Principles of Chemistry Exam Solutions - 3071 Verified Questions

Principles of Chemistry

Exam Solutions

Course Introduction

Principles of Chemistry is an introductory course that explores the fundamental concepts and theories underlying the science of chemistry. Topics include atomic and molecular structure, chemical bonding, stoichiometry, states of matter, thermochemistry, chemical equilibrium, and the behavior of acids and bases. Through a combination of lectures, laboratory experiments, and problem-solving exercises, students gain an understanding of how chemical principles apply to real-world phenomena and develop analytical skills essential for further studies in science and engineering.

Recommended Textbook

Chemistry Structure and Properties 2nd Edition by Nivaldo J. Tro

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Page 2

Chapter 1: Essentials: Units, Measurements, and Problem Solving

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Sample Questions

Q1) What is the volume (in cm³)of a 9.37 g piece of metal with a density of 4.66 g/cm³?

A) 2.01

B) 19.5

C) .425

D) 6.65

E) none of the above

Answer: A

Q2) How many mm are in 3.20 cm?

A) 3.20 × 10¹ mm

B) 3.20 × 10^-1 mm

C) 3.20 × 10^-2 mm

D) 3.20 × 10² mm

E) 3.20 × 10³ mm

Answer: A

Q3) Describe the difference between an intensive and extensive property using examples.

Answer: An intensive property does NOT depend on the amount of the substance present,such as color or density.An extensive property is one that does depend on the amount of the substance,such as mass or volume.

Page 3

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Chapter 2: Atoms

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Sample Questions

Q1) Iodine is an example of

A) a compound.

B) an element.

C) a heterogeneous mixture.

D) a homogeneous mixture.

Answer: B

Q2) an atom that has lost an electron is A) a cation.

B) unlikely to be found in homogeneous mixtures.

C) electrically neutral.

D) likely to behave exactly like the parent atom.

E) an anion.

Answer: A

Q3) How many protons are in nickel?

A) 28

B) 30

C) 31

D) 30.7

E) 58.7

Answer: A

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Chapter 3: The Quantum Mechanical Model of the Atom

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Q1) How many sublevels are contained in the second shell (n = 2)of a given atom?

A) 1

B) 2

C) 9

D) 4

E) 3

Answer: B

Q2) What is the wavelength of an electron (m = 9.11 × 10^-28 g)moving at 1/5 the speed of light?

A) 9.36 × 10^-9 m

B) 1.21 × 10^-11 m

C) 6.73 × 10^-8 m

D) 2.42 × 10^-12 m

E) 4.85 × 10^-10 m

Answer: B

Q3) Why don't we observe the wavelength of everyday macroscopic objects?

Answer: Due to the large mass of macroscopic objects,the de Broglie wavelength is extremely small.The wavelength is so small that it is impossible to detect compared to the size of the object.

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Page 5

Chapter 4: Periodic Properties of the Elements

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Sample Questions

Q1) Choose the ground state electron configuration for Cr³ .

A) [Ar]4s¹3d²

B) [Ar]

C) [Ar]4s²3d⁶

D) [Ar]3d³

E) [Ar]4s²3d¹

Q2) Place the following in order of increasing metallic character. Rb Cs K Na

A) K < Cs < Na < Rb

B) Na < K < Rb < Cs

C) Cs < Rb < K < Na

D) K < Cs < Rb < Na

E) Na < Rb < Cs < K

Q3) Which ion has the largest radius?

A) Na⁺

B) Ga³⁺

C) K⁺

D) Mg²⁺

E) Ca²⁺

Q4) Why does the size of the transition elements stay roughly the same across a period?

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Chapter 5: Molecules and Compounds

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Sample Questions

Q1) What is the charge on the Fe ions in Fe₂O₃?

A) 2-

B) 1+

C) 2+

D) 3+

Q2) Identify the compound with ionic bonding.

