Preparatory Chemistry Exam Answer Key - 2309 Verified Questions

Course Introduction

Preparatory Chemistry

Exam Answer Key

Preparatory Chemistry is designed to provide students with a foundational understanding of basic chemical concepts and principles essential for success in general chemistry courses. Covering topics such as atomic structure, chemical bonding, stoichiometry, states of matter, and basic thermochemistry, this course emphasizes problem-solving skills and the application of scientific reasoning. Through a combination of lectures, demonstrations, and hands-on laboratory activities, students will develop the confidence and competence needed to progress in more advanced chemistry studies. This course is ideal for those with limited prior exposure to chemistry or those seeking a refresher before enrolling in higher-level coursework.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) The numerical value of any temperature expressed in Celsius is always different from the numerical value of the same temperature in Fahrenheit.

A)True

B)False

Answer: False

Q2) Acetone, which is used as a solvent and as a reactant in the manufacture of Plexiglas^®, boils at 56.1°C. What is the boiling point in degrees Fahrenheit?

A) 159°F

B) 133°F

C) 101°F

D) 69.0°F

E) 43.4°F

Answer: B

Q3) An evacuated 276 mL glass bulb weighs 129.6375 g. Filled with an unknown gas, the bulb weighs 130.0318 g. Calculate the gas density in g/L, and express it with an appropriate number of significant figures.

Answer: 1.43 g/L

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3

Chapter 2: The Components of Matter

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Sample Questions

Q1) Which of the following compounds is covalent?

A) CaCl₂

B) MgO

C) Al₂O₃

D) Cs₂S

E) PCl₃

Answer: E

Q2) Modern studies have shown that the Law of Multiple Proportions is not valid.

A)True

B)False

Answer: False

Q3) The compound, P₄S₁₀, is used in the manufacture of safety matches. What is its name?

A) phosphorus sulfide

B) phosphoric sulfide

C) phosphorus decasulfide

D) tetraphosphorus decasulfide

E) phosphorus pentasulfide

Answer: D

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Q1) Terephthalic acid, used in the production of polyester fibers and films, is composed of carbon, hydrogen, and oxygen. When 0.6943 g of terephthalic acid was subjected to combustion analysis it produced 1.471 g CO₂and 0.226 g H₂O. If its molar mass is between 158 and 167 g/mol, what is its molecular formula?

A) C₄H₆O₇

B) C₆H₈O₅

C) C₇H₁₂O₄

D) C₄H₃O₂

E) C₈H₆O₄

Answer: E

Q2) Aluminum oxide, Al₂O₃, is used as a filler for paints and varnishes as well as in the manufacture of electrical insulators. Calculate the number of moles in 47.51 g of Al₂O₃.

A) 2.377 mol

B) 2.146 mol

C) 1.105 mol

D) 0.4660 mol

E) 0.4207 mol

Answer: D

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Page 5

Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) A particular reaction may be both a precipitation and an acid-base (neutralization) reaction.

A)True

B)False

Q2) a. Explain or define what is meant by the term "oxidation."

b. Write down the oxidation numbers of all the atoms in the following formulas.

K₂Cr₂O₇NaH

Q3) The oxidation numbers of P, S, and Cl in H₂PO₂^-, H₂S, and KClO₄ are, respectively

A) -1, -1, +3.

B) +1, -2, +7.

C) +1, +2, +7.

D) -1, -2, +7.

E) -1, -2, +3.

Q4) Which one of the following substances is the best electrolyte?

A) CO

B) CH₃Cl

C) CH₄

D) C₂H₅OH

E) HCl

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Q1) For an ideal gas, a plot of PV/nRT versus P gives a straight line with a positive slope.

A)True

B)False

Q2) A weather balloon was initially at a pressure of 0.950 atm, and its volume was 35.0 L. The pressure decreased to 0.750 atm, without loss of gas or change in temperature. What was the change in the volume?

A) increased by 44.3 L

B) increased by 9.3 L

C) increased by 7.4 L

D) decreased by 27.6 L

E) decreased by 7.4 L

Q3) For a gas obeying Boyle's Law, a plot of V versus 1/P will give a straight line passing through the origin.

A)True

B)False

Q4) State Charles's Law and illustrate it with a graph, using standard x-y coordinate axes. Be sure to label the axes unambiguously with the correct gas variables.

Q5) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650. torr.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) What is the final temperature when 20.0 g of water at 25°C is mixed with 30.0 g of water at 80°C?

