Fundamentals of Chemistry Textbook Exam Questions - 3679 Verified Questions

Fundamentals of Chemistry

Textbook Exam Questions

Course Introduction

Fundamentals of Chemistry provides an essential introduction to the principles and concepts that form the basis of the chemical sciences. Students will explore atomic and molecular structure, chemical bonding, stoichiometry, states of matter, solutions, and basic thermodynamics. The course emphasizes laboratory techniques, problem-solving skills, and real-world applications of chemistry, preparing students for advanced study in science and related fields.

Recommended Textbook Chemistry The Central Science 13th Edition by Theodore E. Brown

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24 Chapters

3679 Verified Questions

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Chapter 1: Introduction: Matter and Measurement

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Sample Questions

Q1) The number with the most significant zeros is ________.

A)0.00002510

B)0.02500001

C)250000001

D)2.501 × 10^-⁷

E)2.5100000

Answer: C

Q2) "Absolute zero" refers to ________.

A)0 Kelvin

B)0° Fahrenheit

C)0° Celsius

D)°C + 9/5(°F - 32)

E)273.15 °C

Answer: A

Q3) The symbol for the element phosphorus is ________.

Answer: P

Q4) Temperature is a physical property that determines the direction of heat flow.

A)True

B)False

Answer: True

Page 3

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Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) Which pair of elements would you expect to exhibit the greatest similarity in their physical and chemical properties?

A)H, Li

B)Cs, Ba

C)Ca, Sr

D)Ga, Ge

E)C, O

Answer: C

Q2) Elements in Group 2A are known as the ________.

A)alkaline earth metals

B)alkali metals

C)chalcogens

D)halogens

E)noble gases

Answer: A

Q3) Which element in Group IA is the most electropositive?

Answer: francium

Q4) What is the name of an alcohol derived from hexane? Answer: hexanol

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Chapter 3: Stoichiometry: Calculations With Chemical

Formulas and Equations

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Sample Questions

Q1) Silver nitrate and aluminum chloride react with each other by exchanging anions: 3AgNO₃ (aq) + AlCl₃ (aq) Al(NO₃)₃ (aq) + 3AgCl (s)

What mass in grams of AgCl is produced when 4.22 g of AgNO₃ react with 7.73 g of AlCl₃?

A)17.6

B)4.22

C)24.9

D)3.56

E)11.9

Answer: D

Q2) The formula weight of Zn(ClO₄)₂ is ________ amu.

A)164.7

Answer: D

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Chapter 4: Aqueous Reactions and Solution Stoichiometry

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Sample Questions

Q1) When 0.500 mol of HC₂H₃O₂ is combined with enough water to make a 300.0 mL solution, the concentration of HC₂H₃O₂is ________ M.

A)3.33

B)1.67

C)0.835

D)0.00167

E)0.150

Q2) How many moles of BaCl₂ are formed in the neutralization of 393 mL of 0.171 M Ba(OH)₂ with aqueous HCl?

Q3) Which of the following is insoluble in water at 25 °C?

A)Mg₃(PO₄)₂

B)Na₂S

C)(NH₄)₂CO₃

D)Ca(OH)₂

E)Ba(C₂H₃O₂)₂

Q4) Ca(OH)₂ is a strong base.

A)True

B)False

Q5) The solvent in an aqueous solution is ________.

Page 6

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Chapter 5: Thermochemistry

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Sample Questions

Q1) The value of E for a system that performs 139 kJ of work on its surroundings and gains 54kJ of heat is ________ kJ.

A)-85

B)193

C)7506

D)85

E)-193

Q2) The internal energy of a system ________.

A)is the sum of the kinetic energy of all of its components

B)is the sum of the rotational, vibrational, and translational energies of all of its components

C)refers only to the energies of the nuclei of the atoms of the component molecules

D)is the sum of the potential and kinetic energies of the components

E)none of the above

Q3) The standard enthalpy change of a reaction is the enthalpy change when all reactants and products are at ________ pressure and a specific temperature.

