Survey of Chemistry Exam Bank - 1830 Verified Questions

Survey of Chemistry Exam Bank

Course Introduction

Survey of Chemistry provides a comprehensive introduction to the core principles of chemistry, incorporating both inorganic and organic topics. Designed for non-science majors and those seeking a foundational understanding, the course explores atomic structure, chemical bonding, stoichiometry, states of matter, solutions, acids and bases, and basic biochemistry. Emphasis is placed on conceptual understanding and real-world applications, enabling students to appreciate the role of chemistry in everyday life, environmental issues, and health. Laboratory activities complement theoretical concepts, fostering critical thinking and scientific literacy.

Recommended Textbook

Principles of General Chemistry 2nd Edition by Martin Silberberg

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23 Chapters

1830 Verified Questions

1830 Flashcards

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Page 2

Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) The appropriate number of significant figures in the result of 15.234 - 15.208 is

A)1

B)2

C)3

D)4

E)5

Answer: B

Q2) You prepare 1000.mL of tea and transfer it to a 1.00 quart pitcher for storage.Which of the following statements is true?

A)The pitcher will be filled to 100% of its capacity with no tea spilled.

B)The pitcher will be filled to about 95% of its capacity.

C)The pitcher will be filled to about 50% of its capacity.

D)The pitcher will be completely filled and a small amount of tea will overflow.

E)The pitcher will be completely filled and most of the tea will overflow.

Answer: D

Q3) The number 6.0448,rounded to 3 decimal places,becomes 6.045.

A)True

B)False

Answer: True

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3

Chapter 2: The Components of Matter

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Sample Questions

Q1) The formula C₉H₂₀ is an empirical formula.

A)True

B)False

Answer: True

Q2) What is the formula for magnesium sulfide?

A)MgS

B)MgS₂

C)Mg₂S

D)Mg₂S₃

E)MgSO₄

Answer: A

Q3) Which one of the following combinations of names and formulas of ions is incorrect?

A)NH₄⁺ ammonium

B)S^2- sulfide

C)CN^- cyanide

D)S₂O₃^2- thiosulfate

E)ClO₃^- perchlorate

Answer: E

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4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Propane,C₃H₈,is commonly provided as a bottled gas for use as a fuel.In 0.200 mol of propane, A)what is the mass of propane?

B)7.21 g

C)1.20 \(\times\) 10²³ C₃H₈ molecules

D)9.64 \(\times\) 10²³ H atoms

Answer: A

Q2) Aluminum metal dissolved in hydrochloric acid as follows: 2Al(s)+ 6HCl(aq) \(\to\) 2AlCl₃(aq)+ 3H₂(g)

a.What is the minimum volume of 6.0 M HCl(aq)needed to completely dissolve 3.20 g of aluminum in this reaction?

b.What mass of AlCl₃ would be produced by complete reaction of 3.20 g of aluminum?

Answer: a.59.3 mL

b.15.8 g

Q3) You are provided with a 250 mL volumetric flask,deionized water and solid NaOH.How much NaOH should be weighed out in order to make 250.mL of 0.100 M solution?

Answer: 1.000 g

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5

Chapter 4: Three Major Classes of Chemical Reactions

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Sample Questions

Q1) Select the classification for the following reaction. Fe(s)+ 2Fe³⁺(aq) \(\to\) 3Fe²⁺(aq)

A)precipitation

B)acid-base

C)redox

D)decomposition

E)None of these choices is correct.

Q2) Potassium chloride,KCl,sodium sulfate,Na₂SO₄,glucose,C₆H₁₂O₆,carbon dioxide,CO₂ and ammonium phosphate, (NH₄)₃PO₄,are soluble in water.Which one produces the largest number of dissolved particles per mole of dissolved solute?

A)KCl

B)Na₂SO₄

C)C₆H₁₂O₆

D)CO₂

E)(NH₄)₃PO₄

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) A 1.00-L sample of a pure gas weighs 0.785 g and is at 0.965 atm and 29.2°C.

a.What is the molar mass of the gas?

b.If the volume and temperature are kept constant while 0.400 g of the same gas are added to that already in the container,what will the new pressure be?

Q2) A 250.0-mL sample of ammonia,NH₃(g),exerts a pressure of 833 torr at 42.4°C.What mass of ammonia is in the container?

