Chemistry for Engineers Exam Review - 3452 Verified Questions

Chemistry for Engineers Exam Review

Course Introduction

Chemistry for Engineers is designed to provide engineering students with a foundational understanding of chemical principles relevant to their field. The course covers topics such as atomic and molecular structure, chemical bonding, reaction stoichiometry, thermodynamics, kinetics, and electrochemistry, emphasizing their applications in engineering processes and materials. Students will explore how chemical properties influence the design and functionality of engineering systems, materials, and devices, preparing them to address real-world challenges that require interdisciplinary problem-solving skills. Laboratory components may be included to reinforce theoretical concepts through hands-on experiments.

Recommended Textbook

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

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25 Chapters

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Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) 5.621 lb = ________ grams.(1 lb = 454 g)

A) 2550 × 10³

B) 1.24 × 10^-2

C) 1.28 × 10³

D) 8.00 × 10³

E) 1.70 × 10³

Answer: A

Q2) Describe the difference between an intensive and extensive property using examples.

Answer: An intensive property does NOT depend on the amount of the substance present,such as color or density.An extensive property is one that does depend on the amount of the substance,such as mass or volume.

Q3) How many significant figures are in the measurement 5.3 g?

A) 1

B) 4

C) 3

D) 2

E) 5

Answer: D

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Give an example of an halogen.

Answer: F,Br,I,or Cl

Q2) How many moles are in 2.16 × 10²⁴ atoms of magnesium?

A) 35.9 moles

B) 3.59 moles

C) 0.359 moles

D) 7.18 moles

E) 0.90 moles

Answer: B

Q3) The mass number of an atom of ⁷⁹Br is ________.

A) 35

B) 114

C) 44

D) 79

E) 72

Answer: D

Q4) Why do elements in the same group tend to have similar chemical properties?

Answer: Since elements in the same group tend to react similarly,they form the ions with the same charge.

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) A sample of pure lithium nitrate contains 7.99% lithium by mass.What is the % lithium by mass in a sample of pure lithium nitrate that has twice the mass of the first sample?

A) 4.00%

B) 7.99%

C) 16.0%

D) 32.0%

Answer: B

Q2) Balance the following equation.

_____ C₉H₂₀ + _____ O₂ _____

H₂O + _____ CO₂

Answer: C₉H₂₀ + 14 O₂ 10 H₂O + 9 CO₂

Q3) What is the empirical formula for Hg₂(NO₃)₂?

A) Hg₂(NO₃)₂

B) HgNO₃

C) Hg(NO₃)₂

D) Hg₂NO₃

E) Hg₄(NO₃)₄

Answer: B

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) What is the concentration of barium ions in a 0.750 M BaCl₂ solution?

A) 0.750 M

B) 0.375 M

C) 2.25 M

D) 1.50 M

E) 0.160 M

Q2) What volume (in mL)of 0.0887 M MgCl₂ solution is needed to make 275.0 mL of 0.0224 M MgCl₂solution?

A) 72.3 mL

B) 91.8 mL

C) 10.9 mL

D) 69.4 mL

E) 14.4 mL

Q3) How many moles of LiI are contained in 258.6 mL of 0.0296 M LiI solution?

A) 1.31 × 10^-3 mol

B) 8.74 × 10^-3 mol

C) 1.14 × 10^-3 mol

D) 3.67 × 10^-3 mol

E) 7.65 × 10^-3 mol

Q4) Define an electrolyte.

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Chapter 5: Gases

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Sample Questions

Q1) Rank the following in order of decreasing rate of effusion. F₂

SF₆ Ne Kr

A) SF₆> Kr > F₂ > Ne

B) Kr > Ne > SF₆ > F₂

C) F₂ > Kr > Ne > SF₆

D) Ne > F₂ > Kr > SF₆

E) Ne > F₂ > SF₆ > Kr

Q2) The volume of a gas is inversely proportional to the pressure of a gas is known as

A) Avogadro's Law.

B) Ideal Gas Law.

C) Charles's Law.

D) Boyle's Law.

E) Dalton's Law.

Q3) Identify the element or molecule that has the lowest diffusion rate.

A) Ne

B) CO

C) HCl

D) SO₂

E) Xe

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Which of the following signs on q and w represent a system that is doing work on the surroundings,as well as losing heat to the surroundings?

A) q = -, w = -

B) q = +, w = +

C) q = -, w = +

D) q = +, w = -

E) None of these represent the system referenced above.

