Course Introduction
Chemical Principles
Test Questions
Chemical Principles provides a foundational understanding of the core concepts in chemistry, including atomic and molecular structure, stoichiometry, chemical bonding, thermodynamics, kinetics, and equilibrium. The course emphasizes the quantitative and qualitative aspects of chemical reactions, the periodic properties of elements, and the relationship between molecular structure and chemical behavior. Through lectures, laboratory experiments, and problem-solving exercises, students develop critical skills and scientific reasoning necessary for further study in chemistry, biology, engineering, or related fields.
Recommended Textbook
Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg
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Page 2
Chapter 1: Keys to the Study of Chemistry
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Sample Questions
Q1) The ripening of fruit, once picked, is an example of physical change.
A)True
B)False
Answer: False
Q2) The number 6.0448, rounded to 3 decimal places, becomes 6.045.
A)True
B)False
Answer: True
Q3) A large pizza has a diameter of 15 inches. Express this diameter in centimeters.
A) 38 cm
B) 24 cm
C) 18 cm
D) 9.3 cm
E) 5.9 cm
Answer: A
Q4) When applying the scientific method, a model or theory should be based on experimental data.
A)True
B)False
Answer: True
Page 3
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Chapter 2: The Components of Matter
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Sample Questions
Q1) Which one of the following combinations of names and formulas is incorrect?
A) H<sub>3</sub>PO<sub>4</sub> phosphoric acid
B) HNO<sub>3</sub> nitric acid
C) NaHCO<sub>3</sub> sodium carbonate
D) H<sub>2</sub>CO<sub>3</sub> carbonic acid
E) KOH potassium hydroxide
Answer: C
Q2) Ionic compounds may carry a net positive or negative charge.
A)True
B)False
Answer: False
Q3) State the two important experimental results (and the names of the responsible scientists) which enabled the mass of the electron to be determined. Answer: Thomson measured m/e, the mass-to-charge ratio. Millikan measured e, the charge. Thus, the mass m could be calculated.
Q4) Modern studies have shown that the Law of Multiple Proportions is not valid.
A)True
B)False Answer: False
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Chapter 3: Stoichiometry of Formulas and Equations
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Sample Questions
Q1) Calculate the number of oxygen atoms in 29.34 g of sodium sulfate, Na<sub>2</sub>SO<sub>4</sub>.
A) 1.244 × 10<sup>23</sup> O atoms
B) 4.976 × 10<sup>23</sup> O atoms
C) 2.409 × 10<sup>24</sup> O atoms
D) 2.915 × 10<sup>24</sup> O atoms
E) 1.166 × 10<sup>25</sup> O atoms
Answer: B
Q2) In a correctly balanced equation, the number of reactant molecules must equal the number of product molecules.
A)True
B)False
Answer: False
Q3) Gaseous methanol (CH<sub>4</sub>O) reacts with oxygen gas to produce carbon dioxide gas and liquid water. Write a balanced equation for this process.
Answer: 2CH<sub>4</sub>O(g) + 3O<sub>2</sub>(g) \(\to\) 2CO<sub>2</sub>(g) + 4H<sub>2</sub>O(l)
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Page 5
Chapter 4: The Major Classes of Chemical Reactions
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Sample Questions
Q1) Which of the following is a strong acid?
A) H<sub>3</sub>PO<sub>4</sub>
B) HNO<sub>3</sub>
C) HF
D) CH<sub>3</sub>COOH
E) H<sub>2</sub>O
Q2) Which one of the following ionic compounds is insoluble in water?
A) Na<sub>3</sub>PO<sub>4</sub>
B) AgNO<sub>3</sub>
C) NaCl
D) CaCO<sub>3</sub>
E) MgCl<sub>2</sub>
Q3) The oxidation numbers of P, S, and Cl in H<sub>2</sub>PO<sub>2</sub><sup>-</sup>, H<sub>2</sub>S, and KClO<sub>4</sub> are, respectively
A) -1, -1, +3.
B) +1, -2, +7.
C) +1, +2, +7.
D) -1, -2, +7.
E) -1, -2, +3.
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Chapter 5: Gases and the Kinetic-Molecular Theory
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Sample Questions
Q1) For real gases, PV < nRT, always.
A)True
B)False
Q2) If 0.750 L of argon at 1.50 atm and 177°C and 0.235 L of sulfur dioxide at 95.0 kPa and 63.0°C are added to a 1.00-L flask and the flask's temperature is adjusted to 25.0°C, what is the resulting pressure in the flask?
