Physical Science Pre-Test Questions - 3207 Verified Questions

Physical Science

Pre-Test Questions

Course Introduction

Physical Science is an interdisciplinary course that explores the fundamental concepts and principles governing the natural world, focusing primarily on physics and chemistry. The course covers topics such as motion, forces, energy, matter, atomic structure, chemical reactions, and the laws that describe physical interactions in the universe. Through a combination of theoretical learning and hands-on experiments, students gain a foundational understanding of scientific inquiry, problem-solving skills, and how physical science impacts everyday life and various technological advancements. This course serves as a stepping stone for further studies in scientific and engineering fields.

Recommended Textbook Chemistry A Molecular Approach 2nd Canadian Edition by Nivaldo J. Tro

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Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) How many significant figures are in the measurement 0.0005890 g?

A)4

B)5

C)6

D)7

E)8

Answer: A

Q2) What answer should be reported,with the correct number of significant figures,for the following calculation?

1.682 × 10^-3+ 4.311 × 10^-2

A)4.44792 × 10^-2

B)4.44792 × 10^-3

C)4.48 × 10^-2

D)4.479 × 10^-3

E)4.479 × 10^-2

Answer: E

Q3) Define the law of the conservation of energy.

Answer: Energy is neither created nor destroyed,only transferred.

Q4) Define matter.

Answer: Matter is anything that occupies space and has mass.

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Bromine has two naturally occurring isotopes,Br-79 and Br-81.The mass of Br-79 is 78.92 u with a percent abundance of 50.69%.If the atomic mass of bromine is 79.904 u,calculate the mass and percent abundance of Br-81.

A)80.916 u and 50.69%

B)78.921 u and 49.31%

C)79.904 u and 49.31%

D)80.916 u and 49.31%

E)80.028 u and 50.31%

Answer: D

Q2) What element is defined by the following information? p⁺ = 20 n° = 20

e^- = 20

A)zirconium

B)calcium

C)potassium

D)neon

E)argon

Answer: B

Q3) The number 6.022 × 10²³ is known as ________.

Answer: Avogadro's number

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Page 4

Chapter 3: Molecules,compounds,and Nomenclature

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Sample Questions

Q1) A certain alcohol contains only three elements,carbon,hydrogen,and oxygen.Combustion of a 50.00 g sample of the alcohol produced 95.50 g of CO₂ and 58.70 g of H₂O.What is the empirical formula of the alcohol?

Answer: C₂H₆O

Q2) Which one of the following compounds contains ionic bonds?

A)SrO

B)HBr

C)PBr₃

D)SiO₂

Answer: A

Q3) Which of the following is a molecular element?

A)Mg

B)Ar

C)Xe

D)I

E)Li

Answer: D

Q4) Define empirical formula.

Answer: An empirical formula gives relative numbers of atoms of each element.

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Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) What are the required coefficients to properly balance the following chemical reaction? N₂H₄(l) NH₃(g)+ N₂(g)

A)3,4,1

B)2,3,2

C)3,4,4

D)1,4,3

E)2,2,1

Q2) What element is undergoing oxidation (if any)in the following reaction? Zn(s)+ 2AgNO₃(aq) Zn(NO₃)₂(aq)+ 2Ag(s)

A)Zn

B)N

C)O

D)Ag

E)This is not an oxidation-reduction reaction.

Q3) How many grams of CH₃ OH must be added to water to prepare 150 mL of a solution that is 2.0 mol L^-1 CH₃ OH?

Q4) Describe the difference between complete ionic and net ionic equations.

Q5) Give an example of an inert electrode.

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Chapter 5: Gases

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Sample Questions

Q1) A gas is at 35.0 °C and 4.50 L.What is the temperature at 9.00 L?

A)343°C

B)70.0°C

C)616°C

D)1.16°C

E)17.5°C

Q2) A fixed amount of gas at 25.0 °C occupies a volume of 20.0 L when the pressure is 0.8386 bar.Use Charles's law to calculate the volume (L)the gas will occupy when the temperature is increased to \(131 ^ { \circ } \mathrm { C }\) while maintaining the pressure at 0.8386 bar.

