Chemistry for Engineers Test Questions - 3452 Verified Questions

Course Introduction

Chemistry for Engineers Test Questions

Chemistry for Engineers introduces the foundational concepts of chemistry with an emphasis on applications relevant to engineering disciplines. This course covers atomic structure, chemical bonding, stoichiometry, thermodynamics, kinetics, and the properties of gases, liquids, and solids. It also explores electrochemistry, materials science, and the role of chemistry in environmental and sustainable engineering practices. Through lectures, laboratory experiments, and problem-solving sessions, students gain a practical understanding of chemical principles and their significance in designing and improving engineering processes and products.

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

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25 Chapters

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) The width,length,and height of a large,custom-made shipping crate are 1.32 m,2.50 m,and 0.83 m,respectively.The volume of the box using the correct number of significant figures is ________ m³.

A) 2.73900

B) 2.7

C) 2.74

D) 2.739

E) 2.7390

Answer: B

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) A flash drive contains 4 gigabytes.How many bytes does it contain?

Answer: 4,000,000,000 bytes or 4,292,967,296 bytes,if someone is computer literate

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Sodium belongs to the ________ group of the periodic table.

A) alkali metal

B) transition element

C) noble gas

D) nonmetal

Answer: A

Q2) Calculate the atomic mass of gallium if gallium has 2 naturally occurring isotopes with the following masses and natural abundances: Ga-69 68.9256 amu 60.11%

Ga-71 70.9247 amu 39.89%

A) 69.72 amu

B) 69.93 amu

C) 70.00 amu

D) 69.80 amu

E) 70.68 amu

Answer: A

Q3) Describe the difference between ions and isotopes.

Answer: Ions have the loss or gain of electrons; isotopes differ in the number of neutrons.

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) How many moles are there in 1.50 g of ethanol,CH₃CH₂OH?

A) 0.0145 mol

B) 0.0326 mol

C) 30.7 mol

D) 69.0 mol

Answer: B

Q2) Why aren't prefixes used in naming ionic compounds?

Answer: The charges on the ions dictate how many must be present to form a neutral unit.Molecular compounds do not have such constraints and therefore must use prefixes to denote the number of atoms present.

Q3) How much calcium is contained in 35.0 g of calcium fluoride?

A) 1.99 g

B) 7.70 g

C) 15.0 g

D) 18.0 g

Answer: D

Q4) Describe the difference between an atomic element and a molecular element.

Answer: Atomic elements exist with a single atom at their basic unit; molecular elements have more than one atom in their basic unit.

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) How many grams of H₃PO₄ are in 255 mL of a 4.50 M solution of H₃PO₄?

A) 112 g

B) 1.15 g

C) 441 g

D) 4.50 g

E) 400 g

Q2) How many milliliters of 0.111 M HClO₄ solution are needed to neutralize 50.00 mL of 0.0789 M NaOH?

A) 0.438

B) 0.0142

C) 70.3

D) 0.0281

E) 35.5

Q3) Which one of the following compounds behaves as an acid when dissolved in water?

A) SrO

B) C₃H₈

C) HBr

D) RbOH

Q4) Determine the oxidation state of Mn in KMnO₄.

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Chapter 5: Gases

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Sample Questions

Q1) How many grams of KCl(s)are produced from the thermal decomposition of KClO₃(s)which produces 100.0 mL of O₂(g)at 25°C and 1.00 atm pressure according to the chemical equation shown below? 2 KClO₃(s) 2 KCl(s)+ 3 O₂(g)?

A) 0.203 g

B) 0.305 g

C) 0.334 g

D) 0.608 g

Q2) A 0.529-g sample of gas occupies 125 mL at 60.cm of Hg and 25°C.What is the molar mass of the gas?

A) 11 g/mol

B) 81 g/mol

C) 110 g/mol

D) 130 g/mol

Q3) Calculate the ratio of effusion rates of Cl₂ to F₂.

A) 51.9

B) 0.535

C) 1.87

D) 1.37

E) 0.732

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Which of the following is NOT a standard state?

A) for a solid, it is 25°F

B) for a gas, it is 25°C

C) for a gas, it is 760 torr

D) for a solution, it is 1 M

E) for a liquid, it is 760 torr

Q2) Give the units of heat capacity.

A) J/°C

B) Jg/ °F

C) J/mole°C

D) g/°F

E) mole/°F

Q3) Which of the following is NOT a major contributor to energy consumption?

A) residential

B) commercial

C) transportation

D) industrial

E) atmospheric

Q4) Define chemical energy.

Q5) Explain the difference between H and DE.

Page 8

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) How many orbitals are contained in the third principal level (n = 3)of a given atom?

