Course Introduction
Physical Science
Chapter Exam Questions
Physical Science is an interdisciplinary course that explores the fundamental principles governing the physical world, focusing on the areas of physics and chemistry. Students examine key concepts such as matter, energy, motion, forces, atomic structure, chemical reactions, and the laws of nature. The course emphasizes scientific inquiry and problem-solving skills through hands-on experiments, demonstrations, and real-world applications. By connecting scientific theories to everyday phenomena, students develop a foundational understanding of how physical principles shape the universe and influence technological advancements.
Recommended Textbook Principles of Chemistry A Molecular Approach 3rd Edition by Nivaldo J. Tro
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Page 2
Chapter 1: Matter, Measurement, and Problem Solving
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Sample Questions
Q1) Which of the following is an example of physical change?
A)Sugar is dissolved in water.
B)Coffee is brewed.
C)Dry ice sublimes.
D)Ice (solid water)melts.
E)All of these are examples of physical change.
Answer: E
Q2) A physical change
A)occurs when iron rusts.
B)occurs when sugar is heated into caramel.
C)occurs when glucose is converted into energy within your cells.
D)occurs when water is evaporated.
E)occurs when propane is burned for heat.
Answer: D
Q3) Give the composition of water.
A)two hydrogen atoms and two oxygen atoms
B)one hydrogen atom and one oxygen atom
C)two hydrogen atoms and one oxygen atom
D)one hydrogen atom and two oxygen atoms
Answer: C
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Chapter 2: Atoms and Elements
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Sample Questions
Q1) Why do the isotopes of the same element have the same atomic size?
Answer: Isotopes only differ in the number of neutrons contained within the nucleus.Since the size of an atom is determined by the electrons,isotopes of the same element should be the same size.
Q2) Predict the charge that a calcium ion would have.
A)6-
B)2-
C)3+
D)2+
E)1+
Answer: D
Q3) How many electrons are in nickel?
A)28
B)30
C)31
D)30.7
E)58.7
Answer: A
Q4) Give the name of the element whose symbol is Na.
Answer: sodium
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Chapter 3: Molecules, Compounds and Chemical
Equations
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Sample Questions
Q1) What is the empirical formula for C<sub>4</sub>H<sub>10</sub>O<sub>2</sub>?
A)C<sub>2</sub>H<sub>5</sub>O
B)CHO
C)C<sub>2</sub>H<sub>4</sub>O
D)CHO<sub>2</sub>
E)CH<sub>2</sub>O
Answer: A
Q2) Which of the following is a molecular element?
A)neon
B)lithium
C)selenium
D)magnesium
E)titanium
Answer: C
Q3) An aqueous solution of H<sub>2</sub>S is named
A)hydrosulfuric acid.
B)hydrosulfurous acid.
C)sulfuric acid.
D)sulfurous acid.
Answer: A
Page 5
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Chapter 4: Chemical Quantities and Aqueous Reactions
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Sample Questions
Q1) Calcium oxide reacts with water in a combination reaction to produce calcium hydroxide:
CaO(s)+ H<sub>2</sub>O(l) Ca(OH)<sub>2</sub>(s)
A 4.50-g sample of CaO is reacted with 4.34 g of H<sub>2</sub>O.How many grams of water remain after the reaction is complete?
A)0.00
B)0.00892
C)2.90
D)1.04
E)0.161
Q2) How many milliliters of a 9.0 M H<sub>2</sub>SO<sub>4</sub> solution are needed to make 0.35 L of a 3.5 M solution?
A)0.14 mL
B)0.90 mL
C)140 mL
D)900 mL
Q3) How many grams of H<sub>3</sub>PO<sub>4</sub> are in 265 mL of a 1.50 M solution of H<sub>3</sub>PO<sub>4</sub>?
Q4) How can you tell if a reaction is an oxidation-reduction reaction?
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Chapter 5: Gases
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Q1) How many molecules of CO<sub>2</sub> are contained in a 10.0 L tank at 7.53 atm and 485 K?
A)1.89 × 10<sup>24</sup> molecules
B)1.14 × 10<sup>24</sup> molecules
C)8.32 × 10<sup>24</sup> molecules
D)4.89 × 10<sup>24</sup> molecules
E)3.63 × 10<sup>24</sup> molecules
Q2) How many liters of hydrogen gas can be generated by reacting 9.25 grams of barium hydride with water at 20°C and 755 mm Hg pressure according to the chemical equation shown below?
BaH<sub>2</sub>(s)+ 2
H<sub>2</sub>O(l) Ba(OH)<sub>2</sub>(aq)+ 2 H<sub>2</sub>(g)
A)0.219 L
B)0.799 L
C)1.60 L
D)3.21 L
Q3) Why does the rate of effusion increase with a decrease in the molar mass?
Q4) Why does hot air rise?
