Fundamentals of Chemistry Exam Practice Tests - 3207 Verified Questions

Fundamentals of Chemistry Exam Practice Tests

Course Introduction

Fundamentals of Chemistry provides an introduction to the basic principles underlying chemical science, including atomic and molecular structure, chemical bonding, stoichiometry, states of matter, and chemical reactions. The course explores key concepts such as the periodic table, chemical nomenclature, thermochemistry, and the properties of solutions. Through a combination of lectures, laboratory experiments, and problem-solving exercises, students develop critical thinking and analytical skills essential for understanding the composition, properties, and transformations of matter, laying the foundation for further study in chemistry and related scientific fields.

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Chemistry A Molecular Approach 2nd Canadian Edition by Nivaldo J. Tro

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Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) How many litres of wine can be held in a wine barrel whose capacity is 22.0 gal? \[( 1 \mathrm { gal } = 4 \mathrm { qt } = 3.7854 \mathrm {~L} )\]

A)1.72 × \(10 ^ { - 4 }\)

B)0.172

C)83.3

D)5.81 × \(10 ^ { 3 }\)

E)5.81

Answer: C

Q2) If a car gets 43.0 miles per gallon (mpg)fuel efficiency,calculate the fuel efficiency in litres per 100 km.(1 US gal = 3.78541 L; 1 mile = 1609.34 m)

A)5.47 L (100 km)^-1

B)8.17 L (100 km)^-1

C)3.19 L (100 km)^-1

D)9.40 L (100 km)^-1

E)4.43 L (100 km)^-1

Answer: A

Q3) Define matter.

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

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Chapter 2: Atoms and Elements

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Q1) Which of the following is an example of the law of multiple proportions?

A)A sample of chlorine is found to contain three times as much Cl-35 as Cl-37.

B)Two different compounds formed from carbon and oxygen have the following mass ratios: 1.33 g O:1 g C and 2.66 g O:1 g C.

C)Two different samples of table salt are found to have the same ratio of sodium to chlorine.

D)The atomic mass of bromine is found to be 79.90 amu.

E)Nitrogen dioxide always has a mass ratio of 2.28 g O:1 g N.

Answer: B

Q2) What mass (in kg)does 5.84 moles of titanium (Ti)have?

A)0.352 kg

B)0.122 kg

C)0.820 kg

D)0.280 kg

E)0.632 kg

Answer: D

Q3) The number 6.022 × 10²³ is known as ________.

Answer: Avogadro's number

Q4) The atomic number is equal to the number of ________.

Answer: protons

Page 4

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Chapter 3: Molecules,compounds,and Nomenclature

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Sample Questions

Q1) Which of the following is an acid in aqueous solution?

A)CaCO₃

B)HClO₂

C)CH₃OCH₃

D)NaCl

E)KNO₃

Answer: B

Q2) How many anions are there in 2.50 g of MgBr₂?

A)8.18 × 10²¹ anions

B)1.64 × 10²² anions

C)4.43 × 10²⁵ anions

D)8.87 × 10²⁵ anions

Answer: B

Q3) Describe the difference between a molecular formula and an empirical formula.Give an example.

Answer: A molecular formula is the exact number of each type of atom necessary to build a specific molecule.An empirical formula is simply the smallest whole number ratio between atoms in a compound.For example,C₂H₄ is the molecular formula for ethene.The empirical formula for ethene is CH₂,the smallest whole number ratio between the elements.

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Chapter 4: Chemical Reactions and Stoichiometry

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Q1) Define a spectator ion.

Q2) How many molecules of sucrose (C₁₂H₂₂O₁₁,molar mass = 342.30 g mol^-1)are contained in 14.3 mL of 0.140 mol L^-1 sucrose solution?

A)8.29 × 10²² molecules C₁₂H₂₂O₁₁

B)1.21 × 10²¹ molecules C₁₂H₂₂O₁₁

C)6.15 × 10²² molecules C₁₂H₂₂O₁₁

D)1.63 × 10²³ molecules C₁₂H₂₂O₁₁

E)5.90× 10²⁴ molecules C₁₂H₂₂O₁₁

Q3) What element is undergoing reduction (if any)in the following reaction? Zn(s)+ 2AgNO₃(aq) Zn(NO₃)₂(aq)+ 2Ag(s)

A)Zn

B)N

C)O

D)Ag

E)This is not an oxidation-reduction reaction.

Q4) Define an electrolyte.

Q5) How can you tell if a reaction is an oxidation-reduction reaction?

