Chemistry for Engineers
Midterm Exam
Course Introduction
Chemistry for Engineers provides a foundational understanding of chemical principles and their application in engineering contexts. The course covers topics such as atomic structure, chemical bonding, thermodynamics, kinetics, equilibrium, and basic organic and inorganic chemistry. Emphasis is placed on how these concepts impact material properties, energy transformations, and industrial processes relevant to engineering disciplines. Through lectures, problem-solving exercises, and laboratory experiments, students develop the skills necessary to analyze and solve chemistry-related engineering problems, promote safety and sustainability, and understand the chemical behavior of materials commonly used in engineering practice.
Recommended Textbook Principles of Chemistry The Molecular Science 1st Edition by John W. Moore
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Chapter 1: The Nature of Chemistry
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Q1) Which of the following is a metal?
A) helium (He, atomic number 2)
B) nitrogen (N, atomic number 7)
C) sodium (Na, atomic number 11)
D) selenium (Se, atomic number 34)
E) carbon (C, atomic number 6)
Answer: C
Q2) Burning of hydrogen fuel is a(n) _____________ change.
Answer: chemical
Q3) Three length scales ordered from smallest to largest are:
A) microscale, nanoscale, macroscale.
B) microscale, macroscale, nanoscale.
C) macroscale, nanoscale, microscale.
D) nanoscale, microscale, macroscale.
E) none of these.
Answer: D
Q4) A mixture that is nonuniform in composition is a(n) _____________ mixture.
Answer: heterogeneous
Q5) How many atoms are in a diatomic molecule?.
Answer: two
Page 3
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Chapter 2: Atoms and Elements
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Q1) Which of the following statements concerning the atom is false?
A) Chemical reactions involve only the electrons.
B) The atom contains protons, neutrons and electrons.
C) Electrons are found in orbits about the nucleus.
D) The nucleus contains only neutrons.
E) Most of the mass of the atom is found in the nucleus.
Answer: D
Q2) Suppose the isotopic ratio of the two boron isotopes <sup>10</sup>B (10.013 amu) and <sup>11</sup>B (11.009 amu) in a sample has been altered from the ratio found in nature and now contains 58.73% <sup>10</sup>B in the sample. Determine the atomic weight of this sample of boron.
A) 10.511 amu
B) 10.013 amu
C) 11.009 amu
D) 10.424 amu
E) 10.498 amu
Answer: D
Q3) There are _____________ centiliters in a liter.
Answer: 100
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Chapter 3: Chemical Compounds
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Q1) Which of the following is not an ionic compound?
A) MgCl<sub>2</sub>
B) NaOH
C) NaF
D) FCl<sub>3</sub>
E) CaO
Answer: D
Q2) An organic compound has an empirical formula of C<sub>2</sub>H<sub>3</sub>O and had an approximate molecular weight of 86. What is its molecular formula?
A) C<sub>2</sub>H<sub>3</sub>O
B) C<sub>4</sub>H<sub>6</sub>O<sub>2</sub>
C) C<sub>6</sub>H<sub>8</sub>O<sub>3</sub>
D) C<sub>8</sub>H<sub>12</sub>O<sub>4</sub>
E) C<sub>10</sub>H<sub>14</sub>O<sub>5</sub>
Answer: B
Q3) Name the compound with the formula VCl<sub>3</sub>. Answer: vanadium(III) chloride
Q4) The molar mass of calcium hydroxide, Ca(OH)<sub>2</sub> is _____________. Answer: 74.09 g/mol
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Chapter 4: Quantities of Reactants and Products
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Q1) What is the percent yield when 0.750 mol of calcium reacts with sufficient oxygen to produce 34.4 g of CaO?
Q2) The complete reaction of 16.12 g of titanium with 23.88 g of chlorine (Cl<sub>2</sub>) produces a compound with the formula Ti<sub>x</sub>Cl<sub>y</sub>. What is the empirical formula of the compound?
A) TiCl
B) Ti<sub>2</sub>Cl
C) Ti<sub>4</sub>Cl
D) TiCl<sub>2</sub>
E) TiCl<sub>4</sub>
Q3) What is the theoretical yield (in grams of NaCl) when 0.750 mol of sodium reacts with excess chlorine to produce NaCl?
Q4) A decomposition reaction occurs when _____________ reactant(s) produces two or more products.
Q5) A reaction has a theoretical yield of 48.23 g. It produces 28.81 g of product. Calculate the percent yield.
Q6) Stoichiometric coefficients found in a balanced equation can be used to derive _____________ ratios.
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Chapter 5: Chemical Reactions
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Sample Questions
Q1) What is the oxidation number of N in NO?