Q2) A 35.6 g sample of ethanol (C₂H₅OH)is burned in a bomb calorimeter,according to the following reaction.If the temperature rose from 35.0 to 76.0°C and the heat capacity of the calorimeter is 23.3 kJ/°C,what is the value of DH°_rxn? The molar mass of ethanol is 46.07 g/mol.

C₂H₅OH(l)+ 3 O₂(g) 2 CO₂(g)+ 3

H₂O(g) H°_rxn= ?

A) -1.24 × 10³ kJ/mol

B) +1.24 × 10³ kJ/mol

C) -8.09 × 10³ kJ/mol

D) -9.55 × 10³ kJ/mol

E) +9.55 × 10³ kJ/mol

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Page 8

Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Place the following types of electromagnetic radiation in order of increasing frequency. x-rays radio waves gamma rays

A) radio waves < x-rays < gamma rays

B) gamma rays < x-rays < radio waves

C) radio waves < gamma rays < x-rays

D) gamma rays < radio waves < x-rays

E) x-rays < gamma rays < radio waves

Q2) What is the maximum number of p orbitals that are possible?

A) 0

B) 3

C) 6

D) 12

E) 9

Q3) Identify the color that has a wavelength of 480 nm.

A) blue

B) green

C) red

D) yellow

E) violet

Q4) Define diffraction.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Which of the following have their valence electrons in the same shell?

A) Rb, Sb, I

B) Al, Ge, Sb

C) P, As, Sb

D) He, Ne, F

E) Zn, Cd, Hg

Q2) How many unpaired electrons are present in the ground state Si atom?

A) 5

B) 8

C) 3

D) 2

E) 4

Q3) How many electrons are in the outermost shell of the Al³⁺ ion in its ground state?

A) 2

B) 3

C) 6

D) 8

Q4) Define ionization energy.

Q5) 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 fluorine atom in F₂ would be expected to have a A) charge of 1-.

B) partial charge -.

C) partial charge +. D) charge of 0.

Q2) Use the bond energies provided to estimate H°_rxn for the reaction below.

C₂H₄(g)+ H₂(g) C₂H₆(g) H°_rxn= ?

Bond Bond Energy (kJ/mol)

C-C 347

C-H 414

C=C 611

C C 837

H-H 436

A) -128 kJ

B) +98 kJ

C) +700 kJ

D) -102 kJ

E) -166 kJ

Q3) List the most electronegative atom.

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) Give the approximate bond angle for a molecule with a linear shape.

A) 109.5°

B) 180°

C) 120°

D) 105°

E) 60°

Q2) If an oil tanker leaks oil in the ocean,where does the oil go?

A) The oil floats on the seawater; water is polar and oil is nonpolar.

B) The oil mixes with the water.

C) The oil sinks below the seawater; water is nonpolar and oil is polar.

D) The oil floats on the seawater; water is nonpolar and oil is polar.

E) The oil sinks below the seawater; water is polar and oil is nonpolar.

Q3) Give the hybridization for the O in H₃O⁺.

A) sp

B) sp³

C) sp²

D) sp³d

E) sp³d²

Q4) Explain why oil and water do not mix.

Q5) According to molecular orbital theory,what is an antibonding orbital?

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) Which of the following compounds has the highest boiling point?

A) H₂O

B) H₂S

C) H₂Se

D) AsH₃

Q2) Explain why ice floats in water.

A) Ice is more dense than water.

B) Steam is less dense than water.

C) Ice is less dense than water.

D) Steam is more dense than water.

E) Ice is the same density as water.

Q3) Place the following substances in order of increasing boiling point. He I₂O₂

A) He < I₂ < O₂

B) I₂ < O₂ < He

C) O₂ < I₂ < He

D) I₂ < He < O₂

E) He < O₂ < I₂

Q4) Define volatile.

Q5) Define dynamic equilibrium.

Page 13

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Rubidium crystallizes in a body-centered cubic structure.What is the coordination number of each atom?

A) 2

B) 6

C) 8

D) 12

E) 16

Q2) Identify the type of solid for gold.

A) metallic atomic solid

B) ionic solid

C) nonbonding atomic solid

D) molecular solid

E) networking atomic solid

Q3) Identify the type of interaction between atoms in a networking atomic solid.

A) polar bonding

B) ionic bonding

C) covalent bonding

D) hydrogen bonding

E) weak dispersion forces

Q4) 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 solution is prepared by dissolving 96.4 g of RbI in enough water to form 865 mL of solution.Calculate the mass % of the solution if the density of the solution is 1.06 g/mL.

A) 11.3%

B) 12.7%

C) 9.4%

D) 10.5%

E) 11.9%

Q2) What mass of LiF is contained in 18.66 g of a 20.0% by mass solution of LiF in water?

