Pre-Medical Chemistry Exam Preparation Guide - 2309 Verified Questions

Pre-Medical Chemistry Exam Preparation Guide

Course Introduction

Pre-Medical Chemistry is an essential foundational course designed for students planning to pursue careers in medicine or health-related professions. The course covers fundamental principles of general and organic chemistry, including atomic structure, chemical bonding, stoichiometry, chemical reactions, thermodynamics, acid-base chemistry, and the structure and function of biologically relevant molecules. Emphasis is placed on developing problem-solving skills and understanding the chemical processes vital to biological systems. Through lectures, laboratory experiments, and interactive discussions, students gain the knowledge and practical skills necessary to progress to advanced courses in chemistry, biology, and medical sciences.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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24 Chapters

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Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) In an average year, the American chemical industry produces more than 9.5 million metric tons of sodium carbonate. Over half of this is used in the manufacture of glass while another third is used in the production of detergents and other chemicals. How many pounds of sodium carbonate are produced annually?

A) 2.1 × 10¹⁰ lb

B) 4.3 × 10⁹ lb

C) 1.1 × 10⁷ lb

D) 2.2 × 10⁶ lb

E) 2.1 × 10⁴lb

Answer: D

Q2) Which of the following is not an SI base unit?

A) meter

B) ampere

C) second

D) gram

E) kelvin

Answer: D

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3

Chapter 2: The Components of Matter

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Q1) Sodium chromate is used to protect iron from corrosion and rusting. Determine its molecular mass.

A) 261.97 amu

B) 238.98 amu

C) 161.97 amu

D) 138.98 amu

E) 74.99 amu

Answer: C

Q2) Rutherford bombarded gold foil with alpha (\(\alpha\)) particles and found that a small percentage of the particles were deflected. Which of the following was not accounted for by the model he proposed for the structure of atoms?

A) the small size of the nucleus

B) the charge on the nucleus

C) the total mass of the atom

D) the existence of protons

E) the presence of electrons outside the nucleus

Answer: C

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4

Chapter 3: Stoichiometry of Formulas and Equations

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Q1) Gaseous methanol (CH₄O) reacts with oxygen gas to produce carbon dioxide gas and liquid water. Write a balanced equation for this process.

Answer: 2CH₄O(g) + 3O₂(g) \(\to\) 2CO₂(g) + 4H₂O(l)

Q2) Which of the following samples has the most moles of the compound?

A) 50.0 g of Li₂O

B) 75.0 g of CaO

C) 200.0 g of Fe₂O₃

D) 50.0 g of CO₂

E) 100.0 g of SO₃

Answer: A

Q3) Ammonia, NH₃, is produced industrially from nitrogen and hydrogen as follows:

N₂(g) + 3H₂(g) \(\to\) 2NH₃(g)

What mass, of which starting material, will remain when 30.0 g of N₂ and 10.0 g of H₂react until the limiting reagent is completely consumed?

Answer: 3.52 g of hydrogen gas remains.

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Chapter 4: The Major Classes of Chemical Reactions

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Q1) Aqueous potassium iodate (KIO₃) and potassium iodide (KI) react in the presence of dilute hydrochloric acid, as shown below.

KIO₃(aq) + 5KI(aq) + 6HCl(aq) \(\to\) 3I₂(aq) + 6KCl(aq) + 3H₂O(l)

What mass of iodine (I₂) is formed when 50.0 mL of 0.020 M KIO₃ solution reacts with an excess of KI and HCl?

A) 0.13 g I₂

B) 0.25 g I₂

C) 0.38 g I₂

D) 0.76 g I₂

E) none of the above

Q2) Which one of the following substances is a strong acid?

A) HNO₃

B) H₂CO₃

C) NH₃

D) CH₃COOH

E) H₃PO₄

Q3) Complete and balance the equation for the following acid-base reaction. Ca(OH)₂(aq) + HCl(aq) \(\to\)

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Page 6

Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) Which of the following gases will be the slowest to diffuse through a room?

A) methane, CH₄

B) hydrogen sulfide, H₂S

C) carbon dioxide, CO₂

D) water, H₂O

E) neon, Ne

Q2) Lithium oxide is an effective absorber of carbon dioxide and can be used to purify air in confined areas such as space vehicles. What volume of carbon dioxide can be absorbed by 1.00 kg of lithium oxide at 25°C and 1.00 atm? Li₂O(aq) + CO₂(g) \(\to\) Li₂CO₃(s)

A) 687 mL

B) 819 mL

C) 687 L

D) 819 L

E) 22.4 L

Q3) According to the postulates of kinetic-molecular theory, the molecules of all gases at a given temperature have the same average kinetic energy.

