General Chemistry II Exam Materials - 1830 Verified Questions

General Chemistry II

Exam Materials

Course Introduction

General Chemistry II builds upon the foundational principles introduced in General Chemistry I, delving deeper into key chemical concepts and phenomena. The course explores topics such as kinetics, chemical equilibrium, acid-base equilibria, thermodynamics, electrochemistry, and coordination compounds. Emphasis is placed on quantitative problem-solving, laboratory techniques, and the application of theoretical principles to real-world chemical systems. Through lectures, laboratory experiments, and collaborative activities, students will develop a deeper understanding of chemical reactions, measurement of chemical changes, and their significance in biological, environmental, and industrial contexts.

Recommended Textbook Principles of General Chemistry 2nd Edition by Martin Silberberg

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Chapter 1: Keys to the Study of Chemistry

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Q1) The area of a 15-inch pizza is 176.7 in².Express this area in square centimeters.

A)1140.cm²

B)448.8 cm²

C)96.8 cm²

D)69.57 cm²

E)27.39 cm²

Answer: A

Q2) An evacuated 276 mL glass bulb weighs 129.6375 g.Filled with an unknown gas,the bulb weighs 130.0318 g.Calculate the gas density in g/L,and express it with an appropriate number of significant figures.

Answer: 1.43 g/L

Q3) Which of the following activities is not a part of good science?

A)proposing a theory

B)developing a hypothesis

C)making quantitative observations

D)designing experiments

E)indulging in speculation

Answer: E

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Page 3

Chapter 2: The Components of Matter

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Q1) State the two important experimental results (and the names of the responsible scientists)which enabled the mass of the electron to be determined.

Answer: Thomson measured m/e,the mass-to-charge ratio.Millikan measured e,the charge.Thus,the mass m could be calculated.

Q2) Rutherford bombarded gold foil with alpha (\(\alpha\))particles and found that a small percentage of the particles were deflected.Which of the following was not accounted for by the model he proposed for the structure of atoms?

A)the small size of the nucleus

B)the charge on the nucleus

C)the total mass of the atom

D)the existence of protons

E)the presence of electrons outside the nucleus

Answer: C

Q3) All neutral atoms of tin have 50 protons and 50 electrons.

A)True

B)False

Answer: True

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4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Hydroxylamine nitrate contains 29.17 mass % N,4.20 mass % H,and 66.63 mass %

O.Determine its empirical formula.

A)HNO

B)H₂NO₂

C)HN₆O₁₆

D)HN₁₆O₇

E)H₂NO₃

Answer: B

Q2) In the combustion analysis of 0.1127 g of glucose (C₆H₁₂O₆),what mass,in grams,of CO₂ would be produced?

A)0.0451 g

B)0.0825 g

C)0.1652 g

D)0.4132 g

E)1.466 g

Answer: C

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Page 5

Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) The oxidation numbers of P,S and Cl in H₂PO₂^-,H₂S and KClO₄ are,respectively

A)-1,-1,+3

B)+1,-2,+7

C)+1,+2,+7

D)-1,-2,+7

E)-1,-2,+3

Q2) Select the classification for the following reaction.

H₂CO₃(aq) \(\to\) H₂O(l)+ CO₂(g)

A)combination

B)decomposition

C)displacement

D)acid-base

E)None of these choices is correct.

Q3) A combination reaction may also be a displacement reaction.

A)True

B)False

Q4) The combustion of an element is always a combination reaction.

A)True

B)False

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Q1) From the postulates of kinetic-molecular theory,it follows that the molecules of all gases at a given temperature have the same average kinetic energy.

A)True

B)False

Q2) Aluminum metal shavings (10.0 g)are placed in 100.mL of 6.00 M hydrochloric acid.What is the maximum volume of hydrogen,measured at STP,which can be produced?

