General Chemistry II Solved Exam Questions - 3326 Verified Questions

General Chemistry II

Solved Exam Questions

Course Introduction

General Chemistry II is a continuation of foundational chemical principles, focusing on topics such as chemical kinetics, chemical equilibrium, acid-base chemistry, thermodynamics, electrochemistry, and introductory concepts in coordination and nuclear chemistry. This course emphasizes problem-solving, laboratory techniques, and critical analytical thinking, providing students with a comprehensive understanding of the interactions and energy changes that govern chemical systems. Through lectures, discussions, and laboratory experiments, students deepen their knowledge of theoretical concepts and develop practical skills applicable to advanced studies in chemistry and related fields.

Recommended Textbook

Chemistry A Molecular Approach 1st Canadian Ediiton by Nivaldo J. Tro

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25 Chapters

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Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) What answer should be reported, with the correct number of significant figures, for the following calculation?

10^-4.68

A)2.1 × 10^-5

B)2 09 × 10^-5

C)2 × 10^-5

D)2.089 × 10^-5

E)2.0893 × 10^-5

Answer: A

Q2) How many square kilometres (km²)are in 246 square miles? (1 km = 0.621371 miles)

A)153 km²

B)481 km²

C)396 km²

D)95.0 km²

E)637 km²

Answer: E

Q3) Define matter.

Answer: Matter is anything that occupies space and has mass.

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following is a transition element?

A)Pd

B)Sn

C)K

D)U

E)Pr Answer: A

Q2) Predict the charge that an aluminum ion would have.

A)5-

B)1+

C)1-

D)2+

E)3+ Answer: E

Q3) Which one of the following elements is a poor conductor of heat and electricity?

A)copper

B)fluorine

C)iron

D)lead

Answer: B

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) Why aren't prefixes used in naming ionic compounds?

Answer: The charges on the ions dictate how many must be present to form a neutral unit. Molecular compounds do not have such constraints and therefore must use prefixes to denote the number of atoms present.

Q2) Identify the formula for barium hydroxide octahydrate.

A)Ba(OH)₂8H₂O

B)BaOH8H₂O

C)B(OH)₃8H₂O

D)8Ba(OH)₂H₂O

E)8BaOHH₂O

Answer: A

Q3) The solid compound Na₂CO₃ contains A)Na⁺, C⁴⁺, and O^2- ions.

B)Na⁺ and CO₃^-2 ions.

C)Na₂⁺ and CO₃^-2 ions.

D)Na₂CO₃ molecules.

Answer: B

Q4) Determine the mass percent (to the hundredths place)of H in sodium bicarbonate (NaHCO₃).

Answer: 1.20

Page 5

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Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) What reagent could be used to separate Br^- from CH₃CO₂^-when added to an aqueous solution containing both?

A)AgNO₃(aq)

B)Ba(OH)₂(aq)

C)CuSO₄(aq)

D)NaI₍_aq₎

Q2) What is the concentration of magnesium ions in a 0.125 mol L^-1 Mg(NO₃)₂ solution?

A)0.125 mol L^-1

B)0.0625 mol L^-1

C)0.375 mol L^-1

D)0.250 mol L^-1

E)0.160 mol L^-1

Q3) What is the concentration (M)of CH₃OH in a solution prepared by dissolving 16.8 g of CH₃OH in sufficient water to give exactly 230 mL of solution?

Q4) Balance the following equation: _____ C₁₀H₁₂ + _____ O₂ _____ H₂O + _____ CO₂

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Chapter 5: Gases

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Sample Questions

Q1) Give the formula for ozone.

A)O₃

B)O₂

C)O

D)H₂O₂

E)H₂O

Q2) How many grams of F₂ gas are there in a 5.00 L cylinder at 5.327 bar and 23 °C?

A)19.5 g

B)41.1 g

C)500 g

D)2.96 × 10⁴ g

Q3) You have three cylinders containing O₂ gas at the same volume and pressure. Cylinder A is at -15 °C, cylinder B is at -5 °F, cylinder C is at 255 K. Which cylinder contains the largest mass of oxygen?

