (b) no; (c) yes; (d) yes You could measure the mass of the bromine liquid, but that would be tricky because it is very reactive and volatile. A better option is to measure the pressure of the gas produced. When the partial pressure has stabilized, the system is at equilibrium. Finally, bromine gas has a dark brown appearance. By measuring how much light the gas absorbs, we can determine if its concentration remains constant and the system is at equilibrium. The position of the equilibrium tells us whether a reaction is product-favored or reactant-favored.


Appendix D Answers to Selected Questions and Problems

12.57 12.59 12.61

The coefficients of the reactants and products are used as exponents in the equilibrium constant expression. The brackets represent molar concentrations. (a) Keq = HF H2 F2 CS2 H2

4 2 2


In the following table, the larger arrows () indicate the change made to the equilibrium reaction and the smaller arrows indicate what the concentration of the products and reactants do following the change. Finally, the central arrow indicates which direction the equilibrium shifts. NO(g) SO3(g) NO2 SO2

(b) Keq =

(a) (b) (c) (d) 12.95

CH4 H2S NO2 N2O4


(c) Keq = 12.63 12.65 12.67 12.69 12.71

12.73 12.75 12.77 12.79 12.81 12.83


(a) C(g) A(g) + B(g) (b) 2A(g) 4B(g) + C(g) + D(g) (a) The equations are inverses of each other. (b) 2NO2 2NO + O2; 2NO + O2 2NO2 0.25 An equilibrium constant only changes with temperature. (a) 3.1 × 10-3; (b) Because the value of Keq is much less than one, the position of equilibrium is described as "reactant-favored." (a) 1.3 × 10-9; (b) The reaction is reactant-favored since Keq < 1. The reaction is at equilibrium, so shifting will not take place. The reaction will proceed in the reverse direction. The reaction will proceed in the forward direction. Heterogeneous reactions are those in which the products and reactants are not all in the same physical state. The concentrations of pure liquids and solids do not change during chemical reactions; only the amount of the pure substance changes. Because the concentrations of pure substances do not change during reactions, their values are essentially incorporated into the equilibrium constant. (a) Keq = NH3 HCl (b) Keq = Ca2+ CO32­

+ ­ (c) Keq = H3O F HF

12.97 12.99 12.101 12.103

12.105 12.107 12.109


12.87 12.89


0.27 M Le Chatelier's principle describes how reactions that are at equilibrium respond to changes in reaction conditions. If stress is put on a reaction, the reaction responds to counteract that stress. (a) The reaction shifts to the left. (b) The concentrations of the products, CO2 and H2, will decrease and the reactant, H2O, will increase in concentration.


(a) This has no effect since the concentration of the solid does not change when more solid is added. (b) This decreases the concentration of Pb2+ by shifting the equilibrium to the left. (c) This causes an increase in Pb2+ concentration. When the concentration of lead ion is increased, the equilibrium shifts to the left. However, it never shifts enough to completely remove all the added material. (a) left; (b) right; (c) no effect (a) no; (b) yes (a) left; (b) right If the equilibrium shifts to the right, the concentration of products at equilibrium is higher. The value of Keq increases. (a) decrease; (b) increase exothermic An increase in the concentration of NO requires that the equilibrium shift to the right. (a) Water is in the liquid state; removing it does not affect the equilibrium. (b) There are fewer moles of gas on the right. The reaction shifts right increasing the concentration of NO. (c) Since the reaction is exothermic, the reaction will shift right with decreasing temperature producing a higher concentration of NO. (d) Adding oxygen causes the equilibrium to shift right producing a higher concentration of NO. (e) Catalysts do not have an effect on equilibrium concentrations. The reaction that produces light slows down, and the reactants take longer to be consumed. The lower light output of the cold light sticks is a result of fewer photons being produced per second (lower reaction rate). It provides a different reaction mechanism with lower activation energy. In general it splits one difficult step (high activation energy) into two or more steps that have lower activation energies. H2 I2 HI


12.115 (a) Keq =

(b) Keq = 0.0199 (c) The position of the equilibrium favors the reactants. 12.117 The effect of temperature depends on whether the reaction is endothermic or exothermic. 12.119 The reaction must be exothermic since the reaction shifted to the right to form the powder.

Appendix D Answers to Selected Questions and Problems


12.121 The simplest treatment would be to increase the concentration of oxygen that the patient is breathing. By increasing the O2 concentration, the second equilibrium will shift to the left and the carbon monoxide can be displaced and exhaled through the blood.

Chapter 13 ­ Acids and Bases

13.1 (a) strong base; (b) Brønsted-Lowry theory; (c) conjugate acid; (d) polyprotic acid; (e) acidic solution; (f) weak base; (g) amphoteric substance; (h) ion-product constant of water, Kw; (i) self-ionization (a) -9; (b) -11.0; (c) 3.87; (d) 5; (e) 0.0 (a) 16; (b) 6.3 × 10-7; (c) 1 Acids have a sour taste (but you should never taste anything to see if it's an acid!), are corrosive to many metals, turn blue litmus red, and neutralize bases. Common "household" bases and their uses: Formula Mg(OH)2 Al(OH)3 NH3 Ca(OH)2 13.11 Name magnesium hydroxide aluminum hydroxide ammonia calcium hydroxide Use antacid antacid cleaning solutions, smelling salts concrete