2Al(s)+ 6HCl(aq) \(\to\) 2AlCl₃(aq)+ 3H₂(g)

Q3) For an ideal gas,a plot of PV/nRT versus P gives a straight line with a positive slope. A)True B)False

Q4) At a temperature of absolute zero,the volume of an ideal gas is zero. A)True B)False

Q5) The rate of diffusion of a gas is inversely proportional to its molar mass. A)True

B)False

Q6) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650.torr.

Q7) State Avogadro's Law.

Page 7

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Calculate q when 28.6 g of water is heated from 22.0°C to 78.3°C.

A)0.385 kJ

B)1.61 kJ

C)6.74 kJ

D)9.37 kJ

E)1.61 \(\times\) 10³ kJ

Q2) If,as a pioneer,you wished to warm your room by taking an object heated on top of a pot-bellied stove to it,which of the following 15-pound objects,each heated to 100°C,would be the best choice? The specific heat capacity (in J/(g·K))for each substance is given in parentheses.Iron (0.450),copper (0.387),granite (0.79),gold (0.129),water (4.18).

A)iron

B)copper

C)granite

D)gold

E)water

Q3) a.Explain fully what is meant by the term "state function."

b.(i)Give two examples of thermodynamic quantities which are state functions. (ii)Give two examples of thermodynamic quantities which are not state functions.

Page 8

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) In the quantum mechanical treatment of the hydrogen atom,the energy depends on the principal quantum number n but not on the values of l or m_l.

A)True

B)False

Q2) According to the Rydberg equation,the line with the shortest wavelength in the emission spectrum of atomic hydrogen is predicted to lie at a wavelength (in nm)of

A)91.2 nm

B)1.10 \(\times\) 10^-2 nm

C)1.10 \(\times\) 10² nm

D)1.10 \(\times\) 10¹⁶ nm

E)None of these choices is correct.

Q3) Green light has a wavelength of 5200 Å.Calculate the energy of one photon of green light.

A)3.4 \(\times\) 10^-40 J

B)3.4 \(\times\) 10^-30 J

C)3.8 \(\times\) 10^-29 J

D)3.4 \(\times\) 10^-27 J

E)3.8 \(\times\) 10^-19 J

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9

Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) "Electrons added to atomic orbitals of the same energy will remain unpaired with parallel spins until the subshell is more than half-filled" is a statement of A)the aufbau principle.

B)Hund's rule.

C)the Pauli exclusion principle.

D)the periodic law.

E)the singularity rule.

Q2) In a single atom,what is the maximum number of electrons which can have quantum number n = 4?

A)16

B)18

C)32

D)36

E)None of these choices is correct.

Q3) State Hund's rule,and show how it applies to the ground state of phosphorus atoms.

Q4) Write down the full electron configuration for ground state atoms of a.S (element 16).

b.Pd (element 46).

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Chapter 9: Models of Chemical Bonding

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Q1) Arrange the following bonds in order of increasing bond strength.

A)C-I < C-Br < C-Cl < C-F

B)C-F < C-Cl < C-Br < C-I

C)C-Br < C-I < C-Cl < C-F

D)C-I < C-Br < C-F< C-Cl

E)None of these orders is correct.

Q2) Which of the following contains covalent bonds?

A)BaO

B)IBr

C)Mg

D)LiBr

E)Cu

Q3) Analysis of an unknown substance showed that it has a high boiling point and is brittle.It is an insulator as a solid but conducts electricity when melted.Which of the following substances would have those characteristics?

A)HCl

B)Al

C)KBr

D)SiF₄

E)I₂

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the actual bond angle in SeCl₂ using the VSEPR theory.

A)more than 120°

B)between 109° and 120°

C)between 90° and 109°

D)exactly 90°

E)less than 90°

Q2) What is the molecular shape of NO₂^- as predicted by the VSEPR theory?

A)linear

B)trigonal planar

C)bent

D)tetrahedral

E)resonant

Q3) According to VSEPR theory,a molecule with the general formula AX₆ will have a ______ molecular shape.