A) Na F

B) K

C) H₂O

D) He

E) S

Q3) Combustion analysis of an unknown compound containing only carbon and hydrogen produced 2.277 g of CO₂ and 1.161 g of H₂O.What is the empirical formula of the compound?

A) CH₂

B) C₂H₅

C) C₄H₁₀

D) C₅H₂

Q4) Describe the difference between a molecular formula and an empirical formula.Give an example.

Page 7

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Chapter 6: Chemical Bonding I

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Q1) Determine the electron geometry,molecular geometry and polarity of SF₆.

A) eg = trigonal bipyramidal, mg = trigonal bipyramidal, nonpolar

B) eg = tetrahedral, mg = tetrahedral, polar

C) eg = trigonal bipyramidal, mg = see-saw, polar

D) eg = octahedral, mg = trigonal bipyramidal, nonpolar

E) eg = octahedral, mg = octahedral, nonpolar

Q2) Determine the electron geometry (eg)and molecular geometry (mg)of XeF₂.

A) eg = trigonal bipyramidal, mg = bent

B) eg = linear, mg = linear

C) eg = tetrahedral, mg = linear

D) eg = trigonal bipyramidal, mg = linear

E) eg = tetrahedral, mg = bent

Q3) Give the approximate bond angle for a molecule with an octahedral shape.

A) 109.5°

B) 180°

C) 120°

D) 105°

E) 90°

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Chapter 7: Chemical Bonding Ii

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Sample Questions

Q1) A molecule containing a central atom with sp hybridization has a(n)________ electron geometry.

A) linear

B) trigonal pyramidal

C) square planar

D) octahedral

E) bent

Q2) Give the electron geometry (eg),molecular geometry (mg),and hybridization for H₂O.

A) eg = tetrahedral, mg = bent, sp³

B) eg = trigonal pyramidal, mg = trigonal pyramidal, sp³

C) eg = tetrahedral, mg = trigonal pyramidal, sp³

D) eg = bent, mg = bent, sp²

E) eg = trigonal planar, mg = trigonal planar, sp²

Q3) Give the hybridization for the O in OF₂.

A) sp

B) sp³

C) sp²

D) sp³d

E) sp³d²

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Chapter 8: Chemical Reactions and Chemical Quantities

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Sample Questions

Q1) How many moles of nitrogen are formed when 58.6 g of KNO₃ decomposes according to the following reaction? The molar mass of KNO₃ is 101.11 g/mol.

4 KNO₃(s) 2 K₂O(s)+ 2 N₂(g)+ 5 O₂(g)

A) 0.290 mol N₂

B) 0.580 mol N₂

C) 18.5 mol N₂

D) 0.724 mol N₂

E) 1.73 mol N₂

Q2) Strontium phosphate reacts with sulfuric acid to form strontium sulfate and phosphoric acid.What is the coefficient for sulfuric acid when the equation is balanced using the lowest,whole-numbered coefficients?

A) 1

B) 2

C) 3

D) none of these

Q3) Describe the greenhouse effect.

Q4) Define a limiting reagent.

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Page 10

Chapter 9: Introduction to Solutions and Aqueous Reactions

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Sample Questions

Q1) 94.20 mL of 0.800 M potassium hydroxide reacts with 125.0 mL of a solution containing oxalic acid,which is a diprotic species.What is the concentration of the acid solution?

A) 0. 542 M

B) 0. 603 M

C) 0. 301 M

D) 1.38 M

Q2) How many grams of NaCl are required to make 300.0 mL of a 2.500 M solution?

A) 58.40 g

B) 175.3 g

C) 14.60 g

D) 43.83 g

Q3) 6.74 g of the monoprotic acid KHP (MW = 204.2 g/mol)is dissolved into water.The sample is titrated with a 0.703 M solution of calcium hydroxide to the equivalence point.What volume of base was used?

A) 8.64 mL

B) 11.8 mL

C) 23.5 mL

D) 47.0 mL

E) 93.9 mL

Page 11

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Chapter 10: Thermochemistry

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Sample Questions

Q1) Which one of the following elements is NOT in its standard state?

