Science Foundation Course Midterm Exam - 3326 Verified Questions

Science Foundation Course

Midterm Exam

Course Introduction

The Science Foundation Course is designed to provide students with a comprehensive introduction to the core concepts and principles across the major scientific disciplines, including biology, chemistry, physics, and earth sciences. Through a blend of theoretical lessons and practical laboratory work, students will develop essential scientific skills such as observation, experimentation, and critical analysis. The course emphasizes the scientific method, fostering an understanding of how scientific knowledge is generated, validated, and applied to solve real-world problems. By the end of the course, students will have established a strong foundation to support further studies in science and related fields.

Recommended Textbook

Chemistry A Molecular Approach 1st Canadian Ediiton 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) Which of the following is the smallest volume?

A)44 cm³

B)1.0 dL

C)5.5 × 10³ mL

D)1.0 × 10⁸ nL

Answer: A

Q2) Which of the following is the lowest temperature?

A)42 °C

B)57 °F

C)318 K

D)All of these temperatures are equal.

Answer: B

Q3) The mass of a single zinc atom is 1.086 × 10^-22 g. This is the same mass as

A)1.086 × 10^-16 mg.

B)1.086 × 10^-25 kg.

C)1.086 × 10^-28 g.

D)1.086 × 10^-³¹ ng.

Answer: B

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following is a transition element?

A)Pd

B)Sn

C)K

D)U

E)Pr

Answer: A

Q2) How many atoms are in 2.50 moles of CO₂?

A)4.52 × 10²⁴ atoms

B)1.52 × 10²⁴ atoms

C)5.02 × 10²³ atoms

D)3.01 × 10²⁴ atoms

E)7.53 × 10²³ atoms

Answer: A

Q3) How many protons (p)and neutrons (n)are in an atom of barium-130?

A)56 p, 74 n

B)56 p, 130 n

C)74 p, 56 n

D)130 p, 56 n

Answer: A

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) Determine the name for aqueous HBr.

A)bromic acid

B)bromous acid

C)hydrobromous acid

D)hydrogen bromate

E)hydrobromic acid

Answer: E

Q2) How many molecules of N₂O₄ are in 76.3 g N₂O₄? The molar mass of N₂O₄ is 92.02 g mol^-1.

A)5.54 × 10²⁵ N₂O₄ molecules

B)7.26 × 10²³ N₂O₄ molecules

C)1.38 × 10²⁴ N₂O₄ molecules

D)4.59 × 10²⁵N₂O₄ molecules

E)4.99 × 10²³ N₂O₄ molecules

Answer: E

Q3) Calculate the mass percent composition of oxygen in Al₂(SO₄)₃.

Answer: 56.11%

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Chapter 4: Chemical Reactions and Stoichiometry

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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 3.00 g of K₃PO₄ in enough water to make 350. mL of solution?

A)0.0101 mol L^-1

B)0.0404 mol L^-1

C)0.162 mol L^-1

D)0.323 mol L^-1

Q2) Metallic aluminum reacts with MnO₂ at elevated temperatures to form manganese metal and aluminum oxide. A mixture of these two reactants is 67.2% mole percent Al. Find the theoretical yield (in grams)of manganese from the reaction of 250 g of this mixture.

A)87.3 g

B)193 g

C)18.13 g

D)6.72 g

E)96.6 g

Q3) How many moles of Co²⁺ are present in 0.150 L of a 0.200 mol L^-1 solution of CoI₂?

Q4) Determine the oxidation state of Mn in KMnO₄.

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Chapter 5: Gases

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Sample Questions

Q1) What is the pressure in a gas container that is connected to an open-end U-tube manometer if the pressure of the atmosphere is 742 Torr and the level of mercury in the arm connected to the container is 8.60 cm higher than the level of mercury open to the atmosphere?