13.3 13.5 13.7



13.15 13.17 13.19 13.21

13.23 13.25 13.27



An Arrhenius acid produces hydrogen ions, H+(aq), when dissolved in water. An Arrhenius base produces hydroxide ions, OH-(aq), when dissolved in water. Brønsted-Lowry theory refers to the transfer of hydrogen ions from an acid to a base. Arrhenius similarly describes the release of hydrogen ions into the solution, but identifies OH- as a component of the base. In Brønsted-Lowry theory, a base is any substance that can accept a hydrogen ion (including OH-). (a) acid; (b) base; (c) acid (a) NO2-; (b) F -; (c) H2BO3- (a) H2O; (b) C6H5NH3+; (c) H2CO3 (a) HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) (b) HClO4(l) + H2O(l) H3O+(aq) + ClO4-(aq) - (c) CH3CO2 (aq) + H2O(l) HCH3CO2(aq) + OH-(aq) Both water (a) and hydrogen sulfite (b) are amphoteric. Carbonic acid has two acidic hydrogen atoms. Strong acids and bases ionize or dissociate completely (100%). For weak acids and bases, the percent of molecules that ionize is usually much lower (typically a few percent). (a) strong acid; H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq) (b) strong base; Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq) (c) weak base; Na2CO3(aq) + H2O(l) HCO3-(aq) + 2Na+(aq) + OH-(aq) (d) weak acid; H3C6H5O7(aq) + H2O(l) H3O+(aq) + H2C5O7-(aq) (e) weak base; C6H5NH2 + H2O(l) C6H5NH3+(aq) + OH-(aq) A is NH3, B is HF, and C is HCl.

KF; HF is a weak acid, so its conjugate base, F-, is a weak base. 13.35 Acid A is a stronger acid. The stronger the acid, the greater the extent of ionization. 13.37 0.50 M HCO2H 13.39 NaCH3CO2 13.41 In a strong acid solution the acid molecules are fully ionized. It does not necessarily mean that the concentration is high. A concentrated acid solution has a high percentage of acid, but this does not necessarily mean that the acid is ionized to any great extent. 13.43 A polyprotic acid is an acid that contains more than one acidic hydrogen atom. 13.45 Hydrosulfuric acid is a weak acid that has two acidic hydrogen atoms. The corresponding acid-base reactions are. H2S(aq) + H2O(l) HS-(aq) + H3O+(aq) Ka1 = 8.9 × 10-8 - HS (aq) + H2O(l) S2-(aq) + H3O+(aq) Ka2 = 1.0 × 10-19 13.47 The following substances can all be found in a solution of oxalic acid: H2C2O4, HC2O4-, C2O42-, H3O+, H2O. The substance with the lowest concentration is the oxalate ion C2O42-. 13.49 Like other equilibrium constants, it is constant only under conditions of constant temperature. 13.51 In pure water, [OH-] = [H3O+] = 1.0 × 10-7 M. 13.53 (a) basic; (b) acidic; (c) basic 13.55 (a) 1.0 × 10-11 M, basic; (b) 1.0 × 10-3 M, acidic; (c) 3.1 × 10-7 M, acidic 13.57 (a) 0.010 M; (b) 0.020 M; (c) 6.7 × 10-13 M 13.59 A decrease in pH represents an increase in the acidity. 13.61 A 10-fold increase in hydrogen ion concentration is 1 pH unit change. The difference in pH units is 1. 13.63 Practically speaking, most solutions of acids and bases are lower in concentration than 1 M so the pH scale typically ranges between 0 and 14. However, the pH values can exceed this range. 13.65 No. A solution is defined as acidic if the pH is less than 7. 13.67 (a) 3.00, acidic; (b) 13.00, basic; (c) 9.47, basic 13.69 (a) 10.00, basic; (b) 7.00, neutral; (c) 4.91, acidic 13.71 (a) 2.00, acidic; (b) 1.70, acidic; (c) 12.18, basic 13.73 Underlined values were given in the problem. 13.33 [H3O+] (a) (b) (c) (d) (e) 13.75 13.77 13.79 1.0 × 10-5 1.0 × 10-10 1.0 × 10-6 2.9 × 10-9 1.1 × 10-5 [OH-] 1.0 × 10-9 1.0 × 10-4 1.0 × 10-8 3.5 × 10-6 9.0 × 10-10 pH 5.00 10.00 6.00 8.54 4.95 pOH 9.00 4.00 8.00 5.46 9.05 Acidic or Basic? acidic basic acidic basic acidic

pH = 2.30, pOH =11.70. The pH and pOH always add to 14.00 at 25°C. (a) 1.0 × 10-5 M, acidic; (b) 1.0 × 10-12 M, basic; (c) 1.3 × 10-6 M, acidic (a) 1.0 × 10-11 M, basic; (b) 4.0 × 10-8 M, basic; (c) 1.3 × 10-2 M, acidic