A)tetrahedral

B)trigonal planar

C)trigonal bipyramidal

D)hexagonal

E)octahedral

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Chapter 11: Theories of Covalent Bonding

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Q1) According to molecular orbital (MO)theory,the twelve outermost electrons in the O₂ molecule are distributed as follows:

A)12 in bonding MOs,0 in antibonding MOs.

B)10 in bonding MOs,2 in antibonding MOs.

C)9 in bonding MOs,3 in antibonding MOs.

D)8 in bonding MOs,4 in antibonding MOs.

E)7 in bonding MOs,5 in antibonding MOs.

Q2) In the valence bond treatment,a \(\pi\) bond is formed when two 2p orbitals overlap side to side.

A)True

B)False

Q3) Valence bond theory predicts that bromine will use _____ hybrid orbitals in BrF₅.

A)sp²

B)sp³

C)sp³d

D)sp³d²

E)None of these choices is correct.

Q4) In one sentence state how molecular orbitals are usually obtained.

Q5) In one sentence state the basic principle of valence bond theory.

Page 13

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

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Sample Questions

Q1) Which of the following liquids is likely to have the highest surface tension?

A)Br₂

B)C₈H₁₈

C)CH₃OCH₃

D)CH₃OH

E)Pb

Q2) In cubic closest packing,the unit cell is body-centered cubic.

A)True

B)False

Q3) Use molecular orbital band diagrams to explain why metals are good conductors but semiconductors are not.

Q4) In an ionic solid MX consisting of the monatomic ions,M⁺ and X^-,the coordination number of M⁺ is _______________.

A)1

B)2

C)6

D)8

E)impossible to predict without knowing the crystal structure of MX

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Q5) How do the electrical properties of semiconductors differ from those of metals?

Chapter 13: The Properties of Solutions

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Q1) The mole fraction of potassium nitrate in an aqueous solution is 0.0194.The solution's density is 1.0627 g/mL.Calculate the molarity of the solution.

A)0.0194 M

B)0.981 M

C)1.05 M

D)1.96 M

E)19.4 M

Q2) Which of the following sets of conditions could exist when two liquids which are completely miscible in one another are mixed?

A)(\(\Delta\)H_soln > 0,entropy of system decreases)

B)(\(\Delta\)H_soln \(\approx\) 0,entropy of system decreases)

C)(\(\Delta\)H_soln \(\approx\) 0,entropy change of system \(\approx\) 0)

D)(\(\Delta\)H_soln \(\approx\) 0,entropy of system increases)

E)None of these choices is correct.

Q3) A 0.137 m solution of sodium ferrocyanide (NaFe(CN)₆.10H₂O)in water has a freezing point depression of 0.755°C.Calculate the van't Hoff factor (i)under these conditions.(K_f = 1.86°C/m)

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Chapter 14: The Main-Group Elements: Applying Principles of Bonding and Structure

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Sample Questions

Q1) Phosphoric acid (H₃PO₄)is a strong acid.

A)True

B)False

Q2) Write balanced equations,showing all reactants and products,to represent a.the reaction of calcium metal with water.

b.the reaction of aluminum metal with oxygen gas.

Q3) Consider the oxides XO₂,where X is one of the first 20 elements.Identify X in each of the following cases:

a.XO₂ is a natural component of the atmosphere,whose buildup is implicated in climate change.

b.XO₂ is a toxic,brown gas which is a serious air pollutant.

c.XO₂ is a colorless,toxic gas formed during the combustion of coal.

d.XO₂ is network solid;the element X itself is also a network solid.

Q4) If an alkali metal is represented by the symbol M,what is the formula of a.the oxide of M?

b.the peroxide of M?

c.the superoxide of M?

Q5) Ionic hydrides do not have exact (stoichiometric)formulas. A)True B)False

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Q1) Ethane (C₂H₆)is much more reactive than disilane (Si₂H₆).