A)656 mmHg

B)733 mmHg

C)751 mmHg

D)828 mmHg

Q2) How many litres of oxygen are needed to exactly react with 19.8 g of methane at STP? CH₄(g)+ 2O₂(g) CO₂(g)+ 2H₂O(l)

A)13.9 L

B)27.8 L

C)56.0 L

D)60.5 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 5.00 g sample of liquid water at 25.0 °C is heated by the addition of 145 J of energy. The final temperature of the water is ________ °C. The specific heat capacity of liquid water is 4.184 J g^-1 K^-1.

A)146

B)25.1

Q2) An unknown metal alloy, specific heat capacity = 0.372 J g^-1 °C^-1, has a temperature change of +21.7 °C after a heat transfer of 216.4 J. Calculate the mass of the alloy in question.

A)14.8 g

B)19.0 g

C)36.2 g

D)21.5 g

E)26.8 g

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

Q4) Define chemical energy.

Q5) Where does the energy absorbed during an endothermic reaction go?

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) It is possible to determine the ionization energy for hydrogen using the Bohr equation. Calculate the ionization energy for an atom of hydrogen, making the assumption that ionization is the transition from n = 1 to n = .

A)-2.18 × 10^-18 J

B)+2 .18 × 10^-18 J

C)+4.59 × 10^-18 J

D)-4.59 × 10^-18 J

E)+4.36 × 10^-18 J

Q2) Each of the following sets of quantum numbers is supposed to specify an orbital.

Choose the one set of quantum numbers that does not contain an error.

A)n = 4, l = 4, m_l = 0

B)n = 3, l = 2, m_l = +3

C)n = 4, l = 0, m_l = -1

D)n = 3, l = 1, m_l = -2

E)n = 5, l = 3, m_l = -3

Q3) Define diffraction.

Q4) 2

A)number of unpaired electrons in Zn²⁺

B)number of unpaired electrons in Ti²⁺

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Which of the following have the same number of valence electrons?

A)Rb, Sb, I

B)Ga, Sn, Bi

C)As, Sb, Bi

D)Ar, Kr, Br

Q2) How many valence electrons do the noble gases possess?

A)1

B)2

C)7

D)6

E)8

Q3) Define ionization energy.

Q4) Which reaction below represents the first ionization of O?

A)O (g)+ e O(g)

B)O(g)+ e O (g)

C)O (g) O(g)+ e

D)O(g) O (g)+ e

E)O (g)+ e O² (g)

Q5) Why do Li, Na, and K have similar chemical properties?

Q6) Define electron affinity.

10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Which of the following processes is endothermic?

A)the reaction associated with the lattice energy of LiCl

B)the reaction associated with the ionization energy of potassium

C)the reaction associated with the heat of formation of CaS

D)the formation of F₂ from its elements in their standard states

E)None of the above processes is endothermic.

Q2) Use Lewis theory to determine the chemical formula for the compound formed between Ca and N.

A)CaN

B)Ca₃N₂

C)CaN₂

D)Ca₂N

E)Ca₂N₃

Q3) Draw the best Lewis structure for BrO₄ . What is the formal charge on bromine?

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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) Identify the number of electron groups around a molecule with a tetrahedral shape.

A)1

B)2

C)3

D)4

E)5

Q2) Determine the electron geometry for the molecule XeF₄.

A)Linear

B)Trigonal planar

C)Tetrahedral

D)Trigonal bipyramidal

E)Octahedral

Q3) Determine the electron geometry (eg), molecular geometry (mg), and polarity of XeO₃.

A)eg = trigonal planar, mg = trigonal planar, nonpolar

B)eg = tetrahedral, mg = trigonal pyramidal, polar

C)eg = trigonal planar, mg = trigonal pyramidal, polar

D)eg = trigonal bipyramidal, mg = trigonal planar, nonpolar

E)eg = octahedral, mg = tetrahedral, nonpolar

Page 12

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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Define deposition.

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.

Q2) Choose the substance with the highest vapour pressure at a given temperature.

A)SiS₂

B)RbCl

C)CH₃SCH₃

D)BF₃

E)SbH₃

Q3) Identify the compound that does not have dipole-dipole forces as its strongest force.

A)CH₂Cl₂

B)CH₃OCH₃

C)CH₃Br

D)HCCl₃

E)CO₂

Q4) Describe the difference between the conduction band and the valence band.

