Chemical Science Textbook Exam Questions - 3115 Verified Questions

Chemical Science

Textbook Exam Questions

Course Introduction

Chemical Science explores the fundamental principles and applications of chemistry, focusing on the molecular composition, structure, properties, and transformations of matter. This course covers topics such as atomic theory, chemical bonding, thermodynamics, kinetics, equilibria, and organic and inorganic chemistry. Through lectures, laboratory experiments, and real-world examples, students gain a solid understanding of chemical processes and their relevance to industries, medicine, the environment, and daily life. The course also emphasizes the development of analytical and problem-solving skills essential for further studies and careers in chemistry and related fields.

Recommended Textbook Chemistry 11th Edition by Raymond Chang

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Chapter 1: Chemistry: the Study of Change

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Sample Questions

Q1) Classify the following as a mixture, a compound, or an element: Brewed coffee, ready to drink.

Answer: Mixture

Q2) The diameter of Earth is 12.7 Mm.Express this diameter in centimeters.

A)1.27 × 10⁵ cm

B)1.27 × 10⁶ cm

C)1.27 × 10⁷ cm

D)1.27 × 10⁸ cm

E)1.27 × 10⁹ cm

Answer: E

Q3) Give the correct number of significant figures and units to the problem below 5.67

m/ 2.04 m² =

A)2.78 m³

B)2.78 m²

C)2.78 m

D)2.78

E)2.78 m^-1

Answer: E

Q4) Define pure substance.

Answer: Matter that has a definite composition

Page 3

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Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) Define allotrope.

Answer: An allotrope is one of the two or more distinct forms of an element.

Q2) Give the number of protons (p), neutrons (n), and electrons (e)in one atom of ²³⁸U.

A)146 p, 92 n, 92 e

B)92 p, 92 n, 92 e

C)92 p, 146 n, 92e

D)146 p, 28 n, 146 e

E)238 p, 146 n, 238 e

Answer: C

Q3) Name the compound CH₃CH₂NH₂

Answer: Ethylamine

Q4) Name the compound Co(NO₃)₂

A)Cobalt (I)nitrate

B)Cobalt (II)nitrate

C)Cobalt (I)nitride

D)Cobalt nitrite

E)Cobalt (II)nitride

Answer: B

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Chapter 3: Mass Relationships in Chemical Reactions

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Sample Questions

Q1) What is the coefficient of H₂SO₄ when the following equation is properly balanced?

___ Ca₃(PO₄)₂ + ___ H₂SO₄ \(\rarr\) ___ CaSO₄ + ___ H₃PO₄

A)3

B)8

C)10

D)11

E)none of these

Answer: A

Q2) Balance the following and list the coefficients in order from left to right. ___ Cr + ___ H₂SO₄ \(\rarr\) ___ Cr₂(SO₄)₃+ ___ H₂

A)2, 3, 1, 2

B)2, 3, 1, 3

C)1, 3, 1, 3

D)4, 6, 2, 6

E)1, 3, 1, 2

Answer: B

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Chapter 4: Reactions in Aqueous Solutions

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Sample Questions

Q1) Which of the following is an example of a disproportionation reaction?

A)2C₂H₆(g)+

7O₂(g)\(\rarr\)4CO₂(g)+ 6H₂O(l)

B)2KBr(aq)+ Cl₂(g)\(\rarr\)2KCl(aq)+ Br₂(l)

C)2H₂O₂(aq)\(\rarr\)

2H₂O(l)+ O₂(g)

D)CaBr₂(aq)+

H₂SO₄(aq)\(\rarr\) CaSO₄(s)+ 2HBr(g)

E)2Al(s)+

3H₂SO₄(aq)\(\rarr\)Al₂(SO₄)₃(aq)+ 3H₂(g)

Q2) Zinc dissolves in hydrochloric acid to yield hydrogen gas: Zn(s)+ 2HCl(aq)\(\rarr\) ZnCl₂(aq)+ H₂(g)

What mass of hydrogen gas is produced when a 7.35 g chunk of zinc dissolves in 500.mL of 1.200M HCl?

A)0.605 g

B)0.113 g

C)0.302 g

D)0.453 g

E)0.227 g

Q3) Define the terms solution, solute, and solvent.

