Chemistry for Engineers Exam Practice Tests - 2734 Verified Questions

Chemistry for Engineers Exam Practice Tests

Course Introduction

This course introduces fundamental concepts of chemistry tailored for engineering students, providing a strong foundation in chemical principles relevant to engineering applications. Topics include atomic and molecular structure, stoichiometry, thermodynamics, chemical kinetics, electrochemistry, materials science, and the properties of solids, liquids, and gases. Emphasis is placed on the practical use of chemistry in solving engineering problems, understanding material behavior, and optimizing processes in industrial and technological settings. Through lectures, laboratory experiments, and problem-solving sessions, students will develop the skills necessary to apply chemical knowledge in various engineering fields.

Recommended Textbook Chemistry 10th Edition by Raymond Chang

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25 Chapters

2734 Verified Questions

2734 Flashcards

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Chapter 1: Chemistry: The Study of Change

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Sample Questions

Q1) Identify the following as a physical or chemical change: Ripening of fruit. Answer: Chemical

Q2) An archeologist finds a huge monolith in the desert. In order to estimate the weight of this object; he estimates the dimensions of the monolith and removes some chips from the rock with his hammer, collecting the following data: dimensions of the monolith = 1.5 m * 5.2 m * 13 m mass of rock chips = 41.73 g volume of rock chips = 15.2 cm³ Determine the mass of the monolith in pounds, assuming it is of uniform composition.(1 lb = 453.6 g) Answer: 6.1 * 10⁵ lb

Q3) An automobile engine has a piston displacement of 1,600 cm³. Express this volume in liters.

Answer: 1.6 L

Q4) The density of lead is 11.4 g/cm³. Express this density in pounds per cubic foot.

Answer: 711 lbs/ft³

Q5) Classify the following as a pure substance or a mixture: Seven-Up®. Answer: Mixture

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Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) The number of neutrons in all atoms of an element is the same.

A)True

B)False

Answer: False

Q2) A phosphide ion has:

A)10 protons and 13 electrons

B)12 protons and 15 electrons

C)15 protons and 15 electrons

D)15 protons and 18 electrons

E)18 protons and 21 electrons

Answer: D

Q3) The correct name for Ba(OH)₂ is

A)barium hydrogen oxide.

B)boron hydroxide.

C)barium hydrate.

D)beryllium hydroxide.

E)barium hydroxide.

Answer: E

Q4) Name the following binary compound: AgCl.

Answer: silver chloride; may accept silver(I)chloride.

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Chapter 3: Mass Relationships in Chemical Reactions

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Sample Questions

Q1) Calculate the molecular mass of ethylene glycol, C₂H₆O₂, a compound frequently used as automobile antifreeze.

Answer: 62.1 g

Q2) What is the mass of 7.80 * 10¹⁸ carbon atoms?

A)1.30 * 10^-5 g

B)6.43 * 10³ g

C)7.80 * 10¹⁸ g

D)1.56 * 10^-4 g

E)12.01 g

Answer: D

Q3) How many moles of Cl atoms are there in 65.2 g CHCl₃?

A)0.548 mol

B)1.09 mol

C)3.3 * 10²³ mol

D)1.64 mol

E)3.0 mol

Answer: D

Q4) How many Mg atoms are present in 170 g of Mg?

Answer: 4.2 * 10²⁴

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Chapter 4: Reactions in Aqueous Solution

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Sample Questions

Q1) A piece of copper metal was added to an aqueous solution of silver nitrate, and within a few minutes it was observed that a grey crystalline solid formed on surface of the copper and the solution turned a blue color characteristic of copper(II)ions. Write the net ionic equation for this reaction.

Q2) When 38.0 mL of 0.1250 M H₂SO₄ is added to 100.mL of a solution of PbI₂, a precipitate of PbSO₄ forms. The PbSO₄ is then filtered from the solution, dried, and weighed. If the recovered PbSO₄ is found to have a mass of 0.0471 g, what was the concentration of iodide ions in the original solution?

A)3.10 * 10^-4 M

B)1.55 * 10^-4 M

C)6.20 * 10^-3 M

D)3.11 * 10^-3 M

E)1.55 * 10^-3 M

Q3) During a titration the following data were collected.A 50.0 mL portion of an HCl solution was titrated with 0.500 M NaOH; 200.mL of the base was required to neutralize the sample.How many grams of HCl are present in 500.mL of this acid solution?

Q4) Give an example of a diprotic acid.

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Page 6

Chapter 5: Gases

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Sample Questions

Q1) Give five examples of compounds that exist as gases at room temperature and pressure.

