Chemistry for Physical Sciences Final Exam Questions - 2309 Verified Questions

Chemistry for Physical Sciences

Final Exam Questions

Course Introduction

This course provides a foundational understanding of chemistry tailored for students in the physical sciences. Emphasizing the principles of atomic and molecular structure, chemical bonding, thermodynamics, kinetics, and equilibrium, the course explores how chemical phenomena underpin broader concepts in physics and related fields. Students will develop problem-solving skills as they examine the quantitative relationships in chemical reactions, properties of gases, solutions, and chemical systems. Laboratory sessions reinforce theoretical knowledge, introduce essential experimental techniques, and cultivate analytical thinking, ensuring students acquire the chemical literacy necessary for interdisciplinary scientific study.

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) The difference between a student's experimental measurement of the density of sodium chloride and the known density of this compound reflects the ___________ of the student's result.

A) accuracy

B) precision

C) random error

D) systematic error

E) indeterminate error

Answer: A

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 number 6.0448, rounded to 2 decimal places, becomes 6.05.

A)True

B)False

Answer: False

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Chapter 2: The Components of Matter

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Sample Questions

Q1) Iron (III) chloride hexahydrate is used as a coagulant for sewage and industrial wastes. What is its formula?

A) Fe(Cl.6H₂O)₃

B) Fe₃Cl.6H₂O

C) FeCl₃(H₂O)₆

D) Fe₃Cl(H₂O)₆

E) FeCl₃.6H₂O

Answer: E

Q2) Give the common name of the group in the periodic table to which each of the following elements belongs:

a. Rb

b. Br

c. Ba

d. Ar

Answer: a. alkali metals

b. halogens

c. alkaline earth metals

d. noble gases

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Balance the following equation:

Ca₃(PO₄)₂(s) + SiO₂(s) + C(s) \(\to\)CaSiO₃(s) + CO(g) + P₄(s)

A) Ca₃(PO₄)₂(s) + 3SiO₂(s) + 8C(s) \(\to\) 3CaSiO₃(s) + 8CO(g) + P₄(s)

B) Ca₃(PO₄)₂(s) + 3SiO₂(s) + 14C(s) \(\to\) 3CaSiO₃(s) + 14CO(g) + P₄(s)

C) Ca₃(PO₄)₂(s) + 3SiO₂(s) + 8C(s) \(\to\) 3CaSiO₃(s) + 8CO(g) + 2P₄(s)

D) 2Ca₃(PO₄)₂(s) + 6SiO₂(s) + 10C(s) \(\to\) 6CaSiO₃(s) + 10CO(g) + P₄(s)

E) 2Ca₃(PO₄)₂(s) + 6SiO₂(s) + 10C(s) \(\to\) 6CaSiO₃(s) + 10CO(g) + 4P₄(s)

Answer: D

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Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) Select the correct name and chemical formula for the precipitate that forms when the following reactants are mixed. CoSO₄(aq) + (NH₄)₃PO₄(aq) \(\to\)

A) cobalt(II) phosphate, Co₃(PO₄)₂

B) cobalt(III) phosphate, Co₃(PO₄)₂

C) cobalt(II) phosphate, CoPO₄

D) cobalt(III) phosphate, CoPO₄

E) ammonium sulfate, (NH₄)₂SO₄

Q2) A 0.00100 mol sample of Ca(OH)₂ requires 25.00 mL of aqueous HCl for neutralization according to the reaction below. What is the concentration of the HCl? Equation: Ca(OH)₂(s) + 2HCl(aq) \(\to\) CaCl₂(aq) + H₂O(l)

A) 0.0200 M

B) 0.0400 M

C) 0.0800 M

D) 4.00 × 10^-5 M

E) none of the above

Q3) The combustion of an element is always a combination reaction. A)True

B)False

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Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) The temperature of the carbon dioxide atmosphere near the surface of Venus is 475°C. Calculate the average kinetic energy per mole of carbon dioxide molecules on Venus.

A) 2520 J/mol

B) 4150 J/mol

C) 5920 J/mol

D) 9330 J/mol

E) 5920 kJ/mol

Q2) "The pressure of an ideal gas is inversely proportional to its volume at constant temperature and number of moles" is a statement of __________________ Law.

