Introductory Chemistry Final Test Solutions - 2309 Verified Questions

Introductory Chemistry

Final Test Solutions

Course Introduction

Introductory Chemistry provides a fundamental overview of chemical principles and concepts, serving as a foundation for further study in chemistry and related sciences. The course covers essential topics such as atomic and molecular structure, chemical bonding, stoichiometry, states of matter, chemical reactions, periodic trends, and basic thermodynamics. Emphasis is placed on problem-solving skills, laboratory techniques, and real-world applications of chemistry in everyday life. Students will develop an understanding of how scientific reasoning and quantitative analysis are used to explore the properties and behaviors of matter.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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24 Chapters

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Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) The potential energy of a car moving on a level road does not depend on its speed.

A)True

B)False

Answer: True

Q2) The appropriate number of significant figures in the result of 15.234 × 15.208 is

A) 2.

B) 3.

C) 4.

D) 5.

E) 6.

Answer: D

Q3) Given that 1 inch = 2.54 cm, 1 cm³ is equal to

A) 16.4 in³.

B) 6.45 in³.

C) 0.394 in³.

D) 0.155 in³.

E) 0.0610 in³.

Answer: E

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3

Chapter 2: The Components of Matter

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Sample Questions

Q1) Zinc acetate is used in preserving wood and in manufacturing glazes for porcelain. What is its formula?

A) ZnAc₂

B) ZnCH₃COO

C) Zn(CH₃COO)₂

D) Zn₂CH₃COO

E) ZnCH₃COCH₃

Answer: C

Q2) Which one of the following combinations of names and formulas of ions is incorrect?

A) O^2- oxide

B) Cd²⁺ cadmium

C) ClO₃^- chlorate

D) HCO₃^- hydrogen carbonate

E) NO₂^- nitrate

Answer: E

Q3) Lead (Pb) is a main-group element.

A)True

B)False

Answer: True

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Balance the following equation for partial oxidation of ammonia, an important reaction in the production of nitric acid: NH₃(g) + O₂(g) \(\to\) NO(g) + H₂O(l)

Answer: 4NH₃(g) + 5O₂(g) \(\to\) 4NO(g) + 6H₂O(l)

Q2) Ammonia will react with fluorine to produce dinitrogen tetrafluoride and hydrogen fluoride (used in production of aluminum, in uranium processing, and in frosting of light bulbs). 2NH₃(g) + 5F₂(g) \(\to\) N₂F₄(g) + 6HF(g)

How many moles of NH₃are needed to react completely with 13.6 mol of F₂?

A) 34.0 mol

B) 27.2 mol

C) 6.80 mol

D) 5.44 mol

E) 2.27 mol

Answer: D

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Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) Which one of the following ionic compounds is soluble in water?

A) Na₂S

B) PbI₂

C) AgCl

D) CuS

E) Ca₃(PO₄)₂

Q2) Automobile batteries use 3.0 M H₂SO₄as an electrolyte. How much 1.20 M NaOH will be needed to neutralize 225 mL of battery acid? H₂SO₄(aq) + 2NaOH(aq) \(\to\)2H₂O(l) + Na₂SO₄(aq)

A) 0.045 L

B) 0.28 L

C) 0.56 L

D) 0.90 L

E) 1.1 L

Q3) Complete and balance the equation for the following acid-base reaction. Ca(OH)₂(aq) + HCl(aq) \(\to\)

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) A 750-mL sample of hydrogen exerts a pressure of 822 torr at 325 K. What pressure does it exert if the temperature is raised to 475 K at constant volume?

A) 188 torr

B) 562 torr

C) 1.11 × 10³ torr

D) 1.20 × 10³ torr

E) 1.90 × 10³ torr

Q2) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650. torr.

Q3) Arrange the following gases in order of increasing rate of effusion.

C₂H₆Ar HCl PH₃

A) Ar < HCl < PH₃< C₂H₆

B) C₂H₆< PH₃< HCl < Ar

C) Ar < PH₃< C₂H₆< HCl

D) C₂H₆< HCl < PH₃< Ar

E) Ar < PH₃< HCl< C₂H₆

Q4) For a pure gas sample, the average kinetic energy is also the most probable kinetic energy.

A)True

B)False

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) The standard heat (enthalpy) of formation of graphite is zero.

A)True

B)False

Q2) Cold packs, whose temperatures are lowered when ammonium nitrate dissolves in water, are carried by athletic trainers when transporting ice is not possible. Which of the following is true of this reaction?