A)14.8 L

B)105 L

C)3.82 L

D)22.4 L

E)27.1 L

Q3) Calculate the root mean square velocity of nitrogen molecules at 25 °C.

A)729 m s^-1

B)515 m s^-1

C)149 m s^-1

D)297 m s^-1

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Chapter 6: Thermochemistry

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Sample Questions

Q1) A 21.8 g sample of ethanol (C₂H₅OH)is burned in a bomb calorimeter according to the following reaction.If the temperature rises from 25.0 °C to 62.3 °C,determine the heat capacity of the calorimeter.The molar mass of ethanol is 46.07 g \(\mathrm { mol } ^ { - 1 }\) . C₂H₅OH(l)+ 3O₂(g)? 2CO₂(g)+ 3H₂O(g) \(D _ { \mathrm { r } }\) U= -1235 kJ mol^-1

A)4.99 kJ °C^-1

B)5.65 kJ °C^-1

C)63.7 kJ °C^-1

D)33.1 kJ °C^-1

E)15.7 kJ °C^-1

Q2) Explain the difference between H and U.

Q3) Calculate the work,w,gained or lost by the system when a gas expands from 15 L to 40 L against a constant external pressure of 1.5199 bar.

A)-6.1 kJ

B)-3.8 kJ

C)+3.8 kJ

D)+6.1 kJ

Q4) The standard state of carbon is ________.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Identify the correct values for a 2p sublevel.

A)n = 3,l = 1,m_l = 0

B)n = 2,l = 1,m_l = -2

C)n = 1,l = 0,m_l = 0

D)n = 2,l = 1,m_l = 0

E)n = 4,l = -1,m_l = -2

Q2) Identify the correct values for a 1s sublevel.

A)n = 3,l = 1,m_l = 0

B)n = 2,l = 1,m_l = -2

C)n = 1,l = 0,m_l = 0

D)n = 2,l = 0,m_l = 0

E)n = 4,l = -1,m_l = -2

Q3) The element that corresponds to the electron configuration

1s²2s²2p² is ________.

A)nitrogen

B)lithium

C)hydrogen

D)phosphorus

E)carbon

Q4) How are 3p orbitals different from the 2p orbitals?

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) How many valence electrons does an atom of Ti possess?

A)2

B)4

C)6

D)8

E)0

Q2) What period 3 element has the following ionization energies (all in kJ mol^-1)? IE₁ = 1012 IE₂ = 1900 IE₃= 2910 IE₄= 4960 IE₅= 6270 IE₆= 22 200

A)Si

B)S

C)P

D)Cl

E)Mg

Q3) Why is the first ionization energy of sulfur smaller than the first ionization energy of phosphorus?

Q4) List the noble gas that has the highest ionization energy.

Q5) Why does the size of the transition elements stay roughly the same as you move across a period?

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Identify the longest bond.

A)single covalent bond

B)double covalent bond

C)triple covalent bond

D)All of the above bonds are the same length.

Q2) Explain why the lattice energy of MgS is approximately four times as large as that of NaCl.

Q3) Place the following elements in order of increasing electronegativity. K\(\quad\) Cs\(\quad\) P

A)P < K < Cs

B)K < P < Cs

C)Cs < P < K

D)Cs < K < P

E)P < Cs < K

Q4) Identify the compound with the lowest magnitude of lattice energy.

A)KCl

B)KBr

C)SrO

D)CaO

Q5) Define bond energy.

11

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Chapter 10: Chemical Bonding II: Molecular Shapes, valence

Bond Theory, and Molecular Orbital Theory

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Sample Questions

Q1) Determine the electron geometry (eg)and molecular geometry (mg)of CH₃⁺.

A)eg = tetrahedral,mg = tetrahedral

B)eg = tetrahedral,mg = trigonal pyramidal

C)eg = trigonal planar,mg = bent

D)eg = trigonal planar,mg = trigonal planar

E)eg = tetrahedral,mg = trigonal planar

Q2) What is the molecular geometry of TeCl₄?

A)seesaw

B)square planar

C)square pyramidal

D)tetrahedral

Q3) Determine the electron geometry (eg),molecular geometry (mg),and polarity of HBrO₂.

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

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

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Chapter 11: Liquids, solids, and Intermolecular Forces

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Sample Questions

Q1) Choose the substance with the lowest viscosity in the liquid phase.