Appendix D Answers to Selected Questions and Problems

13.81 13.83

13.85 13.87 13.89


13.93 13.95 13.97


13.101 13.103

13.105 13.107 13.109

13.111 13.113



(a) 1.0 × 10-13 M; (b) 3.2 × 10-4 M; (c) 4.0 × 10-11 M Acetic acid is a weak acid, so the concentration of hydronium ion would be less than 0.010 M, and the pH would be higher than 2.0. 0.1 M A pH meter typically gives measurements to two decimal places (Figure 13.17). Acid-base indicators are typically colored compounds that are also weak acids or bases. The colors of these compounds depend on whether they are in their acid or base forms. The color range of an indicator indicates the pH values over which the indicator changes color. The hue of the indicator can be used to estimate the pH of a solution. The range over which an indicator changes color is about 1.5 pH units. A selection of indicators would include thymolphthalein. It should be red with a small amount of yellow-orange. Red (or the presence of red color) indicates the pH is less than 2.5. Yellow indicates a pH between about 2.5 and 8.0. Blue (or a blue hue) indicates pH values above 8.0. It will be completely blue above pH 9.5. Universal indicator is a mixture of different indicators. The color of an unknown solution with the indicator present gives a good estimate of the pH. It is similar to pH paper in that pH paper contains universal indicator. The solution would be yellow at the start of the titration and turn blue at the endpoint. A buffer contains both a weak acid and its conjugate base (or a weak base and its conjugate acid). As a result, when small amounts of acid or base are added, the buffer can absorb the excess hydronium or hydroxide with only a small change in pH. The buffer system of the blood helps to maintain a pH between 7.35 and 7.45. any salt containing CH3CO2- (e.g., NaCH3CO2) (a) Hydrogen phosphate, HPO42-, is the acid and the conjugate base is phosphate, PO43-. HPO42-(aq) + H2O(l) PO43-(aq) + H3O+(aq) (b) When acid is added, hydronium, H3O+, is produced and the pH momentarily drops (more acidic). However, the excess hydronium quickly reacts with the PO43- shifting the equilibrium to the left to produce the weak acid HPO42-. Solutions described in (b) and (c) produce buffered solutions. As the concentration of carbon dioxide increases, the equilibrium of the reaction would shift to the right. The formation of carbonic acid would cause a decrease in the pH of the extracellular fluid. Ammonia, NH3, has a lone electron pair in its Lewis structure. Methane does not have a lone electron pair, so it cannot accept a hydrogen ion (proton) in an acidbase reaction. When NaF dissolves in water, sodium and fluoride ions are produced. The fluoride ions react with water to produce a weak base HF by the reaction: F -(aq) + H2O(l) HF(aq) + OH-(aq)





In general the conjugate base (F-) of a weak acid will react with water to produce hydroxide ions. The molecular species HF is stable, so its formation drives the formation of hydroxide. This solution will be basic. The sodium ion does not react with water. No. A neutral solution has a pH = pOH = 7.00 (at 25°C). If sodium hydroxide is added, the hydroxide ion concentration will have to increase. The solution will have a pH greater than 7.00. Both H2SO4 and HNO3 are strong acids. Both will dissociate completely for the first ionization. However, H2SO4 is a polyprotic acid. The additional hydronium ion produced by the second ionization will make the pH of the sulfuric acid lower. (a) HOCl(aq) + H2O(l) OCl-(aq) + H3O+(aq). (b) The HCl reacts with the water to produce hydronium ion: HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq). The + added H3O causes the equilibrium to shift to the left and some of the excess hydronium is consumed. The HOCl concentration will increase and the OCl- concentration will go down. While the pH may not drop below 7.00, any time an acid is added to a solution, the solution becomes more acidic (pH is lowered). Sulfuric acid is a strong acid and is completely ionized.

H3O+ SO42­


13.127 sodium formate and formic acid

Chapter 14 ­ Oxidation-Reduction Reactions

14.1 (a) electrolytic cell; (b) cathode; (c) oxidation; (d) halfreaction; (e) electrochemistry; (f) electrode; (g) oxidizing agent; (h) oxidation-reduction reaction; (i) corrosion; (j) oxidation number reaction where electrons are transferred when there is an increase in charge (oxidation number) (a) Mg(s) + Cu(NO3)2(aq) Cu(s) + Mg(NO3)2(aq) (b) The charge of magnesium increases by two; 2 electrons transferred; electrons are lost. (c) The charge of copper decreases by two; 2 electrons transferred; electrons are gained. (a) Mg; (b) Sn2+; (c) Sn2+; (d) Mg (e)

Mg Mg2+

2+ Sn Sn

14.3 14.5 14.7




An oxidation number is a bookkeeping method for tracking where electrons go during a chemical reaction.

Appendix D Answers to Selected Questions and Problems


14.13 14.15 14.17 14.19 14.21 14.23 14.25 14.27 14.29


14.33 14.35


(a) 0; (b) 0; (c) 3+; (d) 1- S = 6+, O = 2- (a) 5+; (b) 3+; (c) 4+ (a) 5+; (b) 5+; (c) 5+; (d) 1+; (e) 3-; (f) 3+ S = 6+, O = 2- (a) 7+; (b) 5+; (c) 1-; (d) 3+; (e) 5+ (a) Cl = 4+, O = 2-; (b) Ca = 2+, F = 1-; (c) H = 1+, Te = 4+, O = 2-; (d) Na = 1+, H = 1- (a) N = 3+, O = 2-; (b) Cr = 6+, O = 2-; (c) Ag = 1+, Cl = 1-; (d) S = 4+, O = 2-; (e) C = 4+, O = 2- (a) not redox; (b) redox, N2 is the oxidizing agent, H2 is the reducing agent; (c) not redox; (d) not redox; (e) redox, O2 is the oxidizing agent, C2H6 is the reducing agent (a) N2(g) + O2(g) 2NO(g). (b) The oxidation number increases. Each nitrogen atom loses two electrons to change from an oxidation number of 0 to 2+. (c) The oxidation number decreases. Each oxygen atom gains two electrons to change from an oxidation number of 0 to 2-. (a) V2+; (b) Cr2O72-; (c) Cr2O72-; (d) V2+ (a) I2 is both the oxidizing agent and reducing agent, 1 electron transferred. (b) Cr = reducing agent, H+ = oxidizing agent, 2 electrons transferred. (c) Cr2O72- is both the oxidizing and reducing agent, 12 electrons transferred. (d) Fe3+ = oxidizing agent, Al = reducing agent, 3 electrons transferred. In the anode half-cell the half-reaction is Fe(s) Fe2+(aq) + 2e- In the cathode half-cell the half-reaction is Ni2+(aq) + 2e- Ni(s)





14.49 14.51






Salt bridge



14.59 14.61




The iron electrode will lose mass as Fe is oxidized and becomes Fe2+(aq). The nickel electrode will gain mass as Ni2+ is reduced to Ni(s).

