Chemistry for Engineers Midterm Exam - 3504 Verified Questions

Course Introduction

Chemistry for Engineers

Midterm Exam

Chemistry for Engineers provides a foundational understanding of chemical principles and their practical applications within various engineering disciplines. Topics cover atomic and molecular structure, chemical bonding, thermodynamics, kinetics, materials chemistry, and electrochemistry, with an emphasis on real-world engineering contexts such as material selection, reaction design, and environmental considerations. Through lectures, laboratory experiments, and problem-solving exercises, students gain the analytical skills and scientific knowledge necessary to address chemical challenges in engineering projects, ensuring safe, efficient, and innovative solutions to complex technical problems.

Recommended Textbook

Chemistry The Central Science 12th Edition by

Theodore E. Brown

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Page 2

Chapter 1: Introduction: Matter and Measurement

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Sample Questions

Q1) If an object, beginning at rest, is moving at a speed of 700 m/s after 2.75 min, its rate of acceleration (in m/s²)is __________. (Assume that the rate of acceleration is constant.)

A)1.16 × 10⁵

B)255

C)193

D)4.24

E)1.53 × 10⁴

Answer: D

Q2) Cu is the symbol for the element __________. Answer: Copper

Q3) How many significant figures are in the number 0.0034050?

A)3

B)4

C)5

D)6

E)7

Answer: C

Q4) 1 milligram = __________ micrograms

Answer: 1,000

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Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) What is the molecular formula for heptane?

A)C₆H₁₂

B)C₆H₁₄

C)C₇H₁₄

D)C₇H₁₆

E)C₇H₁₈

Answer: D

Q2) The charge on an electron was determined in the __________.

A)cathode ray tube, by J. J. Thompson

B)Rutherford gold foil experiment

C)Millikan oil drop experiment

D)Dalton atomic theory

E)atomic theory of matter

Answer: C

Q3) What is the name of an alcohol derived from hexane?

Answer: hexanol

Q4) H₂SeO₄ is called selenic acid.

A)True

B)False

Answer: True

Page 4

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Chapter 3: Stoichiometry: Calculations With Chemical

Formulas and Equations

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Sample Questions

Q1) A sample of CH₄O with a mass of 32.0 g contains __________ molecules of CH₄O.

A)5.32 × 10^-23

B)1.00

C)1.88 × 10²²

D)6.02 × 10²³

E)32.0

Answer: D

Q2) When the following equation is balanced, the coefficient of O₂ is __________. C₂H₄O (g)+ O₂ (g)

CO₂ (g)+ H₂O (g)

A)2

B)3

C)4 D)5

E)1

Answer: D

Q3) A compound was found to contain 90.6% lead (Pb)and 9.4% oxygen. The empirical formula for this compound is __________. Answer: Pb₃O₄

Page 5

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Chapter 4: Aqueous Reactions and Solution Stoichiometry

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Sample Questions

Q1) Combining aqueous solutions of BaI₂ and Na₂SO₄ affords a precipitate of BaSO₄. Which ion(s)is/are spectator ions in the reaction?

A)Ba²⁺ only

B)Na⁺ only

C)Ba²⁺ and SO₄^2-

D)Na⁺ and I^-

E)SO₄^2- and I^-

Q2) A 17.5 mL sample of an acetic acid (CH₃CO₂H)solution required 29.6 mL of 0.250 M NaOH for neutralization. The concentration of acetic acid was __________ M.

A)0.158

B)0.423

C)134

D)6.88

E)0.214

Q3) The solvent in an aqueous solution is __________.

Q4) How many grams of CH₃OH must be added to water to prepare 150mL of a solution that is 2.0 M CH₃OH?

Q5) What is aqua regia?

6

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Chapter 5: Thermochemistry

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Sample Questions

Q1) For a given process at constant pressure, H is negative. This means that the process is __________.

A)endothermic

B)equithermic

C)exothermic

D)a state function

E)energy

Q2) A sample of calcium carbonate [CaCO₃ (s)] absorbs 45.5 J of heat, upon which the temperature of the sample increases from 21.1 °C to 28.5 °C. If the specific heat of calcium carbonate is 0.82 J/g-K, what is the mass (in grams)of the sample?

A)3.7

B)5.0

C)7.5

D)410

E)5.0 x 10³

Q3) The standard enthalpy change of a reaction is the enthalpy change when all reactants and products are at ________ pressure and a specific temperature.

Q4) __________ is defined as the energy used to move an object against a force.

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Chapter 6: Electronic Structure of Atoms

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Sample Questions

Q1) The correct ground-state electron configuration for silver is __________.

