General Chemistry II Exam Review - 1830 Verified Questions

General Chemistry II Exam Review

Course Introduction

General Chemistry II is a continuation of foundational chemical principles, focusing on topics such as chemical kinetics, chemical equilibrium, acids and bases, thermodynamics, electrochemistry, and an introduction to coordination chemistry. This course provides in-depth exploration of molecular interactions, reaction mechanisms, solution chemistry, and the quantitative analysis of chemical systems. Emphasis is placed on developing problem-solving skills and understanding the applications of chemical principles in real-world and laboratory contexts, setting the stage for advanced studies in chemistry and related scientific disciplines.

Recommended Textbook

Principles of General Chemistry 2nd Edition by Martin Silberberg

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23 Chapters

1830 Verified Questions

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Page 2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) When applying the scientific method,it is important to avoid any form of hypothesis.

A)True

B)False

Answer: False

Q2) The density of mercury,the only metal to exist as a liquid at room temperature,is 13.6 g/cm³.What is that density in pounds per cubic inch?

A)849 lb/in³

B)491 lb/in³

C)376 lb/in³

D)0.491 lb/in³

E)1.83 \(\times\) 10^-3 lb/in³

Answer: D

Q3) Which measurement is expressed to 4 significant figures?

A)0.00423 kg

B)24.049 cm

C)1300 K

D)82,306 m

E)62.40 g

Answer: E

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Page 3

Chapter 2: The Components of Matter

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Sample Questions

Q1) Modern studies have shown that the Law of Multiple Proportions is not valid.

A)True

B)False

Answer: False

Q2) What is the name of the acid formed when HBr gas is dissolved in water?

A)bromic acid

B)bromous acid

C)hydrobromic acid

D)hydrobromous acid

E)hydrobromidic acid

Answer: C

Q3) Which one of the following combinations of names and formulas of ions is incorrect?

A)NH₄⁺ ammonium

B)S^2- sulfide

C)CN^- cyanide

D)S₂O₃^2- thiosulfate

E)ClO₃^- perchlorate

Answer: E

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Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) A compound of bromine and fluorine is used to make UF₆,which is an important chemical in processing and reprocessing of nuclear fuel.The compound contains 58.37 mass percent bromine.Determine its empirical formula.

A)BrF

B)BrF₂

C)Br₂F₃

D)Br₃F

E)BrF₃

Answer: E

Q2) a.A solution of common salt,NaCl,in water has a concentration of 0.0921 M.Calculate the number of moles of HCl contained in 50.0 mL of this solution.

b.If,instead,an NaCl solution is prepared by dissolving 10.0 g of solid NaCl in enough water to make 250.mL of solution,what is the molarity?

Answer: a.0.00461 g

b.0.684 M

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Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) Calculate the oxidation number of the chlorine in perchloric acid,HClO₄,a strong oxidizing agent.

A)-1

B)+4

C)+5

D)+7

E)None of these is the correct oxidation number.

Q2) A particular reaction may be both a precipitation and an acid-base (neutralization)reaction.

A)True

B)False

Q3) In an acid-base (neutralization)reaction the equivalence point is the point where the indicator changes color.

A)True

B)False

Q4) Complete and balance the equation for the following acid-base reaction. Ca(OH)₂(aq)+ HCl(aq) \(\to\)

Q5) All combustion reactions are classified as combination reactions.

A)True

B)False

6

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) Magnesium metal (0.100 mol)and a volume of aqueous hydrochloric acid that contains 0.500 mol of HCl are combined and react to completion.How many liters of hydrogen gas,measured at STP,are produced? Mg(s)+ 2HCl(aq) \(\to\) MgCl₂(aq)+ H₂(g)

A)2.24 L of H₂

B)4.48 L of H₂

C)5.60 L of H₂

D)11.2 L of H₂

E)22.4 L of H₂

Q2) State Boyle's Law and illustrate it with a graph,using standard x-y coordinate axes.Be sure to label the axes unambiguously with the correct gas variables.

Q3) A 500-mL sample of argon at 800 torr has its absolute temperature quadrupled.If the volume remains unchanged,what is the new pressure?

A)200 torr

B)400 torr

C)800 torr

D)2400 torr

E)3200 torr

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7

Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) An important step in the synthesis of nitric acid is the conversion of ammonia to nitric oxide.\(\Delta\)H°_f [NH₃(g)] = -45.9 kJ/mol;\(\Delta\)H°_f [NO(g)] = 90.3 kJ/mol;\(\Delta\)H°_f [H₂O(g)] = -241.8 kJ/mol 4NH₃(g)+ 5O₂(g) \(\to\) 4NO(g)+ 6H₂O(g)

Calculate \(\Delta\)H°_rxn for this reaction.

