Course Introduction
Chemistry for Engineers
Chapter Exam Questions
Chemistry for Engineers provides a foundational understanding of chemical principles and their practical applications in engineering contexts. The course covers essential topics such as atomic structure, chemical bonding, stoichiometry, thermodynamics, chemical kinetics, and equilibria, with a focus on materials and processes relevant to engineering disciplines. Through lectures and laboratory sessions, students develop skills in analyzing chemical reactions, understanding material properties, and applying chemical knowledge to solve engineering problems in areas such as energy production, environmental management, and materials design.
Recommended Textbook
General Chemistry The Essential Concepts 7th Edition by Raymond Chang
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Page 2
Chapter 1: Introduction
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Sample Questions
Q1) How many significant figures does the difference 218.7201 - 218.63 contain? A)1 B)2 C)3
D)5
E)7
Answer: A
Q2) You just measured a sugar cube and obtained the following information: mass = 3.48 g
height = length = width = 1.3 cm
Determine the volume and density of the cube. Suppose the sugar cube was added to a cup of water. Before it dissolves, will the sugar cube float or sink to the bottom?
Answer: Volume of the sugar cube = 2.2 cm<sup>3</sup>; density of the sugar cube = 1.6 g/cm<sup>3</sup>. Before dissolving, the sugar cube will sink in a cup of water.
Q3) Identify the following as a physical or chemical change: Formation of snowflakes.
Answer: Physical
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3
Chapter 2: Atoms, Molecules, and Ions
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Q1) The formula for isopropyl alcohol is sometimes written as (CH<sub>3</sub>)<sub>2</sub>CHOH to better indicate how the atoms are connected. How many hydrogen atoms would be contained in 3 dozen isopropyl alcohol molecules? Answer: 288
Q2) Give the formula of carbonic acid. Answer: H<sub>2</sub>CO<sub>3</sub>
Q3) Give the formula of magnesium nitrate. Answer: Mg(NO<sub>3</sub>)<sub>2</sub>
Q4) Which pair of elements would be most likely to form an ionic compound?
A)P and Br
B)Zn and K
C)F and Al
D)C and S
E)Al and Rb
Answer: C
Q5) Give the formula for carbon disulfide. Answer: CS<sub>2</sub>
Q6) Use the periodic table above to show where the alkali metals are located. Answer: Group 1A
Page 4
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Chapter 3: Stoichiometry
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Sample Questions
Q1) Which of the following statements are true about a sample of sulfur and a sample of oxygen if the two samples are of equal mass?
I. The number of electrons in the two samples is about the same.
II. The number of protons in the two samples is about the same.
III. The number of atoms in the two samples is about the same.
IV. There are roughly twice as many sulfur atoms as oxygen atoms.
V. There are roughly twice as many oxygen atoms as sulfur atoms.
Answer: I, II, and V
Q2) How many ICl<sub>3</sub> molecules are present in 1.75 kg of ICl<sub>3</sub>?
Answer: 4.52 * 10<sup>24</sup>
Q3) What is the limiting reagent when 27.0 g of P and 68.0 g of I<sub>2</sub> react according to the following chemical equation?
2P(s)+ 3I<sub>2</sub>(s) \(\rarr\)2PI<sub>3</sub>(s)
Answer: I<sub>2</sub>
Q4) Calculate the molecular mass, in g/mol, of P<sub>4</sub>O<sub>10</sub>.
Answer: 283.9 g/mol
Q5) What percent by mass of oxygen is present in carbon monoxide, CO?
Answer: 57%
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Chapter 4: Reactions in Aqueous Solution
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Q1) In the following chemical reaction the oxidizing agent is: 5S + 6KNO<sub>3</sub> + 2CaCO<sub>3</sub> \(\rarr\)3K<sub>2</sub>SO<sub>4</sub> + 2CaSO<sub>4</sub> + CO<sub>2</sub> + 3N<sub>2</sub>
A)S
B)N<sub>2</sub>
C)KNO<sub>3</sub>
D)CaSO<sub>4</sub>
E)CaCO<sub>3</sub>
Q2) Batteries in our cars generate electricity by the following chemical reaction. Pb + PbO<sub>2</sub> + 2H<sub>2</sub>SO<sub>4</sub> \(\rarr\) 2PbSO<sub>4</sub> + 2H<sub>2</sub>O
What is the reducing agent in this process?
