Chemistry for Health Sciences Midterm Exam - 3326 Verified Questions

Chemistry for Health Sciences

Midterm Exam

Course Introduction

Chemistry for Health Sciences offers an introduction to the fundamental concepts of chemistry as they apply to the health and medical fields. The course covers topics such as atomic and molecular structure, chemical bonding, solutions, acids and bases, chemical reactions, and the properties of gases and liquids. Emphasis is placed on the chemical processes that underpin physiological functions, drug interactions, and diagnostic procedures. Designed for students pursuing careers in nursing, allied health, or pre-medical studies, this course equips learners with the essential chemical knowledge needed to understand and interpret biological systems and health-related phenomena.

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) Identify the unit of measurement that is a SI base unit of measurement. A)second

B)Celsius

C)cup

D)pound

E)yard

Answer: A

Q2) Ethanol has a boiling point of 351 K. What is the boiling point in °C?

A)78 °C

B)-195 °C

C)122 °C

D)62 °C

E)86 °C

Answer: A

Q3) The factor 10^-3 corresponds to which prefix?

A)deka

B)deci

C)milli

D)centi

Answer: C

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Why do elements in the same group tend to have similar chemical properties?

Answer: Since elements in the same group have the same number of valence electrons (similar electron configurations)they tend to have similar chemical reactivity because chemical reactions typically involve valence electrons.

Q2) All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. Which law does this refer to?

A)Law of definite proportions

B)Law of the conservation of mass

C)Law of modern atomic theory

D)Law of multiple proportions

E)First law of thermodynamics

Answer: A

Q3) Which of the following describes a nonmetal?

A)Tend to gain electrons

B)Group 2 elements on the periodic table

C)Good conductor of electricity

D)Generally unreactive

E)Good conductor of heat

Answer: A

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) 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^-1.

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

Q2) Which of the following is an acid in aqueous solution?

A)NH₃

B)NaF

C)Ca(C₂H₃O₂)₂

D)HI

E)KCl

Answer: D

Q3) Give the name for HNO₂.

Answer: nitrous acid

Q4) Define empirical formula.

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

Page 5

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Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) What is the oxidation number of the chromium atom in K₂CrO₄?

A)-2

B)+2

C)+6

D)+7

Q2) How many grams of calcium chloride are needed to produce 1.50 g of potassium chloride? CaCl₂(aq)+ K₂CO₃(aq) 2KCl(aq)+ CaCO₃(aq)

A)0.896 g

B)1.12 g

C)2.23 g

D)4.47 g

Q3) How many grams of CH₃OH must be added to water to prepare 150 mL of a solution that is 2.0 mol L^-1 CH₃OH?

Q4) Define a limiting reagent.

Q5) How many grams of NaOH (MW = 40.0)are there in 250.0 mL of a 0.275 mol L^-1 NaOH solution?

Q6) Determine the oxidation state of Mn in KMnO₄.

Page 6

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Chapter 5: Gases

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Sample Questions

Q1) Which of the following samples will have the greatest average speed at 355 K?

A)Ne

B)C₂H₄

C)Cl₂

D)CH₄

E)All of these samples will have the same average speed at the same temperature.

Q2) What is the temperature of NO₂ gas if the average speed (actually the root mean square speed)of the molecules is 750 m s^-1?

A)1.38 K

B)1.04 × 10³ K

C)1.38 × 10³ K

D)1.04 × 10⁶ K

Q3) What is the definition of diffusion?

A)gas molecules mix equally

B)gas molecules spread out in a concentration gradient

C)gas molecules escape from a container into a vacuum through a small hole

D)average distance between collisions

E)gas molecules mix unequally

Q4) Why does the rate of effusion increase with a decrease in the molar mass?

Q5) Give the major gas in dry air.

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) How much energy is evolved during the formation of 98.7 g of Fe according to the reaction below?

