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OFFICE HOURS FOR FINALS WEEK Monday Dec 15 11:30 am - 12:30 pm 12:30 pm - 3:15 pm SM 255 SM 206 Tuesday Dec 16 8:15 am - 12:00 pm SM 206 TEST 1 REVIEW 1. For each of the following subatomic particles, give their (1) location in the atom, (2) charge, and (3) relative mass (a) proton (b) neutron (c) electron 2. Define atomic number and mass number. 3. Determine the number of protons, neutrons and electrons in a (a) silver-110 atom (b) silver-110 ion 4. Determine the number of atoms in a 2.50 g sample of beryllium metal. 5. Give the electron dot notation for each atom in the third period of the periodic table. 6. Give the electron configuration notation (a) for Pt and (b) for Pt2+. 7. Explain the atomic radii trend in a group and in a period. 8. Define ionization energy and write the equation for the ionization of a sodium atom. 9. Explain the ionization energy trend in a group and in a period. Explain why Groups 13 and 16 are exceptions to the trend in a period. 10. Explain the trend for successive ionization energies for an atom 11. Define electron affinity and write the equation for the addition of an electron to a sulfur atom. 12. Explain the electron affinity trend in a group and in a period. Explain why Groups 2, 7, 12, 15, and 18 are exceptions to the trend in a period. 13. Give (a) the expected ionic charges for each of the representative elements in the fourth period of the periodic table, and (b) the expected ionic charges for silver, zinc, tin, lead, iron, copper, and mercury. 14. Explain the trend for metallic activity. Give physical and chemical properties of metals. 15. Explain the trend for nonmetal activity. Give the physical and chemical properties of nonmetals. 16. Define molar mass. TEST 2 REVIEW 17. Define ionic bond, covalent bond and metallic bond. 18. Name the following ionic compounds: CaBr2, Cr2S3, Cs3PO4, Co(NO2)2, Cu2CO3.3H2O 19. Name the following binary covalent compounds: NO, NO2, N2O, N2O4, N2O5 20. Name the following acids: HI, HIO, HIO2, HIO3, HIO4 21. Draw a correct Lewis structure for each ionic compound: KF, Li2O, Mg(OH)2 22. Draw a correct Lewis structure for each covalent compound: NF3, HCP, O3 23. Explain what is meant when a molecule is said to possess resonance, and tell how the real electronic structure of the molecule is related to the molecule's resonance structures. 24. Explain the principle behind the VSEPR theory. 25. Give the shape for a molecule with each of the following conditions: (a) SN = 2 (b) SN = 3, 0 lone pairs (c) SN = 3, 1 lone pair (d) SN = 4, 0 lone pairs (e) SN = 4, 1 lone pair (f) SN = 4, 2 lone pairs (g) SN = 4, 3 lone pairs (h) SN = 5, 0 lone pairs (i) SN = 5, 1 lone pair (j) SN = 5, 2 lone pairs (k) SN = 5, 3 lone pairs (l) SN = 6, 0 lone pairs (m) SN = 6, 1 lone pair (n) SN = 6, 2 lone pairs 26. Explain how a molecule is determined to be polar or nonpolar. 27. Explain how the hybridization of an atom is determined. 28. Define (a) a sigma bond and (b) a pi bond. 29. What type of bonds (sigma or pi) constitute a (a) single bond (b) double bond (c) triple bond 30. Draw a bond orbital model for each of the following, labeling each of the bonds: (a) C2H6 (b) C2H4 (c) C2H2 31. What is a delocalized pi system, and what type of molecules possess a delocalized pi system? 