Fundamentals of Chemistry Solved Exam Questions - 3326 Verified Questions

Fundamentals of Chemistry Solved Exam Questions

Course

Introduction

Fundamentals of Chemistry introduces students to the basic principles and concepts that underpin the science of chemistry. The course covers atomic structure, chemical bonding, states of matter, chemical reactions, stoichiometry, and the properties of solutions. Emphasis is placed on understanding the scientific method, applying mathematical problem-solving techniques, and developing laboratory skills through hands-on experiments. This foundational course prepares students for further studies in chemistry and related scientific disciplines by building a solid conceptual framework and practical experience.

Recommended Textbook

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

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25 Chapters

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Page 2

Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) mega

A)10^-6

B)10^-⁹

C)10^-1

D)10^-3

E)10³

F)10^-2

G)10⁶

Answer: G

Q2) If 1.4% of the mass of a human body is calcium, how many kilograms of calcium are there in a 185-pound man?

A)1.2 kg Ca

B)5.7 kg Ca

C)1.2 × 10² kg Ca

D)5.7 × 10² kg

Answer: A

Q3) Define matter.

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

Q4) Define the law of the conservation of energy.

Answer: Energy is neither created nor destroyed, only transferred.

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following elements is a nonmetal?

A)Zn

B)Cs

C)Ca

D)Co

E)P

Answer: E

Q2) What is the chemical symbol for copper?

A)Co

B)Cr

C)Cu

D)C

Answer: C

Q3) Which of the following elements is classified as a semimetal?

A)calcium

B)boron

C)fluorine

D)uranium

Answer: B

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Page 4

Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) Calculate the number of oxygen atoms in 1.187 g of nickel(II)carbonate.

A)1.807 × 10²⁶

B)5.535 × 10²³

C)6.022 × 10²¹

D)1.282 × 10²²

E)1.807 × 10²²

Answer: E

Q2) How many millimoles of Ca(NO₃)₂ contain 4.78 × 10²² formula units of Ca(NO₃)₂? The molar mass of Ca(NO₃)₂ is 164.10 g mol^-1.

A)12.6 mmol Ca(NO₃)₂

B)13.0 mmol Ca(NO₃)₂

C)20.7 mmol Ca(NO₃)₂

D)79.4 mmol Ca(NO₃)₂

E)57.0 mmol Ca(NO₃)₂

Answer: D

Q3) Describe a structural formula

Answer: In a structural formula, lines are used to represent covalent bonds to show how the atoms in the molecule are connected to each other.

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Page 5

Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) Based on the balanced chemical equation shown below, what volume of 0.250 mol L^-1 K₂S₂O₃(aq)is needed to completely react with 24.88 mL of 0.125 mol L^-1 KI₃(aq), according to the following chemical equation: 2S₂O₃^2-(aq)+ I₃^-(aq) S₄O₆^2-(aq)+ 3I^-(aq).

A)6.22 mL

B)12.4 mL

C)24.9 mL

D)99.5 mL

Q2) Which one of the following compounds is insoluble in water?

A)CaCl₂

B)NaNO₃

C)PbCl₂

D)K₂CO₃

Q3) Which pair of compounds is soluble in water?

A)HgBr₂ and Hg₂I₂

B)CuS and Na₂S

C)LiI and Ca(NO₃)₂

D)KNO₃ and CaCO₃

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Chapter 5: Gases

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Sample Questions

Q1) Explain what the term "mean free path" describes. How does it change with decreasing pressure?

Q2) What is the volume of 5.60 g of O₂ at 7.78 bar and 415 K?

A)1.53 L

B)565 L

C)24.5 L

D)25.0 L

E)0.776 L

Q3) A mixture of 1.0 mol He and 1.0 mol Ne are at STP in a rigid container. Which of the following statements is TRUE?

A)Both gases have the same average kinetic energy.

B)Both gases contribute equally to the density of the mixture under these conditions.

C)Both gases have the same molecular speed.

D)The mixture has a volume of 22.4 L.

E)All of the above are true.

Q4) Why does hot air rise?

Q5) Define pressure.

Q6) Give the major gas in dry air.

Q7) Define effusion.

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) A 5.00 g sample of copper metal at 25.0 °C is heated by the addition of 96.0 J of energy. The final temperature of the copper is ________°C. The specific heat capacity of copper is 0.38 J g^-1 K^-1.

A)32.3

B)25.0

C)25.5

D)75.5

E)50.5

Q2) 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.

Q3) Why are the standard heats of formation for elements in their most stable form assigned a value of 0?

Q4) 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.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) An electron initially in the n = 2 orbital of a hydrogen atom absorbs a photon with a wavelength of 656 nm. Calculate the final orbital, n_f.

A)3

B)5

C)2

D)4

E)6

Q2) Calculate the wavelength of light emitted from a hydrogen atom when an electron undergoes an n = 8 to n = 4 transmission.