A) -2
B) -1
C) 0
D) +1
E) +2
Q2) Which of the following is a reducing agent?
A) I<sub>2</sub>
B) Na
C) O<sub>2</sub>
D) F<sub>2</sub>
E) Br<sub>2</sub>
Q3) What is the oxidation number of O in Fe<sub>2</sub>O<sub>3</sub>?
A) +2
B) -2
C) -3
D) -5
E) +5
Q4) Molarity is a unit of solution concentration expressed in moles per
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Chapter 6: Energy and Chemical Reactions
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Q1) The temperature of a 21.6 g sample of a metal rises 6.04°C when 58.8 J of energy is applied to it. What is the identity of the metal?
A) silver (specific heat = 0.232 J g<sup>-1</sup> °C<sup>-1</sup>)
B) copper (specific heat = 0.385 J g<sup>-1</sup> °C<sup>-1</sup>)
C) iron (specific heat = 0.451 J g<sup>-1</sup> °C<sup>-1</sup>)
D) lead (specific heat = 0.128 J g<sup>-1</sup> °C<sup>-1</sup>)
E) aluminum (specific heat = 0.902 J g<sup>-1</sup> °C<sup>-1</sup>)
Q2) Which statement about energy is false?
A) Thermodynamics is the science of working and heating processes.
B) Working and heating are processes that transfer energy.
C) Chemical reactions are processes that transfer energy
D) The chemical energy of foods and fuels is a type of potential energy.
E) The thermal energy of nanoscale objects is a type of potential energy.
Q3) How many joules are there in one glass of milk containing 110 Calories?
A) 4.6 × 10<sup>5</sup> J
B) 460 kJ
C) 2.6 × 10<sup>4</sup> J
D) 26 J
E) 0.46 J
Q4) Enthalpy change is equal to heat transfer at constant _______________.
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Chapter 7: Electron Configurations and the Periodic Table
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Q1) Which set of quantum numbers is not allowed?
A) n = 0, l = 0, m<sub>l</sub> = 0, m<sub>s</sub> = +1/2
B) n = 1, l = 0, m<sub>l</sub> = 0, m<sub>s</sub> = +1/2
C) n = 2, l = 1, m<sub>l</sub> = 1, m<sub>s</sub> = +1/2
D) n = 3, l = 1, m<sub>l</sub> = 0, m<sub>s</sub> = +1/2
E) n = 4, l = 3, m<sub>l</sub> = 0, m<sub>s</sub> = -1/2
Q2) Xenon atoms, iodide ions, and cesium ions are:
A) paramagnetic.
B) diamagnetic.
C) ferromagnetic.
D) lanthanides.
E) isoelectronic.
Q3) Which statement regarding atomic orbitals is false?
A) Atomic orbitals are three dimensional.
B) Only one electron is allowed per orbital.
C) An electron shell consists of a collection of orbitals with the same principal quantum number.
D) An orbital may be designated with the letters s, p, d, f.
E) An orbital describes the location of the electron 90% of the time.
Q4) Each electron in an atom or ion must have a unique set of four _____________.
Page 9
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Chapter 8: Covalent Bonding
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Q1) Which of the following compound can exhibit cis-trans isomerism?
A) CH<sub>2</sub>=CH<sub>2</sub>
B) CH<sub>3</sub>CH<sub>3</sub>
C) H<sub>2</sub>C=O
D) ClHC=CHCl
E) Cl<sub>2</sub>C=CH<sub>2</sub>
Q2) Which bond is least polar?
A) C-C
B) C-N
C) N-H
D) C-F
E) C-O
Q3) Which bond is shortest?
A) carbon-oxygen single bond
B) carbon-hydrogen single bond
C) hydrogen-hydrogen single bond
D) carbon-carbon double bond
E) carbon-oxygen triple bond
Q4) In a polar covalent bond, _____________ are _____________ shared.
Q5) Hydrocarbons that contain double bonds are called _____________.
Page 10
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Chapter 9: Molecular Structure
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Q1) Which accurately describes a molecule with all polar bonds?
A) It is always a polar molecule.
B) It is never a polar molecule.
C) It may be a polar molecule.
D) It may be an ionic molecule.
E) It cannot be a nonpolar molecule.
Q2) The carbon-carbon p bond in ethylene, CH<sub>2</sub>CH<sub>2</sub>, results from the overlap of which of the following?
A) s atomic orbitals
B) p atomic orbitals
C) sp hybrid orbitals
D) sp<sup>2</sup> hybrid orbitals
E) sp<sup>3</sup> hybrid orbitals
Q3) Samples of which substance experience only London intermolecular forces?
