Unit 8
Chapter 20 and Final Review
Oxidation States and Redox
Reactions
Balancing Redox Equations
Voltaic Cells
Free Energy and Redox
Reactions
Cell EMF Under Standard
Conditions
Exam Review
Chapter 2
 Atoms
 Subatomic Particles
 Groups of the periodic table
 Moles
 Empirical and Molecular formulas
 Ionic Compounds
 Nomenclature
 The formula weight of the compound, Al2(SO4)3 18H2O
is:
 (a) 394.4 g
 (b) 666.4 g
 (c) 110,900 g
 (d) 466.8 g
 (e) 561.2 g
 The formula weight of the compound, Al2(SO4)3 18H2O
is:
 (a) 394.4 g
 (b) 666.4 g
 (c) 110,900 g
 (d) 466.8 g
 (e) 561.2 g
 A compound contains, by mass, 40.0% carbon, 6.71%
hydrogen, and 53.3% oxygen. A 0.320 mole sample of
this compound weighs 28.8 g. The molecular formula of
this compound is:
 (a) C2H4O2
 (b) C3H6O3
 (c) C2H4O
 (d) CH2O
 (e) C4H7O2
 Which one of the following salts is insoluble?
 (a) NH4Cl
 (b) Ca(NO3)2
 (c) BaCO3
 (d) Na2S
 (e) Zn(CH3COO)2
 Consider the species 72Zn, 75As and 74Ge. These
species have:
 (a) the same number of electrons.
 (b) the same number of protons.
 (c) the same number of neutrons.
 (d) the same number of protons and neutrons.
 (e) the same mass number.
 The species that contains 24 protons, 26 neutrons and
22 electrons would be represented by the symbol:
 (a) 50V3+
 (b) 26Cr2+
 (c) 50Cr2+
 (d) 50Mn2+
 (e) none of these
Chapter 3
 Balancing Equations
 Stoichiometry
 Types of reactions
 Balance the following equation with the smallest whole number
coefficients. Choose the answer that is the sum of the coefficients
in the balanced equation. Do not forget coefficients of "one."
 PtCl4 + XeF2  PtF6 + ClF + Xe
 (a) 16
 (b) 22
 (c) 24
 (d) 26
 (e) 32
 Which of the following statements is FALSE for the chemical equation given
below in which nitrogen gas reacts with hydrogen gas to form ammonia gas
assuming the reaction goes to completion?
 N2 + 3H2  2NH3
 (a) The reaction of one mole of H2 will produce 2/3 moles of NH3.
 (b) One mole of N2 will produce two moles of NH3.
 (c) One molecule of nitrogen requires three molecules of hydrogen for
complete reaction.
 (d) The reaction of 14 g of nitrogen produces 17 g of ammonia.
 (e) The reaction of three moles of hydrogen gas will produce 17 g of ammonia.
 Calculate the mass of hydrogen formed when 25 grams of
aluminum reacts with excess hydrochloric acid.
 2Al + 6HCl  Al2Cl6 + 3H2
 (a) 0.41 g
 (b) 1.2 g
 (c) 1.8 g
 (d) 2.8 g
 (e) 0.92 g
 If 5.0 g of each reactant were used for the the following
process, the limiting reactant would be:
 2KMnO4 +5Hg2Cl2 + 16HCl  10HgCl2 + 2MnCl2 + 2KCl +
8H2O
 (a) KMnO4
 (b) HCl
 (c) H2O
 (d) Hg2Cl2
 (e) HgCl2
Chapter 4
 Aqueous Reactions
 Solution Stoichiometry
 Strong Electrolytes
 Precipitation Reactions
 Acid Base Reactions
 Net Ionic Equations
 Redox Reactions
 Concentration
 Basic Titrations
 The precipitate formed when barium chloride is treated
with sulfuric acid is _______ .
 (a) BaS2O4
 (b) BaSO3
 (c) BaSO2
 (d) BaSO4
 (e) BaS
 Which one of the following salts is insoluble?