Q4) ________ is defined as the energy used to move an object against a force.

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Chapter 6: Electronic Structure of Atoms

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Sample Questions

Q1) The condensed electron configuration of argon, element 18, is ________.

A)[Ne]3s⁴

B)[Ar]3s²3p²

C)[Ne]3s²3p⁶

D)[He]2s⁴2p¹⁰

E)[He]3s⁴

Q2) If an electron has a principal quantum number (n)of 7 and an angular momentum quantum number (l)of 3, the subshell designation is ________.

A)7f

B)7s

C)7p

D)3f

E)3d

Q3) The electron configuration of a ground-state Ag atom is ________.

A)[Ar]4s²4d⁹

B)[Kr]5s¹4d¹⁰

C)[Kr]5s²3d⁹

D)[Ar]4s¹4d¹⁰

E)[Kr]5s²4d¹⁰

Q4) The ground state electron configuration of scandium is ________.

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Chapter 7: Periodic Properties of the Elements

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Sample Questions

Q1) The first noble gas to be incorporated into a compound was ________.

A)Ar

B)Kr

C)He

D)Ne

E)Xe

Q2) The ion with the largest diameter is ________.

A)Po^2-

B)S^2-

C)Se^2-

D)Te^2-

E)O^2-

Q3) The ion with the smallest diameter is ________.

A)Li⁺

B)Na⁺

C)K⁺

D)Rb⁺

E)Cs⁺

Q4) Which alkali metals can react with oxygen to form either the peroxide or the superoxide?

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Chapter 8: Basic Concepts of Chemical Bonding

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Sample Questions

Q1) In ionic bond formation, the lattice energy of ions ________ as the magnitude of the ion charges _______ and the radii ________.

A)increases, decrease, increase

B)increases, increase, increase

C)decreases, increase, increase

D)increases, increase, decrease

E)increases, decrease, decrease

Q2) Electron affinity is a measure of how strongly an atom can attract additional electrons.

A)True

B)False

Q3) The strength of a covalent bond is measured by its ________.

Q4) The ________ ion has eight valence electrons.

A)Sc³⁺

B)Ti³⁺

C)V³⁺

D)Cr³⁺

E)Mn³⁺

Q5) Which halogen, bromine or iodine, will form the more polar bond with phosphorus?

Page 10

Q6) Polyatomic ions with an odd number of electrons will ________ the octet rule.

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Chapter 9: Molecular Geometry and Bonding Theories

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Sample Questions

Q1) The basis of the VSEPR model of molecular bonding is ________.

A)regions of electron density on an atom will organize themselves so as to maximize s-character

B)regions of electron density in the valence shell of an atom will arrange themselves so as to maximize overlap

C)atomic orbitals of the bonding atoms must overlap for a bond to form

D)electron domains in the valence shell of an atom will arrange themselves so as to minimize repulsions

E)hybrid orbitals will form as necessary to, as closely as possible, achieve spherical symmetry

Q2) A typical double bond ________.

A)is stronger and shorter than a single bond

B)consists of one bond and one bond

C)imparts rigidity to a molecule

D)consists of two shared electron pairs

E)All of the above answers are correct.

Q3) In the valence shell of an atom there are six electron domains. They will be arranged in a(n)________ geometry.

Q4) What are the three bond angles in the trigonal bipyramidal structure?

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Chapter 10: Gases

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Sample Questions

Q1) According to kinetic-molecular theory, in which of the following gases will the root-mean-square speed of the molecules be the highest at 200 °C?

A)NH₃

B)N₂O

C)O₂

D)UF₆

E)None. The molecules of all gases have the same root-mean-square speed at any given temperature.

Q2) Of the following, ________ has a slight odor of bitter almonds and is toxic.

A)NH₃

B)N₂O

C)CO

D)CH₄

E)HCN

Q3) The volume of HCl gas required to react with excess Ca to produce 11.4 L of hydrogen gas at 1.62 atm and 62.0 °C is ________ L.