A) 35°C

B) 42°C

C) 53°C

D) 58°C

E) 70°C

Q2) Two solutions (the system), each of 25.0 mL volume and at 25.0°C, are mixed in a beaker. A reaction occurs between them, and the temperature rises to 35.0°C. After the products have equilibrated with the surroundings, the temperature is again 25.0°C and the total volume is 50.0 mL. No gases are involved in the reaction. Which one of the following relationships concerning the change from initial to final states (both at 25.0°C) is correct?

A) .\(\Delta\)E = 0

B) .\(\Delta\)H = 0

C) .\(\Delta\)E > 0

D) q = 0

E) w = 0

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Chapter 7: Quantum Theory and Atomic Structure

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Q1) Which one of the following sets of quantum numbers can correctly represent a 3p orbital?

A) n = 3 l = 1 m_l = 2

B) n = 1 l = 3 m_l = 3

C) n = 3 l = 2 m_l = 1

D) n = 3 l = 1 m_l = -1

E) n = 3 l = 0 m_l = 1

Q2) Who was the first scientist to propose that the atom had a dense nucleus which occupied only a small fraction of the volume of the atom?

A) Planck

B) Bohr

C) Rydberg

D) Rutherford

E) Thomson

Q3) The shape of an atomic orbital is associated with

A) the principal quantum number (n).

B) the angular momentum quantum number (l).

C) the magnetic quantum number (m_l).

D) the spin quantum number (m_s).

E) the magnetic and spin quantum numbers, together.

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) According to the exclusion principle, two is the maximum number of electrons in an atom which can share the same four quantum numbers.

A)True

B)False

Q2) Describe how the atomic radii of main group elements vary with position in the periodic table.

Q3) Select the correct electron configuration for Cu (Z = 29).

A) [Ar]4s²3d⁹

B) [Ar]4s¹3d¹⁰

C) [Ar]4s²4p⁶3d³

D) [Ar]4s²4d⁹

E) [Ar]5s²4d⁹

Q4) Select the diamagnetic ion.

A) Cu²⁺

B) Ni²⁺

C) Cr³⁺

D) Sc³⁺

E) Cr²⁺

Q5) 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) Most of the copper sold in major metal markets is highly purified, typically to 99.99%. Why is this?

Q2) Which of the following is a covalent compound?

A) Na₂O

B) CaCl₂

C) Cl₂O

D) CsCl

E) Al₂O₃

Q3) Which of the following contains ionic bonding?

A) CO

B) SrF₂

C) Al

D) OCl₂

E) HCl

Q4) Select the correct formula for a compound formed from calcium and chlorine.

A) CaCl

B) CaCl₂

C) Ca₂Cl

D) Ca₂Cl₂

E) CaCl₃

Page 11

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the actual bond angles in SF₃⁺ 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°

Q2) For the chlorate ion, ClO₃^-, draw two different valid Lewis structures, as follows:

a. a structure in which the octet rule is obeyed

b. a structure in which formal charges are minimized

Q3) In the COCl₂ molecule, carbon is the central atom. Based on the best Lewis structure for COCl₂, what is the formal charge on carbon?

A) 0

B) +1

C) -1

D) +2

E) -2

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Chapter 11: Theories of Covalent Bonding

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Q1) In not more than two sentences, explain when and why chemists make use of the concept of hybridization.

Q2) In applying molecular orbital theory, the bond order is calculated in the same way as with Lewis structures.

A)True

B)False

Q3) Sketch the shapes of the \(\sigma\)₁_s and \(\sigma\)₁_s* molecular orbitals formed by the overlap of two hydrogen 1s atomic orbitals.

Q4) A carbon-carbon triple bond in a molecule may give rise to the existence of cis and trans isomers.

A)True B)False

Q5) Explain what is meant by a node (or nodal plane) in a molecular orbital and draw sketches of the following orbitals, indicating at least one nodal plane in each one. a. a \(\sigma\)* orbital b. a \(\pi\)* orbital

Q6) 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) Liquid sodium can be used as a heat transfer fluid. Its vapor pressure is 40.0 torr at 633°C and 400.0 torr at 823°C. Calculate its heat of vaporization.

A) 43.4 kJ/mol

B) 52.5 kJ/mol

C) 70.6 kJ/mol

D) 1.00 × 10² kJ/mol

E) none of the above

Q2) List the three common classes of liquid crystals in order of decreasing degree of order.

Q3) 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.

Q4) How do the electrical properties of semiconductors differ from those of metals?

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Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) Which of the following pairs of ions is arranged so that the ion with the larger (i.e., more negative) heat of hydration is listed first?

A) Br^-, K⁺

B) Mg²⁺, Sr ²⁺

C) Ca²⁺, Sc³⁺

D) Na⁺, Li⁺

E) None of the above are arranged in this way.