A)Cl₃CCCl₃

B)Cl₂CHCH₂Cl

C)Cl₂CHCHCl₂

D)ClCH₂CH₂Cl

E)Cl₃CCHCl₂

Q2) Define boiling.

A)A liquid becomes a gas.

B)A gas becomes a liquid.

C)A gas becomes a solid.

D)A solid becomes a gas.

E)A solid becomes a liquid.

Q3) Give the change in condition to go from a liquid to a gas.

A)increase heat or reduce pressure

B)increase heat or increase pressure

C)cool or reduce pressure

D)cool or increase pressure

E)none of the above

Q4) Define dynamic equilibrium between the liquid and gaseous phases.

Q5) Define volatile.

13

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Chapter 12: Solutions

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Sample Questions

Q1) Choose the aqueous solution with the highest vapour pressure.These are all solutions of nonvolatile solutes and you should assume ideal van't Hoff factors where applicable.

A)0.50 mol kg^-1 C₃H₈O₃

B)0.50 mol kg^-1C₂H₆O₂

C)0.50 mol kg^-1C₆H₁₂O₆

D)0.50 mol kg^-1C₁₂H₂₂O₁₁

E)They all have about the same vapour pressure.

Q2) Which cation in each set is expected to have the larger (more negative)hydration energy?

I \(\quad\)Be²⁺ or Ca²⁺

II \(\quad\)Rb⁺ or Zn²⁺

A)Be²⁺ in set I and Rb⁺ in set II

B)Be²⁺ in set I and Zn²⁺ in set II

C)Ca²⁺ in set I and Rb⁺ in set II

D)Ca²⁺ in set I and Zn²⁺ in set II

Q3) Define osmosis.

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) Give the preparation of rock candy.

Page 14

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) Given the following balanced equation,determine the rate of reaction with respect to [NOCl].If the rate of Cl₂ loss is 4.84 × 10^-2 mol L^-1 s^-1,what is the rate of formation of NOCl?

2NO(g)+ Cl₂(g) 2NOCl(g)

A)4.84 × 10^-2 mol L^-1 s^-1

B)2.42 × 10^-2 mol L^-1 s^-1

C)1.45 × 10^-1 mol L^-1 s^-1

D)9.68 × 10^-2 mol L^-1 s^-1

E)1.61 × 10^-2 mol L^-1 s^-1

Q2) Given the following balanced equation,determine the rate of reaction with respect to [NOCl].If the rate of NO loss is 1.68 × 10^-2 mol L^-1 s^-1,what is the rate of formation of NOCl?

2NO(g)+ Cl₂(g) 2NOCl(g)

A)1.68 × 10^-2 mol L^-1 s^-1

B)3.81 × 10^-2 mol L^-1 s^-1

C)5.05 × 10^-2 mol L^-1 s^-1

D)6.14 × 10^-2 mol L^-1 s^-1

E)1.45 × 10^-2 mol L^-1 s^-1

Q3) Define activation energy.

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) For the following reaction,what is n required in the conversion of K_c to K_p?

P₄(s)+

P₄O₁₀(s)

A)3

B)4

C)-4

D)5

E)-5

5O₂(g)

Q2) Consider the following reaction: NO(g)+ SO₃(g) NO₂(g)+ SO₂(g)

A reaction mixture initially contains 2.0 bar NO₂ and 1.0 bar SO₂.Determine the equilibrium pressure of SO₂ if K_p for the reaction at this temperature is 0.0118.

A)0.011 bar

B)0.037 bar

C)0.10 bar

D)0.0038 bar

E)0.067 bar

Q3) Explain dynamic equilibrium.Use the generic reaction A(g) B(g)to explain.

Q4) Why is a thermodynamic equilibrium constant unitless?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) What does the term amphoteric mean?

Q2) Which of the following bases is the strongest? The base is followed by its K_b.

A)(CH₃CH₂)₂NH,8.6 × 10^-4

B)CH₃NH₂,4.4 × 10^-4

C)C₆H₅NH₂,4.0 × 10^-10

D)NH₃,1.76 × 10^-5

E)C₅H₅N,1.7 × 10^-9

Q3) Determine the pOH of a 0.741 mol L^-1 KOH solution at 25 °C.