A)0.0787 g

B)0.180 g

C)8.04 g

D)17.0 g

E)59.8 g

Q3) A 20.0-L container holds 15.3 mol of Cl₂ gas at 227°C.

a.Calculate the pressure in atmospheres,assuming ideal behavior.

b.Calculate the pressure in atmospheres,assuming van der Waals behavior.The van der Waals constants for Cl₂ are a = 6.49 atm·L²/mol² and b = 0.0562 L/mol.

Q4) A 255-mL gas sample weighing 0.292 g is at 52,810 Pa and 127°C.

a.How many moles of gas are present?

b.What is the molar mass of the gas?

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) A common laboratory reaction is the neutralization of an acid with a base.When 50.0 mL of 0.500 M HCl at 25.0°C is added to 50.0 mL of 0.500 M NaOH at 25.0°C in a coffee cup calorimeter,the temperature of the mixture rises to 28.2°C.What is the heat of reaction per mole of acid? Assume the mixture has a specific heat capacity of 4.18 J/(g·K)and that the densities of the reactant solutions are both 1.00 g/mL.

A)670 J

B)1300 J

C)27 kJ

D)54 kJ

E)> 100 kJ

Q2) A system contracts from an initial volume of 15.0 L to a final volume of 10.0 L under a constant external pressure of 0.800 atm.The value of w,in J,is

A)-4.0 J

B)4.0 J

C)-405 J

D)405 J

E)4.05 \(\times\) 10³ J

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Page 8

Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) What are the possible values for the following quantum numbers in an atom?

a.n

b.l

c.m_l

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) Use the Bohr equation to calculate the energy of

a.the largest energy absorption or emission process involving the n = 2 state of the hydrogen atom.

b.the smallest energy absorption or emission process involving the n = 2 state of the hydrogen atom.

Q4) Explain the context and meanings of the terms "orbit" and "orbital",making a clear distinction between them.

Q5) What is the speed of an electron in m/s if its wavelength is 0.155 nm?

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Which of the following elements has the smallest atomic size?

A)Na

B)Ar

C)K

D)Ca

E)Kr

Q2) In moving down a group in the periodic table,the oxides of the elements become more acidic in nature.

A)True

B)False

Q3) Which one of the following equations correctly represents the process involved in the electron affinity of X?

A)X(g) \(\to\) X⁺(g)+ e^-

B)X⁺(g) \(\to\) X⁺(aq)

C)X⁺(g)+ e^- \(\to\) X(g)

D)X(g)+ e^- \(\to\) X^-(g)

E)X⁺(g)+ Y^-(g) \(\to\) XY(s)

Q4) Define what is meant by electron affinity,and write a balanced chemical equation to represent the relevant process for element Y.

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) In not more than three sentences,describe the electron arrangement responsible for bonding in solid SrCl₂.

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 more C-O and O-H bonds there are in a substance,the greater will be the amount of heat released when a fixed mass of the substance is burned.

A)True

B)False

Q4) The electrostatic energy of two charged particles is inversely proportional to the square of the distance between them.

A)True

B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Predict the ideal bond angles in AsCl₃ using the molecular shape given by the VSEPR theory.

A)90°

B)109°

C)120°

D)180°

E)between 110 and 120°

Q2) What is the molecular shape of NO₂^- as predicted by the VSEPR theory?

A)linear

B)trigonal planar

C)bent

D)tetrahedral

E)resonant

Q3) Draw Lewis structures,showing all valence electrons,for the following species:

a.S^2- b.CO

c.SO₂

d.CH₃OH

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Page 12

Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) Valence bond theory predicts that iodine will use _____ hybrid orbitals in ICl₂^-.

A)sp²

B)sp³

C)sp³d

D)sp³d²

E)None of these choices is correct.

Q2) One can safely assume that the 3s- and 3p-orbitals will form molecular orbitals similar to those formed when 2s- and 2p-orbitals interact.According to molecular orbital theory,what will be the bond order for the Cl₂⁺ ion?

A)0.5

B)1

C)1.5

D)2

E)None of these choices is correct.

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

A)True

B)False

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

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

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Sample Questions

Q1) Which of the following is true about kinetic energy,E_k,and potential energy,E_p,when ethyl alcohol at 40°C is compared with ethyl alcohol at 20°C?

