General Chemistry II Exam Bank - 2309 Verified Questions

General Chemistry II Exam Bank

Course Introduction

General Chemistry II builds upon the foundational principles introduced in General Chemistry I, delving deeper into topics such as chemical kinetics, chemical equilibrium, acid-base theory, thermodynamics, electrochemistry, and coordination chemistry. The course emphasizes quantitative problem solving and practical laboratory techniques, helping students develop a deeper understanding of the molecular basis of chemical reactions. Through lectures, laboratory experiments, and discussion, students gain skills and knowledge that are essential for advanced studies in chemistry, biology, engineering, and related fields.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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24 Chapters

2309 Verified Questions

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Page 2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) You prepare 1000. mL of tea and transfer it to a 1.00 quart pitcher for storage. Which of the following statements is true?

A) The pitcher will be filled to 100% of its capacity with no tea spilled.

B) The pitcher will be filled to about 95% of its capacity.

C) The pitcher will be filled to about 50% of its capacity.

D) The pitcher will be completely filled and a small amount of tea will overflow.

E) The pitcher will be completely filled and most of the tea will overflow.

Answer: D

Q2) The average distance between the Earth and the Moon is 240,000 miles. Express this distance in kilometers.

A) 6.1 × 10⁵km

B) 5.3 × 10⁵km

C) 3.9 × 10⁵km

D) 1.5 × 10⁵km

E) 9.4 × 10⁴km

Answer: A

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3

Chapter 2: The Components of Matter

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Sample Questions

Q1) Which of the following is a metalloid?

A) carbon, C, Z = 6

B) sulfur, S, Z = 16

C) germanium, Ge, Z = 32

D) iridium, Z = 77

E) bromine, Br, Z = 35

Answer: C

Q2) The mass of a neutron is equal to the mass of a proton plus the mass of an electron.

A)True

B)False

Answer: False

Q3) Which one of the following is a polyatomic cation?

A) nitrate

B) chromate

C) permanganate

D) hydronium

E) potassium

Answer: D

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Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Phosphine, an extremely poisonous and highly reactive gas, will react with oxygen to form tetraphosphorus decaoxide and water. PH₃(g) + O₂(g) \(\to\) P₄O₁₀(s) + H₂O(g) [unbalanced]

Calculate the mass of P₄O₁₀(s) formed when 225 g of PH₃reacts with excess oxygen.

A) 1880 g

B) 940. g

C) 900. g

D) 470 g

E) 56.3 g

Answer: D

Q2) The number of hydrogen atoms in 0.050 mol of C₃H₈O₃ is

A) 3.0 × 10²² H atoms.

B) 1.2 × 10²³ H atoms.

C) 2.4 × 10²³ H atoms.

D) 4.8 × 10²³ H atoms.

E) none of the above.

Answer: C

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Page 5

Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) Which one of the following is not a redox reaction?

A) 2Na(s) + 2H₂O(l) \(\to\) 2NaOH(aq) + H₂(g)

B) H₂(g) + Cl₂(g) \(\to\) 2HCl(g)

C) 2H₂O₂(aq) \(\to\) 2H₂O(l) + O₂(g)

D) Fe₂O₃(s) + 3H₂SO₄(aq) \(\to\) Fe₂(SO₄)₃(aq) + 3H₂O(l)

E) 2KMnO₄(aq) + 10FeSO₄(aq) + 8H₂SO₄(aq) \(\to\)K₂SO₄(aq) + 2MnSO₄(aq) + 5Fe₂(SO₄)₃(aq) + 8H₂O(l)

Q2) Vinegar is a solution of acetic acid, CH₃COOH, dissolved in water. A 5.54-g sample of vinegar was neutralized by 30.10 mL of 0.100 M NaOH. What is the percent by weight of acetic acid in the vinegar?

A) 0.184%

B) 1.63%

C) 3.26%

D) 5.43%

E) 9.23%

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) When a closed-ended manometer is used for pressure measurements, and the closed end is under vacuum, the level of manometer liquid in the closed arm can never be lower than that in the other arm.

A)True

B)False

Q2) A gas consists of 85.7 % carbon and 14.3 % hydrogen, by weight. A sample of this gas weighing 0.673 g occupies 729 mL at a pressure of 720.0 mmHg and a temperature of 77°C. Calculate its empirical and molecular formulas.

A) CH, C₂H₂

B) CH₂, C₂H₄

C) CH₂, C₃H₆

D) CH₃, C₂H₆

E) CH₄, CH₄

Q3) According to the postulates of kinetic-molecular theory, the molecules of all gases at a given temperature have the same average kinetic energy.

A)True

B)False

Q4) State Avogadro's Law.

Q5) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650. torr.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Which of the following is not a state function?

