General Chemistry II Exam Practice Tests - 3452 Verified Questions

General Chemistry II Exam Practice Tests

Course Introduction

General Chemistry II is a continuation of introductory chemistry, focusing on the fundamental principles that govern chemical reactions and properties of matter. This course covers topics including chemical kinetics, chemical equilibrium, acid-base chemistry, thermodynamics, electrochemistry, and an introduction to coordination compounds. Emphasis is placed on quantitative problem-solving, laboratory techniques, and critical thinking skills needed to understand the behavior of matter at the molecular and atomic levels. Laboratory sessions complement the lectures by providing hands-on experience with experiments that illustrate core concepts and encourage scientific inquiry.

Recommended Textbook

Chemistry A Molecular Approach 4th Edition by

Nivaldo J. Tro

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25 Chapters

3452 Verified Questions

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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169 Verified Questions

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Sample Questions

Q1) What symbol is used to represent the factor 10^-3?

A) M

B) m

C)

D) n

Answer: B

Q2) Identify the TRUE statement.

A) Hydrogen peroxide is used as rocket fuel and removes color from hair.

B) Hydrogen peroxide can be used as a drink and can be used to bathe in.

C) Both water and hydrogen peroxide are stable molecules.

D) Water reacts with skin.

E) Hydrogen peroxide is smaller than water.

Answer: A

Q3) How many significant figures are there in the answer to the following problem? (9.992 × 3.200)+ 0.610 = ?

A) one

B) two

C) three

D) four

Answer: D

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following does NOT describe a nonmetal?

A) tends to gain electrons

B) found in the upper right hand corner of the periodic table

C) poor conductor of electricity

D) nonmetals are generally unreactive

E) poor conductor of heat

Answer: D

Q2) How many electrons are in the ion,Ni²⁺?

A) 26

B) 28

C) 30

D) 59

E) 52

Answer: A

Q3) What group of elements in the periodic table are the most unreactive and why?

Answer: The noble gases are the most unreactive since they do not combine with other elements to form compounds.

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

Answer: protons

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) How many SO₃^2- ions are contained in 99.6 mg of Na₂SO₃? The molar mass of Na₂SO₃is 126.05 g/mol.

A) 1.52 × 10²⁷ SO₃ ions

B) 4.76 × 10²⁰ SO₃ions

C) 2.10 × 10²¹ SO₃ ions

D) 1.05 × 10²¹ SO₃ ions

E) 9.52 × 10²⁰ SO₃ ions

Answer: B

Q2) How many molecules of N₂O₄ are in 76.3 g N₂O₄? The molar mass of N₂O₄ is 92.02 g/mol.

A) 5.54 × 10²⁵ N₂O₄ molecules

B) 7.26 × 10²³ N₂O₄ molecules

C) 1.38 × 10²⁴ N₂O₄ molecules

D) 4.59 × 10²⁵N₂O₄ molecules

E) 4.99 × 10²³ N₂O₄ molecules

Answer: E

Q3) Give the name for HNO₂.

Answer: nitrous acid

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) How many of the following compounds are soluble in water? Pb(OH)₂

LiNO₃ NH₄Br K₂S

A) 0

B) 1

C) 2

D) 3

E) 4

Q2) Give the spectator ions for the reaction that occurs when aqueous solutions of MgSO₃ and HI are mixed.

A) H⁺ and SO₃^2-

B) Mg²⁺ and I^-

C) H⁺ and I^-

D) SO₃^2- and Mg²⁺

E) No spectator ions are present.

Q3) When dissolved in water,KOH behaves as

A) an acid that forms K⁺ and OH^- ions.

B) an acid that forms KO^- and H⁺ ions.

C) a base that forms K⁺ and OH^- ions.

D) a base that forms KO^- and H⁺ ions.

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

Page 6

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Chapter 5: Gases

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Sample Questions

Q1) The volume of a gas is proportional to number of moles of a gas is known as

A) Avogadro's Law.

B) Ideal Gas Law.