A)True

B)False

Q2) a.Draw two different structures with the molecular formula C₂H₆O.

b.Name the functional group in each structure.

c.Which one will have the higher boiling point,and why?

Q3) Describe and contrast fibrous and globular proteins in terms of their amino acid composition,their structure and their function.

Q4) Draw and name four alkenes with the formula C₄H₈.

Q5) In one sentence,what is the general requirement for a molecule to be optically active?

Q6) Explain what is meant by "complementary" in the context of DNA strands.

Q7) A characteristic reaction of alkenes is addition. A)True B)False

Q8) Draw and name all stable molecules with the formula C₅H₁₂. Page 17

Q9) Draw and name two non-alkenes with the formula C₄H₈.

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Q1) Butadiene,C₄H₆ (used to make synthetic rubber and latex paints)dimerizes to C₈H₁₂ with a rate law of rate = 0.014 L/(mol·s)[C₄H₆]².What will be the concentration of C₄H₆ after 3.0 hours if the initial concentration is 0.025 M?

A)0.0052 M

B)0.024 M

C)43 M

D)190 M

E)0.0000 M

Q2) The decomposition of SOCl₂ is first-order in SOCl₂.If the half-life for the reaction is 4.1 hr,how long would it take for the concentration of SOCl₂ to drop from 0.36 M to 0.045 M?

A)0.52 hr

B)1.4 hr

C)12 hr

D)33 hr

E)> 40 hr

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Page 19

Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) For some gas-phase reactions,K_p = K_c.

A)True

B)False

Q2) The two equilibrium constants for the same reaction,K_c and K_p,will always equal one another when

A)all of the reactants and products are gases.

B)in the reaction equation,the number of moles of gaseous products equals the number of moles of gaseous reactants.

C)in the reaction equation,the number of moles of gaseous products is greater than the number of moles of gaseous reactants.

D)in the reaction equation,the number of moles of gaseous products is smaller than the number of moles of gaseous reactants.

E)in the reaction equation,the total number of moles of reactants equals that of the products.

Q3) 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

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Chapter 18: Acid-Base Equilibria

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Q1) A solution is prepared by adding 0.10 mol of lithium nitrate,LiNO₃,to 1.00 L of water.Which statement about the solution is correct?

A)The solution is basic.

B)The solution is neutral.

C)The solution is weakly acidic.

D)The solution is strongly acidic.

E)The values for K_a and K_b for the species in solution must be known before a prediction can be made.

Q2) The substance H₂SO₃ is considered

A)a weak Arrhenius base.

B)a strong Arrhenius acid.

C)a strong Arrhenius base.

D)a neutral compound.

E)a weak Arrhenius acid.

Q3) Define an acid according to the Arrhenius theory,and write a balanced equation to support this definition.

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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Q1) When a weak acid is titrated with a weak base,the pH at the equivalence point

A)is greater than 7.0.

B)is equal to 7.0.

C)is less than 7.0.

D)is determined by the sizes of K_a and K_b.

E)is no longer affected by addition of base.

Q2) Buffer solutions with the component concentrations shown below were prepared.Which of them should have the lowest pH?

A)[CH₃COOH] = 0.25 M,[CH₃COO^-] = 0.25 M

B)[CH₃COOH] = 0.75 M,[CH₃COO^-] = 0.75 M

C)[CH₃COOH] = 0.75 M,[CH₃COO^-] = 0.25 M

D)[CH₃COOH] = 0.25 M,[CH₃COO^-] = 0.75 M

E)[CH₃COOH] = 1.00 M,[CH₃COO^-] = 1.00 M

Q3) Equal volumes of the following pairs of solutions are mixed.Which pair will produce a buffer solution?