Fe₂O₃(s)+ 2Al(s)

Al₂O₃(s)+ 2Fe(s) _rH°= -852 kJ

A)753 kJ

B)1.51 × 10³ kJ

C)4.20 × 10³ kJ

D)482 kJ

E)241 kJ

Q2) Choose the reaction that illustrates _fH° for NaHCO₃.

A)Na(s)+ H₂(g)+ C(s)+ O₂(g) NaHCO₃ (s)

B)Na⁺(aq)+ HCO₃ ^-1 (aq) NaHCO₃ (s)

C)Na⁺(aq)+ H2O (l)+ CO₂ (g) NaHCO₃(s)

D)Na(s)+ 1/2 H₂(g)+ C(s)+ 3/2 O₂(g) NaHCO₃ (s)

E)Na(s)+ 2 H(g)+ C(s)+ 3 O(g) NaHCO₃(s)

Q3) Identify the greatest source of energy worldwide.

A)petroleum

B)natural gas

C)wood

D)coal

E)nuclear power

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) How much energy (in kJ)do 3.0 moles of photons, all with a wavelength of 655 nm, contain?

A)183 kJ

B)303 kJ

C)394 kJ

D)548 kJ

E)254 kJ

Q2) n = 5 to n = 3

A)657 nm

B)7460 nm

C)103 nm

D)1280 nm

E)122 nm

Q3) Give all possible values of l for a 3 sublevel.

A)0

B)1, 2

C)0, 1, 2

D)-1, 0, 1

E)0, 1, 2, 3

Q4) Define constructive interference.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Use the periodic table to determine the ground-state electron configuration for the following element: Ga

A)[Ar]4s²3d¹⁰4p²

B)[Ar]4s²3d⁹

C)[Ar]4s²3d¹⁰

D)[Ar]4s²4p²

E)[Ar]4s²3d¹⁰4p¹

Q2) An element that has the valence electron configuration 5s²4d⁵ belongs to which period and group?

A)period 4; group 8

B)period 5; group 9

C)period 5; group 7

D)period 5; group 3

E)period 5; group 13

Q3) Of the following, which atom has the largest atomic radius?

A)Rb

B)I

C)Cs

D)At

Q4) Define electron affinity.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Identify the complete electronic configuration for Ca²⁺.

A)1s²2s²2p⁶3s²4p⁶

B)1s²2s²2p⁶3s²3p⁶

C)1s²2s²2p⁶3s²3p⁵

D)1s²2s²3p⁶4s²5p⁶

E)1s²2s²p⁶3s²p⁶

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

A)Al₃O₂

B)Al₂O₃

C)AlO₂

D)Al₂O

E)AlO

Q3) Place the following elements in order of increasing electronegativity. Sr N Na

A)Sr < Na < N

B)Na < N < Sr

C)Sr < N < Na

D)N < Sr < Na

E)N < Na < Sr

Q4) How are electron affinity and electronegativity different?

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

Q2) Determine the electron geometry (eg)and molecular geometry (mg)of CH₃⁺¹.

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

Q3) Determine the electron geometry (eg)and molecular geometry (mg)of the underlined carbon in CH₃CN.

A)eg = tetrahedral, mg = tetrahedral

B)eg = linear, mg = trigonal planar

C)eg = trigonal planar, mg = bent

D)eg = linear, mg = linear

E)eg = trigonal planar, mg = trigonal planar

Q4) Is it possible for a molecule to be nonpolar even though it contains polar bonds? Explain your answer and give an example.

Q5) 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) Lithium crystallizes in a body-centred cubic structure. What is the coordination number of each atom?

A)4

B)6

C)8

D)12

Q2) LiI

A)dispersion forces

B)dipole-dipole forces

C)ion-dipole forces

D)hydrogen bonding

E)ionic bond

F)H₂ + H₂O

Q3) Choose the pair of substances that are most likely to form a homogeneous solution.