32. Draw the Lewis structure and bond orbital model for PS2-, labeling each of the bonds. (continued on next page) 33. Draw a molecular orbital electron energy level diagram for the F2+ ion. 34. Based upon the molecular orbital energy level diagram in problem 33, give the following for the (1) F2 molecule, (2) the F2+ ion, and (3) the F22+ ion. (a) valence electron configuration notation (b) bond order (c) magnetism (paramagnetic or diamagnetic) 35. Rank the three diatomic species, the F2 molecule, the F2+ ion, and the F22+, from highest to lowest for the following properties (a) bond energy (b) bond length 36. Determine the empirical formula of a compound that is found to be 31.5% calcium, 24.4% phosphorus, and 44.1% oxygen by mass. 37. Determine the molecular formula of a compound that is found to be 85.6% carbon and 14.4% hydrogen by mass, and has a molar mass of between 80 and 90 grams. TEST 3 REVIEW 38. For nonpolar molecular substances, give (a) three examples, (b) the major attractive force between the particles, (c) high or low melting points, (d) electrical conductivity as a solid and as a liquid, (e) water solubility, and (f) how the relative melting points of two or more members of this category can be determined. 39. For polar molecular substances, give (a) three examples, (b) the major attractive force between the particles, (c) high or low melting points, (d) electrical conductivity as a solid and as a liquid, (e) water solubility, and (f) how the relative melting points of two or more members of this category can be determined. 40. For ionic substances, give (a) three examples, (b) the major attractive force between the particles, (c) high or low melting points, (d) electrical conductivity as a solid and as a liquid, (e) water solubility, and (f) how the relative melting points of two or more members of this category can be determined. 41. For metallic substances, give (a) three examples, (b) the major attractive force between the particles, (c) high or low melting points, (d) electrical conductivity as a solid and as a liquid, (e) water solubility, and (f) how the relative melting points of two or more members of this category can be determined. 42. For macromolecular substances, give (a) three examples, (b) the major attractive force between the particles, (c) high or low melting points, (d) electrical conductivity as a solid and as a liquid, (e) water solubility, and (f) how the relative melting points of two or more members of this category can be determined. (continued on next page) 43. Calculate the amount of heat it would take to convert 50.0 g of naphthalene (C10H8) at 20.0C to 130.0C. The specific heat capacity for solid naphthalene is 1.51 J/gC, the specific heat capacity for liquid naphthalene is 2.48 J/gC, and the specific heat capacity for gaseous naphthalene is 1.63 J/gC. The heat of fusion for naphthalene is 136 J/g and the heat of vaporization for naphthalene is 1,167 J/g. Naphthalene has a normal melting point of 80.0ºC and a normal boiling point of 218.0ºC 44. Define (a) the boiling point and (b) the freezing point of a substance. 45. Explain the rule likes dissolve likes. 46. What is (a) a strong electrolyte, (b) a weak electrolyte, and (c) a nonelectrolyte? Give two examples of each. 47. Why do electrolyte solutions conduct electricity? 48. Define the concentration units (a) mass percent, (b) molarity, (c) molality, and (d) mole fraction. 49. Explain why nonvolatile solutes lower a solvent's vapor pressure. 50. Explain why nonvolatile solutes raise a solvent's boiling point. 51. Explain why nonvolatile solutes lower a solvent's freezing point. 52. The cooling system of a car is filled with a solution formed by mixing equal volumes of water (density = 1.00 g/mL) and ethylene glycol, C2H6O2 (density = 1.12 g/mL). Calculate (a) the freezing point and (b) the boiling point of this aqueous solution. 53. Safrole is contained in oil of sassafras, and was once used to flavor root beer. A 2.39 g sample of safrole was dissolved in 103.0 g of diphenyl ether. The solution had a freezing point of 25.70ºC. Calculate the molar mass of safrole. The freezing point of pure diphenyl ether is 26.84ºC, and its molal freezing point constant is 8.00 Cºkg/mol. 54. A 0.50 m aqueous solution of a monoprotic acid has a boiling point of 100.36ºC. Determine if the monoprotic acid a strong acid or a weak acid. TEST 4 REVIEW 55. List the five solubility rules to determine if an ionic compound is soluble in water. 