Page 6

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Chapter 5: Gases

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Sample Questions

Q1) Define Charles's Law

Q2) At constant pressure, the density of a gas depends on temperature.Does the density increase or decrease as the temperature increases?

Q3) Which statement is false?

A)The average kinetic energies of molecules from samples of different "ideal" gases is the same at the same temperature.

B)The molecules of an ideal gas are relatively far apart.

C)All molecules of an ideal gas have the same kinetic energy at constant temperature.

D)Molecules of a gas undergo many collisions with each other and the container walls.

E)Molecules of greater mass have a lower average speed than those of less mass at the same temperature.

Q4) Calculate the volume occupied by 56.5 g of argon gas at STP.

A)22.4 L

B)31.7 L

C)34.6 L

D)1,270 L

E)1,380 L

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Find the heat absorbed from the surroundings when 15 g of O₂ reacts according to the equation O + O₂ \(\rarr\) O₃, \(\Delta\)H°_rxn= -103 kJ/mol.

A)4.6 × 10^-3 kJ

B)32 kJ

C)48 kJ

D)96 kJ

E)110 kJ

Q2) Thermal 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) A 0.3423 g sample of pentane, C₅H₁₂, was burned in a bomb calorimeter.The temperature of the calorimeter and the 1.000 kg of water contained therein rose from 20.22°C to 22.82°C.The heat capacity of the calorimeter is 2.21 kJ/°C.The heat capacity of water = 4.184 J/g·°C.How much heat was given off during combustion of the sample of pentane?

Q4) Define specific heat.

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Chapter 7: Quantum Theory and the Electronic Structure of Atoms

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Sample Questions

Q1) Calculate the energy, in joules, required to excite a hydrogen atom by causing an electronic transition from the n = 1 to the n = 4 principal energy level.Recall that the energy levels of the H atom are given by E_n = -2.18 × 10^-18 J(1/n²)

A)2.07 × 10^-29 J

B)2.25 × 10^-18 J

C)2.04 × 10^-18 J

D)3.27 × 10^-17 J

E)2.19 × 10⁵ J

Q2) The colors of the visible spectrum are blue, green, orange, red, violet, and yellow.Of these colors, ______ has the least energy.

Q3) Write the ground state electron configuration for Ni.

Q4) What is the energy in joules of one photon of x-ray radiation with a wavelength of 0.120 nm?

A)2.50 x 10⁹ J

B)1.66 x 10^-24 J

C)1.66 x 10^-33 J

D)2.50 x 10¹⁸J

E)1.66 x 10^-15J

Page 9

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Chapter 8: Periodic Relationships Among the Elements

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Sample Questions

Q1) Why is the Mg²⁺ ion smaller than F^-, even though they are isoelectronic?

Q2) The first ionization energy of sodium is 495.9 kJ/mol.The energy change for the reaction Na(s)\(\rarr\)Na⁺(g)+ e^- is therefore

A)495.9 kJ/mol.

B)less than 495.9 kJ/mol.

C)greater than 495.9 kJ/mol.

D)is equal to the electron affinity of sodium.

E)is equal to the second ionization energy of sodium.

Q3) The electron configuration of the outermost electrons of atoms of the halogen group is ns²np⁷.

A)True

B)False

Q4) Write the ground-state electron configuration for O^2-.

Q5) Write the ground-state electron configuration for I^-.

Q6) Write the ground-state electron configuration for Br^-.

Q7) Write the ground-state electron configuration for Ca²⁺.

Q8) Write the ground-state electron configuration for K⁺.

Q9) Write the ground-state electron configuration for Al³⁺.

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Chapter 9: Chemical Bonding I: Basic Concepts

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Sample Questions

Q1) Arrange the following bonds in order of increasing ionic character H - Cl

C - H

H - H

O - H

A)H - Cl < C - H < H - H < O - H

B)H - H < C - H < O - H < H - Cl

C)H - H < C - H < H - Cl < O - H

D)C - H < O - H < H - Cl < H - H

E)C - H < H - Cl < O - H < H - H

Q2) Define electronegativity:

A)an atoms ability to attract electrons that are shared in a chemical bond

B)an atoms ability to form an ionic bond with another atom

C)an atoms ability to donate valence electrons to another atom

D)an atoms ability to form a cation

E)an atoms ability to form double and triple bonds

Q3) Classify the O - H bond in CH₃OH as ionic, polar covalent, or nonpolar covalent.

