Introductory Chemistry Textbook Exam Questions - 3071 Verified Questions

Introductory Chemistry

Textbook Exam Questions

Course Introduction

Introductory Chemistry provides a comprehensive foundation in the basic principles of chemistry, including the structure of matter, chemical bonding, stoichiometry, states of matter, and chemical reactions. Designed for students with little or no prior background in chemistry, this course emphasizes conceptual understanding and practical applications to everyday life. Through a combination of lectures, laboratory experiments, and problem-solving exercises, students develop essential skills in scientific reasoning, analytical thinking, and laboratory techniques, preparing them for more advanced studies in chemistry and related fields.

Recommended Textbook

Chemistry Structure and Properties 2nd Edition by Nivaldo J. Tro

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Chapter 1: Essentials: Units, Measurements, and Problem

Solving

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Sample Questions

Q1) Define energy.

Answer: Energy is the capacity to do work.

Q2) The average distance between nitrogen and oxygen atoms is 115 pm in a compound called nitric oxide.What is this distance in centimeters?

A) 1.15 × 10^-⁹cm

B) 1.15 × 10^-⁸cm

C) 1.15 × 10¹²cm

D) 1.15 × 10¹⁶ cm

Answer: B

Q3) How many significant figures are in 4.500 × 10⁴ m?

A) 3

B) 4

C) 5

D) 7

E) 8

Answer: B

Q4) Define random error.

Answer: Random error has an equal probability of being too high or too low.

Q5) Define the law of the conservation of energy.

Answer: Energy is neither created or destroyed.

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Chapter 2: Atoms

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Sample Questions

Q1) Identify the crystalline solid.

A) plastic

B) glass

C) table salt

D) water

E) bleach

Answer: C

Q2) How many neutrons are in magnesium?

A) 12

B) 13

C) 14

D) 12.3

E) 24.3

Answer: C

Q3) A wooden baseball bat is an example of

A) a compound.

B) an element.

C) a heterogeneous mixture.

D) a homogeneous mixture.

Answer: C

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Chapter 3: The Quantum Mechanical Model of the Atom

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Q1) How much energy (in kJ)is required to ionize 1.97 moles of hydrogen atoms?

A) 1.29 × 10³ kJ

B) 1.19 × 10³ kJ

C) 4.29 × 10³ kJ

D) 2.59 × 10³ kJ

E) 5.89 × 10³ kJ

Answer: D

Q2) For n = 3,what are the possible sublevels?

A) 0

B) 0, 1

C) 0, 1, 2

D) 0, 1, 2, 3

Answer: C

Q3) For a hydrogen atom,which electronic transition would result in the emission of a photon with the highest energy?

A) 1s 2p

B) 3p 7d

C) 3p 1s

D) 6f 4d

Answer: C

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Chapter 4: Periodic Properties of the Elements

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Q1) Describe the reaction of the noble gases with metals.

A) inert

B) vigorous

C) mild reaction

D) forms water

E) dissolves

Q2) Which of the following elements is a metalloid?

A) Sb

B) N

C) Br

D) Co

E) Kr

Q3) Which of the following does NOT describe a metal?

A) good conductor of heat

B) good conductor of electricity

C) tends to gain electrons

D) forms ionic compounds with nonmetals

E) found on the left side of the periodic table

Q4) Why do successive ionization energies increase?

Q5) Give the name of the element whose symbol is Na.

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Chapter 5: Molecules and Compounds

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Sample Questions

Q1) A double covalent bond contains ________ of electrons.

A) 0 pairs

B) 1 pair

C) 2 pairs

D) 3 pairs

E) 4 pairs

Q2) Calculate the mass percent composition of lithium in Li₃PO₄.

A) 26.75 %

B) 17.98 %

C) 30.72 %

D) 55.27 %

E) 20.82 %

Q3) Identify an amine.

A) CH₃CH₂SCH₂CH₃

B) CH₃CH₂Cl

C) CH₃CH₂NH₂

D) CH₃CH₂CH₂CH₃

E) CH₃COOH

Q4) Why aren't prefixes used in naming ionic compounds?

