General Chemistry II Exam Answer Key - 2075 Verified Questions

General Chemistry II

Exam Answer Key

Course Introduction

General Chemistry II builds upon the foundational principles introduced in General Chemistry I, delving deeper into chemical concepts such as thermodynamics, kinetics, chemical equilibrium, acid-base chemistry, electrochemistry, and coordination compounds. The course emphasizes both qualitative and quantitative problem-solving skills, introduces advanced laboratory techniques, and explores the applications of these chemical principles in real-world contexts. Through lectures, discussions, and hands-on experiments, students will develop a comprehensive understanding of the behavior of matter, the energy changes driving chemical reactions, and the dynamic interactions that govern chemical systems.

Recommended Textbook Chemistry and Chemical Reactivity 9th Edition by John C. Kotz\New folder

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26 Chapters

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Page 2

Chapter 1: Basic Concepts of Chemistry

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Q1) A pure substance composed of two or more different elements is a(n)________.

A) ion

B) heterogeneous mixture

C) chemical compound

D) solid

E) solution

Answer: C

Q2) Which of the following are likely to form a homogeneous mixture?

1) milk and ice cream blended together with chocolate syrup

2) an egg combined with milk and mixed with a whisk

3)1 gram table salt combined with 250 mL of water

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Answer: C

Q3) Potential energy possessed by water at the top of a waterfall is known as ________ energy.

Answer: gravitational

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Chapter 2: Mathlets Review: The Tools of Quantitative Chemistry

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Q1) Convert 5.100 \(\times\)10^-8 grams to nanograms and express the answer in fixed notation using the correct number of significant figures.

A) 51 ng

B) 51.0 ng

C) 51.00 ng

D) 5.10 ng

E) 0.5100 ng

Answer: C

Q2) What volume of a pure liquid (density 0.710 g/mL)has a mass of 0.290 kg?

A) 4.08 \(\times\) 10² mL

B) 2.45 \(\times\) 10^-3 mL

C) 2.06 \(\times\) 10^-1 mL

D) 2.45 mL

E) 4.08 \(\times\) 10^-1 mL

Answer: A

Q3) Assuming the density of water is 1.00 g/cm³,the mass of 1.0 cubic meter (m³) of water is ________ grams.

Answer: 1.0 \(\times\) 10⁶

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Chapter 3: Atoms, molecules, and Ions

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Q1) How many protons,neutrons,and electrons are in a neutral atom of ⁵⁵Fe?

A) 26 protons,29 neutrons,55 electrons

B) 26 protons,29 neutrons,29 electrons

C) 26 protons,29 neutrons,26 electrons

D) 55 protons,26 neutrons,55 electrons

E) 55 protons,26 neutrons,26 electrons

Answer: C

Q2) Which of the following atoms contains the fewest protons?

A) (²³²Th)

B) (²³¹Pa)

C) (²⁴⁵Pu)

D) (²³⁸U)

E) (²³²Pa)

Answer: A

Q3) William Crookes was this first to observe particles produced from a cathode ray tube.These particles eventually became known as ________.

Answer: electrons

Q4) Oxygen and ________ are the two most abundant elements in the Earth's crust. Answer: silicon

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Chapter 4: Chemical Reactions

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Q1) Which anion will form a precipitate with Ca²⁺?

A) Cl^-

B) OH^-

C) C₂H₃O₂^-

D) Br^-

E) none

Q2) Which of the following statements is/are CORRECT?

1)A solution is a homogeneous mixture of two or more substances.

2)A solute is a mixture of a solvent and a soluble compound.

3)Aqueous solutions are solutions in which water is a solvent.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1 and 3

Q3) A(n)________ agent gains electrons in an oxidation-reduction reaction.

Q4) A solution is a homogeneous mixture composed of one or more ____________ dissolved in a solvent.

Q5) Ionic and molecular compounds that form ions in aqueous solution are called _____.

Page 6

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Chapter 5: Stoichiometry: Quantitative Information About Chemical Reactions

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Q1) How many grams of dioxygen are required to completely burn 4.3 g of C₂H₅OH?

A) 19 g

B) 15 g

C) 3.0 g

D) 42 g

E) 58 g

Q2) Chlorine was passed over 2.02 g of heated titanium,and 6.50 g of a chloride-containing compound of Ti was obtained.What is the empirical formula of the chloride-containing compound?

A) TiCl₂

B) TiCl₄

C) TiCl

D) TiCl₃

E) Ti₂Cl₃

Q3) The percent yield of a chemical reaction is calculated by dividing the actual yield by the ________ yield and multiplying by 100%.