 (a) NH4Cl
 (b) Ca(NO3)2
 (c) BaCO3
 (d) Na2S
 (e) Zn(CH3COO)2
 The spectator ion(s) in the following reaction is/are:
 Na2CO3(aq) + Ba(NO3)2(aq)  BaCO3(s) +
2NaNO3(aq)
 (a) Na+ and Ba2+
 (b) Ba2+ and CO32 (c) CO32- and NO3-
 (d) Na+ only
 (e) Na+ and NO3-
 Which of the following statements is FALSE given the
following net ionic equation?
 2H+(aq) + Cu(OH)2(s)  Cu2+(aq) + 2H2O(l)
 (a) If all the water evaporated away, the salt remaining could
possibly be CuS.
 (b) The acid involved must be a strong electrolyte.
 (c) The base, Cu(OH)2, is an insoluble base.
 (d) This could be the net ionic equation for HNO3 reacting
with Cu(OH)2.
 (e) This is classified as a neutralization reaction.
 Determine the oxidation number of carbon in K2CO3.
 (a) 0
 (b) +2
 (c) +4
 (d) -2
 (e) some other value
 How many grams of Ca(OH)2 are contained in 1500 mL
of 0.0250 M Ca(OH)2 solution?
 (a) 3.17 g
 (b) 2.78 g
 (c) 1.85 g
 (d) 2.34 g
 (e) 4.25 g
 What volume of 12.6 M HCl must be added to enough
water to prepare 5.00 liters of 3.00 M HCl?
 (a) 1.19 L
 (b) 21.0 L
 (c) 0.840 L
 (d) 7.56 L
 (e) 2.14 L
 What is the molarity of the salt produced in the reaction
of 200 mL of 0.100 M HCl with 100 mL of 0.500 M
KOH?
 (a) 0.0325 M
 (b) 0.0472 M
 (c) 0.0667 M
 (d) 0.0864 M
 (e) 0.0935 M
 A 0.6745 gram sample of KHP reacts with 41.75 mL of KOH
solution for complete neutralization. What is the molarity of
the KOH solution? (Molecular weight of KHP = 204 g/mol.
KHP has one acidic hydrogen.)
 (a) 0.158 M
 (b) 0.099 M
 (c) 0.139 M
 (d) 0.079 M
 (e) 0.061 M
Chapters 5 and 19
 Energy
 Heat (q = mCΔT)
 Enthalpy
 Hess’s Law
 Spontaneous Process
 Entropy
 Gibbs Free Energy
 A system suffers an increase in internal energy of 80 J
and at the same time has 50 J of work done on it. What
is the heat change of the system?
 (a) +130 J
 (b) +30 J
 (c) -130 J
 (d) -30 J
 (e) 0 J
 For water (m.p. 0oC, b.p. 100oC)
 Heat of fusion = 333 J/g @ 0oC, Heat of vaporization = 2260 J/g @ 100oC,
Specific Heat (solid) = 2.09 J/goC, Specific Heat (liquid) = 4.18 J/goC, Specific
Heat (gas) = 2.03 J/goC
 Calculate the amount of heat (in kJ) that must be absorbed to convert 108 g of
ice at 0oC to water at 70oC.
 (a) 77
 (b) 68
 (c) 64
 (d) 57
 (e) 50
 A 5.000 g sample of methanol, CH3OH, was combusted in the
presence of excess oxygen in a bomb calorimeter conaining 4000
g of water. The temperature of the water increased from 24.000 oC
to 29.765 oC. The heat capacity of the calorimeter was 2657 J/oC.
The specific heat of water is 4.184 J/goC. Calculate ΔH for the
reaction in kJ/mol.