14.63 14.65 14.67

Iron electrode

Nickel electrode

oxidation half-reaction: Cd(s) + 2OH-(aq) Cd(OH)2(s) + 2e- reduction half-reaction: 2e- + NiO2(s) + 2H2O(l) Ni(OH)2(s) + 2OH-(aq) - 3+ (a) 3e + Fe (aq) Fe(s) (b) Zn(s) Zn2+(aq) + 2e- (c) 2Cl-(aq) Cl2(g) + 2e- (d) Fe2+(aq) Fe3+(aq) + e- (a) Zn(s) + 2NO3-(aq) Zn(NO3)2(aq) + 2e- - 3e + Fe(NO3)3(aq) Fe(s) + 3NO3-(aq) 3Zn(s) + 2Fe(NO3)3(aq) 3Zn(NO3)2(aq) + 2Fe(s) (b) Mn(s) + 2Cl-(aq) MnCl2(aq) + 2e- 2e- + 2HCl(aq) H2(g) + 2Cl-(aq) Mn(s) + 2HCl(aq) MnCl2(aq) + H2(g) Fe2(SO4)3(aq) + 2KI(aq) 2FeSO4(aq) + K2SO4(aq) + I2(aq) (a) Ba(s) Ba2+(aq) + 2e- (b) e- + H+(aq) + HNO2(aq) NO(g) + H2O(l) (c) 2e- + 2H+(aq) + H2O2(aq) 2H2O(l) (d) 2Cr3+(aq) + 7H2O(l) Cr2O72-(aq) + 14H+ + 6e- (a) 3OH-(aq) + La(s) La(OH)3(s) + 3e- - - (b) 2e + H2O(l) + NO3 (aq) NO2-(aq) + 2OH-(aq) (c) H2O2(aq) + 2OH-(aq) O2(g) + 2e- + 2H2O(l) (d) 8e- + 3H2O(l) + Cl2O7(aq) 2ClO2-(aq) + 6OH-(aq) (a) 8H+(aq) + 3H2S(aq) + Cr2O72-(aq) 3S(s) + 2Cr3+(aq) + 7H2O(l) + 2+ (b) 4H (aq) + 5V (aq) + 3MnO4-(aq) 5VO2+(aq) + 3Mn2+(aq) + 2H2O(l) (c) 6H+(aq) + 6Fe2+(aq) + ClO3-(aq) 6Fe3+(aq) + Cl-(aq) + 3H2O(l) - (a) 2NH3(aq) + ClO (aq) N2H4(aq) + Cl-(aq) + H2O(l) (b) 2Cr(OH)4-(aq) + 3HO2-(aq) 2CrO42-(aq) + 5H2O(l) + OH-(aq) - (c) 2OH (aq) + Br2(aq) Br-(aq) + BrO-(aq) + H2O(l) 24H+(aq) + 5C6H12O6(aq) + 24NO3-(aq) 30CO2(g) + 12N2(g) + 42H2O(l) The anode could be any metal that is more active than iron (e.g., aluminum or zinc). The supporting electrolyte could be any soluble salt that does not precipitate with Fe3+ (e.g., potassium nitrate since all nitrates are soluble). The anodic half-cell would also contain a salt that includes the metal used for the anode (e.g., aluminum or zinc nitrate). Au < Bi < Ni < Zn < Al < Ca Electrolysis is the process of using electricity to make a nonspontaneous redox reaction occur. (a) Reduction will begin to take place at the cathode to form sodium metal. (b) Oxidation will begin to take place at the anode to form I2(g). (c) Na+(l) + e- Na(l) (reduction) 2I-(l) I2(g) + 2e- (oxidation) (d) 2Na+ + 2I-(l) 2Na(l) + I2(g)


The Pb(s) gets oxidized and the PbO2(s) is reduced. The oxidizing agent is PbO2(s) and the reducing agent is Pb(s).


Appendix D Answers to Selected Questions and Problems Anode Cathode




I2(g) Anode reaction Cathode reaction Na(l)




14.75 14.77 14.79


14.83 14.85

14.87 14.89

Magnesium is more active than iron. The magnesium will reduce the O2(g) and keep the Fe in the reduced state. When the Mg(s) is used up, the iron will begin to corrode. One of the methods of preventing corrosion is to prevent contact with oxygen and moisture. The tin coating serves as a barrier to moisture and oxygen. When the tin coating is scratched, the more active metal, iron, is more readily oxidized. The chromium should corrode first because it is a more active metal than iron. (a) 3-; (b) 2-; (c) 3+; (d) 1-; (e) 5+; (f) 3+ (a) 2H2O(l) + 2OH-(aq) + 2CoCl2(s) + Na2O2(aq) 2Co(OH)3(s) + 4Cl-(aq) + 2Na+(aq) The oxidizing agent is Na2O2 (oxygen is being reduced from 1- to 2-). The reducing agent is CoCl2 (cobalt is being oxidized from 2+ to 3+). (b) 2OH-(aq) + Bi2O3(s) + 2ClO-(aq) 2BiO3-(aq) + 2Cl-(aq) + H2O(l) The oxidizing agent is ClO- (chlorine is being reduced from 1+ to 1-). The reducing agent is Bi2O3 (bismuth is being oxidized from 3+ to 5+). NH4+(aq) + NO3-(aq) N2O(g) + 2H2O(l) The oxidizing agent is NO3- (N is being reduced from 5+ to 1+). The reducing agent is NH4+ (N is being oxidized from 3- to 1+). (a) Al; (b) Zn; (c) Mn; (d) Mg Brass is an alloy of copper and zinc. Cu doesn't react with acid. There must be a protein in sweat that facilitates the oxidation of the copper to form Cu2+. There are several compounds of copper that are green (including CuCO3). 2H+(aq) + Cu(s) + H2SO4(aq) Cu2+(aq) + SO2(g) + 2H2O(l) + (a) reducing agent, Tl ; oxidizing agent, Ce4+; oxidation half-reaction: Tl+(aq) Tl3+(aq) + 2e4+ reduction half-reaction: 2Ce (aq) + 2e2Ce3+(aq) (b) reducing agent, Br ; oxidizing agent, NO3 oxidation half-reaction: 2Br-(aq) Br2(l) + 2ereduction half-reaction: 2NO3 (aq) + 4H+(aq) + 2e2NO2(g) + 2H2O(l)