A) 9

B) 1

C) 12

D) 6

E) 10

Q2) Calculate the frequency of the red light emitted by a neon sign with a wavelength of 680 nm.

A) 2.27 × 10¹⁴s^-1

B) 3.80 × 10¹⁴s^-1

C) 4.41 × 10¹⁴s^-1

D) 5.05 × 10¹⁴s^-1

E) 3.40 × 10¹⁴s^-1

Q3) How many different values of m_l are possible in the 4d sublevel?

A) 2

B) 0

C) 3

D) 5

E) 6

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Page 9

Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) For a particular element,identify the species that has the smallest radius.

A) cation

B) anion

C) neutral

D) radical

E) They are all the same size.

Q2) How many of the following species are paramagnetic? Sc³ Cl

Ca² Se

A) 0

B) 2

C) 1

D) 4

E) 3

Q3) Give the number of valence electrons for Cl.

A) 8

B) 9

C) 2

D) 3

E) 7

Q4) Define paramagnetic.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Using Lewis structures and formal charge,which of the following ions is most stable,?

OCN ONC NOC

A) OCN

B) ONC

C) NOC

D) None of these ions are stable according to Lewis theory.

E) All of these compounds are equally stable according to Lewis theory.

Q2) List the least electronegative atom.

Q3) Identify an ionic bond.

A) Electrons are pooled.

B) Electrons are shared.

C) Electrons are transferred.

D) Neutrons are gained.

E) Electrons are lost.

Q4) Of the following elements,which has the highest electronegativity?

A) Br

B) C

C) Nb

D) Sb

E) Cs

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) List the number of sigma bonds and pi bonds in a double bond.

A) 1 sigma, 1 pi

B) 2 sigma, 1 pi

C) 2 sigma, 2 pi

D) 1 sigma, 2 pi

Q2) Determine the electron geometry (eg),molecular geometry (mg),and polarity of N₂O (N central).

A) eg = linear, mg = linear, nonpolar

B) eg = tetrahedral, mg = linear, polar

C) eg = tetrahedral, mg = bent, nonpolar

D) eg = linear, mg = linear, polar

E) eg = trigonal planar, mg = trigonal pyramidal, polar

Q3) The bond angle in H₂Se is

A) 107°.

B) <104.5°.

C) 109.5°.

D) 190°.

E) 145°.

Q4) Is it possible for a molecule to be nonpolar even though it contains polar bonds? Explain your answer and give an example.

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) Which of the following statements is TRUE?

A) Intermolecular forces are generally stronger than bonding forces.

B) The potential energy of molecules decreases as they get closer to one another.

C) Energy is given off when the attraction between two molecules is broken.

D) Increasing the pressure on a solid usually causes it to become a liquid.

E) None of the above is true.

Q2) Choose the substance with the lowest surface tension.

A) CH₃SeCH₂CH₃

B) CH₃CH₂CH₂CH₃

C) C₆H₆

D) H₂Se

E) (CH₃)₂SO

Q3) Which one of the following has a low density?

A) gas

B) liquid

C) solid

D) none of the above

E) all of the above

Q4) Why is the H_vap higher than H_fusfor a given compound?

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Identify the element with the smallest band gap.

A) lead

B) silicon

C) germanium

D) carbon

E) tin

Q2) A common use of polyurethane is

A) tire cord.

B) fishing line.

C) foam cups.

D) spray-on insulation.

E) films.

Q3) Gold crystallizes in a face-centered cubic structure.What is the edge length of the unit cell if the atomic radius of gold is 144 pm?

A) 204 pm

B) 288 pm

C) 333 pm

D) 407 pm

Q4) Describe the difference between the conduction band and the valence band.

Q5) Give the edge length in terms of r for a simple cubic cell.

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Chapter 13: Solutions

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Sample Questions

Q1) A 150.0 mL sample of an aqueous solution at 25°C contains 15.2 mg of an unknown nonelectrolyte compound.If the solution has an osmotic pressure of 8.44 torr,what is the molar mass of the unknown compound?

A) 223 g/mol

B) 294 g/mol

C) 341 g/mol

D) 448 g/mol

E) 195 g/mol

Q2) Give the preparation of rock candy.

Q3) The boiling point of an aqueous 1.83 m (NH₄)₂SO₄ (molar mass = 132.15 g/mol)solution is 102.5°C.Determine the value of the van't Hoff factor for this solute if the K_b for water is 0.512°C/m.

A) 3.0

B) 3.6

C) 1.8

D) 2.7

E) 2.3

Q4) How does a solution of two volatile components with strong solute-solvent attractions deviate from Raoult's law? Why?

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) How many half-lives are required for the concentration of reactant to decrease to 1.56% of its original value?

