Honors General Chemistry Final Test Solutions - 1830 Verified Questions

Honors General Chemistry

Final Test Solutions

Course Introduction

Honors General Chemistry offers an in-depth and rigorous exploration of fundamental chemical principles, including atomic and molecular structure, chemical bonding, stoichiometry, thermodynamics, kinetics, and equilibrium. Designed for students with strong backgrounds and keen interest in science, this course integrates advanced laboratory techniques, critical analysis, and problem-solving skills. Emphasis is placed on both theoretical understanding and practical applications, preparing students for future study and research in chemistry, medicine, engineering, and related fields. Collaborative projects and discussions further enhance students analytical and experimental abilities in a challenging and supportive environment.

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

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23 Chapters

1830 Verified Questions

1830 Flashcards

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Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) Given that 1 inch = 2.54 cm,1 cm³ is equal to

A)16.4 in³

B)6.45 in³

C)0.394 in³

D)0.155 in³

E)0.0610 in³

Answer: E

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) The potential energy of a car moving on a level road does not depend on its speed.

A)True

B)False

Answer: True

Q4) The ripening of fruit,once picked,is an example of physical change.

A)True

B)False Answer: False

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Chapter 2: The Components of Matter

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Sample Questions

Q1) Silver chloride is used in photographic emulsions.What is its formula?

A)Ag₂Cl₃

B)Ag₂Cl

C)AgCl₃

D)AgCl₂

E)AgCl

Answer: E

Q2) The compound,NaH₂PO₄,is present in many baking powders.What is its name?

A)sodium biphosphate

B)sodium hydrogen phosphate

C)sodium dihydrogen phosphate

D)sodium hydrophosphate

E)sodium dihydride phosphate

Answer: C

Q3) The molecular formula of a compound provides more information than its structural formula.

A)True

B)False

Answer: False

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Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) a.A solution of common salt,NaCl,in water has a concentration of 0.0921 M.Calculate the number of moles of HCl contained in 50.0 mL of this solution.

b.If,instead,an NaCl solution is prepared by dissolving 10.0 g of solid NaCl in enough water to make 250.mL of solution,what is the molarity?

Answer: a.0.00461 g

b.0.684 M

Q2) Gadolinium oxide,a colorless powder which absorbs carbon dioxide from the air,contains 86.76 mass % Gd.Determine its empirical formula.

A)Gd₂O₃

B)Gd₃O₂

C)Gd₃O₄

D)Gd₄O₃

E)GdO

Answer: A

Q3) Analysis of a white solid produced in a reaction between chlorine and phosphorus showed that it contained 77.44% chlorine and 22.56% phosphorus.What is its empirical formula?

Answer: PCl₃

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Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) Which of the following will be least soluble in water?

A)potassium sulfate,K₂SO₄

B)ammonium nitrate,NH₄NO₃

C)chloromethane,CH₃Cl

D)calcium chloride,CaCl₂

E)ethanol,C₂H₆O

Q2) a.Define "oxidation".

b.Identify the oxidizing and reducing agents in the following (unbalanced)equation. HNO₃(aq)+ C₂H₆O(aq)+

K₂Cr₂O₇(aq) \(\to\)KNO₃(aq)+

C₂H₄O(aq)+ H₂O(l)+ Cr(NO₃)₃(aq)

Q3) Select the classification for the following reaction. Fe(s)+ 2Fe³⁺(aq) \(\to\) 3Fe²⁺(aq)

A)precipitation

B)acid-base

C)redox

D)decomposition

E)None of these choices is correct.

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) A carbon dioxide sample weighing 44.0 g occupies 32.68 L at 65°C and 645 torr.What is its volume at STP?

A)22.4 L

B)31.1 L

C)34.3 L

D)35.2 L

E)47.7 L

Q2) In the fermentation process,yeast converts glucose to ethanol and carbon dioxide.What volume of carbon dioxide,measured at 745 torr and 25.0°C,can be produced by the fermentation of 10.0 g of glucose?

C₆H₁₂O₆(aq) \(\to\) 2C₂H₅OH(aq)+ 2CO₂(g)

Q3) For a gas obeying Boyle's Law,a plot of V versus 1/P will give a straight line passing through the origin.

A)True

B)False

Q4) For real gases,PV > nRT,always.

A)True

B)False

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) 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

Q2) For a reaction in a sealed,rigid container,\(\Delta\)H is always greater than \(\Delta\)E.