A)-906.0 kJ

B)-197.4 kJ

C)-105.6 kJ

D)197.4 kJ

E)906.0 kJ

Q2) A system that does no work but which receives heat from the surroundings has

A)q < 0,\(\Delta\)E > 0

B)q > 0,\(\Delta\)E < 0

C)q = \(\Delta\)E

D)q = -\(\Delta\)E

E)w = \(\Delta\)E

Q3) Consider the equation \(\Delta\)E = q + w

Explain fully the meaning of all three terms in the equation,and also the implied sign convention for q and w.

Page 8

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) Use the Rydberg equation to calculate the frequency of a photon absorbed when the hydrogen atom undergoes a transition from n₁ = 2 to n₂ = 4.(R = 1.096776 \(\times\) 10⁷ m^-1)

A)2.056 \(\times\) 10⁶ s^-1

B)2.742 \(\times\) 10⁶ s^-1

C)6.165 \(\times\) 10¹⁴ s^-1

D)8.226 \(\times\) 10¹⁴ s^-1

E)> 10¹⁵ s^-1

Q2) A photon has an energy of 5.53 \(\times\) 10^-17 J.What is its frequency in s^-1?

A)3.66 \(\times\) 10^-50 s^-1

B)1.20 \(\times\) 10^-17 s^-1

C)3.59 \(\times\) 10^-9 s^-1

D)2.78 \(\times\) 10⁸ s^-1

E)8.35 \(\times\) 10¹⁶ s^-1

Q3) In the quantum mechanical treatment of the hydrogen atom,the probability of finding an electron at any point is proportional to the wave function \(\varPsi\) .

A)True

B)False

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Page 9

Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) In a single atom,what is the maximum number of electrons which can have quantum number n = 4?

A)16

B)18

C)32

D)36

E)None of these choices is correct.

Q2) The difference in energies between the 1s and 2s orbitals is due to the penetration effect.

A)True

B)False

Q3) The electronic structure

1s²2s²2p⁶3s²3p⁶4s²3d⁸ refers to the ground state of

A)Kr

B)Ni

C)Fe

D)Pd

E)None of these choices is correct.

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Page 10

Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) The lattice energy of large ions is greater in magnitude than that of small ions of the same charge.

A)True

B)False

Q2) Give a clear and concise definition of the term "electronegativity";i.e. ,what does it measure?

Q3) Based on electronegativity trends in the periodic table,predict which of the following compounds will have the greatest % ionic character in its bonds.

A)H₂O

B)LiI

C)CaO

D)RbF

E)HCl

Q4) The more C-O and O-H bonds there are in a substance,the greater will be the amount of heat released when a fixed mass of the substance is burned.

A)True

B)False

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11

Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) According to VSEPR theory,a molecule with the general formula

AX₂E₃ will have a _____ molecular shape.

A)bent

B)linear

C)trigonal planar

D)T-shaped

E)trigonal pyramidal

Q2) Which one of the following molecules and ions will have a planar geometry?

A)PCl₃

B)BF₄^-

C)XeF₄

D)BrF₅

E)H₃O⁺

Q3) Which of the following molecules has a net dipole moment?

A)BeCl₂

B)SF₂

C)KrF₂

D)CO₂

E)CCl₄

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) The angles between sp² hybrid orbitals are 109.5°.

A)True

B)False

Q2) In molecular orbital theory,combination of two 2p atomic orbitals may give rise to either \(\sigma\) or \(\pi\) type molecular orbitals.

A)True

B)False

Q3) In molecular orbital theory,combination of two atomic orbitals produces two molecular orbitals.

A)True

B)False

Q4) In the context of molecular orbital (MO)theory,explain how atomic p orbitals can give rise to:

a.a \(\sigma\) MO

b.a \(\pi\) MO

Q5) Describe the bonding in ethylene (ethene,C₂H₄)according to valence bond theory.Be sure to indicate the orbital hybridization on the carbon atoms.Draw a diagram clearly showing at least one bond of each type which occurs in the molecule.

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Chapter 12: Intermolecular Forces: Liquids,solids,and Phase Changes

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Sample Questions

Q1) Comparing the energies of the following intermolecular forces on a kJ/mol basis,which would normally have the highest energy (i.e. ,be the strongest force)?

A)ion-induced dipole

B)dipole-induced dipole

C)ion-dipole

D)dipole-dipole

E)dispersion

Q2) a.Name the two unit cells which occur in close packing of identical atoms. b.Briefly explain how the two types of close-packed lattices of identical atoms differ,in terms of atomic arrangements.

Q3) Which of the following atoms should have the smallest polarizability?

A)Si

B)S

C)Te

D)Bi

E)Br

Q4) The surface tension of water is lowered when a detergent is present in solution.