56. Give the products for the decomposition of metallic chlorates. 57. When will (a) an elemental metal replace a metal ion from a compound in solution, (b) an elemental metal replace a hydrogen ion from an acid, and (c) an elemental metal replace a hydrogen ion from water? (continued on next page) 58. When will an elemental nonmetal replace a nonmetal ion from a compound in solution? 59. When will an exchange reaction in solution occur? 60. What happens during a Bronsted acid-base reaction? 61. Use appropriate ionic and molecular formulas to show the reactants and products for the following, each of which results in a reaction occurring. Balance all equations. For reactions in aqueous solution, give the balanced equation in net ionic form. (a) Crystals of strontium chlorate are heated strongly. (b) Magnesium metal is added to acetic acid. (c) Aluminum metal is added to copper (II) chloride solution. (d) Fluorine gas is added to sodium chloride solution (e) Solutions of barium chloride and sodium sulfate are mixed. (f) Solutions of lead (II) nitrate and hydrochloric acid are mixed. (g) Sulfuric acid is added to calcium carbonate. (h) Liquid benzene (C6H6) is burned in air. (i) Solutions of sodium hydroxide and potassium bisulfate are mixed. (j) Potassium permanganate reacts with sodium bromide in an acidic solution. (k) Potassium permanganate reacts with iron (II) nitrate in a basic solution. (l) Sodium dichromate reacts with tin (II) sulfate in an acidic solution. (m) Copper metal is added to concentrated nitric acid. 62. Phosphorus and chlorine molecules, shown in the reaction chamber to the right, react to form phosphorus trichloride. Draw a picture of the chamber after the reaction. 63. Determine the mass of copper metal that can be produced from the reaction of a copper (II) chloride solution with 2.50 grams of aluminum metal. 64. Determine the mass of calcium phosphate that can be produced from the reaction of 3.75 grams of calcium chloride with 4.25 grams of potassium phosphate. 65. A 0.440 gram sample known to contain chloride ions was dissolved in 50 mL of water. The solution was treated with excess silver nitrate, a precipitate was collected, washed, dried, and weighed. If the mass of the precipitate was 0.925 grams, determine the percentage of chloride ions in the sample. 66. What is a standard solution? 67. A lithium hydroxide solution was to be standardized with the solid acid, potassium biphthalate, KHC8H4O4. If 22.85 mL of the lithium hydroxide solution were required to neutralize 0.8725 grams of the solid potassium biphthalate, calculate the molarity of the lithium hydroxide solution. 68. A 25.00 mL sample of a sulfuric acid solution was titrated with the lithium hydroxide solution from question 67. If 37.10 mL of the lithium hydroxide solution were required to reach neutralization, calculate the molarity of the sulfuric acid solution. (continued on next page) 69. Lime-Away™ contains phosphoric acid. A 10.00 mL sample of Lime-Away with a mass of 10.260 grams was titrated with the lithium hydroxide solution from question 67. If 47.75 mL of the lithium hydroxide solution were required to reach neutralization, calculate the mass percent of phosphoric acid in Lime- Away™. 70. A 0.3503 gram sample of a solid diprotic acid was dissolved in 50 mL of water and titrated with the lithium hydroxide solution from question 67. If 26.60 mL of the lithium hydroxide solution were required to reach neutralization, calculate the molar mass of the diprotic acid. TEST 5 REVIEW 71. Give the pressure of dry oxygen gas is each of the following gas collecting devices. Assume atmospheric pressure is 753 torr and the temperature is 20C. Water vapor pressure at 20C is 18 torr. 72. A sample of air with a pressure of 1.000 atm is liquefied and found to consist of 33.66 g nitrogen, 10.35 g oxygen, and 0.59 g argon. Find the partial pressures of each gas in the air. 73. Reynolds Wrap™ was analyzed for its aluminum content. A 0.0240 gram sample of Reynolds Wrap™ was reacted with excess hydrochloric acid, and 33.7 mL of hydrogen gas was collected at a temperature of 22ºC and a pressure of 726 torr. Determine the mass percent of aluminum in Reynolds Wrap™. 