Q4) The numbers preceding the formulas in chemical equations are referred to as the ________ coefficients.

Page 7

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Chapter 6: Principles of Chemical Reactivity: Energy and Chemical Reactions

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Q1) Exactly 212.2 J will raise the temperature of 10.0 g of a metal from 25.0 °C to 60.0 °C.What is the specific heat capacity of the metal?

A) 0.606 J/(g·°C)

B) 1.65 J/(g·°C)

C) 14.5 J/(g·°C)

D) 50.8 J/(g·°C)

E) none of these

Q2) When 66.0 g of an unknown metal at 28.5 S1U1P1\(\circ\)S1S1P0C is placed in 83.0 g H₂O at 78.5 S1U1P1\(\circ\)S1S1P0C,the water temperature decreases to 75.9 S1U1P1\(\circ\)S1S1P0C.What is the specific heat capacity of the metal? The specific heat capacity of water is 4.184 J/g.K.

A) 0.055 J/g.K

B) 0.29 J/g.K

C) 0.69 J/g.K

D) 0.18 J/g.K

E) 2.6 J/g.0K

Q3) In thermodynamics,a(n)________ is defined as the object,or collection of objects,being studied.The surroundings include everything else.

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Chapter 7: The Structure of Atoms

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Q1) Which of the following regions of the electromagnetic spectrum has the lowest frequency?

A) x-ray

B) gamma ray

C) ultraviolet

D) infrared

E) visible

Q2) Which of the following statements is INCORRECT?

A) It is not possible to know the exact location of an electron and its exact energy simultaneously.

B) The energies of an atom's electrons are quantized.

C) Quantum numbers define the energy states and the orbitals available to an electron.

D) The behavior of an atom's electrons can be described by circular orbits around a nucleus.

E) Electrons have both wave and particle properties.

Q3) According to Heisenberg's ________ principle,it is impossible to simultaneously measure the exact location and energy of an electron.

Q4) A point in a standing wave that has zero amplitude is called a(n)________.

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Chapter 8: The Structure of Atoms and Periodic Trends

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Q1) What is the maximum number of electrons that can occupy the n = 1 shell?

A) 2

B) 8

C) 18

D) 32

E) 50

Q2) What is the maximum number of electrons that can occupy \(\underline{\text{ one }}\) d orbital?

A) 1

B) 2

C) 6

D) 10

E) 14

Q3) Choose the statement that is \(\textbf{ true: }\)

A) Outer electrons efficiently shield one another from nuclear charge

B) Core electrons effectively shield outer electrons from nuclear charge

C) Valence electrons are most difficult of all electrons to remove

D) Core electrons have the lowest ionization energies of all electrons

E) Two of the above are true.

Q4) Explain the difference between paramagnetic and ferromagnetic.

Page 10

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Chapter 9: Bonding and Molecular Structure

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Q1) What is the total number of electrons (both lone and bond pairs)in the Lewis structure of C₂O₄2-</sup>?

A) 30

B) 34

C) 32

D) 46

E) 44

Q2) What are the approximate H-N-H bond angles in NH₄⁺?

A) 109.5S1U1P1\(\circ\)S1S1P0

B) 120S1U1P1\(\circ\)S1S1P0

C) 109.5S1U1P1\(\circ\)S1S1P0 and 120S1U1P1\(\circ\)S1S1P0

D) 90S1U1P1\(\circ\)S1S1P0

E) 90S1U1P1\(\circ\)S1S1P0 and 120S1U1P1\(\circ\)S1S1P0

Q3) Which of the following species has the shortest carbon-nitrogen bond?

A) CH₃CN

B) CH₃N₂

C) CH₂NH

D) CH₃CONH₂

E) NO₃^-

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Chapter 10: Bonding and Molecular Structure: Orbital

Hybridization and Molecular Orbitals

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Q1) Ammonia reacts with oxygen and water to produce nitric acid.What change in hybridization of the nitrogen atom occurs in this reaction?

A) sp³ to sp²

B) sp³ to sp

C) sp² to sp³

D) sp² to sp

E) no change

Q2) Which of the following concerning \(\sigma\) and \(\pi\) bonds is/are correct?

1)A sigma bond may be formed from the sideways overlap of two parallel p orbitals.

2)No more than two pi bonds are possible between adjacent carbon atoms.