A) 0.0851 atm
B) 0.244 atm
C) 0.946 atm
D) 1.74 atm
E) 1.86 atm
Q3) For a pure gas sample, the average kinetic energy is also the most probable kinetic energy.
A)True
B)False
Q4) Calculate the density in g/L of gaseous SF<sub>6</sub> at 50.0°C and 650. torr.
Q5) For real gases, PV > nRT, always.
A)True
B)False
Q6) State Avogadro's Law.
Page 7
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Chapter 6: Thermochemistry: Energy Flow and Chemical Change
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Sample Questions
Q1) The standard state of a substance in aqueous solution is a 1 M solution. A)True
B)False
Q2) A Snickers<sup>®</sup> candy bar contains 280 Calories, of which the fat content accounts for 120 Calories. What is the energy of the fat content, in kJ?
A) 5.0 × 10<sup>-1 </sup>kJ
B) 29 kJ
C) 5.0 × 10<sup>2</sup> kJ
D) 1.2 × 10<sup>3</sup> kJ
E) 5.0 × 10<sup>5</sup> kJ
Q3) a. Write a balanced equation for the combustion of benzene, C<sub>6</sub>H<sub>6</sub>(l) in oxygen.
b. The standard heat of combustion of benzene is -3271 kJ/mol. Calculate its standard heat of formation, \(\Delta\)H°<sub>f</sub> , given the data: \(\Delta\)H°<sub>f</sub> [CO<sub>2</sub>(g)] = -394 kJ; \(\Delta\)H°<sub>f</sub> [H<sub>2</sub>O(l)] = -286 kJ
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Chapter 7: Quantum Theory and Atomic Structure
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Sample Questions
Q1) Who developed an empirical equation from which the wavelengths of lines in the spectrum of hydrogen atoms can be calculated?
A) Planck
B) de Broglie
C) Bohr
D) Rutherford
E) Rydberg
Q2) The Rydberg equation is an example of an empirical equation.
A)True
B)False
Q3) Electromagnetic radiation can be specified by its wavelength (\(\lambda\)), its frequency (v) or its period (\(\tau\)). The period \(\tau\) is the time it takes one complete wavelength to pass a point in space. Based on this information, what is the mathematical relationship between v and \(\tau\)?
A) (v = 1/\(\tau\))
B) (v = \(\tau\))
C) (\(\tau\) = v/c)
D) c = \(\tau\) x v
E) (\(\tau\) = c/v)
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Chapter 8: Electron Configuration and Chemical Periodicity
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Q1) The difference in energies between the 1s and 2s orbitals is due to the penetration effect.
A)True
B)False
Q2) Which of the following electron configurations is impossible?
A) 1s<sup>2</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>2</sup>3p<sup>4</sup>
B) 1s<sup>2</sup>2s<sup>2</sup>2p<sup>5</sup>3s<sup>1</sup>3p<sup>4</sup>
C)
1s<sup>2</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>2</sup>3p<sup>6</sup>3d<sup >10</sup>4s<sup>2</sup>
D) 1s<sup>2</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>3</sup>3p<sup>4</sup>
E) 1s<sup>1</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>2</sup>3p<sup>4</sup>
Q3) Which of the following elements has the largest first ionization energy?
A) Na
B) Cl
C) Ca
D) Te
E) Br
Q4) State Hund's rule, and show how it applies to the ground state of phosphorus atoms.
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Chapter 9: Models of Chemical Bonding
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Sample Questions
Q1) Arrange oxygen, sulfur, calcium, rubidium, and potassium in order of decreasing electronegativity.
A) O > S > Ca > K > Rb
B) O > S > Ca > Rb > K
C) O > S > Rb > K > Ca
D) O > S > Rb > Ca > K
E) None of the above orders is correct.
Q2) In not more than three sentences, describe the electron arrangement responsible for bonding in Cl<sub>2</sub> molecules.
Q3) The electrostatic energy of two charged particles is inversely proportional to the square of the distance between them.
A)True
B)False
Q4) Covalently bonded substances do not necessarily exist as separate molecules.
A)True
B)False
Q5) In not more than three sentences, describe the electron arrangement responsible for bonding in solid SrCl<sub>2</sub>.
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Chapter 10: The Shapes of Molecules
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Sample Questions
Q1) In a Lewis structure for a molecule or ion, the sum of the formal charges on the atoms is equal to the charge on the molecule or ion.
A)True
B)False
Q2) What is the molecular shape of ClCN as predicted by the VSEPR theory? (Carbon is the central atom.)
