Chemical Structure and Properties Exam Review - 3071 Verified Questions

Chemical Structure and Properties

Exam Review

Course Introduction

Chemical Structure and Properties explores the fundamental concepts underlying the structure of atoms, molecules, and solids, and how these structures determine the physical and chemical properties of materials. Topics include atomic theory, bonding models (ionic, covalent, metallic, and molecular), molecular geometry, intermolecular forces, and the relationship between structural features and chemical reactivity. The course emphasizes both qualitative and quantitative approaches to understanding how molecular architecture affects phenomena such as boiling and melting points, solubility, conductivity, and reactivity, providing a foundation for further studies in chemistry and related disciplines.

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

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23 Chapters

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Chapter 1: Essentials: Units, Measurements, and Problem

Solving

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Sample Questions

Q1) Since metals expand when heated,what happens to the density of a sample of iron metal as it is heated from room temperature to 100°C?

(This is below the melting point of iron.)

Answer: Since the mass of the iron stays constant,but the volume increases as the temperature is raised,the density of the iron decreases upon heating.

Q2) An alligator is 213.4 cm long.How long is he in feet?

A) 7.00 feet

B) 84.0 ft

C) 17.8 ft

D) 45.2 ft

E) 1009 ft

Answer: A

Q3) Identify the unit of measurement which is a SI base unit of measurement.

A) second

B) Celsius

C) cup

D) ounce

E) yard

Answer: A

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Chapter 2: Atoms

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Q1) An ion has 12 protons,^{14 neutrons,and 10 electrons.The symbol for the ion is}

A) 26_Mg2+

B) 26_Mg2-

C) 24_Si2+

D) 24_Si2-

E) 26_Ne2-

Answer: A

Q2) How many atoms of carbon are in 2.50 moles of CO₂?

A) 4.52 × 10²⁴ atoms

B) 1.51 × 10²⁴ atoms

C) 5.02 × 10²³ atoms

D) 3.01 × 10²⁴ atoms

E) 7.53 × 10²³ atoms

Answer: B

Q3) The atomic number is equal to the number of ________. Answer: protons

Q4) The number 6.022 × 10²³ is called ________. Answer: Avogadro's number

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Chapter 3: The Quantum Mechanical Model of the Atom

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Sample Questions

Q1) Consider a 3p orbital.How is it different from a 2p orbital?

Answer: It is larger in size and contain additional nodes.

Q2) Which of the following transitions (in a hydrogen atom)represent emission of the longest wavelength photon?

A) n = 1 to n = 3

B) n = 3 to n = 1

C) n = 3 to n = 5

D) n = 4 to n = 2

E) n = 5 to n = 4

Answer: E

Q3) The distance between adjacent crests is called A) wavelength.

B) amplitude.

C) frequency.

D) area.

E) median.

Answer: A

Q4) Give an example of a d orbital.

Answer: d_yz,d_xy,d_xz,d_x2_-y2,or d_z2

Page 5

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Chapter 4: Periodic Properties of the Elements

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Sample Questions

Q1) Which of the following statements is FALSE?

A) Halogens are very reactive elements.

B) The alkali metals are fairly unreactive.

C) Sulfur is a main group element.

D) Noble gases do not usually form ions.

E) Zn is a transition metal.

Q2) Choose the ground state electron configuration for Zn² .

A) [Ar]4s²3d⁸

B) [Ar]3d¹⁰

C) [Ar]4s²3d⁶

D) [Ar]

E) [Ar]3d⁸

Q3) Place the following in order of decreasing metallic character. P As K

A) P > As > K

B) As > P > K

C) K > P > As

D) As > K > P

E) K > As > P

Q4) Why is the first ionization energy of sulfur smaller than the first ionization energy of phosphorus?

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Chapter 5: Molecules and Compounds

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Sample Questions

Q1) A sample of pure calcium fluoride with a mass of 15.0 g contains 7.70 g of calcium.How much calcium is contained in 30.0 g of calcium fluoride?

