Chemistry for Life Sciences Exam Materials - 1830 Verified Questions

Chemistry for Life Sciences

Exam Materials

Course Introduction

Chemistry for Life Sciences provides an introduction to the fundamental concepts of chemistry with a focus on their applications in biological systems. The course covers topics such as atomic and molecular structure, chemical bonding, thermodynamics, solutions, acids and bases, and chemical reactions that are essential for understanding processes in living organisms. Emphasis is placed on the relevance of chemistry to topics in health, medicine, physiology, and cellular biology, with practical laboratory components reinforcing theoretical knowledge and developing experimental skills relevant to the life sciences.

Recommended Textbook

Principles of General Chemistry 2nd Edition by Martin Silberberg

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23 Chapters

1830 Verified Questions

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Page 2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) The density of mercury,the only metal to exist as a liquid at room temperature,is 13.6 g/cm³.What is that density in pounds per cubic inch?

A)849 lb/in³

B)491 lb/in³

C)376 lb/in³

D)0.491 lb/in³

E)1.83 \(\times\) 10^-3 lb/in³

Answer: D

Q2) Which of the following activities is not a part of good science?

A)proposing a theory

B)developing a hypothesis

C)making quantitative observations

D)designing experiments

E)indulging in speculation

Answer: E

Q3) The potential energy of a car moving on a level road does not depend on its speed.

A)True

B)False

Answer: True

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Page 3

Chapter 2: The Components of Matter

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Sample Questions

Q1) In nature,some elements exist as molecules,while others do not.

A)True

B)False

Answer: True

Q2) A column of the periodic table is called a A)group

B)period

C)isotopic mixture

D)pillar

E)shell

Answer: A

Q3) Which one of the following formulas of ionic compounds is the least likely to be correct?

A)NH₄Cl

B)Ba(OH)₂

C)Na₂SO₄

D)Ca₂NO₃

E)Cu(CN)₂

Answer: D

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) 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 H₂ react until the limiting reagent is completely consumed?

Answer: 3.52 g of hydrogen gas remains.

Q2) One mole of methane (CH₄)contains a total of 3 \(\times\) 10²⁴ atoms.

A)True

B)False

Answer: True

Q3) Calcium fluoride,CaF₂,is a source of fluorine and is used to fluoridate drinking water.Calculate its molar mass.

A)118.15 g/mol

B)99.15 g/mol

C)78.07 g/mol

D)59.08 g/mol

E)50.01 g/mol

Answer: C

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Page 5

Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) Select the classification for the following reaction.

Fe²⁺(aq)+ 2OH^-(aq) \(\to\) Fe(OH)₂(s)

A)precipitation

B)acid-base

C)redox

D)decomposition

E)None of these choices is correct.

Q2) Select the precipitate that forms when aqueous lead(II)nitrate reacts with aqueous sodium sulfate.

A)NaNO₃

B)Na₂NO₃

C)PbSO₄

D)Pb₂SO₄

E)PbS

Q3) Which one of the following substances is a strong acid?

A)HNO₃

B)H₂CO₃

C)NH₃

D)CH₃COOH

E)H₃PO₄

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) Magnesium metal (0.100 mol)and a volume of aqueous hydrochloric acid that contains 0.500 mol of HCl are combined and react to completion.How many liters of hydrogen gas,measured at STP,are produced?

Mg(s)+ 2HCl(aq) \(\to\) MgCl₂(aq)+ H₂(g)

A)2.24 L of H₂

B)4.48 L of H₂

C)5.60 L of H₂

D)11.2 L of H₂

E)22.4 L of H₂

Q2) Ima Chemist found the density of Freon-11 (CFCl₃)to be 5.58 g/L under her experimental conditions.Her measurements showed that the density of an unknown gas was 4.38 g/L under the same conditions.What is the molar mass of the unknown?

A)96.7 g/mol

B)108 g/mol

C)127 g/mol

D)165 g/mol

E)175 g/mol

Q3) Calculate the density in g/L of gaseous SF₆ at 50.0°C and 650.torr.

Q4) State Avogadro's Law.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) The reaction 2NaOH(aq)+ H₂SO₄(aq) \(\to\) Na₂SO₄(aq)+ 2H₂O(l) was studied in a coffee cup calorimeter.100.mL portions of 1.00 M aqueous NaOH and H₂SO₄,each at 24.0°C,were mixed.The maximum temperature achieved was 30.6°C.Neglect the heat capacity of the cup and the thermometer,and assume that the solution of products has a density of exactly 1 g/mL and a specific heat capacity of 4.18 J/(g·K).

a.Calculate the heat of reaction,q,in J.

b.Calculate \(\Delta\)H,the heat (enthalpy)of reaction,in kJ/mol of Na₂SO₄ produced.

