Chemistry for Engineers Final Exam Questions - 2075 Verified Questions

Course Introduction

Chemistry for Engineers

Final Exam Questions

Chemistry for Engineers introduces fundamental concepts of general chemistry with a focus on applications relevant to engineering disciplines. The course covers atomic and molecular structure, chemical bonding, stoichiometry, thermodynamics, chemical kinetics, and equilibrium. Special emphasis is placed on materials science, corrosion, environmental chemistry, and the role of chemical principles in the design and optimization of engineering processes. Through lectures and laboratory experiments, students develop problem-solving skills and gain practical experience in chemical analysis, preparing them to address chemical challenges in various engineering fields.

Recommended Textbook Chemistry and Chemical Reactivity 9th Edition by John C. Kotz\New folder

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Page 2

Chapter 1: Basic Concepts of Chemistry

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Q1) Which one of the following statements is not a comparison of physical properties?

A) Mercury and gallium are both liquids at 50\(^{\circ}\)C.

B) Oxygen is more soluble in water than helium.

C) Silver and gold are malleable metals.

D) Oxygen and nitrogen are both liquids at -200\(^{\circ}\)C.

E) Calcium dissolves more quickly than iron in acids.

Answer: E

Q2) A(n)________ is the smallest particle of an element that retains the characteristic chemical properties of that element.

Answer: atom

Q3) The element whose symbol is Pb is

A) lead.

B) antimony.

C) lanthanum.

D) phosphorus.

E) none of these.

Answer: A

Q4) The law of ________ states that the total energy of the universe is constant. Answer: conservation of energy

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Chapter 2: Mathlets Review: The Tools of Quantitative Chemistry

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Q1) Calculate the mass of tin that occupies the same volume as 86.9 g of cobalt.The density of cobalt is 8.90 g/cm³ and the density of tin is 7.31 g/cm³.

A) 0.749 g

B) 71.4 g

C) 5.65 \(\times\) 10³ g

D) 0.00945 g

E) 1.34 g

Answer: B

Q2) How many 350-mg aspirin tablets can be made from 35.0 kg of aspirin?

A) 10,000,000

B) 1,000,000

C) 1000

D) 10,000

E) 100,000

Answer: E

Q3) Assuming the density of water is 1.00 g/cm³,the mass of 1.0 cubic meter (m³) of water is ________ grams.

Answer: 1.0 \(\times\) 10⁶

Page 4

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Chapter 3: Atoms, molecules, and Ions

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Q1) A certain element consists of two stable isotopes.The first has a mass of 14.0031 amu and a percent natural abundance of 99.63%.The second has a mass of 15.001 amu and a percent natural abundance of 0.37%.What is the atomic weight of the element?

A) 13.95 u

B) 14.00 u

C) 14.01 u

D) 14.50 u

E) 19.50 u

Answer: C

Q2) What element is in the fourth period in Group 3A?

A) Sb

B) Ga

C) In

D) Si

E) Tl

Answer: B

Q3) Elements that have the same number of protons,but differ in their number of neutrons are called ________.

Answer: isotopes

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Page 5

Chapter 4: Chemical Reactions

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Q1) Which anion will form a precipitate with Ca²⁺?

A) Cl^-

B) OH^-

C) C₂H₃O₂^-

D) Br^-

E) none

Q2) Which of the following would \(\underline{\text{ not }}\) be depicted as the individual ions on the reactant side of a complete ionic reaction?

A) NaOH

B) HCl

C) Cd(NO₃)₂

D) CdCO₃

E) CrCl₃

Q3) All of the following compounds are \(\underline{\text{insoluble }}\) in water except ____.

A) BaCO₃

B) PbF₂

C) Fe(OH)₃

D) Ni(ClO₄)₂

E) PbCrO₄

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Chapter 5: Stoichiometry: Quantitative Information About Chemical Reactions

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Q1) Sulfur hexafluoride is produced by reacting elemental sulfur with fluorine gas.S₈(s)+ 24 F₂(g)\(\to\) 8 SF₆(g)

What is the percent yield if 18.3 g SF₆ is isolated from the reaction of 10.0 g S₈ and 30.0 g F₂?

A) 40.2%

B) 45.8%

C) 47.6%

D) 54.6%

E) 61.0%

Q2) In order to dilute 35.5 mL of 0.533 M HCl to 0.100 M,the volume of water that must be added is

A) 28.8 mL.

B) 6.66 mL.

C) 1.89 \(\times\) 10² mL.

D) 1.50 \(\times\) 10^-3 mL.