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Chapter 12: Solutions

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Sample Questions

Q1) A 125.0 mL sample of an aqueous solution at 25 °C contains 29.4 mg of an unknown nonelectrolyte compound. If the solution has an osmotic pressure of 24.4 mbar, what is the molar mass of the unknown compound?

A)116 g mol^-1

B)281 g mol^-1

C)239 g mol^-1

D)98.6 g mol^-1

E)311 g mol^-1

Q2) An aqueous solution is 0.387 mol L^-1 in HCl. What is the molality of the solution if the density is 1.23 g mL^-1?

A)0.115 mol kg^-1

B)0.387 mol kg^-1

C)0.315 mol kg^-1

D)0.411 mol kg^-1

E)0.318 mol kg^-1

Q3) Explain why the van't Hoff factor for MgCl₂ is less than its predicted value.

Q4) Explain why water does not dissolve in gasoline.

Q5) Define the Tyndall effect.

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) What is the activation energy for a reaction that has a rate constant of 0.0883 L mol^-1 s^-1at 333 K and 0.469 L mol^-1 s^-1 at 366 K?

A)34.8 kJ mol^-1

B)28.0 kJ mol^-1

C)53.1 kJ mol^-1

D)39.1 kJ mol^-1

E)60.9 kJ mol^-1

Q2) The decomposition of dinitrogen pentoxide is described by the chemical equation 2N₂O₅(g) 4NO₂(g)+ O₂(g).

If the rate of disappearance of N₂O₅ is equal to 1.60 mol min^-1 at a particular moment, what is the rate of appearance of NO₂ at that moment?

A)0.800 mol min^-1

B)1.60 mol min^-1

C)3.20 mol min^-1

D)6.40 mol min^-1

Q3) Explain what the exponential factor in the Arrhenius equation represents.

Q4) What function do enzymes serve?

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Determine the value of K_p for the following reaction if the equilibrium concentrations are as follows: P(NOCl)_eq = 0.22 atm, P(NO)_eq = 0.10 atm, P(Cl₂)_eq = 0.081 atm. 2NOCl(g) 2NO(g)+

Cl₂(g)

A)3.7 × 10^-2atm

B)60 atm

C)27 atm

D)1.7 × 10^-2atm

E)1.8 × 10^-3atm

Q2) The equilibrium constant is given for one of two reactions below. Determine the value of the missing equilibrium constant. 2SO₂(g)+ O₂(g) 2SO₃(g)K = 1.7 × 10⁶

SO₃(g) 1/2 O₂(g)+ SO₂(g)K = ?

A)3.4 × 10²

B)8.5

C)1.3 × 10³

D)1.2 × 10^-6

E)7.7 × 10^-4

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

Q4) Why aren't solids or liquids included in an equilibrium expression?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Order the following acids in order of increasing strength: HBrO HBrO3 ^HBrO4HBrO₂

A)HBrO < HBrO₃ < ^HBrO4< HBrO₂

B)HBrO₄ < HBrO₃ < ^HBrO2< HBrO

C)HBrO < HBrO₃ < ^HBrO2< HBrO₄

D)HBrO < HBrO₂ < ^HBrO3< HBrO₄

E)HBrO₂ < HBrO₄ < ^HBrO< HBrO₂

Q2) Which of the following correctly describes the characteristics of a strong acid?

A)ionizes only partially in aqueous solutions

B)has a very electron positive atom attached to the oxygen

C)has a nonpolar bond

D)is completely ionized in aqueous solution

E)is always monoprotic

Q3) Which of the following is a Lewis acid?

A)BCl₃

B)CH₄

C)NH₃

D)CHCl₃

E)CCl₄

Q4) List the compound that is formed from the reaction of carbon dioxide with water.

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) A 100.0 mL sample of 0.10 mol L^-1 NH₃ is titrated with 0.10 mol L^-1 HNO₃. Determine the pH of the solution after the addition of 200.0 mL of HNO₃. The K_b of NH₃ is 1.8 × 10^-5.

