Download SOLUBILITY RULES FOR IONIC COMPOUNDS IN WATER

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 20.0C to
130.0C. 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 20C. Water vapor pressure at 20C 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