Principles of Chemistry Final Exam - 3679 Verified Questions

Principles of Chemistry

Final Exam

Course Introduction

Principles of Chemistry provides a comprehensive introduction to the fundamental concepts underlying the science of chemistry. The course covers atomic and molecular structure, periodic trends, chemical bonding, stoichiometry, states of matter, solutions, thermochemistry, and basic chemical reactions. Students will develop problem-solving skills and gain an understanding of the principles that govern chemical processes in both laboratory and real-world contexts. The course serves as a foundation for advanced study in chemistry and related scientific fields, emphasizing both conceptual understanding and practical application.

Recommended Textbook

Chemistry The Central Science 13th Edition by Theodore E. Brown

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24 Chapters

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Page 2

Chapter 1: Introduction: Matter and Measurement

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Sample Questions

Q1) A certain liquid has a density of 2.67 g/cm³. 1340 g of this liquid would occupy a volume of ________ L.

A)1.99 × 10^-³

B)50.2

C)3.58

D)35.8

E)0.502

Answer: E

Q2) The density of silver is 10.5 g/cm³. A piece of silver with a mass of 61.3 g would occupy a volume of ________ cm³.

A)0.171

B)644

C)10.5

D)0.00155

E)5.84

Answer: E

Q3) Sn is the symbol for the element ________.

Answer: Tin

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3

Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) The formula for the compound formed between aluminum ions and phosphate ions is ________.

A)Al₃(PO₄)₃

B)AlPO₄

C)Al(PO₄)₃

D)Al₂(PO₄)₃

E)AlP

Answer: B

Q2) ________ are found uncombined, as monatomic species in nature.

A)Noble gases

B)Chalcogens

C)Alkali metals

D)Alkaline earth metals

E)Halogens

Answer: A

Q3) The correct name for Na₃N is sodium azide.

A)True

B)False

Answer: False

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Page 4

Chapter 3: Stoichiometry: Calculations With Chemical

Formulas and Equations

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Sample Questions

Q1) A compound that is composed of only carbon and hydrogen contains 80.0% C and 20.0% H by mass. What is the empirical formula of the compound?

A)C₂₀H₆₀

B)C₇H₂₀

C)CH₃

D)C₂H₆

E)CH₄

Answer: C

Q2) The combustion of propane (C₃H₈)produces CO₂ and H₂O: C₃H₈ (g) + 5O₂ (g) 3CO₂ (g) + 4H₂O (g)

The reaction of 2.5 mol of O₂ with 4.6 mol of C₃H₈ will produce ________ mol of H₂O.

A)4.0

B)3.0

C)2.5

D)2.0

E)1.0

Answer: D

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Chapter 4: Aqueous Reactions and Solution Stoichiometry

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Sample Questions

Q1) An aqueous ethanol solution (400 mL)was diluted to 4.00 L, giving a concentration of 0.0400 M. The concentration of the original solution was ________ M.

A)0.400

B)0.200

C)2.00

D)1.60

E)4.00

Q2) The spectator ions in the reaction between aqueous perchloric acid and aqueous barium hydroxide are ________.

A)OH^- and ClO₄^-

B)H⁺, OH^- , ClO₄^-, and Ba²⁺

C)H⁺ and OH^-

D)H⁺ and Ba²⁺

E)ClO₄^- and Ba²⁺

Q3) The compound HClO₄ is a weak acid.

A)True

B)False

Q4) The solvent in an aqueous solution is ________.

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Chapter 5: Thermochemistry

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Sample Questions

Q1) The value of H° for the reaction below is -126 kJ. The amount of heat that is released by the reaction of 10.0 g of Na₂O₂ with water is ________ kJ. 2Na₂O₂ (s) + 2H₂O (l) 4NaOH (s) + O₂ (g)

A)8.08

B)16.2

C)67.5

D)32.3

E)-126

Q2) A sample of iron absorbs 81.0 J of heat, upon which the temperature of the sample increases from 19.7 °C to 28.2 °C. If the specific heat of iron is 0.450 J/g-K, what is the mass (in grams)of the sample?

A)21.2 g

B)4.29 g

C)-21.2 g

D)0.0472 g

E)3.83 g

Q3) The standard enthalpy change of a reaction is the enthalpy change when all reactants and products are at ________ pressure and a specific temperature.

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Page 7

Chapter 6: Electronic Structure of Atoms

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Sample Questions

Q1) The ________ quantum number defines the shape of an orbital.

A)spin

B)magnetic

C)principal

D)angular momentum

E)psi

Q2) The angular momentum quantum number for the outermost electrons in a manganese atom in the ground state is ________.

