Preparatory Chemistry Solved Exam Questions - 3071 Verified Questions

Preparatory Chemistry

Solved Exam Questions

Course Introduction

Preparatory Chemistry is designed to provide students with a foundational understanding of basic chemical principles and laboratory skills necessary for success in college-level chemistry courses. Topics covered include the structure of matter, atomic theory, chemical bonding, stoichiometry, states of matter, and basic thermochemistry. The course also emphasizes essential problem-solving strategies, quantitative reasoning, and the safe handling of chemicals in the laboratory. Through a combination of lectures, discussions, and hands-on experiments, students gain the confidence and competence required to progress to more advanced chemistry studies.

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

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23 Chapters

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Chapter 1: Essentials: Units, Measurements, and Problem Solving

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Sample Questions

Q1) Which of the following are examples of extensive properties?

A) volume

B) color

C) density

D) temperature

E) solubility

Answer: A

Q2) Convert 125°C into Kelvin

A) 0K

B) -148K

C) 583K

D) 249K

E) 398K

Answer: E

Q3) Define energy.

Answer: Energy is the capacity to do work.

Q4) The correct answer (reported to the proper number of significant figures)to the following is ( 1712 - 1615)× ( 8.66 × 7.66)= ________

Answer: 6400

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Chapter 2: Atoms

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Sample Questions

Q1) What mass (in mg)does 2.63 moles of nickel have?

A) 44.8 mg

B) 2.23 × 10⁴ mg

C) 129 mg

D) 3.56 × 10⁵ mg

E) 1.54 × 10⁵mg

Answer: E

Q2) Identify the characteristics of a gas.

A) definite volume and definite shape

B) definite volume and no definite shape

C) no definite shape and definite volume

D) no definite shape and no definite volume

Answer: D

Q3) What mass (in mg)does 0.04518 moles of silver have?

A) 8402 mg

B) 154.2 mg

C) 2284 mg

D) 4874 mg

E) 843.0 mg

Answer: D

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Chapter 3: The Quantum Mechanical Model of the Atom

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Sample Questions

Q1) Which of the following quantum numbers describes the orientation of an orbital?

A) magnetic quantum number

B) principal quantum number

C) angular momentum quantum number

D) spin quantum number

E) Schrödinger quantum number

Answer: A

Q2) How many orbitals are there in the seventh shell?

A) 6

B) 7

C) 21

D) 49

Answer: D

Q3) What is the photoelectric effect?

Answer: Many metals emit electrons when light of high enough energy is shone on them.This observation brought the classical view of light into question.

Q4) Define constructive interference.

Answer: Waves that are in phase combine with each other.

Q5) Consider a 3p orbital.How is it different from a 2p orbital?

Answer: It is larger in size and contain additional nodes.

Page 5

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Chapter 4: Periodic Properties of the Elements

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Sample Questions

Q1) Identify a characteristic of halogens.

A) powerful reducing agents

B) forms water in reactions

C) powerful oxidizing agents

D) absorbs water in reactions

E) inert

Q2) Give the ground state electron configuration for Cd.

A) [Kr]5s²5d¹⁰

B) [Kr]5s²4d¹⁰5p²

C) [Kr]4d¹⁰

D) [Kr]5s²4d⁸

E) [Kr]5s²4d¹⁰

Q3) Which element has the chemical symbol,N?

A) nickel

B) niobium

C) nitrogen

D) nobelium

Q4) Why do Li,Na,and K have similar chemical properties?

Q5) Why do elements in the same group tend to have similar chemical properties?

Q6) Define electron affinity.

Page 6

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Chapter 5: Molecules and Compounds

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Sample Questions

Q1) Place the following in order of decreasing magnitude of lattice energy. KF Mg S RbI

A) RbI > KF > Mg S

B) RbI > Mg S > KF

C) Mg S > RbI > KF

D) KF > RbI > Mg S

E) Mg S > KF > RbI

Q2) What is the mass of 0.500 mol of trichlorofluoro methane,C Cl₃F?