C) Charles's Law.

D) Boyle's Law.

E) Dalton's Law.

Q2) A 0.529-g sample of gas occupies 125 mL at 60.cm of Hg and 25°C.What is the molar mass of the gas?

A) 11 g/mol

B) 81 g/mol

C) 110 g/mol

D) 130 g/mol

Q3) Identify the element or molecule that has the highest effusion rate.

A) He

B) O₂

C) HCl

D) SO₂

E) Rn

Q4) Give the major gas in dry air.

Q5) Define pressure.

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Chapter 6: Thermochemistry

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Sample Questions

Q1) The temperature rises from 25.00°C to 29.00°C in a bomb calorimeter when 3.50 g of sucrose undergoes combustion in a bomb calorimeter.Calculate E_rxn for the combustion of sucrose in kJ/mol sucrose.The heat capacity of the calorimeter is 4.90 kJ/°C.The molar mass of sugar is 342.3 g/mol.

A) -1.92 × 10³ kJ/mole

B) 1.92 × 10³ kJ/mole

C) -1.23 × 10³ kJ/mole

D) 2.35 × 10⁴ kJ/mole

Q2) Using the following equation for the combustion of octane,calculate the heat of reaction for 50.00 g of octane.The molar mass of octane is 114.33 g/mole. 2 C₈H₁₈ + 25 O₂ 16 CO₂ + 18 H₂O H°_rxn= -11018 kJ

A) -2410 kJ

B) -2.410 kJ

C) -4820kJ

D) -6300kJ

Q3) Explain the difference between H and DE.

Q4) Describe the energy changes that occur when a book is held 6 ft off the floor and then dropped.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Give an example of a d orbital.

Q2) Identify the correct values for a 4f orbital.

A) n = 3, l = 2, m_l = 1

B) n = 2, l = 1, m_l = 1

C) n = 1, l = 0, m_l = 0

D) n = 2, l = 0, m_l = 0

E) n = 4, l = 3, m_l = -2

Q3) An FM radio station broadcasts electromagnetic radiation at a frequency of 101.2 MHz.The wavelength of this radiation is ________ m.

A) 2.964 × 10⁶

B) 2.964

C) 3.036 × 10¹⁶

D) 3.036 × 10¹⁰

E) 0.3373

Q4) Describe how a neon light works.

Q5) Sketch one of the 3p orbitals below.How are they different from the 2p orbitals?

Q6) What is the photoelectric effect?

Q7) Why don't we observe the wavelength of everyday macroscopic objects?

Q8) Define diffraction.

Q9) Give an example of a p orbital.

Page 9

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Page 10

Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) What is the general valence-electron ground-state electron configuration for neutral alkali metals?

A) ns²

B) ns¹

C) (n-1)s²ns¹

D) (n-1)s²ns²np²

E) np⁵

Q2) How many valence electrons does an atom of Na possess?

A) 1

B) 3

C) 4

D) 7

E) 0

Q3) Give the electron configuration for O.

A) 1s²2s²2p⁴

B) 1s²2p⁴

C) 1s²2s²2p³

D) 1s²2s¹2p⁵

E) 1s²2s²2p²

Q4) Define ionization energy.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Using periodic trends,place the following bonds in order of increasing ionic character.

Si-P Si-Cl Si-S

A) Si-P < Si-Cl < Si-S

B) Si-P < Si-S < Si-Cl

C) Si-S < Si-Cl < Si-P

D) Si-Cl < Si-P < Si-S

E) Si-Cl < Si-S < Si-P

Q2) Place the following elements in order of increasing electronegativity. Ba C Sc

A) Ba < Sc < C

B) Sc < C < Ba

C) Ba < C < Sc

D) C < Ba < Sc

E) C < Sc < Ba

Q3) Place the following in order of decreasing magnitude of lattice energy. KF BaO RbI

A) RbI > KF > BaO

B) RbI > BaO > KF

C) BaO > RbI > KF

D) KF > RbI > BaO

E) BaO > KF > RbI

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Page 12

Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) What geometric arrangement of charge clouds is expected for an atom that has three charge clouds?