3)The considerable energy required to rotate pi bonded atoms is the primary reason for geometrical isomerism in some pi bonded molecules.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 2 and 3

Q3) The hybridization of the carbon atom in CF₃⁺ is

Page 12

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Chapter 11: Gases and Their Properties

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Q1) A balloon is filled with He gas to a volume of 2.10 L at 35 S1U1P1\(\circ\)S1S1P0C.The balloon is placed in liquid nitrogen until its temperature reaches -196 S1U1P1\(\circ\)S1S1P0C.Assuming the pressure remains constant,what is the volume of the cooled balloon?

A) (-0.375) L

B) 0.375 L

C) 0.525 L

D) 0.00909 L

E) 8.40 L

Q2) A particular gas exerts a pressure of 7.54 \(\times\) 10⁴ Pa.What is this pressure in units of atmospheres? (1 atm = 760 mm Hg = 101.3 kPa = 1.013 bar)

A) 0.744 atm

B) 7.44 \(\times\) 10⁹ atm

C) 0.764 atm

D) 0.754 atm

E) 7.64 \(\times\) 10⁹ atm

Q3) Robert Boyle observed that the volume of a fixed amount of gas is inversely proportional to its ________.

Q4) What are two ways in which real gases differ from ideal gases?

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Chapter 12: Intermolecular Forces and Liquids

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Q1) What intermolecular force(s)is/are present in solid NH₃?

1)induced dipole/induced dipole

2)dipole-dipole

3)hydrogen bonding

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1 and 3

Q2) Which of the following nonpolar molecules has the highest boiling point?

A) N₂

B) C₂H₄

C) F₂

D) O₂

E) CS₂

Q3) Make a sketch to show the hydrogen bonding between two acetic acid molecules (HC₂H₃O₂).

Q4) Above its critical temperature and pressure,a substance becomes a(n)________ fluid.

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Chapter 13: The Chemistry of Solids

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Q1) Gold (atomic mass 197.0 g/mol),with an atomic radius of 144.2 pm,crystallizes in a face-centered cubic lattice.What is the density of gold?

A) 9.65 g/cm³

B) 1.21 g/cm³

C) 4.82 g/cm³

D) 2.41 g/cm³

E) 19.3 g/cm³

Q2) For a metal that crystallizes in a body-centered cubic unit cell,what percentage of the space in the cell is occupied by the metal atoms?

A) 47%

B) 52%

C) 68%

D) 74%

E) 87%

Q3) A salt with a 1:1 ratio of anions to cations may pack in a face-centered cubic unit cell with the anions at the lattice points and the cations occupying one-half of the ________ holes.Zinc sulfide is an example of this structure.

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Chapter 14: Solutions and Their Behavior

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Q1) What is the freezing point of an aqueous 1.36 m NaCl solution? (K_fp for water is 1.858°C/m.)Assume no ion pairing occurs.

A) 5.1°C

B) 2.5°C

C) -5.1°C

D) 0.0°C

E) -2.5°C

Q2) What concentration of silver nitrate (in ppm)is present in 7.1 \(\times\) 10^-7 M AgNO₃(aq)? For very dilute aqueous solutions,you can assume the solution's density is 1.0 g/mL.The molar mass of AgNO₃ is 169.9 g/mol.

A) 0.0071 ppm

B) 0.12 ppm

C) 0.71 ppm

D) 1.7 ppm

E) 8.3 ppm

Q3) ________ are colloidal dispersions of one liquid in another liquid.

Q4) A surfactant used for cleaning is called a(n)________.

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Page 16

Chapter 15: Chemical Kinetics: the Rates of Chemical Reactions

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Q1) A first-order chemical reaction is observed to have a rate constant of 34 min^-1.What is the corresponding half-life for the reaction?

A) 1.2 s

B) 1.2 min

C) 49 min

D) 1.8 s

E) 48.6 s

Q2) Which of the following statements is correct for the first-order reaction: A \(\to\) 2B?

A) The concentration of A decreases linearly with respect to time.

B) The concentration of A is constant with respect to time.

C) The natural logarithm of the concentration of A decreases linearly with respect to time.

D) The rate of reaction is constant with respect to time.

E) The rate constant,k,of the reaction decreases linearly with respect to time.

Q3) The pre-exponential,A,in the Arrhenius equation is called the ________ factor.

Q4) Elementary steps in a reaction mechanism often include reaction ________.These (usually)short-lived species,which are at one point produced and then later consumed,do not appear in the overall chemical reaction.

Page 17

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Chapter 16: Principles of Reactivity: Chemical Equilibria

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Q1) Which of the following statements is/are CORRECT?

1)For a chemical system,if the reaction quotient (Q)is greater than K,reactant must be converted to products to reach equilibrium.

2)For a chemical system at equilibrium,the forward and reverse rates of reaction are equal.