15.13 15.15 15.17 15.19 15.21 15.23 15.25 15.27

Of the elements with atomic number less than 84, technetium (Tc), promethium (Pm), and astatine (At) have no stable isotopes. No isotopes of atomic number 84 or greater are stable. Dense materials are better at shielding gamma radiation. Lead, iron, and aluminum (metals) can be used, but concrete and water are effective only if thick layers are used. Nuclear decay reactions and bombardment reactions result in the production of radiation. In addition, new isotopes are produced by both. During nuclear decay alpha particles, beta particles, gamma rays, and positrons are produced by the nucleus. During nuclear reactions, nuclear charge and mass numbers are conserved. 222 86 Rn

99 43 Tc 15 7N

beta particle 7 4 Be (a) 238U 92 (b) 234Th 90 (c) 234U 92 (a) 227 Fr 87 (b) 18 F 9

0 (d) 210Bi+ -1e83 234 90Th

+ 4 2

0 + -1-

234 91Pa 4 2

+ 230 Th 90

210 82Pb


4 2

+ 223 At 85

+ +0+ + 18O 1 8


15.33 15.35

0 14 (c) 14 C 6 -1 + 7N In a nuclear bombardment reaction a particle accelerator is used to accelerate particles or nuclides to very high velocities. The beam of particles produced is aimed at a target element. If the particle strikes the target with sufficient energy to fuse the target and beam particles, a newly created nucleus is formed. The newly created nucleus spontaneously decays to produce new substances and decay particles. 257 103 Lr

(a) 10 B + 2 H 5 1 (b) 209 Bi + 4 83 2 (c) 12 C + 1 p 6 1 (a) 96 Mo + 1 p 42 1 (b) 209 Bi + 64 Ni 83 28

11 6C

+ 1n 0 + 21n 0

211 85 At 13 7N 96 43 Tc

15.37 15.39 15.41

+ 1n 0 + 1n 0

272 111Rg

15.43 15.45 15.47 15.49

Chapter 15 ­ Nuclear Chemistry

15.1 (a) critical mass; (b) fission; (c) alpha particle; (d) radioactive; (e) gamma radiation; (f) nuclear bombardment; (g) nucleon; (h) radiation A substance that contains unstable nuclides of an element will exhibit radioactivity by the emission of radiation. Radioactivity produces photons (usually gamma rays) and particulate radiation (i.e., beta rays, positrons, neutrons, and alpha particles).

15.3 15.5


approximately 1:1 If the mass numbers of all the stable isotopes are graphed versus their atomic number, a pattern emerges (Figure 15.14). The band of stability is the outline of that pattern. The isotope 16 O should be the most stable. It has a N/Z 8 ratio of 1. (a) beta; (b) alpha and beta; (c) alpha and beta; (d) positron; (e) beta 234 0 54 + 2-1- + 214 Bi 91Pa 2 83 In a sample of radioactive material, the half-life is the time it takes for half of the radioactive nuclei to decay. For nuclear decay, each half-life has exactly the same duration. Film, Geiger-Müller counter, and a scintillation counter can be used to detect radiation.

Appendix D Answers to Selected Questions and Problems


15.53 15.55 15.57 15.59 15.61 15.63 15.65

15.67 15.69


15.73 15.75

15.77 15.79




15.87 15.89 15.91 15.93

One-sixteenth (6.25%) of the original material will be present. three half-lives or 42 days 22.5 hours (three half-lives) 5730 years Twenty-four hours represents four half-lives. The half-life is 6 hours. Radiation-powered devices can last much longer than battery-powered devices. 99m Tc is used because it produces radiation that can be used to image tumors and its short half-life ensures that it won't stay in the body for a long period of time. Positron emission occurs for N/Z ratios < 1. 11 C has a low 6 N/Z ratio, whereas 14 C has a high N/Z ratio. 6 Radiation can (1) have no effect, (2) cause reparable damage, (3) cause damage leading to formation of mutant cells (cancer cells), (4) kill cells (subsequently absorbed by the body). Mutant cells can also be recognized as foreign entities and be destroyed by an immune response. This is a difficult question to answer. How you answer the question stems from your view of our governmental system and your perception of the risk of radon and the benefit of regulations. Clearly, the best protection against unnecessary radon exposure is knowledge of what causes enhanced levels of radon and what measures can be used to reduce radon exposure. The Internet can be a good source of information if you use it carefully (i.e., You would most likely find 210 Pb. 82 Fission occurs when unstable nuclei decay into smaller nuclei and decay products. Fusion results when smaller particles combine to make larger ("heavier") nuclei. three In principle, a chain reaction needs the terminating step of the reaction to be a necessary component of the initiating step. For example, if the products of a reaction (initiating step) start a second reaction (terminating step) and the products of the second reaction start the first reaction, you have a chain reaction. A fission reactor has fuel rods (containing the radioactive material), control rods (to absorb the excess neutrons), a moderator (i.e., water, heavy water, or graphite), and energy converters (convert water to steam which powers electric generators). The water is not directly converted to steam by the nuclear reaction as a safety precaution. The steam that would be produced by the nuclear reaction would be highly radioactive and is also made of poisonous heavy water (water with deuterium rather than hydrogen is highly toxic). If there were a leak of the steam, the released products would be very dangerous. The waste contains a mixture of fission by-products as well as materials contaminated by radioactive waste (such as the spent fuel rods). Isotopes of hydrogen (H-1, H-2, H-3) and lithium are most likely to be used. 2 3 4 1 1H + 1H 2 He + 0 n