Q5) Why doesn't Dalton's Law of Partial Pressures depend on the identity of the gases present?
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Chapter 6: Thermochemistry
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Sample Questions
Q1) The specific heat capacity of liquid mercury is 0.14 J/g K.How many joules of heat are needed to raise the temperature of 5.00 g of mercury from 15.0°C to 36.5°C?
A)7.7 × 10<sup>2</sup> J
B)15 J
C)36 J
D)0.0013 J
E)1.7 J
Q2) A 21.8 g sample of ethanol (C<sub>2</sub>H<sub>5</sub>OH)is burned in a bomb calorimeter,according to the following reaction.If the temperature rises from 25.0 to 62.3°C,determine the heat capacity of the calorimeter.The molar mass of ethanol is 46.07 g/mol.
C<sub>2</sub>H<sub>5</sub>OH(l)+ 3 O<sub>2</sub>(g) 2 CO<sub>2</sub>(g)+ 3
H<sub>2</sub>O(g) H°<sub>rxn </sub>= -1235 kJ
A)4.99 kJ/°C
B)5.65 kJ/°C
C)63.7 kJ/°C
D)33.1 kJ/°C
E)15.7 kJ/°C
Q3) Give the equation to calculate the enthalpy of a system.
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Chapter 7: The Quantum-Mechanical Model of the Atom
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Sample Questions
Q1) Calculate the frequency of light associated with the transition from n = 2 to n = 3 in the hydrogen atom.
A)2.19 × 10<sup>14</sup> s<sup>-1</sup>
B)5.59 × 10<sup>14</sup> s<sup>-1</sup>
C)4.57 × 10<sup>14</sup> s<sup>-1</sup>
D)1.79 × 10<sup>14</sup> s<sup>-1</sup>
E)3.28 × 10<sup>14</sup> s<sup>-1</sup>
Q2) n = 5 to n = 3
A)103 nm
B)657 nm
C)7460 nm
D)122 nm
E)1280 nm
Q3) Which statement is correct about orbital?
A)An orbital can have maximum two electrons.
B)A p orbital can have six electrons.
C)An s orbital has no shape.
D)All p orbitals have same orientation.
Q4) Why do atoms only emit certain wavelengths of light when they are excited? (Why do line spectra exist?)
Page 9
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Chapter 8: Periodic Properties of the Elements
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Sample Questions
Q1) valence electrons
A)0
B)2
C)electrons in the outermost shell
D)1
E)electrons in completed shells
Q2) Of the following,which element has the highest first ionization energy?
A)Sr
B)Rb
C)Na
D)Ca
Q3) Place the following in order of increasing metallic character.
Br Cs Se
A)Br < Se < Cs
B)Se < Br < Cs
C)Cs < Br < Se
D)Cs < Se < Br
E)Br < Cs < Se
Q4) Give the ground state electron configuration for Cd<sup>+2</sup>.
Q5) Define electron affinity.
10
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Chapter 9: Chemical Bonding I: Lewis Theory
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Sample Questions
Q1) Liquid water when boiled at 100<sup>°</sup>C gets converted to water vapor.Which of the following statements is correct for this phase transition?
A)The covalent bonds between hydrogen and oxygen atoms are broken.
B)The intermolecular hydrogen bonds between water molecules are broken.
C)The covalent bonds between hydrogen and oxygen atoms become weaker.
D)The bond distance between hydrogen and oxygen atoms increases.
Q2) Identify an ionic bond.
A)Electrons are pooled.
B)Electrons are shared.
C)Electrons are transferred.
D)Protons are gained.
E)Electrons are lost.
Q3) Give the number of valence electrons for XeI<sub>2</sub>.
A)22
B)20
C)18
D)24
Q4) Define electronegativity.
Q5) Describe a covalent bond.
Q6) Describe the difference between a pure covalent bond and a polar covalent bond.
Page 11
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Chapter 10: Chemical Bonding II: Molecular Shapes,
Valence Bond Theory, and
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Sample Questions
Q1) Describe a square planar shape.
Q2) What geometric arrangement of charge clouds is expected for an atom that has four charge clouds?
A)trigonal bipyramidal
B)octahedral
C)tetrahedral
D)square planar
Q3) What is the molecular geometry of SBr<sub>4</sub>?
A)seesaw
B)square planar
C)square pyramidal
D)tetrahedral
Q4) Choose the compound below that contains at least one polar covalent bond,but is nonpolar.
A)HCN
B)CF<sub>4</sub>
C)SeBr<sub>4</sub>
D)ICl<sub>3</sub>
E)Both B and C are nonpolar and contain a polar covalent bond.
Page 12
Q5) Explain why oil and water do not mix.
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Chapter 11: Liquids, Solids, and Intermolecular Forces
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Q1) Choose the molecule or compound that exhibits dipole-dipole forces as its strongest intermolecular force.