3)For a chemical system at equilibrium,the concentrations of products divided by the concentrations of reactants equals one.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q2) In 1913,the Haber-Bosch process was patented.The product of the Haber-Bosch process is ________.

Q3) The symbol Q is called the ________.

Q4) If a stress is applied to an equilibrium system,the system will respond in such a way as to relieve that stress.This is a statement of ________ principle.

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Chapter 17: The Chemistry of Acids and Bases

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Q1) A 0.10 M solution of a weak monoprotic acid has a hydronium-ion concentration of 4.6 \(\times\) 10^-4 M.What is the acid-ionization constant,K_a,for this acid?

A) 2.1 \(\times\) 10^-2

B) 3.2 \(\times\) 10^-3

C) 4.6 \(\times\) 10^-4

D) 2.1 \(\times\) 10^-6

E) 5.5 \(\times\) 10^-5

Q2) What is the pOH of 0.067 M HI(aq)at 25 S1U1P1\(\circ\)S1S1P0C? (K_w = 1.01 \(\times\) 10^-14)?

A) 2.70

B) 12.83

C) 11.30

D) 15.17

E) 1.17

Q3) When a Lewis acid combines with a Lewis base,the base supplies both electrons to the bond.This type of chemical bond is called a(n)________ covalent bond.

Q4) A molecule that can behave as either a Brønsted-Lowry acid or base is termed

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Chapter 18: Principles of Reactivity: Other Aspects of Aqueous Equilibria

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Q1) What is the molar solubility of Mn(OH)₂(s)in a solution that is buffered at pH 8.00 at 25 S1U1P1\(\circ\)S1S1P0C? The K_sp of Mn(OH)₂ is 1.9 \(\times\) 10^-13 at 25 S1U1P1\(\circ\)S1S1P0C.

A) 3.6 \(\times\) 10^-8 mol/L

B) 1.9 \(\times\) 10^-7 mol/L

C) 3.6 \(\times\) 10^-5 mol/L

D) 1.9 \(\times\) 10^-1 mol/L

E) 1.9 \(\times\) 10³ mol/L

Q2) A 50.00-mL solution of 0.0797 M ammonia (K_b = 1.8 \(\times\) 10^-5)is titrated with a 0.0315 M solution of hydrochloric acid as the titrant.What is the pH of the base solution after 23.88 mL of titrant have been added? (K_w = 1.00 \(\times\) 10^-14)

A) 11.08

B) 9.88

C) 4.12

D) 4.74

E) 12.90

Q3) If the ratio of acid to base in a buffer is increased by a factor of 10,the buffer pH decreases by ________.

Page 20

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Chapter 19: Entropy and Free Energy

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Q1) Which of the following changes lead to a \(\underline{\text{ decrease }}\) in entropy?

A) the sublimation (vaporization)of dry ice (solid carbon dioxide)

B) sugar dissolving in coffee

C) evaporation of water from a lake

D) diffusion of perfume throughout a room

E) halving the volume of a gas

Q2) When a chemical process occurs under standard conditions,which of the following conditions always apply?

1)Gaseous species are at a pressure of 1 bar.

2)Solution concentrations species are 1 molal.

3)The temperature is 298.15 K.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q3) At what temperature (in kelvin units)is the entropy of a pure crystal 0.0 J/K.

Q4) In any chemical process,energy must be conserved.This is a statement of the ________ law of thermodynamics.

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Chapter 20: Principles of Reactivity: Electron Transfer Reactions

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Q1) Write a balanced chemical equation for the overall reaction represented by the cell notation below.Mn(s)| Mn²⁺(aq)|| Zn²⁺(aq)| Zn(s)

A) 2Mn(s)+ Zn²⁺(aq)\(\to\) Zn(s)+ 2Mn²⁺(aq)

B) Mn(s)+ Mn²⁺(aq)\(\to\) Zn(s)+ Zn²⁺(aq)

C) Mn(s)+ Zn²⁺(aq)\(\to\) Zn(s)+ Mn²⁺(aq)

D) 2Zn(s)+ Mn²⁺(aq)\(\to\)Mn(s)+ 2Zn²⁺(aq)

E) Zn(s)+ Mn²⁺(aq)\(\to\)Mn(s)+ Zn²⁺(aq)

Q2) Write a balanced chemical equation for the overall reaction represented by the cell notation below.

