General Chemistry I Exam Review - 3326 Verified Questions

General Chemistry I Exam Review

Course Introduction

General Chemistry I is an introductory course that explores the fundamental principles of chemistry, including atomic and molecular structure, periodic trends, stoichiometry, chemical reactions, thermochemistry, and the properties of gases, liquids, and solids. Students learn to apply quantitative reasoning to chemical problems, understand the nature of chemical bonding, and develop laboratory skills through hands-on experiments that reinforce theoretical concepts. This course lays the essential foundation for advanced studies in chemistry and related scientific disciplines.

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) What wavelength of light would you report in units of nm if the light had a wavelength of 7.60 × 10^-10 m?

A)7.60 × 10^-3 nm

B)7.60 × 10^-19nm

C)1.32 nm

D)0.760 nm

E)760 nm

Answer: D

Q2) An alligator is 244.0 cm long. How long is it in feet?

A)8.00 ft

B)96.0 ft

C)20.3 ft

D)6.00 ft

E)9.00 ft

Answer: A

Q3) Define matter.

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

Q4) 38.325 lbs = ________ grams. (1 lb = 454 g)

Answer: 17400

3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) How many neutrons are in arsenic-42?

A)33

B)41

C)42

D)41.9

E)75

Answer: C

Q2) Calculate the mass (in g)of 1.9 × 10²⁴ atoms of Pb.

A)3.9 × 10² g

B)2.4 × 10² g

C)3.2 × 10² g

D)1.5 × 10² g

E)6.5 × 10² g

Answer: E

Q3) Argon belongs to the ________ group of the periodic table.

A)alkali metal

B)alkaline earth metal

C)halogen

D)noble gas

Answer: D

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) What is the mass of 0.500 mol of dichlorodifluoromethane, CCl₂F₂?

A)4.14 × 10^-3 g

B)60.5 g

C)121 g

D)242 g

Answer: B

Q2) What is the empirical formula for C₄H₁₀O₂?

A)C₂H₅O

B)CHO

C)C₂H₄O

D)CHO₂

E)CH₂O

Answer: A

Q3) Which of the following elements has the LEAST tendency to form an ion?

A)Ca

B)K

C)Kr

D)Se

Answer: C

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Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) When 7.00 × 10²² molecules of ammonia react with 6.00 × 10²² molecules of oxygen according to the chemical equation shown below, how many grams of nitrogen gas are produced? 4NH₃(g)+ 3O₂(g) 2N₂(g)+ 6H₂O(g)

A)1.63 g

B)1.86 g

C)4.19 g

D)6.51 g

Q2) How many millilitres of 0.550 mol L^-1 hydriodic acid are needed to react with 15.00 mL of 0.217 mol L^-1 CsOH? HI(aq)+ CsOH(aq) CsI(aq)+ H₂O(l)

A)0.0263 mL

B)0.169 mL

C)5.92 mL

D)38.0 mL

Q3) Balance the following equation: _____ C₉H₂₀ + _____ O₂ _____ H₂O + _____ CO₂

Q4) What causes a precipitation reaction to occur between two soluble compounds?

Q5) Determine the oxidation state of Mn in KMnO₄.

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Chapter 5: Gases

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Sample Questions

Q1) Which of the following gases has the lowest average speed at 25 °C?

A)C₃H₈

B)Kr 

C)CH₃NH₂

D)SO₂

Q2) Calculate the temperature if 0.0139 moles of Ne has a pressure of 1.237 bar in a 250 mL container.

A)260 °C

B)61.8 °C

C)-23.8 °C

D)15.4 °C

E)-5.57 °C

Q3) A mixture of 0.220 moles CO, 0.350 moles H₂, and 0.640 moles He has a total pressure of 2.95 bar. What is the pressure of CO?

A)1.86 bar

B)0.649 bar

C)0.536 bar

D)1.54 bar

E)0.955 bar

Q4) Why does hot air rise?