 (a) -314 kJ/mol
 (b) -789 kJ/mol
 (c) -716 kJ/mol
 (d) -121 kJ/mol
 (e) -69.5 kJ/mol
 Calculate Ho for the reaction:
 Na2O(s) + SO3(g) Na2SO4(g)
 given the following information:
 Na(s) + H2O(l)  NaOH(s) + ½ H2(g) ΔH = -146 kJ
 Na2SO4(s) + H2O(l)  2 NaOH(s) + SO3(g) ΔH = +418 kJ
 2 Na2O(s) + 2 H2(g)  4 Na(s) + 2 H2O(l) ΔH = +259 kJ
 (A) +255 kJ
 (B) -435 kJ
 (C) -581 kJ
 (D) +531 kJ
 (E) – 452 kJ
 The entropy will usually increase when
 I. a molecule is broken into two or more smaller molecules.
 II. a reaction occurs that results in an increase in the number
of moles of gas.
 III. a solid changes to a liquid.
 IV. a liquid changes to a gas.
 (a) I only
 (b) II only
 (c) III only
 (d) IV only
 (e) I, II, III, and IV
 For the following reaction at 25oC, ΔHo = +115 kJ and
ΔSo = +125 J/K. Calculate ΔGo for the reaction at 25o.
SBr4(g) S(g) + 2Br2(l)
 (a) +152 kJ
 (b) -56.7 kJ
 (c) +77.8 kJ
 (d) +37.1 kJ
 (e) -86.2 kJ
Chapter 6 and 7
 Light
 Electron Configuration
 Effective Nuclear Charge
 Size of atoms and ions
 Ionization Energy
 What is the frequency of light having a wavelength of
4.50 x 10-6 cm?
 (a) 2.84 x 10-12 s-1
 (b) 2.10 x 104 s-1
 (c) 4.29 x 1014 s-1
 (d) 1.06 x 1022 s-1
 (e) 6.67 x 1015 s-1
 The ground state electron configuration for arsenic is:
 (a) [Ar] 4s2 4p13
 (b) [Kr] 4s2 4p1
 (c) 1s2 2s2 2p6 3s2 3p6 3d12 4s2 4p1
 (d) 1s2 2s2 2p6 3s2 3p6 4s2 3d8 4p5
 (e) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
 Which element has the largest atomic radius?
 (a) Li
 (b) Na
 (c) Rb
 (d) F
 (e) I
 Which element has the lowest first ionization energy?
 (a) He
 (b) Ne
 (c) Ar
 (d) Kr
 (e) Xe
 Which of these isoelectronic species has the smallest
radius?
 (a) Br-
 (b) Sr2+
 (c) Rb+
 (d) Se2 (e) They are all the same size because they have the
same number of electrons.
Chapters 8 and 9
 Lewis Symbols
 Ionic Bonding
 Lewis Structures
 Bond Polarity and Electronegativity
 Formal Charge and Resonance
 Molecular shapes (VSEPR)
 Molecular Shapes and Polarity
 Which of the following pairs of elements and valence
electrons is incorrect?
 (a) Al - 3
 (b) Br - 7
 (c) S - 4
 (d) Sr - 2
 (e) Tl – 3
 Choose the molecule that is incorrectly matched with
the electronic geometry about the central atom.
 (a) CF4 - tetrahedral
 (b) BeBr2 - linear
 (c) H2O - tetrahedral
 (d) NH3 - tetrahedral
 (e) PF3 – pyramidal
 Which molecule has a linear arrangement of all
component atoms?
 (a) CH4
 (b) H2O
 (c) CO2
 (d) NH3
 (e) BF3
 Which of the following four molecules are polar: PH3
OF2 HF SO3?
 (a) all except SO3
 (b) only HF
 (c) only HF and OF2
 (d) none of these
 (e) all of these
Chapter 10
 Pressure
 Ideal Gas Law
 Gas Mixtures and Partial Pressure
 Kinetic Molecular Theory
 Effusion and Diffusion
 Real Gasses
 What pressure (in atm) would be exerted by 76 g of
fluorine gas in a 1.50 liter vessel at -37oC?
