Introduction to Chemistry Exam Practice Tests - 1830 Verified Questions

Introduction to Chemistry Exam Practice Tests

Course Introduction

Introduction to Chemistry provides a foundational understanding of the basic principles and concepts of chemistry. This course covers the structure and properties of matter, atomic theory, chemical bonding, stoichiometry, states of matter, chemical reactions, and the fundamentals of thermochemistry. Students will develop problem-solving skills and scientific reasoning through laboratory experiments, demonstrations, and real-world applications. Designed for students with little or no prior background in chemistry, this course prepares learners for more advanced studies in science and related fields.

Recommended Textbook

Principles of General Chemistry 2nd Edition by Martin Silberberg

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23 Chapters

1830 Verified Questions

1830 Flashcards

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Page 2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) Briefly explain the relationship between hypothesis and experiment in the scientific method.

Answer: A hypothesis should be capable of leading to a prediction which is testable by experiment.If the experimental result differs from the prediction,the hypothesis should be modified.

Q2) Which of the following represents the largest volume?

A)10,000 µL

B)1000 pL

C)100 mL

D)10 nL

E)10 cm³

Answer: C

Q3) The S.I.prefix mega- (M)means

A)10^-6

B)10^-3

C)10³

D)10⁶

E)10⁹

Answer: D

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Page 3

Chapter 2: The Components of Matter

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Sample Questions

Q1) Name the three important "laws" that were accounted for by Dalton's atomic theory. Answer: Laws of conservation of mass;definite composition;multiple proportions

Q2) For each of the following names,write down the corresponding formula,including charge where appropriate (atomic numbers and mass numbers are not required):

a.zinc ion

b.nitrite ion

c.carbonic acid

d.cyanide ion

Answer: a.Zn²⁺

b.NO₂^- c.H₂CO₃ d.CN^-

Q3) Copper (Cu)is a transition metal.

A)True

B)False

Answer: True

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4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Potassium chloride is used as a substitute for sodium chloride for individuals with high blood pressure.Identify the limiting reactant and determine the mass of the excess reactant remaining when 7.00 g of chlorine gas reacts with 5.00 g of potassium to form potassium chloride.

A)Potassium is the limiting reactant;2.47 g of chlorine remain.

B)Potassium is the limiting reactant;7.23 g of chlorine remain.

C)Chlorine is the limiting reactant;4.64 g of potassium remain.

D)Chlorine is the limiting reactant;2.70 g of potassium remain.

E)No limiting reagent: the reactants are present in the correct stoichiometric ratio.

Answer: A

Q2) Hydroxylamine nitrate contains 29.17 mass % N,4.20 mass % H,and 66.63 mass O.If its molar mass is between 94 and 98 g/mol,what is its molecular formula?

A)NH₂O₅

B)N₂H₄O₄

C)N₃H₃O₃

D)N₄H₈O₂

E)N₂H₂O₄

Answer: B

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Page 5

Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) How many moles of ions are released when 1.6 mol of ammonium phosphate, (NH₄)₃PO₄,is dissolved in water?

A)0.40 mol

B)1.6 mol

C)3.2 mol

D)4.8 mol

E)6.4 mol

Q2) How many moles of ions are released when 0.27 mol of cobalt(II)chloride,CoCl₂,is dissolved in water?

A)0.81 mol

B)0.54 mol

C)0.27 mol

D)0.18 mol

E)0.090 mol

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

A)True

B)False

Q4) Complete and balance the equation for the following acid-base reaction. Ca(OH)₂(aq)+ HCl(aq) \(\to\)

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) A sample of oxygen gas has its absolute temperature halved while the pressure of the gas remained constant.If the initial volume is 400 mL,what is the final volume?

A)20 mL

B)133 mL

C)200 mL

D)400 mL

E)800 mL

Q2) Nitrogen dioxide is a red-brown gas that is responsible for the color of photochemical smog.A sample of nitrogen dioxide has a volume of 28.6 L at 45.3°C and 89.9 kPa.What is its volume at STP?

A)21.8 L

B)27.6 L

C)29.6 L

D)37.6 L

E)153 L

Q3) Dry air is approximately 78% nitrogen,21% oxygen and 1% argon,by number of molecules.What is the partial pressure of oxygen in a sample of dry air,if atmospheric pressure is 751 torr?

