Chemistry for Engineering Students Exam Review - 3452 Verified Questions

Chemistry for Engineering Students Exam

Review

Course Introduction

Chemistry for Engineering Students introduces the fundamental principles of chemistry with a focus on their applications in engineering disciplines. The course covers atomic structure, chemical bonding, stoichiometry, thermodynamics, kinetics, and equilibrium, emphasizing the relevance of these concepts to materials, energy, and processes encountered in engineering practice. Students will develop problem-solving skills and laboratory techniques essential for analyzing chemical processes in various engineering contexts, enabling them to understand and address real-world challenges in industry, technology, and the environment.

Recommended Textbook

Chemistry A Molecular Approach 4th Edition by Nivaldo J. Tro

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) Identify the smallest volume.

A) 2.30 qt

B) 2.11 L

C) 1950 mL

D) 2015 cm³

E) 124 in³

Answer: C

Q2) A substance composed of two or more elements in a fixed,definite proportion is

A) a homogeneous mixture.

B) a heterogeneous mixture.

C) a compound.

D) a solution.

E) an alloy.

Answer: C

Q3) Which of the following is the lowest temperature?

A) 48°C

B) 75°F

C) 313 K

D) All of these temperatures are all equal.

Answer: B

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following elements is NOT a metal?

A) Mg

B) Na

C) He

D) Au

E) Mn

Answer: C

Q2) Identify the equipment that Thomson utilized to do his research.

A) cathode ray tube

B) light bulb

C) oscilloscope

D) oil atomizer, light source, microscope

E) electromagnetic field generator

Answer: A

Q3) How many protons (p)and neutrons (n)are in an atom of calcium-46?

A) 20 p, 26 n

B) 20 p, 46 n

C) 26 p, 20 n

D) 46 p, 20 n

Answer: A

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) Which one of the following is NOT an empirical formula?

A) CHO

B) CH₂O

C) C₂H₄O

D) C₂H₆O₂

E) C₃H₈O

Answer: D

Q2) Determine the mass percent (to the hundredths place)of H in sodium bicarbonate (NaHCO₃).

A) 1.20

B) 14.30

C) 57.14

D) 27.36

E) 19.05

Answer: A

Q3) Describe a structural formula.

Answer: In a structural formula,lines are used to represent covalent bonds to show how the atoms in the molecule are connected to each other.

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) The titration of 25.0 mL of an unknown concentration H₂SO₄ solution requires 83.6 mL of 0.12 M LiOH solution.What is the concentration of the H₂SO₄ solution (in M)?

A) 0.20 M

B) 0.40 M

C) 0.10 M

D) 0.36 M

E) 0.25 M

Q2) How many H⁺ ions can the acid CH₃CO₂H donate per molecule?

A) 3

B) 2

C) 1

D) 0

Q3) Which of the following compounds is soluble in water?

A) SrS

B) SrCO₃

C) PbBr₂

D) SrSO₄

E) None of these compounds is soluble in water.

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Chapter 5: Gases

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Sample Questions

Q1) Explain what the term "mean free path" describes.How does it change with decreasing pressure?

Q2) Convert 2.00 atm to torr.

A) 760 torr

B) 575 torr

C) 1520 torr

D) 1000 torr

E) 2400 torr

Q3) What is the volume of 20.0 g of argon gas at 157°C and 2.50 kPa pressure?

A) 2.58 L

B) 7.06 L

C) 262 L

D) 716 L

Q4) A gas occupies 22.4 L at STP and 19.0 L at 100°C and 1.50 atm pressure.How many moles of gas did the system gain or lose?

A) 0.08 moles gained

B) 0.01 moles gained

C) 0.01 moles lost

D) 0.07 moles lost

E) 1.40 moles gained

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Chapter 6: Thermochemistry

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Sample Questions

Q1) A 6.55 g sample of aniline (C₆H₅NH₂,molar mass = 93.13 g/mol)was combusted in a bomb calorimeter.If the temperature rose by 32.9°C,use the information below to determine the heat capacity of the calorimeter. 4 C₆H₅NH₂(l)+ 35 O₂(g) 24 CO₂(g)+ 14 H₂O(g)+ 4 NO₂(g) H°_rxn= -1.28 × 10⁴ kJ

A) 97.3 kJ/°C

B) 38.9 kJ/°C

C) 5.94 kJ/°C

D) 6.84 kJ/°C

E) 12.8 kJ/°C

Q2) Given w = 0,an endothermic reaction has the following.