A) internal energy

B) volume

C) work

D) pressure

E) enthalpy

Q2) The compound carbon suboxide, C₃O₂, is a gas at room temperature. Use the data supplied to calculate the heat of formation of carbon suboxide.

(Data: 2CO(g) + C(s) \(\to\) C₃O₂(g) \(\Delta\)H° = 127.3 kJ/mol

And: \(\Delta\)H_f^° of CO(g) = -110.5 kJ/mol)

A) 116.8

B) -93.7

C) 227.8

D) -348.3

E) 93.7

Q3) For all processes, both q and \(\Delta\)E will have the same sign. A)True B)False

Page 8

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) Which scientist first proposed that particles of matter could have wave properties?

A) Einstein

B) Planck

C) de Broglie

D) Compton

E) Heisenberg

Q2) Interference of light waves

A) separates light into its component colors.

B) creates a pattern of light and dark regions.

C) focuses a broad beam of light into a point.

D) bends light as it passes the edge of an object.

E) creates a laser beam.

Q3) Use the Rydberg equation to calculate the wavelength, in nm, of the least energetic (longest wavelength) line in the visible series (n₁ = 2) of the spectrum of atomic hydrogen.

Q4) Continuous spectra are characteristic of molecules in the gas phase.

A)True

B)False

Q5) What is the speed of an electron in m/s if its wavelength is 0.155 nm?

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Select the correct electron configuration for Te (Z = 52).

A) [Kr]5s²5p⁶4d⁸

B) [Kr]5s²5d¹⁰5p⁴

C) [Kr]5s²4d¹⁰5p⁶

D) [Kr]5s²4f¹⁴

E) [Kr]5s²4d¹⁰5p⁴

Q2) State Hund's rule, and show how it applies to the ground state of phosphorus atoms.

Q3) "Each electron in an atom must have its own unique set of quantum numbers" is a statement of

A) the aufbau principle.

B) the Pauli exclusion principle.

C) Hund's rule.

D) the periodic law.

E) Heisenberg's principle.

Q4) In moving down a group in the periodic table, the oxides of the elements become more acidic in nature.

A)True

B)False

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Select the compound with the lowest (i.e., least negative) lattice energy.

A) CsBr(s)

B) NaCl(s)

C) SrO(s)

D) CaO(s)

E) KBr(s)

Q2) Describe, with appropriate explanations, the key factors which affect the magnitude of the lattice energy of an ionic substance.

Q3) The more C-O and O-H bonds there are in a substance, the greater will be the amount of heat released when a fixed mass of the substance is burned.

A)True

B)False

Q4) Bond energy increases as bond order increases, for bonding between a given pair of atoms.

A)True

B)False

Q5) A hypothetical ionic substance will not form merely because it has a high lattice energy. Explain why, using energy-based arguments.

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) In which one of the following species is the central atom (the first atom in the formula) an exception to the octet rule?

A) NH₃

B) NH₄⁺

C) I₂

D) BH₄^-

E) SF₆

Q2) In which one of the following molecules are all the bonds single?

A) O₃

B) POCl₃

C) CO

D) COCl₂

E) N₂H₄

Q3) When resonance occurs, the bond lengths in a molecule fluctuate rapidly.

A)True

B)False

Q4) What is the shape of the PF₃ molecule? Explain your answer, using VSEPR theory.

Q5) Draw all important resonance structures of the nitrate ion, NO₃^-

Page 12

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Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) A carbon-carbon triple bond in a molecule may give rise to the existence of cis and trans isomers.

A)True

B)False

Q2) Valence bond theory predicts that xenon will use _____ hybrid orbitals in XeOF₄.

A) sp

B) sp²

C) sp³

D) sp³d

E) sp³d²

Q3) In molecular orbital theory, combination of two atomic orbitals produces two molecular orbitals.

A)True

B)False

Q4) In molecular orbital theory, combination of two 2p atomic orbitals may give rise to either \(\sigma\) or \(\pi\) type molecular orbitals.

A)True

B)False

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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes

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Sample Questions

Q1) A cubic unit cell has an edge length of 400. pm. The length of its body diagonal (internal diagonal) in pm is therefore

A) 512.

B) 566. C) 631.

D) 693.

E) 724.

Q2) Which of the following statements concerning a face-centered cubic unit cell and the corresponding lattice, made up of identical atoms, is incorrect?

A) The coordination number of the atoms in the lattice is 8.

B) The packing in this lattice is more efficient than for a body-centered cubic system.

C) If the atoms have radius r, then the length of the cube edge is\(\surd\)8 × r.

D) There are four atoms per unit cell in this type of packing.

E) The packing efficiency in this lattice and hexagonal close packing are the same.

Q3) Liquid ammonia boils at -33.4°C and has a heat of vaporization of 23.5 kJ/mol.

Calculate its vapor pressure at -50.0°C.

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Page 14

Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) The Tyndall effect

A) is observed in concentrated solutions.