Q4) State Avogadro's Law.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Nitric acid,which is among the top 15 chemicals produced in the United States,was first prepared over 1200 years ago by heating naturally occurring sodium nitrate (called saltpeter)with sulfuric acid and collecting the vapors produced.Calculate \(\Delta\)H°_rxn for this reaction.\(\Delta\)H°_f [NaNO₃(s)] = -467.8 kJ/mol;\(\Delta\)H°_f [NaHSO₄(s)] = -1125.5 kJ/mol;\(\Delta\)H°_f [H₂SO₄(l)= -814.0 kJ/mol;\(\Delta\)H°_f [HNO₃(g)] = -135.1 kJ/mol

NaNO₃(s)+ H₂SO₄(l) \(\to\)

NaHSO₄(s)+ HNO₃(g)

A)-644.2 kJ

B)-291.4 kJ

C)-21.2 kJ

D)21.2 kJ

E)644.2 kJ

Q2) The only way in which a system can do work on the surroundings is by expansion against the external pressure.

A)True

B)False

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) The energy of an electron in the hydrogen atom is determined by

A)the principal quantum number (n)only.

B)the angular momentum quantum number (l )only.

C)the principal and angular momentum quantum numbers (n & l ).

D)the principal and magnetic quantum numbers (n & m_l).

E)the principal,angular momentum and magnetic quantum numbers.

Q2) Who proposed the principle which states that one cannot simultaneously know the exact position and velocity of a particle?

A)Einstein

B)Planck

C)Heisenberg

D)Compton

E)de Broglie

Q3) In not more than three lines for each answer,briefly outline one important scientific contribution of each of the following.

a.Planck

b.de Broglie

c.Heisenberg

Q4) What is the speed of an electron in m/s if its wavelength is 0.155 nm?

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) The electronic structure

1s²2s²2p⁶3s²3p⁶4s²3d⁸ refers to the ground state of A)Kr

B)Ni

C)Fe

D)Pd

E)None of these choices is correct.

Q2) Electron affinities of neutral atoms may be positive or negative.

A)True

B)False

Q3) Select the diamagnetic ion.

A)Cu²⁺

B)Ni²⁺

C)Cr³⁺

D)Sc³⁺

E)Cr²⁺

Q4) Describe how the ionization energies of main group elements vary across a period in the periodic table,using a diagram if necessary.

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Select the correct formula for a compound formed from barium and nitrogen.

A)BaN

B)BaN₂

C)Ba₂N₃

D)Ba₂N

E)Ba₃N₂

Q2) For which of the following elements (in their normal,stable forms)would it be correct to describe the bonding as involving "electron pooling"?

A)hydrogen

B)helium

C)sulfur

D)iodine

E)aluminum

Q3) Which of the following compounds displays the greatest ionic character in its bonds?

A)NO₂

B)CO₂

C)H₂O

D)HF

E)NH₃

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the ideal bond angles around carbon in C₂I₂ using the molecular shape given by the VSEPR theory.

A)90°

B)109°

C)120°

D)180°

E)None of these choices is correct.

Q2) What is the molecular shape of SCl₃F as predicted by the VSEPR theory?

A)linear

B)bent

C)see-saw

D)T-shaped

E)trigonal pyramidal

Q3) What is the molecular shape of BCl₃ as predicted by the VSEPR theory?

A)linear

B)trigonal planar

C)bent

D)tetrahedral

E)trigonal pyramidal

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) A molecule with the formula AX₄E uses _________ to form its bonds.

A)sp² hybrid orbitals

B)sp³ hybrid orbitals

C)sp³d hybrid orbitals

D)sp³d² hybrid orbitals

E)None of these choices is correct.

Q2) According to molecular orbital (MO)theory,the twelve outermost electrons in the O₂ molecule are distributed as follows:

A)12 in bonding MOs,0 in antibonding MOs.

B)10 in bonding MOs,2 in antibonding MOs.

C)9 in bonding MOs,3 in antibonding MOs.

D)8 in bonding MOs,4 in antibonding MOs.

E)7 in bonding MOs,5 in antibonding MOs.

Q3) In one sentence state how molecular orbitals are usually obtained.

Q4) In molecular orbital theory,combination of two atomic orbitals produces two molecular orbitals.

A)True

B)False

Q5) In one sentence state the basic principle of valence bond theory.

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

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Sample Questions

Q1) Ammonia's unusually high melting point is the result of

A)dipole-dipole forces.

B)London dispersion forces.

C)hydrogen bonding.

D)covalent bonding.

E)ionic bonding.

Q2) The strongest intermolecular interactions between hydrogen fluoride (HF)molecules arise from

A)dipole-dipole forces.

B)London dispersion forces.

C)hydrogen bonding.

D)ion-dipole interactions.

E)ionic bonds.

