Fundamentals of Chemistry Exam Practice Tests

Course Introduction

Fundamentals of Chemistry introduces students to the basic principles underlying the science of chemistry. The course covers atomic and molecular structure, chemical bonding, stoichiometry, states of matter, and an introduction to thermodynamics and chemical reactions. Emphasis is placed on developing problem-solving skills, understanding the periodic trends, and exploring the role of chemistry in everyday life and various scientific fields. Laboratory work complements theoretical learning by allowing students to observe chemical phenomena and practice essential techniques used in chemical analysis.

Recommended Textbook

Chemistry A Molecular Approach 4th Edition by Nivaldo J. Tro

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25 Chapters

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) The mass of a single arsenic atom is 1.244 × 10-22 g.This is the same mass as

A) 1.244 × 10-16 mg.

B) 1.244 × 10-25 kg.

C) 1.244 × 10-28 g.

D) 1.244 × 10-31 ng.

Answer: B

Q2) How many significant figures are in 4.930 × 104 m?

A) 2

B) 4

C) 3

D) 1

E) 5

Answer: B

Q3) A flash drive contains 4 gigabytes.How many bytes does it contain?

Answer: 4,000,000,000 bytes or 4,292,967,296 bytes,if someone is computer literate

Q4) Define the law of the conservation of energy.

Answer: Energy is neither created nor destroyed.

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Page 3

Chapter 2: Atoms and Elements

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Sample Questions

Q1) How many protons does the As3- ion possess?

A) 30

B) 36

C) 4

D) 8

E) 33

Answer: E

Q2) Rutherford proposed the

A) atomic bomb.

B) hydroelectric.

C) First Law of Conservation.

D) theory of explosives.

E) nuclear theory.

Answer: E

Q3) Helium belongs to the ________ group of the periodic table.

A) alkali metal

B) metal

C) transition metal

D) noble gas

Answer: D

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) What is the mass (in kg)of 6.89 × 1025 molecules of CO2? The molar mass of CO2 is 44.01 g/mol.

A) 3.85 kg

B) 5.04 kg

C) 2.60 kg

D) 3.03 kg

E) 6.39 kg

Answer: B

Q2) How many millimoles of Ca(NO3)2 contain 4.78 × 1022 formula units of Ca(NO3)2? The molar mass of Ca(NO3)2 is 164.10 g/mol.

A) 12.6 mmol Ca(NO3)2

B) 13.0 mmol Ca(NO3)2

C) 20.7 mmol Ca(NO3)2

D) 79.4 mmol Ca(NO3)2

E) 57.0 mmol Ca(NO3)2

Answer: D

Q3) Give the name for HNO2.

Answer: nitrous acid

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Page 5

Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) What is the concentration of barium ions in a 0.750 M BaCl2 solution?

A) 0.750 M

B) 0.375 M

C) 2.25 M

D) 1.50 M

E) 0.160 M

Q2) Define an electrolyte.

Q3) How many moles of Co2+ are present in 0.500 L of a 0.450 M solution of CoI2?

A) 0.950 moles

B) 0.500 moles

C) 1.10 moles

D) 0.525 moles

E) 0.225 moles

Q4) The mixing of which pair of reactants will result in a precipitation reaction?

A) Mg(NO3)2(aq) + K2CO3(aq)

B) Na2SO4(aq) + Cu(NO3)2(aq)

C) LiClO4(aq) + (NH4)2S(aq)

D) NH4Cl(aq) + NH4I(aq)

Q5) Describe the difference between complete ionic and net ionic equations.

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Chapter 5: Gases

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Sample Questions

Q1) Determine the density of NH3 gas at 435 K and 1.00 atm.

A) 2.10 g/L

B) 0.477 g/L

C) 0.321 g/L

D) 2.24 g/L

E) 0.851 g/L

Q2) While traveling,an airplane passenger placed a bag of potato chips in his checked luggage.When he opened the luggage,the bag had exploded and potato chips were everywhere in his luggage.Which gas law is this an example of?

A) Avogadro's Law

B) Ideal Gas Law

C) Charles's Law

D) Boyle's Law

E) Dalton's Law

Q3) An instrument used to measure the pressure of a gas in a laboratory is called a

A) barometer.

B) manometer.

C) sphygmomanometer.

D) spectrophotometer.

E) spectrometer.

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Which of the following statements is TRUE?

A) State functions do not depend on the path taken to arrive at a particular state.

B) DErxn can be determined using constant volume calorimetry.

C) Energy is neither created nor destroyed, excluding nuclear reactions.

D) Hrxn can be determined using constant pressure calorimetry.

E) All of the above are true.

Q2) When 0.500 mol of CH4(g)reacts with excess Cl2(g)at constant pressure according to the chemical equation shown below,354 kJ of heat are released.Calculate the value of H for this reaction,as written. 2 CH4(g)+ 3 Cl2(g) 2 CHCl3(l)+ 3 H2(g) H = ?