74. Calculate the root-mean-square-velocity of carbon dioxide gas molecules at 27ºC. 75. Carbon dioxide molecules effuse at a rate of 27.5 mL/min, while an unknown gas effuses at a rate of 32.2 mL/min under the same conditions. Calculate the molar mass of the unknown gas. (continued on next page) 76. The United States produces more than 7 billion kilograms of vinyl chloride annually. Most is converted into polymer polyvinyl chloride (PVC), which is used to make piping, siding, gutters, floor tiles, clothing, and toys. Vinyl chloride is made from a two-step process, in which a chlorine atom replaces a hydrogen atom in ethane according to the balanced equation: 2C2H4(g) + 2HCl(g) + O2(g) 2C2H3Cl(s) + 2H2O(g) Using the bond energies in your text book, calculate the energy change for the above reaction. 77. State the First Law of Thermodynamics in words and mathematically. 78. What is the sign convention for heat and work? 79. A 1.365 g sample of sodium carbonate was added to 25.0 mL of a 1.00 M hydrochloric acid solution in a calorimeter, and the solution’s temperature increased from 21.2ºC to 32.6ºC. If the solution had a mass of 26.834 g and a heat capacity of 4.184 J/gCº, calculate the heat of reaction in kilojoules per mole of sodium chloride. 80. In some liquid fuel rockets, such as the lunar lander module of the Apollo moon missions, the fuels are liquid hydrazine (N2H4) and dinitrogen tetroxide (N2O4). The two chemicals ignite on contact to release very large amounts of energy by the following reaction: 2N2H4(l) + N2O4(l) 3N2(g) + 4H2O(g) Using the standard heat of formations in your text book, calculate the standard enthalpy change for the above reaction in kilojoules per gram of hydrazine. 81. “Strike anywhere” matches contain tetraphosphorus trisulfide, a compound that ignites when heated by friction. It vigorously reacts with oxygen as follows: P4S3(s) + 8O2(g) P4O10(s) + 3SO2(g) H = -3677 kJ Using the standard heat of formations in your text book, calculate the standard heat of formation of solid tetraphosphorus trisulfide. ANSWERS 1. (a) nucleus positive (b) nucleus neutral (c) around the nucleus negative 1 1 1/1836 2. atomic number – the number of protons in an atom mass number – the sum of the protons and neutrons in an atom 3. (a) 47 protons, 63 neutrons, 47 electrons (b) 47 protons, 63 neutrons, 46 electrons 4. 1.67 x 1023 Be atoms 5. Na · Mg : . Al : (continued on next page) . · Si : . ·P: · .. ·S: · .. : Cl : · .. : Ar : ·· 6. (a) [Xe]6s24f145d8 (b) [Xe]4f145d8 7. group – radii increase going down due to increasing number of energy levels, which are shielding the increasing nuclear charge period – radii decrease moving right due to increasing nuclear charge, while the number of shielding energy levels remain the same – 8. energy required to remove an electron from a gaseous atom; Na (g) → Na+ (g) + e 9. group – ionization energy decreases going down due to increasing number of energy levels, which are shielding the increasing nuclear charge, resulting in less attraction for the valence electrons period – ionization energy increases moving right due to increasing nuclear charge, while the number of shielding energy levels remain the same, resulting in greater attraction for the valence electrons in a period ionization energies are lower than expected when the electron is the first one removed from a new, more shielded sublevel where the electron experiences less attraction by the nucleus, or, when the electron is the first paired electron removed from a multiple-orbital sublevel, where it experiences electron-electron repulsion 