Chemistry for Engineers Test Bank - 2309 Verified Questions

Chemistry for Engineers Test

Bank

Course Introduction

Chemistry for Engineers is an essential course designed to provide engineering students with a foundational understanding of chemical principles and their practical applications in engineering contexts. The course covers fundamental concepts including atomic structure, chemical bonding, thermodynamics, kinetics, electrochemistry, materials chemistry, and environmental aspects of chemistry. Emphasis is placed on real-world examples and problem-solving related to materials selection, corrosion, reaction engineering, and energy conversion. Through lectures, laboratory exercises, and case studies, students gain the knowledge and skills necessary to integrate chemical concepts into the design, operation, and analysis of engineering systems.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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24 Chapters

2309 Verified Questions

2309 Flashcards

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Chapter 1: Keys to the Study of Chemistry

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79 Verified Questions

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Sample Questions

Q1) Select the best statement about chemistry before 1800.

A) Alchemy focused on objective experimentation rather than mystical explanations of processes.

B) The phlogiston theory laid a valuable theoretical basis for modern chemistry.

C) Lavoisier's quantitative work on the role of oxygen in combustion was the beginning of modern chemistry.

D) The interpretation of data by alchemists was not biased by their overall view of life.

E) Alchemists failed because they did not develop any practical chemical methods.

Answer: C

Q2) The speed needed to escape the pull of Earth^'s gravity is 11.3 km/s. What is this speed in mi/h?

A) 65,500 mi/h

B) 25,300 mi/h

C) 18,200 mi/h

D) 1,090 mi/h

E) 5.02 × 10^-3mi/h

Answer: A

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Chapter 2: The Components of Matter

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Sample Questions

Q1) Sodium peroxide is an oxidizer used to bleach animal and vegetable fibers. What is its formula?

A) NaO

B) NaO₂

C) Na₂O₂

D) Na₂O

E) NaH₂O₂

Answer: C

Q2) Zinc acetate is used in preserving wood and in manufacturing glazes for porcelain. What is its formula?

A) ZnAc₂

B) ZnCH₃COO

C) Zn(CH₃COO)₂

D) Zn₂CH₃COO

E) ZnCH₃COCH₃

Answer: C

Q3) Copper (Cu) is a transition metal.

A)True

B)False

Answer: True

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Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Hydroxylamine hydrochloride is a powerful reducing agent which is used as a polymerization catalyst. It contains 5.80 mass % H, 20.16 mass % N, 23.02 mass % O, and 51.02 mass % Cl. What is its empirical formula?

A) H₂N₇O₈Cl₁₈

B) H₂N₂O₂Cl

C) HN₃O₄Cl₉

D) H₄NOCl

E) H₄NOCl₂

Answer: D

Q2) Determine the percent composition of potassium dichromate, K₂Cr₂O₇.

A) 17.5% K, 46.6% Cr, 35.9% O

B) 29.8% K, 39.7% Cr, 30.5% O

C) 36.5% K, 48.6% Cr, 14.9% O

D) 37.2% K, 24.7% Cr, 38.1% O

E) none of the above

Answer: B

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Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) Which of the following is a strong acid?

A) H₃PO₄

B) HNO₃

C) HF

D) CH₃COOH

E) H₂O

Q2) Select the classification for the following reaction. KOH(aq) + HCl(aq) \(\to\) KCl(aq) + H₂O(l)

A) precipitation

B) acid-base

C) redox

D) combination

E) none of the above

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

A) Na₂S

B) PbI₂

C) AgCl

D) CuS

E) Ca₃(PO₄)₂

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Page 6

Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) If the molecular mass of a gas increases by a factor of 4 at constant temperature, its rms speed will

A) decrease by a factor of 4.

B) increase by a factor of 4.

C) decrease by a factor of 16.

D) increase by a factor of 16.

E) decrease by a factor of 2.

Q2) A sample container of carbon monoxide occupies a volume of 435 mL at a pressure of 785 torr and a temperature of 298 K. What would its temperature be if the volume were changed to 265 mL at a pressure of 785 torr?

A) 182 K

B) 298 K

C) 387 K

D) 489 K

E) 538 K

Q3) Aluminum metal shavings (10.0 g) are placed in 100. mL of 6.00 M hydrochloric acid. What is the maximum volume of hydrogen, measured at STP, which can be produced? 2Al(s) + 6HCl(aq) \(\to\) 2AlCl₃(aq) + 3H₂(g)

Q4) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650. torr.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Sand is converted to pure silicon in a three step process. The third step is SiCl₄(g) + 2Mg(s) \(\to\) 2MgCl₂(s) + Si(s) \(\Delta\)H = -625.6 kJ

What is the enthalpy change when 25.0 mol of silicon tetrachloride is converted to elemental silicon?