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) According to the following thermochemical equation, what mass of H₂O (in g)must form to produce 975 kJ of energy? SiO₂(s)+ 4HF(g) SiF₄(g)+ 2H₂O(l) _rH°= -184 kJ

A)68.0 g

B)102 g

C)54.1 g

D)191 g

E)95.5 g

Q2) -w

A)the system loses thermal energy to the surroundings

B)work done on the system by the surroundings

C)energy flows into the system from the surroundings

D)the system gains thermal energy from the surroundings

E)energy flows out of the system into the surroundings

F)work done by the system on the surroundings

Q3) Describe the energy changes that occur when a book is held 1 m off the floor and then dropped.

Q4) Explain the difference between H and U.

Q5) 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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Sample Questions

Q1) When filling degenerate orbitals, electrons fill them singly first, with parallel spins. What is this known as?

A)Pauli exclusion principle

B)Hund's rule

C)aufbau principle

D)Heisenberg uncertainty principle

Q2) Which of the following quantum numbers describes the size and energy of an orbital?

A)magnetic quantum number

B)principal quantum number

C)angular momentum quantum number

D)spin quantum number

E) Schrödinger quantum number

Q3) The number of cycles that pass through a stationary point is called

A)wavelength.

B)amplitude.

C)frequency.

D)area.

E)median.

Q4) Describe how a neon light works.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Choose the statement that is 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 the most difficult of all electrons to remove.

D)Core electrons are the easiest of all electrons to remove.

E)All of the above are true.

Q2) An element that has the valence electron configuration

6s²5d⁴ belongs to which period and group?

A)period 6; group 16

B)period 6; group 12

C)period 6; group 4

D)period 6; group 6

E)period 6; group 2

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

Q4) List the noble gas that has the highest ionization energy.

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Use the bond energies provided to estimate _rH° for the reaction below.

XeF_{2 +}2F₂ XeF₆ _rH°= ?

Bond Bond Energy (kJ mol^-1)

Xe-F 147

F-F 159

A)-429 kJ mol^-1

B)+159 kJ mol^-1

C)-660 kJ mol^-1

D)+176 kJ mol^-1

E)-270 kJ mol^-1

Q2) Place the following in order of decreasing bond length. H-F H-I H-Br

A)H-F > H-Br > H-I

B)H-I > H-F > H-Br

C)H-I > H-Br > H-F

D)H-Br > H-F > H-I

E)H-F > H-I > H-Br

Q3) Describe the difference between a pure covalent bond and a polar covalent bond.

Q4) Define bond energy.

Q5) Define formal charge.

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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 approximate bond angle for a molecule with tetrahedral electron geometry and bent molecular geometry.

A)180°

B)<180°

C)<109.5°

D)109.5°

E)<120°

Q2) What is the molecular geometry of ClF₄ ^-?

A)seesaw

B)square planar

C)square pyramidal

D)tetrahedral

Q3) How many of the following molecules have sp³^d² hybridization on the central atom? SeCl₆ XeF₄ IF₅ AsCl₅

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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) The enthalpy change for converting 1.00 mol of ice at -50.0 °C to water at 70.0 °C is ________ kJ. The specific heats of ice, water, and steam are 2.09 J g^-1 °C^-1, 4.18 J g^-1 °C^-1, and 1.84 J g^-1 °C^-1, respectively. For H₂O, _fusH = 6.01 kJ mol^-1, and _vapH = 40.67 kJ mol^-1.

A)12.28

B)6.41

C)13.16

D)7154

E)9.40

Q2) The normal boiling point for H₂Seis higher than the normal boiling point for H₂S. This can be explained by

A)larger dipole-dipole forces for H₂Se.

B)larger dispersion forces for H₂Se.

C)larger hydrogen-bond forces for H₂Se.

D)larger dipole-dipole forces, larger dispersion forces, and larger hydrogen-bond forces for H₂Se.