A)6.44

B)1.48

C)2.00

D)12.52

E)12.00

Q2) Calculate the pH of a buffer that is 0.225 mol L^-1 CH₃COOH and 0.162 mol L^-1 CH₃COOK. The K_a for CH₃COOH is 1.8 × 10^-5.

A)4.89

B)9.11

C)4.74

D)9.26

E)4.60

Q3) Identify the compound that forms stalactites and stalagmites.

Q4) Give the name of the compound that is in antifreeze and is toxic to pets.

Q5) Define a buffer.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Determine _rG° at 298 K using the following information: H₂(g)+ CO(g) CH₂O(g) _rH°= +1.9 kJ mol^-1; _rS°= -109.6 J K^-1mol^-1

A)+57.7 kJ mol^-1

B)-30.8 kJ mol^-1

C)+34.6 kJ mol^-1

D)-41.5 kJ mol^-1

E)+17.3 kJ mol^-1

Q2) Determine _rG° at 298 K using the following information: FeO(s)+ CO(g) Fe(s)+ CO₂(g) _rH°= -11.0 kJ mol^-1; _rS°= -17.4 J K^-1 mol^-1

A)+191.0 kJ mol^-1

B)-5.8 kJ mol^-1

C)+1.6 kJ mol^-1

D)-6.4 kJ mol^-1

E)+89.5 kJ mol^-1

Q3) Why can't we say that a spontaneous reaction is a fast reaction?

Q4) Why is the quantity of energy required to recharge a battery greater than the quantity of work done?

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Why are iron nails coated with zinc?

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

A)Sn²⁺(aq)

B)Cr³⁺(aq)

C)Sn⁴⁺(aq)

D)Cr(s)

E)Sn(s)

Q3) Identify the battery that is used as a common flashlight battery.

A)dry-cell battery

B)lithium-ion battery

C)lead-acid storage battery

D)NiCad battery

E)fuel cell

Q4) Which of the following reactions would have the smallest value of K at 298 K?

A)A + B C; E°_cell = +1.22 V

B)A + 2 B C; E°_cell = +0.98 V

C)A + B 2 C; E°_cell = -0.030 V

D)A + B 3 C; E°_cell = +0.15 V

E)A + B C; E°_cell = -0.015 V

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

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Explain the concept of "magic numbers."

Q2) Determine the binding energy per nucleon of a Mg-24 nucleus. The Mg-24 nucleus has a mass of 24.30506. A proton has a mass of 1.00728 amu, a neutron has a mass of 1.008665 amu, and 1 amu is equivalent to 931 MeV of energy.

A)0.3050 MeV

B)8.83 MeV

C)0.113 MeV

D)106 MeV

E)4.41 MeV

Q3) Describe what changes occur in the atomic nucleus during positron emission.

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 is unchanged and the atomic number increases.

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

Q4) What is the mass defect?

Q5) How does a dosimeter measure exposure to radioactivity?

Q6) Describe what is meant by the "valley of stability."

Q7) Define radioactivity.

Q8) Define chain reaction in terms of the fission of uranium nucleus.

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Chapter 20: Organic Chemistry I: Structures

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Sample Questions

Q1) Which of the following condensed general formulas represents ethers?

A)ROR

B)RCOR

C)RCHO

D)ROH

E)RCOOH

Q2) Identify the generic formula for alkanes.

A)C_nH_2n+2

B)C_nH_2n_-₂

C)C_nH_2n

D)C_nH_2n_-4

E)C_nH₂_n₊₄

Q3) How many isomers are there for C₅H₁₂ ? A)3 B)4 C)5 D)6

Q4) What does the term "racemic mixture" mean?

Q5) What is the index of hydrogen deficiency?

Q6) What is the difference between primary, secondary, and tertiary amines?

Page 22

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Chapter 21: Organic Chemistry II: Reactions

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Sample Questions

Q1) Identify the most oxidized compound.

A)CH₃CH₂COOH

B)CH₃CH₂CHO

C)CH₃CH₂CH₃

D)CH₃CH₂OCH₃

E)CH₃CH₂CH₂OH

Q2) Which of the following statements is TRUE?