A) 1.71 g

B) 7.70 g

C) 15.0 g

D) 15.4 g

Q2) Calculate the mass percent composition of oxygen in Al₂(SO₄)₃.

Q3) Determine the name for H₂CO₃.

A) carbonous acid

B) dihydrogen carbonate

C) carbonic acid

D) hydrocarbonic acid

E) hydrocarbide acid

Q4) The chemical formula for barium nitride is

A) Ba(NO₃)₂.

B) Ba(NO₂)₂.

C) Ba₃N₂.

D) BaN₂.

Q5) Define empirical formula.

Page 7

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Chapter 6: Chemical Bonding I

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Sample Questions

Q1) Identify the number of electron groups around a molecule with an octahedral shape.

A) 6

B) 2

C) 3

D) 4

E) 5

Q2) Give the approximate bond angle for a molecule with an octahedral shape.

A) 109.5°

B) 180°

C) 120°

D) 105°

E) 90°

Q3) How many of the following elements can form compounds with an expanded octet?

S Kr Xe B

A) 0

B) 1

C) 2

D) 3

E) 4

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Chapter 7: Chemical Bonding Ii

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Sample Questions

Q1) Identify the number of electron groups around a molecule with sp² hybridization.

A) 1

B) 2

C) 3

D) 4

E) 5

Q2) Which of the following statements is TRUE?

A) The total number of molecular orbitals formed doesn't always equal the number of atomic orbitals in the set.

B) A bond order of 0 represents a stable chemical bond.

C) When two atomic orbitals come together to form two molecular orbitals, one molecular orbital will be lower in energy than the two separate atomic orbitals and one molecular orbital will be higher in energy than the separate atomic orbitals.

D) Electrons placed in antibonding orbitals stabilize the ion/molecule.

E) All of the above are true.

Q3) Give the electron geometry,molecular geometry,and hybridization for both carbons in CH₃COOH.

Q4) According to molecular orbital theory,what is an antibonding orbital?

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Chapter 8: Chemical Reactions and Chemical Quantities

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Sample Questions

Q1) Determine the theoretical yield of HCl if 60.0 g of BCl₃and 37.5 g of H₂O are reacted according to the following balanced reaction.A possibly useful molar mass is BCl₃ = 117.16 g/mol. BCl₃(g)+ 3 H₂O(l) H₃BO₃(s)+ 3 HCl(g)

A) 75.9 g HCl

B) 132 g HCl

C) 187 g HCl

D) 56.0 g HCl

E) 25.3 g HCl

Q2) How many grams of oxygen are formed when 5.30 moles of KOH are formed? 4 KO(s)+ 2 H₂O(l) 4 KOH(s)+ O₂(g)

A) 5.30 g O₂

B) 21.2 g O₂

C) 42.4 g O₂

D) 170 g O₂

E) 84.8 g O₂

Q3) Balance the following equation. ________ C₁₀H₁₂ + ________ O₂ ________ H₂O + ________ CO₂

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Chapter 9: Introduction to Solutions and Aqueous Reactions

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Sample Questions

Q1) Determine the molarity of a solution formed by dissolving 97.7 g LiBr in enough water to yield 750.0 mL of solution.

A) 1.50 M

B) 1.18 M

C) 0.130 M

D) 0.768 M

E) 2.30 M

Q2) What is the concentration of magnesium ions in a 0.125 M Mg SO₄ solution?

A) 0.125 M

B) 0.0625 M

C) 0.375 M

D) 0.250 M

E) 0.160 M

Q3) Determine the oxidation state of Cl in ClO₃^-.