Q3) Strontium metal crystallizes in a cubic unit cell which has an edge length of 612 pm.If the mass of an atom of Sr is 87.62 amu,and the density of Sr metal is 2.54 g/cm³,calculate the number of atoms per unit cell.(1 amu = 1.661 \(\times\) 10^-24 g. )

Q4) In cubic closest packing,the unit cell is body-centered cubic.

A)True B)False

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Chapter 13: The Properties of Solutions

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Sample Questions

Q1) If a solute dissolves in an endothermic process

A)H bonds must exist between solvent and solute.

B)strong ion-dipole forces must exist in the solution.

C)the solute must be a gas.

D)the entropy of the solution is immaterial.

E)the entropy of the solution must be greater than that of its pure components.

Q2) Predict the van't Hoff factor (i)for each of the following solutes:

a.CaCl₂

b.Fe₂(SO₄)₃ c.NH₄Cl

Q3) Select the weakest electrolyte from the following set.

A)Na₂SO₄

B)KCl

C)CH₃CH₂COOH,propionic acid

D)CaCl₂

E)LiOH

Q4) Heats of solution may be either positive or negative.

A)True

B)False

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Chapter 14: The Main-Group Elements: Applying Principles of Bonding and Structure

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Sample Questions

Q1) Most of the alkali metal salts are soluble in water while many alkaline earth salts have very low solubilities.Why is this so?

A)The alkali metal cations are smaller than the alkaline earth cations and are more easily hydrated.

B)The alkali metals have lower ionization energies than alkaline earth elements.

C)The alkaline earth salts have much greater lattice energies than the alkali metal salts.

D)The alkaline earth metals have greater heats of atomization than the alkali metals.

E)Alkaline earth cations have very low heats of hydration.

Q2) Write balanced equations,showing all reactants and products,to represent a.the reaction of calcium metal with water.

b.the reaction of aluminum metal with oxygen gas.

Q3) Name two different classes (types)of hydride.

Q4) In the gas phase,AlCl₃ exists as molecules of Al₂Cl₆.

A)True

B)False

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Page 16

Chapter 15: Organic Compounds and the Atomic Properties of

Carbon

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Sample Questions

Q1) One characteristic of the monomers that form condensation polymers that is not common in monomers which form addition polymers is

A)the presence of pi bonds.

B)the presence of two functional groups.

C)the presence of alkyl side chains.

D)the ability to form free radicals.

E)the presence of sulfur atoms.

Q2) Carboxylic acids are weak acids.

A)True

B)False

Q3) Explain what is meant by "complementary" in the context of DNA strands.

Q4) The carbon atoms in a molecule of cyclohexane lie in the same plane. A)True B)False

Q5) The backbone of protein molecules consists of repeating N-C-C-O units. A)True B)False

Q7) Draw and name four alkenes with the formula C₄H₈. Page 17

Q6) Draw and name all stable molecules with the formula C₅H₁₂.

Q8) Name the three component parts of a nucleotide.

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) An elementary reaction is a simple,one-step process.

A)True

B)False

Q2) The gas-phase reaction CH₃NC \(\to\) CH₃CN has been studied in a closed vessel,and the rate equation was found to be: rate = -\(\Delta\)[CH₃NC]/\(\Delta\)t = k[CH₃NC].Which one of the following actions is least likely to cause a change in the rate of the reaction?

A)lowering the temperature

B)adding a catalyst

C)using a larger initial amount of CH₃NC in the same vessel

D)using a bigger vessel,but the same initial amount of CH₃NC

E)continuously removing CH₃CN as it is formed

Q3) Which of the following sets of units could be appropriate for a zero-order rate constant?

A)s^-1

B)L mol^-1 s^-1

C)L² mol^-2 s^-1

D)L³ mol^-3 s^-1

E)mol L^-1 s^-1

Page 19

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) For a solution equilibrium,a change in concentration of a reactant or product does not change K_c.

A)True

B)False

Q2) When a chemical system is at equilibrium,

A)the concentrations of the reactants are equal to the concentrations of the products.

B)the concentrations of the reactants and products have reached constant values.

C)the forward and reverse reactions have stopped.

D)the reaction quotient,Q,has reached a maximum.

E)the reaction quotient,Q,has reached a minimum.

Q3) Once a reaction system reaches equilibrium,the concentrations of reactions and products no longer change.

A)True

B)False

Q4) Unless \(\Delta\)H°_rxn = 0,a change in temperature will affect the value of the equilibrium constant K_c.

A)True

B)False

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Page 20

Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) What is the [H₃O⁺] for a solution at 25°C that has pOH = 5.640?