A)ionic

B)polar covalent

C)nonpolar covalent

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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization

of Atomic

Orbitals

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Sample Questions

Q1) The bond angles in SCl₂ are expected to be

A)a little less than 109.5°.

B)109.5°.

C)a little more than 109.5°.

D)120°.

E)180°.

Q2) Give the number of lone pairs around the central atom and the geometry of the ion IBr₂^-.

A)0 lone pairs, linear

B)1 lone pair, bent

C)2 lone pairs, bent

D)3 lone pairs, bent

E)3 lone pairs, linear

Q3) According to the VSEPR theory, the actual F -As -F bond angles in the AsF₄^- ion are predicted to be

A)109.5°

B)90° and 120°

C)180°

D)< 109.5°

E)< 90° and < 120°

Page 12

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Chapter 11: Intermolecular Forces and Liquids and Solids

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Sample Questions

Q1) Which one of the following substances crystallizes as a covalent crystal?

A)CaO

B)SiO₂

C)CO₂

D)Pb

E)KMnO₄

Q2) Which one of the following crystallizes in a metallic lattice?

A)C

B)NaMnO₄

C)K

D)LiClO₄

E)K₂Cr₂O₇

Q3) The molar enthalpy of vaporization of boron tribromide is 30.5 kJ/mol, and its normal boiling point is 91°C.What is the vapor pressure of BBr₃ at 20°C?

A)11.5 torr

B)311 torr

C)5.31 torr

D)143 torr

E)66.1 torr

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Chapter 12: Physical Properties of Solutions

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Sample Questions

Q1) Arrange the following aqueous solutions in order of increasing boiling points: 0.300m C₆H₁₂O₆, 0.110m K₂CO₃, and 0.050m Al(ClO₄)₃

A)C₆H₁₂O₆ < K₂CO₃ < Al(ClO₄)₃

B)Al(ClO₄)₃< C₆H₁₂O₆ < K₂CO₃

C)C₆H₁₂O₆ < Al(ClO₄)₃ < K₂CO₃

D)K₂CO₃ < C₆H₁₂O₆ < Al(ClO₄)₃

E)K₂CO₃< Al(ClO₄)₃< C₆H₁₂O₆

Q2) Using your knowledge of osmosis or osmotic pressure explain the following statement: Drinking salt water actually dehydrates our tissues.

Q3) The solubility of a solid always increases with increasing solvent temperature.

A)True

B)False

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) For what order reaction does the half-life get longer as the initial concentration increases?

A)zero order

B)first order

C)second order

D)none of them because half-life is always independent of the initial concentration

Q2) For the chemical reaction A \(\rarr\) C, a plot of 1/[A]_t versus time was found to give a straight line with a positive slope.What is the order of reaction?

A)zeroth

B)first

C)second

D)Such a plot cannot reveal the order of the reaction.

Q3) For a second order reaction, the half-life is equal to

A)t_1/2 = 0.693/k

B)t_1/2 = k/0.693

C)t_1/2 = 1/k[A]_o

D)t_1/2 = k

E)t_1/2 = [A]_o/2k

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) A reaction with an equilibrium constant K_c = 1.5 x 10²¹ would consist of which of the following at equilibrium:

A)approximately equal reactants and products

B)some reactants and products with reactants slightly favored

C)some reactants and products with products slightly favored

D)essentially all reactants

E)essentially all products

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) A reaction with an equilibrium constant K_c = 1.5 x 10^-25 would consist of which of the following at equilibrium:

A)approximately equal reactants and products

B)some reactants and products with reactants slightly favored

C)some reactants and products with products slightly favored

D)essentially all reactants

E)essentially all products

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Chapter 15: Acids and Bases

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Sample Questions

Q1) The pH of a 0.095 M solution of an unknown monoprotic acid is 5.42.Calculate the K_a of the acid.

A)3.8 x 10^-6

B)3.6 x 10^-7

C)1.5 x 10^-10

D)2.6 x 10^-9

E)2.8 x 10^-8

Q2) The hydronium ion and the hydroxide ion, in that order, are:

A)H₃O⁺, OH⁺

B)OH^-, H₃O^-

C)OH^-, H⁺

D)H₃O⁺, OH^-

E)H₃O^-, OH^-

Q3) Arrange the acids HOCl, HClO₃, and HClO₂ in order of increasing acid strength.