A) Charles's

B) Boyle's

C) Amontons's

D) Avogadro's

E) Gay-Lussac's

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

Q4) State Boyle's Law and illustrate it with a graph, using standard x-y coordinate axes. Be sure to label the axes unambiguously with the correct gas variables.

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Page 7

Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Use Hess's Law to calculate the enthalpy change for the reaction WO₃(s) + 3H₂(g) \(\to\) W(s) + 3H₂O(g)

From the following data: 2W(s) + 3O₂(g) \(\to\) 2WO₃(s) \(\Delta\)H = -1685.4 kJ 2H₂(g) + O₂(g) \(\to\) 2H₂O(g) \(\Delta\)H = -477.84 kJ

A) 125.9 kJ

B) 252.9 kJ

C) 364.9 kJ

D) 1207.6 kJ

E) none of the above

Q2) The only way in which a system can do work on the surroundings is by expansion against the external pressure.

A)True

B)False

Q3) Although internal energy (E) is more fundamental and conceptually easier than enthalpy (H), in most chemical applications \(\Delta\)H is more relevant and useful than \(\Delta\)E. Why?

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Chapter 7: Quantum Theory and Atomic Structure

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Sample Questions

Q1) The AM station KBOR plays your favorite music from the 20's and 30's at 1290 kHz. Find the wavelength of these waves.

A) 4.30 × 10^-2m

B) 0.144 m

C) 6.94 m

D) 232 m

E) > 10³ m

Q2) Electromagnetic radiation of 500 nm wavelength lies in the __________ region of the spectrum.

A) infrared

B) visible

C) ultraviolet

D) X-ray

E) (\(\gamma\))-ray

Q3) In the quantum mechanical treatment of the hydrogen atom, the probability of finding an electron at any point is proportional to the wave function \(\varPsi\) .

A)True

B)False

Q4) 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) 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 the above statements is generally correct.

Q2) The maximum number of electrons in an atom with the same value of n is 2n².

A)True

B)False

Q3) In neutral atoms, the 3d orbitals have higher energy than the 4s orbitals. A)True

B)False

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) Which one of the following properties is least characteristic of typical ionic compounds?

A) high melting point

B) high boiling point

C) brittleness

D) poor electrical conductor when solid

E) poor electrical conductor when molten

Q2) Which of the following compounds displays the greatest ionic character in its bonds?

A) NO₂

B) CO₂

C) H₂O

D) HF

E) NH₃

Q3) Select the most polar bond amongst the following.

A) C-O

B) Si-F

C) Cl-F

D) C-F

E) C-I

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Draw Lewis structures, showing all valence electrons, for the following species:

a. S^2-

b. CO

c. SO₂

d. CH₃OH

Q2) Which one of the following molecules and ions will have a planar geometry?

A) PCl₃

B) BF₄^-

C) XeF₄

D) BrF₅

E) H₃O⁺

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) List the three important ways in which molecules can violate the octet rule, and in each case draw one Lewis structure of your choice as an example.

Page 12

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Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) Select the correct statement about \(\pi\)-bonds in valence bond theory.

A) A \(\pi\) bond is stronger than a sigma bond.

B) A \(\pi\) bond can hold 4 electrons, two above and two below the \(\pi\)-bond axis.

C) A carbon-carbon double bond consists of two \(\pi\) bonds.

D) A \(\pi\) bond is the same strength as a \(\sigma\) bond.

E) A \(\pi\) bond between two carbon atoms restricts rotation about the C-C axis.

Q2) Valence bond theory predicts that bromine will use _____ hybrid orbitals in BrF₅.

A) sp²

B) sp³

C) sp³d

D) sp³d ²

E) none of the above

Q3) Atoms of period 3 and beyond can undergo sp³d ² hybridization, but atoms of period 2 cannot.

A)True

B)False

Q4) In one sentence state how molecular orbitals are usually obtained.

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Page 13

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

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Sample Questions

Q1) Use molecular orbital band diagrams to explain why metals are good conductors but semiconductors are not.

Q2) 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.

Q3) What types of forces exist between molecules of CO₂?

A) hydrogen bonding only.

B) hydrogen bonding and dispersion forces.

C) dipole-dipole forces only.

D) dipole-dipole and dispersion forces.

E) dispersion forces only.

Q4) Iron has a body-centered cubic unit cell, and a density of 7.87 g/cm³.