Q4) What is the partial pressure (in mm Hg)of neon in a 4.00 L vessel that contains 0.838 mol of methane, 0.184 mol of ethane, and 0.755 mol of neon at a total pressure of 928 mm Hg?

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Chapter 11: Liquids and Intermolecular Forces

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Sample Questions

Q1) Cholesteric liquid crystals are colored because ________.

A)each molecule is a chromophore

B)of the slight twist between layers

C)of the large spacing between layers

D)of the large number of conjugated bonds

E)all of the molecules contain multiple benzene rings

Q2) The shape of a liquid's meniscus is determined by ________.

A)the viscosity of the liquid

B)the type of material the container is made of

C)the relative magnitudes of cohesive forces in the liquid and adhesive forces between the liquid and its container

D)the amount of hydrogen bonding in the liquid

E)the volume of the liquid

Q3) Based on the figure above, the boiling point of ethyl alcohol under an external pressure of 0.658 atm is ________ °C.

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Semiconductors are less conductive than metals because of ________ gap.

Q2) Which of the following is not a type of solid?

A)ionic

B)molecular

C)supercritical

D)metallic

E)covalent network

Q3) The molecular-orbital model for Ge shows it to be

A)a conductor, because all the lower energy band orbitals are filled and the gap between the lower and higher bands is large.

B)an insulator, because all the lower energy band orbitals are filled and the gap between the lower and higher bands is large.

C)a semiconductor, because the gap between the filled lower and empty higher energy bands is relatively small.

D)a semiconductor, because the gap between the filled lower and empty higher energy bands is large.

E)a conductor, because its lower energy band orbitals are only partially filled.

Q4) Write the chemical formulas for both polyethylene and the monomer from which it is formed.

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Chapter 13: Properties of Solutions

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Sample Questions

Q1) Which of the following liquids will have the highest freezing point?

A)pure H₂O

B)aqueous glucose (0.050 m)

C)aqueous CoI₂ (0.030 m)

D)aqueous FeI₃ (0.030 m)

E)aqueous NaI (0.030 m)

Q2) An aqueous solution of a soluble compound (a nonelectrolyte)is prepared by dissolving 33.2 g of the compound in sufficient water to form 250 mL of solution. The solution has an osmotic pressure of 1.2 atm at 25 °C. What is the molar mass (g/mole)of the compound?

A)1.0 × 10³

B)2.7 × 10³

C)2.3 × 10²

D)6.8 × 10²

E)28

Q3) After swimming in the ocean for several hours, swimmers noticed that their fingers appeared to be very wrinkled. This is an indication that seawater is supertonic relative to the fluid in cells.

A)True

B)False

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The mechanism for formation of the product X is: A + B C + D (slow)

B + D X (fast)

The intermediate reactant in the reaction is ________.

A)A

B)B

C)C

D)D

E)X

Q2) At elevated temperatures, dinitrogen pentoxide decomposes to nitrogen dioxide and oxygen: 2N₂O₅(g) 4NO₂(g) + O₂ (g) When the rate of formation of NO₂ is 5.5 × 10^-4 M/s, the rate of decomposition of N₂O₅ is ________ M/s.

A)2.2 × 10^-3

B)1.4 × 10^-4

C)10.1 × 10^-4

D)2.8 × 10^-4

E)5.5 × 10^-4

Q3) The Earth's ozone layer is located in the ________.

Q4) The minimum energy to initiate a chemical reaction is the ________.

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Exactly 3.5 moles if N₂O₄ is placed in an empty 2.0-L container and allowed to reach equilibrium described by the equation N₂O₄ (g) 2NO₂(g) If at equilibrium the N₂O₄ is 25% dissociated, what is the value of the equilibrium constant for the reaction?

Q2) At 900.0 K, the equilibrium constant (K_p)for the following reaction is 0.345. 2SO₂ + O₂ (g) 2SO₃ (g) At equilibrium, the partial pressure of SO₂ is 36.9 atm and that of O₂ is 16.8 atm. The partial pressure of SO₃ is ________ atm.