A)cylinder A

B)cylinder B

C)cylinder C

D)All cylinders contain the same mass of O₂.

Q4) Give the major gas in dry air.

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Which of the following processes is endothermic?

A)an atom emits a photon

B)the condensation of water

C)an atom absorbs a photon

D)the electron affinity of a fluorine atom

E)None of the above processes is endothermic.

Q2) The specific heat of copper is 0.385 J g^-1 K^-1. If 34.2 g of copper, initially at 21.0 °C, absorbs 4.689 kJ, what will be the final temperature of the copper?

A)21.4 °C

B)23.8 °C

C)356 °C

D)377 °C

Q3) Describe the energy changes that occur when a book is held 1 m off the floor and then dropped.

Q4) Explain the difference between H and U.

Q5) Write the reaction illustrating _fH° of CaCO₃ and draw the enthalpy diagram showing the relative positions of the reactants and products if it is an exothermic reaction.

Q6) Where does the energy absorbed during an endothermic reaction go?

Page 8

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) What is the maximum number of f orbitals that are possible?

A)1

B)3

C)7

D)5

E)9

Q2) n = 1 to n = 2

A)657 nm

B)7460 nm

C)103 nm

D)1280 nm

E)122 nm

Q3) Give the number of valence electrons for Mg²⁺.

Q4) Calculate the wavelength of an electron (m = 9.11 × 10^-2⁸

g)moving at 3.66 × 10⁶ m s^-1.

A)1.99 × 10^-10 m

B)5.03 × 10^-10 m

C)1.81 × 10^-10 m

D)5.52 × 10^-9 m

E)2.76 × 10^-9 m

Page 9

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Identify the element that has a ground-state electronic configuration of [Kr]5s²4d⁵.

A)Tc

B)Mn

C)Nb

D)Ru

Q2) An element that has the valence electron configuration 5s²5p³ belongs to which period and group?

A)period 5; group 15

B)period 5; group 2

C)period 5; group 3

D)period 5; group 5

E)period 4; group 13

Q3) How many valence electrons do the alkali earth metals possess? A)1 B)2

Q4) Why do successive ionization energies increase?

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Place the following in order of decreasing bond length. H-F H-I H-Br

A)H-F > H-Br > H-I

B)H-I > H-F > H-Br

C)H-I > H-Br > H-F

D)H-Br > H-F > H-I

E)H-F > H-I > H-Br

Q2) Define dipole moment.

Q3) Using Lewis structures and formal charge, which of the following ions is most stable?

OCN ONC NOC

A)OCN

B)ONC

C)NOC

D)None of these ions is stable according to Lewis theory.

E)All of these compounds are equally stable according to Lewis theory.

Q4) Which of the following compounds is hypercoordinate?

A)CsI

B)SnO₂

C)ClF₅

D)ClF

Q5) Describe the difference between a pure covalent bond and a polar covalent bond.

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,

Valence Bond Theory, and Molecular Orbital Theory

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Sample Questions

Q1) The hybrid orbital set used by the central atom in KrF₂ is ________.

A)sp

B)sp²

C)sp³

D)sp³d

E)sp³d²

Q2) Draw the Lewis structure for OF₂. What is the hybridization on the O atom?

A)sp

B)sp³

C)sp²

D)sp³^d

E)sp³^d²

Q3) How many of the following molecules are polar? XeO₂

SiCl₂Br₂ C₂Br₂ SeCl₆ A)1

Page 12

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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Define deposition.

A)A liquid becomes a gas.

B)A gas becomes a liquid.

C)A gas becomes a solid.

D)A solid becomes a gas.

E)A solid becomes a liquid.

Q2) What type of intermolecular force causes the dissolution of NaCl in water?

A)hydrogen bonding

B)dipole-dipole forces

C)ion-dipole force

D)dispersion forces

E)none of the above

Q3) Choose the substance with the highest boiling point.