A) 3.64 × 10^-3 g

B) 68.7 g

C) 137 g

D) 275 g

Q3) Give the formula for calcium bisulfate.

A) CaHSO₄

B) Ca₂(HSO₄)₂

C) Ca₂HSO₄

D) Ca(HSO₄)₂

Q4) Describe the difference between a molecular formula and an empirical formula.Give an example.

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Chapter 6: Chemical Bonding I

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Sample Questions

Q1) Give the number of valence electrons for SO₂.

A) 16

B) 18

C) 20

D) 22

E) 12

Q2) Identify the compound with the smallest dipole moment in the gas phase.

A) Cl₂

B) ClF

C) HF

D) LiF

Q3) Which molecule or compound below contains a polar covalent bond?

A) C₂H₄

B) MgS

C) K F

D) NI₃

E) Ag Cl

Q4) Describe a covalent bond.

Q5) Define bond energy.

Q6) Describe the difference between a pure covalent bond and a polar covalent bond.

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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 H₂O.

A) eg = tetrahedral, mg = bent, sp³

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

C) eg = tetrahedral, mg = trigonal pyramidal, sp³

D) eg = bent, mg = bent, sp²

E) eg = trigonal planar, mg = trigonal planar, sp²

Q2) 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

Q3) Identify the number of electron groups around a molecule with sp hybridization.

A) 1

B) 2

C) 3

D) 4

E) 5

Q4) According to molecular orbital theory,what is an antibonding orbital?

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Chapter 8: Chemical Reactions and Chemical Quantities

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Sample Questions

Q1) How many moles of oxygen are formed when 58.6 g of KNO₃ decomposes according to the following reaction? The molar mass of KNO₃ is 101.11 g/mol.

4 KNO₃(s) 2 K₂O(s)+ 2 N₂(g)+ 5 O₂(g)

A) 0.290 mol O₂

B) 0.580 mol O₂

C) 18.5 mol O₂

D) 0.724 mol O₂

E) 1.73 mol O₂

Q2) If the density of ethanol,C₂H₅OH,is 0.789 g/mL.How many milliliters of ethanol are needed to produce 20.0 g of CO₂ according to the following chemical equation?

C₂H₅OH(l)+ 3 O₂(g) 2 CO₂(g)+ 3 H₂O(l)

A) 8.26 mL

B) 13.3 mL

C) 26.5 mL

D) 53.1 mL

Q3) Define a limiting reagent.

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Chapter 9: Introduction to Solutions and Aqueous Reactions

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Sample Questions

Q1) If the reaction of phosphate ion with water is ignored,what is the total concentration of ions in a solution prepared by dissolving 7.00 g of K₃PO₄ in enough water to make 350.mL of solution?

A) 0.0 236 M

B) 0. 0943 M

C) 0. 377 M

D) 0. 754 M

Q2) According to the balanced equation shown below,2.00 moles of oxalic acid,H₂C₂O₄,reacts with ________ moles of permanganate,MnO₄^-. 5

H₂C₂O₄(aq)+ 2 MnO₄^-(aq)+ 6 H⁺(aq) 10 CO₂(g)+ Mn²⁺(aq)+ 8 H₂O(l)

A) 0.800

B) 2.00

C) 4.00

D) 4.50

Q3) Describe the difference between complete ionic and net ionic equations.

Q4) Define an electrolyte.

Q5) How can you tell if a reaction is an oxidation-reduction reaction?

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Chapter 10: Thermochemistry

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Sample Questions

Q1) Using the following equation for the combustion of octane,calculate the amount of moles of carbon dioxide formed from 100.0 g of octane.The molar mass of octane is 114.33 g/mole.The molar mass of carbon dioxide is 44.0095 g/mole. 2 C₈H₁₈ + 25 O₂ 16 CO₂ + 18 H₂O H°_rxn= -11018 kJ

A) 18.18 moles

B) 6.997 moles

C) 14.00 moles

D) 8.000 moles

E) 10.93 moles

Q2) Give the units of specific heat capacity.