A)1940 nm

B)2620 nm

C)1.245 × 10³ nm

D)3.178 × 10³ nm

E)714.8 nm

Q3) What value of l is represented by a d orbital? A)1 B)2 C)0 D)3

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Which reaction below represents the second ionization of Sr?

A)Sr(g) Sr (g)+ e

B)Sr² (g)+ e Sr (g)

C)Sr (g)+ e Sr(g)

D)Sr(g)+ e Sr² (g)

E)Sr (g) Sr² (g)+ e

Q2) Using the periodic table, identify the element with the following exception electron configuration: [Xe]6s¹5d¹⁰

A)Hg

B)Au

C)Cd

D)Ir

E)Rg

Q3) How many valence electrons does an atom of Ti possess?

A)2

B)4

C)6

D)8

E)0

Q4) Define ionization energy.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Identify the complete electronic configuration for Ca²⁺.

A)1s²2s²2p⁶3s²4p⁶

B)1s²2s²2p⁶3s²3p⁶

C)1s²2s²2p⁶3s²3p⁵

D)1s²2s²3p⁶4s²5p⁶

E)1s²2s²p⁶3s²p⁶

Q2) Ca-O

A)longest covalent bond

B)metallic bond

C)strongest covalent bond

D)weakest ionic bond

E)highest melting point

Q3) Which compound has the longest carbon-carbon bond length?

A)CH₃CH₃

B)CH₂CH₂

C)HCCH

D)All bond lengths are the same.

Q4) Define dipole moment.

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

Q6) Define bond energy.

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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) Give the electron geometry (eg), molecular geometry (mg), and hybridization for CH₃^-.

A)eg = tetrahedral; mg = trigonal pyramidal; sp³

B)eg = tetrahedral; mg = tetrahedral; sp³

C)eg = trigonal pyramidal; mg = trigonal pyramidal; sp³

D)eg = trigonal planar; mg = trigonal planar; sp²

Q2) Determine the molecular geometry for the molecule PF₃.

A)Trigonal pyramidal

B)Trigonal planar

C)Tetrahedral

D)T-shaped

E)Bent

Q3) Using the VSEPR model, the molecular geometry of the central atom in CF₄ is ________.

A)linear

B)trigonal planar

C)tetrahedral

D)bent

E)trigonal pyramidal

Q4) According to molecular orbital theory, what is an antibonding orbital?

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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) CH₃CH₃

A)dispersion forces

B)dipole-dipole forces

C)ion-dipole forces

D)hydrogen bonding

E)ionic bond

F)H₂ + H₂O

Q2) How many H^- ions are around each Na⁺ ion in NaH, which has a cubic unit cell with H^- ions on each corner and each face?

A)1

B)4

C)6

D)8

Q3) Which of the following compounds exhibits hydrogen bonding?

A)CH₃I

B)H₂

C)C₆H₅OC

D)CH₃CH

Q4) Give the edge length in terms of r for a simple cubic cell.

Q5) Describe the difference between the conduction band and the valence band.

Page 13

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Chapter 12: Solutions

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Sample Questions

Q1) A solution is prepared by dissolving 16.2 g of benzene (C₆H₆)in 282 g of carbon tetrachloride (CCl₄)The concentration of benzene in this solution is ________ molal. The molar masses of C₆H₆ and CCl₄ are 78.1 g mol^-1 and 154 g mol^-1, respectively.

A)7.36 × 10^-4

B)0.736

C)0.102

D)0.0543

E)5.43

Q2) 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

Q3) Define the Tyndall effect.

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) The first-order reaction, 2N₂O(g) 2N₂(g)+ O₂(g), has a rate constant equal to 0.76 s^-1 at 1000 K. How long will it take for the concentration of N₂O to decrease to 12% of its initial concentration?

A)0.62 s

B)2.8 s

C)6.3 s

D)8.4 s

Q2) A reaction with an activation energy of 34.0 kJ mol^-1 has a rate constant of 26.8 s^-1 at 354 K. What is the rate constant at 314 K?

A)6.15 s^-1

B)2.08 s^-1

C)7.68 s^-1

D)9.80 s^-1

E)4.00 s^-1

Q3) Is the activation energy for a forward reaction the same as the activation energy for the reverse of the same reaction? Use a sketch of a reaction energy diagram.

Q4) Define the frequency factor.

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Page 15

Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Which of the following statements is TRUE?

A)If Q < K, it means the reverse reaction will proceed to form more reactants.

B)If Q > K, it means the forward reaction will proceed to form more products.

C)If Q = K, it means the reaction is at equilibrium.

D)If Q << K, it means the reaction is shifted to the left.

E)If Q >> K, it means the reaction is shifted to the right.