A) +1

B) +3

C) +5

D) +7

E) 0

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Chapter 10: Thermochemistry

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Sample Questions

Q1) Using the following thermochemical equation,determine the amount of heat produced from the combustion of 24.3 g benzene (C₆H₆).The molar mass of benzene is 78.11 g/mole. 2 C₆H₆(l)+ 15 O₂(g) 12 CO₂(g)+ 6 H₂O(g) H°_rxn= -6278 kJ

A) 3910 kJ

B) 1950 kJ

C) 977 kJ

D) 40.1 kJ

E) 0.302 kJ

Q2) Describe the energy changes that occur when a book is held 6 ft off the floor and then dropped.

Q3) Which of the following processes is endothermic?

A) mixing water and acid

B) rusting iron

C) photosynthesis

D) the electron affinity of a fluorine atom

E) None of the above processes are endothermic.

Q4) Why are the standard heats of formation for elements in their most stable form assigned a value of "0"?

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Chapter 11: Gases

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Sample Questions

Q1) Define hypoxia.

A) oxygen starvation

B) increased oxygen concentration in body tissues

C) increased nitrogen concentration in body tissues and fluids

D) nitrogen starvation

Q2) Determine the volume of O₂ (at STP)formed when 50.0 g of KClO₃ decomposes according to the following reaction.The molar mass for KClO₃is 122.55 g/mol. 2 KClO₃(s) 2 KCl(s)+ 3 O₂(g)

A) 9.14 L

B) 8.22 L

C) 12.3 L

D) 13.7 L

E) 14.6 L

Q3) What volume will 0.780 moles of Xe occupy at STP?

A) 22.4 L

B) 70.0 L

C) 43.7 atm

D) 17.5 L

E) 15.6 L

Q4) Why does hot air rise?

Page 13

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Chapter 12: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) The heat required to melt 1 mol of a solid is known as the A) melting point.

B) heat of fusion.

C) freezing point.

D) heat of vaporization.

E) critical point.

Q2) Which substance below has the strongest intermolecular forces?

A) A₂X, H_vap= 39.6 kJ/mol

B) BY₂, H_vap= 26.7 kJ/mol

C) C₃X₂, H_vap= 36.4 kJ/mol

D) DX₂, H_vap= 23.3 kJ/mol

E) EY₃, H_vap= 21.5 kJ/mol

Q3) The freezing point of water is

A) 32^oF

B) 32^oC

C) 212^oC

D) 100°C

E) 273°C

Q4) Define volatile.

Q5) Define boiling point of a liquid.

Page 14

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Chapter 13: Crystalline Solids and Modern Materials

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Sample Questions

Q1) Which of the following substances should have the highest melting point?

A) Fe

B) Ne

C) Xe

D) N₂

E) CO

Q2) A metal crystallizes in a face centered cubic structure and has a density of 11.9 g/cm³.If the radius of the metal atom is 138 pm,what is the identity of the metal?

A) At

B) Pd

C) Mn

D) Fe

E) Cr

Q3) 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

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Chapter 14: Solutions

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Sample Questions

Q1) Dissolving can be defined as

A) rate of dissolution = rate of deposition.

B) rate of dissolution < rate of deposition.

C) rate of dissolution > rate of deposition.

D) rate of bubbling > rate of dissolving

E) rate of condensing > rate of bubbling

Q2) What is the mole fraction of I₂ in a solution made by dissolving 55.6 g of I₂ in 245 g of hexane,C₆H₁₄?

A) 0.0715

B) 0.0770

C) 0.133

D) 0.154

Q3) Which of the following ions should have the most exothermic H_hydration?

A) Na

B) Ba²

C) Al³

D) Ca²

E) Sr²

Q4) Define osmosis.

Page 16

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Chapter 15: Chemical Kinetics

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Sample Questions

Q1) Biological catalysts that increase the rates of biochemical reactions are known as A) substrates.

B) inhibitors.

C) enzymes.

D) binders.

E) trumanettes.

Q2) Given the following balanced equation,determine the rate of reaction with respect to [Cl₂].If the rate of Cl₂ loss is 4.24 × 10^-2 M/s,what is the rate of formation of NOCl?