A)CCl₄ and SCl₂

B)NF₃ and SO₂

C)CO and C₆H₆

D)NH₂CH₃ and CH₄

E)None of the pairs above will form a homogeneous solution.

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Page 13

Chapter 12: Solutions

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Sample Questions

Q1) Choose the aqueous solution that has the highest boiling point. These are all solutions of nonvolatile solutes and you should assume ideal van't Hoff factors where applicable.

A)0.100 mol kg^-1 AlCl₃

B)0.100 mol kg^-1NaCl

C)0.100 mol kg^-1MgCl₂

D)0.100 mol kg^-1C₆H₁₂O₆

E)They all have the same boiling point.

Q2) Commercial grade HCl solutions are typically 39.0% (by mass)HCl in water. Determine the molality of the HCl if the solution has a density of 1.20 g mL^-1.

A)39.0 mol kg^-1

B)17.5 mol kg^-1

C)6.39 mol kg^-1

D)10.7 mol kg^-1

E)9.44 mol kg^-1

Q3) Give the preparation of rock candy.

Q4) Explain why the van't Hoff factor for MgCl₂ is less than its predicted value.

Q5) Explain why water does not dissolve in gasoline.

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) A reaction with an activation energy of 71.6 kJ mol^-1 has a rate constant of 5.30 × 10 ^-6 L mol^-1 s^-1 at 298 K. At what temperature will the rate constant be 0.530 L mol^-1 s^-1?

A)461 K

B)410 K

C)310 K

D)345 K

E)328 K

Q2) k

A)activation energy

B)half-life

C)reaction order

D)rate constant

E)frequency factor

Q3) What is the difference between average reaction rate and instantaneous reaction rate?

Q4) Is the activation energy for a forward reaction the same as the activation energy for the reverse of the same reaction? Use a sketch of a reaction energy diagram.

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Page 15

Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Which of the following statements is TRUE?

A)If Q < K, it means the reverse reaction will proceed to form more reactants.

B)If Q > K, it means the reverse reaction will proceed to form more reactants.

C)If Q = K, it means the reaction is not at equilibrium.

D)Because the equilibrium is always dynamic, Q can never be equal to K.

E)If Q < K, the reaction proceeding to the left is much faster than the reaction proceeding to the right.

Q2) Q = K

A)reaction will favour formation of reactants

B)reaction has a larger amount of products than reactants

C)reaction favours formation of more products

D)reverse reaction is favoured

E)reaction does not strongly favour reactants or products

F)reaction is at equilibrium

Q3) Why aren't solids or liquids included in an equilibrium expression?

Q4) How is the reaction quotient different from an equilibrium constant for a given reaction?

Q5) How will equilibrium be affected for a reaction having the same number of moles of gaseous reactants and products if the pressure is increased?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Calculate the pOH of a solution that contains 3.9 × 10^-4 mol L^-1 H₃O at 25 °C.

A)4.59

B)3.31

C)10.59

D)9.14

E)0.59

Q2) Identify the monoprotic acid.

A)H₂CO₃

B)NH₃

C)H₃PO₄

D)H₂SO₄

E)HBr

Q3) Describe a molecule that can be a Lewis acid.

Q4) What is the autoionization of water?

Q5) Do both protons ionize instantaneously from a diprotic acid such as H₂CO₃? Explain your answer.

Q6) List the major source of energy in North America.

Q7) List the compound that is formed from the reaction of carbon dioxide with water.

Q8) Describe the relationship between molecular structure and acid strength.

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) pH = pK_a2

A)equivalence point of a weak base-strong acid titration

B)equivalence point of a weak acid-strong base titration

C)three-quarters of the way to second equivalence point of a diprotic acid-strong base titration

D)halfway to equivalence point of a weak acid-strong base titration

E)equivalence point of a strong acid-strong base titration

Q2) A 100.0 mL sample of 0.20 mol L^-1 HF is titrated with 0.10 mol L^-1 KOH. Determine the pH of the solution after the addition of 200.0 mL of KOH. The K_a of HF is 3.5 × 10^-4.

