General Chemistry I Pre-Test Questions - 2309 Verified Questions

General Chemistry I Pre-Test Questions

Course Introduction

General Chemistry I is an introductory course that explores the fundamental principles of chemistry, including atomic structure, chemical bonding, stoichiometry, states of matter, thermochemistry, and the properties of solutions. The course emphasizes the development of problem-solving skills and laboratory techniques, providing a foundation for understanding chemical reactions, molecular interactions, and the behavior of matter at the macroscopic and microscopic levels. Through lectures, discussions, and hands-on experiments, students gain a comprehensive understanding of the core concepts and applications of chemistry in scientific and everyday contexts.

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) In each of the sets below, choose the one quantity or number which is exact.

a. i. the human population

ii. the distance in light years from the sun to Alpha Centauri, a nearby star

iii. the winning time for the 100 m dash in the Olympic Games

b. i. the weight of a particular one cent coin in g

ii. the boiling point of lead, in °C

iii. the number of cm in 1 yd

c. i. the measured value of the speed of light (2.998... × 10⁸ m/s)

ii. \(\pi\) (3.141...)

iii. the volume of milk in a 1-gallon jug

Answer: a. i

b. iii

c. ii

Q2) 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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Chapter 2: The Components of Matter

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Sample Questions

Q1) Which of the following ions occurs commonly?

A) N³⁺

B) S⁶⁺

C) O^2-

D) Ca⁺

E) Cl⁺

Answer: C

Q2) a. Give the names of the following ions: (i) O₂^2-

(ii) SO₄^2-

b. Write down the formulas of the following ions: (i) ammonium

(ii) nitrate

Answer: a. (i) peroxide (ii) sulfate

b. (i) NH₄⁺ (ii) NO₃^-

Q3) Name the three important "laws" that were accounted for by Dalton's atomic theory.

Answer: laws of conservation of mass; definite composition; multiple proportions

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Constitutional (structural) isomers have the same molecular formula but different structural formulas.

A)True

B)False

Answer: True

Q2) In combustion analysis, the carbon and hydrogen contents of a substance are determined from the CO₂ and H₂O, respectively, which are collected in the absorbers.

A)True

B)False

Answer: True

Q3) Calcium fluoride, CaF₂, is a source of fluorine and is used to fluoridate drinking water. Calculate its molar mass.

A) 118.15 g/mol

B) 99.15 g/mol

C) 78.07 g/mol

D) 59.08 g/mol

E) 50.01 g/mol

Answer: C

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Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) Which of the following is a weak acid?

A) H₂SO₄

B) HNO₃

C) HF

D) HBr

E) HCl

Q2) The amount of calcium present in milk can be determined by adding oxalate to a sample and measuring the mass of calcium oxalate precipitated. What is the mass percent of calcium if 0.429 g of calcium oxalate forms in a 125-g sample of milk when excess aqueous sodium oxalate is added?

Na₂C₂O₄(aq) + Ca²⁺(aq) \(\to\) CaC₂O₄(s) + 2Na⁺(aq)

A) 0.107%

B) 0.202%

C) 0.343%

D) 1.10%

E) 1.37%

Q3) In a redox reaction, the oxidizing agent undergoes loss of electrons.

A)True

B)False

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) Arrange the following gases in order of increasing rate of effusion. C₂H₆Ar HCl PH₃

A) Ar < HCl < PH₃< C₂H₆ B) C₂H₆< PH₃< HCl < Ar

C) Ar < PH₃< C₂H₆< HCl

D) C₂H₆< HCl < PH₃< Ar

E) Ar < PH₃< HCl< C₂H₆

Q2) 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? Li₂O(aq) + CO₂(g) \(\to\) Li₂CO₃(s)

A) 687 mL

B) 819 mL

C) 687 L

D) 819 L

E) 22.4 L

Q3) State Avogadro's Law.

Q4) 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) The compound carbon suboxide, C₃O₂, is a gas at room temperature. Use the data supplied to calculate the heat of formation of carbon suboxide.