A) J/°C

B) (^J/g °C)

C) Jmole °C

D) g/°C

E) gmole °C

Q3) Give the temperature and pressure for the standard state for a liquid.

Q4) Where does the energy absorbed during an endothermic reaction go?

Q5) Explain the difference between H and DE.

Q6) Define chemical energy.

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Chapter 11: Gases

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Sample Questions

Q1) What is the volume of 30.0 g of argon gas at 157°C and 2.50 kPa pressure?

A) 3.87 L

B) 10.6 L

C) 393 L

D) 1070 L

Q2) A gas mixture consists of N₂,O₂,and Ne,where the mole fraction of N₂ is 0.55 and the mole fraction of Ne is 0.25.If the mixture is at STP in a 5.0 L container,how many molecules of O₂ are present?

A) 4.5 × 10²²molecules O₂

B) 2.7 × 10²² molecules O₂

C) 3.7 × 10²³ molecules O₂

D) 1.1 × 10²³ molecules O₂

E) 9.3 × 10²⁴ molecules O₂

Q3) Which of the following would have a density of 1.37 g/L at 7.0°C and 0.987 atm?

A) N₂

B) O₂

C) Kr

D) Rn

Q4) Why does the rate of effusion increase with a decrease in the molar mass?

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Chapter 12: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Determine the vapor pressure (in mm Hg)of a substance at 29°C,whose normal boiling point is 76°C and has a H_vapof 38.7 kJ/mol.

A) 80 mm Hg

B) 13 mm Hg

C) 21 mm Hg

D) 48 mm Hg

E) 96 mm Hg

Q2) Define fusion.

A) the phase transition from solid to liquid

B) the phase transition from gas to solid

C) the phase transition from gas to liquid

D) the phase transition from liquid to gas

E) the phase transition from liquid to solid

Q3) The resistance of a liquid to flow is known as

A) surface tension.

B) meniscus.

C) viscosity.

D) capillary action.

E) intermolecular forces.

Q4) Why is the H_vap higher than H_fusfor a given compound?

Page 14

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Chapter 13: Crystalline Solids and Modern Materials

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Sample Questions

Q1) Which of the following substances should have the highest melting point?

A) Fe

B) Ne

C) Xe

D) N₂

E) CO

Q2) A common use of polyvinyl chloride is

A) plumbing.

B) paint.

C) plastic bottles.

D) styrofoam.

E) all of the above.

Q3) When an X-ray beam strikes the surface of a crystal with a lattice spacing of 145 nm,it produces a maximum reflection at an angle of 41.0°.What is the wavelength of the X-ray with a value of n=1?

A) 130°

B) 164°

C) 158°

D) 190°

E) 167°

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Chapter 14: Solutions

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Sample Questions

Q1) Determine the vapor pressure of a solution at 25°C that contains 85.3 g of naphthalene (C₁₀H₈)in 540 g of benzene (C₆H₆).The vapor pressure of benzene at 25°C is 74.6 torr.

A) 74.6 toff.

B) 68.0 torr.

C) 54.9 torr.

D) 7.52 torr.

E) 22.4 toff.

Q2) A solution is prepared by dissolving 1.928 g of KNO3 into enough water to make 651

mL.What is the molarity of the solution?

A) 0.159 M

B) 0.381 M

C) 0.0293 M

D) 2.48 M

E) 0.0832 M

Q3) Give the preparation of rock candy.

Q4) What happens to a supersaturated solution of potassium acetate once it is cooled and a small crystal of solid potassium acetate is added?

Q5) Define osmosis.