A)[Kr]5s²4d⁹

B)[Kr]5s¹4d¹⁰

C)[Kr]5s²4d¹⁰

D)[Xe]5s²4d⁹

E)[Xe]5s¹4d¹⁰

Q2) When the value of n is greater than or equal to 3, electrons can reside in d orbitals. A)True

B)False

Q3) At what speed (m/s)must a 10.0-mg object be moving to have a de Broglie wavelength of 3.3 × 10^-41 m?

A)4.1

B)1.9 × 10^-11

C)2.0 × 10¹²

D)3.3 × 10^-42

E)1.9 × 10¹³

Q4) The largest principal quantum number in the ground state electron configuration of francium is __________.

Q5) The ground state electron configuration of copper is __________.

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Chapter 7: Periodic Properties of the Elements

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Sample Questions

Q1) Which of the following is an isoelectronic series?

A)B^5-, Si^4-, As^3-, Te^2-

B)F^-, Cl^-, Br^-, I^-

C)S, Cl, Ar, K

D)Si^2-, P^2-, S^2-, Cl^2-

E)O^2-, F^-, Ne, Na⁺

Q2) Which periodic table group contains only metals?

A)8A

B)2A

C)6A

D)7A

E)5A

Q3) The atomic radius of iodine is one-half the distance separating the iodine nuclei. A)True

B)False

Q4) As successive electrons are removed from an element, the ionization energy

Q5) Which noble gas has the highest first ionization energy?

Q6) [Kr]5s² is the electron configuration for __________.

Page 9

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Chapter 8: Basic Concepts of Chemical Bonding

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Sample Questions

Q1) In which of the molecules below is the carbon-carbon distance the shortest?

A)H₂C-CH₂

B)H-C C-H

C)H₃C-CH₃

D)H₂C-C-CH₂

E)H₃C-CH₂-C H₃

Q2) The only noble gas without eight valence electrons is __________.

A)Ar

B)Ne

C)He

D)Kr

E)All noble gases have eight valence electrons.

Q3) Bond enthalpy is __________.

A)always positive

B)always negative

C)sometimes positive, sometimes negative

D)always zero

E)unpredictable

Q4) Write the balanced chemical equation for the reaction for which H_rxn is the lattice energy for potassium bromide.

Page 10

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Chapter 9: Molecular Geometry and Bonding Theories

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Sample Questions

Q1) The central iodine atom in IF₅ has __________ unbonded electron pairs and __________ bonded electron pairs in its valence shell.

A)1, 5

B)0, 5

C)5, 1

D)4, 1

E)1, 4

Q2) Based on molecular orbital theory, the bond order of the N-N bond in the N₂²⁺ ion is __________.

A)0

B)3

C)1

D)2

E)1/2

Q3) Each molecular orbital can accommodate, at most, two electrons with their spins paired. This is called the __________.

Q4) In molecular orbital theory the stability of a covalent body is related to its

Q5) What are the three bond angles in the trigonal bipyramidal structure?

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Chapter 10: Gases

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Sample Questions

Q1) The molecular weight of a gas that has a density of 5.75 g/L at STP is __________ g/mol.

A)3.90

B)129

C)141

D)578

E)1.73 × 10^-3

Q2) A vessel contained N₂, Ar, He, and Ne. The total pressure in the vessel was 987 torr. The partial pressures of nitrogen, argon, and helium were 44.0, 486, and 218 torr, respectively. The partial pressure of neon in the vessel was __________ torr.

A)42.4

B)521

C)19.4

D)239

E)760

Q3) Abnormally high blood pressure is called __________.

Q4) What is the density (in g/L)of oxygen gas at 77.0°C and 700.0 torr?

Q5) The temperature and pressure specified by STP are __________ °C and __________ atm.

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Chapter 11: Liquids and Intermolecular Forces

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Sample Questions

Q1) Based on the figure above, the boiling point of diethyl ether under an external pressure of 1.32 atm is __________°C.

A)10

B)20

C)30

D)40

E)0

Q2) What is the predominant intermolecular force in CH₃CH₂OH?

A)London-dispersion forces

B)ion-dipole attraction

C)ionic bonding

D)induced dipole-dipole attraction

E)hydrogen-bonding

Q3) What type(s)of intermolecular forces exist between Cl₂ and CO₃^-2?

A)dispersion forces

B)dispersion forces and ion-dipole

C)dispersion forces, ion-dipole, and induced dipole - induced dipole

D)dispersion forces and ion-induced dipole

E)dispersion forces, ion-dipole, dipole-dipole, and ion-induced dipole

Page 13

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) When lattice points occur at the center of each face, as well as each corner of a unit cell, the cell is called __________.

Q2) An alloy is a

A)heterogeneous mixture of two metals.

B)pure metal.

C)metallic material that is composed of two or more elements.

D)nonmetal with some properties of a metal.

E)a mineral containing two or more metals.

Q3) Semiconductors are less conductive than metals because of __________ gap.

Q4) The type of solid that is characterized by low melting point, softness, and low electrical conduction is a covalent-network solid.