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Chapter 5: Gases

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Q1) How many molecules of N₂ are in a 400.0 mL container at 1.0399 bar and 135 °C?

A)7.01 × 10^{21 molecules}

B)7.38 × 10^{21 molecules}

C)2.12 × 10^{22 molecules}

D)2.23 × 10^{22 molecules}

Q2) A syringe contains 0.65 moles of He gas that occupy 750.0 mL.What volume (in L)of gas will the syringe hold if 0.35 moles of Ne is added?

A)0.87 L

B)4.9 L

C)1.2 L

D)2.1 L

E)1.9 L

Q3) What is the Celsius temperature of 100.0 g of chlorine gas in a 55.0 L container at 1.069 bar?

A)-23°C

B)228°C

C)250°C

D)500°C

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Chapter 6: Thermochemistry

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Q1) Which of the following processes is exothermic?

A)the formation of dew in the morning

B)the melting of ice

C)the chemical reaction in a "cold pack" often used to treat injuries

D)the vaporization of water

E)None of the above is exothermic.

Q2) At 1.0133 bar,the heat of sublimation of gallium is 277 kJ \(\mathrm { mol } ^ { - 1 }\) and the heat of vaporization is 271 kJ \(\mathrm { mol } ^ { - 1 }\) .To the correct number of significant figures,how much heat is required to melt 5.50 mol of gallium at 1.0133 bar?

A)6 kJ

B)33 kJ

C)244 kJ

D)274 kJ

Q3) The standard state of sulfur is ________.

Q4) Explain the difference between H and U.

Q5) Define chemical energy.

Q6) Give the temperature and pressure for the standard state for a liquid.

Q7) The standard state of carbon is ________.

Q8) Where does the energy absorbed during an endothermic reaction go?

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Q1) An electron ends in orbital n = 4 after a hydrogen atom emits a photon with a wavelength of 2165 nm.Calculate the initial orbital,n_i.

A)6

B)7

C)1

D)3

E)2

Q2) Calculate the energy of the violet light emitted by a hydrogen atom with a wavelength of 410 nm.

A)4.84 × 10^-19 J

B)2.06 × 10^-19 J

C)1.23 × 10^-19 J

D)8.13 × 10^-19 J

E)5.27 × 10^-19 J

Q3) How many orbitals are contained in the n = 2 level? Give thel and m_l values of each of them.

Q4) How are 3p orbitals different from the 2p orbitals?

Q5) Give an example of a p orbital.

Q6) What is the photoelectric effect?

Page 9

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) How many valence electrons do the noble gases possess?

A)1

B)3

C)7

D)6

E)8

Q2) Using the periodic table,identify the element with the following exception electron configuration: [Kr]5s¹4d⁵

A)Zr

B)Ni

C)Re

D)Tc

E)Mo

Q3) Identify the number of core electrons for O.

A)0

B)1

C)2

D)3

E)4

Q4) Why do successive ionization energies increase?

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) A triple covalent bond contains ________ of electrons.

A)0 pairs

B)1 pair

C)2 pairs

D)3 pairs

E)4 pairs

Q2) Identify the number of valence electrons for XeI₂.

A)22

B)20

C)18

D)24

Q3) Which element can act as the centre atom in a hypercoordinate compound?

A)N

B)O

C)Br

D)He

Q4) Describe a covalent bond.

Q5) Define bond energy.

Q6) Define formal charge.

Q7) How are electron affinity and electronegativity different?

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Chapter 10: Chemical Bonding II: Molecular Shapes, valence

Bond Theory, and Molecular Orbital Theory

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Q1) Place the following in order of decreasing dipole moment. I.cis-CHCl=CHCl \(\quad\)II.trans-CHCl=CHCI \(\quad\)III.cis-CHF=CHF

A)III > I > II

B)II > I > III

C)I > III > II

D)II > III > I

E)I = III > II

Q2) List the number of sigma bonds and pi bonds in a single bond.

A)1 sigma,0 pi

B)0 sigma,1 pi

C)1 sigma,1 pi

D)1 sigma,2 pi

Q3) Using the VSEPR model,the electron-domain geometry of the central atom in BrF₄^- is ________.

A)tetrahedral

B)trigonal bipyramidal

C)trigonal planar

D)octahedral

E)linear

Q4) Explain why oil and water do not mix.