10. successive ionization energies increase due to a decrease in shielding with the removal of each electron, resulting in greater attraction for the remaining electrons a large increase in ionization energy is seen when an electron is removed from a new, inner energy level, which is less shielded from the nuclear charge – – 11. energy change when a gaseous atom gains an electron; S (g) + e → S (g) 12. group – electron affinities become less exothermic going down due to increasing number of energy levels, which are shielding the increasing nuclear charge, resulting in less attraction for a free electron period – electron affinities become more exothermic moving right due to increasing nuclear charge, while the number of shielding energy levels remain the same, resulting in greater attraction for a free electron electron affinities are less exothermic than expected when the electron is the first one added to a new sublevel, where the electron experiences less attraction to the nucleus because the shape of the new orbital does not allow the electron as much probability close to the nucleus inside the shielding, or, when the electron is the first to be paired in a multiple-orbital sublevel, where it experiences electronelectron repulsion – 13. (a) K+; Ca2+; Ga3+; Ge4+ and Ge2+; As3–; Se2–; Br (b) Ag+; Zn2+; Sn4+ and Sn2+; Pb4+ and Pb2+; Fe3+ and Fe2+; Cu2+ and Cu+; Hg2+ and Hg22+ 14. metals are very active at the left of the Periodic Table, and activity decreases moving right to the Coinage Metals; past the Coinage Metals metals are active again, and then activity decreases moving downand to the right physical: luster, malleable, ductile, conductors of heat and electricity; chemical: lose electrons to form positive ions (continued on next page) 15. except for the noble Gases which are essentially inert, nonmetals are very active at the top right of the Periodic Table, and activity decreases moving down and to the left physical: opposite of metals; chemical: gain electrons to form negative ions 16. the mass of one mole of atoms of an elemental substance, the mass of one mole of molecules of a molecular substance, or the mass of one mole of formula units of an ionic substance 17. ionic bond – attraction between positive and negative ions; formed between metals and nonmetals covalent bond – attraction between nuclei and the electrons they share; formed between nonmetals metallic bond – delocalized covalent bond; formed between metals 18. calcium bromide, chromium (III) sulfide, cesium phosphate, cobalt (II) nitrite, copper (I) carbonate trihydrate 19. nitrogen monoxide, nitrogen dioxide, dinitrogen monoxide, dinitrogen tetroxide, dinitrogen pentoxide 20. hydroiodic acid, hypoiodous acid, iodous acid, iodic acid, periodic acid 21. (a) (b) (c) 22. (a) (b) (c) 23. a condition when more than one valid Lewis structure can be drawn to represent the bonding in the molecule; real electronic structure is the average of the resonance structures 24. for each atom in a molecule, all electron pairs in the the outer shell of the atom (bonded atoms and lone pairs) will spread themselves out as far apart as possible to minimize repulsion 25. Handout 2D 26. determine if the bonds in the molecule are polar or nonpolar by subtracting the electronegativities of the bonding atoms; if the bonds are nonpolar, the molecule is nonpolar; if the bonds are polar, determine the geometry of the polar bonds in the molecule; if the polar bonds are arranged symmetrically, their dipole moments will cancel out and the molecule is nonpolar, if the polar bonds are arranged asymmetrically, their dipole moments will not cancel out and the molecule is polar 27. determine the steric number of the atom and that will be the number of atomic orbitals that must be hybridized: SN = 2, sp; SN = 3, sp2; SN = 4, sp3; SN = 5, sp3d; SN = 6, sp3d2 28. (a) a bond that is completely symmetrical around the internuclear axis (b) a bond that is only symmetrical upon a 180º rotation around the internuclear axis (continued on next page) 29. (a) 1 σ (b) 1 σ and 1 π (c) 1 σ and 2 π 30. (a) (b) (c) a = σ(1s + sp3) b = σ(sp3 + sp3) a = σ(sp2 + sp2) b = π(2p + 2p) c = σ(1s + sp2) a = σ(sp + sp) b = π(2p + 2p) c = σ(1s + sp) 31. Pi bonds in which the electrons are not localized between two atoms, but are shared by more than two atoms; molecules possessing resonance, with parallel p orbitals on neighboring atoms 32. 