3 + 2H 1 2 He In radioactive dating, a sample is taken and analyzed for the quantity of a particular radioisotope. Based on the amount of the isotope, and assuming you know how much 1 1H

of the isotope was initially present in the sample, you can calculate the sample age (using the half-life). For example, 1 if the radioisotope is at -- the original concentration, two 4 1 1 half-lives have passed (i.e., -- × --). 2 2 15.95 15.97 The sample contains about -- the carbon-14 in a modern 64th sample. This is about 6 half-lives or 34,000 years. 1 1 2 0 + 1H + 1H 1H + +1

2 1H 1 1H 1

+ 1H 1

3 2He

4 0 + + 3He 2 2He + +1 15.99 The initial product of the bombardment reaction is car13 bon-13 (9 Be + 4He 4 2 6 C). Loss of a neutron would produce carbon-12 (N/Z =1): 9 4 12 1 4 Be + 2He 6 C + 0 n. 15.101 With a half-life of 138.4 days, at the end of one year approximately 3 half-lives have already passed. Only oneeighth of the material would be remaining. You would be better off buying brushes every year (or as you needed them). 15.103 Even though the element chlorine has only been known to exist since 1774, chloride-containing compounds were used well before the discovery of chlorine.

Chapter 16 ­ Organic Chemistry

16.1 (a) isomer; (b) alcohol; (c) ester; (d) aldehyde; (e) cyclic hydrocarbon; (f) saturated hydrocarbon; (g) alkyl group; (h) cycloalkane; (i) aliphatic hydrocarbon; (j) amino acid; (k) organic chemistry; (l) unsaturated hydrocarbon (a) seven single bonds; (b) one double bond, four single bonds ; (c) one triple bond, two single bonds (a) amine; (b) ketone; (c) alcohol; (d) alkene; (e) carboxylic acid; (f) alkyne; (g) aldehyde; (h) ether; (i) ester molecular formula: C20H28O structural formula: H H H HH H H C C H H C H H H C H H H C HC C C C C C C H C C C O C H H H H H H C C C C H H H H H line structure: O 16.9 16.11 (a) C7H14O; (b) C8H18 (a) (b) O

16.3 16.5 16.7

(c) (d) O C H


Appendix D Answers to Selected Questions and Problems




(a) (CH3)2CHCH2CH3; (b) (CH2OH)2CHOH; (c) CH3OCH2CH3 line structures: (a) OH HO (c) O H H (b) H H C H H C C


16.39 HO 16.41 16.43 16.45


(a) CH3CH(CH3)CH3; (b) CH3CH(CH3)C(CH2CH3)2CH2CH3 (c) CH3CH(CH3)CH2CH(CH3)CH2CH2CH2CH3 (d) CH3CH2C(CH2CH3)(CH3)CH2CH2CH3 heat (a) CH3CH2CH3 + Cl2 CH3CHClCH3 + HCl heat or CH3CH2CH3 + Cl2 CH3CH2CH2Cl + HCl heat (b) (CH3)3CCH2CH2CH3 + 11O2 7CO2 + 8H2O heat (c) CH3CH3 + Br2 CH3CH2Br + HBr An alkene is a hydrocarbon containing one or more carbon-carbon double bonds. (a) 5-methyl-3-heptene; (b) 2,3-dimethyl-2-hexene; (c) trans-3,4-dimethyl-3-hexene only (c) because it can be made to be asymmetric along the carbon-carbon double bond axis (a) 3-methyl-1-butene (b) propyne trans-2-butene




(c) (d) 16.47


16.19 16.21


16.25 16.27 16.29

H H H (c) Cl Cl C C H H H H An alkane hydrocarbon has only carbon-carbon single bonds, an alkene contains at least one carbon-carbon double bond, and an alkyne has at least one carbon-carbon triple bond. (a) unsaturated; (b) unsaturated; (c) unsaturated Aromatic hydrocarbons are cyclic hydrocarbons that can be drawn with alternating double and single bonds between the carbon atoms. When the bonds are distributed in this way, the electrons are said to be delocalized. The London dispersion forces must be stronger in butane than in 2-methyl propane (isobutane). London dispersion forces increase with the length of the molecule because the molecules have more opportunities to interact. Because the London dispersion forces are stronger in butane, it boils at a higher temperature. (a) CH3CH2CH3; (b) CH3(CH2)4CH3 (a) ethane; (b) n-octane or octane; (c) propane; (d) n-nonane or nonane There are five isomers of C6H14.

cis-3-hexene (a) CH3CH=CHCH3 + HBr (b) CH3CH2CH=CHCH3 + Br2 (c) +

H2 Pd


16.49 16.51 16.53

Aromatic hydrocarbons are cyclic and have alternating single and double bonds (delocalized electrons). (a) 1,2-dibromobenzene; (b) 1-chloro-4-ethylbenzene H H (a) H H C C C C H H H C C C C C H H H H H 1,2,4-trimethylbenzene H H (b) C C H C C Cl C C H H chlorobenzene