A)1

B)2

C)3

D)-1

E)0

Q3) When the electron in a hydrogen atom moves from n = 6 to n = 1, light with a wavelength of ________ nm is emitted.

A)487

B)411

C)434

D)93.8

E)657

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Chapter 7: Periodic Properties of the Elements

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Sample Questions

Q1) The degree of interaction between two electrical charges depends on the ________ and the ________ of the charges and the distance between them.

Q2) Metals can be ________ at room temperature.

A)liquid only

B)solid only

C)solid or liquid

D)solid, liquid, or gas

E)liquid or gas

Q3) Which equation correctly represents the third ionization of aluminum?

A)Al²⁺(g) Al³⁺(g)+ e^-

B)Al(g) Al⁺(g)+ e^-

C)Al^2-(g) Al^3-(g)+ e^-

D)Al³⁺(g)+ e^- Al²⁺(g)

E)Al⁺(g) Al²⁺(g)+ e^-

Q4) Which alkaline earth metal will not react with liquid water or with steam?

A)Be

B)Mg

C)Ca

D)Ba

E)They all react with liquid water and with steam.

Page 9

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Chapter 8: Basic Concepts of Chemical Bonding

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Sample Questions

Q1) A valid Lewis structure of ________ cannot be drawn without violating the octet rule.

A)NF₃

B)BeH₂

C)SO₂

D)CF₄

E)SO₃^2-

Q2) Draw the Lewis structure of ICl₂⁺.

Q3) When a metal gains an electron, the process is endothermic.

A)True

B)False

Q4) Which of the following would have to gain two electrons in order to achieve a noble gas electron configuration? O Sr Na Se Br

A)Br

B)Sr

C)Na

D)O, Se

E)Sr, O, Se

Q5) Using the noble gas shorthand notation, write the electron configuration for Fe³⁺.

Page 10

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Chapter 9: Molecular Geometry and Bonding Theories

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Sample Questions

Q1) The F-N-F bond angle in the NF₃molecule is slightly less than

A)90°

B)109.5°

C)120°

D)180°

E)60°

Q2) The Lewis structure of carbon monoxide is given below. The hybridizations of the carbon and oxygen atoms in carbon monoxide are ________ and ________, respectively. : C O :

A)sp, sp³

B)sp², sp³

C)sp³, sp²

D)sp, sp

E)sp², sp²

Q3) Each molecular orbital can accommodate, at most, two electrons with their spins paired. This is called the ________.

Q4) The sensation of vision results from a nerve impulse that is triggered by the separation of retinal from ________.

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Chapter 10: Gases

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Sample Questions

Q1) Sodium bicarbonate is reacted with concentrated hydrochloric acid at 25.0 °C and 1.50 atm. The reaction of 7.75 kg of bicarbonate with excess hydrochloric acid under these conditions will produce ________ L of CO₂.

A)1.82 × 10³

B)2.85 × 10⁴

C)1.82 × 10⁴

D)1.50 × 10³

E)8.70 × 10²

Q2) One significant difference between gases and liquids is that ________.

A)a gas is made up of molecules

B)a gas expands to fill its container

C)a gas may consist of both elements and compounds

D)gases are always mixtures

E)All of the above answers are correct.

Q3) The van der Waals equation corrects the ideal gas law for the finite volume and ________ of gas molecules.

Q4) What is the density (in g/L)of oxygen gas at 77.0 °C and 700.0 torr?

Q5) The temperature and pressure specified by STP are ________ °C and ________ atm.

Q6) The rms speed of methane molecules at 45.0 °C is ________ m/sec.

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Chapter 11: Liquids and Intermolecular Forces

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Sample Questions

Q1) Volatility and vapor pressure are ________.

A)inversely proportional to one another

B)directly proportional to one another

C)not related

D)the same thing

E)both independent of temperature

Q2) Which statements about viscosity are true? (i)Viscosity increases as temperature decreases.

(ii)Viscosity increases as molecular weight increases.

(iii)Viscosity increases as intermolecular forces increase.

A)(i)only

B)(ii)and (iii)

C)(i)and (iii)

D)none

E)all

Q3) Heats of vaporization are greater than heats of fusion.

A)True

B)False

Q4) The direct conversion of a solid to a gas is called ________.

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Which element is typically not added to steel to modify its properties?

A)carbon

B)vanadium

C)chromium

D)nitrogen

E)nickel

Q2) When lattice points occur at the corners and at the center of a unit cell, the cell is called ________.

Q3) Of the following, only ________ is not a polymer.

A)dental amalgam

B)cellulose

C)Teflon

D)plastic

E)protein

Q4) 24 karat gold contains ________% gold.