E) 1.54 \(\times\) 10² mL.

Q3) The percent yield of a chemical reaction is calculated by dividing the actual yield by the ________ yield and multiplying by 100%.

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Chapter 6: Principles of Chemical Reactivity: Energy and Chemical Reactions

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Q1) If 35.0 g H₂O at 22.7 S1U1P1\(\circ\)S1S1P0C is combined with 65.0 g H₂O at 87.5 S1U1P1\(\circ\)S1S1P0C,what is the final temperature of the mixture? The specific heat capacity of water is 4.184 J/g.K.

A) 25.1 S1U1P1\(\circ\)S1S1P0C

B) 45.4 S1U1P1\(\circ\)S1S1P0C

C) 50.8 S1U1P1\(\circ\)S1S1P0C

D) 64.8 S1U1P1\(\circ\)S1S1P0C

E) 48.9 S1U1P1\(\circ\)S1S1P0C

Q2) When 0.236 mol of a weak base (A^-)is reacted with excess HCl,6.91 kJ of energy is released as heat.What is \(\Delta\)H for this reaction per mole of A^- consumed?

A) -34.2 kJ/mol

B) -59.4 kJ/mol

C) -29.3 kJ/mol

D) 34.2 kJ/mol

E) 29.3 kJ/mol

Q3) Dry ice converts directly from a solid to a gas when heated.This process is called

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Chapter 7: The Structure of Atoms

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Q1) If a hydrogen atom in the excited n = 4 state relaxes to the ground state,what is the maximum number of possible emission lines?

A) 1

B) 3

C) 6

D) 8

E) infinite

Q2) How many p orbitals are in the n = 3 shell?

A) 1

B) 6

C) 0

D) 5

E) 3

Q3) The ____ of a photon of light is ____ proportional to its frequency and ____ proportional to its wavelength.

A) energy,directly,inversely

B) energy,inversely,directly

C) velocity,directly,inversely

D) intensity,inversely,directly

E) amplitude,directly,inversely

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Chapter 8: The Structure of Atoms and Periodic Trends

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Q1) Which of the following atoms is diamagnetic in its ground state?

A) ytterbium (Yb)

B) holmium (Ho)

C) tantalum (Ta)

D) aluminum (Al)

E) germanium (Ge)

Q2) How many valence electrons does a nitrogen atom have?

A) 5

B) 8

C) 7

D) 3

E) 15

Q3) Which of the following orbital occupancy designations is \(\underline{\text{ incorrect? }}\)

A) 2s²

B) 3s²

C) 1s²

D) 4d³

E) 1s⁶

Q4) Explain the difference between paramagnetic and ferromagnetic.

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Chapter 9: Bonding and Molecular Structure

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Q1) Use VSEPR theory to predict the molecular geometry around the carbon atom in formaldehyde,H₂CO.

A) linear

B) bent

C) trigonal-planar

D) tetrahedral

E) octahedral

Q2) The molecule H₂S has

A) 2 bonding pairs and 3 lone pairs

B) 2 bonding pairs and 2 lone pairs

C) 3 bonding pairs and 1 lone pair

D) 3 bonding pairs and 3 lone pairs

E) none of these

Q3) Which of the following elements is most likely to form a molecular structure that disobeys the octet rule?

A) B

B) C

C) N

D) O

E) F

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Chapter 10: Bonding and Molecular Structure: Orbital

Hybridization and Molecular Orbitals

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Q1) What is the molecular geometry around an atom that has 2 \(\sigma\) bonds,0 \(\pi\) bonds,and 1 lone pair of electrons?

A) linear

B) trigonal planar

C) tetrahedral

D) bent

E) trigonal pyramidal

Q2) Ammonia reacts with oxygen and water to produce nitric acid.What change in hybridization of the nitrogen atom occurs in this reaction?

A) sp³ to sp²

B) sp³ to sp

C) sp² to sp³

D) sp² to sp

E) no change

Q3) In valence bond theory,each sigma bond in CH₄ is formed from the overlap of a hydrogen atom's 1s orbital with a(n)________ hybridized orbital on the carbon atom.

Q4) The hybridization of the carbon atom in CF₃⁺ is

Page 12

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Chapter 11: Gases and Their Properties

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Q1) A 2.00-L glass soda bottle filled only with air is tightly capped at 19°C and 680.0 mmHg.If the bottle is placed in water at 81°C,what is the pressure in the bottle?