A) .\(\Delta\)H < 0, process is exothermic

B) .\(\Delta\)H > 0, process is exothermic

C) .\(\Delta\)H < 0, process is endothermic

D) .\(\Delta\)H > 0, process is endothermic

E) .\(\Delta\)H = 0, since cold packs are sealed

Q3) For a reaction in a sealed, rigid container, \(\Delta\)H is always greater than \(\Delta\)E.

A)True

B)False

Q4) Starting from equations relating pressure to force and force to work, derive the relationship w = - P\(\Delta\)V, explaining the steps in your argument.

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) a. Use the Bohr equation to calculate the energy needed to ionize a hydrogen atom from its ground state.

b. What is the minimum wavelength of a photon needed for it to have the energy needed in (a)?

Q2) Which of the following frequencies of electromagnetic radiation has the shortest wavelength?

A) 1 kilohertz

B) 1 terahertz

C) 1 dekahertz

D) 1 gigahertz

E) 1 megahertz

Q3) Which scientist demonstrated that photons transferred momentum during collisions with matter?

A) Bohr

B) de Broglie

C) Planck

D) Compton

E) Billiard

Q4) What is the speed of an electron in m/s if its wavelength is 0.155 nm?

Q5) The following combinations of quantum numbers are not allowed.

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Which of the following electron configurations is not possible for an atom in an excited state?

A)

1s²2s²2p⁶3s²3p⁶3d<sup >10</sup>4s¹4p¹

B) 1s²2s²2p⁶3s¹3p⁵

C)

1s²2s²2p⁶3s²3p²3d²

D)

1s²2s²2p⁶3s²3p⁶3d<sup >10</sup>4s²

E)

1s²2s²2p⁶3s²3p⁶3d<sup >10</sup>4s¹⁴p³

Q2) Select the most acidic compound from the following:

A) SO₂

B) Al₂O₃

C) CaO

D) PbO

E) H₂O

Q3) State Hund's rule, and show how it applies to the ground state of phosphorus atoms.

Page 10

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) When two atoms form a covalently-bonded diatomic molecule, the distance between the nuclei at which the potential energy is at a minimum is called

A) the bond energy.

B) the bond length.

C) the molecular diameter.

D) the covalent radius.

E) the covalent diameter.

Q2) Which of the following elements is the most electronegative?

A) Ne

B) Rb

C) P

D) I

E) Cl

Q3) Give a clear and concise definition of the term "electronegativity"; i.e., what does it measure?

Q4) Covalently bonded substances do not necessarily exist as separate molecules.

A)True

B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Which of the following molecules has a net dipole moment?

A) BeCl₂

B) SF₂

C) KrF₂

D) CO₂

E) CCl₄

Q2) In which one of the following is the best Lewis structure a resonance structure?

A) SO₃

B) BF₃

C) I₃^-

D) SCO (C = central atom)

E) SO₃^2-

Q3) According to VSEPR theory, a molecule with the general formula AX₃ will have a ______ molecular shape.

A) linear

B) bent

C) trigonal planar

D) tetrahedral

E) trigonal pyramidal

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) Valence bond theory predicts that tin will use _____ hybrid orbitals in SnCl₃^-.

A) sp

B) sp²

C) sp³

D) sp³d

E) sp³d²

Q2) Overlap of two sp² hybrid orbitals produces a \(\pi\) bond.

A)True

B)False

Q3) Explain what is meant by the term "bond order" and describe how it can be calculated using the information in a molecular orbital energy level diagram.

Q4) According to molecular orbital theory, what is the bond order in the O₂⁺ ion?

A) 5.5

B) 5

C) 4

D) 2.5

E) 1.5

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

A)True

B)False

Q3) What types of forces exist between molecules of CO₂?

A) hydrogen bonding only.

B) hydrogen bonding and dispersion forces.

C) dipole-dipole forces only.

D) dipole-dipole and dispersion forces.

E) dispersion forces only.

Q4) Germanium is a semiconductor. With the aid of diagrams showing bands of molecular orbital, explain why it is a poor conductor and how doping it with phosphorus increases its conductivity.

Page 14

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Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) List five important intermolecular forces that operate in globular proteins, and explain what parts of the molecule are involved in these forces.

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) The mole fraction of potassium nitrate in an aqueous solution is 0.0194. The solution's density is 1.0627 g/mL. Calculate the molarity of the solution.