A)E_k(40°C)< E_k(20°C);E_p(40°C) \(\approx\) E_p(20°C)

B)E_k(40°C)> E_k(20°C);E_p(40°C) \(\approx\)

E_p(20°C)

C)E_p(40°C)< E_p(20°C);E_k(40°C) \(\approx\)

E_k(20°C)

D)E_p(40°C)> E_p(20°C);E_k(40°C) \(\approx\) E_k(20°C)

E)E_p(40°C)> E_p(20°C);E_k(40°C)> E_k(20°C)

Q2) The strongest intermolecular interactions between hydrogen fluoride (HF)molecules arise from

A)dipole-dipole forces.

B)London dispersion forces.

C)hydrogen bonding.

D)ion-dipole interactions.

E)ionic bonds.

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Chapter 13: The Properties of Solutions

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Sample Questions

Q1) A 0.100 m K₂SO₄ solution has a freezing point of -0.43°C.What is the van't Hoff factor for this solution?

K_f = 1.86°C/m

A)0.77

B)1.0

C)2.3

D)3.0

E)>3.0

Q2) 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

Q3) In general,water is a good solvent for both polar and non-polar compounds.

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 for the following set of reactants.

PCl₃(l)+ H₂O(l) \(\to\)

A)H₃PO₃(aq)+ HCl(aq)

B)H₃PO₄(aq)+ Cl₂(g)

C)PH₃(g)+ HCl(aq)+ O₂(g)

D)P₂O₅(s)+ HCl(aq)

E)PCl₅(l)+ PH₃(g)+ O₂(g)

Q2) Which,if any,of the following generalized formulas does not exist for interhalogen compounds? (X and Y represent different halogens. )

A)XY

B)XY₂

C)XY₃

D)XY₅

E)All the above can represent stable interhalogen compounds.

Q3) Write balanced equations,showing all reactants and products,to represent a.roasting of limestone (CaCO₃)to give lime.

b.tetraphosphorus decaoxide (P₄O₁₀)reacting with water to produce phosphoric acid.

Q4) 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) A characteristic reaction of alkanes is addition.

A)True

B)False

Q2) Secondary amines have the general formula RNH₂.

A)True

B)False

Q3) Explain what is meant by "complementary" in the context of DNA strands.

Q4) 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?

Q5) Draw and name all stable molecules with the formula C₅H₁₂.

Q6) The backbone of protein molecules consists of repeating N-C-C-O units.

A)True

B)False

Q7) Esters can be formed by the dehydration-condensation of a carboxylic acid and an alcohol.

A)True

B)False

Q8) In one sentence,what is the general requirement for a molecule to be optically Page 17

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) At 25.0°C,a rate constant has the value 5.21 \(\times\) 10^-8 L mol^-1 s^-1.If the activation energy is 75.2 kJ/mol,calculate the rate constant when the temperature is 50.0°C.

Q2) The kinetics of the decomposition of dinitrogen pentaoxide is studied at 50°C and at 75°C.Which of the following statements concerning the studies is correct?

A)The rate at 75°C will be greater than the rate at 50°C because the activation energy will be lower at 75°C than at 50°C.

B)The rate at 75°C will be greater than the rate at 50°C because the activation energy will be higher at 75°C than at 50°C.

C)The rate at 75°C will be less than the rate at 50°C because the molecules at higher speeds do not interact as well as those at lower speeds.

D)The rate at 75°C will be greater than at 50°C because the concentration of a gas increases with increasing temperature.

E)The rate at 75°C will be greater than the rate at 50°C because the number of molecules with enough energy to react increases with increasing temperature.

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) When a chemical system is at equilibrium,

A)the concentrations of the reactants are equal to the concentrations of the products.

B)the concentrations of the reactants and products have reached constant values.

C)the forward and reverse reactions have stopped.

D)the reaction quotient,Q,has reached a maximum.

E)the reaction quotient,Q,has reached a minimum.

Q2) 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

Q3) When a reaction system reaches equilibrium,the forward and reverse reactions stop. A)True

B)False

Q4) For some gas-phase reactions,K_p = K_c.

A)True

B)False

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20

Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) All strong acids have weak conjugate bases.

A)True

B)False

Q2) 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

Q3) Iodine trichloride,ICl₃,will react with a chloride ion to form ICl₄^-.Which species,if any,acts as a Lewis acid this reaction?

A)ICl₄^-

B)ICl ₃

C)Cl^-

D)the solvent

E)None of the species acts as a Lewis acid in this reaction.

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) When a strong acid is titrated with a strong base,the pH at the equivalence point

A)is greater than 7.0.