A)HOCl < HClO₃ < HClO₂

B)HOCl < HClO₂ < HClO₃

C)HClO₂ < HOCl < HClO₃

D)HClO₃< HOCl < HClO₂

E)HClO₃< HClO₂ < HOCl

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Chapter 16: Acid-Base Equilibria and Solubility Equilibria

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Q1) For which type of titration will the pH be basic at the equivalence point?

A)Strong acid vs.strong base.

B)Strong acid vs.weak base.

C)Weak acid vs.strong base.

D)All of the above.

E)None of the above.

Q2) The solubility of lead(II)iodide is 0.064 g/100 mL at 20ºC.What is the solubility product for lead(II)iodide?

A)1.1 × 10^-8

B)3.9 × 10^-6

C)1.1 × 10^-11

D)2.7 × 10^-12

E)1.4 × 10^-3

Q3) The K_sp for silver(I)phosphate is 1.8 × 10^-18.Calculate the molar solubility of silver(I)phosphate.

A)1.6 × 10^-5 M

B)2.1 × 10^-5 M

C)3.7 × 10^-5 M

D)7.2 × 10^-1 M

E)1.8 × 10^-1 M

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Chapter 17: Entropy Free Energy and Equilibrium

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Q1) Predict the signs (-, +, or 0)of \(\Delta\)H and \(\Delta\)S, in that order, for the expansion of an ideal gas into a vacuum.

Q2) Ozone (O₃)in the atmosphere can reaction with nitric oxide (NO): O₃(g)+ NO(g)\(\rarr\)NO₂(g)+ O₂(g).

Calculate the \(\Delta\)G° for this reaction at 25°C.(\(\Delta\)H° = -199 kJ/mol, \(\Delta\)S° = -4.1 J/K·mol)

A)1020 kJ/mol

B)-1.22 × 10³ kJ/mol

C)2.00 × 10³ kJ/mol

D)-1.42 × 10³ kJ/mol

E)-198 kJ/mol

Q3) Under what conditions (always, never, high temperature only, low temperature only)is the expansion of an ideal gas into a vacuum expected to be spontaneous?

Q4) The entropy change \(\Delta\)S° for the reaction NH₄Cl(s)\(\rarr\)NH₃(g)+ HCl(g)will be negative.

A)True

B)False

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Calculate E°_cell for the following reaction: 2Fe²⁺(aq)+ Cd²⁺(aq)\(\rarr\)2Fe³⁺(aq)+ Cd(s)

A)-0.37 V

B)0.37 V

C)-1.17 V

D)1.17 V

E)None of these.

Q2) For the electrochemical cell, Fe(s)| Fe²⁺(aq)|| Cu²⁺(aq)| Cu⁺(aq)| Pt(s), determine the equilibrium constant (K_eq)at 25°C for the reaction that occurs.

A)1.6 x 10^-10

B)8.6 x 10¹⁹

C)1.2 x 10^-20

D)9.3 x 10⁹

E)1.3 x 10^-5

Q3) Complete and balance the following redox reaction under basic conditions: CrO₄^2-(aq)+ SO₃^2-(aq)\(\rarr\) Cr(OH)₃(s)+ SO₄^2-(aq)

Q4) Will H₂(g)form when Ag is placed in 1.0 M HCl?

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Chapter 19: Nuclear Chemistry

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Sample Questions

Q1) The half-life of ⁹⁰Sr is 29 years.What fraction of the atoms in a sample of ⁹⁰Sr would remain 175 years later?

A)0.17

B)0.12

C)0.062

D)0.015

E)0.50

Q2) A sample of a radioisotope shows an activity of 999 disintegrations per minute due to beta decay.If after 1.10 years the activity is 952 disintegrations per minute, what is the half-life of this radioisotope?

A)4.38 × 10^-2 yr

B)11.4 yr

C)0.25 yr

D)15.8 yr

E)9.1 yr

Q3) Compare nuclear reactions with chemical reactions in terms of the conservation of mass.

Q4) When a ⁸⁷Br nucleus emits a beta particle, the nuclear species that results is ___________.

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Chapter 20: Chemistry in the Atmosphere

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Q1) The normal process occurring on Earth that warms its surface and the lower atmosphere is called

A)photosynthesis.