A)0.10 mol L^-1 HCl and 0.05 mol L^-1 NaOH

B)0.10 mol L^-1 HCl and 0.15 mol L^-1 NH₃

C)0.10 mol L^-1 HCl and 0.05 mol L^-1 NH₃

D)0.10 mol L^-1 HCl and 0.20 mol L^-1 CH₃COOH

E)0.10 mol L^-1 HCl and 0.20 mol L^-1 NaCl

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Chapter

Direction of Chemical Reactions

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Q1) a.Explain what is meant by a spontaneous process.

b.Is a spontaneous process necessarily a rapid one? Explain,and provide a real reaction as an example to illustrate your answer.

Q2) In order for a process to be spontaneous,

A)(\(\Delta\)H must be less than zero.)

B)(\(\Delta\)S must be greater than zero.)

C)(\(\Delta\)G must be greater than zero.)

D)it should be rapid.

E)(\(\Delta\)S_sys + \(\Delta\)S_surr must be greater than zero.)

Q3) For a chemical reaction to be spontaneous at all 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)(It is not possible for a reaction to be spontaneous at all temperatures.

Q4) State the second and third laws of thermodynamics.

Page 23

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) Cryolite,Na₃AlF₆ is used in the electrolysis of aluminum oxide because

A)it is a good source of fluoride ions.

B)it provides a source of fluorine,an oxidizing agent.

C)it provides a source of sodium,a reducing agent.

D)it reduces the energy requirement of the process,by lowering the melting point.

E)it is very soluble in water.

Q2) Calculate E°_cell for the reaction of nickel(II)ions with cadmium metal at 25°C.K = 1.17 \(\times\) 10⁵ Ni²⁺(aq)+ Cd(s) \(\to\) Cd²⁺(aq)+ Ni(s)

A)0.075 V

B)0.10 V

C)0.12 V

D)0.15 V

E)0.30 V

Q3) Oxidation occurs at the cathode of a galvanic cell,but at the anode of an electrolytic cell.

A)True

B)False

Page 24

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Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) a.Explain how the crystal field theory can use the magnitude of the splitting energy \(\Delta\) to provide an explanation of the color and magnetic properties of octahedral complexes.

b.In promoting an electron from the t₂_g set of orbitals to the e_g set,an octahedral complex absorbs a photon with a wavelength \(\lambda\) of 523 nm.Calculate the value of \(\Delta\) in the complex,in kJ/mol.

Q2) The ground state electron configuration of Zn²⁺ is:

A)[Ar]4s²3d⁸

B)[Ar]4s²3d¹⁰

C)[Ar]4s¹3d⁹

D)[Ar]3d¹⁰

E)[Ar]3d⁸

Q3) Write the formula for pentaamminechlorocobalt(III)chloride.

A)[Co(NH₃)₅Cl]Cl

B)[Co(NH₃)₅Cl]Cl₂

C)[Co(NH₃)₅Cl]Cl₃

D)[Co(NH₃)₅Cl]Cl₄

E)None of these choices is correct.

Q4) Why is the +2 oxidation state so common among transition elements?

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) A certain isotope has a specific activity of 7.29 \(\times\) 10^-4 Ci/g.How many \(\alpha\) particles will a 75.0 mg sample emit in one hour?

A)9.99 \(\times\) 10⁴

B)2.02 \(\times\) 10⁶

C)7.28 \(\times\) 10⁹

D)1.29 \(\times\) 10¹²

E)None of these choices is correct.

Q2) Iodine-131,t_1/2 = 8.0 days,is used in diagnosis and treatment of thyroid gland diseases.If a laboratory sample of iodine-131 initially emits 9.95 \(\times\) 10¹⁸ \(\beta\) particles per day,how long will it take for the activity to drop to 6.22 \(\times\) 10¹⁷ \(\beta\) particles per day?

A)2.0 days

B)16 days

C)32 days

D)128 days

E)None of these choices is correct.

Q3) Gamma rays are not deflected by an electric field.

A)True

B)False

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