2 NO(g)+ Cl₂(g) 2 NOCl(g)

A) 4.24 × 10^-2 M/s

B) 2.12 × 10^-2 M/s

C) 1.06 × 10^-1 M/s

D) 8.48 × 10^-2 M/s

E) 1.61 × 10^-2 M/s

Q3) Define activation energy.

Q4) What happens to the concentration of reactants and products during a chemical reaction?

Q5) What function do enzymes serve?

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Chapter 16: Chemical Equilibrium

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Sample Questions

Q1) The equilibrium constant is equal to 5.00 at 1300 K for the reaction: 2 SO₂(g)+ O₂(g) 2 SO₃(g).

If initial concentrations are [SO₂] = 1.20 M,[O₂] = 0.45 M,and [SO₃] = 1.80 M,the system is

A) at equilibrium.

B) not at equilibrium and will remain in an unequilibrated state.

C) not at equilibrium and will shift to the left to achieve an equilibrium state.

D) not at equilibrium and will shift to the right to achieve an equilibrium state.

Q2) What is n for the following equation in relating K_c to K_p?

CH₄(g)+ H₂O(g) CO(g)+ 3 H₂(g)

A) -5

B) -1

C) -2

D) 2

E) 1

Q3) Why aren't solids or liquids included in an equilibrium expression?

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Chapter 17: Acids and Bases

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Q1) Calculate the hydronium ion concentration in an aqueous solution with a pH of 4.33 at 25°C.

A) 2.1 × 10^-10 M

B) 9.7 × 10^-10M

C) 4.7 × 10^-5 M

D) 3.8 × 10^-5 M

E) 6.3 × 10^-6 M

Q2) Determine the ammonia concentration of an aqueous solution that has a pH of 11.00.The equation for the dissociation of NH₃ (K_b= 1.8 × 10^-5)is

NH₃(aq)+ H₂O(l) NH₄⁺(aq)+ OH^-(aq).

A) 3.0 M

B) 0.056 M

C) 1.8 × 10^-² M

D) 1.0 × 10^-3 M

Q3) Which Brønsted-Lowry acid is not considered to be a strong acid in water?

A) HI

B) HBr

C) H₂S O₃

D) HNO₃

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Chapter 18: Aqueous Ionic Equilibrium

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Q1) What is the pH at the equivalence point of a weak base-strong acid titration if 20.00 mL of NaOCl requires 28.30 mL of 0.20 M HCl? K_a = 3.0 × 10^-8 for HOCl.

A) 0.70

B) 3.39

C) 3.76

D) 4.23

Q2) Determine the molar solubility for Cd(OH)₂ in pure water.K_sp for Cd(OH)₂ is 2.0 × 10^-14.

A) 2.0 × 10^-14 M

B) 1.7 × 10^-5 M

C) 6.6 × 10^-8 M

D) 3.5 × 10^-6 M

E) 2.9 × 10^-6 M

Q3) Which of the following solutions is a good buffer system?

A) a solution that is 0.10 M NaCl and 0.10 M HCl

B) a solution that is 0.10 M HCN and 0.10 M LiCN

C) a solution that is 0.10 M NaOH and 0.10 M HNO₃

D) a solution that is 0.10 M HNO₃ and 0.10 M NaNO₃

E) a solution that is 0.10 M HCN and 0.10 M K Br

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Chapter 19: Free Energy and Thermodynamics

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Sample Questions

Q1) Calculate S°_rxn for the following reaction.The S° for each species is shown below the reaction. C₂H₂(g)+ H₂(g) C₂H₄(g)

S°(J/molK)200.9 130.7 219.3

A) +112.3 J/K

B) +550.9 J/K

C) -112.3 J/K

D) +337.1 J/K

E) -550.9 J/K

Q2) Why can't we say that a spontaneous reaction is a fast reaction?