A)CH₄

B)KI

C)CS₂

D)HF

E)I₂

Q4) Describe the difference between the conduction band and the valence band.

Q5) Define viscosity.

13

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Chapter 12: Solutions

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Sample Questions

Q1) Calculate the mass of maltose, C₁₂H₂₂O₁₁, contained in 725.0 mL of solution if the measured vapour pressure of the solution is 41.6 mbar at 30 °C. The vapour pressure of pure water at 30 °C is 42.40 mbar; assume no change in volume occurs upon dissolution.

A)482 g

B)329 g

C)265 g

D)163 g

E)94.1 g

Q2) Determine the freezing point of a solution that contains 78.8 g of naphthalene (C₁₀H₈, molar mass = 128.16 g mol^-1)dissolved in 722 mL of benzene (d = 0.877 g mL^-1). Pure benzene has a melting point of 5.50 °C and a freezing point depression constant of 4.90 °C m^-1.

A)4.76 °C

B)4.17 °C

C)0.74 °C

D)1.33 °C

E)1.68 °C

Q3) Define the Tyndall effect.

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) What are the units of k in a zero-order reaction?

A)mol L^-1 s^-1

B)mol L^-1 s

C)L mol^-1s^-1

D)M² s^-1

E)s^-1

Q2) The first-order rearrangement of CH₃NC is measured to have a rate constant of 3.61 × 10^-15 s^-1 at 298 K and a rate constant of 8.66 × 10^-7 s^-1 at 425 K. Determine the activation energy for this reaction.

A)160. kJ mol^-1

B)240. kJ mol^-1

C)417 kJ mol^-1

D)127 kJ mol^-1

E)338 kJ mol^-1

Q3) What function do enzymes serve?

Q4) What happens to the concentration of reactants and products during a chemical reaction?

Q5) Define half-life.

Q6) Explain what the exponential factor in the Arrhenius equation represents.

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction, equilibrium concentrations, and equilibrium constant at a particular temperature. Determine the equilibrium concentration of SO₃(g). 2SO₂(g)+ O₂(g) 2SO₃(g) K_c = 1.7 × 10⁸M^-1

[SO₂]_eq = 0.0034 mol L^-1 [O₂]_eq = 0.0018 mol L^-1

A)1.9 mol L^-1

B)1.0 × 10³ mol L^-1

C)0.53 mol L^-1

D)9.6 × 10^-4 mol L^-1

E)0.73 mol L^-1

Q2) For the isomerization reaction: butane isobutane

K_p equals 25.0 at 500 °C. If the initial pressures of butane and isobutane are 20.0 atm and 0.0 atm, respectively, what are the pressures of the two gases at equilibrium?

A)P(butane)= 0.77 atm and P(isobutane)= 19.2 atm

B)P(butane)= 0.80 atm and P(isobutane)= 20. atm

C)P(butane)= 19.2 atm and P(isobutane)= 0.77 atm

D)P(butane)= 20 atm and P(isobutane)= 0.80 atm

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Which Bronsted-Lowry acid is not considered to be a strong acid in water?

A)HBr

B)HI

C)HIO₄

D)HNO₃

Q2) Which one of the following will form a basic solution in water?

A)NaBr

B)LiCN

C)KClO₂

D)KClO₄

E)NH₄Cl

Q3) What is the difference between a strong and weak acid?

Q4) The stronger the acid:

A)the stronger the conjugate acid.

B)the stronger the conjugate base.

C)the weaker the conjugate base.

D)the weaker the conjugate acid.

E)the stronger the conjugate acid-base pair.

Q5) Describe the relationship between molecular structure and acid strength.

Q6) Describe a molecule that can be a Lewis acid.

Page 17

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Calculate the pH of a buffer that is 0.105 mol L^-1 CH₃COOH and 0.146 mol L^-1 CH₃COOK. The K_a for CH₃COOH is 1.8 × 10^-5.

A)4.89

B)9.11

C)4.74

D)9.26

E)4.60

Q2) Which of the following compounds will be more soluble in acidic solution than in pure water?

