General Chemistry II Practice Questions - 3452 Verified Questions

Course Introduction

General Chemistry II Practice Questions

General Chemistry II continues the foundational study of chemistry, building on concepts introduced in General Chemistry I. This course covers topics such as chemical kinetics, chemical equilibrium, acid-base chemistry, thermodynamics, electrochemistry, and introductory topics in nuclear and organic chemistry. Through lectures, laboratory experiments, and problem-solving exercises, students develop a deeper understanding of molecular interactions, energy changes, and reaction mechanisms, equipping them with the knowledge to analyze and solve more complex chemical problems.

Recommended Textbook Chemistry A Molecular Approach 4th Edition by Nivaldo J. Tro

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25 Chapters

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) Identify the type of energy that is NOT chemical energy.

A) battery

B) gasoline in a car

C) light bulb

D) ball rolling down the hill

E) food

Answer: D

Q2) Choose the compound from the list below.

A) silver

B) methanol

C) helium

D) tin

E) sodium

Answer: B

Q3) Identify a gas.

A) definite volume and definite shape

B) definite volume and no definite shape

C) definite shape and no definite volume

D) no definite shape and no definite volume

Answer: D

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) The average atomic mass for silver is ________.

A) 28.09

B) 14

C) 107.87

D) 47

Answer: C

Q2) Which of the following elements is a good conductor of heat and electricity?

A) carbon

B) chlorine

C) xenon

D) aluminum

Answer: D

Q3) An atom of ¹³C contains ________ protons.

A) 6

B) 19

C) 7

D) 9

E) 13

Answer: A

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Page 4

Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) How many chloride ions are there in 7.50 mol of aluminum chloride?

A) 3.00 chloride ions

B) 22.5 chloride ions

C) 4.52 × 10²⁴ chloride ions

D) 1.35 × 10²⁵ chloride ions

Answer: D

Q2) Balance the following equation.

_____ C₉H₂₀ + _____ O₂ _____ H₂O + _____ CO₂

Answer: C₉H₂₀ + 14 O₂ 10 H₂O + 9 CO₂

Q3) Calculate the molar mass for Ba(ClO₄)₂.

A) 336.23 g/mol

B) 236.78 g/mol

C) 99.45 g/mol

D) 137.33 g/mol

E) 188.78 g/mol

Answer: A

Q4) Give the name for HNO₂.

Answer: nitrous acid

Page 5

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) Magnesium burns in air with a dazzling brilliance to produce magnesium oxide: 2 Mg(s)+ O₂(g) 2 MgO(s)

How many moles of O₂ are consumed when 2.10 mol of magnesium burns?

A) 0.0864

B) 0.952

C) 2.10

D) 4.20

E) 1.05

Q2) Which of the following compounds is NOT an Arrhenius acid?

A) CH₃CO₂H

B) C₆HNH

C) HNO₂

D) H₂SO₄

Q3) If 100.mL of 0.100 M Na₂SO₄ is added to 200.mL of 0.150 M NaCl,what is the concentration of Na⁺ ions in the final solution? Assume that the volumes are additive.

A) 0.133 M

B) 0.167 M

C) 0.250 M

D) 0.350 M

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Chapter 5: Gases

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Sample Questions

Q1) The density of a gas is 1.43 g/L at STP.What is the gas?

A) H₂

B) F₂

C) Kr

D) O₂

E) N₂

Q2) How many liters of O₂ gas at 25°C and 1.00 atm pressure are needed to react with 30.12 grams of potassium hydride according to the chemical equation shown below? 2 KH(s)+ O₂(g) H₂O(l)+ K₂O(s)

A) 4.59 L

B) 9.19 L

C) 18.4 L

D) 36.8 L

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

A) H₂

B) N₂

C) Ar

D) Cl₂

E) Rn

Q4) Give the major gas in dry air.

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Use the H°_f and H°_rxninformation provided to calculate H°_f for SO₃(g): H°_{f (kJ/mol)} 2 SO₂(g)+ O₂(g) 2 SO₃(g) H°_rxn= -198 kJ SO₂(g)-297

A) -792 kJ/mol

B) -248 kJ/mol

C) -495 kJ/mol

D) -578 kJ/mol

E) -396 kJ/mol

Q2) Which of the following statements is TRUE?