Q2) What is standard temperature and standard pressure?

Q3) At constant pressure, the density of a gas depends on temperature.Does the density increase or decrease as the temperature increases?

Q4) Which of these properties is/are characteristic(s)of gases?

A)High compressibility

B)Relatively large distances between molecules

C)Formation of homogeneous mixtures regardless of the nature of gases

D)A and B.

E)A, B, and C.

Q5) A sample of carbon monoxide gas was collected in a 2.0 L flask by displacing water at 28°C and 810 mmHg. Calculate the number of CO molecules in the flask. The vapor pressure of water at 28°C is 28.3 mmHg.

A)5.0 * 10²²

B)5.2 * 10²²

C)3.8 * 10²³

D)5.4 * 10²³

E)3.8 * 10²⁵

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Chapter 6: Thermo-Chemistry

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Sample Questions

Q1) Calculate the enthalpy of reaction for H₂(g)+ C₂H₄(g)\(\rarr\) C₂H₆(g). [\(\Delta\)H°_f(C₂H₄(g))= 52.3 kJ/mol; \(\Delta\)H°_f(C₂H₆(g))= -84.7 kJ/mol]

Q2) Chemical energy is

A)the energy stored within the structural units of chemical substances. B)the energy associated with the random motion of atoms and molecules. C)solar energy, i.e.energy that comes from the sun. D)energy available by virtue of an object's position.

Q3) Find the standard enthalpy of formation of ethylene, C₂H₄(g), given the following data: heat of combustion of C₂H₄(g)= -1411 kJ/mol; \9\Delta\)H°_f[CO₂(g)] = -393.5 kJ/mol; \(\Delta\)H°_f[H₂O(l)] = -285.8 kJ/mol.

A)52 kJ/mol

B)87 kJ/mol

C)731 kJ/mol

D)1.41 * 10³ kJ/mol

E)2.77 * 10³ kJ/mol

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Page 8

Chapter 7: Quantum Theory and the Electronic Structure of Atoms

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Sample Questions

Q1) What is the wavelength, in meters, of an alpha particle with a kinetic energy of 8.0 * 10^-13 J.(mass of an alpha particle = 4.00150 amu; 1 amu = 1.67 * 10^-27 kg)

Q2) Which element has the following ground-state electron configuration? [Kr]5s¹4d⁵

A)Mn

B)Mo

C)Nb

D)Re

E)Tc

Q3) What is the total number of electrons possible in the 2p orbitals?

Q4) According to de Broglie's equation, the wavelength associated with the motion of a particle increases as the particle mass decreases.

A)True

B)False

Q5) What is the electron configuration of calcium?

Q6) If one electron is added to the outer shell of chlorine, it would have the same electron configuration as what element?

Page 9

Q7) Write the ground state electron configuration for Ni.

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Chapter 8: Periodic Relationships Among the Elements

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Sample Questions

Q1) Consider the element with the electron configuration [Kr]5s²4d⁷. This element is

A)a representative element.

B)a transition metal.

C)a nonmetal.

D)an actinide element.

E)a noble gas.

Q2) Which of the following is the electron configuration for the aluminum ion?

A)1s²2s²2p⁶3s²

B)1s²2s²2p⁶3s²3p²

C)1s²2s²2p⁶3s²3p¹ D)1s²2s²2p⁶

E)1s²2s²2p⁶3s²3p⁴

Q3) Which of the following elements has the greatest metallic character?

A)Br

B)Se

C)Ni

D)As

E)Si

Q4) Write the ground-state electron configuration for Ni²⁺.

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Chapter 9: Chemical Bonding I: Basic Concepts

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Sample Questions

Q1) The standard enthalpy of formation of ammonia at 25°C is -46.3 kJ/mol. Estimate the N-H bond enthalpy at this temperature. (Given: BE(N\(\equiv\)N)=941.4 kJ/mol, BE(H-H)= 436.4 kJ/mol)

A)383 kJ/mol

B)475 kJ/mol

C)360 kJ/mol

D)391 kJ/mol

E)459 kJ/mol

Q2) Which one of the following is most likely to be a covalent compound?

A)KF

B)CaCl₂

C)SF₄

D)Al₂O₃

E)CaSO₄

Q3) The electron dot structure for AsCl₃ shows

A)a total of 84 electron dots

B)three single bonds and 10 lone pairs

C)two single bonds, one double bond, and 9 lone pairs

D)one single bond, two double bonds, and 8 lone pairs

E)three single bonds and one lone pair

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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization of Atomic Orbitals

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Sample Questions

Q1) More energy is required to break a bond with an order of 3/2 than is required to break a bond of order 2.