A)True

B)False

Q3) a.Starting from the equation H = E + PV,show how the relationship \(\Delta\)H = q_p is derived.Clearly indicate any necessary assumptions or conditions. b.In one sentence,state in full what is meant by the equation: \(\Delta\)H = q_p.

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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) What are the possible values for the following quantum numbers in an atom?

a.n

b.l

c.m_l

Q3) In the Rydberg equation,for a fixed value of n₁,the longest wavelength line has n₂ = \(\infty\)

A)True

B)False

Q4) The energy of a photon is directly proportional to the wavelength of the radiation.

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) Write down the maximum number of electrons in an atom which can have a.quantum number n = 4. b.orbital designation 3d. c.orbital designation 2p_z.

Q2) Which one of the following equations correctly represents the process relating to the ionization energy of X?

A)X(s) \(\to\) X⁺(g)+ e^-

B)X₂(g) \(\to\) X⁺(g)+ X^-(g)

C)X(g)+ e^- \(\to\) X^-(g)

D)X^-(g) \(\to\) X(g)+ e^-

E)X(g) \(\to\) X⁺(g)+ e^-

Q3) Which of the following elements has the largest first ionization energy?

A)Na

B)Cl

C)Ca

D)Te

E)Br

Q4) Electron affinities of neutral atoms may be positive or negative.

A)True

B)False

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Give a clear and concise definition of the term "electronegativity";i.e. ,what does it measure?

Q2) Select the most polar bond amongst the following.

A)C-O

B)Si-F

C)Cl-F

D)C-F

E)C-I

Q3) Select the compound with the lowest lattice energy.

A)CsBr(s)

B)NaCl(s)

C)SrO(s)

D)CaO(s)

E)KBr(s)

Q4) In not more than three sentences,describe the electron arrangement responsible for bonding in solid SrCl₂.

Q5) When an atom is represented in a Lewis electron dot symbol,the element symbol represents ______________ and the dots represent ______________.

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the actual bond angles in SF₃⁺ using the VSEPR theory.

A)more than 120°

B)exactly 120°

C)between 109° and 120°

D)between 90° and 109°

E)less than 90°

Q2) Predict the ideal bond angles around carbon in C₂I₂ using the molecular shape given by the VSEPR theory.

A)90°

B)109°

C)120°

D)180°

E)None of these choices is correct.

Q3) Explain what is meant by "dipole moment",and give an example of a molecule which has polar bonds but which does not itself have a dipole moment.

Q4) Draw the Lewis structure of XeF₄.Use this structure,in conjunction with VSEPR theory,to predict the shape of this molecule.Outline your reasoning.

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) Explain what is meant by the term "bond order" and describe how it can be calculated using the information in a molecular orbital energy level diagram.

Q2) A molecule with the formula AX₄ uses _________ to form its bonds.

A)sp hybrid orbitals

B)sp² hybrid orbitals

C)sp³ hybrid orbitals

D)sp³d hybrid orbitals

E)sp³d² hybrid orbitals

Q3) In the context of molecular orbital (MO)theory,explain how atomic p orbitals can give rise to:

a.a \(\sigma\) MO

b.a \(\pi\) MO

Q4) Valence bond theory predicts that xenon will use _____ hybrid orbitals in XeOF₄.

A)sp

B)sp²

C)sp³

D)sp³d

E)sp³d²

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

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Sample Questions

Q1) Which of the following substances will have hydrogen bonds between molecules?

A)(CH₃)₃N

B)CH₃-O-CH₃

C)CH₃CH₂-OH

D)CH₃CH₂-F

E)HI

Q2) Assuming that atoms are spherical,calculate the fraction of space which is occupied by atoms (i.e. ,the packing efficiency)in a metal with a simple cubic unit cell.

Q3) Of the five major types of crystalline solid,which would you expect each of the following to form? (e.g. ,H₂O: molecular)

a.Sn

b.Si

c.KCl

d.Xe

e.F₂

Q4) Mercury melts at -39°C and boils at 357°C.Draw a diagram of the heating curve of mercury.Label all lines and axes,and clearly indicate the melting and boiling points on your diagram.

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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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Sample Questions

Q1) Calculate the molarity of a solution prepared by diluting 1.85 L of 6.5 M KOH to 11.0 L.

A)0.28 M

B)0.91 M

C)1.1 M

D)3.1 M

E)3.9 M

Q2) From the following list of aqueous solutions and water,select the one with the lowest freezing point.