Q3) Describe the difference between the conduction band and the valence band.

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Chapter 12: Solutions

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Sample Questions

Q1) A nonelectrolyte compound with a molar mass of 862 g mol^-1 was dissolved in enough water to make 100 mL of solution at 25 °C and produced an osmotic pressure of 38.4 mbar. Calculate the mass of the compound, in mg, used to make this solution.

A)116 mg

B)185 mg

C)134 mg

D)75.3 mg

E)84.3 mg

Q2) A solution containing less than the equilibrium amount is called

A)an unsaturated solution.

B)a dilute solution.

C)a supersaturated solution.

D)a concentrated solution.

E)a saturated solution.

Q3) Explain why water does not dissolve in gasoline.

Q4) Define a colloid.

Q5) Explain why the van't Hoff factor for MgCl₂ is less than its predicted value.

Q6) Define the Tyndall effect.

Page 14

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) The isomerization of methylisonitrile to acetonitrile CH₃NC(g) CH₃CN(g)

Is first order in CH₃NC. The half-life of the reaction is 5.20 × 10¹ s at 545 K. The rate constant when the initial [CH₃NC] is 0.030 mol L^-1 is ________ s^-1.

A)75.1

B)0.641

C)0.0133

D)1.56

E)2.84 × 10⁴

Q2) Given the following balanced equation, determine the rate of reaction with respect to [SO₃]. If the rate of O₂ loss is 3.56 × 10^-3 mol L^-1 s^-1, what is the rate of formation of SO₃? 2SO₂(g)+ O₂(g) 2SO₃(g)

A)3.56 × 10^-3 mol L^-1 s^-1

B)1.19 × 10^-3 mol L^-1 s^-1

C)1.78 × 10^-3 mol L^-1 s^-1

D)1.42 × 10^-2 mol L^-1 s^-1

E)7.12 × 10^-3 mol L^-1 s^-1

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) For which reaction will K_p = K_c?

A)S(s)+ O₂(g) SO₂(g)

B)2HgO(s) Hg(l)+ O₂(g)

C)CaCO₃(s) CaO(s)+ CO₂(g)

D)H₂CO₃(s) H₂O(l)+ CO₂(g)

E)2H₂O(l) 2H₂(g)+ O₂(g)

Q2) Consider the following reaction at equilibrium. What effect will reducing the pressure of the reaction mixture have on the system? Xe(g)+ 2F₂(g) XeF₄(g)

A)The reaction will shift in the direction of reactants.

B)The equilibrium constant will increase.

C)No effect will be observed.

D)The reaction will shift in the direction of products.

E)The equilibrium constant will decrease.

Q3) Can the K_p and K_c for a reaction ever have the same value? Why or why not?

Q4) Define Le Châtelier's principle.

Q5) Why aren't solids or liquids included in an equilibrium expression?

Q6) How is the reaction quotient different from an equilibrium constant for a given reaction?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Order the following acids in order of increasing strength: HBrO HBrO3 ^HBrO4HBrO₂

A)HBrO < HBrO₃ < ^HBrO4< HBrO₂ B)HBrO₄ < HBrO₃ < ^HBrO2< HBrO

C)HBrO < HBrO₃ < ^HBrO2< HBrO₄

D)HBrO < HBrO₂ < ^HBrO3< HBrO₄ E)HBrO₂ < HBrO₄ < ^HBrO< HBrO₂

Q2) Calculate the pH of a 1.60 mol L^-1 CH₃NH₃Cl solution. K_b for methylamine, CH₃NH₂, is 3.7 × 10^-4.

A)1.61

B)5.18

C)8.82

D)12.39

Q3) Which of the following correctly describes the characteristics of a strong acid?

A)ionizes only partially in aqueous solutions

B)has a very electron positive atom attached to the oxygen

C)has a nonpolar bond

D)is completely ionized in aqueous solution

E)is always monoprotic

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Calculate the percent ionization of nitrous acid in a solution that is 0.249 mol L^-1 in nitrous acid. The acid dissociation constant of nitrous acid is 4.50 × 10^-4.