Q3) Batteries in our cars generate electricity by the following chemical reaction. Pb + PbO<sub>2</sub> + 2H<sub>2</sub>SO<sub>4</sub> \(\rarr\) 2PbSO<sub>4</sub> + 2H<sub>2</sub>O
What is the oxidizing agent in this process?
Q4) Write the net ionic equation for the following reaction. Aqueous iron(III)sulfate is added to aqueous sodium sulfide to produce solid iron(III)sulfide and aqueous sodium sulfate.
Page 6
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Chapter 5: Gases
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Sample Questions
Q1) Baking powder is made up of sodium hydrogen carbonate and calcium hydrogen phosphate. When baking powder is wet, these components react to produce carbon dioxide. The equation for this reaction is given below. NaHCO<sub>3</sub>(aq)+ CaHPO<sub>4</sub>(aq)\(\rarr\)NaCaPO<sub>4</sub>(aq)+ CO<sub>2</sub>(g)+ H<sub>2</sub>O(l)
Assuming all of the carbon dioxide was released as a gas, how many liters of CO<sub>2</sub>(g)would be formed at room temperature from 4.00 g of NaHCO<sub>3</sub> and excess CaHPO<sub>4</sub>?
Q2) The temperature of an ideal gas in a 5.00 L container originally at 1 atm pressure and 25°C is lowered to 220 K. Calculate the new pressure of the gas.
A)1.0 atm
B)1.35 atm
C)8.8 atm
D)0.738 atm
E)0.114 atm
Q3) How many grams of N<sub>2</sub>O, nitrous oxide, are contained in 500. mL of the gas at STP?
Q4) What is the significance of the magnitude of the van der Waals "a" constant?
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Chapter 6: Energy Relationships in Chemical Reactions
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Sample Questions
Q1) Naphthalene combustion can be used to calibrate the heat capacity of a bomb calorimeter. The heat of combustion of naphthalene is -40.1 kJ/g. When 0.8210 g of naphthalene was burned in a calorimeter containing 1,000. g of water, a temperature rise of 4.21°C was observed. What is the heat capacity of the bomb calorimeter excluding the water?
A)32.9 kJ/°C
B)7.8 kJ/°C
C)3.64 kJ/°C
D)1.76 kJ/°C
E)15.3 kJ/°C
Q2) Which of the following processes is endothermic?
A)O<sub>2</sub>(g)+ 2H<sub>2</sub>(g)\(\rarr\)2H<sub>2</sub>O(g)
B)H<sub>2</sub>O(g)\(\rarr\)H<sub>2</sub>O(l)
C)3O<sub>2</sub>(g)+ 2CH<sub>3</sub>OH(g)\(\rarr\)2CO<sub>2</sub>(g)+ 2H<sub>2</sub>O(g)
D)H<sub>2</sub>O(s)\(\rarr\)H<sub>2</sub>O(l)
Q3) The heat absorbed by a system at constant pressure is equal to \(\Delta\)E + P\(\Delta\)V.
A)True
B)False
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Chapter 7: The Electronic Structure of Atoms
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Sample Questions
Q1) Complete this sentence: Atoms emit visible and ultraviolet light
A)as electrons jump from lower energy levels to higher levels.
B)as the atoms condense from a gas to a liquid.
C)as electrons jump from higher energy levels to lower levels.
D)as they are heated and the solid melts to form a liquid.
E)as the electrons move about the atom within an orbit.
Q2) The Bohr model of the hydrogen atom found its greatest support in experimental work on the photoelectric effect.