Q2) When Karl Kaveman adds chilled grog to his new granite mug,he removes 10.9 kJ of energy from the mug.If it has a mass of 625 g and was at 25°C,what is its new temperature? Specific heat capacity of granite = 0.79 J/(g·K)

A)3°C

B)14°C

C)22°C

D)47°C

E)None of these choices is correct.

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) Which word best describes the phenomenon which gives rise to a rainbow?

A)reflection

B)dispersion

C)diffraction

D)interference

E)deflection

Q2) a.Use the Bohr equation to calculate the energy needed to ionize a hydrogen atom from its ground state.

b.What is the minimum wavelength of a photon needed for it to have the energy needed in (a)?

Q3) Atomic orbitals developed using quantum mechanics

A)describe regions of space in which one is most likely to find an electron.

B)describe exact paths for electron motion.

C)give a description of the atomic structure which is essentially the same as the Bohr model.

D)allow scientists to calculate an exact volume for the hydrogen atom.

E)are in conflict with the Heisenberg Uncertainty Principle.

Q4) The Rydberg equation is an example of an empirical equation.

A)True

B)False

Page 9

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Select the element with the greatest metallic character.

A)Li

B)Ca

C)Al

D)Pb

E)Cs

Q2) Energy states of atoms containing more than one electron arise from nucleus-electron and electron-electron interactions.Which of the following statements correctly describes these effects?

A)Larger nuclear charge lowers energy,more electrons in an orbital lowers energy.

B)Larger nuclear charge lowers energy,more electrons in an orbital increases energy.

C)Smaller nuclear charge lowers energy,more electrons in an orbital lowers energy.

D)Smaller nuclear charge lowers energy,more electrons in an orbital increases energy.

E)None of these statements is generally correct.

Q3) State Hund's rule,and show how it applies to the ground state of phosphorus atoms.

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Which of the following contains ionic bonding?

A)CO

B)SrF₂

C)Al

D)OCl₂

E)HCl

Q2) Arrange the following bonds in order of increasing bond strength.

A)C-I < C-Br < C-Cl < C-F

B)C-F < C-Cl < C-Br < C-I

C)C-Br < C-I < C-Cl < C-F

D)C-I < C-Br < C-F< C-Cl

E)None of these orders is correct.

Q3) The majority of elements are good electrical conductors when in solid form. A)True

B)False

Q4) The lattice energy of large ions is greater in magnitude than that of small ions of the same charge.

A)True

B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) For the chlorate ion,ClO₃^-,draw two different valid Lewis structures,as follows:

a.a structure in which the octet rule is obeyed

b.a structure in which formal charges are minimized

Q2) Draw all important resonance structures of the nitrate ion,NO₃

Q3) 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.

Q4) According to VSEPR theory,a molecule with the general formula AX₄ will have a ______ molecular shape.

A)bent

B)trigonal planar

C)trigonal pyramidal

D)square planar

E)tetrahedral

Q5) Boron never achieves an octet in any of its compounds.

A)True

B)False

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Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) Valence bond theory predicts that carbon will use _____ hybrid orbitals in the carbonate anion,CO₃^2-.

A)sp

B)sp²

C)sp³

D)sp³d

E)sp³d²

Q2) Explain what is meant by a node (or nodal plane)in a molecular orbital and draw sketches of the following orbitals,indicating at least one nodal plane in each one.

a.a \(\sigma\)* orbital

b.a \(\pi\)* orbital

Q3) Overlap of two sp² hybrid orbitals produces a \(\pi\) bond.

A)True

B)False

Q4) Valence bond theory explains the bonding in diatomic molecules such as HCl without resorting to the use of hybrid orbitals.

A)True

B)False

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Page 13

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

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Sample Questions

Q1) In the packing of identical atoms with cubic unit cells,the packing efficiency increases as the coordination number increases.

A)True

B)False

Q2) Which of the following statements about the packing of monatomic solids with different unit cells is incorrect?

A)The coordination number of atoms in hcp and fcc structures is 12.

B)The coordination number of atoms in simple cubic structures is 6.

C)The coordination number of atoms in bcc structures is 8.