Q3) Consider a reaction that has a negative H and a positive S.Which of the following statements is TRUE?

A) This reaction will be spontaneous only at high temperatures.

B) This reaction will be spontaneous at all temperatures.

C) This reaction will be nonspontaneous at all temperatures.

D) This reaction will be nonspontaneous only at high temperatures.

E) It is not possible to determine without more information.

Q4) Define the third law of thermodynamics.

Q5) What is "free" energy?

Give a fictitious example.

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Chapter 20: Electrochemistry

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Q1) What is electrolysis?

Q2) What is the shorthand notation that represents the following galvanic cell reaction? Sn²⁺(aq)+ F₂(g) Sn³⁺(aq)+ 2 F^-(aq)

A) Sn²⁺(aq) Sn³⁺(aq) F₂(g) F^-(aq)

B) Sn(s) Sn²⁺(aq) Sn⁴⁺(aq) F₂(g) F^-(aq) C(s)

C) Pt(s) Sn⁴⁺(aq), Sn²⁺(aq), F₂(g) F^-(aq) C(s)

D) Pt(s) Sn²⁺(aq), Sn³⁺(aq) F₂(g) F^-(aq) C(s)

Q3) Determine which of the following pairs of reactants will result in a spontaneous reaction at 25°C.

A) I^-(aq) + Fe²⁺(aq)

B) Ca(s) + Mg²⁺(aq)

C) H₂(g) + Ni²⁺(aq)

D) Ag(s) + Sn²⁺(aq)

E) All of the above pairs will react.

Q4) What is the difference between a voltaic cell and an electrolytic cell?

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Chapter 21: Radioactivity and Nuclear Chemistry

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Q1) Define mass defect.

A) the difference in mass between an atom and the sum of its separate components

B) an atom with too many neutrons

C) the difference in mass between a radioactive atom and a nonradioactive atom

D) energy released in a radioactive reaction

E) energy absorbed in a radioactive reaction

Q2) What is the "mass defect"?

Q3) Calculate the mass defect in Mo-96 if the mass of a Mo-96 nucleus is 95.962 amu.The mass of a proton is 1.00728 amu and the mass of a neutron is 1.008665 amu.

A) 0.197 amu

B) 0.795 amu

C) 0.212 amu

D) 0.812 amu

E) 0.188 amu

Q4) The combination of two light nuclei to form a heavier nuclei is called

A) radioactive cleavage.

B) nuclear fission.

C) nuclear fusion.

D) radioactive merge.

E) half life.

Page 23

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Chapter 22: Organic Chemistry

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Q1) Write a balanced chemical equation to represent the reaction of 3-ethyl-2-methylhexane with Br₂.

A) C₇H₁₄+Br₂ C₇H₁₂Br + HBr

B) C₉H₂₀+Br₂

C₉H₁₉Br + HBr

C) C₆H₁₂+Br₂

C₆H₁₀Br₂ + H₂

D) C₉H₁₈+Br₂ C₉H₁₆Br₂ + H₂

E) C₆H₁₂+Br₂

C₆H₁₁Br + HBr

Q2) How many isomers are there for C₈H₁₈ ?

A) 18

B) 19

C) 20

D) 21

Q3) Which part of HCN,the hydrogen or the cyano group,adds to the O in a C=O bond and why?

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Chapter 23: Transition Metals and Coordination Compounds

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Q1) Why is +2 a common oxidation state for transition elements?

Q2) Identify the transition metal that is used in oxygen transport in the human body. A) chromium

B) iron

C) potassium

D) copper

E) silver

Q3) How many unpaired electrons would you expect for the complex ion: [Co(NO₂)₆]^3-?

A) 0

B) 1

C) 3

D) 4

E) 5

Q4) What is the difference between a weak-field complex and a strong-field complex?

Q5) What is the Lanthanide contraction?

Q6) Explain how EDTA is used to treat lead poisoning.

Q7) What is a coordinate covalent bond?

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