A)True

B)False

14

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Chapter 13: The Properties of Solutions

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Sample Questions

Q1) Which of the following aqueous solutions will have the lowest osmotic pressure?

A)0.10 m KOH

B)0.10 m RbCl

C)0.05 m CaSO₄

D)0.05 m BaCl₂

E)0.10 m K₂SO₄

Q2) What volume of concentrated (14.7 M)phosphoric acid is needed to prepare 25.0 L of 3.0 M H₃PO₄?

A)0.20 L

B)0.57 L

C)1.8 L

D)3.6 L

E)5.1 L

Q3) When two pure substances are mixed to form a solution

A)heat is released.

B)heat is absorbed.

C)there is an increase in entropy.

D)there is a decrease in entropy.

E)entropy is conserved.

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Page 15

Chapter 14: The Main-Group Elements: Applying Principles of Bonding and Structure

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Sample Questions

Q1) Which of the following elements exists in allotropic forms?

A)silicon

B)germanium

C)carbon

D)lead

E)xenon

Q2) Most of the alkali metal salts are soluble in water while many alkaline earth salts have very low solubilities.Why is this so?

A)The alkali metal cations are smaller than the alkaline earth cations and are more easily hydrated.

B)The alkali metals have lower ionization energies than alkaline earth elements.

C)The alkaline earth salts have much greater lattice energies than the alkali metal salts.

D)The alkaline earth metals have greater heats of atomization than the alkali metals.

E)Alkaline earth cations have very low heats of hydration.

Q3) Potassium nitrate is a strong reducing agent.

A)True

B)False

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Page 16

Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) Carboxylic acids are weak acids.

A)True

B)False

Q2) Helical and sheet-like segments in proteins arise from

A)disulfide bridges.

B)salt bridges.

C)crosslinking via covalent bonds.

D)dispersion forces within the protein's interior.

E)hydrogen bonding.

Q3) Draw and name two non-alkenes with the formula C₄H₈.

Q4) a.Draw two different structures with the molecular formula C₂H₆O.

b.Name the functional group in each structure.

c.Which one will have the higher boiling point,and why?

Q5) Esters can be formed by the dehydration-condensation of a carboxylic acid and an alcohol.

A)True

B)False

Page 17

Q6) Describe and contrast fibrous and globular proteins in terms of their amino acid composition,their structure and their function.

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) The rate of a reaction is determined by the rate of the fastest step in the mechanism. A)True

B)False

Q2) The decomposition of dinitrogen pentaoxide to nitrogen dioxide and oxygen follows first-order kinetics and has an activation energy of 102 kJ/mol.By what factor will the fraction of collisions with energy greater than or equal to the activation energy increase if the reaction temperature goes from 30°C to 60°C?

A)1.00

B)1.10

C)2.00

D)4.00

E)38.4

Q3) A transition state is a species (or state)corresponding to an energy maximum on a reaction energy diagram.

A)True

B)False

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) Which of the following has an effect on the magnitude of the equilibrium constant?

A)activation energy of the forward reaction

B)concentrations of the reactants and products

C)presence of a catalyst

D)change in volume of container

E)change in temperature

Q2) For a solution equilibrium,a change in concentration of a reactant or product does not change K_c.

A)True

B)False

Q3) Although a system may be at equilibrium,the rate constants of the forward and reverse reactions will in general be different.

A)True

B)False

Q4) Unless \(\Delta\)H°_rxn = 0,a change in temperature will affect the value of the equilibrium constant K_c.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) Arrhenius bases raise the hydroxide ion concentration when dissolved in water.

A)True

B)False

Q2) It is not possible to have a pH lying outside the range 0 to 14.

A)True

B)False

Q3) All Brønsted-Lowry bases contain the hydroxide ion,OH^-.

A)True

B)False

Q4) The substance Ca(OH)₂ is considered

A)a weak Arrhenius acid.

B)a weak Arrhenius base.

C)a strong Arrhenius acid.

D)a strong Arrhenius base.

E)a neutral compound.

Q5) (a)Write a balanced equation representing the reaction of the acid,H ₂PO₄^- with the base,water.

(b)Write the expression for K_a of H₂PO₄^- in terms of concentrations of relevant species.

Page 21

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Q1) Calculate the solubility of copper(II)carbonate,CuCO₃,in 1.00 mol L^-1 NH₃.K_sp = 3.0 \(\times\) 10^-12 for CuCO₃,K_f = 5.6 \(\times\) 10¹¹ for Cu(NH₃)₄²⁺

Q2) The indicator propyl red has K_a = 3.3 \(\times\) 10^-6.What would be the approximate pH range over which it would change color?