Q2) In a reaction mixture containing reactants and products, each at a concentration of 1 mol L^-1, what is the value of Q?

A)-1

B)1

C)

D)0

E)Q = K

Q3) Give the direction of the reaction if K >> 1.

A)The forward reaction is favoured.

B)The reverse reaction is favoured.

C)Neither direction is favoured.

D)If the temperature is raised, then the forward reaction is favoured.

E)If the temperature is raised, then the reverse reaction is favoured.

Q4) 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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Sample Questions

Q1) Which of the following is a strong base?

A)Cl^-

B)NH₃

C)CH₃OH

D)NO₃

E)KOH

Q2) Determine the pH of a 0.62 mol L^-1 NH₄NO₃ solution at 25 °C. The K_b for NH₃ is 1.76 × 10^-5.

A)2.48

B)9.27

C)11.52

D)4.73

E)9.45

Q3) Determine the pH of a 0.023 mol L^-1 HNO₃ solution.

A)12.36

Q4) 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) Describe the information necessary (and why)to choose a good indicator for a weak acid-strong base titration.

Q2) Which of the following compounds will have the highest molar solubility in pure water?

A)PbSO₄,K_sp = 1.82 × 10^-8

B)MgCO₃,K_sp = 6.82 × 10^-6

C)AgI, K_sp = 8.51 × 10^-17

D)PbS, K_sp = 9.04 × 10^-29

E)FeS, K_sp = 3.72 × 10^-19

Q3) What is the pH of a solution made by mixing 25.00 mL of 0.100 mol L^-1

HCl with 40.00 mL of 0.100 mol L^-1 KOH? Assume that the volumes of the solutions are additive.

A)0.64

B)1.64

C)12.36

D)13.36

Q4) Sketch the titration curve for a monoprotic weak acid titrated with a strong base. Make sure to indicate the equivalence point (and whether it is acidic, basic, or neutral)and the buffer region.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Place the following in order of decreasing standard molar entropy. NaCl(s)Na₃PO₄(aq)NaCl(aq)

A)NaCl(s)> NaCl(aq)> Na₃PO₄(aq) B)NaCl(aq)> NaCl(s)> Na₃PO₄(aq) C)Na₃PO₄(aq)> NaCl(aq)> NaCl(s) D)NaCl(s)> Na₃PO₄(aq)> NaCl(aq) E)NaCl(aq)> Na₃PO₄(aq)> NaCl(s)

Q2) Use Hess's law to calculate _rG° using the following information: NO(g)+ O(g) NO₂(g) _rG°= ?

2O₃(g) 3O₂(g) _rG°= +489.6 kJ mol^-1

O₂(g) 2O(g) _rG°= +463.4 kJ mol^-1 NO(g)+ O₃(g) NO₂(g)+ O₂(g) _rG°= - 199.5 kJ mol^-1

A)+753.5 kJ mol^-1

B)+277.0 kJ mol^-1

C)-676.0 kJ mol^-1

D)-1152.5 kJ mol^-1

E)-225.7 kJ mol^-1

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Page 19

Chapter 18: Electrochemistry

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Sample Questions

Q1) Determine which of the following pairs of reactants will result in a spontaneous reaction at 25 °C.

A)I^-(aq)+

Zn²⁺(aq); if E^o(I₂/I^-)= + 0.54V and E^o(Zn²⁺/Zn)= -0.76V

B)Ca(s)+ Mg²⁺(aq); if E^o(Ca²⁺/Ca)= -2.76V and E^o(Mg²⁺/Mg)= -2.37V

C)H₂(g)+ Cd²⁺(aq); if E^o(Cd²⁺/Cd)= -0.40V

D)Ag(s)+ Sn²⁺(aq); if E^o(Ag⁺/Ag)= + 0.80V and E^o(Sn²⁺/Sn)= -0.14V

E)Ag⁺(aq)+ Mn²⁺(aq); if E^o(Ag⁺/Ag)= + 0.80V and E^o(MnO₄^-/Mn²⁺)= +1.51V

Q2) Sketch a voltaic cell that contains the following half reactions and label all relevant components.

Zn²⁺(aq)+ 2e Zn(s) E°= -0.76 V Cu²⁺(aq)+ 2e Cu(s) E°= +0.34 V

Q3) Why are iron nails coated with zinc?

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Page 20

Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Which of the following nuclides are most likely to decay via beta decay?

A)I-131

B)Ar-40

C)F-18

D)Zr-90

E)Pb-206

Q2) If we start with 1.000 g of cobalt-60, 0.675 g will remain after 3.00 y. This means that the half-life of cobalt-60 is ________ y.

A)3.08

B)4.44

C)2.03

D)5.29

E)7.65

Q3) The combination of two light nuclei to form a heavier nuclei is called ________.