A) Br₂(g)

B) F₂(g)

C) Hg(l)

D) Na(s)

E) Cl₂(g)

Q2) For a particular process that is carried out at constant pressure,q = 145 kJ and w = -35 kJ.Therefore

A) E = 110 kJ and H = 145 kJ.

B) E = 145 kJ and H = 110 kJ.

C) E = 145 kJ and H = 180 kJ.

D) E = 180 kJ and H = 145 kJ.

Q3) Give the units of specific heat capacity.

A) J/°C

B) (^J/g °C)

C) Jmole °C

D) g/°C

E) gmole °C

Q4) Give the equation with the elements in MgSO₄ in their standard state as the reactants and MgSO₄ as the product.

Page 12

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Chapter 11: Gases

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Sample Questions

Q1) A gas mixture consists of N₂,O₂,and Ne,where the mole fraction of N₂ is 0.55 and the mole fraction of Ne is 0.25.If the mixture is at STP in a 5.0 L container,how many molecules of O₂ are present?

A) 4.5 × 10²²molecules O₂

B) 2.7 × 10²² molecules O₂

C) 3.7 × 10²³ molecules O₂

D) 1.1 × 10²³ molecules O₂

E) 9.3 × 10²⁴ molecules O₂

Q2) The total pressure of a gas mixture is the sum of the partial pressure of its components is known as

A) Avogadro's Law.

B) Ideal Gas Law.

C) Charles's Law.

D) Boyle's Law.

E) Dalton's Law.

Q3) Consider a container of gas under a particular P,V,T set of conditions.Describe how the pressure would change if the volume were doubled while the absolute temperature was increasing by a factor of two.

Q4) Define effusion.

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Chapter 12: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Propanol has a normal boiling point of 97.8°C.At 400 torr,it has a boiling point of 82.0°C.What is the heat of vaporization?

A) 24.9 kJ/mol

B) 38.7 kJ/mol

C) 12.3 kJ/mol

D) 52.7 kJ/mol

E) 44.4 kJ/mol

Q2) Define fusion.

A) the phase transition from solid to liquid

B) the phase transition from gas to solid

C) the phase transition from gas to liquid

D) the phase transition from liquid to gas

E) the phase transition from liquid to solid

Q3) Choose the compound that exhibits hydrogen bonding as its strongest intermolecular force.

A) NBr₃

B) C₈H₁₈

C) CH₃CH₂SeH

D) CH₂Br₂

E) None of the above compounds exhibit hydrogen bonding.

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Chapter 13: Crystalline Solids and Modern Materials

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Sample Questions

Q1) Identify the type of solid for argon.

A) metallic atomic solid

B) ionic solid

C) nonbonding atomic solid

D) molecular solid

E) networking atomic solid

Q2) A polymer that eliminates an atom or small group of atoms is known as a(n)

A) addition polymer

B) condensation polymer

C) dimer

D) copolymer

E) substituted polymer

Q3) What is the edge length of a face-centered cubic unit cell made up of atoms having a radius of 128 pm?

A) 181 pm

B) 362 pm

C) 512 pm

D) 1020 pm

Q4) Give the edge length in terms of r for a simple cubic cell.

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Chapter 14: Solutions

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Sample Questions

Q1) Calculate the mole fraction of Ba I₂ in an aqueous solution prepared by dissolving 0.400 moles of Ba I₂ in 850.0 g of water.

A) 0.00841

B) 0.0270

C) 0.00900

D) 0.0248

E) 0.0167

Q2) Determine the vapor pressure of a solution at 25°C that contains 85.3 g of naphthalene (C₁₀H₈)in 540 g of benzene (C₆H₆).The vapor pressure of benzene at 25°C is 74.6 torr.

A) 74.6 toff.

B) 68.0 torr.

C) 54.9 torr.

D) 7.52 torr.

E) 22.4 toff.