A)True
B)False
Q3) The electron in a hydrogen atom falls from an excited energy level to the ground state in two steps, causing the emission of photons with wavelengths of 2624 and 97.2 nm. What is the quantum number of the initial excited energy level from which the electron falls? A)2
Q4) Write the ground state electron configuration for a lead atom.
Page 9
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Chapter 8: The Periodic Table
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Sample Questions
Q1) Which of the elements listed below has the highest first ionization energy?
A)C
B)Ge
C)P
D)O
E)Se
Q2) The electron configuration of the outermost electrons of atoms of the halogen group is ns<sup>2</sup>np<sup>7</sup>.
A)True
B)False
Q3) What is the charge on the monatomic ion that calcium forms in its compounds?
A)+2
B)+1
C)-1
D)-2
E)-3
Q4) Write the ground-state electron configuration for Ca<sup>2+</sup>.
Q5) Briefly explain why the atomic radius decreases within a period when moving from left to right.
Page 10
Q6) Write the ground-state electron configuration for I<sup>-</sup>.
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Chapter 9: Chemical Bonding I: the Covalent Bond
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Sample Questions
Q1) Write a Lewis structure for the phosphate ion, PO<sub>4</sub><sup>3-</sup>, that obeys the octet rule, showing all non-zero formal charges, and give the total number of resonance structures for PO<sub>4</sub><sup>3-</sup> that obey the octet rule.
Q2) Which one of the following molecules has an atom with an expanded octet?
A)HCl
B)AsCl<sub>5</sub>
C)ICl
D)NCl<sub>3</sub>
E)Cl<sub>2</sub>
Q3) Which of the elements listed below has the greatest electronegativity?
A)Na
B)As
C)Ga
D)Cs
E)Sb
Q4) The polarity of covalent bonds increases as the percent ionic character increases.
A)True
B)False
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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization of
Atomic
Orbitals
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Sample Questions
Q1) Which of the following correctly lists species in order of increasing bond length?
A)C<sub>2</sub><sup>-</sup> < C<sub>2</sub> < C<sub>2</sub><sup>+</sup>
B)C<sub>2</sub> < C<sub>2</sub><sup>+</sup> < C<sub>2</sub><sup>-</sup>
C)C<sub>2</sub><sup>-</sup> < C<sub>2</sub><sup>+</sup> < C<sub>2</sub>
D)C<sub>2</sub><sup>+</sup> < C<sub>2</sub> < C<sub>2</sub><sup>-</sup>
E)C<sub>2</sub><sup>+</sup> < C<sub>2</sub><sup>-</sup> < C<sub>2</sub>
Q2) Give the number of lone pairs around the central atom and the geometry of the ion PCl<sub>4</sub><sup>-</sup>.
A)0 lone pairs, tetrahedral
B)1 lone pair, distorted tetrahedron (seesaw)
C)1 lone pair, square pyramidal
D)1 lone pair, tetrahedral
E)2 lone pairs, square planar
Q3) Ozone (O<sub>3</sub>)is an allotropic form of oxygen. Use VSEPR theory to predict the shape of the ozone molecule.
Q4) The hybridization of B in the BF<sub>3</sub> molecule is sp<sup>3</sup>.
A)True
B)False
Page 12
Q5) Which should have the longer bond, B<sub>2</sub> or B<sub>2</sub><sup></sup>?
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Chapter 11: Introduction to Organic Chemistry
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Sample Questions
Q1) How many structural isomers are there of C<sub>4</sub>H<sub>10</sub>?
A)4
B)6
C)2
D)8
E)10
Q2) The expected product from the addition of HCl to CH<sub>3</sub>-CH<sub>2</sub>-CH=CH<sub>2</sub> is A)CH<sub>3</sub>-CH<sub>2</sub>-CH=CHCl.
B)CH<sub>3</sub>-CH<sub>2</sub>-CCl=CH<sub>2</sub>.
C)CH<sub>3</sub>-CHCl-CH=CH<sub>2</sub>.
D)CH<sub>3</sub>-CH<sub>2</sub>-CH<sub>2</sub>-CH<sub>2</sub>Cl.