A)0.75 M (NH₄)₃PO₄

B)l.0 M CaSO₄

C)l 0 M LiClO₄

D)1.5 M CH₃OH,methyl alcohol

E)pure water

Q3) The solubility of gases in water increases with increase in the pressure of the gas.

A)True

B)False

Q4) The density of pure water at 25°C is 0.997 g/mL.Considering water as being both solvent and solute,calculate its molarity.

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Page 15

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

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Sample Questions

Q1) Which of the following oxides is most basic?

A)As₂O₃

B)P₄O₁₀

C)Sb₂O₃

D)Sb₂O₅

E)NO₂

Q2) Dinitrogen monoxide,N₂O,is

A)a brown poisonous gas that is one of the chemicals involved in the production of photochemical smog.

B)a colorless gas used in the production of nitric acid.

C)a colorless gas used as an aerosol propellant.

D)a colorless gas that disproportionates into nitrogen and oxygen.

E)None of these choices is correct.

Q3) Ionic hydrides do not have exact (stoichiometric)formulas.

A)True

B)False

Q4) The halogens act as oxidizing agents in most of their reactions.

A)True

B)False

Q5) Name two different classes (types)of hydride.

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) Esters can be formed by the dehydration-condensation of a carboxylic acid and an alcohol.

A)True

B)False

Q2) All ketones are capable of hydrogen bonding.

A)True

B)False

Q3) Ethane (C₂H₆)is much more reactive than disilane (Si₂H₆).

A)True

B)False

Q4) All the carbon atoms in a molecule of benzene lie in the same plane.

A)True

B)False

Q5) Helical and sheet-like segments in proteins arise from A)disulfide bridges. B)salt bridges.

C)crosslinking via covalent bonds.

D)dispersion forces within the protein's interior.

E)hydrogen bonding.

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) Briefly list the features/properties common to all catalysts and how they work.Draw a labeled reaction energy diagram as part of your answer.

Q2) Tetrafluoroethylene,C₂F₄,can be converted to octafluorocyclobutane which can be used as a refrigerant or an aerosol propellant.A plot of 1/[C₂F₄] vs.time gives a straight line with a slope of 0.0448 L mol^-1s^-1.What is the rate law for this reaction?

A)Rate = 0.0448 (L mol^-1s^-1)[C₂F₄]

B)Rate = 22.3 (mol L^-1s)[C₂F₄]

C)Rate = 0.0448 (L mol^-1s^-1)[C₂F₄]²

D)Rate = 22.3 (mol L^-1s)[C₂F₄]²

E)Rate = 0.0448 s^-1 [C₂F₄]

Q3) Which one of the following sets of units is appropriate for a second-order rate constant?

A)s^-1

B)mol L^-1 s^-1

C)L mol^-1 s^-1

D)mol² L^-2 s^-1

E)L² mol^-2 s^-1

Page 18

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) Which of the following has an effect on the magnitude of the equilibrium constant?

A)activation energy of the forward reaction

B)concentrations of the reactants and products

C)presence of a catalyst

D)change in volume of container

E)change in temperature

Q2) Changing the amount of reactant or product in an equilibrium reaction will always change the equilibrium position,regardless of the physical state of the substance involved.

A)True

B)False

Q3) Unless \(\Delta\)H°_rxn = 0,a change in temperature will affect the value of the equilibrium constant K_c.

A)True

B)False

Q4) When a reaction system reaches equilibrium,the forward and reverse reactions stop. A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) The chloride ion,Cl^-,is a typical Lewis acid.

A)True

B)False

Q2) (a)Write a balanced equation representing the reaction of the acid,H ₂PO₄^- with the base,water. (b)Write the expression for K_a of H₂PO₄^- in terms of concentrations of relevant species.

Q3) Which one of the following pairs is not a conjugate acid-base pair?

A)H₂O/OH^-

B)H₂O₂/HO₂^-

C)OH^-/O^2-

D)H₂PO₄^-/HPO₄^2-

E)HCl/H⁺

Q4) All strong acids have weak conjugate bases.

A)True

B)False

Q5) It is not possible to have a pH lying outside the range 0 to 14.