A)0.13

B)13.87

C)0.17

D)12.65

E)11.88

Q4) What is the difference between a strong and weak acid?

Q5) Which Bronsted-Lowry acid is not considered to be a strong acid in water?

A)HBr

B)HCl

C)HNO₂

D)HClO₄

Page 17

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Calculate the pH of a buffer that is 0.158 mol L^-1 HClO and 0.099 mol L^-1 NaClO.The K_a for HClO is 2.9 × 10^-8.

A)7.54

B)6.67

C)3.77

D)6.46

E)7.33

Q2) What is the pH of a solution prepared by mixing 25.00 mL of 0.10 mol L^-1

CH₃COOH with 25.00 mL of 0.010 mol L^-1

CH₃COONa? Assume that the volume of the solutions are additive and that K_a = 1.8 × 10^-5 for CH₃COOH.

A)2.87

B)3.74

C)4.75

D)5.74

Q3) Define a buffer.

Q4) Name the chemical compound that forms stalactites and stalagmites.

Q5) Give the name of the compound that is in antifreeze and is toxic to pets.

Q6) Explain the common ion effect with respect to molar solubility.

Page 18

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Identify the compound with the lowest standard Gibbs energy of formation.

A)NaCl(s)

B)N₂(g)

C)NO(g)

D)O₃(g)

E)Cl₂₍_g₎

Q2) Why can endothermic reactions be spontaneous?

Q3) Consider a reaction that has a negative _rH and a negative _rS.Which of the following statements is TRUE?

A)This reaction will be spontaneous only at high temperatures.

B)This reaction will be spontaneous at all temperatures.

C)This reaction will be nonspontaneous at all temperatures.

D)This reaction will be nonspontaneous only at high temperatures.

E)This reaction will be spontaneous only at absolute zero (0 Kelvin).

Q4) Under which of the following conditions would one mole of He have the highest entropy,S?

A)17 °C and 15 L

B)127 °C and 15 L

C)17 °C and 25 L

D)127 °C and 25 L

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Which of the following metals will dissolve in nitric acid but not hydrochloric acid?

A)Fe; E°(Fe²⁺/Fe)= -0.45V

B)Pb; E°(Pb²⁺/Pb)= -0.13V

C)Cu; E°(Cu²⁺/Cu)= +0.34V

D)Sn; E°(Sn²⁺/Sn)= -0.14V

E)Ni; E°(Ni²⁺/Ni)= -0.23V

Q2) Which of the following is the weakest reducing agent?

A)Br₂(l)

B)Au³⁺(aq)

C)Ag(s)

D)Br^- (aq)

E)Au(s)

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

Q4) What is electrolysis?

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Which of the following equations shows the correct relationship between the half-cell of a nuclide and the radioactive decay rate constant?

A)ln \(\frac { N _ { t } } { N _ { 0 } }\) = -\(\frac{k}{t}\)

B) \(\frac { N _ { t } } { N _ { 0 } }\) = - ln\(\frac{k}{t}\)

C) \(t 1 / 2\) = 0.693 × k

D) \(t 1 / 2\) =\(\frac { k } { 0.693 }\)

E) \(t 1 / 2\) =\(\frac { 0.693 } { k }\)

Q2) A geological sample is found to have a Pb-206/U-238 mass ratio of 0.337/1.00.Assuming there was no Pb-206 present when the sample was formed,how old is it? The half-life of U-238 is 4.5 × 10⁹ years.

A)7.3 × 10¹¹ years

B)1.4 × 10¹⁰ years

C)2.4 × 10¹⁰ years

D)2.1 × 10⁹ years

E)7.1 × 10⁹ years

Q3) List two problems associated with nuclear power.

Q4) Explain the concept of "magic numbers."

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Chapter 20: Organic Chemistry I: Structures

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Q1) What are the possible combinations of multiple bonds and rings if a compound's index of hydrogen deficiency equals 2?

A)two rings; one ring and a triple bond; two triple bonds

B)two rings and a double bond; three rings; one ring and two double bonds; two triple bonds

C)three rings; two rings and a double bond; one ring and two double bonds; two triple bonds; one ring and one triple bond

D)two rings; one ring and a double bond; two double bonds; one triple bond

E)two rings; two double bonds; a ring and a triple bond; one double bond and one triple bond

Q2) Which one of the following molecules is the most polar?