A)True

B)False

Q5) Blue LEDs are usually made of __________.

A)GaAs

B)GaP

C)GaO

D)GaS

E)GaN

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Chapter 13: Properties of Solutions

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Sample Questions

Q1) The freezing point of ethanol (C₂H₅OH)is -114.6°C. The molal freezing point depression constant for ethanol is 2.00°C/m. What is the freezing point (°C)of a solution prepared by dissolving 40.0 g of glycerin (C₃H₈O₃, a nonelectrolyte)in 200.0 g of ethanol?

A)-115.0

B)-4.34

C)-132.3

D)-118.9

E)-114.6

Q2) The vapor pressure of pure ethanol at 60°C is 0.459 atm. Raoult's Law predicts that a solution prepared by dissolving 10.0 mmol naphthalene (nonvolatile)in 90.0 mmol ethanol will have a vapor pressure of __________ atm.

A)0.498

B)0.413

C)0.790

D)0.367

E)0.0918

Q3) The formula weight of FeCl₃^.6H₂O is

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Page 15

Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The following reaction is second order in [A] and the rate constant is 0.025 M^-1s^-1: A B

The concentration of A was 0.65 M at 33 s. The initial concentration of A was __________ M.

A)2.4

B)0.27

C)0.24

D)1.4

E)1.2 × 10^-2

Q2) The average rate of disappearance of A between 10 s and 20 s is __________ mol/s.

A)2.2 × 10^-3

B)1.1 × 10^-3

C)4.4 × 10^-3

D)454

E)9.90 × 10^-3

Q3) The uptake of molecules into the interior of another substance is referred to as

Q4) Reaction rate data showing temperature dependence obey an equation devised by

Page 16

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) For an exothermic reaction, increasing the reaction temperature results in a(an)__________ in K.

Q2) Which of the following statements is true?

A)Q does not change with temperature.

B)K_eq does not change with temperature, whereas Q is temperature dependent.

C)K does not depend on the concentrations or partial pressures of reaction components. D)Q does not depend on the concentrations or partial pressures of reaction components.

E)Q is the same as K_eq when a reaction is at equilibrium.

Q3) Pure __________ and pure __________ are excluded from equilibrium-constant expressions.

Q4) The effect of a catalyst on an equilibrium is to __________.

A)increase the rate of the forward reaction only

B)increase the equilibrium constant so that products are favored

C)slow the reverse reaction only

D)increase the rate at which equilibrium is achieved without changing the composition of the equilibrium mixture

E)shift the equilibrium to the right

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Chapter 16: Acid-Base Equilibria

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Sample Questions

Q1) A 0.1 M aqueous solution of __________ will have a pH of 7.0 at 25.0°C. NaOCl KCl

NH₄Cl Ca(OAc)₂

A)NaOCl

B)KCl

C)NH₄Cl

D)Ca(OAc)₂

E)KCl and NH₄Cl

Q2) What is the conjugate acid of NH₃?

A)NH₃

B)NH₂⁺

C)NH₃⁺

D)NH₄⁺

E)NH₄OH

Q3) A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor.

A)True

B)False

Q4) An acid containing the COOH group is called a carbo-oxy acid.

A)True

B)False

Page 18

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Chapter 17: Additional Aspects of Aqueous Equilibria

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Q1) A 25.0 mL sample of a solution of an unknown compound is titrated with a 0.115 M NaOH solution. The titration curve above was obtained. The unknown compound is

A)a strong acid

B)a strong base

C)a weak acid

D)a weak base

E)neither an acid nor a base

Q2) The pH of a solution prepared by mixing 50.0 mL of 0.125 M NaOH and 40.0 mL of 0.125 M HNO₃ is __________.

A)13.29

B)7.00

C)8.11

D)11.00

E)none of the above

Q3) In general, the solubility of a slightly soluble salt is __________ by the presence of a second solute that furnishes a common ion.

Q4) An assembly of a metal ion and the Lewis bases bonded to it is called a

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Chapter 18: Chemistry of the Environment

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Sample Questions

Q1) The greenhouse effect of a methane molecule far exceeds the effect of a carbon dioxide molecule.

A)True

B)False

Q2) What compound in limestone and marble is attacked by acid rain?

A)hydroxyapatite

B)calcium carbonate

C)gypsum

D)graphite

E)potassium hydroxide

Q3) The concentration of Br^- in a sample of seawater is 9.3 × 10^-4 M. If a liter of seawater has a mass of 1.0 kg, the concentration of Br^- is __________ ppm.

A)0.074

B)74

C)0.93

D)9.3

E)0.0093

Q4) The production of dead and decaying plant material is a process called

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Chapter 19: Chemical Thermodynamics

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Sample Questions

Q1) Which reaction produces a decrease in the entropy of the system?