(Data: 2CO(g) + C(s) \(\to\) C₃O₂(g) \(\Delta\)H° = 127.3 kJ/mol

And: \(\Delta\)H_f^° of CO(g) = -110.5 kJ/mol)

A) 116.8

B) -93.7

C) 227.8

D) -348.3

E) 93.7

Q2) A mass of 1.250 g of benzoic acid (C₇H₆O₂) was completely combusted in a bomb calorimeter. If the heat capacity of the calorimeter was 10.134 kJ/K and the heat of combustion of benzoic acid is -3226 kJ/mol, calculate (to three decimal places) the temperature increase that should have occurred in the apparatus.

Q3) Starting from equations relating pressure to force and force to work, derive the relationship w = - P\(\Delta\)V, explaining the steps in your argument.

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) Green light has a wavelength of 5200 Å. Calculate the energy of one photon of green light.

A) 3.4 × 10^-40 J

B) 3.4 × 10^-30J

C) 3.8 × 10^-29 J

D) 3.4 × 10^-27 J

E) 3.8 × 10^-19 J

Q2) The size 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 angular momentum and magnetic quantum numbers, together.

Q3) As the frequency of electromagnetic radiation increases, its wavelength also increases.

A)True

B)False

Q4) The following combinations of quantum numbers are not allowed.

Q5) What is the minimum uncertainty in the position of a neutron if the uncertainty in its speed is 0.0250 m/s?

Page 9

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Which of the following elements has the largest second ionization energy (IE₂)?

A) Li

B) B

C) O

D) F

E) Na

Q2) Select the element with the greatest metallic character.

A) Li

B) Ca

C) Al

D) Pb

E) Cs

Q3) Select the element with the most negative electron affinity (i.e., accepts an electron most readily).

A) H

B) Li

C) C

D) F

E) Ne

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Chapter 9: Models of Chemical Bonding

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Q1) Electronegativity is a measure of

A) the energy needed to remove an electron from an atom.

B) the energy released when an electron is added to an atom.

C) the magnitude of the negative charge on an electron.

D) the attraction by an atom for electrons in a chemical bond.

E) the magnitude of the negative charge on a molecule.

Q2) The melting points of metals are only moderately high because A) metallic bonding is weak.

B) metals have fewer bonding electrons than non-metals.

C) metals also have relatively low boiling points.

D) the melting process does not break the metallic bonds.

E) metals prefer to be bonded to non-metals.

Q3) Covalently bonded substances do not necessarily exist as separate molecules.

A)True

B)False

Q4) Bond energy increases as bond order increases, for bonding between a given pair of atoms.

A)True

B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Using SO₂ as an example, describe the sort of experimental data which might suggest that no single Lewis structure is an accurate representation of its bonding.

Q2) 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.

Q3) What is the shape of the PF₃ molecule? Explain your answer, using VSEPR theory.

Q4) 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

Q5) Name and outline the concept which is introduced when more than one valid Lewis structure can be drawn for a given molecule or ion. Use appropriate diagrams of the formate ion (HCO₂^-, carbon is the central atom) to illustrate.

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Chapter 11: Theories of Covalent Bonding

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Q1) In the valence bond treatment, a \(\pi\) bond is formed when two 2p orbitals overlap side to side.

A)True

B)False

Q2) a. In the context of valence bond theory, explain the difference in geometry between a \(\sigma\) and a \(\pi\) bond. Use a real molecule to illustrate your answer.

b. What two important differences are there in the properties of \(\sigma\) and \(\pi\) bonds, in terms of how they affect the structure and reactivity of molecules?

Q3) According to valence bond theory, which of the following molecules involves sp² hybridization of orbitals on the central atom? (central atom is bold)

A) C₂H₂

B) C₂H₄

C) C₂H₆

D) CO₂

E) H₂O

Q4) In one sentence state the basic principle of valence bond theory.

Q5) In one sentence state how molecular orbitals are usually obtained.

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Page 13

Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes

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Sample Questions

Q1) Which one of the following substances will have hydrogen bonds between molecules?

A) (CH₃)₃N

B) CH₃-O-CH₃

C) CH₃CH₂-OH

D) CH₃CH₂-F

E) HI

Q2) The phase diagram for xenon has a solid-liquid curve with a positive slope. Which of the following is true?