Q2) The vapor pressure of pure acetone (propanone) is 266 torr. When a non-volatile solute is added, the vapor pressure of acetone above the solution falls to 232 torr. What is the mole fraction of the non-volatile solute in the solution?

A) 0.87

B) 0.69

C) 0.32

D) 0.13

E) 0.045

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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Sample Questions

Q1) Boron has 3 valence electrons. Which of the following processes is involved in boron's achieving a complete outer shell?

A) formation of a B³⁺ cation

B) formation of bridge bonds

C) formation of \(\pi\)-bonds using its d-orbitals

D) formation of \(\pi\)-bonds using sp³-orbitals

E) none of the above

Q2) Predict the products for the following set of reactants. Ca₃As₂(s) + H₂O(l) \(\to\)

A) As³⁺(aq) + Ca(OH)₂(aq)

B) As(OH)₃(s) + Ca²⁺(aq) + H₂(g)

C) As(OH)₃(s) + Ca(OH)₂(aq)

D) AsH₃(g) + Ca(OH)₂(aq)

E) CaH₂(aq) + As₂O₃(aq)

Q3) Name the three different classes (types) of hydride, and list some of their important characteristics.

Q4) The halogens act as oxidizing agents in most of their reactions.

A)True

B)False

Page 16

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Chapter 15: Organic Compounds and the Atomic Properties

of Carbon

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Q1) As a result of delocalization of the pi-electrons, all the carbon-carbon bonds in benzene are identical.

A)True

B)False

Q2) Because organic compounds are associated with living organisms, their properties are governed by fundamentally different chemical principles from those of inorganic compounds.

A)True

B)False

Q3) Because of delocalization, all the carbon-carbon bond lengths in 1,3-cyclopentadiene are identical.

A)True

B)False

Q4) Helical and sheet-like segments in proteins arise from

A) disulfide bridges.

B) salt bridges.

C) crosslinking via covalent bonds.

D) dispersion forces within the protein's interior.

E) hydrogen bonding.

Q5) Draw the structure of trans-5-methyl-3-heptene.

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Q1) In the collision theory of reaction rates, the rate constant for a bimolecular reaction can be written as k = z.p.exp(-E_a/RT)

In one sentence each, clearly explain the physical meaning (interpretation) of the following three factors which appear in the above expression: a. z

b. p c. exp(-E_a/RT)

Q2) Butadiene, C₄H₆ (used to make synthetic rubber and latex paints) reacts to C₈H₁₂ with a rate law of rate = 0.014 L/(mol·s) [C₄H₆]². What will be the concentration of C₄H₆after 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

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Page 18

Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) Hydrogen iodide, HI, is formed in an equilibrium reaction when gaseous hydrogen and iodine gas are heated together. If 20.0 g of hydrogen and 20.0 g of iodine are heated, forming 10.0 g of hydrogen iodide, what mass of hydrogen remains unreacted?

A) 10.0 g hydrogen remains

B) 10.9 g hydrogen remains

C) 15.0 g hydrogen remains

D) 19.9 g hydrogen remains

E) Need to know the equilibrium constant in order to calculate the answer.

Q2) 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

Q3) Once a reaction system reaches equilibrium, the concentrations of reactions and products no longer change.

A)True

B)False

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19

Chapter 18: Acid-Base Equilibria

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Q1) A solution is prepared by adding 0.10 mol of sodium sulfide, Na₂S, to 1.00

L of water. Which statement about the solution is correct?

A) The solution is basic.

B) The solution is neutral.

C) The solution is acidic.

D) The concentration of sodium ions and sulfide ions will be identical.

E) The concentration of sulfide ions will be greater than the concentration of sodium ions.

Q2) Which of the following pairs has the stronger acid listed first?

A) HBr, HI

B) HClO₂, HClO₃

C) H₂SeO₄, H₂SeO₃

D) HNO₂, HNO₃

E) HF, HCl

Q3) Which of the following aqueous systems has the highest pH?

A) 0.1 M HA, pK_a = 11.89

B) 0.1 M HMO, pK_a = 8.23

C) 0.1 M HA, pK_a = 4.55

D) 0.1 M HBO, pK_a = 2.43

E) pure water

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Page 20

Chapter 19: Ionic Equilibria in Aqueous Systems

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Q1) Which of the following substances has the greatest solubility in water?

A) MgCO₃, K_sp = 3.5 × 10^-8

B) NiCO₃, K_sp = 1.3 × 10^-7

C) AgIO₃, K_sp = 3.1 × 10^-8

D) CuBr, K_sp = 5.0 × 10^-9

E) AgCN, K_sp = 2.2 × 10^-16

Q2) At the equivalence point in an acid-base titration

A) the [H₃O⁺] equals the K_a of the acid.