Al(s)| Al³⁺(aq)|| Br^-(aq)| Br₂(g)| Pt(s)

A) 2Al(s)+ 3Br₂(g)\(\to\) 2Al³⁺(aq)+ 6Br^-(aq)

B) Al(s)+ Al³⁺(aq)\(\to\)Br^-(aq)+ Br₂(g)

C) 2Al³⁺(aq)+ 6Br^-(aq)\(\to\) 2Al(s)+ 3Br₂(g)

D) Al(s)+ 3Br₂(g)\(\to\) Al³⁺(s)+ 2Br^-(aq)

E) Al(s)+ 2Br^-(aq)\(\to\) Br₂(g)+ Al³⁺(aq)

Q3) When a secondary battery provides electrical energy,it is acting as a voltaic cell.When the battery is recharging,it is operating as a(n)________ cell.

Q4) Explain the function of a salt bridge in a voltaic cell.

Page 22

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Chapter 21: Environmental Chemistry Earths

Environment,energy,and Sustainability

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Q1) Which ion occurs in the highest concentration in seawater?

A) Cl^-

B) Na⁺

C) HCO₃^-CO₃^2-

D) Mg²⁺

E) Ca²⁺

Q2) The majority of electricity generated in the United States is produced from ___.

A) wind and solar

B) moving water

C) nuclear fission

D) natural gas

E) coal

Q3) The majority of crude oil ends up being used

A) to make synthetic fibers.

B) to make fine chemicals.

C) to make plastics.

D) as fuel.

E) to make rubber.

Q4) Explain why fossil fuels,hydrocarbons produced by natural processes,are considered a nonrenewable energy resource.

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Chapter 23: The Chemistry of the Transition Elements

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Q1) The white in most white paints comes from ________.

A) TiO₂

B) CdS

C) Fe₄[Fe(CN)₆]₃.14 H₂O

D) KMnO₄

E) AgI

Q2) What is the coordination number of the transition metal ion in [Ag(NH₃)₂]Cl?

A) 2

B) 3

C) 4

D) 6

E) 8

Q3) Which of the following is \(\underline{\text{ not }}\) a transition-metal?

A) vanadium

B) silver

C) iron

D) chromium

E) tin

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Chapter 24: Carbon: Not Just Another Element

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Q1) What are the products of a condensation reaction between a primary amine and a carboxylic acid?

A) water and an amide

B) water and an ester

C) ammonia and an amide

D) ammonia and an ester

E) an ammonium cation and a carboxylate anion

Q2) Which one of the following hydrocarbons has cis and trans isomers?

A) hexane

B) 1-pentene

C) 2-methylpropane

D) 2-ethylbutane

E) 2-pentene

Q3) The reaction of methane with chlorine in the presence of ultraviolet radiation can produce ____.

A) chloromethane

B) dichloromethane

C) trichloromethane

D) carbon tetrachloride

E) all of these products

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Chapter 25: Biochemistry

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Q1) What is the complementary base which hydrogen bonds with thymine in double stranded DNA?

A) adenine

B) cytosine

C) guanine

D) thymine

E) uracil

Q2) Which statement concerning the twenty different \(\alpha\)-amino acids widely found in nature is true?

A) Each \(\alpha\)-amino acid has a different number of carbon atoms.

B) Each \(\alpha\)-amino acid has one amino (-NH₂)group.

C) Each \(\alpha\)-amino acid has one carboxylic acid (-COOH)group.

D) These \(\alpha\)-amino acids differ in the R groups attached to the \(\alpha\)-carbon.

E) None of these \(\alpha\)-amino acids are aromatic compounds.

Q3) The process by which a double stranded of DNA is copied is called ___.In the process,a ___ strand of DNA is produced from each single strand of the original double-stranded DNA.

Q4) The oxygen that bridges two sugars in a disaccharide is known as a ___ bond.

Q5) A compound which contains both a positive and negative charge is called a ____.

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Chapter 26: Nuclear Chemistry

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Q1) Which of the following types of radiation has the greatest quality factor?

A) beta particles

B) alpha particles emitted within the body

C) gamma rays

D) low energy protons

E) low energy neutrons

Q2) What is the percent activity of a radioactive sample (relative to its original activity)that has undergone two half-lives of decay?

A) 25.0%

B) 12.5%

C) 6.25%

D) 3.13%

E) 75.0%

Q3) Explain the difference between 1 rad and 1 rem of radiation.

Q4) If a nucleus undergoes positron particle emission

A) its atomic number decreases by four and its mass number decreases by two.

B) its atomic number decreases by two and its mass number decreases by four.

C) its atomic number increases by one and its mass number is unchanged.

D) its atomic number decreases by one and its mass number is unchanged.

E) its atomic number is unchanged and its mass number decreases by one.

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