Page 7

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Chapter 6: Chemical Bonding I

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Sample Questions

Q1) Identify the compound with the largest dipole moment in the gas phase.

A) Cl₂

B) ClF

C) HF

D) LiF

Q2) Determine the electron geometry,molecular geometry and polarity of HCN.

A) eg = trigonal bipyramidal, mg = trigonal planar, nonpolar

B) eg = octahedral, mg = square planar, nonpolar

C) eg = tetrahedral, mg = bent, polar

D) eg = tetrahedral, mg = linear, nonpolar

E) eg = linear, mg = linear, polar

Q3) Describe a covalent bond.

Q4) Choose the bond below that is the strongest.

A) C-F

B) C=O

C) C-Br

D) Br-Br

E) C C

Q5) Draw the Lewis structure for BrO₃^-.Make sure to include any important resonance structures.

Page 8

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Chapter 7: Chemical Bonding Ii

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Sample Questions

Q1) Give the electron geometry (eg),molecular geometry (mg),and hybridization for IF₄⁺.

A) eg = tetrahedral, mg = tetrahedral, sp³

B) eg = trigonal pyramidal, mg = trigonal pyramidal, sp³

C) eg = octahedral, mg = square planar, sp³d²

D) eg = octahedral, mg = octahedral, sp³d²

E) eg = trigonal bipyramidal, mg = seesaw, sp³d

Q2) Consider the following compound.How many sigma and pi bonds does it contain? NCCH₂COOH

A) 9 sigma, 4 pi

B) 11 sigma, 2 pi

C) 9 sigma, 2 pi

D) 6 sigma, 2 pi

E) 8 sigma, 3 pi

Q3) Describe a pi bond.

A) side by side overlap of p orbitals

B) end to end overlap of p orbitals

C) s orbital overlapping with the end of a d orbital

D) overlap of two d orbitals

E) p orbital overlapping with an f orbital

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Chapter 8: Chemical Reactions and Chemical Quantities

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Sample Questions

Q1) Give the percent yield when 28.16 g of CO₂ are formed from the reaction of 8.000 moles of C₈H₁₈ with 4.000 moles of O₂. 2 C₈H₁₈ + 25 O₂ 16 CO₂ + 18 H₂O

A) 20.00%

B) 25.00%

C) 50.00%

D) 12.50%

Q2) How many grams of the excess reagent are left over if 37.8g of Cl₂ react with 39.4 g of NaF?

Cl₂(g)+ 2 NaF(aq) 2 NaCl(aq)+ F₂(g)

A) 8.40 g

B) 29.4 g

C) 1.43 g

D) 12.6 g

E) 4.47 g

Q3) If 588 grams of FeS₂ is allowed to react with 352 grams of O₂ according to the following unbalanced equation,how many grams of Fe₂O₃ are produced?

FeS₂ + O₂ Fe₂O₃ + SO₂

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Chapter 9: Introduction to Solutions and Aqueous Reactions

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Sample Questions

Q1) 94.20 mL of 0.800 M potassium hydroxide reacts with 125.0 mL of a solution containing oxalic acid,which is a diprotic species.What is the concentration of the acid solution?

A) 0. 542 M

B) 0. 603 M

C) 0. 301 M

D) 1.38 M

Q2) What precipitate is most likely formed from a solution containing Ba⁺²,K⁺¹,OH^-1,and CO₃^-2.

A) KOH

B) BaCO₃

C) K₂CO₃

D) Ba(OH)₂

Q3) Identify the polyprotic acid.

A) H₂SO₄

B) HCl

C) Li I

D) NaOH

E) Ba(OH)₂

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Q4) What causes a precipitation reaction to occur between two soluble compounds?

Chapter 10: Thermochemistry

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Sample Questions

Q1) Two aqueous solutions are both at room temperature and are then mixed in a coffee cup calorimeter.The reaction causes the temperature of the resulting solution to fall below room temperature.Which of the following statements is TRUE?