A)H<sub>2</sub>
B)SO<sub>2</sub>
C)NH<sub>3</sub>
D)CF<sub>4</sub>
E)BCl<sub>3</sub>
Q2) Choose the substance with the highest viscosity.
A)(CH<sub>3</sub>CH<sub>2</sub>)<sub>2</sub>CO
B)C<sub>2</sub>H<sub>4</sub>Cl<sub>2</sub>
C)HOCH<sub>2</sub>CH<sub>2</sub>CH<sub>2</sub>CH<sub>2</sub>OH
D)CF<sub>4</sub>
E)C<sub>6</sub>H<sub>14</sub>
Q3) Why is the H<sub>vap</sub> higher than H<sub>fus </sub>for any given compound?
Q4) Choose the substance with the lowest boiling point.
A)H<sub>2</sub>S
B)NBr<sub>3</sub>
C)F<sub>2</sub>
D)CF<sub>2</sub>H<sub>2</sub>
E)H<sub>2</sub>O<sub>2</sub>
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Chapter 12: Solutions
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Q1) Which aqueous solution has the lowest boiling point?
A)0.10 m AlCl<sub>3</sub>
B)0.10 m NaCl
C)0.10 m BaCl<sub>2</sub>
D)0.10 m C<sub>6</sub>H<sub>12</sub>O<sub>6</sub>
Q2) Which of the following compounds will be most soluble in pentane (C<sub>5</sub>H<sub>12</sub>)?
A)pentanol
(CH<sub>3</sub>CH<sub>2</sub>CH<sub>2</sub>CH<sub>2</sub>CH<sub>2</sub>OH)
B)benzene (C<sub>6</sub>H<sub>6</sub>)
C)acetic acid (CH<sub>3</sub>CO<sub>2</sub>H)
D)ethyl methyl ketone (CH<sub>3</sub>CH<sub>2</sub>COCH<sub>3</sub>)
E)None of these compounds should be soluble in pentane.
Q3) 0.050 m C<sub>6</sub>H<sub>12</sub>O<sub>6 </sub>(aqueous)
A)largest van't Hoff factor
B)solution of ionic compound with highest freezing point
C)solution with T<sub>b</sub> = 0.026°C
D)highest boiling point
E)solution that is most strongly dependent upon pressure
Q4) Define molality.
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Chapter 13: Chemical Kinetics
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Q1) The rate of disappearance of HBr in the gas phase reaction 2HBr(g) H<sub>2</sub>(g)+ Br<sub>2</sub>(g)
Is 0.130 M s<sup>-1</sup> at 150°C.The rate of reaction is ________ M s<sup>-1</sup>.
A)3.85
B)0.0650
C)0.0169
D)0.260
E)0.0860
Q2) The first-order reaction,SO<sub>2</sub>Cl<sub>2</sub> SO<sub>2</sub> + Cl<sub>2</sub>,has a rate constant equal to 2.20 × 10<sup>-5</sup> s<sup>-1</sup> at 593 K.What percentage of the initial amount of SO<sub>2</sub>Cl<sub>2</sub> will remain after 6.00 hours?
A)1.00%
B)37.8%
C)40.2%
D)62.2%
Q3) Is the activation energy for a forward reaction the same as the activation energy for the reverse of the same reaction?
Q4) Define half-life.
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Chapter 14: Chemical Equilibrium
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Q1) Determine the value of K<sub>p</sub> for the following reaction if the equilibrium concentrations are as follows: P(CO)<sub>eq</sub> = 6.8 × 10<sup>-11</sup> atm,P(O<sub>2</sub>)<sub>eq</sub> = 1.3 × 10<sup>-3</sup> atm,P(CO<sub>2</sub>)<sub>eq</sub> = 0.041 atm.
2 CO(g)+ O<sub>2</sub>(g) 2 CO<sub>2</sub>(g)
A)3.6 × 10<sup>-21</sup>
B)2.8 × 10<sup>20</sup>
C)4.6 × 10<sup>11</sup>
D)2.2 × 10<sup>-12</sup>
E)3.6 × 10<sup>-15</sup>
Q2) The equilibrium constant is given for two of the reactions below.Determine the value of the missing equilibrium constant.
A(g)+ 2B(g) AB<sub>2</sub>(g)Kc = 59
AB<sub>2</sub>(g)+ B(g) AB<sub>3</sub>(g)K<sub>c</sub> = ?
A(g)+ 3B(g) AB<sub>3</sub>(g)K<sub>c</sub> = 478
A)3.5 × 10<sup>-5</sup>
B)2.8 × 10<sup>4</sup>
C)8.1
D)0.12
E)89
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Chapter 15: Acids and Bases
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Q1) The pOH of pure water at 40<sup>o</sup>C is 6.8.What is the hydronium ion concentration in pure water [H<sub>3</sub>O<sup>+</sup>] at this temperature?