B) is observed only in dilute solutions.

C) is observed in colloidal dispersions.

D) is caused by Brownian motion.

E) is used to determine the osmotic pressure of solutions.

Q2) Which of the following aqueous solutions will have the lowest osmotic pressure?

A) 0.10 m KOH

B) 0.10 m RbCl

C) 0.05 m CaSO₄

D) 0.05 m BaCl₂

E) 0.10 m K₂SO₄

Q3) Predict the ideal van't Hoff factor (i) for each of the following solutes:

a. CaCl₂

b. Fe₂(SO₄)₃

c. NH₄Cl

Q4) Raoult's Law relates the vapor pressure of a solvent above a solution to the mole fraction of the solvent and the vapor pressure of the pure solvent.

A)True

B)False

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) The halogens are

A) strong oxidizing agents.

B) strong reducing agents.

C) strong acids.

D) strong bases.

E) amphoteric.

Q2) When elements from a group exhibit more than one oxidation state,

A) the higher oxidation state is more important as one goes down a group.

B) the lower oxidation state is more important as one goes down a group.

C) both oxidation states are equally important throughout the group.

D) the oxidation state will be affected by the elements on either side of it in the period.

E) None of the above conclusions is true.

Q3) The smallest ionization energies are found in the ____________ ___________ region of the periodic table.

A) upper right

B) upper left

C) lower right

D) lower left

E) transition element (d-block)

16

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Chapter 15: Organic Compounds and the Atomic Properties

of Carbon

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Sample Questions

Q1) A characteristic reaction of alkenes is addition.

A)True

B)False

Q2) Secondary amines have the general formula RNH₂.

A)True

B)False

Q3) Which of the following polymers is a condensation polymer?

A) polystyrene

B) Teflon

C) polyvinylchloride

D) polypropylene

E) Dacron

Q4) The backbone of protein molecules consists of repeating N-C-C-O units.

A)True

B)False

Q5) Both nylons and proteins are polyamides.

A)True

B)False

Q7) Name the three component parts of a nucleotide. Page 17

Q6) Draw and name two non-alkenes with the formula C₄H₈.

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) The rate law for the reaction 3A \(\to\) 2B is rate = k[A] with a rate constant of 0.0447 hr^-1. What is the half-life of the reaction?

A) 0.0224 hr

B) 0.0645 hr

C) 15.5 hr

D) 22.4 hr

E) 44.7 hr

Q2) Briefly list the features/properties common to all catalysts and how they work. Draw a labeled reaction energy diagram as part of your answer.

Q3) The decomposition of dinitrogen pentaoxide to nitrogen dioxide and oxygen follows first-order kinetics and has an activation energy of 102 kJ/mol. By what factor will the fraction of collisions with energy greater than or equal to the activation energy increase if the reaction temperature goes from 30°C to 60°C?

A) 1.00

B) 1.10

C) 2.00

D) 4.00

E) 38.4

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) The reaction quotient for a gas phase reaction has a value of 2000. If the number of moles of reactants in the reaction equation is equal to that of the products, which of the following statements is definitely true?

A) The reaction must proceed to the left to establish equilibrium.

B) The reaction must proceed to the right to establish equilibrium.

C) When the system is at equilibrium, the concentrations of the products will be much larger than the concentrations of the reactants.

D) The concentrations of the products are generally larger than the concentrations of the reactants.

E) None of the above statements is true.

Q2) Although a system may be at equilibrium, the rate constants of the forward and reverse reactions will in general be different.

A)True

B)False

Q3) A good catalyst for a reaction will speed up the forward reaction and slow down the reverse reaction.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) In order to be a Brønsted-Lowry base, a species must contain a proton.

A)True

B)False

Q2) Lactic acid has a pK_a of 3.08. What is the approximate degree of dissociation of a 0.35 M solution of lactic acid?

A) 1.1%

B) 2.2%

C) 4.8%

D) 14%

E) none of the above

Q3) What is the pOH of a 0.0250 M HI solution?

A) 0.944

B) 1.602

C) 12.398

D) 13.056

E) none of the above

Q4) All strong acids have weak conjugate bases.

A)True

B)False

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) A 25.0-mL sample of 1.00 M NH₃is titrated with 0.15 M HCl. What is the pH of the solution after 15.00 mL of acid have been added to the ammonia solution? K_b = 1.8 × 10^-5

A) 10.26

B) 9.30

C) 9.21

D) 8.30

E) 8.21

Q2) Calculate the solubility of silver oxalate, Ag₂C₂O₄, in pure water. K_sp = 1.0 × 10^-11

A) 1.4 × 10^-4M

B) 8.2 × 10^-5M

C) 5.4 × 10^-5M

D) 3.2 × 10^-6M

E) 2.5 × 10^-12M

Q3) What is the pH of 375 mL of solution containing 0.150 mol of propenoic acid (HA) and 0.250 mol of sodium propenoate (NaA)? (K_a for propenoic acid is 5.52 × 10^-5.)

