Principles of Chemistry
Exam Answer Key
Course Introduction
Principles of Chemistry provides a foundational understanding of chemical concepts and theories essential to the study of matter and its interactions. The course explores atomic structure, chemical bonding, stoichiometry, states of matter, thermochemistry, and the properties of solutions. Students also examine the periodic table, chemical reactions, and the underlying principles governing chemical behavior. Emphasis is placed on problem-solving skills, laboratory techniques, and the application of chemical principles in real-world contexts, preparing students for advanced study in science, engineering, health, and related fields.
Recommended Textbook
Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg
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Chapter 1: Keys to the Study of Chemistry
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Sample Questions
Q1) At a pressure of one billionth (10<sup>-9</sup>) of atmospheric pressure, there are about 2.7 × 10<sup>10</sup> molecules in one cubic centimeter of a gas. How many molecules is this per cubic meter?
A) 2.7 × 10<sup>16</sup>
B) 2.7 × 10<sup>14</sup>
C) 2.7 × 10<sup>12</sup>
D) 2.7 × 10<sup>8</sup>
E) 2.7 × 10<sup>4</sup>
Answer: A
Q2) The number 6.0448, rounded to 3 decimal places, becomes 6.045.
A)True
B)False
Answer: True
Q3) The weight of a coin measured as 1.96235 g on one balance is definitely more accurate than a weight measurement of 1.95 g on another balance.
A)True
B)False
Answer: False
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Chapter 2: The Components of Matter
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Sample Questions
Q1) What is the formula for lead (II) oxide?
A) PbO
B) PbO<sub>2</sub>
C) Pb<sub>2</sub>O
D) PbO<sub>4</sub>
E) Pb<sub>2</sub>O<sub>3</sub>
Answer: A
Q2) Barium sulfate is used in manufacturing photographic paper. What is its formula?
A) BaSO<sub>4</sub>
B) Ba(SO<sub>4</sub>)<sub>2</sub>
C) Ba<sub>2</sub>SO<sub>4</sub>
D) Ba<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>
E) BaSO<sub>3</sub>
Answer: A
Q3) When an alkali metal combines with a non-metal, a covalent bond is normally formed.
A)True
B)False
Answer: False
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Chapter 3: Stoichiometry of Formulas and Equations
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Q1) Gadolinium oxide, a colorless powder which absorbs carbon dioxide from the air, contains 86.76 mass % Gd. Determine its empirical formula.
A) Gd<sub>2</sub>O<sub>3</sub>
B) Gd<sub>3</sub>O<sub>2</sub>
C) Gd<sub>3</sub>O<sub>4</sub>
D) Gd<sub>4</sub>O<sub>3</sub>
E) GdO
Answer: A
Q2) Alkanes are compounds of carbon and hydrogen with the general formula C<sub>n</sub>H<sub>2</sub><sub>n</sub><sub>+2</sub>. An alkane component of gasoline has a molar mass of between 125 and 130 g/mol. What is the value of n for this alkane?
A) 4
B) 9
C) 10
D) 13
E) 14
Answer: B
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Chapter 4: The Major Classes of Chemical Reactions
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Q1) Ammonia, NH<sub>3</sub>, produces hydroxide ions in aqueous solution by
A) donating a proton to an acid molecule.
B) donating a proton to a water molecule.
C) donating a hydrogen ion to a water molecule.
D) accepting an electron from an acid molecule.
E) accepting a proton from a water molecule.
Q2) The compound PS1U1B14S1U1B0OS1U1B110S1U1B0is used in refining sugar. Select the classification for the reaction in which it is synthesized. PS1U1B14S1U1B0(s) + 5OS1U1B12S1U1B0(g) \(\rarr\)PS1U1B14S1U1B0OS1U1B110S1U1B0(s)
A) combination
B) decomposition
C) displacement
D) acid-base
E) precipitation
Q3) Chemical reactions generally reach equilibrium because one of the reactants is used up.
A)True
B)False
Q4) Complete and balance the equation for the following acid-base reaction. Ca(OH)<sub>2</sub>(aq) + HCl(aq) \(\to\)
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Chapter 5: Gases and the Kinetic-Molecular Theory
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Sample Questions
Q1) Calculate the root-mean-square speed of methane, CH<sub>4 </sub>(g), at 78°C.
A) 23 m/s
B) 350 m/s
C) 550 m/s
D) 667 m/s
E) 740 m/s
Q2) A 255-mL gas sample weighing 0.292 g is at 52810 Pa and 127°C.
a. How many moles of gas are present?
b. What is the molar mass of the gas?