Q3) Define osmosis.

Q4) Explain why the van't Hoff factor for MgCl₂ is less than it's predicted value.

Q5) Explain why water does not dissolve in gasoline.

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Chapter 15: Chemical Kinetics

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Sample Questions

Q1) Explain what the exponential factor in the Arrhenius equation represents.

Q2) What is the overall order of the following reaction,given the rate law?

2A + 3 B 2C + D Rate = k[A] [B] ⁰

A) 3rd order

B) 2nd order

C) 5th order

D) 1st order

E) 0th order

Q3) The decomposition of dinitrogen pentoxide is described by the chemical equation 2 N₂O₅(g) 4 NO₂(g)+ O₂(g)

If the rate of disappearance of N₂O₅ is equal to 1.60 mol/min at a particular moment,what is the rate of appearance of NO₂ at that moment?

A) 0.800 mol/min

B) 1.60 mol/min

C) 3.20 mol/min

D) 6.40 mol/min

Q4) Define activation energy.

Q5) What function do enzymes serve?

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Chapter 16: Chemical Equilibrium

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Q1) Which of the following statements is TRUE?

A) If Q < K, it means the reverse reaction will proceed to form more reactants.

B) If Q > K, it means the forward reaction will proceed to form more products.

C) If Q = K, it means the reaction is at equilibrium.

D) All of the above are true.

E) None of the above are true.

Q2) Determine the value of K_p for the following reaction if the equilibrium partial pressures are as follows: P(CO₂)_eq =1.8 atm,P(CO)_eq = 0.35 atm,P(O₂)_eq = 0.50 atm. 2 CO₂(g) 2 CO(g)+ O₂(g)

A) 2.4

B) 0.83

C) 1.3

D) 0.019

E) 0.22

Q3) How will equilibrium be affected,with the same number of moles of gas on both sides,if the pressure is increased?

Q4) Can the K_p and K_c for a reaction ever have the same value?

Why or why not?

Q5) Why aren't solids or liquids included in an equilibrium expression?

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Chapter 17: Acids and Bases

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Sample Questions

Q1) Which of the following acids will have the strongest conjugate base?

A) HCl

B) HClO₄

C) HNO₃

D) HCN

E) HI

Q2) Describe a molecule that can be a Lewis acid.

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) Determine the pOH in a 0.235 M NaOH solution.

A) 13.76

B) 0.24

C) 13.37

D) 0.63

E) 12

Q5) What is the difference between a strong and weak acid?

Page 19

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Chapter 18: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Determine the molar solubility for Ag₂C₂O₄ in pure water.The K_sp for Ag₂C₂O₄ is 3.5 × 10^-11.

A) 1.8 × 10^-11 M

B) 1.4 × 10^-6 M

C) 5.9 × 10^-6 M

D) 2.1 × 10^-4 M

E) 3.5 × 10^-6 M

Q2) What is the pH of the resulting solution if 45.00 mL of 0.10 M acetic acid is added to 10.00 mL of 0.10 M NaOH? Assume that the volumes of the solutions are additive.K_a = 1.8 × 10^-5for CH₃CO₂H.

A) 9.80

B) 8.71

C) 5.29

D) 4.20

Q3) Sketch the titration curve for a monoprotic weak acid titrated with a strong base.Make sure to indicate the equivalence point (and whether it is acidic,basic or neutral)and the buffer region.

Q4) Give the name of the compound that is in antifreeze and is toxic to pets.

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Chapter 19: Free Energy and Thermodynamics

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Q1) At what temperature does a perfect crystalline solid have S=0?

A) 0 K

B) 0°C

C) 100 °C

D) 273 K

E) It is not possible for a substance to have S=0.

Q2) What is TRUE if ln K is 1?

A) G°_rxn is positive and the reaction is spontaneous in the forward direction.

B) G°_rxn is negative and the reaction is spontaneous in the forward direction.

C) G°_rxn is negative and the reaction is spontaneous in the reverse direction.