A) 160 mmHg

B) 824 mmHg

C) 2900 mmHg

D) 561 mmHg

E) 422 mmHg

Q2) The partial pressures of CH₄,N₂,and O₂ in a sample of a gas mixture were found to be 183 mmHg,493 mmHg,and 551 mmHg,respectively.Calculate the mole fraction of nitrogen.

A) 0.725

B) 0.359

C) 19.7

D) 0.449

E) 0.402

Q3) The ideal gas law can be modified to correct for the errors arising from nonideality.The modified equation is known as the ________ equation of state for a real gas.

Q4) What are two ways in which real gases differ from ideal gases?

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Chapter 12: Intermolecular Forces and Liquids

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Q1) Which of the following phase transitions is/are endothermic?

1)liquid water freezing to a solid at its normal freezing point of 0.0 S1U1P1\(\circ\)S1S1P0C.

2)the sublimation of solid carbon dioxide into the gas phase.

3)gaseous sulfur dioxide condensing into a liquid at its boiling point of -10.0 S1U1P1\(\circ\)S1S1P0C.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q2) As pure molecular solids,which of the following exhibits dipole-dipole intermolecular forces: HBr,NBr₃,SBr₂,and CBr₄?

A) HBr only

B) CBr₄ only

C) HBr and SBr₂

D) NBr₃ and CBr₄

E) HBr,NBr₃,and SBr₂

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Chapter 13: The Chemistry of Solids

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Q1) Which process requires the greatest endothermic change in enthalpy for water?

A) freezing

B) condensation

C) sublimation

D) melting

E) vaporization

Q2) Gold (atomic mass 197.0 g/mol),with an atomic radius of 144.2 pm,crystallizes in a face-centered cubic lattice.What is the density of gold?

A) 9.65 g/cm³

B) 1.21 g/cm³

C) 4.82 g/cm³

D) 2.41 g/cm³

E) 19.3 g/cm³

Q3) Which of the following is expected to have the most negative lattice enthalpy?

A) LiCl

B) NaCl

C) KCl

D) RbCl

E) CsCl

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Chapter 14: Solutions and Their Behavior

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Q1) Which of the following liquids will be miscible with water in any proportion: ethanol (CH₃CH₂OH),carbon tetrachloride (CCl₄),hexane (C₆H₁₄),and/or formic acid (HCO₂H)?

A) ethanol and carbon tetrachloride

B) carbon tetrachloride and hexane

C) ethanol and formic acid

D) ethanol,carbon tetrachloride,and benzene

E) carbon tetrachloride,and formic acid

Q2) Ideally,colligative properties depend only on the

A) relative numbers of solute and solvent particles in a solution.

B) molar masses of the solute particles in a solution.

C) density of a solution.

D) hydrated radii of the molecules or ions dissolved in a solution.

E) partial pressure of the gases above the surface of a solution.

Q3) All of the following are colloidal dispersions EXCEPT ____.

A) marshmallow

B) white wine

C) milk

D) whipped cream

E) cheese

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Chapter 15: Chemical Kinetics: the Rates of Chemical Reactions

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Q1) The rate constant for a reaction at 40.0°C is exactly 6 times that at 20.0°C.Calculate the Arrhenius energy of activation for the reaction.(R = 8.314 J/K.0mol)

A) 6.00 kJ/mol

B) 8.22 kJ/mol

C) 68.3 kJ/mol

D) 14.9 kJ/mol

E) none of these

Q2) The rate constant for a particular reaction is 0.0040 M.s^-1.What is the overall order of this reaction?

A) 0

B) 1

C) 2

D) 3

E) 4

Q3) If a catalyst is present in a different phase from the reactants and products,it is referred to as a(n)________ catalyst.

Q4) Draw the reaction coordinate diagram for an exothermic reaction involving three mechanistic steps.

Page 17

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Chapter 16: Principles of Reactivity: Chemical Equilibria

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Q1) Which of the following statements is/are CORRECT?

1)For a chemical system,if the reaction quotient (Q)is greater than K,reactant must be converted to products to reach equilibrium.

2)For a chemical system at equilibrium,the forward and reverse rates of reaction are equal.

3)For a chemical system at equilibrium,the concentrations of products divided by the concentrations of reactants equals one.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q2) If a stress is applied to an equilibrium system,the system will respond in such a way as to relieve that stress.This is a statement of ________ principle.

Q3) In 1913,the Haber-Bosch process was patented.The product of the Haber-Bosch process is ________.

Q4) The symbol Q is called the ________.