Q3) Lithium oxide is an effective absorber of carbon dioxide and can be used to purify air in confined areas such as space vehicles. What volume of carbon dioxide can be absorbed by 1.00 kg of lithium oxide at 25°C and 1.00 atm? Li<sub>2</sub>O(aq) + CO<sub>2</sub>(g) \(\to\) Li<sub>2</sub>CO<sub>3</sub>(s)
A) 687 mL
B) 819 mL
C) 687 L
D) 819 L
E) 22.4 L
Q4) State Avogadro's Law.
Q5) Calculate the density in g/L of gaseous SF<sub>6</sub> at 50.0°C and 650. torr.
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Chapter 6: Thermochemistry: Energy Flow and Chemical Change
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Q1) Which of the following is not a state function?
A) internal energy
B) volume
C) work
D) pressure
E) enthalpy
Q2) A piece of copper metal is initially at 100.0°C. It is dropped into a coffee cup calorimeter containing 50.0 g of water at a temperature of 20.0°C. After stirring, the final temperature of both copper and water is 25.0°C. Assuming no heat losses, and that the specific heat (capacity) of water is 4.18 J/(g·K), what is the heat capacity of the copper in J/K?
A) 2.79 J/K
B) 3.33 J/K
C) 13.9 J/K
D) 209 J/K
E) none of the above
Q3) a. Explain fully what is meant by the term "state function."
b. (i) Give two examples of thermodynamic quantities which state functions. (ii) Give two examples of thermodynamic quantities which are not state functions.
Page 8
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Chapter 7: Quantum Theory and Atomic Structure
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Q1) In the quantum mechanical treatment of the hydrogen atom, the probability of finding an electron at any point is proportional to the wave function \(\varPsi\) .
A)True
B)False
Q2) As the frequency of electromagnetic radiation increases, its wavelength also increases.
A)True
B)False
Q3) The Rydberg equation is an example of an empirical equation.
A)True
B)False
Q4) The FM station KDUL broadcasts music at 99.1 MHz. Find the wavelength of these waves.
A) 1.88 × 10<sup>-2 </sup>m
B) 0.330 m
C) 3.03 m
D) 5.33 × 10<sup>2</sup> m
E) > 10<sup>3</sup> m
Q5) What is the speed of an electron in m/s if its wavelength is 0.155 nm?
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Chapter 8: Electron Configuration and Chemical Periodicity
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Q1) Which of the following elements is paramagnetic?
A) Kr
B) Zn
C) Sr
D) V
E) Ar
Q2) Which one of the following statements about orbital energies is incorrect?
A) In the hydrogen atom, the energy of an orbital depends only on the value of the quantum number n.
B) In many-electron atoms the energy of an orbital depends on both n and l.
C) Inner electrons shield outer electrons more effectively than do electrons in the same orbital.
D) The splitting of sublevels in many-electron atoms is explained in terms of the penetration effect.
E) The energy of a given orbital increases as the nuclear charge Z increases.
Q3) Write down the maximum number of electrons in an atom which can have
a. quantum number n = 4.
b. orbital designation 3d.
c. orbital designation 2p<sub>z</sub>.
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Chapter 9: Models of Chemical Bonding
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Q1) Quartz (SiO<sub>2</sub>) is a solid with a melting point of 1550 °C. The bonding in quartz is best described as
A) lattice energy.
B) network attractions.
C) ionic bonding.
D) covalent bonding.
E) metallic bonding.
Q2) As a measure of the strength of metallic bonding, the boiling point of a metal is a better indicator than its melting point.
A)True
B)False
Q3) When an atom is represented in a Lewis electron dot symbol, the element symbol represents ______________ and the dots represent ______________.
Q4) Select the most polar bond amongst the following.
A) C-O
B) Si-F
C) Cl-F
D) C-F
E) C-I
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Chapter 10: The Shapes of Molecules
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Q1) In which one of the following species is the best Lewis structure a resonance structure?
A) NH<sub>3</sub>
B) CO<sub>2</sub>
C) SF<sub>6</sub>
D) O<sub>2</sub>
E) CO<sub>3</sub><sup>2-</sup>
Q2) Predict the actual bond angles in SF<sub>3</sub><sup>+</sup> using the VSEPR theory.
A) more than 120°
B) exactly 120°
C) between 109° and 120°
D) between 90° and 109°
E) less than 90°
Q3) A molecule which contains polar bonds will always have a dipole moment.
A)True
B)False
Q4) List the three important ways in which molecules can violate the octet rule, and in each case draw one Lewis structure of your choice as an example.
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Chapter 11: Theories of Covalent Bonding
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Q1) How many sigma and pi bonds, respectively, are in the molecule below?