A) trigonal bipyramidal

B) linear

C) trigonal

D) tetrahedral

Q2) A molecule containing a central atom with sp³d hybridization has a(n)________ electron geometry.

A) seesaw

B) bent

C) octahedral

D) trigonal planar

E) trigonal bipyramidal

Q3) Determine the molecular geometry about each interior atom in the following structure.Sketch the three-dimensional structure and label the interior atoms with the corresponding molecular geometry.

CH₂CHCCCH₃

Q4) Give the electron geometry,molecular geometry,and hybridization for both carbons in CH₃COOH.

Page 13

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) Choose the substance with the highest surface tension.

A) HOCH₂CH₂OH

B) CH₂F₂

C) CH₃CH₂I

D) CH₃CH₂CH₂OH

E) CH₃CH₂CH₃ 

Q2) Identify the compound that does NOT have hydrogen bonding.

A) (CH₃)₂N(CH₂)₃CH₃

B) H₂O

C) CH₃(CH₂)₅

D) HF

E) CH₃(CH₂)₂NH₂

Q3) Choose the pair of substances that are most likely to form a homogeneous solution.

A) KCl and Hg

B) NaI and C₆H₁₄

C) C₃H₈ and C₂H₅OH

D) F₂ and PF₃

E) NH₃ and CH₃OH

Q4) Define volatile.

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Identify an element that is NOT in silicate ceramics.

A) oxygen

B) silicon

C) aluminum

D) carbon

E) hydrogen

Q2) Identify for the hexagonal unit cell.

A) a b = c

B) a = b c

C) a b c

D) a = b = c

Q3) A common use of polyurethane is

A) tire cord.

B) fishing line.

C) foam cups.

D) spray-on insulation.

E) films.

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.

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Chapter 13: Solutions

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Sample Questions

Q1) Which of the following statements is generally TRUE?

A) The solubility of a solid is not dependent on either temperature or pressure.

B) The solubility of a solid is highly dependent on pressure.

C) The solubility of a solid is highly dependent on both pressure and temperature.

D) The solubility of a solid is highly dependent on temperature.

E) None of the above.

Q2) Solutions having osmotic pressures more than those of body fluids are called A) isosmotic.

B) hyperosmotic.

C) hyposmotic.

D) neosmotic.

E) magnosmotic.

Q3) If two substances are soluble in all proportions then they are ________ in each other.

A) miscible

B) nonpolar

C) immiscible

D) neutral

E) polar

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Page 16

Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The first-order rearrangement of CH₃NC is measured to have a rate constant of 3.61 × 10^-15s^-1 at 298 K and a rate constant of 8.66 × 10^-7s^-1 at 425 K.Determine the activation energy for this reaction.

A) 160. kJ/mol

B) 240. kJ/mol

C) 417 kJ/mol

D) 127 kJ/mol

E) 338 kJ/mol

Q2) For a reaction that follows the general rate law,Rate = k[A][B]²,what will happen to the rate of reaction if the concentration of A is increased by a factor of 5.00? The rate will

A) decrease by a factor of 1/25.0.

B) decrease by a factor of 1/5.00.

C) increase by a factor of 5.00.

D) increase by a factor of 25.0.

Q3) Explain how the order of a reaction can be determined.

Q4) What is the difference between average reaction rate and instantaneous reaction rate?

Q5) Define activation energy.

Page 17

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction at equilibrium.What will happen if the pressure increased? 4 FeS₂(s)+ 11 O₂(g) 2 Fe₂O₃(s)+ 8 SO₂(g)

A) The equilibrium constant will increase.

B) The equilibrium will change in the direction of the reactants.

C) The equilibrium will change in the direction of the products.

D) No change in equilibrium is observed.

E) The equilibrium constant will decrease.