Calculate the edge length of the unit cell, in pm. (The atomic mass of iron is 55.85 amu. Also, 1 amu = 1.661 × 10^-24 g.)

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Q5) How do the electrical properties of semiconductors differ from those of metals?

Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) If the shell of a raw egg is carefully dissolved away, and the egg in its flexible membrane is then placed in distilled water, the egg's volume will expand. Explain.

Q2) The vapor pressure of pure acetone (propanone) is 266 torr. When a non-volatile solute is added, the vapor pressure of acetone above the solution falls to 232 torr. What is the mole fraction of acetone in the solution?

A) 0.87

B) 0.69

C) 0.32

D) 0.13

E) 0.045

Q3) Which of the following pairs of ions is arranged so that the ion with the larger (i.e., more negative) heat of hydration is listed first?

A) Br^-, K⁺

B) Mg²⁺, Sr ²⁺

C) Ca²⁺, Sc³⁺

D) Na⁺, Li⁺

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

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) Which of the following will have the highest boiling point?

A) O₂

B) Cl₂

C) Br₂

D) I₂

E) Xe

Q2) Which of the following formulas does not represent a stable compound?

A) N₂O

B) NO

C) NO₂

D) NO₃

E) N₂O₄

Q3) Write balanced equations, showing all reactants and products, to represent

a. the reaction of calcium metal with water.

b. the reaction of aluminum metal with oxygen gas.

Q4) Carbon monoxide's toxicity is related to its ability to bond to iron in hemoglobin.

A)True

B)False

Page 16

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Chapter 15: Organic Compounds and the Atomic Properties of Carbon

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Sample Questions

Q1) All ketone molecules are capable of hydrogen bonding to other ketone molecules. A)True

B)False

Q2) a. Draw two different structures with the molecular formula C₂H₆O.

b. Name the functional group in each structure.

c. Which one will have the higher boiling point, and why?

Q3) Nylon-type compounds are prepared by the reaction of A) a diol with a diacid.

B) an alcohol with a diacid.

C) a diamine with a diacid.

D) an amine with a diacid.

E) a diamine with an acid.

Q4) Carboxylic acids are weak acids. A)True

B)False

Q5) The carbon atoms in a molecule of cyclohexane lie in the same plane.

A)True

B)False

Q6) Explain how a disulfide bridge can arise a protein molecule. Page 17

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Page 18

Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Q1) A catalyst accelerates a reaction because

A) it increases the number of molecules with energy equal to or greater than the activation energy.

B) it lowers the activation energy for the reaction.

C) it increases the number of collisions between molecules.

D) it increases the temperature of the molecules in the reaction. E) it supplies energy to reactant molecules.

Q2) Sulfur trioxide can undergo decomposition according to the equation 2SO₃ \(\to\) 2SO₂ + O₂

For this reaction, rate = -0 0.5\(\Delta\)[SO₃]/\(\Delta\)t = k[SO₃]². If the reaction rate is 1.75 × 10^-7 mol L^-1 min^-1 when the concentration of sulfur trioxide is 5.4 × 10^-3 mol L^-1, what is the value of the rate constant k?

A) 3.2 × 10^-5 L mol^-1 min^-1

B) 1.6 × 10^-5 L mol^-1 min^-1

C) 6.0 × 10^-3 L mol^-1 min^-1

D) 3.0 × 10^-3 L mol^-1 min^-1

E) 1.6 × 10^-2 L mol^-1 min^-1

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Q1) If all of the coefficients in the balanced equation for an equilibrium reaction are doubled, then the value of the equilibrium constant, K_c, will also be doubled.

A)True

B)False

Q2) An equilibrium is established in which both the forward (fwd) and the reverse (rev) reactions are elementary. If the equilibrium constant K_c = 1.6 × 10^-2 and the rate constant k_fwd = 8.0 × 10^-7s^-1 what is the value of k_rev?

A) 1.3 × 10^-8 s^-1

B) 7.8 × 10⁷ s^-1

C) 2 × 10⁴ s^-1

D) 5.0 × 10^-5 s^-1

E) none of the above

Q3) Although a system may be at equilibrium, the rate constants of the forward and reverse reactions will in general be different.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) 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) ICl₃

C) Cl^-

D) the solvent

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

Q2) The substance HClO₄is considered

A) a weak acid.

B) a weak base.

C) a strong acid.

D) a strong base.

E) a neutral compound.