E)CH<sub>3</sub>-CH<sub>2</sub>-CHCl-CH<sub>3</sub>.
Q3) The octane rating of gasoline refers to its
A)percentage C<sub>8</sub>H<sub>18</sub> by volume.
B)radiation dose.
C)alcohol level.
D)ability to resist engine knocking.
E)percentage of unsaturated hydrocarbons.
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Page 13
Chapter 12: Intermolecular Forces and Liquids and Solids
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Sample Questions
Q1) Given that the heat of vaporization of mercury is 59.0 kJ/mol and the vapor pressure of mercury is 0.0017 torr at 25°C, calculate the normal boiling point of mercury.
Q2) The zincblende structure of ZnS has the relatively large sulfide ions arranged at the lattice points of a face-centered cubic structure. The edge length of this cubic unit cell is 540.9 pm. Determine the density of zincblende.
A)3.081 g/cm<sup>3</sup>
B)1.023 g/cm<sup>3</sup>
C)4.091 g/cm<sup>3</sup>
D)2.046 g/cm<sup>3</sup>
E)2.032 g/cm<sup>3</sup>
Q3) The molar enthalpy of vaporization of carbon disulfide is 26.74 kJ/mol, and its normal boiling point is 46°C. What is the vapor pressure of CS<sub>2</sub> at 0°C?
A)447 torr
B)4160 torr
C)313 torr
D)139 torr
E)5.47 torr
Q4) Indicate all the types of intermolecular forces of attraction in HCl(g).
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Chapter 13: Physical Properties of Solutions
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Sample Questions
Q1) How many grams of water are needed to dissolve 27.8 g of ammonium nitrate
NH<sub>4</sub>NO<sub>3</sub> in order to prepare a 0.452 m solution?
A)769 g
B)36.2 g
C)100. g
D)0.157 g
E)157 g
Q2) For water K<sub>f</sub> = 1.86°C/m. Therefore, the freezing points of 1.0 M aqueous KCl and C<sub>2</sub>H<sub>5</sub>OH (ethanol)solutions are the same.
A)True
B)False
Q3) The solubility of CO<sub>2</sub> gas in water
A)increases with increasing temperature.
B)decreases with decreasing temperature.
C)decreases with increasing temperature.
D)is not dependent on temperature.
Q4) The solubility of gases in water always decreases with increasing temperature.
A)True
B)False
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Chapter 14: Chemical Kinetics
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Sample Questions
Q1) Appropriate units for a second-order rate constant are A)M/s.
B)1/M·s.
C)1/s.
D)1/M<sup>2</sup>·s.
Q2) The activation energy for the following reaction is 60. kJ/mol. Sn<sup>2+</sup> + 2Co<sup>3+</sup> \(\rarr\)Sn<sup>4+</sup> + 2Co<sup>2+</sup> <sup> </sup>By what factor (how many times)will the rate constant increase when the temperature is raised from 10°C to 28°C?
A)1.002
B)4.6
C)5.6
D)2.8
E)696
Q3) The rate law predicted by the following two-step mechanism is rate = k[A][B].
A \(\rarr\) C + B slow
A + B \(\rarr\) C + E fast
A)True
B)False
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Chapter 15: Chemical Equilibrium
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Q1) What conditions are used in the Haber process to enhance the yield of ammonia? Explain why each condition affects the yield in terms of the Le Châtelier principle.
Q2) Describe why addition of a catalyst does not affect the equilibrium constant for a reaction.
Q3) For the common allotropes of carbon (graphite and diamond), C(gr)\(\rarr\)<sub> </sub>C(dia)with equilibrium constant K = 0.32. The molar volumes of graphite and diamond are, respectively, 5.30 cm<sup>3</sup>/mol and 3.42 cm<sup>3</sup>/mol; \(\Delta\)H<sub>f</sub> of diamond is 1.90 kJ/mol. This data suggests that the formation of diamond is favored at A)low temperatures and low pressures. B)high temperatures and low pressures. C)low temperatures and high pressures. D)high temperatures and high pressures.