A) -25.0 kJ

B) -7820 kJ

C) -1.56 × 10⁴ kJ

D) -3.13 × 10⁴kJ

E) none of the above

Q2) Galena is the ore from which elemental lead is extracted. In the first step of the extraction process, galena is heated in air to form lead(II) oxide. 2PbS(s) + 3O₂(g) \(\to\)2SO₂(g) \(\Delta\)H = -827.4 kJ

What mass of galena is converted to lead oxide if 975 kJ of heat are liberated?

A) 203 g

B) 282 g

C) 406 g

D) 478 g

E) 564 g

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) The Rydberg equation is an example of an empirical equation.

A)True

B)False

Q2) As the frequency of electromagnetic radiation increases, its wavelength also increases.

A)True

B)False

Q3) Interference of light waves

A) separates light into its component colors.

B) creates a pattern of light and dark regions.

C) focuses a broad beam of light into a point.

D) bends light as it passes the edge of an object.

E) creates a laser beam.

Q4) Which scientist demonstrated that photons transferred momentum during collisions with matter?

A) Bohr

B) de Broglie

C) Planck

D) Compton

E) Billiard

Page 9

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Elements in which the outermost electron has the same principal quantum number n, show similar chemical properties.

A)True

B)False

Q2) What is the correct order of decreasing size of the following ions?

A) P^3-> Cl^-> K⁺> Ca²⁺

B) Ca²⁺ > K⁺ > Cl^- > P^3-

C) K⁺ > Cl^-> Ca²⁺ > P^3-

D) K⁺> Cl^- > P^3-> Ca²⁺

E) None of the above is correct.

Q3) How many valence electrons are there in an atom with the electron configuration [noble gas]ns²(n - 1)d⁸?

A) 2

B) 6

C) 8

D) 10

E) none of the above

Q4) State Hund's rule, and show how it applies to the ground state of phosphorus atoms.

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Page 10

Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Most of the copper sold in major metal markets is highly purified, typically to 99.99%. Why is this?

Q2) Describe in brief how electronegativity values can be used to predict the percent ionic character of a bond between two atoms.

Q3) The majority of elements are good electrical conductors when in solid form.

A)True B)False

Q4) When two atoms form a covalently-bonded diatomic molecule, the distance between the nuclei at which the potential energy is at a minimum is called A) the bond energy.

B) the bond length.

C) the molecular diameter.

D) the covalent radius.

E) the covalent diameter.

Q5) The lattice energy of large ions is greater in magnitude than that of small ions of the same charge.

A)True B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the ideal bond angles in AsCl₃using the molecular shape given by the VSEPR theory.

A) 90°

B) 109°

C) 120°

D) 180°

E) between 110 and 120°

Q2) Draw the Lewis structure of XeF₄. Use this structure, in conjunction with VSEPR theory, to predict the shape of this molecule. Outline your reasoning.

Q3) What is the molecular shape of ClF₄^-as predicted by the VSEPR theory?

A) square pyramidal

B) square planar

C) see-saw

D) octahedral

E) tetrahedral

Q4) 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.

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) One can safely assume that the 3s- and 3p-orbitals will form molecular orbitals similar to those formed when 2s- and 2p-orbitals interact. According to molecular orbital theory, what will be the bond order for the Cl₂⁺ ion?

A) 0.5

B) 1

C) 1.5

D) 2

E) none of the above

Q2) What type of hybridization is needed to explain why ethyne, C₂H₂, is linear?

Q3) Valence bond theory predicts that xenon will use _____ hybrid orbitals in XeOF₄.

A) sp

B) sp²

C) sp³

D) sp³d

E) sp³d²

Q4) In not more than two sentences, explain when and why chemists make use of the concept of hybridization.

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

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Sample Questions

Q1) A sample of octane in equilibrium with its vapor in a closed 1.0-L container has a vapor pressure of 50.0 torr at 45°C. The container's volume is increased to 2.0 L at constant temperature and the liquid/vapor equilibrium is reestablished. What is the vapor pressure?

A) > 50.0 torr

B) 50.0 torr

C) 25.0 torr

D) The mass of the octane vapor is needed to calculate the vapor pressure.

E) The external pressure is needed to calculate the vapor pressure.

Q2) Which of the following terms refers to the resistance of a liquid to flow?

A) surface tension

B) capillary action

C) viscosity

D) adhesion

E) cohesion

Q3) Ceramic superconductors often contain copper in unusual oxidation states.

A)True

B)False

Q4) How do the electrical properties of semiconductors differ from those of metals?

Page 14

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Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) Raoult's Law relates the vapor pressure of the solvent above the solution to its mole fraction in the solution. Which of the following is an accurate statement?