D)A bcc structure has a higher packing efficiency than a simple cubic structure.

E)A bcc structure has a higher packing efficiency than a fcc structure.

Q3) The strongest intermolecular interactions between ethyl alcohol (CH₃CH₂OH)molecules arise from

A)dipole-dipole forces.

B)London dispersion forces.

C)hydrogen bonding.

D)ion-dipole interactions.

E)carbon-oxygen bonds.

Q4) How do the electrical properties of semiconductors differ from those of metals?

Page 14

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Chapter 13: The Properties of Solutions

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Sample Questions

Q1) The concentration of iodine in sea water is 60.parts per billion by mass.If one assumes that the iodine exists in the form of iodide anions,what is the molarity of iodide in sea water? (The density of sea water is 1.025 g/mL. )

A)4.8 \(\times\) 10^-13 M

B)4.8 \(\times\) 10^-10 M

C)4.8 \(\times\) 10^-7 M

D)4.7 \(\times\) 10^-4 M

E)4.7 \(\times\) 10^-1 M

Q2) How many moles of bromide ions are present in 750.0 mL of 1.35 M MgBr₂?

A)0.506 mol

B)1.01 mol

C)2.03 mol

D)3.04 mol

E)None of these choices is correct.

Q3) The heat of solution is the total enthalpy change when a solution is formed from the separated solute and solvent.

A)True

B)False

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Page 15

Chapter 14: The Main-Group Elements: Applying Principles of Bonding

and Structure

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Sample Questions

Q1) Predict the products formed from the following set of reactants.

FeSe(s)+ HCl(aq) \(\to\)

A)Se(s)+ FeCl₂(aq)

B)SeCl₂(s)+ Fe(s)+ H₂(g)

C)FeH₂(s)+ SeCl₂(l)

D)H₂Se(g)+ Fe(s)+ Cl₂(g)

E)H₂Se(g)+ FeCl₂(aq)

Q2) The chemical that ranks first in production among all industrial chemicals is

A)NH₃,ammonia.

B)H₃PO₄,phosphoric acid.

C)NaOH,sodium hydroxide.

D)Na₂CO₃,sodium carbonate.

E)H₂SO₄,sulfuric acid.

Q3) Neon does not form any known compounds.

A)True

B)False

Q4) Potassium is a strong reducing agent.

A)True

B)False

Q5) Name two different classes (types)of hydride.

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) Name the three component parts of a nucleotide.

Q2) Explain what is meant by "complementary" in the context of DNA strands.

Q3) The backbone of protein molecules consists of repeating N-C-C-O units.

A)True

B)False

Q4) A characteristic reaction of alkenes is addition.

A)True

B)False

Q5) In one sentence,what is the general requirement for a molecule to be optically active?

Q6) All the carbon atoms in a molecule of benzene lie in the same plane.

A)True B)False

Q7) Which of the following polymers is a condensation polymer?

A)polystyrene

B)Teflon®

C)polyvinylchloride

D)polypropylene

E)Dacron®

Q8) Draw and name all stable molecules with the formula Page 17

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) A chemical reaction of the general type A \(\to\) 2B is first-order,with a rate constant of 1.52 \(\times\) 10^-4 s^-1.

a.Calculate the half-life of A.

b.Assuming the initial concentration of A is 0.067 mol L^-1,calculate the time needed for the concentration to fall to 0.010 mol L^-1.

Q2) The decomposition of SOCl₂ is first-order in SOCl₂.If the half-life for the reaction is 4.1 hr,how long would it take for the concentration of SOCl₂ to drop from 0.36 M to 0.045 M?

A)0.52 hr

B)1.4 hr

C)12 hr

D)33 hr

E)> 40 hr

Q3) All second-order reactions are bimolecular reactions.

A)True

B)False

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Page 19

Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) Unless \(\Delta\)H°_rxn = 0,a change in temperature will affect the value of the equilibrium constant K_c.

A)True

B)False

Q2) In order to write the correct mass-action expression for a reaction one must

A)know the kinetic rate law for the reaction.

B)know the mechanism for the reaction.

C)have a properly balanced chemical equation.

D)have values for the concentrations of the reactants.

E)know the limiting reactant.

Q3) For some gas-phase reactions,K_p = K_c.

A)True

B)False

Q4) Which of the following has an effect on the magnitude of the equilibrium constant?