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Chapter 17: The Chemistry of Acids and Bases

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Q1) What is the pH of a 0.28 M solution of sodium propionate,NaC₃H₅O₂,at 25°C? (propionic acid,HC₃H₅O₂,is monoprotic and has a K_a = 1.3 \(\times\) 10^-5at 25°C..K_w = 1.01 \(\times\) 10^-14)

A) 6.26

B) 4.83

C) 11.10

D) 7.74

E) 9.17

Q2) What is the pH of the solution which results from mixing 50.0 mL of 0.30 M HF(aq)and 50.0 mL of 0.30 M NaOH(aq)at 25 S1U1P1\(\circ\)S1S1P0C? (K_a of HF = 7.2 \(\times\) 10^-4)

A) 1.98

B) 5.84

C) 8.16

D) 10.85

E) 12.02

Q3) When a Lewis acid and a Lewis base combine,the product may be referred to as an acid-base ________.

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Chapter 18: Principles of Reactivity: Other Aspects of Aqueous Equilibria

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Q1) Hyperventilation can cause your blood pH to rise.One way to lower your blood pH is to breath into a paper bag,thus recycling the air you exhale.Why does this procedure lower your blood pH?

Q2) What is the pH of a buffer that results when 0.50 mole of H₃PO₄ is mixed with 0.75 mole of NaOH and diluted with water to 1.00 L? (The acid dissociation constants of phosphoric acid are K_a1 = 7.5 \(\times\) 10^-3,K_a2 = 6.2 \(\times\) 10^-8,and K_a3 = 3.6 \(\times\) 10^-13)

A) 1.82

B) 2.12

C) 6.91

D) 7.21

E) 12.44

Q3) When a weak base is titrated with a strong acid,the pH at the equivalence point is ___.

A) equal to 7

B) greater than 7.

C) less than 7.

D) equal to the acid pK_a.

E) equal to the base pK_b.

Page 20

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Chapter 19: Entropy and Free Energy

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Q1) Which of the following is true for the vaporization of a liquid substance?

A) (\(\Delta\)S) = 0 and \(\Delta\)H = 0.

B) (\(\Delta\)S) < 0 and \(\Delta\)H < 0.

C) (\(\Delta\)S) < 0 and \(\Delta\)H > 0.

D) (\(\Delta\)S) > 0 and \(\Delta\)H > 0.

E) (\(\Delta\)S) > 0 and \(\Delta\)H < 0.

Q2) The standard entropy for the formation of SF₆(g)from the elements, S(s)+ 3 F₂(g)\(\to\)SF₆(g)

Is -348.7 J/K.mol-rxn at 298.15 K.Calculate the standard molar entropy of SF₆(g)given SS1U1P1\(\circ\)S1S1P0[S(s)] = 32.1 J/K.mol and SS1U1P1\(\circ\)S1S1P0[F₂(g)] = 202.8 J/K.mol.

A) -988.6 J/K.mol

B) +291.8 J/K.mol

C) -291.8 J/K.mol

D) -113.2 J/K.mol

E) +113.8 J/K.mol

Q3) For any process,the change in entropy of the universe equals the sum of the entropy changes for the system and for the ________.

Q4) At what temperature (in kelvin units)is the entropy of a pure crystal 0.0 J/K.

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Chapter 20: Principles of Reactivity: Electron Transfer Reactions

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Q1) Calculate \(\Delta\)_rGS1U1P1\(\circ\)S1S1P0 for the disproportionation reaction of Cu⁺ at 25 S1U1P1\(\circ\)S1S1P0C, 2 Cu⁺(aq)\(\to\)Cu²⁺(aq)+ Cu(s)

Given the following thermodynamic information. Cu⁺(aq)+ e^- \(\to\) Cu(s)\(~~~~~~~~\)ES1U1P1\(\circ\)S1S1P0 = +0.518 V Cu²⁺(aq)+ 2 e^- \(\to\)Cu(s)\(~~~~~~~~\)ES1U1P1\(\circ\)S1S1P0 = +0.337 V

A) -165 kJ/mol.rxn

B) -135 kJ/mol.rxn

C) -34.9 kJ/mol.rxn

D) +17.5 kJ/mol.rxn

E) +135 kJ/mol.rxn

Q2) In an electrolytic cell,reduction occurs at the ________ and oxidation occurs at the ________.

Q3) When a secondary battery provides electrical energy,it is acting as a voltaic cell.When the battery is recharging,it is operating as a(n)________ cell.

Q4) Explain the function of a salt bridge in a voltaic cell.

Page 22

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Chapter 21: Environmental Chemistry Earths

Environment,energy,and Sustainability

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Q1) Ocean acidification is the result of the increasing atmospheric concentration of what gas?