A) Solid xenon has a higher density than liquid xenon.

B) Solid xenon has the same density as liquid xenon.

C) The phase diagram cannot be used to predict which phase of xenon is denser.

D) Freezing xenon is an endothermic process.

E) None of the above statements is true.

Q3) The surface tension of water is lowered when a detergent is present in solution.

A)True

B)False

Q4) The energy of a hydrogen bond is greater than that of a typical covalent bond.

A)True

B)False

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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 Tyndall effect

A) is observed in concentrated solutions.

B) is observed only in dilute solutions.

C) is observed in colloidal dispersions.

D) is caused by Brownian motion.

E) is used to determine the osmotic pressure of solutions.

Q3) An emulsion is a dispersion consisting of a

A) solid in a liquid.

B) liquid in a liquid.

C) gas in a liquid.

D) liquid in a solid.

E) gas in a solid.

Page 15

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Q1) Silicon carbide is an important industrial chemical. Which of the following statements about silicon carbide is true?

A) It will react with water to form acetylene.

B) It is used in preparation of glass which can be used at high temperature.

C) It is the hardest known substance.

D) It is a naturally occurring mineral source of silicon.

E) It can be doped to form a high temperature semiconductor.

Q2) Which of the following would you predict to have the greatest thermal stability?

A) AsAt₃

B) AsI₃

C) AsBr₃

D) AsCl₃

E) AsF₃

Q3) Sulfuric acid is produced when sulfur dioxide dissolves in water.

A)True

B)False

Q4) Oxygen gas (O₂) is paramagnetic, but ozone (O₃) is diamagnetic.

A)True

B)False

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Chapter 15: Organic Compounds and the Atomic Properties of

Carbon

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Sample Questions

Q1) The most abundant organic chemical on earth is A) glycogen.

B) starch.

C) cellulose.

D) glucose.

E) carbonate.

Q2) All alcohols are capable of hydrogen bonding.

A)True

B)False

Q3) What is the simplest molecular formula for

a. an alkane capable of having structural isomers?

b. an alkane capable of having optical isomers?

c. an alkene capable of having geometrical isomers?

Q4) All the atoms in a molecule of benzene lie in the same plane.

A)True

B)False

Q5) a. Draw two different structures with the molecular formula C₂H₆O.

b. Name the functional group in each structure.

c. Which one will have the higher boiling point, and why?

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Q1) When the reaction A \(\to\) B + C is studied, a plot of ln[A]_t vs. time gives a straight line with a negative slope. What is the order of the reaction?

A) zero

B) first

C) second

D) third

E) More information is needed to determine the order.

Q2) A transition state is a species (or state) corresponding to an energy maximum on a reaction energy diagram.

A)True

B)False

Q3) The units of the rate constant depend on the order of the reaction.

A)True

B)False

Q4) Chlorine atoms act as heterogeneous catalysts in the destruction of ozone in the stratosphere.

A)True

B)False

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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) In a chemical reaction, if the starting concentrations of reactants are increased, then the equilibrium constant K_c will also increase.

A)True

B)False

Q3) There is a direct correlation between the speed of a reaction and its equilibrium constant.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Q1) (a) Write a balanced equation representing the reaction of the acid, H₂PO₄^- with the base, water.

(b) Write the expression for K_a of H₂PO₄^- in terms of concentrations of relevant species.

Q2) The substance NaNO₃is considered

A) a weak Arrhenius acid.

B) a weak Arrhenius base.

C) a strong Arrhenius acid.

D) a strong Arrhenius base.

E) a neutral compound.

Q3) A 1.25 M solution of the weak acid HA is 9.2% dissociated. What is the pH of the solution?

A) 0.64

B) 0.94

C) 1.13

D) 2.16

E) none of the above

Q4) Define an acid according to the Arrhenius theory, and write a balanced equation to support this definition.

Page 20

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Equal volumes of the following pairs of solutions are mixed. Which pair will produce a buffer solution?