B)vaporization of water.

C)incomplete combustion.

D)the greenhouse effect.

E)the ozone hole.

Q2) Heat is radiated from Earth to space predominately in the form of:

A)infrared radiation.

B)radioactivity.

C)visible light.

D)solar flares.

E)UV radiation.

Q3) The compound CFCl₃ is used as a/an

A)enzyme

B)anesthetic

C)gaseous fuel

D)coolant

E)CFC replacement

Q4) Name four "greenhouse gases" besides carbon dioxide.

Page 22

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Chapter 21: Metallurgy and the Chemistry of Metals

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Q1) What is the total number of moles (n)of electrons exchanged between the oxidizing agent and the reducing agent in the reaction below to obtain chromium metal: Cr₂O₃(s)+ 2Al(s)\(\rarr\) 2Cr(l)+ Al₂O₃(s)

A)1

B)2

C)3

D)6

E)21

Q2) A naturally occurring substance with a range of chemical composition is

A)an element.

B)a mineral.

C)gangue.

D)an ore.

E)an amalgam.

Q3) In n-type semiconductors

A)the energy gap between the valence band and the conduction band is very large.

B)impurities that donate electrons are added to provide conduction electrons.

C)a valence band overlaps the empty conduction band.

D)impurities that provide "positive holes" are added to a pure semiconductor.

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Chapter 22: Nonmetallic Elements and Their Compounds

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Q1) Determine all of the ionic hydrides from the compounds listed below. I.CaH<sub>2

</sub>II.H₂S III.SiH₄

A)I only

B)II only

C)I and II

D)I and III

E)II and III

Q2) Alkali metal hydrides are very reactive with water, forming H₂ gas.

A)True

B)False

Q3) A chemical formula of the carbide ion is

A)Si^4-

B)C₂^-

C)C₂^2-

D)CN^-

E)C₃^-

Q4) Write an equation for a laboratory preparation of nitrogen gas.

Q5) What is carborundum, and how is it prepared.What physical property of carborundum makes it a useful material?

Q6) Write the formula for the binary hydride you would expect calcium to form.

Page 24

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Chapter 23: Transition Metal Chemistry and Coordination Compounds

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Q1) The electron configuration of an Fe²⁺ ion is

A)[Ar]4s²4d⁴.

B)[Ar]4s²3d⁶.

C)[Ar]3d³.

D)[Ar]3d⁵.

E)[Ar]3d⁶.

Q2) Which of these ligands produces the weakest crystal field?

A)CN^-

B)I^-

C)OH^-

D)H₂O

E)NH₃

Q3) In the complex ion

[Cr(C₂O₄)₂(H₂O)₂]^-, the oxidation number of Cr is A)+1.

B)+2. C)+3.

D)-2. E)-1.

Q4) Write the chemical formula of diamminedichloroplatinum(II).

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Chapter 24: Organic Chemistry

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Q1) Which one of these hydrocarbons does not have isomers?

A)C₇H₁₆

B)C₆H₁₄

C)C₅H₁₀

D)C₄H₈

E)C₃H₈

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) Bromination of benzene (C₆H₆), an aromatic compound,

A)occurs by substitution rather than addition.

B)occurs by addition rather than substitution.

C)occurs more rapidly than bromination of a nonaromatic compound.

D)results in formation of 1,2,3,4,5,6-hexabromocyclohexane.

E)occurs in the absence of a catalyst.

Q4) A compound with the formula C₆H₁₂ may or may not be a saturated hydrocarbon.Explain.

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Chapter 25: Synthetic and Natural Organic Polymers

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Q1) The intermolecular force between bases on the opposite strands of DNA responsible for its double-helical structure is

A)hydrogen bonding.

B)dispersion force.

C)covalent bonding.

D)ionic force.

E)dipole-dipole force.

Q2) Amides are synthesized from two classes of organic compounds.Those two types of compounds are

A)carboxylic acids and alkenes.

B)amines and alcohols.

C)alcohols and carboxylic acids.

D)amines and carboxylic acids.

E)alkenes and amines.

Q3) Which one of these molecules is part of the make-up of both DNA and RNA?

A)deoxyribose

B)ribose

C)phosphate

D)thymine

E)uracil

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