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Chapter 15: Chemical Kinetics

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Sample Questions

Q1) Given the following balanced equation,determine the rate of reaction with respect to [Cl₂].If the rate of Cl₂ loss is 4.24 × 10^-2 M/s,what is the rate of formation of NOCl?

2 NO(g)+ Cl₂(g) 2 NOCl(g)

A) 4.24 × 10^-2 M/s

B) 2.12 × 10^-2 M/s

C) 1.06 × 10^-1 M/s

D) 8.48 × 10^-2 M/s

E) 1.61 × 10^-2 M/s

Q2) How many half-lives are required for the concentration of reactant to decrease to 25% of its original value?

A) 1

B) 4

C) 1.5

D) 3.5

E) 2

Q3) What is the difference between average reaction rate and instantaneous reaction rate?

Q4) Explain what the exponential factor in the Arrhenius equation represents.

Q5) Define activation energy.

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Chapter 16: Chemical Equilibrium

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Sample Questions

Q1) The equilibrium constant,K_p,equals 3.40 for the isomerization reaction: cis-2-butene trans-2-butene.

If a flask initially contains 0.250 atm of cis-2-butene and 0.145 atm of trans-2-butene,what is the equilibrium pressure of each gas?

A) P(cis-2-butene) = 0.0426 atm and P(trans-2-butene) = 0.145 atm

B) P(cis-2-butene) = 0.0471 atm and P(trans-2-butene) = 0.160 atm

C) P(cis-2-butene) = 0.0735 atm and P(trans-2-butene) = 0.250 atm

D) P(cis-2-butene) = 0.0898 atm and P(trans-2-butene) = 0.305 atm

Q2) Consider the following reaction: 2 NO(g)+ O₂(g) 2 NO₂(g)

At equilibrium,6.00 mol of NO and 1.90 mol of O₂ are present at equilibrium in a 2.50L flask.If the value for Kc is 23.8,how many moles of NO₂ are also present?

A) 25.5 mol

B) 10.2 mol

C) 41.4 mol

D) 0.114 mol

Q3) Define Le Chatelier's Principle.

Q4) Can the K_p and K_c for a reaction ever have the same value?

Why or why not?

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Chapter 17: Acids and Bases

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Sample Questions

Q1) Which one of the following salts,when dissolved in water,produces the solution with the lowest pH?

A) NaI

B) KI

C) MgI₂

D) AlI₃

Q2) Which one of the following salts,when dissolved in water,produces the solution with the highest pH?

A) KHSO₄

B) RbClO₄

C) BaO

D) CH₃CH₃NH₃Br

Q3) Which of the following solutions would have the highest pH? Assume that they are all 0.10 M in acid at 25C.The acid is followed by its K_a value.

A) HF, 3.5 × 10^-4

B) HCN, 4.9 × 10^-10

C) HNO₂, 4.6 × 10^-4

D) HCHO₂, 1.8 × 10^-4

E) HClO₂, 1.1 × 10^-2

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Chapter 18: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Calculate the pH of a buffer that is 0.040 M HF and 0.080 M NaF.The K_a for HF is 3.5 × 10^-4.

A) 2.06

B) 4.86

C) 3.16

D) 3.46

E) 3.76

Q2) Which one of the following statements is TRUE?

A) A buffer is an aqueous solution composed of two strong bases.

B) A buffer can absorb an unlimited amount of base.

C) A buffer resists pH change by neutralizing added acids and bases.

D) A buffer does not change pH when strong base is added.

E) All of the above are true.

Q3) Determine the molar solubility of BaF₂in pure water.K_sp for BaF₂ = 2.45 × 10^-5.