A) The products have a lower potential energy than the reactants.

B) This type of experiment will provide data to calculate E_rxn.

C) The reaction is exothermic.

D) Energy is leaving the system during reaction.

E) None of the above statements are true.

Q2) How much heat is absorbed when 45.00 g of C(s)reacts in the presence of excess SO₂(g)to produce CS₂(l)and CO(g)according to the following chemical equation?

5 C(s)+ 2 SO₂(g) CS₂(l)+ 4 CO(g) H° = 239.9 kJ

A) 179.8 kJ

B) 239.9 kJ

C) 898.5 kJ

D) 2158 kJ

Q3) Where does the energy absorbed during an endothermic reaction go?

Q4) Explain the difference between H and DE.

Q5) Give the temperature and pressure for the standard state for a liquid.

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Chapter 11: Gases

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Sample Questions

Q1) A 55.0-L steel tank at 20.0°C contains acetylene gas,C₂H₂,at a pressure of 1.39 atm.Assuming ideal behavior,how many grams of acetylene are in the tank?

A) 3.17 g

B) 8.20 g

C) 82.9 g

D) 1210 g

Q2) A sample of CO₂ gas fills a container with a volume of 44.5 L.When the volume is decreased to 13.8 L at a constant temperature,the pressure inside the container is 6.25 atm.What was the pressure inside the container originally?

A) 0.035 atm

B) 0.51 atm

C) 1.94 atm

D) 3.84 atm

E) 6.25 atm

Q3) Define pressure.

Q4) Why doesn't Dalton's Law of Partial Pressures depend on the identity of the gases present?

Q5) Why does hot air rise?

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Chapter 12: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Define dynamic equilibrium.

Q2) Place the following substances in order of increasing boiling point. Ne I₂N₂

A) Ne < I₂< N₂

B) I₂ < N₂< Ne

C) N₂ < I₂< Ne

D) I₂ < Ne < N₂

E) Ne < N₂< I₂

Q3) Calculate the energy that is required to change 50.0 g ice at -30.0°C to a liquid at 73.0°C.The heat of fusion = 333 J/g,the heat of vaporization = 2256 J/g,and the specific heat capacities of ice = 2.06 J /gK and liquid water = 4.184 J /gK.

A) 3.50 × 10⁴ J

B) 2.14 × 10⁴ J

C) 1.31 × 10⁵ J

D) 6.59 × 10³ J

E) 1.66 × 10⁴ J

Q4) Define volatile.

Q5) Why does the temperature of a substance stay constant during a phase change such as vaporization?

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Chapter 13: Crystalline Solids and Modern Materials

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Sample Questions

Q1) When an X-ray beam with a wavelength of 180 pm strikes the surface of a crystal,it produces a maximum reflection at an angle of 54.0°.If n=1,what is the separation between layers of atoms in the crystal?

A) 151 nm

B) 38.9 mm

C) 111 nm

D) 55.3 nm

E) 83.5 nm

Q2) All of the following are examples of allotropes of carbon EXCEPT

A) graphite

B) glass

C) diamond

D) graphine

E) all of the above

Q3) Which of the following substances should have the highest melting point?

A) CO₂

B) SrS

C) Kr

D) F₂

E) MgO

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Chapter 14: Solutions

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Sample Questions

Q1) Identify the polar solvent.

A) acetone

B) hexane

C) diethyl ether

D) toluene

E) carbon tetrachloride

Q2) Place the following solutions in order of increasing osmotic pressure. I.0.15 M C₂H₆O₂ II.0.15 M BaCl₂ III.0.15 M Na I

A) III < I < II

B) II < III < I

C) I < II < III

D) II < I < III

E) I < III < II

Q3) What happens to a supersaturated solution of potassium acetate once it is cooled and a small crystal of solid potassium acetate is added?

Q4) Define osmosis.

Q5) Why isn't pentanol (CH₃CH₂CH₂CH₂CH₂OH) very soluble in water?