33. 34. (a) F2 (σ2sb)2(σ2s*)2(π2pb)4(σ2pb)2(π2p*)4 F2+ (σ2sb)2(σ2s*)2(π2pb)4(σ2pb)2(π2p*)3 F22+ (σ2sb)2(σ2s*)2(π2pb)4(σ2pb)2(π2p*)2 (b) F2 = 1, F2+ = 1.5, F22+ = 2 (c) F2 = diamagnetic, F2+ = paramagnetic, F22+ = paramagnetic 35. (a) F2 < F2+ < F22+ (b) F22+ < F2+ < F2 36. Ca2P2O7 37. C6H12 38. (a) N2, C3H8, BF3 (b) London dispersion forces (c) low (d) no conduction as a solid or a liquid (e) low (f) (1) more electrons each molecule has, the more they polarize each other’s electron clouds, the stronger the London dispersion forces, the higher the MP; (2) the greater number of close atoms in neighboring molecules, the more they polarize each other’s electron clouds, the stronger the London dispersion forces, the higher the MP (continued on next page) 39. (a) HBr, SO2, CHF3 (b) London dispersion forces, dipole-dipole interactions, and in some cases hydrogen bonding (c) low (d) no conduction as solid or liquid (e) high (f) (1) more electrons each molecule has, the more they polarize each other’s electron clouds, the stronger the London dispersion forces, the higher the MP; (2) the more area of contact between molecules, the more they polarize each other’s electron clouds, the stronger the London dispersion forces, the higher the MP 40. (a) MgO, Fe(NO3)2, (NH4)2S (b) ionic bonds (c) high (d) no conduction as solid, but conducts as a liquid (e) high (f) (1) the higher the charges of the ions, the stronger the ionic bond, the higher the MP; (2) the smaller the ions, the closer they get, the stronger the ionic bond, the higher the MP 41. (a) Mg, Fe, brass (b) metallic bonds (c) high (d) conducts as a solid or a liquid (e) low (f) (1) the more net bonding electrons, the greater the attraction of the bonding electrons to the nuclei of the bonding atoms, the stronger the metallic bond, the higher the MP; (2) the smaller the atoms, the smaller the bonding molecular orbitals, the closer the bonding electrons are to the nuclei of the bonding atoms, the stronger the metallic bond, the higher the MP 42. (a) C, Si, SiO2 (b) covalent bonds (c) high (d) no conduction as a solid or a liquid (e) low (f) (1) the more net bonding electrons, the greater the attraction of the bonding electrons to the nuclei of the bonding atoms, the stronger the covalent bond, the higher the MP; (2) the smaller the atoms, the smaller the bonding molecular orbitals, the closer the bonding electrons are to the nuclei of the bonding atoms, the stronger the covalent bond, the higher the MP 43. 17,530 J 44. (a) the temperature at which the equilibrium vapor pressure of the liquid equals prevailing atmospheric phessure; (b) the temperature at which the equilibrium vapor pressure of the solid equals the equilibrium vapor pressure of the liquid 45. substance with the same polarity will be soluble, substances with different polarities will be insoluble 46. (a) a compound that completely ionizes or dissociates in water solution, producing many ions in solution; sodium chloride, hydrochloric acid; (b) a compound that partially ionizes or dissociates in water solution, producing few ions in solution; hydrofluoric acid, carbonic acid; (c) a compound that does not ionize or dissociate in water solution, producing no ions in solution; glucose, ethanol 47. electrolytes produce mobile ions in solution, and the movement of charged particles is electrical conduction 48. (a) (b) (c) (d) grams of solute