Q2) The reaction below has a K_p value of 41.What is the value of K_c for this reaction at 400 K? N₂(g)+ 3 H₂(g) 2 NH₃(g)

A) 3.8 × 10^-2

B) 4.4 × 10⁴

C) 26

D) 2.3 × 10^-5

E) 41

Q3) Why aren't solids or liquids included in an equilibrium expression?

Q4) Why is an equilibrium constant unitless?

Q5) Explain dynamic equilibrium.Use the generic reaction A(g) B(g)to explain.

Q6) Define Le Chatelier's Principle.

Page 18

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Chapter 16: Acids and Bases

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Sample Questions

Q1) Which of the following is a polyprotic acid?

A) HBr

B) H₂SO₄

C) KCN

D) CF₄

E) HC₂H₃O₂

Q2) Calculate the pH of a solution that contains 2.4 × 10^-⁵ M H₃O at 25°C.

A) 2.40

B) 9.38

C) 4.62

D) 8.38

E) 5.62

Q3) Which one of the following salts,when 1 mole is dissolved in water,produces the solution with a pH closest to 7.00?

A) NH₄ClO₄

B) Na₂O

C) NaHSO4

D) NaF

E) NaOH

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Chapter 17: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Which of the following is TRUE?

A) An effective buffer has a [base]/[acid] ratio in the range of 10 - 100.

B) A buffer is most resistant to pH change when [acid] = [conjugate base]

C) An effective buffer has very small absolute concentrations of acid and conjugate base.

D) A buffer cannot be destroyed by adding too much strong base. It can only be destroyed by adding too much strong acid.

E) None of the above is true.

Q2) Explain the common ion effect with respect to molar solubility.

Q3) Which of the following compounds will have the highest molar solubility in pure water?

A) PbS, K_sp = 9.04 × 10^-²⁹

B) ZnSe, K_sp = 3.6 × 10^-²⁶

C) Fe(OH)₃, K_sp = 2.79 × 10^-³⁹

D) ZnS, K_sp = 2.0 × 10^-²⁵

E) Ag₂S, K_sp = 6.0 × 10^-51

Q4) Define a buffer.

Q5) Give the name of the compound that is in antifreeze and is toxic to pets.

Q6) Identify the compound that is in stalactites and stalagmites.

Page 20

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) For the following example,identify the following. F₂(l) F₂(g)

A) at low temperature, the reaction is spontaneous and G < 0 and at high temperature, the reaction is spontaneous and G < 0

B) at low temperature, the reaction is nonspontaneous and G > 0 and at high temperature, the reaction is spontaneous and G < 0

C) at low temperature, the reaction is spontaneous and G < 0 and at high temperature, the reaction is nonspontaneous and G > 0

D) at low temperature, the reaction is nonspontaneous and G > 0 and at high temperature, the reaction is nonspontaneous and G > 0

E) It is not possible to determine without more information.

Q2) Identify the process in which the entropy increases.

A) a decrease in the number of moles of a gas during a chemical reaction

B) the phase transition from a gas to a liquid

C) the phase transition from a solid to a gas

D) the phase transition from a gas to a solid

E) the phase transition from a liquid to a solid

Q3) Define the third law of thermodynamics.

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Chapter 19: Electrochemistry

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Sample Questions

Q1) Predict the species that will be reduced first if the following mixture of molten salts undergoes electrolysis. Na⁺,Ca²⁺,Cl ,Br ,F

A) Na⁺

B) Cl^-

C) Ca²⁺

D) Br^-

E) F^-

Q2) Describe how water can be a good conductor of current.

A) use pure water

B) heat the water

C) add salt

D) chill the water

E) vaporize the water

Q3) Identify the location of reduction in an electrochemical cell.

A) the anode

B) the cathode

C) the electrode

D) the salt bridge

E) the socket

Q4) Give an example of an inert electrode.

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Write a nuclear equation to describe the neutron induced fission of U-235 to form Xe-134 and Sr-100.Determine how many neutrons are produced in the reaction.