A) 0.10 mol L^-1 HCl and 0.05 mol L^-1 NaOH

B) 0.10 mol L^-1 HCl and 0.15 mol L^-1 NH₃

C) 0.10 mol L^-1 HCl and 0.05 mol L^-1 NH₃

D) 0.10 mol L^-1 HCl and 0.20 mol L^-1 CH₃COOH

E) 0.10 mol L^-1 HCl and 0.20 mol L^-1 NaCl

Q2) A buffer is prepared by adding 0.5 mol of solid sodium hydroxide to 1.0 L of 1.0 M acetic acid (CH₃COOH). What is the pH of the buffer?

A) The pH will be pK_a - 0.30, where pK_a is that of acetic acid.

B) The pH will be greater than the pK_a for acetic acid.

C) The pH will be less than the value in answer a.

D) The pH will be equal to the pK_a for acetic acid.

E) More information is needed to solve the problem.

Q3) Increasing the concentrations of the components of a buffer solution will increase the buffer range.

A)True

B)False

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

Chemical Reactions

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Q1) Which relationship or statement best describes \(\Delta\)S° for the following reaction?

2NH₃(g) + 2ClF₃(g) \(\to\) 6HF(g) + N₂(g) + Cl₂(g)

A) .\(\Delta\)S° \(\approx\) 0

B) .\(\Delta\)S° < 0

C) .\(\Delta\)S° > 0

D) .\(\Delta\)S° = \(\Delta\)H°/T

E) More information is needed to make a reasonable prediction.

Q2) For a given reaction, a change in the pressure may result in a change in the sign of \(\Delta\)G.

A)True

B)False

Q3) 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.

Page 22

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) Consider the following balanced redox reaction

Mn²⁺(aq) + S₂O₈^2-(aq) + 2H₂O(l) \(\to\) MnO₂(s) + 4H⁺(aq) + 2SO₄^2-(aq)

Which of the following statements is true?

A) Mn²⁺(aq) is the oxidizing agent and is reduced.

B) Mn²⁺(aq) is the oxidizing agent and is oxidized.

C) Mn²⁺(aq) is the reducing agent and is oxidized.

D) Mn²⁺(aq) is the reducing agent and is reduced.

E) Manganese does not change its oxidation number in this reaction.

Q2) Consider the following balanced redox reaction

3CuO(s) + 2NH₃(aq) \(\to\) N₂(g) + 3H₂O(l) + 3Cu(s)

Which of the following statements is true?

A) CuO(s) is the oxidizing agent and copper is reduced.

B) CuO(s) is the oxidizing agent and copper is oxidized.

C) CuO(s) is the reducing agent and copper is oxidized.

D) CuO(s) is the reducing agent and copper is reduced.

E) CuO(s) is the oxidizing agent and N₂(g) is the reducing agent.

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Chapter 22: The Elements in Nature and Industry

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Q1) The Frasch process is used to

A) convert aluminum oxide to metallic aluminum.

B) convert copper ore to copper sulfide.

C) mine magnesium.

D) mine elemental sulfur.

E) produce sodium metal.

Q2) 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.

Q3) The most abundant element in the Earth's crust is

A) H.

B) Fe.

C) Al.

D) Si.

E) O.

Q4) Hydrogen displays the largest kinetic isotope effect of all the elements.

A)True

B)False

Q5) 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.

Page 24

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) Which of the following will be paramagnetic?

A) V⁵⁺

B) Ni²⁺

C) Mn⁷⁺

D) Ti⁴⁺

E) Zn

Q2) In the spectrochemical series, which one of the following ligands has the strongest field?

A) H₂O

B) CN^-

C) NH₃

D) OH^-

E) Cl^-

Q3) Which of the following ions could exist in only the high-spin state in an octahedral complex?

A) Cr²⁺

B) Mn⁴⁺

C) Fe³⁺

D) Co³⁺

E) Ni²⁺

Page 25

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) Most foodstuffs contain natural, radioactive isotopes.

A)True

B)False

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

Q3) A N-14 nucleus is hit by a particle, forming a C-14 nucleus and a proton as the only products. Identify the type of particle which struck the N-14 nucleus.

A) alpha

B) proton

C) electron

D) neutron

E) deuterium

Q4) 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

Q5) Write a complete, balanced equation to represent the beta decay of thallium-207.

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