A)3.5-5.5

B)4.5-6.5

C)5.5-7.5

D)6.5-8.5

E)None of these choices is correct.

Q3) A solution is prepared by adding 500 mL of 0.3 M NaClO to 500 mL of 0.4 M HClO.What is the pH of this solution?

A)The pH will be greater than the pK_a of hypochlorous acid.

B)The pH will be less than the pK_a of hypochlorous acid.

C)The pH will be equal to the pK_a of hypochlorous acid.

D)The pH will equal the pK_b of sodium hypochlorite.

E)The pH will be none of these.

Q4) Make a clear distinction between buffer range and buffer capacity.

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Chapter 20: Thermodynamics: Entropy, free Energy, and the

Direction of Chemical Reactions

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Q1) Which of the following is always true for an exothermic process?

A)q_sys > 0,\(\Delta\)S_surr < 0

B)q_sys < 0,\(\Delta\)S_surr > 0

C)q_sys < 0,\(\Delta\)S_surr < 0

D)q_sys > 0,\(\Delta\)S_surr > 0

E)w < 0

Q2) In the expression,S = k ln W,W is called the number of microstates.Explain clearly the meaning of the word "microstate",and why a system under a given set of conditions normally has many microstates.

Q3) As a chemical reaction proceeds toward equilibrium,the free energy of the system decreases.

A)True

B)False

Q4) a.Explain what is meant by a spontaneous process.

b.Is a spontaneous process necessarily a rapid one? Explain,and provide a real reaction as an example to illustrate your answer.

Q5) A chemical reaction has \(\Delta\)G° = 10.0 kJ and \(\Delta\)S° = 50.0 J/K

a.Calculate \(\Delta\)H° for this reaction at 25°C.

b.Could this reaction ever be spontaneous? Explain your answer.

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) a.Write a balanced equation to represent the overall reaction you would expect to occur in the electrolysis of molten KCl.

b.Write a balanced equation to represent the overall reaction you would expect to occur in the electrolysis of aqueous KCl.

c.Clearly explain why the products of the two processes are not the same.

Q2) A voltaic cell is prepared using copper and silver.Its cell notation is shown below. Cu(s)| Cu²⁺(aq)|| Ag⁺(aq)| Ag(s)

Which of the following processes occurs at the cathode?

A)Cu(s) \(\to\) Cu²⁺(aq)+ 2e^- B)Cu²⁺(aq)+ 2e^- \(\to\) Cu(s)

C)Ag(s) \(\to\) Ag⁺(aq)+ e^- D)Ag⁺(aq)+ e^- \(\to\) Ag(s) E)Cu(s)+ 2Ag⁺(aq) \(\to\) Cu²⁺(aq)+ 2Ag(s)

Q3) A primary battery is one which can be recharged.

A)True

B)False

Q4) Write down equations representing the anode half-reaction,the cathode half-reaction and the overall cell reaction for the lead-acid battery.

Page 24

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Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) Of the 3d transition series of elements,scandium has the greatest atomic radius.

A)True

B)False

Q2) Which of the following is considered a bidentate ligand?

A)cyanide,CN^-

B)thiocyanate,SCN^-

C)oxalate,C₂O₄^2-

D)nitrite,NO₂^-

E)hydroxide,OH^-

Q3) What is the difference between a coordination compound and a complex ion?

Q4) Tetrahedral complexes can exhibit both optical and linkage isomerism.

A)True

B)False

Q5) a.State the requirement for two molecules to be optical isomers.

b.A complex ion MABCD²⁺ (where A,B,C and D are different unidentate ligands)rotates the plane of polarized light.Deduce the geometry of the complex and draw the optical isomers of this ionic formula.

Page 25

Q6) a.How can the formation of a complex ion be described in terms of a theory of acids and bases?

b.What is the essential requirement for a molecule or ion to act as a ligand?

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) Carbon-14 will emit a \(\beta\) particle with an energy of 0.1565 MeV.What is this energy in joules?

A)1.0 \(\times\) 10^-²⁴ J

B)2.5 \(\times\) 10^-20 J

C)1.0 \(\times\) 10^-18 J

D)2.5 \(\times\) 10^-14 J

E)None of these choices is correct.

Q2) No alpha decay is observed for isotopes of elements with Z < 83.

A)True

B)False

Q3) When an electron and its anti-particle,a positron,collide,they annihilate each other.Calculate the energy released in this process,in J.(The positron mass is the same as the electron mass,namely 9.11 \(\times\) 10^-31 kg. )

Q4) Bombardment of uranium-238 nuclei by carbon-12 nuclei produces californium-246 and neutrons.Write a complete,balanced equation for this nuclear process.

Q5) Most foodstuffs contain natural,radioactive isotopes.

A)True

B)False

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