A)9.62

B)7.00

C)3.46

D)10.54

E)8.14

Q3) Sketch the titration curve for a strong acid titrated with a strong base. Make sure to indicate the equivalence point (and whether the equivalence point is acidic, basic, or neutral)and the buffer region.

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

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) What is the change in Gibbs energy of a process?

A)It is a maximum amount of energy available to do work.

B)It is energy that is lost to the surroundings.

C)It is energy that is used to break chemical bonds.

D)It is energy that is converted to heat.

E)It is the total energy of a process.

Q2) Define the third law of thermodynamics.

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

Q4) What is the name of the reaction that achieves the theoretical limits with respect to the change in Gibbs energy in thermodynamics?

A)reversible reaction

B)forward reaction

C)reverse reaction

D)equilibrium reaction

E)irreversible reaction

Q5) Define the second law of thermodynamics.

Q6) How is a nonspontaneous process made spontaneous?

Q7) Give the standard states for a gas, liquid, solid, and solution.

Page 19

Q8) What is "free" energy? Give a fictitious example.

Q9) Define entropy.

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Page 20

Chapter 18: Electrochemistry

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Sample Questions

Q1) Which of the following is the strongest oxidizing agent?

A)MnO₂(s)

B)Cl (aq)

C)Cu (aq)

D)SO₄^2-(aq)

E)MnO₄ (aq)

Q2) Nickel can be plated from aqueous solution according to the following half-reaction. How long would it take (in min)to plate 29.6 g of nickel at 4.7 A? Ni²⁺(aq)+ 2 e Ni(s)

A)1.7 × 10² min

B)5.9 × 10² min

C)3.5 × 10² min

D)4.8 × 10² min

E)6.2 × 10² min

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

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) What is the mass defect?

Q2) A freshly prepared sample of curium-243 undergoes 3312 disintegrations per second. After 6.00 y., the activity of the sample declines to 2755 disintegrations per second. The half-life of curium-243 is ________ y.

A)4.99

B)32.6

C)7.21

D)0.765

E)22.6

Q3) Nuclides below the valley of stability can become more stable through which of the following processes?

A)gamma emission

B)beta emission

C)positron emission

D)neutron emission

E)neutron bombardment

Q4) Define radioactivity.

Q5) Define chain reaction in terms of the fission of uranium nucleus.

Q6) List two problems associated with nuclear power.

Q7) Explain the concept of "magic numbers."

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Chapter 20: Organic Chemistry I: Structures

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Sample Questions

Q1) Which of the following alcohols has the lowest solubility in water?

A)CH₃CH₂OH

B)CH₃CH₂CH₂CH₂OH

C)CH₃CH₂CH₂CH₂CH₂O H

D)CH₃CH₂CH₂CH₂CH₂C H₂OH

E)CH₃CH₂CH₂CH₂CH₂CH ₂CH₂CH₂OH

Q2) Explain the formation of triple bonds in alkynes. Make sure to mention the hybridization of carbon atoms.

Q3) Which of the following compounds is ethanol?

A)C₂H₆

B)C₂H₅OH

C)CH₃CO₂H

D)CH₃CO₂CH₃

E)CH₃-O-CH₃

Q4) What is the difference between primary, secondary, and tertiary amines?

Q5) What is the index of hydrogen deficiency?

Q6) What are constitutional isomers?

Q7) What does the term "racemic mixture" mean?

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Chapter 21: Organic Chemistry II: Reactions

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Sample Questions

Q1) Which is the strongest organic acid?

A)F₃CCOOH

B)F₂HCCOOH

C)H₃COOH

D)ClH₂CCOOH

E)FH₂CCOOH

Q2) Identify the most oxidized compound.