A) State functions do not depend on the path taken to arrive at a particular state.

B) DE_rxn can be determined using constant volume calorimetry.

C) Energy is neither created nor destroyed, excluding nuclear reactions.

D) H_rx_n can be determined using constant pressure calorimetry.

E) All of the above are true.

Q3) Define chemical energy.

Q4) Where does the energy absorbed during an endothermic reaction go?

Q5) Explain the difference between H and DE.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Which of the following occurs as the wavelength of a photon increases?

A) The frequency decreases.

B) The energy increases.

C) The speed decreases.

D) Planck's constant increases.

E) None of the above occurs as the wavelength of a photon increases.

Q2) Calculate the wavelength (in nm)of the blue light emitted by a mercury lamp with a frequency of 6.88 × 10¹⁴ Hz.

A) 229 nm

B) 436 nm

C) 206 nm

D) 485 nm

E) 675 nm

Q3) How many different values of m_l are possible in the 4d sublevel?

A) 2

B) 0

C) 3

D) 5

E) 6

Q4) Why don't we observe the wavelength of everyday macroscopic objects?

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Place the following in order of increasing IE₁. N F As

A) N < As < F

B) As < N < F

C) F < N < As

D) As < F < N

E) F < As < N

Q2) Which of the following have the same number of valence electrons?

A) Cs, Bi, At

B) Al, Ge, Sb

C) P, As, Sb

D) Xe, Rn, At

E) Fr, Ra, Lr

Q3) When filling degenerate orbitals,electrons fill them singly first,with parallel spins.This is known as (the)

A) Pauli exclusion principle.

B) Hund's rule.

C) Aufbau principle.

D) Heisenberg uncertainty principle.

Q4) Why do successive ionization energies increase?

Q5) Why do Li,Na,and K have similar chemical properties?

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Place the following in order of decreasing magnitude of lattice energy. NaF RbBr KCl

A) RbBr > NaF > KCl

B) NaF > KCl > RbBr

C) KCl > NaF > RbBr

D) NaF > RbBr > KCl

E) RbBr > KCl > NaF

Q2) Which of the following contains an atom that does NOT obey the octet rule?

A) LiF

B) SnO₂

C) BrF₃

D) IF

E) O₂

Q3) Define bond energy.

Q4) How many lone pairs are on the Br atom in BrF₂^-?

A) 0

B) 1

C) 2

D) 3

Q5) List the most electronegative atom.

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) Soap works with water because

A) the polar head and the nonpolar tail attracts water.

B) the polar head attracts the grease and the nonpolar tail attracts water.

C) the polar head attracts water and the nonpolar tail attracts the grease.

D) the polar head and the nonpolar tail attracts the grease.

Q2) What is the molecular geometry of NCl₃?

A) T-shaped

B) tetrahedral

C) trigonal planar

D) trigonal pyramidal

E) octahedral

Q3) Give the hybridization for the Br in BrF₅.

A) sp³d²

B) sp³d

C) sp³

D) sp²

E) sp

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

Q5) 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) Identify the characteristics of a liquid.

A) indefinite shape and volume

B) indefinite shape, but definite volume

C) definite shape and volume

D) none of the above

E) all of the above

Q2) Capillary action occurs because

A) cohesive forces are greater than adhesive forces.

B) adhesive forces equal adhesive forces.

C) adhesive forces are greater than cohesive forces.

D) surface tension is high.

E) surface tension is low.

Q3) Identify the place which has the highest boiling point of water.

A) Death Valley, 282 feet below sea level

B) A pressurized passenger jet, 39,000 feet

C) Panama City, Florida, sea level

D) Mt. Everest, 29,035 feet

E) Denver, Colorado, 5,280 feet

Q4) Define volatile.

Q5) Define viscosity.

13

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Describe the difference between the conduction band and the valence band.

Q2) Cesium has a radius of 272 pm and crystallizes in a body-centered cubic structure.What is the edge length of the unit cell?

A) 314 pm

B) 385 pm

C) 544 pm

D) 628 pm

Q3) Identify a use for cement.

A) beakers

B) bridges

C) furniture

D) wooden decks

E) bricks

Q4) Which of the following is considered an atomic solid?

A) F₂

B) CsBr

C) N₂

D) Nb

E) None of these is an atomic solid.