A) 6

B) 4

C) 24

D) 6.5

E) 7.5

Q2) Which rate law has a molecularity of one?

A) Rate = k[A]⁰

B) Rate = k[A][B]

C) Rate = k[A][B][C]

D) Rate = k[A]

E) Rate = k[A]²

Q3) What is a catalyst and what function does it serve?

Q4) Is the activation energy for a forward reaction the same as the activation energy for the reverse of the same reaction? Use a sketch of a reaction energy diagram.

Q5) Explain how the order of a reaction can be determined.

Q6) What is the difference between average reaction rate and instantaneous reaction rate?

Page 16

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Why aren't solids or liquids included in an equilibrium expression?

Q2) Consider the following reaction: 2 H₂O(g)+ 2 SO₂(g) 2 H₂S(g)+ 3 O₂(g)

A reaction mixture initially contains 2.8 M H₂O and 2.6 M SO₂.Determine the equilibrium concentration of H₂S if K_c for the reaction at this temperature is 1.3 × 10^-6.

A) 0.045 M

B) 0.058 M

C) 0.028 M

D) 3.1 × 10^-3 M

E) 0.12 M

Q3) Consider the following reaction at equilibrium.What will happen if FeS₂ is added to the reaction? 4 FeS₂(s)+ 11 O₂(g) 2 Fe₂O₃(s)+ 8 SO₂(g)

A) The equilibrium constant will increase.

B) The equilibrium will change in the direction of the reactants.

C) The equilibrium will change in the direction of the products.

D) No change in equilibrium is observed.

E) The equilibrium constant will decrease.

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Chapter 16: Acids and Bases

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Sample Questions

Q1) A solution with a hydroxide ion concentration of 4.15 × 10^-6 M is ________ and has a hydrogen ion concentration of ________.

A) acidic, 2.41 × 10^-8 M

B) acidic, 2.41 × 10^-9 M

C) basic, 2.41 × 10^-8 M

D) basic, 2.41 × 10^-9 M

Q2) What does the term amphoteric mean?

Q3) Determine the pOH of a 0.348 M Ba(OH)₂ solution at 25°C.

A) 13.08

B) 13.54

C) 13.84

D) 0.157

E) 0.458

Q4) Calculate the pH of a 0.800 M NaCH₃CO₂ solution.K_a for acetic acid,CH₃CO₂H,is 1.8 × 10^-⁵.

A) 2.42

B) 4.68

C) 9.32

D) 11.58

Page 18

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Chapter 17: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Define a buffer.

Q2) A solution containing CaCl₂ is mixed with a solution of Li₂C₂O₄ to form a solution that is 2.1 × 10^-5 M in calcium ion and 4.75 × 10^-5M in oxalate ion.What will happen once these solutions are mixed? K_sp (CaC₂O₄)= 2.3 × 10^-9.

A) A precipitate will form since Q > K_sp for calcium oxalate.

B) Nothing will happen since both calcium chloride and lithium oxalate are soluble compounds.

C) Nothing will happen since calcium oxalate is extremely soluble.

D) Nothing will happen since K_sp > Q for all possible precipitants.

E) There is not enough information to determine.

Q3) Calculate the pH of a solution formed by mixing 120.0 mL of 0.30 M HClO with 1180.0 mL of 0.20 M KClO.The K_a for HClO is 2.9 × 10^-8.

A) 5.99

B) 8.01

C) 7.54

D) 7.06

E) 6.46

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) Which one of the following would be expected to have the lowest standard molar entropy,S°,at 25°C?

A) C₈H₁₈(s)

B) C₈H₁₈(l)

C) C₁₄H₃₀(g)

D) C₁₄H₃₀OH(l)

Q2) Place the following in order of increasing molar entropy at 298 K. NO CO SO

A) NO < CO < SO

B) SO < CO < NO

C) SO < NO < CO

D) CO < SO < NO

E) CO < NO < SO

Q3) What is "free" energy? Give a fictitious example.

Q4) Define allotrope.

Q5) Identify the compound with the most negative standard free energy of formation.

A) NaF(s)

B) N₂(g)

C) NO(g)

D) O₃(g)

E) It is hard to determine.

Page 20

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Chapter 19: Electrochemistry

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Sample Questions

Q1) Which of the following is the strongest oxidizing agent?

A) ClO₂(s)

B) I (aq)

C) Li (aq)

D) SO₄^2-(aq)

E) MnO₄ (aq)

Q2) Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25°C.(The equation is balanced.) Sn(s)+ 2 Ag (aq) Sn²⁺(aq)+ 2 Ag(s)

Sn²⁺(aq)+ 2 e Sn(s)E° = -0.14 V Ag (aq)+ e Ag(s)E° = +0.80 V

A) +1.74 V

B) +0.94 V

C) +1.08 V

D) -1.08 V

E) -1.74 V

Q3) Why,if we multiply a reaction by 2,don't we multiply its E°_red by 2?