B)is equal to 7.0.

C)is less than 7.0,but is not 3.5.

D)is equal to the pK_a of the acid.

E)is equal to 3.5.

Q2) Calculate the solubility of silver chromate,Ag₂CrO₄,in 0.005 M Na₂CrO₄.K_sp = 2.6 \(\times\) 10^-12

A)1.4 \(\times\) 10^-4 M

B)3.4 \(\times\) 10^-5 M

C)1.1 \(\times\) 10^-5 M

D)1.6 \(\times\) 10^-6 M

E)< 1.0 \(\times\) 10^-6 M

Q3) Silver phosphate,Ag₃PO₄,is an ionic compound with a solubility product constant K_sp of 2.6 \(\times\) 10^-18.Calculate the solubility of this compound in a.pure water.

b.0.20 mol L^-1 Na₃PO₄ solution.

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Page 22

Chapter

Direction of Chemical Reactions

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Sample Questions

Q1) For any reaction,if \(\Delta\)G° > 0,then K < 1.

A)True

B)False

Q2) 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)

Q3) Which of the following is always true for an endothermic process?

A)q _sys > 0,\(\Delta\)S_surr < 0

B)q_sys < 0,\(\Delta\)S_surr > 0

C)q_sys < 0,\(\Delta\)S_surr < 0

D)q_sys > 0,\(\Delta\)S_surr > 0

E)w < 0

Q4) a.Explain what is meant by a spontaneous process.

b.Is a spontaneous process necessarily a rapid one? Explain,and provide a real reaction as an example to illustrate your answer.

Q5) 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) The lead-acid battery is an example of a secondary battery. A)True

B)False

Q2) 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

Q3) When metal A is placed in a solution of metal ions B²⁺,a reaction occurs between A and B²⁺,and metal ions A²⁺ appear in the solution.When metal B is placed in acid solution,gas bubbles form on its surface.When metal A is placed in a solution of metal ions C²⁺,no reaction occurs.Which of the following reactions would not occur spontaneously?

A)C(s)+ 2H⁺(aq) \(\to\) H₂(g)+ C²⁺(aq)

B)C(s)+ A²⁺(aq) \(\to\) A(s)+ C²⁺(aq)

C)B(s)+ C²⁺(aq) \(\to\) C(s)+ B²⁺(aq)

D)A(s)+ 2H⁺(aq) \(\to\) H₂(g)+ A²⁺(aq)

E)B(s)+ 2H⁺(aq) \(\to\) H₂(g)+ B²⁺(aq)

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Page 24

Chapter 22: The Transition Elements and Their Coordination Compounds

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Q1) The ground state electron configuration of Cr²⁺ is

A)[Ar]4s¹3d⁵

B)[Ar]4s²3d⁴

C)[Ar]3d⁴

D)[Ar]4s¹3d³

E)[Ar]4s²3d²

Q2) Which of the following ligands is most likely to form a low-spin octahedral complex with iron(III)?

A)Cl^-

B)H₂O

C)NH₃

D)OH^-

E)CO

Q3) The M²⁺ ions of the first transition series of elements all have the general electronic configuration [Ar]4s²3d^x,where x is an integer from 1 to 8.

A)True

B)False

Q4) What is the difference between a coordination compound and a complex ion?

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

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Chapter 23: Nuclear Reactions and Their Applications

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Q1) Calcium-39 undergoes positron decay.Each positron carries 5.49 MeV of energy.How much energy will be emitted when 0.0025 mol of calcium-39 decays?

A)13.2 kJ

B)1.32 \(\times\) 10⁴ kJ

C)1.32 \(\times\) 10⁶ kJ

D)1.32 \(\times\) 10⁹ kJ

E)None of these choices is correct.

Q2) The radiochemist,Will I.Glow,studied thorium-232 and found that 2.82 \(\times\) 10^-7 moles emitted 8.42 \(\times\) 10⁶ \(\alpha\) particles in one year.What is the decay constant for thorium-232?

A)3.35 \(\times\) 10^-14 yr^-1

B)4.96 \(\times\) 10^-11 yr^-1

C)1.40 \(\times\) 10¹⁰ yr^-1

D)2.99 \(\times\) 10¹³ yr^-1

E)None of these choices is correct.

Q3) After 4 half-lives,the fraction of a radioactive isotope which still remains is approximately one eighth.

A)True

B)False

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Page 26