A) Raoult's Law applies exactly to all solutions.

B) Raoult's Law works best when applied to concentrated solutions.

C) Raoult's Law works best when applied to dilute solutions.

D) Raoult's Law applies only to non-ideal solutions.

E) None of the above statements is accurate.

Q2) Which of the following sets of conditions could exist when two liquids which are completely miscible in one another are mixed?

A) .\(\Delta\)H_soln > 0, entropy of system decreases

B) .\(\Delta\)H_soln \(\approx\) 0, entropy of system decreases

C) .\(\Delta\)H_soln \(\approx\) 0, entropy change of system \(\approx\) 0

D) .\(\Delta\)H_soln \(\approx\) 0, entropy of system increases E) none of the above

Q3) A 50.0 % (w/w) solution of sulfuric acid (H₂SO₄) in water has a density of 1.395 g/mL. What is its molarity?

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Page 15

Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) Predict the molecular shape of XeF₃^-.

A) linear

B) trigonal planar

C) trigonal pyramid

D) T-shaped

E) square

Q2) Bromine will form compounds with each of the other elements in Period 4 of the periodic table. How does the type of bonding in the compounds change as one moves from potassium bromide to selenium bromide?

A) ionic to polar covalent

B) polar covalent to ionic

C) polar covalent to non-polar covalent

D) coordinate covalent to polar covalent

E) none of the above

Q3) Potassium nitrate is a strong reducing agent.

A)True

B)False

Q4) The Haber process is the first step in the manufacture of sulfuric acid.

A)True

B)False

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) Dacron is a ___________ compound and is produced by a __________ reaction.

A) polyamide; addition

B) polyamide; condensation

C) polyester; addition

D) polyester; condensation

E) polyene; addition

Q2) A characteristic reaction of alkanes is addition.

A)True

B)False

Q3) Select the correct type for the following reaction. CH₃CH₂CH₂CH₂OH + HBr \(\to\) CH₃CH₂CH₂CH₂Br + H₂O

A) dehydration

B) dehydroxylation

C) addition

D) elimination

E) substitution

Q4) Explain how a disulfide bridge can arise a protein molecule.

Page 17

Q5) Name the three component parts of a nucleotide.

Q6) Draw and name two non-alkenes with the formula C₄H₈.

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) 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]

Q2) An increase in temperature increases the reaction rate because

A) a greater fraction of the collisions have the correct orientation of molecules. B) the activation energy of the reaction will increase. C) the activation energy of the reaction will decrease.

D) temperature acts as a catalyst in chemical reactions.

E) more collisions will have enough energy to exceed the activation energy.

Q3) Briefly list the features/properties common to all catalysts and how they work. Draw a labeled reaction energy diagram as part of your answer.

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) Unless \(\Delta\)H°_rxn = 0, a change in temperature will affect the value of the equilibrium constant K_c.

A)True

B)False

Q2) The two equilibrium constants for the same reaction, K_c and K_p, will always equal one another when

A) all of the reactants and products are gases.

B) in the reaction equation, the number of moles of gaseous products equals the number of moles of gaseous reactants.

C) in the reaction equation, the number of moles of gaseous products is greater than the number of moles of gaseous reactants.

D) in the reaction equation, the number of moles of gaseous products is smaller than the number of moles of gaseous reactants.

E) in the reaction equation, the total number of moles of reactants equals that of the products.

Q3) For some gas-phase reactions, K_p = K_c.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) A 1.25 M solution of the weak acid HA is 9.2% dissociated. What is the pH of the solution?

A) 0.64

B) 0.94

C) 1.13

D) 2.16

E) none of the above

Q2) The ammonium ion, NH₄⁺, is a weak acid.

A)True

B)False

Q3) The substance Ba(OH)₂is considered

A) a weak Arrhenius acid.

B) a weak Arrhenius base.

C) a strong Arrhenius acid.

D) a strong Arrhenius base.

E) a neutral compound.

Q4) Define an acid according to the Arrhenius theory, and write a balanced equation to support this definition.

Q5) Describe what is meant by the "leveling effect". Use a real acid as an example, and write an appropriate equation.

Page 21

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Calculate the solubility of silver oxalate, Ag₂C₂O₄, in pure water. K_sp = 1.0 × 10^-11

A) 1.4 × 10^-4M

B) 8.2 × 10^-5M

C) 5.4 × 10^-5M

D) 3.2 × 10^-6M

E) 2.5 × 10^-12M

Q2) What is the maximum mass of KCl that can be added to1.0 L of a 0.010 M lead(II) chloride solution without causing any precipitation of lead(II) chloride? Assume that addition of KCl does not affect the solution volume. For lead(II) chloride, K_sp = 1.6 × 10^-5.