A)True

B)False

Q2) Give the number of lone pairs around the central atom and the geometry of the ion NO₂^-.

A)0 lone pairs, linear

B)1 lone pair, bent

C)2 lone pair, bent

D)3 lone pairs, bent

E)3 lone pairs, linear

Q3) Give the number of lone pairs around the central atom and the geometry of the ion SeO₄^2-.

A)0 lone pairs, square planar

B)0 lone pairs, tetrahedral

C)1 lone pair, distorted tetrahedron (seesaw)

D)1 lone pair, square pyramidal

E)2 lone pairs, square planar

Q4) Which should have the longer bond, B₂ or B₂^-?

Page 12

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Chapter 11: Intermolecular Forces and Liquids and Solids

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Sample Questions

Q1) Indicate all the types of intermolecular forces of attraction in CHCl₃(l).

Q2) Which of the following is not an endothermic process?

A)melting of a solid

B)vaporization

C)raising the temperature of a gas

D)condensation of water vapor

E)sublimation of dry ice

Q3) Butter melts over a range of temperature, rather than with a sharp melting point. Butter is classified as a/an

A)metallic crystal.

B)covalent solid.

C)molecular crystal.

D)amorphous solid.

E)ionic crystal.

Q4) Of the pair of compounds given, which would have the stronger intermolecular forces of attraction?

SF₄ or C₁₀H₂₂

Q5) Indicate all the types of intermolecular forces of attraction in SO₂(l).

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Chapter 12: Physical Properties of Solutions

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Sample Questions

Q1) Explain the following, on the basis of osmosis or osmotic pressure: Meat that is salted before cooking tends to dry out.

Q2) Calculate the mass of solute in the following solution: 50.0 mL of 0.0300 M C₁₂H₂₂O₁₁.

Q3) How many grams of propanol (C₃H₇OH, 60.10 g/mol)would be needed to make 750 mL of a solution with an osmotic pressure of 25 atm at 25°C? (R = 0.0821 L.atm/K.mol)

Q4) What is the osmotic pressure of a solution prepared from 13.7 g of the electrolyte HCl and enough water to make 0.500 L of solution at 18°C?

A)0.55 atm

B)1.10 atm

C)8.95 atm

D)17.9 atm

E)35.9 atm

Q5) For dilute aqueous solutions, the concentration units molarity and molality have almost the same values.

A)True

B)False

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) Solids cannot react with gases.

A)1 and 2

B)1 and 3

C)1 and 4

D)2 and 3

E)3 and 4

Q2) For the reaction represented below, the experimental rate law is given by rate = k [(CH₃)₃CCl].

(CH₃)₃CCl(aq)+ OH^-F1F1F1S1 F1F1F10 (CH₃)₃COH(aq)+ Cl^- If some solid sodium hydroxide were added to a solution in which [(CH₃)₃CCl] = 0.01 M and [NaOH] = 0.10 M, which of the following would be true? (Assume the temperature and volume remain constant.)

A)Both the reaction rate and k would increase.

B)Both the reaction rate and k would decrease.

C)Both the reaction rate and k would remain the same.

D)The reaction rate would increase but k would remain the same.

E)The reaction rate would decrease but k would remain the same.

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Describe why addition of a catalyst does not affect the equilibrium constant for a reaction.

Q2) What conditions are used in the Haber process to enhance the yield of ammonia? Explain why each condition affects the yield in terms of the Le Châtelier principle.

Q3) Which of these statements is true about chemical equilibria in general?

A)At equilibrium the total concentration of products equals the total concentration of reactants, that is, [products] = [reactants].

B)Equilibrium is the result of the cessation of all chemical change.

C)There is only one set of equilibrium concentrations that equals the K_c value.

D)At equilibrium, the rate constant of the forward reaction is equal to the rate constant for the reverse reaction.

E)At equilibrium, the rate of the forward reaction is equal to as the rate of the reverse reaction.

Q4) The dissociation of solid silver chloride in water to produce silver ions and chloride ions has an equilibrium constant of 1.8 * 10^-18.Based on the magnitude of the equilibrium constant, is silver chloride very soluble in water? Why?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) A 0.10 M HF solution is 8.4% ionized. Calculate the H⁺ ion concentration.

A)0.84 M

B)0.12 M

C)0.10 M

D)0.084 M

E)8.4 * 10^-3 M

Q2) Arrange the acids HBr, H₂Se, and H₃As in order of increasing acid strength.