A) linear
B) bent
C) angular
D) trigonal
E) none of the above
Q3) How many resonance structures are possible for NO<sub>3</sub><sup>-</sup>?
A) 1
B) 2
C) 3
D) 4
E) 5
Q4) Explain what is meant by "dipole moment", and give an example of a molecule which has polar bonds but which does not itself have a dipole moment.
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Chapter 11: Theories of Covalent Bonding
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Sample Questions
Q1) Which one, if any, of the following statements about the MO treatment of the bonding in benzene is correct?
A) MO theory uses resonance to describe the delocalized nature of the bonding in benzene.
B) The lowest energy pi-bonding orbital in benzene can hold up to 6 electrons.
C) MO theory predicts that benzene should be paramagnetic.
D) The lowest energy pi-bonding MO in benzene has two hexagonal lobes, above and below the plane of the carbon atoms.
E) None of the above statements is correct.
Q2) A molecule with the formula AX<sub>4</sub>E uses _________ to form its bonds.
A) sp<sup>2</sup> hybrid orbitals
B) sp<sup>3</sup> hybrid orbitals
C) sp<sup>3</sup>d hybrid orbitals
D) sp<sup>3</sup>d<sup>2</sup> hybrid orbitals
E) none of the above
Q3) Sketch the shapes of the \(\sigma\)<sub>1</sub><sub>s</sub> and \(\sigma\)<sub>1</sub><sub>s</sub>* molecular orbitals formed by the overlap of two hydrogen 1s atomic orbitals.
Q4) In one sentence state the basic principle of valence bond theory.
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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes
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Sample Questions
Q1) Octane is a component of fuel used in internal combustion engines. The dominant intermolecular forces in octane are
A) dipole-dipole forces.
B) London dispersion forces.
C) hydrogen bonding.
D) covalent bonds.
E) carbon-hydrogen bonds.
Q2) Which of the following statements about the packing of monatomic solids with different unit cells is incorrect?
A) The coordination number of atoms in hcp and fcc structures is 12.
B) The coordination number of atoms in simple cubic structures is 6.
C) The coordination number of atoms in bcc structures is 8.
D) A bcc structure has a higher packing efficiency than a simple cubic structure.
E) A bcc structure has a higher packing efficiency than a fcc structure.
Q3) Mercury melts at -39°C and boils at 357°C. Draw a diagram of the heating curve of mercury. Label all lines and axes, and clearly indicate the melting and boiling points on your diagram.
Q4) How do the electrical properties of semiconductors differ from those of metals?
Page 14
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Chapter 13: The Properties of Mixtures: Solutions and Colloids
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Sample Questions
Q1) Calcium nitrite is used as a corrosion inhibitor in lubricants. What is the molality of a solution prepared by dissolving 18.5 g of calcium nitrite in 83.5 g of distilled water?
A) 0.0342 m
B) 0.0855 m
C) 0.222 m
D) 0.444 m
E) 1.68 m
Q2) The Henry's Law constant (k) for carbon monoxide in water at 25°C is 9.71 × 10<sup>-</sup><sup>4</sup> mol/(L·atm). How many grams of CO will dissolve in 1.00 L of water if the partial pressure of CO is 2.75 atm?
A) 3.53 × 10<sup>-4</sup> g
B) 2.67 × 10<sup>-3 </sup>g
C) 9.89 × 10<sup>-3</sup> g
D) 7.48 × 10<sup>-2</sup> g
E) none of the above
Q3) List three important intermolecular forces that operate in DNA, and explain what parts of the molecule are involved in these forces.
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Page 15
Chapter 14: Periodic Patterns in the Main Group Elements:
Bonding, Structure, and Reactivity
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Q1) What are the products of the following reaction of strontium hydride and water? SrH<sub>2</sub>(s) + H<sub>2</sub>O(l) \(\to\)
A) Sr<sup>2+</sup>(aq) + H<sub>2</sub>(g) + O<sub>2</sub>(g)
B) Sr<sup>2+</sup>(aq) + H<sub>2</sub>(g) + OH<sup>-</sup>(aq)
C) Sr(s) + H<sub>2</sub>(g) + OH<sup>-</sup>(aq)
D) Sr(s) + H<sub>3</sub>O<sup>+</sup>(aq)
E) none of the above
Q2) The nitrate anion is
A) a strong acid.
B) a strong base.
C) amphoteric.
D) a strong reducing agent.
E) a strong oxidizing agent.
Q3) Xenon forms several compounds with oxygen and fluorine. It is the most reactive non-radioactive noble gas because
A) its large radius allows oxygen and fluorine to bond without being crowded.