A) 3.0 g

B) 1.5 g

C) 0.8 g

D) 1.0 g

E) 0.2 g

Q3) Make a clear distinction between buffer range and buffer capacity.

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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction

of Chemical Reactions

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Sample Questions

Q1) Hydrogen sulfide decomposes according to the following reaction 2H₂S(g) \(\to\) 2H₂(g) + S₂(g)

For this reaction at 298K \(\Delta\)S°= 78.1 J/K, \(\Delta\)H°= 169.4 kJ, and \(\Delta\)G° = 146.1 kJ. What is the value of \(\Delta\)G° at 900 K?

A) -69881 kJ

B) 48.4 kJ

C) 99.1 kJ

D) 240 kJ

E) 441 kJ

Q2) A chemical reaction has \(\Delta\)G° = 10.0 kJ and \(\Delta\)S° = 50.0 J/K

a. Calculate \(\Delta\)H°for this reaction at 25°C.

b. Could this reaction ever be spontaneous? Explain your answer.

Q3) For the reaction of xenon and fluorine gases to form solid XeF₄, \(\Delta\)H° = -251 kJ and \(\Delta\)G° = -121 kJ at 25°C. Calculate \(\Delta\)S° for the reaction.

Q4) The free energy of a perfect crystal at absolute zero, is zero.

A)True

B)False

Q5) State the second and third laws of thermodynamics.

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Sample Questions

Q1) In the shorthand notation for cells, a double vertical line is used to separate the reduced and oxidized forms of a redox couple.

A)True

B)False

Q2) A concentration cell consists of two Zn/Zn²⁺ electrodes. The electrolyte in compartment A is 0.10 M Zn(NO₃)₂and in compartment B is 0.60 M Zn(NO₃)₂. What is the voltage of the cell at 25°C?

A) 0.010 V

B) 0.020 V

C) 0.023 V

D) 0.046 V

E) none of the above

Q3) Which of the following solids is commonly used as an inactive electrode in electrochemical cells?

A) zinc

B) graphite

C) copper

D) iron

E) sodium

Page 24

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Chapter 22: The Elements in Nature and Industry

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Q1) In the industrial electrolysis of aqueous NaCl (the chlor-alkali process), the modern trend is toward the use of cells incorporating polymer membranes to separate the anode and cathode solutions.

A)True

B)False

Q2) Sulfur trioxide is the anhydride of sulfuric acid. However, SO₃is not added directly to water during the synthesis of sulfuric acid because A) hydration of SO₃is very exothermic and difficult to control.

B) hydration of SO₃is very endothermic and proceeds too slowly to be profitable.

C) at high temperatures water vapor catalyzes polymerization of SO₃.

D) vapor phase sulfuric acid is corrosive and limits the useful life of the reactor.

E) the water would decompose under such conditions.

Q3) Briefly describe the three main pathways for nitrogen fixation.

Q4) Give two important reasons why there is so much more Na⁺ in the oceans than K⁺, despite their similar abundances in the earth's crust.

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Page 25

Chapter 23: The Transition Elements and Their Coordination Compounds

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Sample Questions

Q1) What is the coordination number of cobalt in the complex ion [Co(en)Cl₄]^-? (en = ethylenediamine)

A) 1

B) 2

C) 4

D) 6

E) 8

Q2) The Cu²⁺ ion has 1 unpaired electron.

A)True

B)False

Q3) Which of the following ions could exist in only the high-spin state in an octahedral complex?

A) Cr²⁺

B) Mn⁴⁺

C) Fe³⁺

D) Co³⁺

E) Ni²⁺

Q4) All the actinide series of transition elements are radioactive.

A)True

B)False

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Chapter 24: Nuclear Reactions and Their Applications

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Sample Questions

Q1) When an electron and its anti-particle, a positron, collide, they annihilate each other. Calculate the energy released in this process, in J. (The positron mass is the same as the electron mass, namely 9.11 × 10^-31 kg.)

Q2) Write a complete, balanced equation to represent the electron capture decay of argon-37.

Q3) An 85-kg person exposed to barium-141 receives 2.5 × 10⁵ \(\beta\) particles, each with an energy of 5.2 × 10^-13 J. How many rads does the person receive?

A) 2.4 × 10^-20

B) 1.5 × 10^-7

C) 1.8 × 10^-16

D) 6.1 × 10^-15

E) none of the above

Q4) What is the mechanism by which control rods slow down the fission rate in a nuclear reactor?

Q5) Explain how the number of protons and neutrons in a radioactive nucleus can be used to predict its probable mode of decay. Illustrate your answer with a schematic graph, properly labeled, showing stable nuclides (nuclei) in relation to number of protons and neutrons.

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