 (a) 26 atm
 (b) 4.1 atm
 (c) 19,600 atm
 (d) 84
 (e) 8.2 atm
 A mixture of 90.0 grams of CH4 and 10.0 grams of
argon has a pressure of 250 torr under conditions of
constant temperature and volume. The partial pressure
of CH4 in torr is:
 (a) 143
 (b) 100
 (c) 10.7
 (d) 239
 (e) 26.6
 If helium effuses through a porous barrier at a rate of
4.0 moles per minute, at what rate (in moles per
minute) would oxygen gas diffuse?
 (a) 0.20
 (b) 0.50
 (c) 2.0
 (d) 8.0
 (e) 1.41
 A real gas most closely approaches the behavior of an
ideal gas under conditions of:
 (a) high P and low T
 (b) low P and high T
 (c) low P and T
 (d) high P and T
 (e) STP
Chapter 11
 Intermolecular forces
 Polarizability
 Phase Changes
 Properties of Liquids
 What type of interparticle forces holds liquid N2
together?
 (a) ionic bonding
 (b) London forces
 (c) hydrogen bonding
 (d) dipole-dipole interaction
 (e) covalent bonding
 Which of the following boils at the highest
temperature?
 (a) CH4
 (b) C2H6
 (c) C3H8
 (d) C4H10
 (e) C5H12
 Which of the following changes would increase the vapor pressure of a liquid?
 1. an increase in temperature
 2. an increase in the intermolecular forces in the liquid
 3. an increase in the size of the open vessel containing the liquid
 (a) 1 and 2 only
 (b) 1 and 3 only
 (c) 1 only
 (d) 2 only
 (e) 3 only
Chapter 14
 Reaction Rates
 Rate Law
 Reaction Order
 Activation Energy
 Reaction Mechanism
 The combustion of ethane (C2H6) is represented by the
equation:
 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(l)
 In this reaction:
 (a) the rate of consumption of ethane is seven times faster
than the rate of consumption of oxygen.
 (b) the rate of formation of CO2 equals the rate of formation
of water.
 (c) water is formed at a rate equal to two-thirds the rate of
formation of CO2.
 (d) the rate of consumption of oxygen equals the rate of
consumption of water.
 (e) CO2 is formed twice as fast as ethane is consumed.
 What are the units of k for the rate law: Rate = k[A][B]2,
when the concentration unit is mol/L?
 (a) s-1
 (b) s
 (c) L mol-1 s-1
 (d) L2 mol-2 s-1
 (e) L2 s2 mol-2
 The half-life for a first-order reaction is 32 s. What was
the original concentration if, after 2.0 minutes, the
reactant concentration is 0.062 M?
 (a) 0.84 M
 (b) 0.069 M
 (c) 0.091 M
 (d) 0.075 M
 (e) 0.13 M
 A possible mechanism for the reaction
A + 2B  AB2
is as follows
Step 1: A + B  AB
Slow
Step 2: AB + B  AB2 Fast
The rate law expression for this reaction must be
(A) Rate = k[A]
(B) Rate = k[B]
(C) Rate = k[A][B]
(D) Rate = k[B]2
(E) Rate = k[A][B]2
Chapter 15
 Equilibrium
 Equilibrium constant expressions
 Heterogeneous equilibrium
 Reaction Quotient
 Le Chatelier’s Principle
 2SO3(g)  2SO2(g) + O2(g)
The conventional equilibrium constant expression (Kc)
for the system as described by the above equation is:
 (a) [SO2]2/[SO3]
 (b) [SO2]2[O2]/[SO3]2
 (c) [SO3]2/[SO3]2[O2]
 (d) [SO2][O2]
 (e) none of these
 Consider the following reversible reaction. In a 3.00 liter
container, the following amounts are found in equilibrium at
400 oC: 0.0420 mole N2, 0.516 mole H2 and 0.0357 mole
NH3. Evaluate Kc.