A) + H and - E

B) - H and + E

C) + H and + E

D) - H and - E

Q3) Explain the difference between H and DE.

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

Q5) Give the temperature and pressure for the standard state for a liquid.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Which of the following occurs as the wavelength of a photon increases?

A) The frequency decreases.

B) The energy increases.

C) The speed decreases.

D) Planck's constant increases.

E) None of the above occurs as the wavelength of a photon increases.

Q2) Which of the following quantum numbers describes the shape of an orbital?

A) principal quantum number

B) magnetic quantum number

C) spin quantum number

D) Schrödinger quantum number

E) angular momentum quantum number

Q3) Determine the end (final)value of n in a hydrogen atom transition,if the electron starts in n = 2 and the atom absorbs a photon of light with a frequency of 4.57 × 10¹⁴ Hz.

A) 3

B) 1

C) 4

D) 6

E) 7

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Electrons filling up the orbitals from low to high is known as (the)

A) Pauli exclusion principle.

B) Hund's rule.

C) Aufbau principle.

D) Heisenberg uncertainty principle.

Q2) Give the ground state electron configuration for Br^-.

A) [Ar]4s²3d¹⁰4p⁵

B) [Ar]4s²3d¹⁰4p³

C) [Ar]4s²4p⁶

D) [Ar]4s²3d¹⁰4p⁶

E) [Ar]4s²3d⁸4p⁶

Q3) Which reaction below represents the electron affinity of Li?

A) Li(g) + e Li (g)

B) Li(g) Li (g) + e

C) Li(g) + e Li (g)

D) Li (g) Li(g) + e

E) Li (g) + e Li(g)

Q4) Why do successive ionization energies increase?

Q5) Give the ground state electron configuration for Cd²⁺.

Q6) Give the number of core electrons for Mg²⁺.

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) The antimony atom in SbCl₃ would be expected to have a

A) partial positive ( +) charge.

B) partial negative ( -) charge.

C) 3+ charge.

D) 3- charge.

Q2) How many lone pairs of electrons are on the As atom in AsF₃?

A) 0

B) 1

C) 2

D) 3

Q3) The electronegativity is 2.1 for H and 2.5 for I.Based on these electronegativities,HI would be expected to

A) be ionic and contain H^- ions.

B) be ionic and contain H⁺ ions.

C) have polar covalent bonds with a partial negative charges on the H atoms.

D) have polar covalent bonds with a partial positive charges on the H atoms.

Q4) Define bond energy.

Q5) List the least electronegative atom.

Q6) How are electron affinity and electronegativity different?

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Q7) Describe the difference between a pure covalent bond and a polar covalent bond.

Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

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

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

B) eg = tetrahedral, mg = tetrahedral, polar

C) eg = trigonal bipyramidal, mg = seesaw, polar

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

E) eg = octahedral, mg = octahedral, nonpolar

Q2) Using the VSEPR model,the molecular geometry of the central atom in SO₃ is

A) linear.

B) trigonal planar.

C) tetrahedral.

D) bent.

E) trigonal pyramidal.

Q3) The bond angle in NBr₃ is

A) 107°.

B) 104.5°.

C) 120°.

D) 60°.

E) 95°.

Page 12

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) How much energy must be removed from a 94.4 g sample of benzene (molar mass = 78.11 g/mol)at 322.0 K to solidify the sample and lower the temperature to 205.0 K? The following physical data may be useful.