per grams of solution moles of solute per liter of solution moles of solute per kilogram of solvent moles of one component per total moles of solution 49. a pure solvent’s equilibrium vapor pressure is determined by the number of solvent molecules in the vapor phase that are in equilibrium with the liquid through equal rates of evaporation and condensation; nonvolatile solute particles block solvent molecules on the surface of a solution from evaporating, decreasing the rate of evaporation of the solvent; because of the unequal rates of evaporation and condensation, some of the vapor molecules condense into liquid, reducing the rate of condensation so that equilibrium is again established, but now, with fewer vapor molecules, the equilibrium vapor pressure will be lower (continued on next page) 50. a pure solvent boils when its equilibrium vapor pressure equals prevailing atmospheric pressure; nonvolatile solute particles added to a boiling solvent cause the equilibrium vapor pressure of the solution to be decreased, so the equilibrium vapor pressure is now less than atmospheric pressure and the solution stops boling; to boil the equilibrium vapor pressure of the solution must be increased to equal the prevailing atmospheric pressure by raising the temperature, therefore the boiling point of the solution will now higher than the boiling point of the pure solvent 51. a pure solvent freezes when the liquid solvent’s equilibrium vapor pressure equals the solid solvent’s equilibrium vapor pressure; nonvolatile solute particles added to a freezing solvent cause the equilibrium vapor pressure of the liquid to be decreased, so the liquid’s equilibrium vapor pressure is now less than the solid solvent’s equilibrium vapor pressure; to freeze the temperature of the solution must be lowered in order for the equilibrium vapor pressures of the liquid and the equilibrium vapor pressure of the solid to be equal again, therefore the freezing point of the solution will now be lower than the freezing point of the pure solvent 52. (a) -33.5ºC (b) 109.2ºC 53. 163 g/mol 54. weak 55. Handout 4A 56. metallic chloride and oxygen 57. (a) when the elemental metal is more active than the metal ion (b) when the elemental metal is more active than hydrogen (c) when the elemental metal is very active 58. when the elemental nonmetal is more active than the nonmetal ion 59. when two solutions are mixed and an insoluble product or water is formed 60. a hydrogen ion is transferred from an acid to a base 61. (a) Sr(ClO3)2 (s) SrCl2 (s) + 3O2 (g) – (b) Mg (s) + 2HC2H3O2 (aq) Mg2+ (aq) + 2C2H3O2 (aq) + H2 (g) (c) 2Al (s) + 3Cu2+ (aq) 2Al3+ (aq) + 3Cu (s) – – (d) F2 (g) + 2Cl (aq) 2F (aq) + Cl2 (g) (e) Ba2+ (aq) + SO42- (aq) BaSO4 (s) – (f) Pb2+ (aq) + 2Cl (aq) PbCl2 (s) (g) 2H+ (aq) + SO42- (aq) + CaCO3 (s) CaSO4 (s) + H2O (l) + CO2 (g) (h) 2C6H6 (l) + 15O2 (g) 12CO2 (g) + 6H2O (g) – – (i) OH (aq) + HSO4 (aq) H2O (l) + SO42- (aq) – – (j) 2MnO4 (aq) + 10Br (aq) + 16H+ (aq) 2Mn2+ (aq) + 5Br2 (l) + 8H2O (l) – – (k) 5OH (aq) + MnO4 (aq) + 3Fe2+ (aq) + 2H2O (l) MnO2 (s) + 3Fe(OH)3 (s) (l) Cr2O72- (aq) + 3Sn2+ (aq) + 14H+ (aq) 2Cr3+ (aq) + 3Sn4+ (aq) + 7H2O (l) – (m) 4H+ (aq) + Cu (s) + 2NO3 (aq) Cu2+ (aq) + 2NO2 (g) + 2H2O (l), or – 2H+ (aq) + Cu (s) + NO3 (aq) Cu+ (aq) + NO2 (g) + H2O (l) (continued on next page) 62. 63. 8.83 g 64. 3.11 g 65. 52.0% 66. a solution of known concentration 67. 0.1870 M 68. 0.1387 M 69. 2.842% 70. 140.9 g/mol 71. (a) 121 torr (b) 874 torr (c) 735 torr (d) 729 torr 72. 0.7803 atm N2, 0.2101 atm O2, 0.0096 atmAr 73. 99.6% 74. 412 m/s 75. 32.1 g/mol 76. -371 kJ 77. the energy of the universe is constant, or, energy is neither created nor destroyed; ΔE = q + w 78. positive – heat transferred to the system or work done on the system; negative – heat transferred to the surroundings or work done on the surroundings 79. -51.2 kJ/mol 80. -16.4 kJ/g 81. -198 kJ/mol