A) HF
B) OF<sub>2</sub>
C) PF<sub>3</sub>
D) CCl<sub>4</sub>
E) CH<sub>3</sub>Cl
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Chapter 10: Gases and the Atmosphere
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Q1) Which description of the kinetic-molecular theory of gases is true?
A) The space between two gas molecules is much greater than the size of each gas molecule.
B) The forces of attraction and repulsion do not exist for gas molecules.
C) Colliding gas molecules have the same speed before and after the collision.
D) Gas molecules move randomly in all directions with the same speed.
E) Gas molecules lose energy with every collision and will eventually stop moving.
Q2) If at 29°C and 2.50 atm of pressure, a gas occupies 5.99 L, what volume (in L) will the same gas occupy at 19°C and 1.20 atm of pressure?
A) 2.04 × 10<sup>-4</sup> L
B) 6.04 × 10<sup>-1</sup> L
C) 8.18 L
D) 12.1 L
E) 20.0 L
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Chapter 11: Liquids, Solids, and Materials
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Q1) Viscosity of a liquid ____ as temperature decreases because ____.
A) cannot be predicted; viscosity is determined by more than just temperature
B) increases; the molecules have more kinetic energy at the new temperature
C) increases; the molecules have less kinetic energy at the new temperature
D) decreases; the molecules have more kinetic energy at the new temperature
E) decreases; the molecules have less kinetic energy at the new temperature
Q2) Which choice is an example of a molecular solid?
A) NaCl
B) iodine
C) iron
D) quartz
E) glass
Q3) In an ionic compound such as NaCl, which consists of a fcc lattice of chloride ions, each sodium ion has ____ chloride ions surrounding it as nearest neighbors.
A) 8
B) 6
C) 4
D) 1
E) 0
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Chapter 12: Chemical Kinetics: Rates of Reactions
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Q1) For a reaction that is zeroth-order with respect to reactant A, and nth-order overall
A) Rate = k[0]<sup>m</sup>[B]<sup>n</sup>
B) Rate = k[B]<sup>n</sup>
C) Rate = k[A]<sup>m</sup>[0]<sup>n</sup>
D) Rate = k[A]<sup>m</sup>[B]<sup>0</sup>
E) Rate = k[A]<sup>0</sup>[B]<sup>0</sup>
Q2) Which statement, comparing catalyzed and uncatalyzed versions of a reaction, is incorrect?
A) The initial and final energy states are the same.
B) The reaction mechanisms are different.
C) The reaction progress diagrams are different.
D) The balanced equation for the overall reaction is the same.
E) The activation energy is lower in the uncatalyzed reaction.
Q3) Which of the following units may be used for describing reaction rates?
A) mol L<sup>-1</sup> s<sup>-1</sup>
B) M min<sup>-1</sup>
C) M h<sup>-1</sup>
D) M s<sup>-1</sup>
E) all of these
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Chapter 13: Chemical Equilibrium
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Q1) An endothermic reaction which results in an increase in moles of gas will be most product-favored under conditions of ____ pressure and ____ temperature.
A) high; high
B) high; moderate
C) high; low
D) low; high
E) low; low
Q2) Which of the following is false?
A) a large equilibrium constant is indicative of a reactant-favored reaction
B) an equilibrium constant close to 1.00 always indicates a highly product-favored reaction
C) if K<sub>p</sub> = K<sub>c</sub> the reaction has gone to 100% completion at equilibrium
D) if K<sub>p</sub> = K<sub>c</sub> = 0.50, the reaction has gone to 50% completion at equilibrium
E) all of these
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Chapter 14: The Chemistry of Solutes and Solutions
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Q1) The osmotic pressure of a solution of Mg(NO<sub>3</sub>)<sub>2</sub> at a given temperature is ____ that of a solution of NaNO<sub>3</sub> of identical concentration at the same temperature.
A) half
B) two-thirds
C) the same as
D) one and a half times
E) twice
Q2) Thermal pollution affects fish and other marine wildlife by
A) increasing the growth of marine plants.
B) increasing the uptake of toxic metals.
C) raising the body temperature of fish to a dangerous level.
D) decreasing the solubility of oxygen in water.
E) hindering the growth of marine plants.
Q3) To convert from ____ to ____, one must multiply by ____.
A) parts per million; parts per billion; 10<sup>-3</sup>
B) parts per thousand; parts per billion; 10<sup>3</sup>
C) parts per million; weight percent; 10<sup>-4</sup>
D) parts per billion; parts per trillion; 10<sup>6</sup>
E) parts per thousand; weight percent; 10
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Chapter 15: Acids and Bases
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Q1) In terms of base strength, which base does not belong in a category with the others?