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Chapter 11: Liquids, solids, and Intermolecular Forces

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Q1) Place the following substances in order of decreasing boiling point. N₂ \(\quad\)O₂ \(\quad\)H₂

A)O₂ > H₂> N₂

B)N₂ > H₂ > O₂

C)N₂ > O₂ > H₂

D)O₂ > N₂ > H₂

E)H₂ > N₂ > O₂

Q2) Place the following substances in order of increasing boiling point. Ne \(\quad\)Cl₂\(\quad\)O₂

A)Ne < Cl₂< O₂

B)Cl₂ < O₂< Ne

C)O₂ < Cl₂< Ne

D)Cl₂ < Ne < O₂

E)Ne < O₂< Cl₂

Q3) In liquid ethanol, CH₃CH₂OH

Which intermolecular forces are present?

A)Dispersion,hydrogen bonding and dipole-dipole forces are present.

B)Only dipole-dipole and ion-dipole forces are present.

C)Only dispersion and dipole-dipole forces are present.

D)Only hydrogen bonding forces are present.

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Chapter 12: Solutions

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Q1) A 4.55 L sample of water contains 0.115 g of sodium ions.Determine the concentration of sodium ions in ppm if the density of the solution is 1.00 g mL^-1.

A)52.3 ppm

B)13.2 ppm

C)12.7 ppm

D)25.3 ppm

E)36.5 ppm

Q2) What is the weight percent of a caffeine solution made by dissolving 8.35 g of caffeine,C₈H₁₀N₄O₂,in 75 g of benzene,C₆H₆?

A)0.010%

B)0.011%

C)10%

D)11%

Q3) Calculate the osmotic pressure,in bar,of a solution containing 3.00 mg of a sugar (342 g mol^-1)in 15.0 mL of water at 25 °C.

Q4) Explain why water does not dissolve in gasoline.

Q5) Define the Tyndall effect.

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Chapter 13: Chemical Kinetics

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Q1) A particular first-order reaction has a rate constant of 1.35 × \(10 ^ { 2 }\) \(s ^ { - 1 }\) at 25.0 °C.What is the magnitude of k at \( 95^{\circ} \mathrm{C} \) if \( E_{\mathrm{a}}=55.5 \mathrm{~kJ} \mathrm{~mol}-1 ? \)

A)9.57 \(\times 10 ^ { 3 } \mathrm {~s} ^ { - 1 }\)

B)3.74 ×10⁸⁷ \(s ^ { - 1 }\)

C)1.36 × \(10 ^ { 2 }\) \(s ^ { - 1 }\)

D)575 \(s ^ { - 1 }\)

E)1.05 × \(10 ^ { - 4 }\) \(s ^ { - 1 }\)

Q2) Calculate the activation energy,E_a,for a reaction with a frequency factor,A,of 9.3 × 10¹⁰ s^-1 and a rate constant of 1.06 × 10³ s^-1 at 273 K.

A)42 kJ mol^-1

B)45 kJ mol^-1

C)36 kJ mol^-1

D)49 kJ mol^-1

E)58 kJ mol^-1

Q3) What is a catalyst and what function does it serve?

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Chapter 14: Chemical Equilibrium

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Q1) The reaction below has a K_c value of 1.0 × 10¹²M^-1.What is the value of K_p for this reaction at 500 K? \(\quad\)2SO₂(g)+ O₂(g) 2SO₃(g)

A)4.2 × 10^-11bar^-1

B)1.0 × 10¹²bar^-1

C)2.4 × 10^-12bar^-1

D)4.1 × 10¹³bar^-1

E)2.4 × 10¹⁰bar^-1

Q2) Consider the following reaction:

Xe(g)+ 2F₂(g) XeF₄(g)

A reaction mixture initially contains 2.24 bar Xe and 4.27 bar F₂.If the equilibrium pressure of Xe is 0.34 bar,find the equilibrium constant (K_p)for the reaction.

A)25 bar^-2

B)0.12 bar^-2

C)0.99 bar^-2

D)8.3 bar^-2

E)0.040 bar^-2

Q3) Explain dynamic equilibrium.Use the generic reaction A(g) B(g)to explain.

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Chapter 15: Acids and Bases

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Q1) Which of the following is the correct Arrhenius definition of bases?

A)A base is a substance that produces H⁺ ions in aqueous solutions.

B)A base is a substance that reacts with OH^- ions in aqueous solutions.

C)A base is a substance that decreases the amount of OH^- ions in aqueous solutions.

D)A base is a substance that produces OH^- ions in aqueous solutions.

E)A base is a substance that does not change the concentration of OH^- ions in aqueous solutions.

Q2) Determine the pH of a 0.00598 mol L^-1 HClO₄ solution.