A)butane

B)butanoic acid

C)cyclohexane

D)ethanol

Q3) What are functional groups and how are they related to the concept of organic families? Provide examples.

Q4) What is the difference between the R/ S designation and d/l designation of chiral compounds?

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Chapter 21: Organic Chemistry II: Reactions

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Q1) Which of the following is the balanced chemical equation that represents the addition of Cl₂ to CH₃CH=CH₂?

A)CH₃CH=CH₂+Cl₂ CH₃CHClCH₃ + HCl

B)CH₃CH=CH₂+Cl₂ CH₃CCl=CHCl₊H₂ C)CH₃CH=CH₂+Cl₂

CH₃CHClCH₂Cl

D)CH₃CH=CH₂+2Cl₂ CH₃CHCl₂ + CH₂Cl₂

E)CH₃CH=CH₂+2Cl₂ CH₃CCl₂CHCl₂+ H₂

Q2) Choose the weak base from the compounds below.

A)(CH₃)₂NH

B)CH₃CO₂H

C)CH₃COCH₃ D)CH₃CH₃

E)CH₃CH₂F

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Chapter 22: Biochemistry

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Q1) Which one of the following amino acids does not contain a basic side chain?

A)arginine

B)histidine

C)lysine

D)threonine

Q2) Which one of the following amino acids contains a hydrophobic side chain?

A)valine

B)histidine

C)glutamine

D)glutamic acid

Q3) Identify the disaccharide.

A)glucose

B)cholesterol

C)fructose

D)cellulose

E)sucrose

Q4) Why are lipids well-suited as structural components of cell membranes?

Q5) Why does the melting point of fatty acids decrease with the number of double bonds present?

Q6) What is a codon?

24

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Chapter 23: Chemistry of the Nonmetals

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Q1) Refer to the above table.Use charge balance to determine the value of x in the amphibole

Mg₂(OH)₂Ca_x(Si₄O₁₁)<su b>2</sub>.

A)1

B)4

C)2

D)5

E)3

Q2) Determine the oxidation state of carbon in the ionic carbide CaC₂.

A)-2

B)+4

C)-1

D)+1

E)-4

Q3) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q4) Draw the Lewis structure for IF₄ and determine the geometry of the ion.

Q5) Describe the major production method for obtaining oxygen.

Q6) Why are phosphate compounds added to detergents?

Page 25

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Chapter 24: Metals and Metallurgy

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Q1) Which of the following describes slag?

A)a waste material that is the product of smelting a metal ore

B)the aqueous portion of the gangue separated from a metal ore

C)a low-density waste product,formed from the reaction of gangue with another substance,that can be easily separated from the metal ore of interest

D)a substance that reacts with the gangue to make a low-density liquid compound that can be easily separated from the metal of interest

E)a heavy precipitate containing by-products obtained during the purification process

Q2) Which of the following processes is used to form metal components from very small metal particles?

A)electrometallurgy

B)hydrometallurgy

C)pyrometallurgy

D)powder metallurgy

E)smelting

Q3) Why does a "two-phase" structure occur in a substitutional alloy?

Q4) Describe the difference between a substitutional alloy and an interstitial alloy.

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Which of the following species is diamagnetic?

A)an isolated gas-phase Cr³⁺ ion

B)a high-spin octahedral Cr²⁺ complex

C)an isolated gas-phase Cu²⁺ ion

D)a low-spin octahedral Co³⁺ complex

Q2) Identify the ion that is responsible for the red colour of rubies.

A)Cr⁺³

B)Cr⁺⁴

C)Cr⁺⁵

D)Cr⁺⁶

E)Cr⁺⁷

Q3) Identify the isomers that have ligands that coordinate to a metal in different orientations.

A)linkage isomers

B)cis-trans isomers

C)coordination isomers

D)optical isomers

E)stereoisomers

Q4) Why is +2 a common oxidation state for transition elements?

Q5) What is a coordinate covalent bond?

To view all questions and flashcards with answers, click on the resource link above. Page 27