Q4) The dissociation of solid silver chloride in water to produce silver ions and chloride ions has an equilibrium constant of 1.8 * 10<sup>-18</sup>. Based on the magnitude of the equilibrium constant, is silver chloride very soluble in water? Why?
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Chapter 16: Acids and Bases
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Q1) What mass of sodium formate (HCOONa)must be added to 350. mL of water in order to obtain a solution having a pH of 8.50? [K<sub>a</sub>(HCOOH)= 1.77 * 10<sup>-4</sup>]
A)0.23 g
B)4.3 g
C)35 g
D)12 g
E)130 g
Q2) Which of these lists of molecules is arranged in order of increasing acid strength?
A)H<sub>2</sub>S < H<sub>2</sub>O < H<sub>2</sub>Se
B)H<sub>2</sub>O < H<sub>2</sub>S < H<sub>2</sub>Se C)H<sub>2</sub>Se < H<sub>2</sub>O < H<sub>2</sub>S
D)H<sub>2</sub>S < H<sub>2</sub>Se < H<sub>2</sub>O E)H<sub>2</sub>O < H<sub>2</sub>Se < H<sub>2</sub>S
Q3) In comparing three solutions with pH's of 2.0, 4.8, and 5.2, which is most acidic?
Q4) The pH of a 0.02 M solution of an unknown weak base is 8.1. What is the pK<sub>b</sub> of the unknown base?
Q5) Write the chemical formula for nitric acid.
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Chapter 17: Acid-Base Equilibria and Solubility Equilibria
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Q1) Calculate the silver ion concentration in a saturated solution of silver(I)carbonate (K<sub>sp</sub> = 8.1 * 10<sup>-12</sup>).
A)5.0 * 10<sup>-5</sup> M
B)2.5 * 10<sup>-4</sup> M
C)1.3 * 10<sup>-4</sup> M
D)2.0 * 10<sup>-4</sup> M
E)8.1 * 10<sup>-4</sup> M
Q2) The K<sub>sp</sub> value for lead(II)chloride is 2.4 * 10<sup>-4</sup>. What is the molar solubility of lead(II)chloride?
A)2.4 * 10<sup>-4</sup> mol/L
B)6.2 * 10<sup>-2</sup> mol/L
C)7.7 * 10<sup>-3</sup> mol/L
D)3.9 * 10<sup>-2</sup> mol/L
E)6.0 * 10<sup>-5</sup> mol/L
Q3) What is the effective pH range for a sodium acetate/acetic acid buffer? (For CH<sub>3</sub>COOH, K<sub>a</sub>= 1.8 * 10<sup>-5</sup>)
Q4) Describe how to prepare 500. mL of a cyanic acid (HCNO)/sodium cyanate (NaCNO)buffer having a pH of 4.80. [K<sub>a</sub>(HCNO)= 2.0 * 10<sup>-4</sup>]
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Chapter 18: Thermodynamics
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Q1) Which of these species has the highest entropy (S°)at 25°C?
A)CO(g)
B)CH<sub>4</sub>(g)
C)NaCl(s)
D)H<sub>2</sub>O(l)
E)Fe(s)
Q2) Calculate the free energy of formation of NaBr(s)given the following information: NaBr(s) \(\rarr\)Na(s)+ <sup>1</sup>/<sub>2</sub>Br<sub>2</sub>(l), \(\Delta\)G° = 349 kJ/mol
Q3) For the reaction H<sub>2</sub>O<sub>2</sub>(g) \(\rarr\) H<sub>2</sub>O(g)+ <sup>1</sup>/<sub>2</sub>O<sub>2</sub>(g), \(\Delta\)H° = -106 kJ/mol and \(\Delta\)S° = 58 J/K·mol at 25°C. Is H<sub>2</sub>O<sub>2</sub>(g)stable with respect to dissociation into water vapor and oxygen gas at 25°C?
Q4) Choose the substance with the higher entropy per mole at a given temperature: O<sub>2</sub>(g)at 5 atm or O<sub>2</sub>(g)at 0.5 atm.