A)CH₃CH₂COOH

B)CH₃CH₂CHO

C)CH₃CH₂CH₃

D)CH₃CH₂OCH₃

E)CH₃CH₂CH₂OH

Q3) Which of the following statements is TRUE?

A)Alkenes are as unreactive as alkanes.

B)Electrophilic addition reactions proceed via either an E1 or E2 mechanism.

C)The first step of an H-X electrophilic addition is the formation of a carbanion.

D)Carbon-carbon double bonds are electron rich and are nucleophilic.

E)During the addition of an H-X molecule, the H atom behaves as a nucleophile while the X atom behaves as an electrophile.

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Page 24

Chapter 22: Biochemistry

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Q1) Which one of the following amino acids contains a polar side chain?

A)alanine

B)glycine

C)threonine

D)valine

Q2) Which of the following is a type of nucleic acid?

A)RNA

B)amino acid

C)dipeptide

D)carbohydrate

E)lipid

Q3) Which of the following is an example of a secondary protein structure?

A)amino acid

B)fatty acid

C)lipid

D)(\(\beta\)-pleated sheet)

E)dipeptide

Q4) Why does the melting point of fatty acids decrease with the number of double bonds present?

Q5) What is a codon?

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Chapter 23: Chemistry of the Nonmetals

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Q1) Which of the following statements is TRUE?

A)Ammonium nitrate is insoluble in water.

B)Sodium nitrite is used in the glass industry.

C)NO₂ reacts with water to form ammonia.

D)Hydrazine is a good reducing agent.

E)Nitric acid is a weak acid.

Q2) Which of the following statements is TRUE?

A)All silicates are composed of SiO₄ tetrahedra that bond together through an oxygen to form single chains.

B)The fibrous nature of asbestos comes from its silicate network structure.

C)In some silicates, ionic bonding between a metal ion and an Si_xO_y^z^- ion is what holds the structure together.

D)The flakiness of mica comes from its single silicate chain structure.

E)Silicates are easily soluble in water because they contain silicate anions in their structure.

Q3) Why are phosphate compounds added to detergents?

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

Q5) Describe the major production method for obtaining oxygen.

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Chapter 24: Metals and Metallury

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Q1) What is the difference between ferromagnetism and paramagnetism?

Q2) Which of the following describes gangue?

A)the undesirable product formed during smelting

B)the undesirable portion of a metal-containing ore

C)the volatile product formed during calcination

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

E)a crude metal that still has to be further purified for industrial use

Q3) Why can't chromium and nickel form a miscible solid solution over the entire composition range?

A)The lever rule says that these two metals can't coexist in an alloy.

B)The two metals crystallize into different cubic structures in their pure forms.

C)The tetrahedral holes in the nickel crystal are not large enough to accommodate the chromium.

D)The octahedral holes in the chromium crystal are not large enough to accommodate the nickel.

E)Chromium atoms can only fit inside tetrahedral holes, which are not present in the structure of pure nickel.

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

Q5) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Which of the following complex ions absorbs light of the shortest wavelength?

A)[FeCl₆]^3-

B)[Fe(OH)₆]^3-

C)[Fe(CN)₆]^3-

D)[Fe(H₂O)₆]³⁺

E)[FeI₆]^3-

Q2) Choose the electron configuration for Zr²⁺.

A)[Kr]5s²4d²

B)[Kr]5s²

C)[Kr]4d²

D)[Ar]4s²3d¹⁰4p⁶

E)[Kr]5s²4d⁴

Q3) Which of the following complex ions absorbs light of the longest wavelength?

A)[Cr(CN)₆]^3-

B)[CrCl₆]^3-

C)[Cr(en)₃]³⁺

D)[Cr(NH₃)₆]³⁺

E)[Cr(NO₂)₆]^3-

Q4) What is the difference between a weak-field complex and a strong-field complex?

Q5) Why is +2 a common oxidation state for transition elements?

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