A)88.8

B)3.89 × 10^-³

C)214

D)5.57 × 10^-4

E)42.4

Q3) If the reaction quotient Q for a reaction is greater than the value of the equilibrium constant K for that reaction at a given temperature, ________ must be converted to ________ for the system to reach equilibrium.

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Chapter 16: Acid-Base Equilibria

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Sample Questions

Q1) Calculate the pOH of a solution at 25.0 °C that contains 1.94 × 10^-10 M hydronium ions.

A)1.94

B)4.29

C)7.00

D)14.0

E)9.71

Q2) A 0.1 M aqueous solution of ________ will have a pH of 7.0 at 25.0 °C. LiF RbBr

NaClO₄ NH₄Cl

A)RbBr and NaClO₄

B)LiF and RbBr

C)NaClO₄ only

D)LiF only

E)NH₄Cl only

Q3) Which one of the following is a Br nsted-Lowry acid?

A)(CH₃)₃NH⁺

B)CH₃COOH

C)HF

D)HNO₂

E)all of the above

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Chapter 17: Additional Aspects of Aqueous Equilibria

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Sample Questions

Q1) Which one of the following pairs cannot be mixed together to form a buffer solution?

A)NH₃, NH₄Cl

B)NaC₂H₃O₂, HCl (C₂H₃O₂^- = acetate)

C)RbOH, HBr

D)KOH, HF

E)H₃PO₄, KH₂PO₄

Q2) 200.0 ml of a solution containing 0.5000 moles of acetic acid per liter is added to 200.0 ml of 0.5000 M NaOH. What is the final pH? The K_a of acetic acid is 1.77 × 10^-5.

Q3) The molar solubility of ________ is not affected by the pH of the solution.

A)Na₃PO₄

B)NaF

C)KNO₃

D)AlCl₃

E)MnS

Q4) Although CaCO₃ has a relatively small solubility product, it is quite soluble in the presence of ________.

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Chapter 18: Chemistry of the Environment

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Sample Questions

Q1) How does lime reduce sulfur dioxide emissions from the burning of coal?

A)It reacts with the sulfur dioxide to form calcium sulfite solid that can be precipitated.

B)It reduces the sulfur dioxide to elemental sulfur that is harmless to the environment.

C)It oxidizes the sulfur dioxide to tetrathionate that is highly water soluble so it can be scrubbed from the emission gases.

D)It catalyzes the conversion of sulfur dioxide to sulfur trioxide which is much less volatile and can be removed by condensation.

E)It converts SO₂ to solid, elemental sulfur.

Q2) Which of the following is not a stage in water treatment?

A)coarse filtration

B)aeration

C)chlorination

D)distillation

E)settling

Q3) List two of the three major sources of nitrogen and phosphorus in water.

Q4) Describe the process of reverse osmosis that is used to desalinate seawater.

Q5) Approximately 90% of the earth's ozone is in the ________.

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Chapter 19: Chemical Thermodynamics

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Sample Questions

Q1) If G° for a reaction is less than zero, then ________.

A)K > 1

B)K = 1

C)K < 1

D)K = 0

E)more information is needed.

Q2) The value of S° for the decomposition of phosphorous trichloride into its constituent elements, 2PCl₃(g) P₂(g) + 3Cl₂( g) Is ________ J/K mol.

A)-311.7

B)+311.7

C)+263.6

D)+129.4

E)-129.4

Q3) Which reaction produces an increase in the entropy of the system?

A)Ag⁺ (aq) + Cl^- (aq) AgCl (s)

B)CO₂ (s) CO₂ (g)

C)H₂ (g) + Cl₂ (g) 2 HCl (g)

D)N₂ (g) + 3 H₂ (g) 2 NH₃ (g)

E)H₂O (l) H₂O (s)

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Chapter 20: Electrochemistry

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Sample Questions

Q1) Consider an electrochemical cell based on the reaction: 2H⁺ (aq) + Sn (s) Sn²⁺ (aq) + H₂ (g)

Which of the following actions would not change the measured cell potential?