Q4) Make a clear distinction between buffer range and buffer capacity.

Page 22

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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions

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Q1) For a chemical reaction to be spontaneous only at low temperatures, which of the following conditions must be met?

A) .\(\Delta\)S° > 0, \(\Delta\)H° > 0

B) .\(\Delta\)S° > 0, \(\Delta\)H° < 0

C) .\(\Delta\)S° < 0, \(\Delta\)H° < 0

D) .\(\Delta\)S° < 0, \(\Delta\)H° > 0

E) .\(\Delta\)G° > 0

Q2) For a chemical reaction to be spontaneous only at high temperatures, which of the following conditions must be met?

A) .\(\Delta\)S° > 0, \(\Delta\)H° > 0

B) .\(\Delta\)S° > 0, \(\Delta\)H° < 0

C) .\(\Delta\)S° < 0, \(\Delta\)H° < 0

D) .\(\Delta\)S° < 0, \(\Delta\)H° > 0

E) .\(\Delta\)G° > 0

Q3) A chemical reaction has \(\Delta\)H° = 42.8 kJ and \(\Delta\)S° = 92.5 J/K, at 25°C. Calculate the temperature at which \(\Delta\)G° = 0. State any approximation involved in your calculation.

Q4) State the second and third laws of thermodynamics.

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) A voltaic cell prepared using zinc and iodine has the following cell notation. Zn(s) | Zn²⁺(aq) || I^-(aq) | I₂(s) | C(graphite)

Which of the following equations correctly represents the balanced, spontaneous, cell reaction?

A) 2I^-(aq) + Zn²⁺(aq) \(\to\) I₂(s) + Zn(s)

B) I₂(s) + Zn(s) \(\to\) 2I^-(aq) + Zn²⁺(aq)

C) 2I^-(aq) + Zn(s) \(\to\) I₂(s) + Zn²⁺(aq)

D) I₂(s) + Zn²⁺(aq) \(\to\) 2I^-(aq) + Zn(s)

E) None of the above, since graphite must be in the equation.

Q2) Calculate the potential of a voltaic cell (E°_cell) if it is required to do 5.43 × 10^-3 kJ of work when a charge of 2.50 C is transferred.

A) 2.17 × 10³ V

B) 2.17 × 10^-3 V

C) 2.17 V

D) 13.6 V

E) 1.36 × 10^-2 V

Q3) Write down equations representing the anode half-reaction, the cathode half-reaction, and the overall cell reaction for the lead-acid battery.

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Chapter 22: The Elements in Nature and Industry

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Q1) Plants extract phosphate from the soil

A) by converting it to the dihydrogen phosphate ion by addition of acid to the soil near roots.

B) by converting it to phosphoric acid.

C) by osmosis.

D) by leaching.

E) by fixation.

Q2) Pyrometallurgy uses __________________ to separate a metal from its ore.

A) solid phase chemical properties

B) electrical processes

C) thermal processes

D) aqueous chemical processes

E) explosives

Q3) What gas is produced at the anode in the Downs cell, in which molten NaCl is electrolyzed?

Q4) Aluminum and magnesium form are light metals which are used in many structural applications. Although they are highly active metals (as their electrode potentials suggest), they can often be used without the use of paint or other applied finishes to protect them against corrosion. How is this possible?

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) Which one of the following has the ground state electron configuration [Ar]3d¹⁰4s¹?

A) In⁺

B) Cd²⁺

C) Ag⁺

D) Ag

E) Cu

Q2) The compound Rh(CO)(H)(PH₃)₂ forms cis and trans isomers. Use this information to predict the geometry of this complex, and draw the geometric isomers.

Q3) What is the highest possible oxidation state for molybdenum, Mo?

A) +2

B) +4

C) +6

D) +8

E) none of the above

Q4) Of the 3d transition series of elements, zinc has the greatest atomic radius.

A)True

B)False

Q5) Why is the +2 oxidation state so common among transition elements?

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) No alpha decay is observed for isotopes of elements with Z < 83.

A)True

B)False

Q2) Bombardment of uranium-238 nuclei by carbon-12 nuclei produces californium-246 and neutrons. Write a complete, balanced equation for this nuclear process.

Q3) An isotope with a low value of N/Z will generally decay through

A) .\(\alpha\) decay.

B) .\(\beta\) decay.

C) .\(\gamma\) decay.

D) electron capture.

E) spontaneous fission.

Q4) Which one of the following elements is formed largely in supernova explosions?

A) H

B) He

C) Mg

D) Fe

E) U

Q5) The r-process occurs during supernova explosions.

A)True

B)False

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