A)4 NH₃ (g)+ 5 O₂ (g) 4 NO (g)+ 6 H₂O (g)

B)Na (s)+ 1/2Cl₂ (g) NaCl (s)

C)2 HgO (s) 2 Hg (l)+ O₂ (g)

D)U (s)+ 3F₂ (g) UF₆(s)

E)H₂O (s) H₂O (g)

Q2) The value of S° for the formation of calcium chloride from its constituent elements, Ca (s)+ Cl₂ (g) CaCl₂(s) Is __________ J/K mol.

A)-104.6

B)+104.6

C)+369.0

D)-159.8

E)+159.8

Q3) Which of the following statements is true?

A)Processes that are spontaneous in one direction are spontaneous in the opposite direction.

B)Processes are spontaneous because they occur at an observable rate.

C)Spontaneity can depend on the temperature.

D)All of the statements are true.

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Chapter 20: Electrochemistry

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Q1) A voltaic cell is constructed with two Zn²⁺-Zn electrodes, where the half-reaction is Zn²⁺ + 2e^- Zn (s)E° = -0.763 V

The concentrations of zinc ion in the two compartments are 5.50 M and 1.11 × 10^-2 M, respectively. The cell emf is __________ V.

A)-1.54 × 10^-3

B)-378

C)0.0798

D)0.160

E)-0.761

Q2) Which one of the following reactions is a redox reaction?

A)NaOH + HCl NaCl + H₂O

B)Pb²⁺ + 2Cl^- PbCl₂

C)AgNO₃ + HCl HNO₃ + AgCl

D)None of the above is a redox reaction.

Q3) Calculate the number of grams of aluminum produced in 30.0 minutes by electrolysis of AlCl₃ at a current of 12.0 A.

Q4) At constant temperature and pressure the Gibbs free energy value is a measure of the __________ of a process.

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Chapter 21: Nuclear Chemistry

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Sample Questions

Q1) Positron emission causes a decrease of one in the atomic number.

A)True

B)False

Q2) What is the typical percent of uranium-235 in the enriched UO₂ pellets used in nuclear reactors?

A)0.7

B)1

C)3

D)10

E)14

Q3) The curie is a measure of the

A)number of disintegrations per second of a radioactive substance.

B)total energy absorbed by an object exposed to a radioactive source.

C)lethal threshold for radiation exposure.

D)number of alpha particles emitted by exactly one gram of a radioactive substance.

E)None of the above is correct.

Q4) What happens in the nucleus of an atom that undergoes positron emission?

Q5) What isotope of what element is produced if krypton-81 undergoes beta decay?

Q6) The use of radioisotopes in tracing metabolism is possible because __________.

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Chapter 22: Chemistry of the Nonmetals

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Sample Questions

Q1) Which halogen forms an oxyacid with the formula HXO₂?

A)bromine

B)fluorine

C)chlorine

D)iodine

E)astatine

Q2) What process replenishes O₂?

Q3) Of the following, which is most likely to form interstitial carbides?

A)active metals

B)transition metals

C)boron and silicon

D)alkaline earth metals

E)alkali metals

Q4) The compound whose formula is CaC₂ is __________.

A)calcium carbide

B)carborundum

C)carbon calcide

D)calcium dicarbon

E)limestone

Q5) KNO₃ and NaNO₃ are also known as __________.

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Chapter 23: Transition Metals and Coordination Chemistry

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Q1) How many ligands are there in the coordination sphere of [Co(en)₂Cl₂]⁺?

A)3

B)6

C)4

D)1

E)0

Q2) The chelate effect is enhanced by polydentate ligand binding because of the change in __________.

Q3) A substance with unpaired electrons will be

A)slightly attracted to a magnet.

B)slightly repelled by a magnet.

C)permanently magnetic.

D)brightly colored.

E)nonmetallic.

Q4) A large difference in formation constant (K_f)of a poly- versus monodentate ligand is called __________.

Q5) To separate racemic mixtures the isomers must be in a chiral environment.

A)True

B)False

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Chapter 24: The Chemistry of Life: Organic and Biological Chemistry

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Q1) The melting and boiling points of hydrocarbons are determined by __________.

A)ion-dipole attraction

B)dipole-dipole attraction

C)London forces

D)hydrogen bonding

E)ionic bonding

Q2) Benzene behaves differently from a hydrocarbon which simply contains three C-C bonds in that the latter would be expected to react much more readily with

A)H₂

B)Cl₂

C)Br₂

D)HCl

E)all of the above

Q3) Of the 20 amino acids found in our bodies, __________ of them must be ingested because our bodies cannot synthesize sufficient quantities of them.

Q4) In DNA adenine is always paired with __________.

Q5) The hydrolysis of an ester in the presence of a base is called __________.

Q6) Why is cyclopropane more reactive than propane?

Q7) Non-superimposable mirror-image isomers of a substance are called Page 26

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