B) it has the highest electronegativity of these noble gases.
C) it has the highest electron affinity of these noble gases.
D) its effective nuclear charge is lower than the other noble gases.
Page 16
E) it has the lowest ionization energy of these noble gases.
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Chapter 15: Organic Compounds and the Atomic Properties of Carbon
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Q1) Dacron is a ___________ compound and is produced by a __________ reaction.
A) polyamide; addition
B) polyamide; condensation
C) polyester; addition
D) polyester; condensation
E) polyene; addition
Q2) The most abundant organic chemical on earth is
A) glycogen.
B) starch.
C) cellulose.
D) glucose.
E) carbonate.
Q3) Nylon-type compounds are prepared by the reaction of
A) a diol with a diacid.
B) an alcohol with a diacid.
C) a diamine with a diacid.
D) an amine with a diacid.
E) a diamine with an acid.
Q5) Explain how a disulfide bridge can arise a protein molecule. Page 17
Q4) Draw and name all stable molecules with the formula C<sub>5</sub>H<sub>12</sub>.
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Page 18
Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions
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Sample Questions
Q1) If the activation energy of a reaction decreases by 10.0 kJ/mol, from 100.0 to 90.0 kJ/mol, what effect will this have on the rate of reaction at 298K?
A) The rate will increase, by a factor of more than 50.
B) The rate will decrease, by a factor of more than 50.
C) The rate will increase, by a factor of less than 50.
D) The rate will decrease, by a factor of less than 50.
E) The rate will not change unless temperature changes.
Q2) The rate constant for a reaction is 4.65 L mol<sup>-1</sup>s<sup>-1</sup>. What is the overall order of the reaction?
A) zero
B) first
C) second
D) third
E) More information is needed to determine the overall order.
Q3) In a reversible reaction, a catalyst will speed up the forward reaction but not affect the reverse reaction.
A)True
B)False
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Chapter 17: Equilibrium: the Extent of Chemical Reactions
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Q1) The reaction quotient, Q<sub>c</sub>, for a reaction has a value of 75 while the equilibrium constant, K<sub>c</sub>, has a value of 195. Which of the following statements is accurate?
A) The reaction must proceed to the left to establish equilibrium.
B) The reaction must proceed to the right to establish equilibrium.
C) The concentrations of the products will be much smaller than the concentrations of the reactants when the system is at equilibrium.
D) The concentrations of the products will be about the same as the concentrations of the reactants when the system is at equilibrium.
E) None of the above statements is accurate.
Q2) Unless \(\Delta\)H°<sub>rxn</sub> = 0, a change in temperature will affect the value of the equilibrium constant K<sub>c</sub>.
A)True
B)False
Q3) A good catalyst for a reaction will speed up the forward reaction and slow down the reverse reaction.
A)True
B)False
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Chapter 18: Acid-Base Equilibria
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Q1) Which one of the following substances will give an aqueous solution of pH < 7?
A) KI
B) NH<sub>4</sub>Br
C) Na<sub>2</sub>CO<sub>3</sub>
D) CH<sub>3</sub>COONa
E) CH<sub>3</sub>OH
Q2) Define an acid according to the Arrhenius theory, and write a balanced equation to support this definition.
Q3) Picric acid has been used in the leather industry and in etching copper. However, its laboratory use has been restricted because it dehydrates on standing and can become shock sensitive. It has an acid dissociation constant of 0.42. What is the [H<sub>3</sub>O<sup>+</sup>] for a 0.20 M solution of picric acid?
A) 0.022 M
B) 0.052 M
C) 0.15 M
D) 0.29 M
E) none of the above
Q4) Describe what is meant by the "leveling effect". Use a real acid as an example, and write an appropriate equation.
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Chapter 19: Ionic Equilibria in Aqueous Systems
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Q1) At the equivalence point in an acid-base titration
A) the [H<sub>3</sub>O<sup>+</sup>] equals the K<sub>a</sub> of the acid.
B) the [H<sub>3</sub>O<sup>+</sup>] equals the K<sub>a</sub> of the indicator.
C) the amounts of acid and base which have been combined are in their stoichiometric ratio.
D) the pH is 7.0.
E) the pH has reached a maximum.
Q2) A 20.0-mL sample of 0.30 M HBr is titrated with 0.15 M NaOH. What is the pH of the solution after 40.3 mL of NaOH have been added to the acid?
A) 2.95
B) 3.13
Q3) Make a clear distinction between buffer range and buffer capacity.