A) 1.83 × 10^-2 M

B) 1.23 × 10^-5 M

C) 2.90 × 10^-2 M

D) 4.95 × 10^-3 M

E) 6.13 × 10^-6 M

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Chapter 19: Free Energy and Thermodynamics

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Sample Questions

Q1) Determine G°_rxn for the following reaction at 378 K. CaCO₃(s) CaO(s)+ CO₂(g) H°= +179.2 kJ; DS°= +160.2 J/K

A) +179.2 kJ

B) + 239.8 kJ

C) - 239.8 kJ

D) + 118.6 kJ

E) - 118.6 kJ

Q2) Determine the equilibrium constant for the following reaction at 549 K. CH₂O(g)+ 2 H₂(g) CH₄(g)+ H₂O(g) H°

= -94.9 kJ; S°= -224.2 J/K

A) 481

B) 1.07 × 10⁹

C) 2.08 × 10^-3

D) 9.35 × 10^-10

E) 1.94 × 10^-12

Q3) Define the third law of thermodynamics.

Q4) Define allotrope.

Q5) Why is heating your home with gas more efficient than heating it with electricity?

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Chapter 20: Electrochemistry

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Q1) What is the balanced equation for the galvanic cell reaction expressed using shorthand notation below? .

Zn(s) Zn²⁺(aq) Cl₂(g) Cl^-(aq) C(s)

A) Zn(s) + 2 Cl^-(aq) Zn²⁺(aq) + Cl₂(g)

B) Zn(s) + Cl₂(g) Zn²⁺(aq) + 2 Cl^-(aq)

C) Zn²⁺(aq) + 2 Cl^-(aq) Zn(s) + Cl₂(g)

D) Zn²⁺(aq) + 2 Cl^-(aq) ZnCl₂(s)

Q2) The standard emf for the cell using the overall cell reaction below is +0.48 V: Zn(s)+ Ni²⁺(aq) Zn²⁺(aq)+ Ni(s)

The emf generated by the cell when [Ni²⁺] = 0.100 M and [Zn²⁺] = 2.25 M is ________ V.

A) 0.56

B) 0.50

C) 0.44

D) 0.40

E) 0.52

Q3) Why are iron nails coated with zinc?

Q4) Give an example of an inert electrode.

Q5) What is the difference between a voltaic cell and an electrolytic cell?

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Chapter 21: Radioactivity and Nuclear Chemistry

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Q1) The amount of energy required to break apart the nucleus into its component nucleons is the

A) mass defect.

B) alpha decay.

C) nuclear binding energy.

D) electron capture.

E) beta decay.

Q2) 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

Q3) What is the "mass defect"?

Q4) List two problems associated with nuclear power.

Q5) Explain the concept of "magic numbers."

Q6) Define radioactivity.

Q7) Describe what is meant by the term "valley of stability"?

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Chapter 22: Organic Chemistry

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Q1) Give the product for the dehydration of CH₃CH₂CH₂OH.

A)

CH₃CH₂CH₂OCH₂CH₂CH ₃

B) CH₃CH=CH₂

C) CH₃CH=CHCH₃

D) CH₃CH₂CH=CH₂

E) CH₂=C=CH₂

Q2) Which of the following names is correct?

A) 2-methyl-1-butene

B) 3,3- dipropyl-4- propene

C) 2-ethyl-3- butene

D) 1- ethyl-2-pentene

E) None of the above are correct.

Q3) Which of the following compounds exhibits geometric isomerism?

A) CH₂=CH₂

B) CH₂=CCl₂

C) CBr₂=CHBr

D) CHCl=CHCl

E) (CH₃)₂C=CH-CH₃

Page 24

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Chapter 23: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Identify the geometry of [Pd I₄]^2-.

A) tetrahedral

B) square planar

C) linear

D) octahedral

E) trigonal planar

Q2) How many unpaired electrons would you expect for the complex ion: [Co(OH)₆]^3-?

A) 1

B) 2

C) 4

D) 0

E) 3

Q3) Explain how EDTA is used to treat lead poisoning.

Q4) Identify the geometry of [Ag(NH₃)₂]⁺.

A) tetrahedral

B) square planar

C) linear

D) bent

E) trigonal pyramidal

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