H_vap = 33.9 kJ/mol

H_fus = 9.8 kJ/mol

C_liq = 1.73 J/g°C

C_gas= 1.06 J/g°C

C_sol = 1.51 J/g°C

T_melting= 279.0 K

T_boiling = 353.0 K

A) 17.6 kJ

B) 11.8 kJ

C) 70.2 kJ

D) 10.5 kJ

E) 29.4 kJ

Q2) Why is the H_vap higher than H_fusfor a given compound?

Q3) Define dynamic equilibrium.

Q4) Define boiling point of a liquid.

Q5) Define volatile.

Q6) Define viscosity.

Page 13

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Identify the addition polymer.

A) polyurethane

B) polypropylene

C) polyethylene terephthalate

D) nylon 6,6

E) polyester

Q2) Identify an element that is NOT in silicate ceramics.

A) oxygen

B) silicon

C) aluminum

D) carbon

E) hydrogen

Q3) Nickel has a face-centered cubic structure and has a density of 8.90 g/cm³.What is its atomic radius?

A) 125 pm

B) 249 pm

C) 353 pm

D) 997 pm

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

Q5) Describe the difference between the conduction band and the valence band.

Page 14

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Chapter 13: Solutions

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Sample Questions

Q1) Define colloid.

A) a mixture in which a substance floats on the dispersing medium

B) a mixture in which a substance is dissolved in a solute

C) a mixture in which a substance is dissolved in a solvent

D) a mixture in which a substance settles to the bottom of a dispersing medium

E) a mixture in which a dispersed substance is finely divided in a dispersing medium

Q2) Give the preparation of rock candy.

Q3) A solution is prepared by dissolving 16.2 g of benzene (C₆H₆)in 282 g of carbon tetrachloride (CCl₄).The concentration of benzene in this solution is ________ molal.The molar masses of C₆H₆ and CCl₄ are 78.1 g/mol and 154 g/mol,respectively.

A) 7.36 × 10^-⁴

B) 0.736

C) 0.102

D) 0.0543

E) 5.43

Q4) Calculate the osmotic pressure,in torr,of a solution containing 3.00 mg of a sugar (342 g/mole)in 15.0 mL of water at 25°C.

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The rate constant,k,for a first-order reaction is equal to 4.2 × 10^-4 s^-1.What is the half-life for the reaction?

A) 2.9 × 10^-4 s

B) 1.2 × 10³s

C) 1.7 × 10³ s

D) 2.4 × 10³ s

Q2) The decomposition of dinitrogen pentoxide is described by the following chemical equation 2 N₂O₅(g) 4 NO₂(g)+ O₂(g) If the rate of disappearance of N₂O₅is equal to 1.80 mol/min at a particular moment,what is the rate of appearance of NO₂ at that moment?

A) 0.900 mol/min

B) 1.80 mol/min

C) 3.60 mol/min

D) 7.20 mol/min

E) 1.50 mol/min

Q3) Define activation energy.

Q4) Explain how the order of a reaction can be determined.

Q5) Define half-life.

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction,equilibrium concentrations,and equilibrium constant at a particular temperature.Determine the equilibrium concentration of NO₂(g).

N₂O₄(g) 2 NO₂(g)K_c = 0.21

[N₂O₄]_eq = 0.033 M

A) 1.44 M

B) 8.3 × 10^-2M

C) 12 × 10^-2 M

D) 6.9 × 10^-3M

E) 144 M

Q2) Consider the following reaction at equilibrium.What effect will increasing the temperature have on the system? C₃H₈(g)+ 5 O₂(g) 3 CO₂(g)+ 4 H₂O(l) H° = -2220 kJ

A) The reaction will shift to the right in the direction of products.

B) The reaction will shift to the left in the direction of reactants.

C) The equilibrium constant will increase.

D) The equilibrium constant will decrease.

E) No effect will be observed.

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

Q4) Why is an equilibrium constant unitless?

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Chapter 16: Acids and Bases

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Sample Questions

Q1) Calculate the pH for an aqueous hydrobromic acid solution that contains 4.15 × 10^-3M hydronium ion.