A) 20.00 g

B) 21.15 g

C) 3.73 g

D) 1.50 g

E) 12.44 g

Q3) Give the preparation of rock candy.

Q4) What happens to a supersaturated solution of potassium acetate once it is cooled and a small crystal of solid potassium acetate is added?

Q5) Explain why water does not dissolve in gasoline.

Q6) Define Tyndall effect.

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The decomposition of dinitrogen pentoxide is described by the following chemical equation 2 N₂O₅(g) 4 NO₂(g)+ O₂(g)

If the rate of appearance of O₂ is equal to 3.00 mol/min at a particular moment,what is the rate of disappearance of N₂O₅ at that moment?

A) 0.750 mol/min

B) 1.50 mol/min

C) 6.00 mol/min

D) 12.0 mol/min

Q2) In the hydrogenation of double bonds,a catalyst is needed.In the first step,the reactants must come into contact with a metal surface.This step is known as A) adsorption.

B) diffusion.

C) reaction.

D) desorption.

E) none of the above

Q3) Define half-life.

Q4) What function do enzymes serve?

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Determine the value of K_p for the following reaction if the equilibrium concentrations are as follows: P(CO)_eq = 6.8 × 10^-11 atm,P(O₂)_eq = 1.3 × 10^-3 atm,P(CO₂)_eq = 0.041 atm. 2 CO(g)+ O₂(g) 2 CO₂(g)

A) 3.6 × 10^-21

B) 2.8 × 10²⁰

C) 4.6 × 10¹¹

D) 2.2 × 10^-12

E) 3.6 × 10^-15

Q2) Explain dynamic equilibrium.Use the generic reaction A(g) B(g)to explain.

Q3) Why is an equilibrium constant unitless?

Q4) Consider the following reaction at equilibrium.What will happen if the pressure increased? 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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Q1) Calculate the concentration of bicarbonate ion,HCO₃^-,in a 0.010 M H₂CO₃ solution that has the stepwise dissociation constants K_a1 = 4.3 × 10^-7 and K_a2 = 5.6 × 10^-11.

A) 6.6 × 10^-⁵ M

B) 4.3 × 10^-7 M

C) 4.3 × 10^-⁹ M

D) 5.6 × 10^-11 M

Q2) Determine the pH of a 0.023 M HNO₃ solution.

A) 12.36

B) 3.68

C) 1.64

D) 2.30

E) 2.49

Q3) What is the K_w of pure water at 50.0°C,if the pH is 6.630?

A) 2.34 × 10^-7

B) 5.50 × 10^-14

C) 2.13 × 10^-14

D) 1.00 × 10^-14

E) There is not enough information to calculate the K_w.

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Chapter 17: Aqueous Ionic Equilibrium

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Q1) A 1.00 L buffer solution is 0.250 M in HF and 0.250 M in NaF.Calculate the pH of the solution after the addition of 100.0 mL of 1.00 M HCl.The K_a for HF is 3.5 × 10^-4.

A) 3.09

B) 4.11

C) 3.82

D) 3.46

E) 2.78

Q2) Identify the compound that is acid-insoluble.

A) PbCl₂

B) As₂S₃

C) CoS

D) Ca₃(PO₄)₂

E) NaCl

Q3) Identify the indicator that can be used at the highest pH.

A) alizarin

B) thymol blue

C) crystal violet

D) phenolphthalein

E) alizarin yellow R

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) What is "free" energy? Give a fictitious example.

Q2) Use Hess's law to calculate G°_rxn using the following information. NO(g)+ O(g) NO₂(g) G°_rxn= ?

2 O₃(g) 3 O₂(g) G°_rxn= +489.6 kJ

O₂(g) 2 O(g) G°_rxn= +463.4 kJ NO(g)+ O₃(g) NO₂(g)+ O₂(g) G°_rxn=

-199.5 kJ

A) +753.5 kJ

B) +277.0 kJ

C) -676.0 kJ

D) -1152.5 kJ

E) -225.7 kJ

Q3) For a given reaction, H = -19.9 kJ/mol and S = -55.5 J/Kmol.The reaction will have G = 0 at ________ K.Assume that H and S do not vary with temperature.

A) 359

B) 2789

C) 298

D) 2.79

E) 0.359

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Chapter 19: Electrochemistry

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Sample Questions

Q1) Identify the battery type that has a high overcharge tolerance.

A) NiCad

B) Li ion

C) NiMH

D) lead storage

E) none of the above

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 are iron nails coated with zinc?