CH<sub>3</sub>CH<sub>2</sub>CHCHCH<sub>3</sub>?
A) 16, 1
B) 16, 0
C) 15, 1
D) 15, 0
E) 14, 1
Q2) Valence bond theory is able to explain why molecular oxygen is paramagnetic.
A)True
B)False
Q3) In the context of molecular orbital (MO) theory, explain how atomic p orbitals can give rise to:
a. a \(\sigma\) MO
b. a \(\pi\) MO
Q4) A molecule with the formula AX<sub>4 </sub>uses _________ to form its bonds.
A) sp hybrid orbitals
B) sp<sup>2 </sup>hybrid orbitals
C) sp<sup>3</sup> hybrid orbitals
D) sp<sup>3</sup>d hybrid orbitals
E) sp<sup>3</sup>d<sup>2</sup> hybrid orbitals
Page 13
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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes
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Q1) The phase diagram of a substance shows the energy changes associated with changes of state.
A)True
B)False
Q2) Hexagonal close packing of identical atoms occurs when close-packed layers are stacked in an abcabc.... arrangement.
A)True
B)False
Q3) For the solid forms of the following elements, which one is most likely to be of the molecular type?
A) Xe
B) C
C) Pb
D) S
E) Cr
Q4) Iron has a body-centered cubic unit cell, and a density of 7.87 g/cm<sup>3</sup>.
Calculate the edge length of the unit cell, in pm. (The atomic mass of iron is 55.85 amu. Also, 1 amu = 1.661 × 10<sup>-24</sup> g.)
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Q5) How do the electrical properties of semiconductors differ from those of metals?
Chapter 13: The Properties of Mixtures: Solutions and Colloids
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Q1) Human blood has a molar concentration of solutes of 0.30 M. What is the osmotic pressure of blood at 25°C?
A) 0.012 atm
B) 0.62 atm
C) 6.8 atm
D) 7.3 atm
E) >10. atm
Q2) Based only on the relative lattice energies of the compounds below, which one would be expected to have the lowest solubility in water?
A) NaBr
B) CaS
C) NaOH
D) KI
E) CsCl
Q3) A 0.137 m solution of sodium ferrocyanide (Na<sub>4</sub>Fe(CN)<sub>6</sub>.10H<sub>2</sub>O) in water has a freezing point depression of 0.755°C. Calculate the van't Hoff factor (i) under these conditions. (K<sub>f </sub>= 1.86°C/m)
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Chapter 14: Periodic Patterns in the Main Group Elements:
Bonding, Structure, and Reactivity
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Q1) In gaseous BeCl<sub>2</sub>, Be does not obey the octet rule; in solid BeCl<sub>2</sub> it does.
A)True
B)False
Q2) Which one of the following elements is likely to exhibit the most violent and rapid reaction with water?
A) Na
B) Rb
C) Mg
D) Sr
E) Cl<sub>2</sub>
Q3) Sodium tripolyphosphate is useful
A) as a starting material in the synthesis of organic phosphorus compounds.
B) as a water-softening agent.
C) as an emulsifier in making of processed cheese.
D) as a radiator corrosion inhibitor.
E) as a fertilizer.
Q4) Ionic hydrides do not have exact (stoichiometric) formulas.
A)True
B)False
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Chapter 15: Organic Compounds and the Atomic Properties of Carbon
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Q1) Each amino acid has two functional groups in common and one of 20 other groups attached to the \(\alpha\)-carbon. The two functional groups are
A) carboxyl and amine.
B) ester and amine.
C) carboxyl and amide.
D) alcohol and amine.
E) carboxyl and peptide.
Q2) Explain how a salt bridge can arise a protein molecule.
Q3) The backbone of protein molecules consists of repeating N-C-C-O units.
A)True
B)False
Q4) Draw and name two non-alkenes with the formula C<sub>4</sub>H<sub>8</sub>.
Q5) Both nylons and proteins are polyamides.
A)True
B)False
Q6) Given that each 3-base sequence in DNA is a "code word" for a particular amino acid, how many different code words are possible using a 3-base sequence and the bases available in DNA?
Page 17
Q7) Explain how a disulfide bridge can arise a protein molecule.
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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions
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Q1) Butadiene, C<sub>4</sub>H<sub>6</sub> (used to make synthetic rubber and latex paints) reacts to C<sub>8</sub>H<sub>12</sub> with a rate law of rate = 0.014 L/(mol·s) [C<sub>4</sub>H<sub>6</sub>]<sup>2</sup>. What will be the concentration of C<sub>4</sub>H<sub>6 </sub>after 3.0 hours if the initial concentration is 0.025 M?