A)activation energy of the forward reaction

B)concentrations of the reactants and products

C)presence of a catalyst

D)change in volume of container

E)change in temperature

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) What is the value of K_a for the methylammonium ion,CH₃NH₃⁺?

K_b(CH₃NH₂)= 4.4 \(\times\) 10^-4

A)4.4 \(\times\) 10^-4

B)4.8 \(\times\) 10^-6

C)4.4 \(\times\) 10^-10

D)2.3 \(\times\) 10^-11

E)4.4 \(\times\) 10^-18

Q2) Define a base according to the Brønsted-Lowry theory.Write a balanced equation in support of this definition,in which the base is acting as a Brønsted-Lowry base but not as an Arrhenius base.

Q3) Which one of the following will give a solution with a pH > 7,but is not an Arrhenius base in the strict sense?

A)CH₃NH₂

B)NaOH

C)CO₂

D)Ca(OH)₂

E)CH₄

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Formic acid is a monoprotic acid with a K_a value of 1.8 \(\times\) 10^-4 at 25°C.

a.Calculate the pH of a 0.200 mol L^-1 solution of the acid,making any reasonable approximations.

b.If 0.0050 mol of NaOH is added to 100.mL of the solution in (a),calculate the pH of the resultant buffer.

Q2) What is the maximum amount of sodium sulfate that can be added to 1.00 L of 0.0020 M Ca(NO₃)₂ before precipitation of calcium sulfate begins? K_sp = 2.4 \(\times\) 10^-5 for calcium sulfate

A)1.2 \(\times\) 10^-2 mol

B)4.9 \(\times\) 10^-3 mol

C)3.5 \(\times\) 10^-3 mol

D)1.2 \(\times\) 10^-5 mol

E)4.8 \(\times\) 10^-8 mol

Q3) The equivalence point in a titration is defined as the point when the indicator changes color.

A)True

B)False

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22

Chapter

Direction of Chemical Reactions

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Q1) 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°

Q2) You are given pure samples of ammonia,NH₃(g),and nitrogen trifluoride,NF₃(g).What prediction would you make concerning their standard molar entropies at 298 K?

A)S°_ammonia > S°_{nitrogen trifluoride}

B)S°_ammonia < S°_{nitrogen trifluoride}

C)S°_ammonia \(\approx\) S°_{nitrogen trifluoride}

D)Other conditions need to be specified before a reliable prediction can be made.

E)Even if more conditions are specified,a reliable prediction cannot be made.

Q3) State the second and third laws of thermodynamics.

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Page 23

Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) A battery that cannot be recharged is a A)fuel cell.

B)primary battery.

C)secondary battery.

D)simple battery.

E)flow battery.

Q2) 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.

Q3) In the electrolyte of an electrochemical cell,current is carried by anions moving toward the anode and cations moving in the opposite direction.

A)True

B)False

Q4) Electrolytic cells utilize electrical energy to drive non-spontaneous redox reactions. A)True

B)False

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Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) A certain transition metal complex has the formula MX₄²⁺.If the metal ion has a d⁸ electron configuration,what is the shape of the complex?

A)octahedral

B)square pyramid

C)tetrahedral

D)trigonal pyramid

E)square planar

Q2) Give the systematic name for [Cu(NH₃)₄]Cl₂.

A)dichlorotetraamminecuprate(II)

B)tetraamminecopper(II)chloride

C)copper(II)ammonium chloride

D)tetraaminocopper(II)chloride

E)None of these choices is correct.

Q3) Give the systematic name for Cr(CO)₃(NH₃)₃.

A)chromiumtriaminotricarbonyl

B)triamminechromium carbonate

C)triamminetricarbonylchromate(0)

D)triamminetricarbonylchromium(0)

E)None of these choices is correct.

Page 25

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) An isotope with Z > 83,which lies close to the band of stability,will generally decay through

A)(\(\alpha\) decay.)

B)(\(\beta\) decay.)

C)(\(\gamma\) decay.)

D)positron decay.

E)electron capture.

Q2) Gamma rays are high energy electrons.

A)True

B)False

Q3) Explain how the number of protons and neutrons in a radioactive nucleus can be used to predict its probable mode of decay.Illustrate your answer with a schematic graph,properly labeled,showing stable nuclides (nuclei)in relation to number of protons and neutrons.

Q4) An isotope with a high value of N/Z will tend to decay through

A)(\(\alpha\) decay.)

B)(\(\beta\) decay.)

C)positron decay.

D)electron capture.

E)(\(\gamma\) decay.)

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