A)1.12 × 10^-4

B)0.0450

C)4.25

D)0.342

E)5.53

Q2) What is the molar solubility of silver chloride (AgCl)in water? The solubility-product constant for AgCl is 1.8 × 10^-10 at 25 °C.

A)9.0 × 10^-11 mol L^-1

B)3.6 × 10^-10 mol L^-1

C)9.74 mol L^-1

D)1.9 × 10^-5 mol L^-1

E)1.3 × 10^-5 mol L^-1

Q3) 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.

Q4) Give the name of the compound that is in antifreeze and is toxic to pets.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Calculate _rS° for the following reaction. The S° for each species is shown below the reaction. N₂H₄(l)+ H₂(g) 2NH₃(g)

S° (J K^-1 mol^-1)121.2 130.7 192.8

A)+133.7 J K^-1 mol^-1

B)-59.1 J K^-1 mol^-1

C)+118.2 J K^-1 mol^-1

D)-202.3 J K^-1 mol^-1

E)+178.9 J K^-1 mol^-1

Q2) Why is heating your home with gas more efficient than heating it with electricity?

Q3) Determine the equilibrium constant for the following reaction at 298 K: SO₃(g)+ H₂O(g) H₂SO₄(l) _rG° = -90.5 kJ mol^-1

A)1.37 × 10^-16

B)4.78 × 10¹¹

C)9.11 × 10^-8

D)7.31 × 10¹⁵

E)0.964

Q4) Give the standard states for a gas, liquid, solid, and solution.

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Q > K

A)E°_cell < 0

B)E_cell< 0

C)E°_cell > 0

D)E_cell> 0

E)E_cell = 0

F)E_cell = E°_cell

Q2) What is the shorthand notation that represents the following galvanic cell reaction?

2Fe²⁺(aq)+ F₂(g)

2F^-(aq)

2Fe³⁺(aq)+

A)Fe²⁺(aq) Fe³⁺(aq) F₂(g) F^-(aq)

B)Fe(s) Fe²⁺(aq)

Fe³⁺(aq)F₂(g) F^-(aq) C(s)

C)Pt(s) Fe³⁺(aq), Fe²⁺(aq), F₂(g) F^-(aq) C(s)

D)Pt(s) Fe²⁺(aq), Fe³⁺(aq) F₂(g) F^-(aq) C(s)

Q3) What is electrolysis?

Q4) Why is sugar water not a good conductor of current?

Q5) What is the difference between a voltaic cell and an electrolytic cell?

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) The splitting of the uranium atom is called ________.

A)radioactive cleavage

B)nuclear fission

C)nuclear fusion

D)radioactive merge

E)the half-life

Q2) Describe what is meant by the "valley of stability."

Q3) Which of the following nuclides are most likely to decay via positron emission?

A)Cs-137

B)I-131

C)Al-24

D)K-42

E)N-14

Q4) Explain the concept of "magic numbers."

Q5) 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

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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 is TRUE regarding stereoisomers?

A)Stereoisomers differ in the way atoms are connected but have identical spatial orientation of atoms. They can be conformational or configurational.

B)Stereoisomers differ in the way atoms are connected but have identical spatial orientation of atoms. They can be constitutional or configurational.

C)Stereoisomers have the same atom connectivity but have different spatial orientation of atoms. They can be constitutional or enantiomeric.

D)Stereoisomers have the same atom connectivity but have different spatial orientation of atoms. They can be conformational or configurational.

E)Stereoisomers have the same atom connectivity but have different functional groups on the same positions of the carbon backbone. They can be conformational or configurational.

Q3) What are functional groups and how are they related to the concept of organic families? Provide examples.

Q4) What is the index of hydrogen deficiency?

Q5) What is the difference between primary, secondary, and tertiary amines?