Q5) Give the edge length in terms of r for a simple cubic cell.

Page 14

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Chapter 13: Solutions

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Sample Questions

Q1) Parts per million requires a multiplication factor of ________.

A) 10^-9

B) 10^-3

C) 10³

D) 10¹⁵

E) 10⁶

Q2) 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 and 154 g/mol,respectively.

A) 7.36 × 10^-⁴

B) 0.736

C) 0.102

D) 0.0543

E) 5.43

Q3) Give the preparation of rock candy.

Q4) Define colloid.

Q5) Define osmosis.

Q6) Explain why water does not dissolve in gasoline.

Q7) Define Tyndall effect.

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) Identify the rate-determining step.

A) the slowest step

B) the faster step

C) the fast step

D) always the last step

E) always the second step

Q2) Which of the following statements is FALSE?

A) The half-life of a zero-order reaction is dependent on concentration.

B) The half-life of a second-order reaction is not dependent on concentration.

C) The rate of second-order reactions is dependent on concentration.

D) The rate of a first-order reaction is dependent on concentration.

E) None of the statements is FALSE.

Q3) The rate constant for a first-order reaction is 0.54 s^-1.What is the half-life of this reaction if the initial concentration is 0.33 M?

A) 0.089 s

B) 1.8 s

C) 0.31 s

D) 5.6 s

E) 1.3 s

Q4) What is a catalyst and what function does it serve?

Page 16

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) The reaction below has a K_c value of 1.0 × 10¹².What is the value of K_p for this reaction at 200 K? 2 SO₂(g)+ O₂(g) 2 SO₃(g)

A) 6.1 × 10^-14

B) 1.0 × 10¹²

C) 1.6 × 10^-11

D) 1.6 × 10¹³

E) 6.1 × 10¹⁰

Q2) Consider the following reaction at equilibrium.What will happen if O₂ is added to the reaction? 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.

Q3) Define Le Chatelier's Principle.

Q4) Can the K_p and K_c for a reaction ever have the same value? Why or why not?

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Chapter 16: Acids and Bases

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Sample Questions

Q1) Which one of the following will form a basic solution in water?

A) KC₂H₃O₂

B) KCN

C) NaClO₂

D) LiBrO

E) All of the above will form basic solutions.

Q2) Calculate the pH of a 1.60 M 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) Calculate the pOH in an aqueous solution with a pH of 7.85 at 25°C.

A) 4.15

B) 5.15

C) 6.15

D) 7.15

E) 8.15

Q4) What is the difference between a strong and weak acid?

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Chapter 17: Aqueous Ionic Equilibrium

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Sample Questions

Q1) A sample contains Ba₃(PO₄)₂,HgS,AgCl,NH₄Br,and CoS.Identify the precipitate after the addition of 6 M HCl; then H₂S and 0.2 M HCl; and then OH^- to a pH of 8.

A) Ba₃(PO₄)₂

B) HgS

C) AgCl

D) NH₄Br

E) CoS

Q2) Determine the solubility of the ions that is calculated from the Ksp for Mg₂CO₃.

A) 2S³

B) S³

C) 4S³

D) S²

E) 2S²

Q3) Consider the K_sp values for two compounds: MZ,K_sp = 1.5 × 10^-2⁰ and MZ₂,K_sp = 1.5 × 10^-20.Why don't these compounds have the same molar solubility?

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) Give the name of the reaction that does not achieve the theoretical limits with respect to free energy in thermodynamics.

A) reversible reaction

B) forward reaction

C) reverse reaction

D) equilibrium reaction

E) irreversible reaction

Q2) Calculate G_rxn at 298 K under the conditions shown below for the following reaction. CaCO₃(s) CaO(s)+ CO₂(g) G° = +131.1 kJ

P(CO₂)= 0.033 atm

A) -49.3 kJ

B) -8.32 kJ

C) +122.6 kJ

D) +39.7 kJ

E) +43.3 kJ

Q3) How many microstates are possible in a collection of four particles that are present,with two particles each in two connected flasks? Sketch them below.

Q4) Why can endothermic reactions be spontaneous?