A)radioactive cleavage

B)nuclear fission

C)nuclear fusion

D)radioactive merge

E)the half-life

Q4) Describe what is meant by the "valley of stability."

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Chapter 20: Organic Chemistry I: Structures

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Sample Questions

Q1) What are constitutional isomers?

Q2) Which of the following compounds exhibits optical isomerism?

A)CH₃-CH₂-CH₃

B)CH₃-CH₂-CHBr-CH₃

C)CH₃-CHCl-CH₃

D)CH₃-CH₂-CH₂Br

E)CH₃-CH₂-CBr₂-CH₃

Q3) Identify the generic formula for alkenes.

A)C_nH_2n+2

B)C_nH_2n_-₂

C)C_nH_2n

D)C_nH_2n_-4

E)C_nH₂_n₊₄

Q4) Which compound is a saturated hydrocarbon?

A)1-butyne

B)cycloheptene

C)3-methylheptane

D)2-heptyne

Q5) What does the term "racemic mixture" mean?

Q6) What is the index of hydrogen deficiency?

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Chapter 21: Organic Chemistry II: Reactions

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Q1) What are primary alcohols oxidized to?

A)secondary alcohols

B)ethers

C)esters

D)aldehydes

E)ketones

Q2) Which of the following is the balanced chemical equation for the addition of two equivalents of Br₂ to CH₃C CH?

A)CH₃C CH + 2Br₂

CH₃CBr₂CHBr₂

B)CH₃C CH + 2Br₂ CH₃CBr=CHBr

C)CH₃C CH + 2Br₂

CH₃CH₂CHBr₂

D)CH₃C CH + 2Br₂

CH₃CBr₂CH₃

E)CH₃C CH + 2Br₂ CH₃CHBrCH₂Br

Q3) Why doesn't benzene typically undergo addition reactions like the alkenes do?

Q4) Which part of HCN, the hydrogen or the cyano group, adds to the O in a C=O bond and why?

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Chapter 22: Biochemistry

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Sample Questions

Q1) Identify the polysaccharide.

A)glucose

B)cholesterol

C)fructose

D)cellulose

E)sucrose

Q2) Which of the following statements is TRUE?

A)Monosaccharides are the building blocks of polypeptides.

B)Nucleotides are the building blocks of saccharides.

C)Fatty acids are responsible for DNA replication.

D)Peptides are the building blocks of amino acids.

E)Carbohydrates are water soluble and responsible for short-term energy storage.

Q3) Which of the following has a tertiary protein structure?

A)amino acid

B)fibrous protein

C)nucleotide

D)fatty acid

E)globular protein

Q4) What is a codon?

Q5) How do cis- and trans-fats differ?

24

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Chapter 23: Chemistry of the Nonmetals

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Sample Questions

Q1) Refer to the above table. Use charge balance to determine the value of x in the pyroxene Ca_x(SiO₃)₂.

A)2

B)1

C)4

D)6

E)3

Q2) Write the balanced equation that represents the formation of hydrofluoric acid from the reaction of calcium fluoride with sulfuric acid.

A)CaF₂ + H₂S CaS + 2HF

B)CF₄ + 2H₂S 4HF + CS₂

C)CF₄ + 2H₂SO₄ 4HF + CS₂ + CO₂

D)CaF₂ + H₂SO₄ 2HF + CaSO₄

E)CaF₂ + H₂SO₃ CaSO₃ + 2HF

Q3) Why are phosphate compounds added to detergents?

Q4) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q5) Describe the major production method for obtaining oxygen.

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Chapter 24: Metals and Metallury

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Sample Questions

Q1) An interstitial alloy contains carbon in one-fourth of the octahedral holes of a closest-packed metal, M. What is the formula of this alloy?

A)M₄C

B)MC₄

C)MC

D)M₂C

E)MC₂

Q2) Zn

A)sphalerite

B)rhodochrosite

C)rutile

D)carnotite

E)malachite

F)galena

G)cinnabar

Q3) What is the difference between ferromagnetism and paramagnetism?

Q4) Why does a "two-phase" structure occur in a substitutional alloy?

Q5) What is an advantage of hydrometallurgical processes over pyrometallurgical processes?

Q6) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Determine the chemical formula for the compound tris(ethylenediamine)chromium(III)chloride.

A)[Cr(NH₂CH₂CH₂NH₂)₃] Cl₃

B)[(NH₂CH₂CH₂NH₂)₃]Cr ³⁺

C)[Cr(NH₂CH₂CH₂NH₂)₃ ]³⁺

D)Cl₂[Cr(NH₂CH₂CH₂NH<sub>2</sub >)₄]

E)[Cr(NH₂CH₂CH₂NH₂)₂] Cl₂

Q2) How many moles of aqueous ions will be produced from the dissolution of 1.0 mole of K₃[FeCl₆] in water? A)4

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