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Chapter 21: Organic Chemistry II: Reactions

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Q1) Choose the weak base from the compounds below.

A)(CH₃)₂NH

B)CH₃CO₂H

C)CH₃COCH₃

D)CH₃CH₃

E)CH₃CH₂F

Q2) Which of the following statements is TRUE?

A)In nucleophilic substitution reactions at a saturated carbon atom, a carbon with a positive charge behaves as a nucleophile.

B)Nucleophilic substitution reactions at a saturated carbon atom can follow one of three mechanisms: S_N1, S_N2, and S_N3.

C)A carbon atom with a partial positive charge can react with a nucleophile, a molecule, or ion that has a lone pair of electrons.

D)Alkanes are an example of saturated organic compounds that easily participate in nucleophilic substitution reactions at a saturated carbon atom.

E)In nucleophilic substitution reactions at a saturated carbon atom, the replaced group is called an electrophile.

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Chapter 22: Biochemistry

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Q1) Which of the following hold two strands of DNA in an a-helix?

A)ester linkages

B)disulfide linkages

C)peptide bonds

D)hydrogen bonds

E)glycosidic linkages

Q2) Why does the melting point of fatty acids decrease with the number of double bonds present?

Q3) Which of the following describes a primary protein structure?

A)protein structure maintained by disulfide linkages

B)amino acid sequence maintained by peptide bonds

C)protein chains maintained by interactions of peptide backbones

D)arrangement of multiple protein subunits

E)protein structure maintained through multiple hydrogen bonds

Q4) Identify the polysaccharide.

A)glucose

B)cholesterol

C)fructose

D)cellulose

E)sucrose

Page 24

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Chapter 23: Chemistry of the Nonmetals

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Q1) Which of the following is TRUE?

A)Phosphoric acid is used to produce the "fizziness" in soda.

B)Sodium phosphate is an additive in baked goods to give them a longer shelf life.

C)Phosphorus can range in oxidation state from -3 to +5.

D)The phosphorus in the phosphite ion has an oxidation state of -3.

E)Phosphoric acid is a very strong oxidizing agent, just like nitric acid.

Q2) Determine the number of vertices and faces in the closo-Borane

B₁₂H₁₂^2-.

A)vertices = 12, faces = 24

B)vertices = 12, faces = 20

C)vertices = 8, faces = 16

D)vertices = 6, faces = 8

E)vertices = 8, faces = 12

Q3) pyroxene

A)Si₂O₅^2-

B)SiO₄^4-

C)Si₄O₁₁^6-

D)SiO₃^2-

E)Si₂O₇^6-

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Page 25

Chapter 24: Metals and Metallury

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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) Which of the following is TRUE concerning interstitial alloys?

A)The metal takes on the properties of the foreign atoms once they are present.

B)Small atoms fit in between the metal atoms in the crystal structure.

C)Brass is an example of an interstitial alloy.

D)The foreign atoms present weaken the metallic structure.

E)They are alloys composed of two nonmetallic elements.

Q3) What is an advantage of hydrometallurgical processes over pyrometallurgical processes?

Q4) Describe the difference between a substitutional alloy and an interstitial alloy.

Q5) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Which of the following is a type of structural isomerism?

A)linkage isomers

B)cis-trans isomers

C)fac-mer isomers

D)enantiomers

E)optical isomers

Q2) Choose the electron configuration for Fe³⁺.

A)[Ar]4s²3d³

B)[Ar]3d⁵

C)[Ar]4s²3d⁹

D)[Ar]4s¹3d⁴

E)[Ar]4s²3d⁶

Q3) Determine the chemical formula for the compound tetracarbonylplatinum(IV)chloride.

A)[Pt(CO)₄]Cl₂

B)[Pt(CO)₄]Cl₄

C)[Pt(CO)₄Cl₄]^2-

D)[PtCl₄](CO)₄

E)[Pt(CO)₄Cl₄]^4-

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

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