A)In nucleophilic substitution reactions at a saturated carbon atom, a carbon with a positive charge behaves as a nucleophile.

B)Nucleophilic substitution reactions at a saturated carbon atom can follow one of three mechanisms: S_N1, S_N2, and S_N3.

C)A carbon atom with a partial positive charge can react with a nucleophile, a molecule, or ion that has a lone pair of electrons.

D)Alkanes are an example of saturated organic compounds that easily participate in nucleophilic substitution reactions at a saturated carbon atom.

E)In nucleophilic substitution reactions at a saturated carbon atom, the replaced group is called an electrophile.

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Chapter 22: Biochemistry

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Q1) Which of the following shows all of the tripeptides that can be formed from one molecule each of glycine (Gly), valine (Val), and leucine (Leu)?

A)GlyValLeu, GlyLeuVal, ValLeuGly, ValGlyLeu, LeuGlyVal, LeuValGly

B)GlyValLeu, GlyLeuVal, LeuGlyVal

C)ValGlyLeu, GlyValLeu, GlyLeuVal, LeuGlyVal

D)ValGlyLeu and GlyLeuVal

E)GlyValLeu

Q2) How do cis- and trans-fats differ?

Q3) Which of the following statements is TRUE?

A)Monosaccharides are the building blocks of polypeptides.

B)Nucleotides are the building blocks of saccharides.

C)Fatty acids are responsible for DNA replication.

D)Peptides are the building blocks of amino acids.

E)Carbohydrates are water soluble and responsible for short-term energy storage.

Q4) Which of the following link together into amino acid units?

A)glycosidic linkages

B)disulfide linkages

C)peptide bonds

D)hydrogen bonds

E)ester linkages

Page 24

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Chapter 23: Chemistry of the Nonmetals

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Q1) pyrosilicate

A)Si₂O₅^2-

B)SiO₄^4-

C)Si₄O₁₁^6-

D)SiO₃^2-

E)Si₂O₇^6-

Q2) Write the balanced equation that represents the formation of hydrofluoric acid from the reaction of calcium fluoride with sulfuric acid.

A)CaF₂ + H₂S CaS + 2HF

B)CF₄ + 2H₂S 4HF + CS₂

C)CF₄ + 2H₂SO₄ 4HF + CS₂ + CO₂

D)CaF₂ + H₂SO₄ 2HF + CaSO₄

E)CaF₂ + H₂SO₃ CaSO₃ + 2HF

Q3) Identify the formula of the compound that is in common glass.

A)SiO₂

B)Si₂O₃

C)SiO₃

D)Si₂O₂

E)SiO₄

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Chapter 24: Metals and Metallury

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Q1) Pb

A)sphalerite

B)rhodochrosite

C)rutile

D)carnotite

E)malachite

F)galena

G)cinnabar

Q2) Identify the normal crystal structure at room temperature for Ni.

A)face-centred cubic

B)body-centred cubic

C)tetrahedral

D)hexagonal

E)hexagonal closest packed

Q3) Which substance is the best conductor of electricity?

A)arsenic

B)boron

C)silver

D)tellurium

Q4) Why does a "two-phase" structure occur in a substitutional alloy?

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Which of the following can function as a bidentate ligand?

A)CN^-

B)SH^-

C)HO

D)H₂NCH₂

Q2) What is a coordinate covalent bond?

Q3) linkage

A)[Zn(NH₃)₂Cl₂]Br_{2 and}[Zn(NH₃)₂Br₂]Cl₂

B)[Cr(en)₃]³⁺

C)[Ni(CN)₂(OH)₂]^2-

D)[Cu(SCN)₄]^3-and[Cu(NCS)₄]^3-

E)[Cr(NH₃)₃(CO)₃]³⁺

Q4) What is the ground-state electron configuration for Cr³⁺ (Z = 24)?

A)[Ar] 4s²3d⁷

B)[Ar] 4s²3d¹

C)[Ar] 3d³

D)[Ar] 4s¹3d²

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

Q5) What is the difference between a weak-field complex and a strong-field complex?