Q4) Why is sugar water NOT a good conductor of current?

Q5) What is electrolysis?

Q6) Why,if we multiply a reaction by 2,don't we multiply its E°_red by 2?

Page 21

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Which particle changes the mass of the isotope the least?

A) alpha emission

B) proton capture

C) gamma particle

D) positron capture

E) electron capture

Q2) Stable isotopes,with low atomic numbers,have a N/Z ratio of 1.What does that imply?

A) The number of neutrons equals the number of protons.

B) The number of neutrons equals the number of electrons plus protons.

C) The number of electrons equals the number of protons plus neutrons.

D) The atomic number equals twice the atomic mass.

E) The number of protons equals the number of electrons plus neutrons.

Q3) Identify the part of the body that can be studied with radiotracers.

A) spleen

B) brain

C) heart

D) tumors

E) all of the above

Q4) How does a dosimeter measure exposure to radioactivity?

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Chapter 21: Organic Chemistry

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Q1) Nonsuperimposable mirror images are called

A) structural isomers.

B) achiral.

C) diastereomers.

D) enantiomers.

E) racemic mixture.

Q2) Write the balanced chemical equation for the addition of 2 Br₂ to CH₃C CH.

A) CH₃C CH + 2

Br₂ CH₃CBr₂CHBr₂

B) CH₃C CH + 2 Br₂ CH₃CBr=CHBr

C) CH₃C CH + 2

Br₂ CH₃CH₂CHBr₂

D) CH₃C CH + 2

Br₂ CH₃CBr₂CH₃

E) CH₃C CH + 2 Br₂ CH₃CHBrCH₂Br

Q3) Why are there so many more carbon compounds than the number of compounds made up of all the rest of the elements combined?

Q4) What does the term "racemic mixture" mean?

Q5) Why doesn't benzene typically undergo addition reactions like the alkenes do?

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Chapter 22: Biochemistry

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Q1) Which of the following statements is TRUE?

A) Monosaccharides are the building blocks of amino acids.

B) Nucleotides are the building blocks of polysaccharides.

C) Fatty acids are responsible for DNA replication.

D) Polysaccharides are the building blocks of amino acids.

E) Carbohydrates are water soluble and responsible for short term energy storage.

Q2) Identify the polysaccharide.

A) fructose

B) maltose

C) cholesterol

D) starch

E) estradiol

Q3) How do cis- and trans-fats differ?

Q4) Identify a base that is NOT in DNA.

A) adenine

B) cytosine

C) thymine

D) uracil

E) guanine

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Chapter 23: Chemistry of the Nonmetals

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Q1) Identify a characteristic that natural oxygen does NOT have.

A) odorless

B) colorless

C) pale green liquid

D) diatomic

E) nonpolar

Q2) List the name of powdered form of carbon,which is a component of soot.

A) diamond

B) graphite

C) fullerenes

D) buckyballs

E) carbon black

Q3) Why are phosphate compounds added to detergents?

Q4) Determine the oxidation state of carbon in the ionic carbide,CaC₂.

A) -2

B) +4

C) -1

D) +1

E) -4

Q5) Describe the Frasch process for obtaining sulfur.

Page 25

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Chapter 24: Metals and Metallurgy

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Q1) An interstitial alloy contains carbon in one-fourth of the octahedral holes of a closest packed metal,M.What is the formula of this alloy?

A) M₄C

B) MC₄

C) MC

D) M₂C

E) MC₂

Q2) Identify the metal used as a coating for steel objects such as nails.

A) gold

B) silver

C) zinc

D) palladium

E) manganese

Q3) Why is zinc used to coat steel objects?

Q4) Pyrometallurgy is the

A) refining of metal ores using reactions with aqueous solutions.

B) refining of metal ores using heat.

C) refining of metal ores using reduction reactions.

D) forming metal parts using heat and small crystals of metal.

E) forming metal parts using electricity.

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Which of the following compounds can exhibit fac-mer isomerism?

A) [Cr(H₂O)₄Br₂]⁺

B) [Fe(CO)₅ONO]²⁺

C) [Fe(CO)₃(NH₃)₃]³⁺

D) [Cu(CO)₅Br]⁺

E) [V(NH₃)₄(H₂O)₂]²⁺

Q2) Identify the geometry of the complex ion if the hybridization is d²sp³.

A) square planar

B) octahedral

C) tetrahedral

D) linear

E) trigonal planar

Q3) Identify the geometry of the complex ion if the hybridization is dsp².

A) square planar

B) octahedral

C) tetrahedral

D) linear

E) trigonal planar

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