A) 0.0194 M

B) 0.981 M

C) 1.05 M

D) 1.96 M

E) 19.4 M

Q4) The density of pure water at 25°C is 0.997 g/mL. Considering water as being both solvent and solute, calculate its molarity.

Page 15

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) When elements from a group exhibit more than one oxidation state,

A) the higher oxidation state is more important as one goes down a group.

B) the lower oxidation state is more important as one goes down a group.

C) both oxidation states are equally important throughout the group.

D) the oxidation state will be affected by the elements on either side of it in the period.

E) None of the above conclusions is true.

Q2) Tin reacts to form several organic compounds. Which of the following is a use for organotin compounds?

A) They are starting materials for polymers.

B) They can be used as fungicides.

C) They can be used to prepare semiconductors.

D) They can be used as solvents for reactions involving polar organic compounds.

E) They are used as protective coatings to prevent corrosion of metals.

Q3) Nitric acid, HNO₃

A) is a strong reducing agent.

B) is a weak oxidizing agent.

C) does not react with copper metal.

D) may react with copper metal to produce NO gas.

E) reacts with metal hydroxides to produce nitrite salts.

Page 16

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Chapter 15: Organic Compounds and the Atomic Properties of

Carbon

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Sample Questions

Q1) Condensation polymers are formed by free radical addition reactions.

A)True

B)False

Q2) The backbone of protein molecules consists of repeating N-C-C-O units.

A)True

B)False

Q3) Addition polymers are mostly made from monomers with carbon-carbon double bonds.

A)True

B)False

Q4) What is the difference between 1-butyne and 1-butene?

A) 1 carbon atom

B) 1 carbon atom and 2 hydrogen atoms

C) 2 hydrogen atoms

D) 4 hydrogen atoms

E) 1 carbon atom and 4 hydrogen atoms

Q5) In one sentence, what is the general requirement for a molecule to be optically active?

Q6) Amino acids in solution can undergo an internal acid-base reaction.

A)True

B)False Page 17

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) Ammonium cyanate (NH₄CNO) reacts to form urea (NH₂CONH₂). At 65°C the rate constant, k, is 3.60 L mol^-1s^-1. What is the rate law for this reaction?

A) Rate = 3.60 L mol^-1s^-1[NH₄CNO]

B) Rate = 3.60 L mol^-1s^-1[NH₄CNO]²

C) Rate = 0.28 mol L^-1 s^-1[NH₄CNO]

D) Rate = 0.28 mol L^-1 s^-1[NH₄CNO]²

E) Rate = 3.60 L mol^-1s^-1[NH₂CONH₂]^-1

Q2) Sucrose decomposes to fructose and glucose in acid solution. When ln [sucrose] is plotted vs. time, a straight line with slope of -0.208 hr^-1 results. What is the rate law for the reaction?

A) Rate = 0.208 hr^-1 [sucrose]²

B) Rate = 0.208 hr^-1 [sucrose]

C) Rate = 0.0433 hr [sucrose]²

D) Rate = 0.0433 hr [sucrose]

E) Rate = 0.208 mol L^-1hr^-1[sucrose]⁰

Q3) Briefly list the features/properties common to all catalysts and how they work. Draw a labeled reaction energy diagram as part of your answer.

Page 19

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) There is a direct correlation between the speed of a reaction and its equilibrium constant.

A)True B)False

Q2) A good catalyst for a reaction will speed up the forward reaction and slow down the reverse reaction.

A)True B)False

Q3) When a chemical system is at equilibrium,

A) the concentrations of the reactants are equal to the concentrations of the products.

B) the concentrations of the reactants and products have reached constant values.

C) the forward and reverse reactions have stopped.

D) the reaction quotient, Q, has reached a maximum.

E) the reaction quotient, Q, has reached a minimum.

Q4) For a solution equilibrium, a change in concentration of a reactant or product does not change K_c.

A)True B)False

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Chapter 18: Acid-Base Equilibria

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Q1) If a strong acid such as HCl is diluted sufficiently with water, the pH will be higher than 7.

A)True

B)False

Q2) All weak acids have strong conjugate bases.

A)True

B)False

Q3) A solution is prepared by adding 0.10 mol of sodium sulfide, Na₂S, to 1.00

L of water. Which statement about the solution is correct?

A) The solution is basic.

B) The solution is neutral.

C) The solution is acidic.

D) The concentration of sodium ions and sulfide ions will be identical.