A) -1420 kJ

B) -708 kJ

C) +708 kJ

D) +1420 kJ

E) -710 kJ

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

Q4) Explain the difference between H and DE.

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Page 8

Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Electromagnetic radiation with a wavelength of 640 nm appears as orange light to the human eye.The frequency of this light is ________ s-1.

A) 4.688 × 1014

B) 4.688 × 105

C) 1.920 × 102

D) 1.920 × 1011

E) 2.133 × 10-15

Q2) Give all possible values of l if n = 3.

A) 3

B) +1/2

C) 0, 1, 2

D) -1, 0, 1

E) 1, 2, 3, 4

Q3) Which of the following quantum numbers describes the size and energy of an orbital?

A) magnetic quantum number

B) principal quantum number

C) angular momentum quantum number

D) spin quantum number

E) Schrödinger quantum number

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Choose the diamagnetic species from below.

A) Sn2

B) Br

C) As

D) Cr

E) None of the above is diamagnetic.

Q2) Identify the number of valence electrons in As-3.

A) 6

B) 7

C) 8

D) 5

E) 2

Q3) Identify the d-block element.

A) Th

B) W

C) Na

D) Te

E) None of the above

Q4) Define paramagnetic.

Q5) List the noble gas that has the highest ionization energy.

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Identify the ionic compound that is used as an antacid.

A) silver nitrate

B) barium sulfate

C) lithium carbonate

D) sodium chloride

E) sodium bicarbonate

Q2) Which of the following ionic compounds would be expected to have the highest lattice energy?

A) Li Br

B) Na Br

C) K Br

D) Rb Br

Q3) Identify the number of bonding pairs and lone pairs of electrons in H2O.

A) 2 bonding pairs and 3 lone pairs

B) 1 bonding pair and 3 lone pairs

C) 3 bonding pairs and 1 lone pair

D) 2 bonding pairs and 2 lone pairs

E) 4 bonding pairs and 0 lone pairs

Q4) Describe a covalent bond.

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) Draw the Lewis structure for the molecule CH3CH2CCH.How many sigma and pi bonds does it contain?

A) 11 sigma, 0 pi

B) 9 sigma, 1 pi

C) 8 sigma, 3 pi

D) 9 sigma, 2 pi

E) 8 sigma, 1 pi

Q2) Determine the electron geometry (eg)and molecular geometry (mg)of the underlined atom CH3OH.

A) eg = tetrahedral, mg = tetrahedral

B) eg = linear, eg = trigonal planar

C) eg = tetrahedral, mg = trigonal planar

D) eg = trigonal bipyramidal, mg = trigonal bipyramidal

E) eg = octahedral, mg = square planar

Q3) Is it possible for a molecule to be nonpolar even though it contains polar bonds?

Explain your answer and give an example.

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

Q5) Give the electron geometry,molecular geometry,and hybridization for both carbons in CH3COOH.

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) Define freezing.

A) the phase transition from solid to gas

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

Q2) List the compounds in order of decreasing boiling point. CH3O CH3 Rn CH3CHO

A) CH3O CH3 > CH3CHO > Rn

B) Rn > CH3O CH3 > CH3CHO

C) CH3O CH3 > Rn > CH3CHO

D) CH3CHO > CH3O CH3 > Rn

E) Rn > CH3CHO > CH3O CH3

Q3) Choose the molecule or compound that exhibits dipole-dipole forces as its strongest intermolecular force.

A) N2

B) SO2

C) H2O

D) CBr4

E) BCl3

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Identify the element that is the best insulator.

A) lead

B) silicon

C) germanium

D) carbon

E) tin

Q2) Give the edge length in terms of r for a simple cubic cell.

Q3) Identify an element that is NOT in nonoxide ceramics.

A) nitrogen

B) silicon

C) aluminum

D) boron

E) carbon

Q4) Identify an element that is NOT in oxide ceramics.

A) oxygen

B) silicon

C) aluminum

D) magnesium

Q5) Describe the difference between the conduction band and the valence band.

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Chapter 13: Solutions

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Sample Questions

Q1) A 150.0 mL sample of an aqueous solution at 25°C contains 15.2 mg of an unknown nonelectrolyte compound.If the solution has an osmotic pressure of 8.44 torr,what is the molar mass of the unknown compound?

A) 223 g/mol

B) 294 g/mol

C) 341 g/mol

D) 448 g/mol

E) 195 g/mol

Q2) Calculate the molality of a solution formed by dissolving 34.8 g of LiI in 500.0 mL of water.

A) 0.254 m

B) 0.130 m

C) 0.696 m

D) 0.260 m

E) 0.520 m

Q3) How many grams of KBr are required to make 350.mL of a 0.115 M KBr solution?