A)True

B)False

Page 20

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Propanoic acid (CH₃CH₂COOH)has a K_a of 1.34 \(\times\)10^-5.A 25.00 mL sample of 0.1000 mol L^-1 propanoic acid (in flask)is titrated with 0.1000 mol L^-1 NaOH solution,added from a buret.Carry out the calculations of the quantities indicated below.

a.The pH after 0.00 mL of NaOH are added.

b.The pH after 15.00 mL of NaOH are added.

c.The hydroxide ion concentration after 26.00 mL of NaOH are added.

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

A)1.4 \(\times\) 10^-4 M

B)8.2 \(\times\) 10^-5 M

C)5.4 \(\times\) 10^-5 M

D)3.2 \(\times\) 10^-6 M

E)2.5 \(\times\) 10^-12 M

Q3) Calculate the solubility of copper(II)carbonate,CuCO₃,in 1.00 mol L^-1 NH₃.K_sp = 3.0 \(\times\) 10^-12 for CuCO₃,K_f = 5.6 \(\times\) 10¹¹ for Cu(NH₃)₄²⁺

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Page 21

Chapter

Direction of Chemical Reactions

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Q1) Which relationship or statement best describes \(\Delta\)S° for the following reaction?

2NH₃(g)+ 2ClF₃(g) \(\to\) 6HF(g)+ N₂(g)+ Cl₂(g)

A)(\(\Delta\)S° \(\approx\) 0)

B)(\(\Delta\)S° < 0)

C)(\(\Delta\)S° > 0)

D)(\(\Delta\)S° = \(\Delta\)H°/T)

E)More information is needed to make a reasonable prediction.

Q2) Which relationship or statement best describes \(\Delta\)S° for the following reaction?

Pb(s)+ Cl₂(g) \(\to\) PbCl₂(s)

A)(\(\Delta\)S°\(\approx\) 00

B)9\(\Delta\)S° < 0)

C)(\(\Delta\)S° > 0)

D)(\(\Delta\)S° = \(\Delta\)H°/T)

E)More information is needed to make a reasonable prediction.

Q3) For any reaction,if \(\Delta\)G° > 0,then K < 1.

A)True

B)False

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) The value of the equilibrium constant for the reaction of nickel(II)ions with cadmium metal is 1.17 \(\times\) 10⁵.Calculate \(\Delta\)G° for the reaction at 25°C.

A)-12.6 kJ

B)-28.9 kJ

C)12.6 kJ

D)28.9 kJ

E)None of these choices is correct.

Q2) Which one of the following statements about electrochemical cells is correct?

A)In a salt bridge,current is carried by cations moving toward the anode,and anions toward the cathode.

B)In the external wire,electrons travel from cathode to anode.

C)The anode of a voltaic cell is labeled minus (-).

D)Oxidation occurs at the cathode,in an electrolytic cell.

E)None of these statements is correct.

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

A)True

B)False

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Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) In the spectrochemical series,which one of the following ligands has the strongest field?

A)H₂O

B)CN^-

C)NH₃

D)OH^-

E)Cl^-

Q2) What is the difference between a coordination compound and a complex ion?

Q3) A certain transition metal complex has the formula MX₄²⁺.If the metal ion has a d⁸ electron configuration,what is the shape of the complex?

A)octahedral

B)square pyramid

C)tetrahedral

D)trigonal pyramid

E)square planar

Q4) a.State the requirement for two molecules to be optical isomers.

b.A complex ion MABCD²⁺ (where A,B,C and D are different unidentate ligands)rotates the plane of polarized light.Deduce the geometry of the complex and draw the optical isomers of this ionic formula.

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) A patient's thyroid gland is to be exposed to an average of 5.5 µCi for 16 days as an ingested sample of iodine-131 decays.If the energy of the \(\beta\) radiation is 9.7 \(\times\) 10^-14 J and the mass of the thyroid is 32.0 g,what is the dose received by the patient?

A)0.027 rads

B)1.2 rads

C)37 rads

D)85 rads

E)None of these choices is correct.

Q2) The radiochemist,Will I.Glow,studied thorium-232 and found that 2.82 \(\times\) 10^-7 moles emitted 8.42 \(\times\) 10⁶ \(\alpha\) particles in one year.What is the decay constant for thorium-232?

A)3.35 \(\times\) 10^-14 yr^-1

B)4.96 \(\times\) 10^-11 yr^-1

C)1.40 \(\times\) 10¹⁰ yr^-1

D)2.99 \(\times\) 10¹³ yr^-1

E)None of these choices is correct.

Q3) Write a complete,balanced equation to represent the formation of manganese-55 by the beta decay of another nuclide.

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Page 25