Q4) Give an example of an inert electrode.

Q5) What is electrolysis?

Q6) 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 discovered by Marie Curie.

A) polonium and radium

B) radium and cesium

C) argon and xenon

D) radon and xenon

E) selenium and tungsten

Q2) Which of the following statements is TRUE?

A) Positrons are similar in ionizing power and penetrating power to beta particles.

B) A positron is the antiparticle of the electron.

C) Beta decay occurs when a neutron changes into a proton while emitting an electron.

D) An alpha particle is a helium 2+ ion.

E) All of the above are true.

Q3) The splitting of a heavy nucleus to form two or more lighter ones is called

A) radioactive cleavage.

B) nuclear fission.

C) nuclear fusion.

D) radioactive merge.

E) half-life.

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22

Chapter 21: Organic Chemistry

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Sample Questions

Q1) Give the organic product for the following reaction. CH₃CH₂CH₂NH₂ + HCl

A) ClCH₂CH₂CH₂NH₂

B) CH₃CHClCH₂NH₂

C) CH₃CH₂CHClNH₂

D) CH₃CH₂CH₂NHCl

E) CH₃CH₂CH₂NH₃⁺Cl^-

Q2) Identify the alkane with the highest boiling point.

A) pentane

B) butane

C) ethane

D) methane

E) propane

Q3) Identify the formula for an alkyne.

A) C_nH_2n+2

B) C_nH_2n-2

C) C_nH_2n

D) C_nH_2n-3

E) C_nH_2n+3

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Chapter 22: Biochemistry

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Sample Questions

Q1) Identify the source below that does NOT contain triglycerides.

A) corn oil

B) olive oil

C) whale oil

D) beef fat

E) diesel oil

Q2) Identify the interaction that does NOT maintain tertiary structure.

A) disulfide bond

B) salt bridge

C) carbon-carbon bond

D) hydrophobic interaction

E) hydrogen bond

Q3) Which one of the following amino acids contains an acidic side chain?

A) glutamic acid

B) asparagine

C) cysteine

D) arginine

E) valine

Q4) Why are lipids well-suited as structural components of cell membranes?

Q5) What is a codon?

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Chapter 23: Chemistry of the Nonmetals

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Q1) Determine the oxidation state of nitrogen in KNO₃.

A) +1

B) +6

C) -2

D) +4

E) +5

Q2) Which of the following statements is TRUE?

A) Boron oxide, when added to silicon oxide glass, forms a more heat stable glass called Pyrex.

B) Boron most commonly forms compounds that fulfill the octet rule.

C) Boron forms fairly weak bonds with oxygen, due to their large electronegativity difference.

D) Boron trichloride is an example of a strong Lewis base.

E) None of the above is true.

Q3) Which of the following are major components of Earth's crust?

A) oxygen and silicon

B) calcium and silicon

C) oxygen and hydrogen

D) oxygen, magnesium and carbon

E) oxygen, hydrogen and carbon

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Chapter 24: Metals and Metallurgy

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Q1) Give the components of brass.

A) copper and tin

B) lead and tin

C) nickel and gold

D) copper and zinc

E) zinc and lead

Q2) Chromates are strong

A) oxidizers.

B) bases.

C) acids.

D) metals.

E) reducers.

Q3) Which substance is the best conductor of electricity?

A) astatine

B) diamond

C) tungsten

D) xenon

Q4) What is the difference between ferromagnetism and paramagnetism?

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

Q6) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Determine the chemical formula for the compound,hexammineiron(III)hexacyanoferrate(III).

A) [Fe₂(NH₃)₆(CN)₆]³⁺

B) [Fe(NH₃)₆][Fe(CN)₆]₂

C)

[Fe(NH₃)₆]₃[Fe(CN)₆]₂

D) [Fe(NH₃)₆][Fe(CN)₆]

E) [Fe₂(NH₃)₃(CN)₃]⁶⁺

Q2) Determine the chemical formula for the compound,tetracarbonylplatinum(IV)chloride.

A) [Pt(CO)₄]Cl₂

B) [Pt(CO)₄]Cl₄

C) [Pt(CO)₄Cl₄]^2-

D) [PtCl₄](CO)₄

E) [Pt(CO)₄Cl₄]^4-

Q3) How many d electrons are in WO₄^2-?

A) 0

B) 1

C) 2

D) 3

Q4) What is a coordinate covalent bond?

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