 N2(g) + 3H2(g)  2NH3(g)
 (a) 0.202
 (b) 1.99
 (c) 16.0
 (d) 4.94
 (e) 0.503
 At 445oC, Kc for the following reaction is 0.020.
 2HI(g)  H2(g) + I2(g)
 A mixture of H2, I2, and HI in a vessel at 445oC has the following
concentrations: [HI] = 2.0 M, [H2] = 0.50 M and [I2] = 0.10 M. Which one of the
following statements concerning the reaction quotient, Qc, is TRUE for the
above system?
 (a) Qc = Kc; the system is at equilibrium.
 (b) Qc is less than Kc; more H2 and I2 will be produced.
 (c) Qc is less than Kc; more HI will be produced.
 (d) Qc is greater than Kc; more H2 and I2 will be produced.
 (e) Qc is greater than Kc; more HI will be produced.
 Consider the gas-phase equilibrium system represented by the equation:
 2H2O(g)  2H2(g) + O2(g)
 Given that the forward reaction (the conversion of "left-hand" species to "righthand" species) is endothermic, which of the following changes will decrease
the equilibrium amount of H2O?
 (a) adding more oxygen
 (b) adding a solid phase calalyst
 (c) decreasing the volume of the container (the total pressure increases)
 (d) increasing the temperature at constant pressure
 (e) adding He gas
Chapter 16
 Acids and Bases
 Autoionization of Water
 pH and pOH
 Strong acids and bases
 Weak acids and bases
 In a sample of pure water, only one of the following
statements is always true at all conditions of
temperature and pressure. Which one is always true?
 (a) [H3O+] = 1.0 x 10-7 M
 (b) [OH-] = 1.0 x 10-7 M
 (c) pH = 7.0
 (d) pOH = 7.0
 (e) [H3O+] = [OH-]
 The pOH of a solution of NaOH is 11.30. What is the
[H+] for this solution?
 (a) 2.0 x 10-3
 (b) 2.5 x 10-3
 (c) 5.0 x 10-12
 (d) 4.0 x 10-12
 (e) 6.2 x 10-8
 A 0.10 M solution of a weak acid, HX, is 0.059%
ionized. Evaluate Ka for the acid.
 (a) 3.8 x 10-9
 (b) 6.5 x 10-7
 (c) 7.0 x 10-6
 (d) 4.2 x 10-6
 (e) 3.5 x 10-8
Chapter 17
 Common Ion Effect
 Buffers
 Titrations
 Solubility
 What is the pH of a solution composed of 0.20 M NH3
and 0.15 M NH4Cl?
 (a) 2.15
 (b) 4.62
 (c) 8.26
 (d) 9.38
 (e) 8.89
 Consider a solution which is 0.10 M in CH3COOH and 0.20 M in
NaCH3COO. Which of the following statements is true?
 (a) If a small amount of NaOH is added, the pH decreases very
slightly.
 (b) If NaOH is added, the OH- ions react with the CH3COO- ions.
 (c) If a small amount of HCl is added, the pH decreases very
slightly.
 (d) If HCl is added, the H+ ions react with CH3COOH ions.
 (e) If more CH3COOH is added, the pH increases.
 What is the pH at the equivalence point in the titration
of 100.0 mL of 0.20 M ammonia with 0.10 M
hydrochloric acid?
 (a) 4.6
 (b) 5.2
 (c) 7.0
 (d) 5.5
 (e) 4.9
 The solubility product expression for tin(II) hydroxide,
Sn(OH)2, is
 (a) [Sn2+][OH-]
 (b) [Sn2+]2[OH-]
 (c) [Sn2+][OH-]2
 (d) [Sn2+]3[OH-]
 (e) [Sn2+][OH-]3
 The molar solubility of PbBr2 is 2.17 x 10-3 M at a
certain temperature. Calculate Ksp for PbBr2.
 (a) 6.2 x 10-6
 (b) 6.4 x 10-7
 (c) 4.1 x 10-8
 (d) 3.4 x 10-6
 (e) 1.4 x 10-5