A) 2.41 × 10^-12M

B) 4.15 × 10^-3M

C) 2.38 D) 11.62

12.50

Q2) Determine the pH of a 0.18 M H₂CO₃ solution.Carbonic acid is a diprotic acid whose K_a1 = 4.3 × 10^-7 and K_a2 = 5.6 × 10^-11.

A) 11.00

3.56

Q3) List the major source of energy in the United States.

Q4) What does the term amphoteric mean?

Q5) Do both protons ionize instantaneously from a diprotic acid such as H₂CO₃? Explain your answer.

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Chapter 17: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Describe the solubility of Al(OH)₃ with respect to pH.

A) soluble at low pH, insoluble in pH-neutral solution, and soluble at high pH

B) soluble at low pH, insoluble in pH-neutral solution, and insoluble at high pH

C) insoluble at low pH,insoluble in pH-neutral solution, and soluble at high pH

D) soluble at low pH, in pH-neutral solution, and at high pH

E) pH has no effect on the solubility.

Q2) A sample contains Ca₃(PO₄)₂,HgS,AgCl,NH₄Cl,and MnS.Identify the precipitate after the addition of 6 M HCl; then H₂S and 0.2 M HCl; then OH^-to a pH of 8; and then (NH₄)₂HPO₄ with NH₃.

A) Ca₃(PO₄)₂

B) HgS

C) AgCl

D) NH₄Cl

E) MnS

Q3) Describe the information necessary (and why)to choose a good indicator for a weak acid/strong base titration.

Q4) Define a buffer.

Q5) Identify the compound that is in stalactites and stalagmites.

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) Identify the process that is endothermic.

A) vaporization

B) deposition

C) freezing

D) condensation

Q2) Given the following equation, C₃H₈(g)+ 5 O₂(g) 3 CO₂(g)+ 4 H₂O(g) G°_rxn = -2074 kJ

Calculate G°_rxn for the following reaction.

7 C₃H₈(g)+ 35 O₂(g) 21 CO₂(g)+ 28 H₂O(g)

A) -2074 kJ

B) +14518 kJ

C) -14518 kJ

D) +2074 kJ

E) -296 kJ

Q3) Give the standard states for a gas,liquid,solid,and solution.

Q4) What is "free" energy? Give a fictitious example.

Q5) Why is the quantity of energy required to recharge a battery greater than the quantity of work done?

Page 20

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Chapter 19: Electrochemistry

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Sample Questions

Q1) Define extraction.

A) Free metal is oxidized out of its ore.

B) Undesired reduction of metals from reducing agents in the environment

C) Plating a film on a metal to protect it

D) Undesired oxidation of metals from oxidizing agents in the environment

E) Free metal is reduced out of its ore.

Q2) In a nickel-metal hydride battery,NiO(OH)is reduced to

A) Ni(OH)₄.

B) NiO₂.

C) Ni.

D) Ni(OH)₂.

E) NiOH.

Q3) Describe the reactions during the electrolysis of water.

A) Oxygen is reduced and hydrogen is oxidized.

B) Oxygen and hydrogen are both oxidized.

C) Oxygen and hydrogen are both reduced.

D) Oxygen is oxidized and hydrogen is reduced.

E) Neither oxygen nor hydrogen are oxidized or reduced.

Q4) Give an example of an inert electrode.

Q5) Why are iron nails coated with zinc?

Page 21

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Identify the elements used in radiometric dating.

A) carbon-14 to nitrogen-14

B) uranium-238 to lead-206

C) potassium-40 to argon-40

D) all of the above

E) none of the above

Q2) The age of an ancient tree trunk is estimated using radiocarbon dating.If the trunk has a C-14 decay rate that is 34% of what it is in living plants,how old is the trunk? The half-life of C-14 is 5730 years.

A) 2.92 × 10⁴ years

B) 1.94 × 10⁴ years

C) 8.92 × 10³ years

D) 5.31 × 10³ years

E) 1.74 × 10² years

Q3) Determine the half-life of a nuclide that loses 38.0% of its mass in 407 hours.

A) 204 hours

B) 568 hour

C) 590 hours

D) 281 hours

E) 291 hours

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Chapter 21: Organic Chemistry

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Q1) Write the balanced chemical equation for the addition of HBr to CH₂=CHCH₂CH₃.