16.55 16.57 16.59 16.61 16.63

C6H6 + Cl2 C6H5Cl + HCl Both (b) and (c) are alcohols. There are four isomers. (a) 2-methyl-1-propanol; (b) 2-butanol (a) OH 3-hexanol (b)


16.31 16.33

(a) 2-methylpentane; (b) 2-methylbutane; (c) 2-methylbutane (same as previous structure) (a) 3-methylpentane; (b) 2,2-dimethylpropane; (c) 2-methylbutane

OH 2,2-dimethyl-3-pentanol

Appendix D Answers to Selected Questions and Problems



16.67 16.69 16.71 16.73

The primary difference is that the intermolecular forces between ethanol (CH3CH2OH) molecules are hydrogen bonds, while those between propane molecules are London dispersion forces. The stronger intermolecular forces between ethanol molecules means that a higher temperature (energy) is needed to raise the vapor pressure of ethanol to 1 atm. O Ethers contain the functional group -O- attached to two alkyl groups . R R Only (a) CH3OCH3 is an ether. (a) ethylpropyl ether; (b) dimethyl ether O (a) (b) O ethyl methyl ether



ethyl phenyl ether 16.75 Water is a hydrogen-bonding solvent, so substances that are hydrogen-bonding, polar molecules will be more soluble in water. Ethers (CH3OCH3) do not directly hydrogen-bond because the oxygen is attached to carbon atoms on either side. Alcohols are more water soluble because the alcohol group -OH can hydrogen-bond. O An aldehyde has the functional group R-CHO where R can be an organic group or a hydrogen atom. C R H Compounds (a), (c), and (d) are aldehydes. Both compounds are polar because of the C=O bond, but ethanal (CH3CHO) has stronger London dispersion forces because it has a larger molecular mass. As a result ethanal has a higher boiling point. A carboxylic acid contains the functional group R-COOH or R-CO2H

16.77 16.79 16.81

( )

16.83 16.85 16.87



H where R is an organic group.


16.89 16.91 16.93

One route to the formation of esters is the condensation of a carboxylic acid and an alcohol. Methyl propionate (CH3CH2CO2CH3) can be formed from methanol (CH3OH) and propionic acid (CH3CH2CO2H). Saturated fatty acids are high-molar-mass carboxylic acids which do not have double or triple bonds in the hydrocarbon chain. As a result, it is no longer possible to add any more hydrogen atoms to the hydrocarbon chain and they are said to be saturated with hydrogen. Acetic acid, CH3CO2H, should have a higher boiling point than propanal, CH3CH2CHO, because acetic acid can hydrogen-bond. Aldehydes are polar but do not hydrogen-bond. N An amine is a derivative of ammonia, NH3, with one or more of the H atoms replaced by organic groups R H . H NH2 (b) (c) (a) N N N or H

( )

16.95 16.97 16.99

Methylamine, CH3NH2, can hydrogen-bond, so it has relatively strong intermolecular forces resulting in a higher boiling point. Ethane, CH3CH3, is nonpolar and cannot hydrogen-bond. (CH3CH2)2NH + HCl(aq) (CH3CH2)2NH2+(aq) + Cl-(aq) A simple organization of hydrocarbons is given in the chart below.


Straight or branched













Appendix D Answers to Selected Questions and Problems

16.101 Alkenes and alkynes differ primarily in three ways: carbon-carbon bond type, bond angles, and number of hydrogen atoms in the structure. Alkanes Bond type Bond angle General formula (straight or branched) Single 109.5° #H = 2n + 2 Alkenes Double bonds 120° #H = 2n Alkynes Triple bonds 180° #H = 2n - 2

#H = number of hydrogens; n = number of carbons. The formula assumes only one double or triple bond per structure. 16.103 Aldehydes and ketones are similar functional groups. O A ketone has the general formula O

16.105 16.107


16.111 16.113 16.115 16.117 16.119 16.121

. The aldehyde functional group is . C C R H R R (a) 2,4-dimethylpentane; (b) 1,3-dimethylcyclopentane; (c) 2,2,3,3-tetramethylbutane One possible test is to compare the solubility of each substance in an alcohol (such as ethanol). However, this test is not very specific. Addition of bromine, Br2, which is a reddish brown solution would be a better alternative. Bromine will react with an alkene but not with an alcohol. If an alkene is present, the bromine solution will lose color as the bromine reacts with the alkene. The trend in boiling points is directly related to the strength of the intermolecular forces in each substance. The weakest attractive forces are found in propane which is nonpolar. Dimethyl ether is polar, but ethanol is both polar and has hydrogen bonding. Because hydrogen bonds are so strong, ethanol has the highest boiling point. (a) CH3COCH3; (b) CH3CH2COCH3; (c) CH3CH2CHO; (d) CHOCHCHCH3; (e) CH3CHO (a) CH3CH2CH2CO2CH2CH3; (b) CH3CH2CH2CH2CO2H; (c) CH3CO2CH3; (d) HCO2H; (e) CH3CH2CO2H (a) dimethylamine; (b) ethyldimethylamine (the official IUPAC name is N,N-dimethyl-1-aminoethane, which is why we use the common name); (c) aniline (a) CH3C(CH3)2CH2CH(CH3)2; (b) CH3CHC(CH2CH3)2; (c) CH2CHCHCH2; (d) C4H8; (e) C6H5Cl (a) alkane; (b) ether; (c) carboxylic acid; (d) aldehyde; (e) ester; (f) amine; (g) ketone Only (c) 2-butene and (d) 2-pentene can exist as cis and trans isomers. (c) (d)

16.123 (a) isomer; (b) isomer; (c) same compound

Chapter 17 ­ Biochemistry

17.1 17.3 17.5 (a) transfer RNA; (b) secondary structure; (c) nucleotide; (d) nucleic acids; (e) lipid; (f) primary structure; (g) deoxyribonucleic acid; (h) fatty acid; (i) active site; (j) carbohydrate; (k) peptide; (l) replication; (m) translation regulation of cell process, control reaction rates, transport substances in and out of cells, fight disease COOH H2N C H R 17.7 COOH H COOH CH









17.9 17.11

Alanine has a methyl group and valine has an isopropyl group. O H O Ala-Gly H2N CH C CH3 O H Gly-Ala H2N CH C H N CH CH3 N CH H O C OH ; The methyl group is on the C-terminal amino acid. C OH ; The methyl group is on the N-terminal amino acid.