A) 0.338 g

B) 3.04 g

C) 4.79 g

D) 40.3 g

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) Identify the rate-determining step.

A) the slowest step

B) the faster step

C) the fast step

D) always the last step

E) always the second step

Q2) Nitrogen dioxide decomposes at 300°C via a second-order process to produce nitrogen monoxide and oxygen according to the following chemical equation. 2 NO2(g) 2 NO(g)+ O2(g).

A sample of NO2(g)is initially placed in a 2.50-L reaction vessel at 300°C.If the half-life and the rate constant at 300°C are 11 seconds and 0.54 M-1 s-1,respectively,how many moles of NO2 were in the original sample?

A) 0.17 mol

B) 0.42 mol

C) 5.9 mol

D) 15 mol

Q3) Define half-life.

Q4) What happens in the concentration of reactants and products during a chemical reaction?

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction,equilibrium concentrations,and equilibrium constant at a particular temperature.Determine the equilibrium pressure of CO. CO(g)+ 2 H2(g) CH3OH(l)Kp = 2.25 × 104 P(H2)eq = 0.52 atm

A) 8.3 × 104 atm

B) 1.2 × 10-5 atm

C) 6.25 × 10-3 atm

D) 8.5 × 10-5 atm

E) 1.6 × 10-4 atm

Q2) Determine the value of Kc for the following reaction if the equilibrium concentrations are as follows: [H2]eq = 0.14 M,[Br2]eq = 0.39 M,[HBr]eq = 1.6 M. H2(g)+ Br2(g) 2 HBr(g)

A) 2.1 × 10-2

B) 29

C) 47

D) 3.4 × 10-2

E) 8.7 × 10-2

Q3) Define Le Chatelier's Principle.

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Page 17

Chapter 16: Acids and Bases

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Sample Questions

Q1) If an equal number of moles of the weak acid HF and the strong base NaOH are added to water,is the resulting solution acidic,basic,or neutral?

A) acidic

B) basic

C) neutral

D) There is insufficient information provided to answer this question.

Q2) Which of the following is TRUE?

A) A neutral solution contains [H2O] = [H3O ].

B) An neutral solution does not contain any H3O+ or OH-.

C) An acidic solution has [H3O ] > [OH ].

D) A basic solution does not contain H3O+.

E) None of the above is true.

Q3) Which of the following is a strong acid?

A) H COOH

B) HF

C) HI

D) NH4+

E) H2S

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Page 18

Chapter 17: Aqueous Ionic Equilibrium

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Sample Questions

Q1) What is the pH of a solution prepared by mixing 30.00 mL of 0.10 M CH3CO2H with 30.00 mL of 0.030 M CH3CO2K? Assume that the volume of the solutions are additive and that Ka = 1.8 × 10-5 for CH3CO2H.

A) 2.87

B) 4.22

C) 4.75

D) 5.27

Q2) When titrating a weak monoprotic acid with NaOH at 25°C,the A) pH will be less than 7 at the equivalence point. B) pH will be equal to 7 at the equivalence point.

C) pH will be greater than 7 at the equivalence point.

D) titration will require more moles of base than acid to reach the equivalence point. E) titration will require more moles of acid than base to reach the equivalence point.

Q3) Give the name of the compound that is in antifreeze and is toxic to pets.

Q4) Identify the compound that is in stalactites and stalagmites.

Q5) Explain the common ion effect with respect to molar solubility.

Q6) Define a buffer.

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) The value of G° at 141.0°C for the formation of phosphorous trichloride from its constituent elements, P2(g)+ 3 Cl2(g) 2 PCl3(g) Is ________ kJ/mol.At 25.0°C for this reaction, H° is -720.5 kJ/mol, G° is -642.9 kJ/mol,and S° is -263.7 J/K.

A) -612.3

B) 3.65 × 104

C) 1.08 × 105

D) -683.3

E) -829.7

Q2) Estimate G°rxn for the following reaction at 825 K. 2 Hg(g)+ O2(g) 2 HgO(s) H°= -304.2 kJ; S°= -414.2 J/K

A) -2.63 kJ

B) -37 kJ

C) +37 kJ

D) +645 kJ

E) -645 kJ

Q3) Why can endothermic reactions be spontaneous?

Q4) Define the third law of thermodynamics.

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Page 20

Chapter 19: Electrochemistry

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Sample Questions

Q1) What statement is NOT true about standard electrode potentials?

A) E°cell is positive for spontaneous reactions.

B) Electrons will flow from more negative electrode to more positive electrode.

C) The electrode potential of the standard hydrogen electrode is exactly zero.

D) E°cell is the difference in voltage between the anode and the cathode.

E) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the standard hydrogen electrode and therefore has a positive E°.