B) the [H₃O⁺] equals the K_a 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.

Q3) The solubility of calcium chromate is 1.56 × 10^-3 g/100 mL of solution. What is the K_sp for CaCrO₄?

A) 2.4 × 10^-4

B) 1.5 × 10^-5

C) 7.6 × 10^-6

D) 1.0 × 10^-8

E) < 1.0 × 10^-8

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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of

Chemical Reactions

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Q1) For a chemical reaction to be non-spontaneous at any temperature, which of the following conditions must be met?

A) .\(\Delta\)S° > 0, \(\Delta\)H° > 0

B) .\(\Delta\)S° > 0, \(\Delta\)H° < 0

C) .\(\Delta\)S° < 0, \(\Delta\)H° < 0

D) .\(\Delta\)S° < 0, \(\Delta\)H° > 0

E) All reactions are spontaneous at some temperature.

Q2) In the expression, S = k ln W, W is called the number of microstates. Explain clearly the meaning of the word "microstate", and why a system under a given set of conditions normally has many microstates.

Q3) Given: C₂H₂(g) \(\to\) 2C(graphite) + H₂(g) \(\Delta\)G° = -209 kJ

A sample of gaseous C₂H₂(acetylene, or ethyne) was stored for one year, yet at the end of this period the sample remained unchanged and no graphite or hydrogen gas had been formed. Briefly explain why there is no inconsistency between the sign of \(\Delta\)G° and the apparent stability of the sample.

Q4) State the second and third laws of thermodynamics.

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) When the following redox equation is balanced with smallest whole number coefficients, the coefficient for nitrogen dioxide will be _____.

I₂(s) + HNO₃(aq) \(\to\) HIO₃(aq) + NO₂(g) + H₂O(l)

A) 1

B) 2

C) 4

D) 10

E) none of the above

Q2) The line notation, Al(s) | Al³⁺(aq) || Co²⁺(aq) | Co(s), indicates that

A) Co is the reducing agent.

B) Co²⁺ ions are oxidized.

C) Al is the reducing agent.

D) Al³⁺ is the reducing agent.

E) aluminum metal is the cathode.

Q3) In the absence of oxygen, iron will rust as long as moisture is present.

A)True

B)False

Page 23

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Chapter 22: The Elements in Nature and Industry

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Q1) What are the main mineral sources of phosphorus?

A) phosphides

B) phosphites

C) phosphors

D) phosphates

E) elemental phosphorus

Q2) The purpose of anodizing aluminum is to remove the oxide layer from the metal's surface.

A)True

B)False

Q3) Sulfide ores are frequently treated by flotation in order to concentrate them.

A)True

B)False

Q4) Compared to the asbestos diaphragm-cell used in the chlor-alkali process, list a. one main advantage and one disadvantage of mercury cells. b. two important advantages of polymer membrane-cells.

Q5) Give two important reasons why there is so much more Na⁺ in the oceans than K⁺, despite their similar abundances in the earth's crust.

Q6) Briefly describe the three main pathways for nitrogen fixation.

Page 24

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) The compound K₃[Fe(CN)₆] is used in calico printing and wool dyeing. Give its systematic name.

A) potassium iron(III) hexacyanate

B) tripotassium iron(III) hexacyanate

C) potassium hexacyanoferrate(III)

D) potassium hexacyanideferrate

E) none of the above

Q2) The crystal field splitting energy, \(\Delta\),

A) is larger for tetrahedral complexes than for octahedral complexes.

B) depends on the metal but not on the ligand.

C) determines the color of a complex.

D) is larger for ionic ligands like chloride than for molecular ligands like carbon monoxide, CO.

E) determines the charge of a complex.

Q3) What is the difference between a coordination compound and a complex ion?

Q4) a. How many unpaired 3d electrons will there be in (i) high and (ii) low-spin complexes of Co(II)?

b. How can high and low-spin complexes be recognized and distinguished experimentally?

Page 25

Q5) What geometry is particularly common for complexes of d¹⁰ metal ions?

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) An isotope with a high value of N/Z will tend to decay through

A) .\(\alpha\) decay.

B) .\(\beta\) decay.

C) positron decay.

D) electron capture.

E) .\(\gamma\) decay.

Q2) Write a complete, balanced equation to represent the alpha decay of radon-210.

Q3) Carbon-14 will emit a \(\beta\) particle with an energy of 0.1565 MeV. What is this energy in joules?

A) 1.0 × 10^-24 J

B) 2.5 × 10^-20 J

C) 1.0 × 10^-18 J

D) 2.5 × 10^-14 J

E) none of the above

Q4) 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.

Page 27

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