A)True

B)False

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Page 7

Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) a. Explain fully what is meant by the term "state function." b. (i) Give two examples of thermodynamic quantities which state functions. (ii) Give two examples of thermodynamic quantities which are not state functions.

Q2) Calculate the \(\Delta\)H°_rxn for the decomposition of calcium carbonate to calcium oxide and carbon dioxide. \(\Delta\)H°_f [CaCO₃(s)] = -1206.9 kJ/mol; \(\Delta\)H°_f [CaO(s)] = -635.1 kJ/mol; \(\Delta\)H°_f [CO₂(g)] = -393.5 kJ/mol

CaCO₃(s) \(\to\) CaO(s) + CO₂(g)

A) -2235.5 kJ

B) -1448.5 kJ

C) -178.3 kJ

D) 178.3 kJ

E) 2235.5 kJ

Q3) Calculate q when 28.6 g of water is heated from 22.0°C to 78.3°C.

A) 0.385 kJ

B) 1.61 kJ

C) 6.74 kJ

D) 9.37 kJ

E) 1.61 × 10³ kJ

Page 8

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) In the Bohr model of the hydrogen atom, the electron moves in a circular path which Bohr referred to as an orbital.

A)True

B)False

Q2) Select the arrangement of electromagnetic radiation which starts with the lowest energy and increases to greatest energy.

A) radio, infrared, ultraviolet, gamma rays

B) radio, ultraviolet, infrared, gamma rays

C) gamma rays, infrared, radio, ultraviolet

D) gamma rays, ultraviolet, infrared, radio

E) infrared, ultraviolet, radio, gamma rays

Q3) Which scientist demonstrated that photons transferred momentum during collisions with matter?

A) Bohr

B) de Broglie

C) Planck

D) Compton

E) Billiard

Q4) The following combinations of quantum numbers are not allowed.

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Select the most acidic compound from the following:

A) SO₂

B) Al₂O₃

C) CaO

D) PbO

E) H₂O

Q2) Which of the following elements has the largest second ionization energy (IE₂)?

A) Li

B) B

C) O

D) F

E) Na

Q3) Select the element with the least metallic character.

A) Sn

B) Sr

C) Tl

D) Ge

E) Ga

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Page 10

Chapter 9: Models of Chemical Bonding

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Q1) Electronegativity is a measure of

A) the energy needed to remove an electron from an atom.

B) the energy released when an electron is added to an atom.

C) the magnitude of the negative charge on an electron.

D) the attraction by an atom for electrons in a chemical bond.

E) the magnitude of the negative charge on a molecule.

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) The electrostatic energy of two charged particles is inversely proportional to the square of the distance between them.

A)True

B)False

Q4) Ionic bonding typically occurs when a ______________ bonds with a

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Chapter 10: The Shapes of Molecules

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Q1) Explain what is meant by "dipole moment", and give an example of a molecule which has polar bonds but which does not itself have a dipole moment.

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) Predict the actual bond angles in SF₃⁺ using the VSEPR theory.

A) more than 120°

B) exactly 120°

C) between 109° and 120°

D) between 90° and 109°

E) less than 90°

Q4) Using SO₂ as an example, describe the sort of experimental data which might suggest that no single Lewis structure is an accurate representation of its bonding.

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Chapter 11: Theories of Covalent Bonding

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Q1) According to molecular orbital theory, all diatomic molecules with an even number of electrons will be diamagnetic.

A)True

B)False

Q2) A molecule with the formula AX₃E uses _________ to form its bonds.

A) s and p atomic orbitals

B) sp³ hybrid orbitals

C) sp² hybrid orbitals

D) sp hybrid orbitals

E) sp³d² hybrid orbitals

Q3) According to molecular orbital theory, what is the bond order in the O₂⁺ ion?

A) 5.5

B) 5

C) 4

D) 2.5

E) 1.5

Q4) In one sentence state the basic principle of valence bond theory.

Q5) In not more than two sentences, explain when and why chemists make use of the concept of hybridization.

Page 13

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

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Sample Questions

Q1) Smectic liquid crystals are more highly ordered than either nematic or cholesteric liquid crystals.

A)True

B)False

Q2) Liquid ammonia boils at -33.4°C and has a heat of vaporization of 23.5 kJ/mol. Calculate its vapor pressure at -50.0°C.

Q3) The highest temperature at which superconductivity has been achieved is approximately

A) 4 K.

B) 30 K.

C) 70 K.

D) 100 K.

E) 130 K.

Q4) A metal with a body-centered cubic lattice will have ______ atom(s) per unit cell.

A) 1

B) 2

C) 3

D) 4

E) 9

14

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Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) Which of the following pairs of ions is arranged so that the ion with the smaller charge density is listed first?

A) K⁺, Rb⁺

B) Cl^-, K⁺

C) Cl^-, Br^-

D) Ca²⁺, Ba²⁺

E) None of the above pairs are arranged in this way.