A)1.777

B)6.434

C)7.566

D)2.223

E)3.558

Q3) What does the term amphoteric mean?

Q4) Do both protons ionize instantaneously from a diprotic acid such as H₂CO₃? Explain your answer.

Q5) Describe a molecule that can be a Lewis acid.

Q6) What is the autoionization of water?

Page 17

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Chapter 16: Aqueous Ionic Equilibrium

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Q1) Calculate the pH of a buffer that is 0.225 mol L^-¹ CH₃COOH and 0.162 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) In which of the following solutions would solid PbBr₂ be expected to be the least soluble at 25 °C?

A)0.1 mol L^-1 HBr

B)0.1 mol L^-1 NaBr

C)0.1 mol L^-1 CaBr₂

D)0.1 mol L^-1 KNO₃

Q3) Describe the solubility of Al(OH)₃ with respect to pH.

A)soluble at low pH,insoluble in pH-neutral solution,and soluble at high pH

B)soluble at low pH,insoluble in pH-neutral solution,and insoluble at high pH

C)insoluble at low pH,insoluble in pH-neutral solution,and soluble at high pH

D)soluble at low pH,in pH-neutral solution,and at high pH

E)insoluble at low pH,in pH-neutral solution,and at high pH

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Chapter 17: Gibbs Energy and Thermodynamics

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Q1) Calculate the equilibrium constant for the following reaction at 298 K: 2H₂S(g) 2H₂(g)+ S₂(g)\(\quad\)\(\quad\) _rH° = 169.8 kJ mol^-1; _rS°= 78.0 J K^-1 mol^-1

A)3.28 × 10^-21

B)4.99 × 10^-29

C)1.05 × 10^-25

D)2.06 × 10^-26

E)6.38 × 10^-24

Q2) Define entropy.

Q3) In which of the following processes do the molecules become more ordered?

A)water freezing

B)ice melting

C)water evaporating

D)salt dissolving in water

E)dry ice subliming

Q4) Define the second law of thermodynamics.

Q5) Why is the quantity of energy required to recharge a battery greater than the quantity of work done?

Page 19

Q6) What is "free" energy? Give a fictitious example.

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Chapter 18: Electrochemistry

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Q1) Calculate the cell potential for the following unbalanced reaction that takes place in an electrochemical cell at 25 °C when [ \(\mathrm { Mg } ^ { 2 + }\) ] = 0.914 M and [ \(\mathrm { Fe } ^ { 3 + }\) ] = 0.0230 M \(\quad\)\(\quad\)Mg(s)+ \(\mathrm { Fe } ^ { 3 + }\) (aq)? \(\mathrm { Mg } ^ { 2 + }\) (aq)+ Fe(s)

E°(Mg²⁺/Mg)= -2.37 V and E°(Fe³⁺/Fe)= -0.036 V

A)+2.32 V

B)+2.30 V

C)-2.32 V

D)-2.30 V

E)+1.23 V

Q2) Which of the following is the strongest oxidizing agent?

A)MnO₂(s)

B)Cl^- (aq)

C)Cu⁺(aq)

D)SO₄^2-(aq)

E)MnO₄^- (aq)

Q3) Why are iron nails coated with zinc?

Q4) Why,if we multiply a reaction by 2,don't we multiply its E°_red by 2?

Q5) What is electrolysis?

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Chapter 19: Radioactivity and Nuclear Chemistry

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Q1) The half-life of cobalt-60 is 5.20 y.How many milligrams of a 6.000-mg sample remain after 10.5 years?

A)5.75 mg

B)24.3 mg

C)0.0543 mg

D)0.123 mg

E)1.48 mg

Q2) Determine how many neutrons are produced during the spontaneous fission of 244 95 Am to form I-134 and Mo-107.

A)0

B)1

C)2

D)3

E)4

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) Identify the generic formula for alkanes.

A)C_nH_2n+2

B)C_nH_2n_-₂

C)C_nH_2n

D)C_nH_2n_-4

E)C_nH₂_n₊₄

Q2) Explain the concept of bond polarization and inductive effect using organic halides as an example.

Q3) Which of the following pairs of organic compound families could have constitutional isomers?

A)an ether and an alcohol

B)an amide and an amine

C)an ether and a carboxylic acid

D)a carboxylic acid and an amide

E)an amine and an alcohol

Q4) Which one of the following molecules is the most polar?