A)24

B)25

C)50.

D)75

E)1.0 × 10²

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Chapter 13: Properties of Solutions

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Sample Questions

Q1) Which one of the following substances is more likely to dissolve in CCl₄?

A)CBr₄

B)HBr

C)HCl

D)CH₃CH₂OH

E)NaCl

Q2) Which of the following statements is false?

A)Nonpolar liquids tend to be insoluble in polar liquids.

B)The weaker the attraction between the solute and solvent molecules, the greater the solubility.

C)Substances with similar intermolecular attractive forces tend to be soluble in one another.

D)The solubility of a gas increases in direct proportion to its partial pressure above the solution.

E)The solubility of gases in water decreases with increasing temperature.

Q3) A solution contains 150.8 grams of NaCl in 678.3 grams of water. Calculate the vapor pressure of water (in torr)over the solution at 25.0 °C. (Note: The vapor pressure of pure water at 25.0 °C is 23.76 torr.)

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The rate constant for this reaction is ________ s^-1.

A)6.9 × 10^-2

B)3.0 × 10^-2

C)14

D)0.46

E)4.0 × 10²

Q2) The rate law for this reaction is rate = ________.

A)k[A][B]

B)k[P]

C)k[A]²[B]

D)k[A]²[B]²

E)k[A]²

Q3) A catalyst that is present in the same phase as the reacting molecules is called a(n)________ catalyst.

Q4) The minimum energy to initiate a chemical reaction is the ________.

Q5) The instantaneous rate of a reaction can be read directly from the graph of molarity versus time at any point on the graph.

A)True

B)False

Q6) For the reaction aA + Bb cC + dD, the rate law is ________.

Page 16

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) What role did Karl Bosch play in development of the Haber-Bosch process?

A)He discovered the reaction conditions necessary for formation of ammonia.

B)He originally isolated ammonia from camel dung and found a method for purifying it.

C)Haber was working in his lab with his instructor at the time he worked out the process.

D)He developed the equipment necessary for industrial production of ammonia.

E)He was the German industrialist who financed the research done by Haber.

Q2) In an exothermic equilibrium reaction, increasing the reaction temperature favors the formation of reactants.

A)True

B)False

Q3) Le Châtelier's principle states that if a system at equilibrium is disturbed, the equilibrium will shift to minimize the disturbance.

A)True

B)False

Q4) If the reaction quotient Q for a reaction is greater than the value of the equilibrium constant K for that reaction at a given temperature, ________ must be converted to ________ for the system to reach equilibrium.

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17

Chapter 16: Acid-Base Equilibria

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Sample Questions

Q1) A Br nsted-Lowry acid is defined as a substance that ________.

A)increases K_a when placed in H₂O

B)decreases [H⁺] when placed in H₂O C)increases [OH^-] when placed in H₂O

D)acts as a proton acceptor

E)acts as a proton donor

Q2) K_b for NH₃ is 1.8 × 10^-⁵. What is the pH of a 0.40 M aqueous solution of NH₄Cl at 25.0 °C?

A)2.57

B)11.43

C)9.18

D)4.82

E)11.23

Q3) The conjugate base of H₂PO₄^- is ________.

A)H₃PO₄

B)HPO₄^2-

C)PO₄^3-

D)H₃O⁺

E)OH^-

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Page 18

Chapter 17: Additional Aspects of Aqueous Equilibria

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Q1) The pH of a solution prepared by dissolving 0.350 mol of solid methylamine hydrochloride (CH₃NH₃Cl)in 1.00 L of 1.10 M methylamine (CH₃NH₂)is ________. The K_b for methylamine is 4.40 × 10^-⁴. (Assume the final volume is 1.00 L.)

A)1.66

B)2.86

C)10.28

D)11.14

E)10.61

Q2) Which one of the following will cause hemoglobin to release oxygen?

A)increase in pH

B)decrease in pH

C)decrease in temperature

D)decrease in CO₂ concentration

E)increase in O₂ concentration

Q3) Calculate the pH of a buffer that contains 0.270 M hydrofluoric acid (HF)and 0.180 M cesium fluoride (CsF). The K_a of hydrofluoric acid is 6.80 × 10^-4.

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Page 19

Chapter 18: Chemistry of the Environment

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Q1) In the past, CFCs were not used in ________.

A)spray cans

B)plastic manufacturing

C)air conditioners

D)refrigerators

E)dry cleaning

Q2) Which process in the Global Water Cycle involves the phase transition

H₂O (l) H₂O (g)?

A)evaporation from fresh water

B)fog drip

C)plant uptake

D)freshwater storage

E)snowmelt runoff to streams

Q3) Natural, unpolluted rainwater is typically acidic. What is the source of this natural acidity?