Q4) Use a carefully drawn and labeled diagram of the titration curve to illustrate the titration of a weak diprotic acid, in which K<sub>a1</sub> and K<sub>a2</sub> are substantially different, with a strong base (base is the titrant). Label as many features of the diagram as possible.
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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction
of Chemical Reactions
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Q1) A certain process has \(\Delta\)H° > 0, \(\Delta\)S° < 0, and \(\Delta\)G° > 0. The values of \(\Delta\)H° and \(\Delta\)S° do not depend on the temperature. Which of the following is a correct conclusion about this process?
A) It is non-spontaneous at all T.
B) It is spontaneous at high T.
C) It is spontaneous at low T.
D) It is spontaneous at all T.
E) None of the above conclusions is correct.
Q2) Which one of the following phase changes decreases the entropy of the system?
A) melting
B) heating a gas
C) vaporization
D) condensation
E) sublimation
Q3) Which of the following is always true for an exothermic process?
A) q<sub> sys</sub> > 0, \(\Delta\)S<sub>surr</sub> < 0
B) q<sub> sys</sub> < 0, \(\Delta\)S<sub>surr</sub> > 0
C) q<sub> sys</sub> < 0, \(\Delta\)S<sub>surr</sub> < 0
D) q<sub> sys</sub> > 0, \(\Delta\)S<sub>surr</sub> > 0
E) w < 0
Page 23
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Chapter 21: Electrochemistry: Chemical Change and Electrical Work
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Q1) A battery that cannot be recharged is a
A) fuel cell.
B) primary battery.
C) secondary battery.
D) simple battery.
E) flow battery.
Q2) Which of the following elements could be prepared by electrolysis of the aqueous solution shown?
A) sodium from Na<sub>3</sub>PO<sub>4</sub>(aq)
B) sulfur from K<sub>2</sub>SO<sub>4</sub>(aq)
C) oxygen from H<sub>2</sub>SO<sub>4</sub>(aq)
D) potassium from KCl(aq)
E) nitrogen from AgNO<sub>3</sub>(aq)
Q3) A salt bridge provides a path for electrons to move between the anode and cathode compartments of a voltaic cell.
A)True
B)False
Q4) A primary battery is one which can be recharged.
A)True
B)False
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Chapter 22: The Elements in Nature and Industry
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Q1) What gas is produced at the anode in the Hall-Heroult process for the electrolytic production of aluminum?
Q2) Electrometallurgy uses _________________ to separate a metal from its ore.
A) solid phase chemical properties
B) electrical processes
C) thermal processes
D) aqueous chemical processes
E) molten salt processes
Q3) The debris accompanying a mineral is called A) slag.
B) gangue.
C) ore.
D) halite.
E) ash.
Q4) The leaching of gold can be represented by the equation 4Au(s) + O<sub>2</sub>(s) + 8CN<sup>-</sup>(aq) + 2H<sub>2</sub>O(l) \(\to\)4Au(CN)<sub>2</sub><sup>-</sup>(aq) + 4OH<sup>-</sup>(aq)
In this reaction, identify the oxidant, the reductant and the oxidation numbers of the elements which change.
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Chapter 23: The Transition Elements and Their Coordination Compounds
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Q1) The Cu<sup>2+</sup> ion has 1 unpaired electron.
A)True
B)False
Q2) The permanganate ion (MnO<sub>4</sub><sup>-</sup>) is a powerful reducing agent.
A)True
B)False
Q3) In a coordination compound involving a complex ion of square planar geometry, which of the following types of isomerism is/are never possible?
A) geometric
B) optical
C) linkage
D) coordination
E) more than one of the above
Q4) Which of the following will be diamagnetic?
A) Ni<sup>2+</sup>
B) Cr<sup>2+</sup>
C) Mn<sup>2+</sup>
D) Co<sup>3+</sup>
E) Ti<sup>4+</sup>
Q5) Why is the +2 oxidation state so common among transition elements? Page 26
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Chapter 24: Nuclear Reactions and Their Applications
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Q1) Which one of the following elements is formed largely in supernova explosions?
A) H
B) He
C) Mg
D) Fe
E) U
Q2) Which of the following series of radioactive decays would convert Pa-234 to Ra-226?
A) beta, alpha, beta
B) alpha, alpha
C) beta, alpha, alpha, beta
D) beta, alpha, alpha
E) alpha, beta, gamma
Q3) Which of the following materials is put into a nuclear reactor to slow the chain reaction?
A) heavy water
B) moderators
C) control rods
D) reflectors
E) chlorine
To view all questions and flashcards with answers, click on the resource link above.
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