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Chapter 15: Chemical Kinetics

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Sample Questions

Q1) What data should be plotted to show that experimental concentration data fits a first-order reaction?

A) 1/[reactant] vs. time

B) [reactant] vs. time

C) ln[reactant] vs. time

D) ln(k) vs. 1/T

E) ln(k) vs. E_a

Q2) A radioactive isotope decays with a half life of 645 s.How much of the material remains after 2.00 minutes?

A) 12.1%

B) 87.9%

C) 0.02%

D) 99.8%

Q3) Explain what the exponential factor in the Arrhenius equation represents.

Q4) Define the frequency factor.

Q5) What is the difference between average reaction rate and instantaneous reaction rate?

Q6) What is a catalyst and what function does it serve?

Q7) Explain how the order of a reaction can be determined.

Q8) What function do enzymes serve?

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Chapter 16: Chemical Equilibrium

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Q1) Consider the following reaction,equilibrium concentrations,and equilibrium constant at a particular temperature.Determine the equilibrium concentration of SO₃(g). 2 SO₂(g)+ O₂(g) 2 SO₃(g)K_c = 1.7 × 10⁸ [SO₃]_eq = 0.0034 M [O₂]_eq = 0.0018 M

A) 2.8 × 10¹³ M

B) 1.88 M

C) 6.1 × 10^-6 M

D) 1.0 × 10³ M

E) 1.4 M

Q2) Determine the value of K_p for the following reaction if the equilibrium concentrations are as follows: P(CO)_eq = 6.8 × 10^-11 atm,P(O₂)_eq = 1.3 × 10^-3 atm,P(CO₂)_eq = 0.041 atm. 2 CO(g)+ O₂(g) 2 CO₂(g)

A) 3.6 × 10^-21

B) 2.8 × 10²⁰

C) 4.6 × 10¹¹

D) 2.2 × 10^-12

E) 3.6 × 10^-15

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Chapter 17: Acids and Bases

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Q1) Calculate the hydronium ion concentration in an aqueous solution with a pOH of 4.33 at 25°C.

A) 2.1 × 10^-10 M

B) 9.7 × 10^-10M

C) 4.7 × 10^-5 M

D) 3.8 × 10^-5 M

E) 6.3 × 10^-6 M

Q2) Which one of the following salts,when dissolved in water,produces the solution with a pH closest to 7.00?

A) NH₄I

B) Na₂O

C) KHCO₃ 

D) Cs Cl

Q3) Which of the following acids will have the strongest conjugate base?

A) HCl

B) HClO₄

C) HNO₃

D) HCN

E) HI

Q4) What does the term amphoteric mean?

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Chapter 18: Aqueous Ionic Equilibrium

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Q1) A 25.0-mL sample of 0.150 M hydrocyanic acid is titrated with a 0.150 M NaOH solution.What is the pH after 13.3 mL of base is added? The K_a of hydrocyanic acid is 4.9 × 10^-10.

A) 9.04

B) 1.34

C) 5.32

D) 9.37

E) 9.25

Q2) A 100.0 mL sample of 0.20 M HF is titrated with 0.10 M KOH.Determine the pH of the solution after the addition of 400.0 mL of KOH.The K_a of HF is 3.5 × 10^-4.

A) 13.08

B) 12.60

C) 13.85

D) 12.30

E) 12.78

Q3) Sketch the titration curve for a monoprotic weak acid titrated with a strong base.Make sure to indicate the equivalence point (and whether it is acidic,basic or neutral)and the buffer region.

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Chapter 19: Free Energy and Thermodynamics

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Q1) Which of the following processes have a S > 0?

A) CH₃OH(l) CH₃OH(s)

B) N₂(g) + 3 H₂(g) 2 NH₃(g)

C) CH₄(g) + H₂O(g) CO(g) + 3 H₂(g)

D) Na₂CO₃(s) + H₂O(g) + CO₂(g) 2 NaHCO₃(s)

E) All of the above processes have a DS > 0.