A)PbCl₂

B)FeS

C)Ca(ClO₄)₂

D)CuI

E)AgBr

Q3) Define a buffer.

Q4) Sketch the titration curve for a diprotic acid. Make sure to label the equivalence point(s)and the relative amounts of base that were added to reach them.

Q5) Identify the compound that forms stalactites and stalagmites.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Place the following in order of increasing standard molar entropy. H₂O(l)H₂O(g)H₂O(s)

A)H₂O(g)< H₂O(l)< H₂O(s) B)H₂O(s)< H₂O(l)< H₂O(g)

C)H₂O(g)< H₂O(s)< H₂O(l) D)H₂O(l)< H₂O(s)< H₂O(g) E)H₂O(s)< H₂O(g)< H₂O(l)

Q2) How is a nonspontaneous process made spontaneous?

Q3) Calculate _rG at 298 K under the conditions shown below for the following reaction: Fe₂O₃(s)+ 3CO(g) 2Fe(s)+ 3CO₂(g) G° = -28.0 kJ mol^-1 P(CO)= 1.4 atm, P(CO₂)= 2.1 atm

A)+31.0 kJ mol^-1

B)+2.99 kJ mol^-1

C)-30.7 kJ mol^-1

D)+17.5 kJ mol^-1

E)-25.0 kJ mol^-1

Q4) Why can't we say that a spontaneous reaction is a fast reaction?

Q5) Why is heating your home with gas more efficient than heating it with electricity?

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Identify the location of reduction in an electrochemical cell.

A)the anode

B)the cathode

C)the electrode

D)the salt bridge

E)the socket

Q2) Define a salt bridge.

A)A pathway, composed of salt water, that ions pass through.

B)A pathway in which no ions flow.

C)A pathway between the cathode and anode in which ions are reduced.

D)A pathway between the cathode and anode in which ions are oxidized.

E)A pathway through which counterions can flow between the half-cells without the solutions in the half-cell totally mixing.

Q3) Why are iron nails coated with zinc?

Q4) Identify the characteristics of a spontaneous reaction.

A) _rG° > 0

B) E°_cell < 0

C)K < 0

D) E°_cell > 0

E)K = 0

Page 20

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Determine the binding energy per nucleon of a Mg-24 nucleus. The Mg-24 nucleus has a mass of 24.30506. A proton has a mass of 1.00728 amu, a neutron has a mass of 1.008665 amu, and 1 amu is equivalent to 931 MeV of energy.

A)0.3050 MeV

B)8.83 MeV

C)0.113 MeV

D)106 MeV

E)4.41 MeV

Q2) Describe what changes occur in the atomic nucleus during alpha decay.

A)The mass number and atomic number decrease.

B)The mass number and atomic number increase.

C)The mass number is unchanged and the atomic number decreases.

D)The mass number is unchanged and the atomic number increases.

E)The mass number and atomic number do not change.

Q3) Which particle has the lowest penetrating power?

A)alpha particle

B)beta particle

C)gamma rays

D)positron emission

E)electron capture

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Chapter 20: Organic Chemistry I: Structures

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Q1) Which of the following statements is TRUE?

A)Organic halides are polar molecules with C-X bond polarity increasing in order C-I < C-Br < C-Cl < C-F.

B)Alkenes are hydrocarbons that contain one or more C=C bonds and therefore must have some carbons sp hybridized.

C)Amines are organic compounds that contain nitrogen and have R-C(=O)-NH-R as their general formula.

D)Esters are a family of organic compounds that contain oxygen and have the general formula R-O-R.

E)Because alcohols contain nonpolar C-H bonds, they cannot mix with water.

Q2) What is meant by the term, "structural isomer"? Draw a structural isomer of CH₃-CH₂-CH₂-CH₃.

Q3) Which of the following is a secondary alcohol?

A)CH₃CH₂OH

B)(CH₃)₂(CH₃CH₂)COH

C)(CH₃)₃COH

D)(CH₃)₂CHCH(CH₃)OH

E)CH₃OCH₃

Q4) What is the difference between primary, secondary, and tertiary amines?