Q2) A saturated solution of carbon dioxide in water contains 3.00 g of CO₂ when the CO₂ partial pressure is 8.0 atm. What mass of CO₂ will escape if the partial pressure is lowered to 3.2 atm?

A) 0.90 g

B) 1.20 g

C) 1.40 g

D) 1.80 g

E) 2.20 g

Q3) The density of pure water at 25°C is 0.997 g/mL. Considering water as being both solvent and solute, calculate its molarity.

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) The Haber process is the first step in the manufacture of sulfuric acid. A)True

B)False

Q2) Predict the products for the following set of reactants. Li(s) + CH₃Cl(g) \(\to\)

A) CH₃CH₃(g) + LiCl(s)

B) CH₄(g) + LiCl(s)

C) CH₃Li(s) + LiCl(s)

D) CH₃Li(s) + Cl₂(g)

E) CH₂Li(s) + HCl(g)

Q3) Predict the products for the reaction of the following set of reactants. MgCO₃(s) + heat \(\to\)

A) Mg(s) + CO(g)

B) Mg(s) + CO₂(g)

C) MgO(s) + CO(g)

D) MgO(s) + CO₂(g)

E) no reaction

Q4) Phosphoric acid (H₃PO₄) is a strong acid.

A)True B)False

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) Explain how a salt bridge can arise a protein molecule.

Q2) In one sentence, what is the general requirement for a molecule to be optically active?

Q3) Given that each 3-base sequence in DNA is a "code word" for a particular amino acid, how many different code words are possible using a 3-base sequence and the bases available in DNA?

Q4) Condensation polymers are formed by free radical addition reactions.

A)True

B)False

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

B)False

Q6) Explain how a disulfide bridge can arise a protein molecule.

Q7) All alcohols are capable of hydrogen bonding. A)True

B)False

Page 17

Q8) A characteristic reaction of alkenes is addition. A)True

B)False

Q9) Draw and name two non-alkenes with the formula C₄H₈.

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Q1) The reaction CH₃NC(g) \(\to\) CH₃CN(g) is first-order with respect to methyl isocyanide, CH₃NC. If it takes 10.3 minutes for exactly one quarter of the initial amount of methyl isocyanide to react, what is the rate constant in units of min^-1?

A) -0.135 min^-1

B) 0.0279 min^-1

C) 0.089 min^-1

D) 0.135 min^-1

E) 35.8 min^-1

Q2) A reactant R is being consumed in a first-order reaction. What fraction of the initial R is consumed in 4.0 half-lives?

A) 0.94

B) 0.87

C) 0.75

D) 0.13

E) 0.063

Q3) All bimolecular reactions are second-order reactions.

A)True

B)False

Page 19

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) For a gas-phase equilibrium, a change in the pressure of any single reactant or product will affect the amounts of other substances involved in the equilibrium.

A)True

B)False

Q2) SO₂ reacts with O₂ to produce SO₃. If 86.0 g of SO₂ is placed in a reaction vessel along with excess oxygen gas, how many moles of SO₂ remain when 50.0 g of SO₃ have been formed?

A) 0.56 mol

B) 0.62 mol

C) 0.72 mol

D) 0.78 mol

E) 1.34 mol

Q3) In a chemical reaction, if the starting concentrations of reactants are increased, then the equilibrium constant K_c will also increase.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) Describe what is meant by the "leveling effect". Use a real acid as an example, and write an appropriate equation.

Q2) Which of the following aqueous liquids will have the highest pH?

A) 0.1 M CH₃COOH, pK_a = 4.7

B) 0.1 M CuCl₂, pK_a = 7.5

C) 0.1 M H₃C₆H₅O₇, pK_a = 3.1

D) 0.1 M ZnCl₂, pK_a = 9.0

E) pure water

Q3) What is the pH of a 0.010 M triethanolammonium chloride, (HOC₂H₂)₃NHCl, solution? K_b, ((HOC₂H₂)₃N) = 5.9 × 10^-7

A) 2.75

B) 4.89

C) 9.11

D) 11.25

E) none of the above

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Q1) Barium sulfate (BaSO₄) is a slightly soluble salt, with K_sp = 1.1 × 10^-10. What mass of Ba²⁺ ions will be present in 1.0 L of a saturated solution of barium sulfate?

A) < 10^-7 g

B) 1.0 × 10^-5 g

C) 0.0014 g

D) 0.0024 g

E) > 0.05 g

Q2) Which of the following aqueous mixtures would be a buffer system?