A)CO₂

B)SO₂

C)NO₂

D)HCl

E)chlorofluorocarbons

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Chapter 19: Chemical Thermodynamics

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Sample Questions

Q1) The value of G° at 25 °C for the decomposition of gaseous sulfur trioxide to solid elemental sulfur and gaseous oxygen, 2SO₃ (g) 2S (s, rhombic) + 3O₂ (g)

Is ________ kJ/mol.

A)+740.8

B)-370.4

C)+370.4

D)-740.8

E)+185.2

Q2) The value of G° at 25 °C for the formation of POCl₃ from its constituent elements, P₂ (g) + O₂ (g) + 3Cl₂ (g) 2POCl₃ (g)

Is ________ kJ/mol.

A)-1108.7

B)+1108.7

C)-606.2

D)+606.2

E)-1,005

Q3) Calculate G° for the autoionization of water at 25 °C. K_w = 1.0 × 10^-14

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Chapter 20: Electrochemistry

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Sample Questions

Q1) The more ________ the value of E°_red, the greater the driving force for reduction.

A)positive

B)negative

C)exothermic

D)endothermic

E)extensive

Q2) Which substance is serving as the reducing agent in the following reaction? 14H⁺ + Cr₂O₇^2- + 3Ni 3Ni²⁺ + 2Cr³⁺ + 7H₂O

A)Ni

B)H⁺

C)Cr₂O₇^2-

D)H₂O

E)Ni²⁺

Q3) At constant temperature and pressure the Gibbs free energy value is a measure of the ________ of a process.

Q4) In the formula G = -nFE, F is the ________.

Q5) The anode of the alkaline battery is powdered zinc in a gel that contacts ________.

Page 22

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Chapter 21: Nuclear Chemistry

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Sample Questions

Q1) What happens in the nucleus of an atom that undergoes positron emission?

Q2) Carbon-11 is used in medical imaging. The half-life of this radioisotope is 20.4 min. What percentage of a sample remains after 50.0 min?

A)66.5

B)8.62

C)18.3

D)40.8

E)2.45

Q3) Charged particles are accelerated because the faster they move there is a greater chance of producing a nuclear reaction.

A)True

B)False

Q4) The half-life of carbon-11 is 20.3 minutes. How much of a 100.0 mg sample remains after 1.6 hours?

A)3.77 mg

B)0.0377 mg

C)94.7 mg

D)1.75 × 10^-89_mg

E)99.9 mg

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Chapter 22: Chemistry of the Nonmetals

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Q1) The oxidation number of B in B₂O₃ is ________.

A)+3

B)+2

C)+1

D)+4

E)+5

Q2) The oxidation numbers of sulfur in the sulfate ion, sulfite ion, sulfur trioxide, and hydrogen sulfide are ________, ________, ________, and ________, respectively.

A)+4, -2, +4, +6

B)+6, +2, +4, +6

C)+6, +4, +6, -2

D)+4, +6, +4, -2

E)-2, +6, -2, 0

Q3) B₂O₃ is the anhydride of ________.

A)borous acid

B)diborane

C)tetraboric acid

D)boric acid

E)borax

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Page 24

Chapter 23: Transition Metals and Coordination Chemistry

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Q1) Which one of the following is not true about transition metals?

A)They frequently have more than one common oxidation state.

B)Their compounds are frequently colored.

C)Their compounds frequently exhibit magnetic properties.

D)They are found in the d-block of the periodic table.

E)They typically have low melting points.

Q2) What is the oxidation state of cobalt in [Co(NH₃)₅Cl](NO₃)₂?

A)0

B)2

C)3

D)4

E)6

Q3) Draw a diagram of the short-hand ground state electron configuration of zinc.

Q4) How many bonds can ethylenediamine form to a metal ion?

A)1

B)2

C)3

D)4

E)6

25

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Chapter 24: The Chemistry of Life: Organic and Biological Chemistry

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Q1) The octane number of straight-run gasoline is about ________.

A)0

B)25

C)50

D)75

E)93

Q2) Hydrogenation of what alkyne produces propane?

Q3) What geometric shape will the structure of cyclohexane appear as when drawn as a line structure?

A)hexagon

B)heptagon

C)octagon

D)triangle

E)square

Q4) Predict the product of the catalytic hydrogenation of 6-ethyl-3-decene.

Q5) Humans digest starch but not cellulose because of differences in the type of linkage between the glucose monomers of these substances.

A)True

B)False

Q6) In the oxidation of ethanol the intermediate formed is ________. Page 26

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