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Chapter 21: Organic Chemistry II: Reactions

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Q1) What is the product obtained after a reaction between an aldehyde and Grignard reagent is worked-up with dilute acid?

A)a ketone

B)a primary alcohol

C)a secondary alcohol

D)a tertiary alcohol

E)a carboxylic acid

Q2) What are secondary alcohols oxidized to?

A)tertiary alcohols

B)ethers

C)esters

D)aldehydes

E)ketones

Q3) Which part of HCN, the hydrogen or the cyano group, adds to the O in a C=O bond and why?

Q4) Why doesn't benzene typically undergo addition reactions like the alkenes do?

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Chapter 22: Biochemistry

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Sample Questions

Q1) In the sequence His-Tyr-Val-Pro, which amino acid contains the N-terminal group?

A)valine

B)tyrosine

C)proline

D)histidine

Q2) The following is part of a DNA sequence. What is its complementary sequence?

AGTTCGAGCCT

A)AGTTCGAGCCT

B)TCCGAGCTTGA

C)CAGTCCA

D)AGGCTCGAACT

E)TCAAGCTCGGA

Q3) Which of the following shows all of the tripeptides that can be formed from one molecule each of lysine (Lys), threonine (Thr), and histidine (His)?

A)LysThrHis

B)LysThrHis, LysHisThr, HisThrLys, HisLysThr, ThrLysHis, ThrHisLys

C)LysThrHis, LysHisThr, HisLysThr, ThrHisLys

D)LysThrHis, LysHisThr, HisLysThr

E)LysThrHis and LysHisThr

Q4) Why are lipids well-suited as structural components of cell membranes?

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Chapter 23: Chemistry of the Nonmetals

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Sample Questions

Q1) Determine the formula for an aluminosilicate where one-third of the silicons are replaced by an aluminum atom. The charge is balanced by potassium ions.

A)K₂(AlO₂)(SiO₂)₂

B)K(AlO₂)(SiO₂)

C)K(AlO₂)₂(SiO₂)₂

D)K(AlO₂)(SiO₂)₂ E)K₂(AlO₂)₂(SiO₂)₂

Q2) What is the oxidation number of oxygen in H₂O₂?

A)-4

B)+2 C)-2

D)-1

E)+1

Q3) Determine the oxidation state of phosphorus in P₂O₇⁴ .

A)+5

B)+4 C)+3

D)+7

E)+2

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Chapter 24: Metals and Metallury

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Sample Questions

Q1) Pb

A)sphalerite

B)rhodochrosite

C)rutile

D)carnotite

E)malachite

F)galena

G)cinnabar

Q2) Which substance is the best conductor of electricity?

A)arsenic

B)boron

C)silver

D)tellurium

Q3) Which of the following is TRUE concerning interstitial alloys?

A)The metal takes on the properties of the foreign atoms once they are present.

B)Small atoms fit in between the metal atoms in the crystal structure.

C)Brass is an example of an interstitial alloy.

D)The foreign atoms present weaken the metallic structure.

E)They are alloys composed of two nonmetallic elements.

Q4) Why is zinc used to coat steel objects?

Page 26

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) How many unpaired electrons would you expect for the complex ion [FeF₆]^4-?

A)2

B)6

C)4

D)3

E)0

Q2) What is the ground-state electron configuration for Cr³⁺ (Z = 24)?

A)[Ar] 4s²3d⁷

B)[Ar] 4s²3d¹

C)[Ar] 3d³

D)[Ar] 4s¹3d²

Q3) What is the ground-state electron configuration for the element chromium (Z = 24)?

A)[Ne] 4s²3d⁴

B)[Ar] 4s²3d⁴

C)[Ar] 4s¹3d⁵

D)[Ar] 3d⁶

Q4) What is the difference between a weak-field complex and a strong-field complex?

Q5) Explain how EDTA is used to treat lead poisoning.

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