Q2) Given the following equation, N₂O(g)+ NO₂(g) 3 NO(g) G°_rxn = -23.0 kJ

Calculate G°_rxn for the following reaction.

3 NO(g) N₂O(g)+ NO₂(g)

A) -23.0 kJ

B) 69.0 kJ

C) -69.0 kJ

D) -7.67 kJ

E) 23.0 kJ

Q3) Why can endothermic reactions be spontaneous?

Q4) Define the second law of thermodynamics.

Q5) How is a nonspontaneous process made spontaneous?

Q6) Define the third law of thermodynamics.

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Chapter 20: Electrochemistry

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Q1) Nickel can be plated from aqueous solution according to the following half reaction.How long would it take (in min)to plate 29.6 g of nickel at 4.7 A?

Ni²⁺(aq)+ 2 e Ni(s)

A) 1.7 × 10² min

B) 5.9 × 10² min

C) 3.5 × 10² min

D) 4.8 × 10² min

E) 6.2 × 10² min

Q2) Which of the following is the strongest oxidizing agent?

A) MnO₂(s)

B) Cl (aq)

C) Ag (aq)

D) SO₄^2-(aq)

E) MnO₄ (aq)

Q3) Identify the battery that is in most automobiles.

A) dry-cell battery

B) lithium ion battery

C) lead-acid storage battery

D) NiCad battery

E) fuel cell

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Chapter 21: Radioactivity and Nuclear Chemistry

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Q1) A rock contains 0.112 mg of lead-206 for each milligram of uranium-238.The half-life for the decay of uranium-238 to lead-206 is 4.5 × 10⁹ yr.The rock was formed ________ yr ago.

A) 5.04 × 10⁸

B) 4.17 × 10⁸

C) 5.48 × 10⁸

D) 7.90 × 10⁸

E) 6.01 × 10⁸

Q2) technetium-99m can be used to label antibodies in order to help locate infections.What type of emission is associated with technetium-99m?

A) alpha emission

B) beta emission

C) gamma emission

D) positron emission

E) electron capture

Q3) Why is an alpha emitter much more harmful if is ingested than when applied to the skin?

Q4) What is the "mass defect"?

Q5) Define chain reaction in terms of the fission of uranium nucleus.

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Chapter 22: Organic Chemistry

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Q1) All of the following compounds exhibit geometric isomerism EXCEPT

A) CFCl=CHBr

B) CH₃ClC=CBrCH₃

C) CHBr=CHCl

D) CH₂=CH-CH₃

E) All of the above exhibit geometric isomerism.

Q2) Write the balanced chemical equation for the addition of HBr to CH₂=CHCH₂CH₃.

A) CH₂=CHCH₂CH₃ + HBr CH₃BrCH₂CH₂CH₃

B) CH₂=CHCH₂CH₃ + 2 HBr 2 CH₂BrCH₃

C) CH₂=CHCH₂CH₃ + 2 HBr CH₃Br + CH₂BrCH₂CH₃

D) CH₂=CHCH₂CH₃ + 4 HBr 4 CH₃Br

E) CH₂=CHCH₂CH₃ + HBr CH₃CHBrCH₂CH₃

Q3) Which part of HCN,the hydrogen or the cyano group,adds to the O in a C=O bond and why?

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Chapter 23: Transition Metals and Coordination Compounds

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Q1) Identify the transition metal that is used in hemoglobin synthesis in the human body.

A) lithium

B) iron

C) copper

D) manganese

E) chromium

Q2) What is the difference between a weak-field complex and a strong-field complex?

Q3) Choose the polydentate ligand from the substances below.

A) chloride ion

B) EDTA

C) nitrate ion

D) hydroxide ion

E) carbon monoxide

Q4) Identify the ion that is responsible for the red color of rubies.

A) Cr³⁺

B) Cr⁴⁺

C) Cr⁵⁺

D) Cr⁶⁺

E) Cr⁷⁺

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