The peptide bond is boxed and the side chains are circled.

Appendix D Answers to Selected Questions and Problems

A-37 O C OH





















Leu-Ala O H H2N CH C N CH3 CH3 O CH C OH

17.19 CH3


O H H2N CH C N CH3 CH CH3 Gly-Ala



Ala-Gly O H H2N CH C CH3 N H O CH C OH H



17.23 17.25 17.27

Ser-Gly-Cys; Ser-Cys-Gly; Cys-Ser-Gly; Cys-Gly-Ser; Gly-Ser-Cys; Gly-Cys-Ser legumes, nuts, and seeds O O O H2N CH C H Gly OH H2N CH C CH3 Val OH H2N CH C CH2 OH Ser OH CH CH3


The abbreviated formula of the peptide is Gly-Ala-Val. Upon complete hydrolysis, the structures of the amino acids are O O O H2N CH C H Gly OH H2N CH C CH3 Ala OH H2N CH C CH3 Val OH CH CH3

17.31 O H2N CH C CH3 N





Appendix D Answers to Selected Questions and Problems


17.35 17.37 17.39 17.41 17.43 17.45 17.47 17.49 17.51 17.53 17.55 17.57 17.59 17.61

The primary structure is the amino acid sequence. The secondary structure results from intramolecular attractions such as hydrogen bonding and produces structural units such as the alpha helix and beta pleated sheets. The tertiary structure is the three-dimensional shape obtained from the folding of the peptide chain. Folding is caused by the intramolecular interactions of the secondary structural units. The quaternary structure is used to describe the shape of a protein when it is composed of more than one peptide chain. The quaternary structure is established through intermolecular interactions. In myoglobin, a globular protein, the protein chain folds back on itself. Wool, a fibrous protein, is primarily composed of chains of alpha helices. tertiary structure Hydrogen bonding holds together the twisted-helical shape of the peptide chain. Fibrous proteins have long rod-like shapes. In globular proteins, the peptide chain folds back onto itself giving the protein a "glob"-like shape. Denaturation disrupts the attractive forces which give rise to secondary, tertiary, and quaternary structures and causes the "unfolding" of the protein. Ribonucleic acids and deoxyribonucleic acids. In deoxyribonucleic acids, the ribose unit lacks the hydroxide group at the 2 position. A nucleic acid is composed of a chain of nucleotides. ribose; adenine; RNA The shapes and sizes of these pairs (A-T and C-G) complement each other and provide for the maximum degree of hydrogen bonding. AATCGGTCACGAT UGAUUCUAG Genetic information is encoded as complementary strands of nucleic acids. When genetic information is replicated, it copies one strand using complementary base pairing of DNA nucleotides. AAA AUG-CUU-CCC-CAA-CGA-UUU-CCC-CGA Met - Leu - Pro - Gln - Arg - Phe - Pro - Arg (so the answer is Pro and Arg)











Open chain

17.65 17.67 17.69 17.71 17.73

glucose, C6H12O6; sucrose, C12H22O11 structure A (glucose) monosaccharide Sucrose yields glucose and fructose. Starch yields only glucose. In starches, the main chain repeats with the glucose molecules maintaining the same orientation in the chains. Amylopectin also branches (which is not observed in cellulose).





Appendix D Answers to Selected Questions and Problems


In cellulose, the glucose units invert after every ether bond (circled). Note the position of the ­ CH2OH group.


17.75 17.77 17.79







Before we can taste the sugars in starch, the glucose monomers must be cleaved from the terminal of the amylose or amylopectin. The fatty acids in fat molecules tend to be completely or almost completely saturated (few or no C=C double bonds). The fatty acids in oil molecules are generally polyunsaturated and contain several C=C double bonds on each fatty acid chain. triacyl glycerol (tristearate) O H2C HC C O O C O O

17.81 17.83 17.85 17.87

H2C C O It would be a fat. phospholipid They are both lipids, and all steroids are derived from cholesterol. O O H O H H2N CH C CH2 OH N CH CH2 OH O H H2N CH C CH3 N CH CH2 SH CH CH3 O H C N CH CH2 C N CH CH2 OH O C OH C OH



17.93 17.95

OH Proteins perform many functions; one of the most important is to catalyze the biochemical reactions. They are also responsible for the machinery of the cell which regulates cellular processes such as transcription, translation, immune response, and energy production. Carbohydrates provide energy (i.e., glucose and metabolism) and play important structural roles in plants (cellulose). DNA and RNA each have a sugar molecule, base, and a phosphate unit. The sugar in DNA is deoxyribose. The DNA contains all of the genetic information. RNA, derived from DNA through transcription, contains only information for a specific protein. Genetic therapy can be used to produce novel substances from organisms. An example is human insulin production from bacteria. Plants can be bred to tolerate various chemical agents (i.e., herbicides) using this technique.


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