A) 4

B) 3

C) 1

D) 0

E) 2

Q2) Determine the binding energy per nucleon of an Mg-24 nucleus.The Mg-24 nucleus has a mass of 23.985045.A proton has a mass of 1.00728 amu,a neutron has a mass of 1.008665 amu,and 1 amu is equivalent to 931.5 MeV of energy.

A) 0.3050 MeV

B) 8.83 MeV

C) 0.113 MeV

D) 192 MeV

E) 4.41 MeV

Q3) Describe what is meant by the term "valley of stability"?

Q4) Define radioactivity.

Q5) List two problems associated with nuclear power.

Q6) Define chain reaction in terms of the fission of uranium nucleus.

Page 23

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Chapter 21: Organic Chemistry

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Sample Questions

Q1) Which of the following alkane names is correct?

A) 1,2-dipropylhexane

B) 1-propylnonane

C) 4,5-dipropylpentane

D) 2-methyl-3-ethylheptane

E) All of the above names are correct.

Q2) Which of the following names is correct?

A) 4-bromo-2-methylbenzene

B) 1-chloro-3-propylbenzene

C) 3-bromo-4-iodobenzene

D) 2,2-dibromobenzene

E) 4-bromo-2-iodobenzene

Q3) Name the following compound.

CH₃CH₂OCH₂CH₂CH₂CH ₃

A) ethyl propyl ether

B) methyl propyl ether

C) butyl ethyl ether

D) butyl propyl ether

E) butyl methyl ether

Q4) What does the term "racemic mixture" mean?

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Chapter 22: Biochemistry

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Sample Questions

Q1) Identify the polysaccharide.

A) fructose

B) maltose

C) cholesterol

D) starch

E) estradiol

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

Q3) What is the difference between an -glycosidic linkage and a -glycosidic linkage? What implication does this have for the strength of the intermolecular forces between molecules?

Q4) The following is part of a DNA sequence.What is its complementary sequence? TCGTAAGCTTGCG

A) AGCATTCGAACGC

B) CGCAAGCTTACGA

C) TCATTCAC

D) GCGTTCGAATGCT

E) TCGTAAGCTTGCG

Q5) What is a codon?

Q6) Why are lipids well-suited as structural components of cell membranes?

25

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Chapter 23: Chemistry of the Nonmetals

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Sample Questions

Q1) 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

Q2) What is the major detraction from using bituminous coal as an energy source?

A) It has a very low carbon content.

B) It contains a high amount of sulfur, which when burned contributes to acid rain.

C) It has an extremely high surface area which makes it difficult to burn.

D) Bituminous coal is noncrystalline, which gives it a low energy/gram ratio.

E) There are no detractions from using bituminous coal.

Q3) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q4) Draw the Lewis structure for IF₄ and determine the geometry of the ion.

Q5) Describe the Frasch process for obtaining sulfur.

Q6) Why are phosphate compounds added to detergents?

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Chapter 24: Metals and Metallurgy

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Sample Questions

Q1) Identify the metal most commonly used for wires in the electrical industry.

A) copper

B) silver

C) lead

D) aluminum

E) potassium

Q2) 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

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

Q4) Which substance is the best conductor of electricity?

A) astatine

B) diamond

C) tungsten

D) xenon

Page 27

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Identify the ion that is responsible for the blue of turquoise.

A) Cr²⁺

B) Cu⁺

C) Cu²⁺

D) Cr³⁺

E) Fe²⁺

Q2) Which ion would you expect to have the largest crystal field splitting ?

A) [Rh(CN)₆]^4-

B) [Rh(CN)₆]^3-

C) [Rh(H₂O)₆]²⁺

D) [Rh(H₂O)₆]³⁺

Q3) Identify the geometry of the complex ion if the hybridization is dsp².

A) square planar

B) octahedral

C) tetrahedral

D) linear

E) trigonal planar

Q4) Explain how EDTA is used to treat lead poisoning.

Q5) What is the Lanthanide contraction?

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