E) The concentration of sulfide ions will be greater than the concentration of sodium ions.

Q4) Define an acid according to the Arrhenius theory, and write a balanced equation to support this definition.

Q5) Describe what is meant by the "leveling effect". Use a real acid as an example, and write an appropriate equation.

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Q1) What is the maximum mass of KCl that can be added to1.0 L of a 0.010 M lead(II) chloride solution without causing any precipitation of lead(II) chloride? Assume that addition of KCl does not affect the solution volume. For lead(II) chloride, K_sp = 1.6 × 10^-5.

A) 3.0 g

B) 1.5 g

C) 0.8 g

D) 1.0 g

E) 0.2 g

Q2) Which of the following indicators would be the best to use when 0.050 M benzoic acid (K_a = 6.6 × 10^-5) is titrated with 0.05 M NaOH?

A) bromphenol blue, pH range: 3.0-4.5

B) bromcresol green, pH range: 3.8-5.4

C) alizarin, pH range: 5.7-7.2

D) phenol red, pH range: 6.9-8.2

E) phenolphthalein, pH range: 8.0-10.1

Q3) 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) State the second and third laws of thermodynamics.

Q2) The entropy of one mole of oxygen gas in a 0.5-L container is less than it would be in a 22.4-L container at the same temperature.

A)True

B)False

Q3) Consider the reaction

11ec6d53_97e1_49fd_9cd1_030bf8888a28_TB5833_

If the concentrations of the Cu⁺ and I^- ions in equilibrium at 298 K are both equal to 1.03 × 10^-6 M, what is the value of \(\Delta\)G° for the reaction?

A) -68 kJ

B) 68 kJ

C) -30. kJ

D) 30 kJ

E) 34 kJ

Q4) The free energy of a perfect crystal at absolute zero, is zero.

A)True

B)False

Page 23

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) The line notation, Al(s) | Al³⁺(aq) || Co²⁺(aq) | Co(s), indicates that

A) Co is the reducing agent.

B) Co²⁺ ions are oxidized.

C) Al is the reducing agent.

D) Al³⁺ is the reducing agent.

E) aluminum metal is the cathode.

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 buried iron pipe can be protected against corrosion by connecting it to a rod of copper.

A)True

B)False

Page 24

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Chapter 22: The Elements in Nature and Industry

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Q1) The main effect of the biosphere on the chemistry of the earth's crust has been to

A) create a reducing atmosphere.

B) create an oxidizing atmosphere.

C) increase the relative humidity.

D) decrease the relative humidity.

E) increase the level of atmospheric carbon dioxide.

Q2) The main end use of sulfuric acid is

A) manufacture of iron and steel.

B) processing of petroleum.

C) production of fertilizers.

D) production of explosives.

E) production of various other chemicals.

Q3) The most abundant element in the Earth's crust is

A) H.

B) Fe.

C) Al.

D) Si.

E) O.

Q4) Briefly describe the three main pathways for nitrogen fixation.

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) Square planar complexes can exhibit both geometric and optical isomerism. A)True

B)False

Q2) Apply the valence bond theory to predict the electronic structure and hybridization pattern of chromium in the complex ion Cr(NH₃)₆³⁺.

Q3) What is the highest possible oxidation state for palladium, Pd?

A) +1

B) +2

C) +3

D) +4

E) +6

Q4) The most common oxidation state for ions of the inner transition elements is

A) +2.

B) +3.

C) +4.

D) +5.

E) +7.

Page 26

Q5) Why is the +2 oxidation state so common among transition elements?

Q6) What is the difference between a coordination compound and a complex ion?

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) Which of the following series of radioactive decays would convert Pa-234 to Ra-226?

A) beta, alpha, beta

B) alpha, alpha

C) beta, alpha, alpha, beta

D) beta, alpha, alpha

E) alpha, beta, gamma

Q2) Calculate to four significant figures

a. the mass defect in kg, and b. the energy released in kJ/mol, when a neutron decays to produce a proton and an electron. The neutron, proton, and electron masses are 1.67493 × 10^-27 kg, 1.67262 × 10^-27 kg and 9.10939 × 10^-31 kg, respectively.

Q3) What is the specific activity (in Ci/g) of an isotope if 3.56 mg emits 4.26 × 10⁸ \(\beta\) particles per second?

A) 0.003232 Ci/g

B) 0.0115 Ci/g

C) 0.309 Ci/g

D) 3.23 Ci/g

E) none of the above

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