A)2.34 \(\times\) 10^-4 M

B)2.29 \(\times\) 10^-6 M

C)4.37 \(\times\) 10^-9 M

D)4.27 \(\times\) 10^-11 M

E)8.360 M

Q2) A 0.050 M solution of the weak acid HA has [H₃O⁺] = 3.77 \(\times\) 10^-4 M.What is the K_a for the acid?

A)7.5 \(\times\) 10^-3 M

B)2.8 \(\times\) 10^-6 M

C)7.0 \(\times\) 10^-7 M

D)7.0 \(\times\) 10^-8 M

E)2.6 \(\times\) 10^-11 M

Q3) Define an acid according to the Lewis theory.Write a balanced equation in support of this definition,in which the acid is neither an Arrhenius acid nor a Brønsted-Lowry acid.

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

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Page 21

Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Calculate the solubility of silver phosphate,Ag₃PO₄,in pure water.K_sp = 2.6 \(\times\) 10^-18

A)4.0 \(\times\) 10^-5 M

B)1.8 \(\times\) 10^-5 M

C)4.0 \(\times\) 10^-6 M

D)1.5 \(\times\) 10^-6 M

E)< 1.0 \(\times\) 10^-6 M

Q2) What will be the effect of adding 0.5 mL of 0.1 M HCl to 100 mL of a phosphate buffer in which

[H₂PO₄^-] = [HPO₄^2-] = 0.35 M?

A)The pH will increase slightly.

B)The pH will increase significantly.

C)The pH will decrease slightly.

D)The pH will decrease significantly.

E)Since it is a buffer solution,the pH will not be affected.

Q3) Hydrochloric acid (0.100 mol L^-1,25.00 mL aliquot)is being titrated with sodium hydroxide of the same molarity.Calculate the solution pH after addition of 24.80 mL of the sodium hydroxide.Make any reasonable approximations.

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Page 22

Chapter

Direction of Chemical Reactions

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Sample Questions

Q1) For a chemical reaction to be spontaneous only at low temperatures,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)(\(\Delta\)G° > 0)

Q2) For a chemical reaction to be spontaneous only at high temperatures,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)(\(\Delta\)G° > 0)

Q3) For a given reaction,a change in the pressure may result in a change in the sign of \(\Delta\)G.

A)True

B)False

Q4) State the second and third laws of thermodynamics.

Page 23

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Sample Questions

Q1) Oxidation occurs at the cathode of a galvanic cell,but at the anode of an electrolytic cell.

A)True

B)False

Q2) A voltaic cell consists of a Hg/Hg₂²⁺ electrode (E°= 0.85 V)and a Sn/Sn²⁺ electrode (E°= -0.14 V).Calculate [Sn²⁺] if [Hg₂²⁺] = 0.24 M and E_cell = 1.04 V at 25°C.

A)0.0001 M

B)0.0007 M

C)0.005 M

D)0.03 M

E)0.05 M

Q3) In the electrolyte of an electrochemical cell,current is carried by electrons moving from the anode to the cathode.

A)True

B)False

Q4) For the reaction occurring in a voltaic (galvanic)cell,\(\Delta\)G > 0.

A)True

B)False

Page 24

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Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) In the presence of a strong octahedral ligand field,the number of unpaired electrons in Co(III)will be

A)0.

B)2.

C)4.

D)6.

E)None of these choices is correct.

Q2) What is the coordination number of cobalt in the complex ion [Co(en)Cl₄]^-? (en = ethylenediamine)

A)1

B)2

C)4

D)6

E)8

Q3) Write the formula for sodium tetracyanonickelate(II).

A)Na[Ni(CN)₄]

B)Na[Ni(CN)₄]₂

C)Na₂[Ni(CN)₄]

D)Na₄[Ni(CN)₄]

E)None of these choices is correct.

Page 25

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) When an electron and its anti-particle,a positron,collide,they annihilate each other.Calculate the energy released in this process,in J.(The positron mass is the same as the electron mass,namely 9.11 \(\times\) 10^-31 kg. )

Q2) A 9.52 \(\times\) 10^-5 mol sample of rubidium-86 emits 8.87 \(\times\) 10¹⁶ \(\beta\) particles in one hour.What is the half-life of rubidium-86?

A)2.23 \(\times\) 10^-3 h

B)1.55 \(\times\) 10^-3 h

C)448 h

D)645 h

E)None of these choices is correct.

Q3) An isotope with a low value of N/Z will generally decay through

A)(\(\alpha\) decay.)

B)9\(\beta\) decay.)

C)(\(\gamma\) decay.)

D)electron capture.

E)spontaneous fission.

Q4) Briefly explain the relationship between the rad and the rem as units of radiation dosage.

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