Q3) What is the pH of a 0.050 M LiOH solution?

A) < 1.0

B) 1.30

C) 3.00

D) 11.00

E) 12.70

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Page 21

Chapter 19: Ionic Equilibria in Aqueous Systems

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Q1) A buffer is prepared by adding 1.00 L of 1.0 M HCl to 750 mL of 1.5 M NaHCOO. What is the pH of this buffer? K_a = 1.7 × 10^-4

A) 2.87

B) 3.72

C) 3.82

D) 3.95

E) 4.66

Q2) The indicator propyl red has K_a = 3.3 × 10^-6. What would be the approximate pH range over which it would change color?

A) 3.5-5.5

B) 4.5-6.5

C) 5.5-7.5

D) 6.5-8.5

E) none of the above

Q3) Calculate the solubility of copper(II) carbonate, CuCO₃, in 1.00 mol L^-1 NH₃. K_sp = 3.0 × 10^-12 for CuCO₃, K_f = 5.6 × 10¹¹ for Cu(NH₃)₄²⁺

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22

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

of Chemical Reactions

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Q1) State the second and third laws of thermodynamics.

Q2) For a chemical reaction to be spontaneous only at high temperatures, which of the following conditions must be met?

A) .\(\Delta\)S° > 0, \(\Delta\)H° > 0

B) .\(\Delta\)S° > 0, \(\Delta\)H° < 0

C) .\(\Delta\)S° < 0, \(\Delta\)H° < 0

D) .\(\Delta\)S° < 0, \(\Delta\)H° > 0

E) .\(\Delta\)G° > 0

Q3) 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°

Q4) The term microstate refers to the energy state of a single molecule in a system of many molecules.

A)True B)False

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) A current of 1000. A flows for exactly 1 hour, through a cell in which the following reaction occurs at one of the electrodes.

Mg²⁺ + 2e^- \(\to\) Mg

a. Calculate the charge, in coulombs, which passes through the circuit in this time. b. Calculate the theoretical mass of Mg (magnesium metal) which is produced in this time.

Q2) A concentration cell consists of two Zn/Zn²⁺ electrodes. The electrolyte in compartment A is 0.10 M Zn(NO₃)₂and in compartment B is 0.60 M Zn(NO₃)₂. What is the voltage of the cell at 25°C?

A) 0.010 V

B) 0.020 V

C) 0.023 V

D) 0.046 V

E) none of the above

Q3) Write down equations representing the anode half-reaction, the cathode half-reaction, and the overall cell reaction for the lead-acid battery.

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Page 24

Chapter 22: The Elements in Nature and Industry

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

Q2) The bond energy in the hydrogen molecule (H₂) is greater than that of the tritium molecule (T₂).

A)True

B)False

Q3) The Hall-Heroult process refers to

A) the production of aluminum by electrolysis.

B) the recovery of sulfur from underground deposits.

C) the manufacture of sulfuric acid.

D) the production of ammonia from nitrogen and hydrogen gases.

E) the isolation of Al₂O₃ from bauxite.

Q4) The most common source for commercial production of aluminum is called A) aluminite.

B) hematite.

C) galena.

D) cinnabar.

E) bauxite.

Q5) Sulfide ores are frequently treated by flotation in order to concentrate them. A)True

B)False

Page 25

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) Octahedral complexes can exhibit geometric, optical, and linkage isomerism.

A)True

B)False

Q2) All atoms of the first transition series of elements have the ground state electronic configuration [Ar]4s²3d^x, where x is an integer from 1 to 10.

A)True

B)False

Q3) In complexes of transition metals, the maximum coordination number of the metal is equal to its number of d electrons.

A)True

B)False

Q4) Which of the following will be paramagnetic?

A) V⁵⁺

B) Ni²⁺

C) Mn⁷⁺

D) Ti⁴⁺

E) Zn

Q5) What is the difference between a coordination compound and a complex ion?

Page 26

Q6) What geometry is particularly common for complexes of d¹⁰ metal ions?

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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) Which of the following types of radioactive decay does not produce new element?

A) gamma emission

B) electron capture

C) beta emission

D) alpha emission

E) double beta emission

Q3) Most foodstuffs contain natural, radioactive isotopes.

A)True

B)False

Q4) Radioactive decay follows zero-order kinetics.

A)True

B)False

Page 28

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