A)butane

B)butanoic acid

C)cyclohexane

D)ethanol

Page 22

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Chapter 21: Organic Chemistry II: Reactions

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Q1) Give the organic product for the following reaction. CH₃CH₂CH₂NH₂ + HCl A)ClCH₂CH₂CH₂NH₂ B)CH₃CHClCH₂NH₂ C)CH₃CH₂CHClNH₂ D)CH₃CH₂CH₂NHCl E)CH₃CH₂CH₂NH₃⁺Cl<su p>-</sup>

Q2) Identify the organic product for the following reaction. \(\mathrm { CH } _ { 3 } \mathrm { CH } _ { 2 } \mathrm { COOH } + \mathrm { CH } _ { 3 } \mathrm { CH } _ { 2 } \mathrm { CH } _ { 2 } \mathrm { NH } _ { 2 } \stackrel { \text { with } \mathrm { SOCl } _ { 2 } } { \longrightarrow }\)

A)CH₃CH₂NHCH₂CH₂CH<sub>3</sub >

B)CH₃CH₂CON(CH₂CH₂CH<sub>3</s ub>)₂

C)CH₃CH₂CH₂CONHCH₂CH<sub>3</s ub>

D)CH₃CH₂CONHCH₂CH₂CH<sub>3</s ub>

E)CH₃CH₂N(CH₂CH₂CH₃ )₂

Page 23

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Chapter 22: Biochemistry

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Q1) In the sequence Arg-Pro-Leu-Gln,which amino acid contains the C-terminal group?

A)proline

B)leucine

C)glutamine

D)arginine

Q2) Which of the following form between glycerol and fatty acids to form triglycerides?

A)hydrogen bonds

B)disulfide linkages

C)peptide bonds

D)glycosidic linkages

E)ester linkages

Q3) Which of the following is an example of a secondary protein structure?

A)dipeptide

B)triglyceride

C)a-helix

D)amino acid

E)fatty acid

Q4) Why does the melting point of fatty acids decrease with the number of double bonds present?

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Chapter 23: Chemistry of the Nonmetals

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Q1) Refer to the above table.Predict the silicate structure for KAl(SiO₃)₂.

A)phyllosilicate

B)orthosilicate

C)amphibole

D)pyroxene

E)pyrosilicate

Q2) Determine the number of vertices and faces in the closo-Borane B₇H₇^2-.

A)vertices = 4,faces = 10

B)vertices = 7,faces = 14

C)vertices = 3,faces = 6

D)vertices = 14,faces = 24

E)vertices = 7,faces = 10

Q3) What is the oxidation number of oxygen in Na₂O?

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Chapter 24: Metals and Metallurgy

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Q1) Calcination is

A)heating an ore in the presence of oxygen or another substance to cause a chemical reaction that drives off newly formed volatile compounds.

B)heating an ore in the presence of oxygen or another substance to purify the ore and obtain the liquid metal.

C)selectively dissolving a metal in solution to separate it from its ore.

D)the process of heating an ore to drive off volatile compounds.

E)forming metal parts using heat and small crystals of metal.

Q2) Roasting is

A)forming metal parts using heat and small crystals of metal.

B)the process of heating an ore to drive off volatile compounds.

C)selectively dissolving a metal in solution to separate it from its ore.

D)heating an ore in the presence of oxygen or another substance to purify the ore and obtain the liquid metal.

E)heating an ore in the presence of oxygen or another substance to cause a chemical reaction that drives off newly formed volatile compounds.

Q3) Why does a "two-phase" structure occur in a substitutional alloy?

Q4) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Source URL: https://quizplus.com/quiz/68959

Sample Questions

Q1) The complex ion [Ni(NH₃)₆]²⁺ has a maximum absorption near 580 nm.Calculate the crystal field splitting energy (in kJ mol^-1)for this ion.

A)114 kJ mol^-1

B)292 kJ mol^-1

C)343 kJ mol^-1

D)206 kJ mol^-1

E)485 kJ mol^-1

Q2) Determine the chemical formula for the compound diamminetetraaquairon(II)chloride.

A)[Fe(NH₃)₂(H₂O)₄Cl]

B)[Fe(NH₃)₂][(H₂O)₄Cl]

C)[Fe(NH₃)₂(H₂O)₄]Cl₂

D)[Fe(H₂O)₄][(NH₃)₂Cl]

E)[Fe(NH₃)₂(H₂O)₄]Cl₃

Q3) What is a coordinate covalent bond?

Q4) Why is +2 a common oxidation state for transition elements?

Q5) What is the difference between a weak-field complex and a strong-field complex?

Q6) Explain how EDTA is used to treat lead poisoning.

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