A)lowering the pH in the cathode compartment

B)addition of more tin metal to the anode compartment

C)increasing the tin (II)ion concentration in the anode compartment

D)increasing the pressure of hydrogen gas in the cathode compartment

E)Any of the above will change the measured cell potential.

Q2) Galvanized iron is iron coated with ________.

A)magnesium

B)zinc

C)chromium

D)phosphate

E)iron oxide

Q3) The anode of the alkaline battery is powdered zinc in a gel that contacts

Q4) In the formula G = -nFE, F is the ________.

Q5) The quantity of charge passing a point in a circuit in one second when the current is one ampere is called a ________.

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Chapter 21: Nuclear Chemistry

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Sample Questions

Q1) All atoms of a given element have the same ________.

A)mass number

B)number of nucleons

C)atomic mass

D)number of neutrons

E)atomic number

Q2) Radioactive seeds that are implanted into a tumor are coated with ________ to stop alpha and beta ray penetration.

Q3) The half-life of ²¹⁸Po is 3.1 minutes. How much of a 155 gram sample remains after 0.40 hours?

A)0.00067 g

B)0.0072 g

C)0.72 g

D)0.0047 g

E)none of the above

Q4) What was the purpose of the Manhattan project?

Q5) The SI unit of an absorbed dose of radiation is the gray.

A)True

B)False

Q6) The use of radioisotopes in tracing metabolism is possible because ________.

Page 23

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Chapter 22: Chemistry of the Nonmetals

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Q1) Tetraboric acid, H₂B₄O₇, is prepared by heating boric acid, H₃BO₃ (a condensation reaction involving water loss). If 400.0 mmol H₃BO₃ are used, what mass (g)of H₂O is formed, assuming quantitative stoichiometric conversion?

A)5.77

B)0.500

C)0.320

D)7.21

E)9.01

Q2) Chlorine can have a positive oxidation state ________.

A)if it combines with bromine or iodine

B)if it combines with oxygen or fluorine

C)if it combines with hydrogen

D)if it combines with an alkali metal

E)in its elemental form

Q3) The danger from mixing ammonia with bleach is the production of ________.

Q4) What sulfur gas is used to sterilize wine?

Q5) Ozone is a pale blue poisonous gas with an irritating odor.

A)True

B)False

Page 24

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Chapter 23: Transition Metals and Coordination Chemistry

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Q1) Using the following abbreviated spectrochemical series, determine which complex ion is most likely to absorb light in the red region of the visible spectrum. small splitting Cl^- < H₂O < NH₃ < CN^- large splitting

A)[CuCl₄]^2-

B)[Cu(H₂O)₄]²⁺

C)[Cu(NH₃)₄]²⁺

D)[Cu(CN)₄]^2-

E)not enough information given to determine

Q2) The ligand with the formula Br^- is named ________ in complexes with transition metals.

A)bromo

B)bromide

C)fluoro

D)carbonato

E)azido

Q3) The two more common oxidation states of chromium are ________ and ________.

Q4) How can high-spin and low-spin transition metal complexes be distinguished from each other?

Q5) What is a siderophore?

Page 25

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Chapter 24: The Chemistry of Life: Organic and Biological Chemistry

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Q1) Ethers can be made by condensation of two ________ molecules by splitting out a molecule of water.

A)alkyne

B)alcohol

C)ketone

D)aldehyde

E)olefin

Q2) Starch, glycogen, and cellulose are made of repeating units of ________.

A)lactose

B)glucose

C)fructose

D)sucrose

E)amino acids

Q3) The stability of benzene is a major function of delocalized bonding.

A)True

B)False

Q4) Predict the product of the catalytic hydrogenation of 6-ethyl-3-decene.

Q5) The condensation reaction of a carboxyl group of one amino acid and the amino group of a second amino acid results in the formation of a(n)________.

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