A) CH₂=CHCH₂CH₃ + HBr CH₃BrCH₂CH₂CH₃

B) CH₂=CHCH₂CH₃ + 2 HBr 2 CH₂BrCH₃

C) CH₂=CHCH₂CH₃ + 2 HBr CH₃Br + CH₂BrCH₂CH₃

D) CH₂=CHCH₂CH₃ + 4 HBr 4 CH₃Br E) CH₂=CHCH₂CH₃ + HBr CH₃CHBrCH₂CH₃

Q2) Give the organic product for the following reaction in the presence of sulfuric acid. CH₃CH₂COOH + CH₃CH₂CH₂OH A) CH₃CH₂COOCH₂CH₃ B)

CH₃CH₂COOCH₂CH₂CH₃ C)

CH₃CH₂CH₂COOCH₂CH₃ D)

CH₃CH₂COCH₂CH₂CH₃ E) CH₃CH₂OCH₂CH₂CH₃

Page 23

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Chapter 22: Biochemistry

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Q1) Which of the following link together amino acid units?

A) glycosidic linkages

B) sulfide linkages

C) peptide bonds

D) hydrogen bonds

E) ester linkages

Q2) Identify the type of bond between the two strands of DNA.

A) hydrogen bonds

B) London dispersion forces

C) dipole-dipole bonds

D) amide bonds

E) water bonds

Q3) The following is part of a DNA sequence.What is its complementary sequence?

TCGTAAGCTTGCG

A) AGCATTCGAACGC

B) CGCAAGCTTACGA

C) TCATTCAC

D) GCGTTCGAATGCT

E) TCGTAAGCTTGCG

Q4) How do cis- and trans-fats differ?

Page 24

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Chapter 23: Chemistry of the Nonmetals

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Q1) Which of the following statements is TRUE?

A) Each boron nitride unit contains six valence electrons.

B) Carbon nanotubes and boron nitride nanotubes are not similar.

C) Boron nitride contains BN units that are isoelectric with carbon.

D) Boron nitrides conduct an electrical current.

E) The size and electronegativity of a carbon atom are not close to the average of those properties for a boron atom and a nitrogen atom.

Q2) Describe the Frasch process for obtaining sulfur.

Q3) Determine the number of vertices and faces in the closo-Borane B₇H₇^2-.

A) vertices = 4, faces = 10

B) vertices = 7, faces = 14

C) vertices = 3, faces = 6

D) vertices = 14, faces = 24

E) vertices = 7, faces = 10

Q4) Draw the Lewis structure for IF₄ and determine the geometry of the ion.

Q5) Describe the major production method for obtaining oxygen.

Q6) Why are phosphate compounds added to detergents?

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Chapter 24: Metals and Metallurgy

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Q1) Identify the metal used as a coating for steel objects such as nails.

A) gold

B) silver

C) zinc

D) palladium

E) manganese

Q2) Which of the following is TRUE concerning interstitial alloys?

A) The metal takes on the properties of the foreign atoms once they are present.

B) Small atoms fit in between the metal atoms in the crystal structure.

C) Brass is an example of an interstitial alloy.

D) The foreign atoms present weaken the metallic structure.

E) None of the above is true.

Q3) Metals make up only ________ percent of the earth's crust.

A) 40

B) 50

C) 25

D) 30

E) 20

Q4) Why is zinc used to coat steel objects?

Q5) Why does a "two-phase" structure occur in a substitutional alloy?

Page 26

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Which of the following can function as a chelating agent?

A) OH^-

B) H₂O

C) C₂O₄^2- 

D) NH₃

Q2) How many unpaired electrons would you expect for the complex ion: [FeCl₆]^4-?

A) 2

B) 6

C) 4

D) 3

E) 0

Q3) How many unpaired electrons would you expect for the complex ion: [Co(NH₃ )₆]⁴⁺?

A) 1

B) 0

C) 5

D) 2

E) 3

Q4) What is the Lanthanide contraction?

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