D) G°_rxn is positive and the reaction is spontaneous in the reverse direction. E) G°_rxn is zero and the reaction is at equilibrium.

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.

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Chapter 20: Electrochemistry

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Q1) Why is sugar water not a good conductor of current?

Q2) Which of the following is the weakest oxidizing agent?

A) H₂O₂(aq)

B) Fe³⁺(aq)

C) ClO₂(g)

D) F^-(s)

E) Fe(s)

Q3) Balance the following redox reaction if it occurs in acidic solution.What are the coefficients in front of H and Fe³⁺ in the balanced reaction?

Fe²⁺(aq)+

MnO₄ (aq)

Fe³⁺(aq)+ Mn²⁺(aq)

A) H = 2, Fe³⁺= 3

B) H = 8, Fe³⁺= 5

C) H = 3, Fe³⁺= 2

D) H = 5, Fe³⁺= 1

E) H = 8, Fe³⁺= 1

Q4) Why are iron nails coated with zinc?

Q5) Explain the significance of the standard hydrogen electrode (SHE)in the tabulation of standard reduction potentials of other species.

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Chapter 21: Radioactivity and Nuclear Chemistry

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Q1) Define chain reaction in terms of the fission of uranium nucleus.

Q2) A geological sample is found to have a Pb-206/U-238 mass ratio of 0.337/1.00.Assuming there was no Pb-206 present when the sample was formed,how old is it? The half-life of U-238 is 4.5 × 10⁹ years.

A) 7.3 × 10¹¹ years

B) 1.4 × 10¹⁰ years

C) 2.4 × 10¹⁰ years

D) 2.1 × 10⁹ years

E) 7.1 × 10⁹ years

Q3) Identify the lowest natural radiation.

A) cosmic radiation from outer space

B) terrestrial radiation

C) natural radionuclides in the body

D) a five-hour jet airplane ride

E) radon gas

Q4) Beta decay of ⁶⁰Co produces a beta particle and A) (⁵⁶Mn).

B) (⁵⁹Co).

C) (⁶⁰Fe).

D) (⁶⁰Ni).

Page 23

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Chapter 22: Organic Chemistry

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Q1) Which of the following compounds exhibit geometric isomerism?

A) CH₂=CH-CH₃

B) CCl₂=CBr₂

C) CH₃-CH=CH-CH₃

D) CCl₂=CHBr

E) All of the above exhibit geometric isomerism.

Q2) Why are there so many more carbon compounds than the number of compounds made up of all the rest of the elements combined?

Q3) Which of the following compounds is an alcohol?

A) CH₃CH₂CH₂CO₂H

B) CH₃-O-CH₃

C) CH₃CO₂CH₃

D) CH₃CH₂CH=O

E) CH₂OH-CH₂OH

Q4) Identify the formula for an alkyne.

A) C_nH₂_n₊₂

B) C_nH₂_n-₂

C) C_nH₂_n

D) C_nH₂_n-4

E) C_nH_2n+4

Page 24

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Chapter 23: Transition Metals and Coordination Compounds

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Q1) Identify the ion that is responsible for the red color of rubies.

A) Cr³⁺

B) Cr⁴⁺

C) Cr⁵⁺

D) Cr⁶⁺

E) Cr⁷⁺

Q2) The complex ion,[Cr(NH₃)₆]³⁺,has a maximum absorption of 463 nm.Calculate the crystal field splitting energy (in kJ/mol)for this ion.

A) 299 kJ/mol

B) 538 kJ/mol

C) 372 kJ/mol

D) 258 kJ/mol

E) 193kJ/mol

Q3) What is the Lanthanide contraction?

Q4) What is the ground-state electron configuration for the element nickel (Z = 28)?

A) [Ne] 4s¹ 3d⁹

B) [Ar] 4s¹ 3d⁹

C) [Ar] 4s² 3d⁸

D) [Ar] 3d¹⁰

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