Q5) Choose the substance with the higher entropy per mole at a given temperature: CO<sub>2</sub>(g)or CO<sub>2</sub>(aq).
Q6) How does the entropy change when a liquid is vaporized?
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Chapter 19: Redox Reactions and Electrochemistry
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Q1) Determine the equilibrium constant, K<sub>eq</sub>, at 25°C for the reaction 2Br<sup>- </sup>(aq)+ I<sub>2</sub>(s)\(\rarr\)Br<sub>2</sub>(l)+ 2I<sup></sup>(aq)
A)5.7 * 10<sup>-19</sup>
B)18.30
C)1.7 * 10<sup>54</sup>
D)1.9 * 10<sup>18</sup>
E)5.7 * 10<sup>-55</sup>
Q2) A metal object is to be gold-plated by an electrolytic procedure using aqueous AuCl<sub>3</sub> electrolyte. Calculate the number of moles of gold deposited in 3.0 min by a constant current of 10. A.
A)6.2 * 10<sup>-3</sup> mol
B)9.3 * 10<sup>-3</sup> mol
C)1.8 * 10<sup>-2</sup> mol
D)3.5 * 10<sup>-5 </sup>mol
E)160 mol
Q3) Will H<sub>2</sub>(g)form when Ag is placed in 1.0 M HCl?
Q4) Complete and balance the following redox reaction under acidic conditions: ClO<sub>2</sub><sup>-</sup>(aq)\(\rarr\) ClO<sub>2</sub>(g)+ Cl<sup>-</sup>(aq)
Q5) Will H<sub>2</sub>(g)form when Sn is placed in 1.0 M HCl?
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Chapter 20: The Chemistry of Coordination Compounds
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Q1) A molecule or ion that provides an electron pair for coordinate covalent bond formation is called a Lewis ________.
Q2) The electron configuration of a Ti atom is
A)[Ne]3s<sup>2</sup>3d<sup>2</sup>.
B)[Ne] 3s<sup>2</sup>4d<sup>2</sup>.
C)[Ar]4s<sup>2</sup>3d<sup>2</sup>.
D)[Ar]4s<sup>2</sup>4d<sup>2</sup>.
E)[Ar]3d<sup>4</sup>.
Q3) In the complex ion [ML<sub>6</sub>]<sup>n+</sup>, M<sup>n+</sup> has seven d electrons and L is a strong field ligand. According to crystal field theory, the magnetic properties of the complex ion correspond to how many unpaired electrons?
A)0
Q4) Bidentate and polydentate ligands are also called chelating agents.
A)True
B)False
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Chapter 21: Nuclear Chemistry
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Q1) Protactinium-234 has a half-life of 1 minute. How much of a 400. g sample protactinium would remain after 1 minute?
Q2) The half-life of <sup>90</sup>Sr is 29 years. What fraction of the atoms in a sample of <sup>90</sup>Sr would remain 175 years later?
A)0.17
B)0.12
C)0.062
D)0.015
E)0.50
Q3) Beta particles are identical to A)protons.
B)helium atoms.
C)hydrogen atoms.
D)helium nuclei.
E)electrons.
Q4) When a <sup>87</sup>Br nucleus emits a beta particle, the nuclear species that results is ___________.
Q5) Protactinium-234 has a half-life of 1 minute. How much of a 400. g sample protactinium would remain after 2 minutes?
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Chapter 22: Organic Polymerssynthetic and Natural
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Q1) Which of these molecules is a product of the hydrolysis of DNA?
A)acetic acid
B)glucose
C)adenine
D)ribose
E)water
Q2) A protein is
A)a polysaccharide.
B)a saturated ester of glycerol.
C)one of the units making up a nucleic acid.
D)a polymer of amino acids.
E)an aromatic hydrocarbon.
Q3) The intermolecular force between bases on the opposite strands of DNA responsible for its double-helical structure is
A)hydrogen bonding.
B)dispersion force.
C)covalent bonding.
D)ionic force.
E)dipole-dipole force.
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