Q2) 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

Q3) Why are iron nails coated with zinc?

Q4) Why is sugar water NOT a good conductor of current?

Q5) Give an example of an inert electrode.

Q6) Why,if we multiply a reaction by 2,don't we multiply its E°red by 2?

Q7) What is the difference between a voltaic cell and an electrolytic cell?

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Describe what changes occur during electron capture.

A) The mass number and atomic number decrease.

B) The mass number and atomic number increase.

C) The mass number is unchanged and the atomic number decreases.

D) The mass number decreases and the atomic number is unchanged.

E) The mass number and atomic number do not change.

Q2) Potassium-40 decays to argon-40 with a half-life of 1.27 × 109 yr.The age of a mineral sample that has a mass ratio of 40Ar to 40K of 0.812 is ________ yr.

A) 1.56 × 109

B) 1.02 × 109

C) 1.47 × 109

D) 7.55 × 108

E) 1.09 × 109

Q3) Define radioactivity.

Q4) Describe what is meant by the term "valley of stability"?

Q5) Explain the concept of "magic numbers."

Q6) How does a dosimeter measure exposure to radioactivity?

Q7) List two problems associated with nuclear power.

Q8) What is the "mass defect"?

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Chapter 21: Organic Chemistry

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Sample Questions

Q1) Which of the following names is correct?

A) 2-methyl-1-butene

B) 3,3-diethyl-4-propene

C) 2-ethyl-3-pentene

D) 1-propyl-2-pentene

E) None of the above is correct.

Q2) Nonsuperimposable mirror images are called

A) structural isomers.

B) achiral.

C) diastereomers.

D) enantiomers.

E) racemic mixture.

Q3) Which of the following alkane names is correct?

A) 1,2-dipropylhexane

B) 1-propylnonane

C) 4,5-dipropylpentane

D) 2-methyl-3-ethylheptane

E) All of the above names are correct.

Q4) What is meant by the term "structural isomer"? Draw a structural isomer of CH3-CH2-CH2-CH3.

Page 23

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Chapter 22: Biochemistry

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Sample Questions

Q1) Identify the source below that does NOT contain triglycerides.

A) corn oil

B) olive oil

C) whale oil

D) beef fat

E) diesel oil

Q2) Identify the sugar in RNA.

A) ribose

B) deoxyribose

C) fructose

D) glucose

E) galactose

Q3) How many chromosomes are in the nuclei of human cells?

A) 42

B) 45

C) 46

D) 49

E) 50

Q4) How do cis- and trans-fats differ?

Q5) What is a codon?

Page 24

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Chapter 23: Chemistry of the Nonmetals

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Q1) Which of the following statements is TRUE?

A) Each boron nitride unit contains six valence electrons.

B) Carbon nanotubes and boron nitride nanotubes are not similar.

C) Boron nitride contains BN units that are isoelectric with carbon.

D) Boron nitrides conduct an electrical current.

E) The size and electronegativity of a carbon atom are not close to the average of those properties for a boron atom and a nitrogen atom.

Q2) Identify an example of a nitrate.

A) LiN3

B) NaNO2

C) NaNO3

D) NO2

E) NO

Q3) What is the oxidation number of oxygen in H2O2?

A) -4

B) +2

C) -3

D) -1

E) +1

Q4) Describe the Frasch process for obtaining sulfur.

Page 25

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Chapter 24: Metals and Metallurgy

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Q1) Pyrometallurgy is the

A) refining of metal ores using reactions with aqueous solutions.

B) refining of metal ores using heat.

C) refining of metal ores using reduction reactions.

D) forming metal parts using heat and small crystals of metal.

E) forming metal parts using electricity.

Q2) Identify the substances which are homogeneous,naturally occurring,crystalline inorganic solids.

A) slag

B) minerals

C) ores

D) gangue

E) alloy

Q3) Identify the metal that has the most brightly colored ions.

A) magnesium

B) manganese

C) copper

D) chromium

E) cobalt

Q4) What is the difference between ferromagnetism and paramagnetism?

Page 26

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Which of the following compounds can exhibit fac-mer isomerism?

A) [Cr(H2O)4Br2]+

B) [Fe(CO)5ONO]2+

C) [Fe(CO)3(NH3)3]3+

D) [Cu(CO)5Br]+

E) [V(NH3)4(H2O)2]2+

Q2) What type of geometry (according to valence bond theory)does V exhibit in the complex ion: [V(NH3)4]2+?

A) square bipyramidal

B) see-saw

C) trigonal pyramidal

D) bent

E) square planar

Q3) Identify the geometry of [Zn(NH3)4]+2.

A) tetrahedral

B) octahedral

C) linear

D) square planar

E) trigonal planar

Q4) Explain how EDTA is used to treat lead poisoning.

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