A)6.3 × 10<sup>-8</sup>
B)1 × 10<sup>-14</sup>
C)1.0 × 10<sup>-7</sup>
D)1.6 × 10<sup>-7</sup>
Q2) Calculate the hydroxide ion concentration in an aqueous solution with a pH of 4.33 at 25°C.
A)2.1 × 10<sup>-10</sup> M
B)9.7 × 10<sup>-10 </sup>M
C)4.7 × 10<sup>-5</sup> M
D)3.8 × 10<sup>-5</sup> M
E)6.3 × 10<sup>-6</sup> M
Q3) Determine the pH of a 0.116 M Ba(OH)<sub>2</sub> solution at 25°C.
A)8.62
B)13.06
C)13.37
D)0.63
E)12.56
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Chapter 16: Aqueous Ionic Equilibrium
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Q1) 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 200.0 mL of KOH.The K<sub>a</sub> of HF is 3.5 × 10<sup>-4</sup>.
A)9.62
E)8.14
Q2) Which buffer controls the pH of blood?
A)CH<sub>3</sub>COOH and CH<sub>3</sub>COO<sup>-</sup>
B)H<sub>2</sub>CO<sub>3</sub> and HCO<sub>3</sub>C)NH<sub>4</sub><sup>+</sup> and NH<sub>3</sub>
D)H<sub>3</sub>PO<sub>4</sub> and H<sub>2</sub>PO<sub>4</sub><sup>-</sup>
Q3) Which one of the following statements is true?
A)A buffer is an aqueous solution composed of two weak acids.
B)A buffer can absorb an unlimited amount of acid or base.
C)A buffer resists pH change by neutralizing added acids and bases.
D)A buffer does not change pH when strong acid or base is added.
E)None of the above are true.
Q4) Explain the common ion effect with respect to molar solubility.
Page 18
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Chapter 17: Free Energy and Thermodynamics
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Q1) Why can endothermic reactions be spontaneous?
Q2) Calculate the G°<sub>rxn</sub> using the following information.
2 H<sub>2</sub>S(g)+ 3 O<sub>2</sub>(g) 2 SO<sub>2</sub>(g)+ 2
H<sub>2</sub>O(g) G°<sub>rxn </sub>= ?
H°<sub>f</sub> (kJ/mol)-20.6 296.8 -241.8
S°(J/mol<sup></sup>K)205.8 205.248.2 188.8
A)+196.8 kJ
B)+108.2 kJ
C)-466.1 kJ
D)+676.2 kJ
E)-147.1 kJ
Q3) Calculate G<sub>rxn</sub> at 298 K under the conditions shown below for the following reaction.
CaCO<sub>3</sub>(s) CaO(s)+ CO<sub>2</sub>(g) G° = +131.1 kJ
P(CO<sub>2</sub>)= 0.033 atm
A)-49.3 kJ
B)-8.32 kJ
C)+122.6 kJ
D)+39.7 kJ
E)+43.3 kJ
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Chapter 18: Electrochemistry
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Q1) A voltaic cell is constructed with two silver-silver chloride electrodes,where the half-reaction is
AgCl(s)+ e<sup>- </sup> Ag(s)+ Cl<sup>-</sup>(aq)E° = +0.222 V
The concentrations of chloride ion in the two compartments are 0.0222 M and 2.22 M,respectively.The cell emf is ________ V.
A)0.212
B)0.118
C)0.00222
D)22.2
E)0.232
Q2) How many kilowatt-hours of electricity are used to produce 3.00 kg of magnesium in the electrolysis of molten MgCl<sub>2</sub> with an applied emf of 4.50 V?
A)0.0336
B)0.0298
C)7.4
D)29.8
E)14.9
Q3) What is the difference between a voltaic cell and an electrolytic cell?
Q4) Why,if we multiply a reaction by 2,don't we multiply its E°<sub>red</sub> by 2?
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Chapter 19: Radioactivity and Nuclear Chemistry
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Q1) Identify the symptom that is not from radiation exposure.
A)increased white cell count
B)increased cancer risk
C)death
D)genetic effects
E)weaker immune systems
Q2) Which of the following statements are true?
A)If N/Z ratio is too high,there are too many protons and the nuclide will undergo positron emission or electron capture.
B)If N/Z ratio lies somewhere below 1,the nuclide is stable.
C)If N/Z ratio is too low,there are too many neutrons and the nuclide will undergo beta decay.
D)The valley of stability is the geographic location where many of the known nuclides were first discovered.
E)None of the above are true.
Q3) Why is an alpha emitter much more harmful if ingested?
Q4) Explain how radiation increases cancer risk.
Q5) Explain the concept of "magic numbers."
Q6) Describe what is meant by the "valley of stability."
Page 21
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