A) 0.0052 M
B) 0.024 M
C) 43 M
D) 190 M
E) 0.0000 M
Q2) The greater the energy of activation, E<sub>a</sub>, the faster will be the reaction.
A)True
B)False
Q3) A catalyst lowers the activation energy but does not affect the mechanism of a reaction.
A)True
B)False
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Chapter 17: Equilibrium: the Extent of Chemical Reactions
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Q1) Which of the following has an effect on the magnitude of the equilibrium constant?
A) removing products as they are formed
B) adding more of a reactant
C) adding a catalyst
D) increasing the pressure, in a gas-phase reaction
E) change in temperature
Q2) When a reaction system reaches equilibrium, the forward and reverse reactions stop.
A)True
B)False
Q3) A chemical reaction has an equilibrium constant of 2 × 10<sup>6</sup>. If this reaction is at equilibrium, select the one correct conclusion that can be made about the reaction.
A) The forward and back reactions have stopped.
B) The limiting reactant has been used up.
C) The forward and reverse rate constants are equal.
D) The forward and reverse reaction rates are equal.
E) None of the above conclusions is correct.
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Chapter 18: Acid-Base Equilibria
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Q1) Define an acid according to the Arrhenius theory, and write a balanced equation to support this definition.
Q2) Hydrated metal ions in aqueous solution can act as Brønsted-Lowry acids.
A)True
B)False
Q3) Which, if any, of the following acids is strong?
A) phosphoric
B) carbonic
C) acetic
D) water
E) none of the above
Q4) The substance H<sub>2</sub>SO<sub>3 </sub>is considered
A) a weak Arrhenius base.
B) a strong Arrhenius acid.
C) a strong Arrhenius base.
D) a neutral compound.
E) a weak Arrhenius acid.
Q5) All Brønsted-Lowry bases have at least one lone pair of electrons.
A)True
B)False
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Chapter 19: Ionic Equilibria in Aqueous Systems
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Q1) A popular buffer solution consists of carbonate (CO<sub>3</sub><sup>2-</sup>) and hydrogen carbonate (HCO<sub>3</sub><sup>-</sup>) conjugate acid-base pair.
Which, if any, of the following such buffers has the highest buffer capacity?
A) 0.9 M CO<sub>3</sub><sup>2-</sup> and 0.1 M HCO<sub>3</sub><sup>-</sup>
B) 0.1 M CO<sub>3</sub><sup>2-</sup> and 0.9 M HCO<sub>3</sub><sup>-</sup>
C) 0.5 M CO<sub>3</sub><sup>2-</sup> and 0.5 M HCO<sub>3</sub><sup>-</sup>
D) 0.1 M CO<sub>3</sub><sup>2-</sup> and 0.1 M HCO<sub>3</sub><sup>-</sup>
E) They all have the same buffer capacity.
Q2) When a strong acid is titrated with a weak base, the pH at the equivalence point
A) is greater than 7.0.
B) is equal to 7.0.
C) is less than 7.0.
D) is equal to the pK<sub>a</sub> of the acid.
E) is equal to the pK<sub>b</sub> of the base.
Q3) Which of the following substances has the greatest solubility in water?
A) MgCO<sub>3</sub>, K<sub>sp</sub> = 3.5 × 10<sup>-8</sup>
B) NiCO<sub>3</sub>, K<sub>sp</sub> = 1.3 × 10<sup>-7</sup>
C) AgIO<sub>3</sub>, K<sub>sp</sub> = 3.1 × 10<sup>-8</sup>
D) CuBr, K<sub>sp</sub> = 5.0 × 10<sup>-9</sup>
E) AgCN, K<sub>sp</sub> = 2.2 × 10<sup>-16</sup>
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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions
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Q1) The second law of thermodynamics tells us that
A) the entropy of the universe is constant.
B) entropy is neither created nor destroyed.
C) the universe proceeds toward a state of lower entropy.
D) the universe proceeds toward a state of higher entropy.
E) the universe cannot create entropy.
Q2) A chemical reaction has \(\Delta\)G° = 10.0 kJ and \(\Delta\)S° = 50.0 J/K
a. Calculate \(\Delta\)H°<sup> </sup>for this reaction at 25°C.
b. Could this reaction ever be spontaneous? Explain your answer.
Q3) A reaction has a positive value of \(\Delta\)H° and a positive value of \(\Delta\)S°.
Draw a neat, labeled schematic plot to show how \(\Delta\)G° (y-axis) will depend on absolute temperature (x-axis).