A)HBr < H₂Se < H₃As

B)HBr < H₃As < H₂Se

C)H₂Se < H₃As < HBr

D)H₃As< H₂Se < HBr

E)H₃As< HBr < H₂Se

Q3) Calculate the pH of a solution containing 0.20 g of NaOH in 2,000.mL of solution.

Q4) An unknown substance was added to a solution and the pH decreases. What type of substance was added?

Q5) Will a 0.1 M solution of NH₄CN(aq)be acidic, basic, or neutral?

Q6) Calculate the H⁺ ion concentration in a solution with a pH of 3.85.

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Chapter 16: Acid-Base Equilibria and Solubility Equilibria

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Sample Questions

Q1) An environmental chemist obtained a 200.mL sample of lake water believed to be contaminated with a single monoprotic strong acid.Titrating this sample with a 0.0050 M NaOH(aq)required 7.3 mL of the NaOH solution to reach the endpoint.If the size of the lake can be approximated as 1.1 km long by 2.3 km wide, and has an average depth of 10.m, estimate how many moles of the strong acid are present in the lake?

Q2) At 25 °C, the base ionization constant for NH₃ is 1.8 * 10^-5.

Determine the pH of a solution prepared by adding 0.0500 mol of solid ammonium chloride to 100.mL of 0.150 M ammonia.

Q3) The molar solubility of tin(II)iodide is 1.28 * 10^-2 mol/L. What is K_sp for this compound?

A)8.4 * 10^-6

B)1.28 * 10^-2

C)4.2 * 10^-6

D)1.6 * 10^-4

E)2.1 * 10^-6

Q4) The K_sp of CaF₂ is 4 * 10^-11. What is the maximum concentration of Ca²⁺ possible in a 0.10 M NaF solution?

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Chapter 17: Chemistry in the Atmosphere

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Sample Questions

Q1) Which one of the following reactions is an example of nitrogen fixation?

A)N₂O₅(g)\(\to\) NO₃(g)+ NO₂(g)

B)N₂(g)+ O₂(g)\(\to\) 2NO(g)

C)3NO₂(g)+ H₂O(l)\(\to\) 2HNO₃(aq)+ NO(g)

D)2NO(g)+ O₂(g)\(\to\) 2NO₂(g)

E)2NH₃(g)\(\to\) 3H₂(aq)+ N₂(g)

Q2) The compound CFCl₃ is used as a/an

A)enzyme

B)anesthetic

C)gaseous fuel

D)coolant

E)CFC replacement

Q3) Write out the steps that show how sulfur in coal is converted to sulfuric acid in acid rain.

Q4) Write out the steps in the mechanism of ozone destruction by chlorine atoms.

Q5) Name four regions of the atmosphere in order of increasing altitude.

Q6) Write chemical equations that show what happens when acid rain reacts with limestone rock

Q7) What happens to nitrogen dioxide in sunlight? Write the reaction for this process.

Page 19

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Chapter 18: Entropy, Free Energy, and Equilibrium

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Sample Questions

Q1) For the reaction H₂(g)+ S(s)\( \to\) H₂S(g), \(\Delta\)H° = -20.2 kJ/mol and \(\Delta\)S° = +43.1 J/K·mol.Which of these statements is true?

A)The reaction is only spontaneous at low temperatures.

B)The reaction is spontaneous at all temperatures.

C)(\(\Delta\)G° becomes less favorable as temperature increases.)

D)The reaction is spontaneous only at high temperatures.

E)The reaction is at equilibrium at 25°C under standard conditions.

Q2) Consider the gas phase reaction shown below.

CH₄(g)+ CF₄(g)\(\to\) CH₃F(g)+ CHF₃(g)

Will the standard entropy increase, decrease, or remain if the reaction proceeds from reactants to products, as written? Explain your reasoning.

Q3) The heat of vaporization of water is 2.27 kJ/g. What is \(\Delta\)S_vap per mole at the normal boiling point?

Q4) Choose the substance with the higher entropy per mole at a given temperature: 1 mole of N₂(g)in a 22.4 L container or 1 mole of N₂(g)in a 2.24 L container.

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Page 20

Chapter 19: Electrochemistry

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Sample Questions

Q1) Which one of the following reagents is capable of transforming Br(aq)to Br₂(l)under standard-state conditions?

A)I^-(aq)

B)NO₃^-(aq)

C)Ag⁺(aq)

D)Al³⁺(aq)

E)Au³⁺(aq)

Q2) The half-reaction that occurs at the cathode during electrolysis of an aqueous sodium iodide solution is

A)Na⁺ + e^- \(\to\)Na.