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Page 20

Chapter 19: Electrochemistry

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Sample Questions

Q1) Balance the following redox reaction if it occurs in acidic solution.What are the coefficients in front of H₂C₂O₄and H₂O in the balanced reaction? MnO₄ (aq)+ H₂C₂O₄(aq) Mn²⁺(aq)+ CO₂(g)

A) H₂C₂O₄= 5, H₂O = 8

B) H₂C₂O₄= 1, H₂O = 1

C) H₂C₂O₄= 5, H₂O = 1

D) H₂C₂O₄= 1, H₂O = 4 E) H₂C₂O₄= 3, H₂O = 2

Q2) Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C.(The equation is balanced.) Pb(s)+ Br₂(l) Pb²⁺(aq)+ 2 Br (aq) Pb²⁺(aq)+ 2 e Pb(s)E° = -0.13 V Br₂(l)+ 2 e^- 2 Br^-(aq)E° = +1.07 V

A) +1.20 V

B) +0.94 V

C) -0.94 V

D) -1.20 V

E) -0.60 V

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Atoms with Z > ________ are radioactive and decay in one or more steps involving mostly alpha and beta decay.

A) 72

B) 130

C) 83

D) 150

E) 30

Q2) Why is an alpha emitter much more harmful if is ingested than when applied to the skin?

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

Q4) Identify the common radiotracers used in the diagnosis of medical problems.

A) fluorine-17

B) iodine-132

C) thallium-202

D) iron-58

E) iodine-131

Q5) List two problems associated with nuclear power.

Q6) Explain the concept of "magic numbers."

Q7) Define chain reaction in terms of the fission of uranium nucleus.

Q8) Define radioactivity.

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Chapter 21: Organic Chemistry

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Sample Questions

Q1) Identify the compound that is in nail polish remover.

A) benzaldehyde

B) butanone

C) acetone

D) butanal

E) vanillin

Q2) Identify the alkane with the highest boiling point.

A) pentane

B) butane

C) ethane

D) methane

E) propane

Q3) Identify the alcohol in rubbing alcohol.

A) 2-pentanol

B) isopropyl alcohol

C) methanol

D) 1-butanol

E) ethanol

Q4) What is meant by the term "structural isomer"? Draw a structural isomer of CH₃-CH₂-CH₂-CH₃.

Page 23

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Chapter 22: Biochemistry

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Q1) Which of the following statements is TRUE?

A) Monosaccharides are the building blocks of amino acids.

B) Nucleotides are the building blocks of polysaccharides.

C) Fatty acids are responsible for DNA replication.

D) Polysaccharides are the building blocks of amino acids.

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

Q2) Give the definition of biochemistry.

A) the study of the chemistry of living organisms

B) the study of the biology of dead organisms

C) the study of the chemistry of dead organisms

D) the study of the biology of living organisms

E) the study of polymers

Q3) How do cis- and trans-fats differ?

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

Q5) What is a codon?

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

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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) None of the above is true.

Q2) Which of the following statements is TRUE?

A) Coke is composed of charcoal and air.

B) Activated carbon has a low surface area.

C) Heating wood in the absence of air produces charcoal.

D) Carbon monoxide is nontoxic.

E) None of the above is true.

Q3) Identify the compound that smells like a rotten egg.

A) SO

B) H₂SO₂

C) H₂SO₄

D) S₄

E) H₂S

Q4) Describe the Frasch process for obtaining sulfur.

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Chapter 24: Metals and Metallurgy

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Sample Questions

Q1) Identify the element with the highest thermal conductivity.

A) copper

B) vanadium

C) silver

D) iron

Q2) Identify the metal that has the most brightly colored ions.

A) magnesium

B) manganese

C) copper

D) chromium

E) cobalt

Q3) Identify the element that has the most crystal structures at varying temperatures.

A) nickel

B) copper

C) vanadium

D) cobalt

E) manganese

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

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Identify the geometry of [Zn(NH₃)₄]⁺².

A) tetrahedral

B) octahedral

C) linear

D) square planar

E) trigonal planar

Q2) Identify the isomers that have ligands with different spatial arrangements about the metal ions.

A) linkage isomers

B) geometric isomers

C) coordination isomers

D) optical isomers

E) structural isomers

Q3) What type of geometry (according to valence bond theory)does V exhibit in the complex ion: [V(NH₃)₄]²⁺?

A) square bipyramidal

B) see-saw

C) trigonal pyramidal

D) bent

E) square planar

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