Q4) For what signs of \(\Delta\)H and \(\Delta\)S will a process
a. be spontaneous at high temperatures but not at low temperatures?
b. not be spontaneous at any temperatures?
Q5) a. Explain what is meant by a spontaneous process.
b. Is a spontaneous process necessarily a rapid one? Explain, and provide a real reaction as an example to illustrate your answer.
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Chapter 21: Electrochemistry: Chemical Change and Electrical Work
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Q1) Consider the reaction in the lead-acid cell Pb(s) + PbO<sub>2</sub>(s) + 2H<sub>2</sub>SO<sub>4</sub>(aq) \(\to\) 2PbSO<sub>4</sub>(aq) + 2H<sub>2</sub>O(l)
For which E°<sub>cell</sub> = 2.04 V at 298 K. \(\Delta\)G° for this reaction is
A) -3.94 × 10<sup>5</sup> kJ.
B) -3.94 × 10<sup>2</sup> kJ.
C) -1.97 × 10<sup>5</sup> kJ.
D) -7.87 × 10<sup>2</sup> kJ.
E) none of the above.
Q2) A concentration cell consists of two Zn/Zn<sup>2+</sup> electrodes. The electrolyte in compartment A is 0.10 M Zn(NO<sub>3</sub>)<sub>2 </sub>and in compartment B is 0.60 M Zn(NO<sub>3</sub>)<sub>2</sub>. 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
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Chapter 22: The Elements in Nature and Industry
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Q1) In the commercial production of aluminum metal, the final stage involves the reaction 2Al<sub>2</sub>O<sub>3</sub>(in molten Na<sub>3</sub>AlF<sub>6</sub>) + 3C(graphite) \(\to\) 4Al(l) + 3CO<sub>2</sub>(g)
Briefly outline the method and equipment with which the reaction is accomplished.
Q2) Sulfide ores are frequently treated by flotation in order to concentrate them.
A)True
B)False
Q3) Which of the following reactions is a good small-scale laboratory method for preparation of hydrogen?
A) CH<sub>4</sub>(g) + H<sub>2</sub>O(g) \(\to\) CO(g) + 3H<sub>2</sub>(g)
B) 3FeCl<sub>2</sub>(s) + 4H<sub>2</sub>O(g) \(\to\)Fe<sub>3</sub>O<sub>4</sub>(s) + 6HCl(g) + H<sub>2</sub>(g)
C) NaH(s) + H<sub>2</sub>O(l) \(\to\) NaOH(aq) + H<sub>2</sub>(g)
D) CO(g) + H<sub>2</sub>O (g) \(\to\) CO<sub>2</sub>(g) + H<sub>2</sub>(g)
E) None of the above methods is useful.
Q4) Describe, and give reasons for, the fundamental bonding differences between transition metal sulfides and transition metal oxides.
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Chapter 23: The Transition Elements and Their Coordination Compounds
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Q1) Write the formula for pentaamminechlorocobalt(III) chloride.
A) [Co(NH<sub>3</sub>)<sub>5</sub>Cl]Cl
B) [Co(NH<sub>3</sub>)<sub>5</sub>Cl]Cl<sub>2</sub>
C) [Co(NH<sub>3</sub>)<sub>5</sub>Cl]Cl<sub>3</sub>
D) [Co(NH<sub>3</sub>)<sub>5</sub>Cl]Cl<sub>4</sub>
E) none of the above
Q2) In complexes of transition metals, the maximum coordination number of the metal is equal to its number of d electrons.
A)True
B)False
Q3) Octahedral complexes can exhibit geometric, optical, and linkage isomerism.
A)True
B)False
Q4) Which of the following coordination numbers applies to octahedral complexes?
A) 4
B) 5
C) 6
D) 8
E) none of the above
Page 25
Q5) What geometry is particularly common for complexes of d<sup>10</sup> metal ions?
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Chapter 24: Nuclear Reactions and Their Applications
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94 Verified Questions
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Sample Questions
Q1) Which of the following types of radioactive decay does not produce new element?
A) gamma emission
B) electron capture
C) beta emission
D) alpha emission
E) double beta emission
Q2) No alpha decay is observed for isotopes of elements with Z < 83.
A)True
B)False
Q3) Gamma rays are high energy electrons.
A)True
B)False
Q4) Bombardment of uranium-238 nuclei by carbon-12 nuclei produces californium-246 and neutrons. Write a complete, balanced equation for this nuclear process.
Q5) An isotope with a low value of N/Z will generally decay through A) .\(\alpha\) decay.
B) .\(\beta\) decay.
C) .\(\gamma\) decay.
D) electron capture.
E) spontaneous fission.
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