A) HCl, NaCl

B) HNO₃, NaNO₃

C) H₃PO₄, H₂PO₄^-

D) H₂SO₄, CH₃COOH

E) NH₃, NaOH

Q3) Fe(NO₃)₃ (0.00100 mol) and KSCN (0.200 mol) are added to water to make exactly 1 liter of solution. The red complex ion FeSCN²⁺ is produced. Calculate the concentrations of Fe³⁺(aq) and FeSCN²⁺(aq) at equilibrium, if K_f of the FeSCN²⁺ is 8.9 × 10².

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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction

of Chemical Reactions

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Q1) Which of the following is true for pure oxygen gas, O₂(g) at 25°C?

A) .\(\Delta\)H° _f > 0

B) .\(\Delta\)H° _f < 0

C) .\(\Delta\)G° _f > 0

D) .\(\Delta\)G° _f < 0

E) S° > 0

Q2) Compare one mole of ice with one mole of liquid water, both at 1.0 atm and 0°C. The melting point of ice at 1.0 atm is 0°C. For the process H₂O(s) \(\to\) H₂O(l) under these conditions predict whether each of the following quantities will be greater than, less than, or equal to, zero . Explain each prediction in one sentence. a. \(\Delta\)H° b. \(\Delta\)S° c. \(\Delta\)G°

Q3) Which of the following is necessary for a process to be spontaneous?

A) .\(\Delta\)H_sys < 0

B) .\(\Delta\)S_sys > 0

C) .\(\Delta\)S_surr< 0

D) .\(\Delta\)S_univ > 0

E) .\(\Delta\)G_sys = 0

Q4) State the second and third laws of thermodynamics.

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) Consider the following balanced redox reaction Mn²⁺(aq) + S₂O₈^2-(aq) + 2H₂O(l) \(\to\) MnO₂(s) + 4H⁺(aq) + 2SO₄^2-(aq)

Which of the following statements is true?

A) Mn²⁺(aq) is the oxidizing agent and is reduced.

B) Mn²⁺(aq) is the oxidizing agent and is oxidized.

C) Mn²⁺(aq) is the reducing agent and is oxidized.

D) Mn²⁺(aq) is the reducing agent and is reduced.

E) Manganese does not change its oxidation number in this reaction.

Q2) In the shorthand notation for cells, a single vertical line represents a salt bridge. A)True B)False

Q3) a. Write a balanced equation to represent the overall reaction you would expect to occur in the electrolysis of molten KCl.

b. Write a balanced equation to represent the overall reaction you would expect to occur in the electrolysis of aqueous KCl.

c. Clearly explain why the products of the two processes are not the same.

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Chapter 22: The Elements in Nature and Industry

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Q1) Briefly describe the three main pathways for nitrogen fixation.

Q2) Compared to the asbestos diaphragm-cell used in the chlor-alkali process, list a. one main advantage and one disadvantage of mercury cells. b. two important advantages of polymer membrane-cells.

Q3) The alkali metals are isolated from non-aqueous systems. Why is this necessary?

A) The electrolysis of aqueous solutions of the alkali metals requires more energy than electrolysis of the molten salts.

B) The dissolved alkali earth halides are too reactive to be electrolyzed.

C) The aqueous metal ions are more difficult to reduce than water.

D) The reduction potentials of the alkali metals are more positive than the reduction potential of water.

E) The aqueous metal ions react violently with water.

Q4) What gas is produced at the anode in the Hall-Heroult process for the electrolytic production of aluminum?

Q5) The purpose of anodizing aluminum is to remove the oxide layer from the metal's surface.

A)True

B)False

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) 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

Q2) Apply the valence bond theory to predict the electronic structure and hybridization pattern of chromium in the complex ion Cr(NH₃)₆³⁺.

Q3) A characteristic of ligands is that A) they are Lewis acids.

B) they are Lewis bases.

C) they are ions.

D) they are electron pair acceptors.

E) they are Brønsted-Lowry acids.

Q4) The inner transition series of elements arise from the filling of f orbitals.

A)True

B)False

Q5) Why is the +2 oxidation state so common among transition elements?

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) An 85-kg person exposed to barium-141 receives 2.5 × 10⁵ \(\beta\) particles, each with an energy of 5.2 × 10^-13 J. How many rads does the person receive?

A) 2.4 × 10^-20

B) 1.5 × 10^-7

C) 1.8 × 10^-16

D) 6.1 × 10^-15

E) none of the above

Q2) So-called "magic numbers" of particles are thought to convey extra stability to certain nuclei. These magic numbers refer to which of the following particles?

A) protons only

B) electrons only

C) positrons only

D) neutrons only

E) protons and neutrons

Q3) Write a complete, balanced equation to represent the beta decay of thallium-207.

Q4) Write a complete, balanced equation to represent the formation of manganese-55 by the beta decay of another nuclide.

Q5) Write a complete, balanced equation to represent the alpha decay of radon-210.

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