B)Na \(\to\)Na⁺ + e^-.

C)2H₂O + 2e^- \(\to\) H₂ + 2OH^-.

D)I₂ + 2e^- \(\to\) 2I^-.

E)2I^- \(\to\) I₂ + 2e^-.

Q3) Complete and balance the following redox reaction under acidic conditions: Cu(s)+ NO₃^-(aq)\(\to\) Cu²⁺(aq)+ NO₂(g)

Q4) How many moles of silver metal are produced in 1.2 hours by the electrolysis of AgNO₃(aq)using a current of 6.0 A?

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Chapter 20: Metallurgy and the Chemistry of Metals

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Sample Questions

Q1) The following reaction is used to produce titanium metal at high temperature. TiCl₄(g)+ 2Mg(l)\(\to\) Ti(s)+ 2MgCl₂(l)

Which element is oxidized and which is reduced?

Q2) The Hall process involves the reduction of Al₂O₃ to aluminum by A)carbon (coke).

B)carbon monoxide.

C)molecular hydrogen.

D)sodium.

E)electrolysis.

Q3) Which one of these metals would normally be obtained by electrolytic reduction?

A)aluminum

B)chromium

C)copper

D)iron

E)zinc

Q4) Mercury, magnesium, and zinc have low enough boiling points that they can be purified by distillation.

A)True

B)False

Page 22

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Chapter 21: Nonmetallic Elements and Their Compounds

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Sample Questions

Q1) Which one of these elements is not a gas at room temperature?

A)chlorine

B)fluorine

C)nitrogen

D)oxygen

E)sulfur

Q2) Which of these elements occupies a position on the periodic table that is not entirely consistent with its chemical properties?

A)H

B)He

C)Ar

D)Al

E)U

Q3) Write the chemical formula of the oxide ion.

Q4) The main commercial source of bromine is

A)silver bromide ore.

B)seawater.

C)the Frasch process.

D)phosphate rock.

E)the Ostwald process.

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Chapter 22: Transition Metal Chemistry and Coordination Compounds

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Sample Questions

Q1) Write the chemical formula of diamminedichloroplatinum(II).

Q2) In K₄[Fe(CN)₆], how many 3d electrons does the iron atom have?

A)3

B)4

C)5 D)6

E)7

Q3) How many 3d electrons does a Mn²⁺ ion have? A)1 B)2 C)3 D)4 E)5

Q4) The total number of electrons in the 3d orbitals of Co³⁺ is A)4

B)5 C)6

D)7

E)10

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Chapter 23: Nuclear Chemistry

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Sample Questions

Q1) Which of the following is an advantage of nuclear power plants over coal-burning plants? Nuclear power plants

A)produce radioactive byproducts with very short half-lives, reducing the need for waste storage.

B)do not pollute the air with SO₂, soot, and fly-ash.

C)create no thermal pollution.

D)generate radioactive byproducts that can be sold for use in secondary applications.

Q2) Which type of nuclear process requires an extremely high temperature (millions of degrees)?

A)beta decay

B)fission reaction

C)fusion reaction

D)alpha decay

E)positron emission

Q3) Describe the role of a moderator in a fission reactor, and give examples of two substances that are used as moderators.

Q4) Protactinium-234 has a half-life of 1 minute.How much of a 400.g sample protactinium would remain after 2 minutes?

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Page 25

Chapter 24: Organic Chemistry

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Sample Questions

Q1) Which type of organic compound does not contain a carbonyl group?

A)ethers

B)carboxylic acids

C)ketones

D)aldehydes

E)esters

Q2) Which is the product of the reaction of one mole of HCl with one mole of 1-butyne?

A)1-chloro-1-butene

B)1-chloro-2-butene

C)2-chloro-1-butene

D)ethyl chloride + acetylene

Q3) The group of atoms that is responsible for the characteristic properties of a family of organic compounds is called a/an

A)reaction center.

B)functional group.

C)binding site.

D)enzyme.

E)polyatomic ion.

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Chapter 25: Synthetic and Natural Organic Polymers

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Sample Questions

Q1) Both DNA and RNA have double-helical structures.

A)True

B)False

Q2) Which nitrogen base is found in RNA but not in DNA?

A)adenine

B)cytosine

C)guanine

D)thymine

E)uracil

Q3) The only true hydrocarbon polymer found in nature is

A)rubber

B)nylon

C)Tyvek

D)polystyrene

E)neoprene

Q4) The monomer used to prepare polyvinyl chloride (PVC)is CHCl=CHCl.

A)True

B)False

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