A) Ba(OH)<sub>2</sub>
B) NaOH
C) Sr(OH)<sub>2</sub>
D) LiOH
E) NH<sub>4</sub>OH
Q2) Two common polyprotic acids are
A) sulfuric acid and perchloric acid
B) sulfuric acid and nitric acid
C) sulfuric acid and phosphoric acid
D) perchloric acid and nitric acid
E) nitric acid and phosphoric acid
Q3) Which cannot act as a Bronsted-Lowry acid?
A) HPO<sub>4</sub><sup>2-</sup>
B) H<sub>3</sub>O<sup>+</sup>
C) CH<sub>3</sub>COO<sup>-</sup>
D) CH<sub>3</sub>NH<sub>3</sub><sup>+</sup>
E) HCOOH
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Chapter 16: Additional Aqueous Equilibria
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Q1) A buffer solution is 0.080 M in lactic acid (K<sub>a</sub> = 1.8 × 10<sup>-4</sup>) and 0.070 M in sodium lactate. The pH of the solution is
A) 2.86.
B) 3.68.
C) 3.80.
D) 4.18.
E) 4.62.
Q2) Isotonic saline solution is 0.154 M NaCl(aq). What is the solubility of AgCl (K<sub>sp</sub> = 1.8 × 10<sup>-10</sup>) in such a solution?
A) 2.8 × 10<sup>-11</sup> M
B) 1.2 × 10<sup>-9</sup> M
C) 5.3 × 10<sup>-6</sup> M
D) 3.4 × 10<sup>-5</sup> M
E) 0.077 M
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Chapter 17: Thermodynamics: Directionality of Chemical Reactions
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Q1) For a particular reaction, the value of DH<sup>\(\circ\)</sup>= +98.8 kJ and DS<sup>\(\circ\)</sup> = +141.5 J/K. This reaction is
A) product-favored, because DS<sup>\(\circ\)</sup><sub>universe</sub> is positive
B) reactant-favored, because DS<sup>\(\circ\)</sup><sub>universe</sub> is positive
C) reactant-favored, because DS<sup>\(\circ\)</sup><sub>universe</sub> is negative
D) product-favored, because DS<sup>\(\circ\)</sup><sub>universe</sub> is negative
E) cannot determine without further information
Q2) Which process is reactant-favored?
A) decomposition of iron ore into pure iron
B) the melting of ice at 25<sup>\(\circ\)</sup>C
C) diffusion of the odor of cooking food
D) the freezing of water at -5<sup>\(\circ\)</sup>C
E) dissociation of a strong acid in water
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Chapter 18: Electrochemistry and Its Applications
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Q1) In the salt bridge in an electrochemical cell,
A) negatively charged ions migrate towards the anode compartment
B) positively charged ions migrate towards the anode compartment
C) negatively charged ions migrate towards the cathode compartment
D) both b and c
E) neither a nor c
Q2) Although aluminum is a moderately reactive metal, it can be regarded as unreactive in everyday use because a layer of aluminum oxide coats the surface and prevents further oxidation. This is an example of
A) anodic inhibition.
B) cathodic inhibition.
C) galvanization.
D) cathodic protection.
E) anodic protection.
Q3) Which of the following would require the most Faradays of electricity?
A) converting 2.0 mol of Al<sup>3+</sup> into Al
B) converting 2.0 mol of Cr<sup>3+</sup> into Cr
C) converting 1.5 mol of Cr<sup>6+</sup> into Cr
D) converting 3.0 mol of Sn<sup>4+</sup> into Sn<sup>2+</sup>
E) converting 2.0 mol of Sn<sup>4</sup><sup>+</sup> into Sn
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Chapter 19: Nuclear Chemistry
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Q1) On a plot of number of neutrons as a function of number of protons, the stable (non-radioactive) isotopes form a "peninsula of stability". Isotopes with a greater or lesser number of neutrons than this, and all isotopes beyond a certain number of protons, are radioactive. Briefly describe the most common type of radioactive decay exhibited by each of these groups.
Q2) Gamma radiation is best described as
A) neutral particles that weigh approximately one atomic mass unit.
B) positive particles that are identical to the nucleus of an atom of <sup>4</sup>He.
C) electrons ejected at high speeds from a radioactive nucleus.
D) a form of electromagnetic radiation.
E) high-speed particles similar in size to an electron, but oppositely charged.
Q3) The force holding the protons and neutrons together in a nucleus is known as
A) binding energy
B) gamma radiation
C) electrostatic attraction
D) gravity
E) mass defect
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