A) CH₄

B) NO₂

C) CO₂

D) SO₂

E) SO₃

Q2) What is wrong with the statement,"To stop global worming we must eliminate all greenhouse gases."?

Q3) What mass of sulfur dioxide gas would be produced from 6.00 kg of coal which contains 1.15 % sulfur by mass?

A) 276 g

B) 12.0 g

C) 34.5 g

D) 138 g

E) 103 g

Q4) Explain why fossil fuels,hydrocarbons produced by natural processes,are considered a nonrenewable energy resource.

Q5) Explain why the melting of Arctic sea ice will not raise the sea level.

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Chapter 22: The Chemistry of the Main Group Elements

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Q1) At high temperatures limestone decomposes according to which balanced chemical equation?

A) Ca(OH)₂(s)\(\to\) CaO(s)+ H₂O(g)

B) 2 CaO(s)\(\to\) 2 Ca(s)+ O₂(g)

C) CaCO₃(s)\(\to\) CaO(s)+ CO₂(g)

D) Ca(OH)₂(s)\(\to\) CaH₂(s)+ O₂(g)

E) 2 CaCO₃(s)\(\to\) 2 Ca(s)+ O₂(g)+ CO₂(g)

Q2) How is sodium metal commercially produced?

A) By electrolysis of molten Na₂CO₃.

B) By electrolysis of NaCl(aq).

C) By electrolysis molten NaCl using a Downs Cell.

D) By reaction of potassium metal vapor with molten NaCl.

E) By reduction of NaOH with carbon.

Q3) What is the molecular formula of chlorous acid?

A) HClO

B) HClO₂

C) HClO₃

D) HClO₄

E) HCl

Q4) Give two examples of the allotropes of elemental oxygen.

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Chapter 23: The Chemistry of the Transition Elements

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Q1) Which of the following will not as a ligand to a transition metal cation?

A) O₂

B) H₃O⁺

C) PH₃

D) NO₂^-

E) F^-

Q2) What are the possible geometries of a metal complex with a coordination number of 6?

1)square planar

2)tetrahedral

3)octahedral

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q3) EDTA^4-,a hexadentate ligand,is most effective at chelating metal ions when the pH is greater than 12.Explain why low pH decreases the chelating ability of the ligand.

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Chapter 24: Carbon: Not Just Another Element

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Q1) The reaction of methane with chlorine in the presence of ultraviolet radiation can produce ____.

A) chloromethane

B) dichloromethane

C) trichloromethane

D) carbon tetrachloride

E) all of these products

Q2) What is the molecular formula for heptane?

A) C₆H₁₂

B) C₆H₁₄

C) C₇H₁₄

D) C₇H₁₆

E) C₉H₂₀

Q3) What is the name of the product of the reaction that occurs when a mixture of methanol and propionic acid is heated in the presence of acid?

A) methyl propylether

B) 2-propanone

C) propyl methanoate

D) 2-methylpropanal

E) methyl propanoate

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Chapter 25: Biochemistry

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Q1) Which of the following statements concerning polysaccharides is/are correct?

1)Humans do not have enzymes that will break down cellulose.

2)Amylopectin and glycogen are both branched chains of glucose.

3)Amylose starch is a linear polysaccharide.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

Q2) Nonpolar amino acid side chains contain substituents made mostly of what atoms?

A) carbon and hydrogen

B) nitrogen and oxygen

C) carbon and nitrogen

D) carbon and oxygen

E) nitrogen and hydrogen

Q3) The amide linkages in a protein and also called ___ bonds.

Q4) When a drop of the fatty acid oleic acid (CH₃(CH₂)₇CH=CH(CH₂)<sub>7</sub >CO₂H)is added to water it forms a molecule thick layer on the surface of the water.What is the likely orientation of the acid on the water surface?

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Chapter 26: Nuclear Chemistry

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Source URL: https://quizplus.com/quiz/72721

Sample Questions

Q1) In positron emission tomography (PET),a positron emitted from an unstable isotope travels a short distance before it is annihilated by

A) an electron,creating a proton that is detected by the instrument.

B) a neutron,creating two gamma rays that travel in opposite directions.

C) an electron,creating two gamma rays that travel in opposite directions.

D) an alpha particle,creating two protons that travel in opposite directions.

E) gamma ray,creating an electron that is detected by the instrument.

Q2) Which of the following statements is/are CORRECT?

1)Naturally occurring isotopes account for only a small fraction of known radioactive isotopes.

2)A few radioactive isotopes with long half-lives,such as U-235 and U-238,are found in nature.

3)Trace quantities of some short-lived radioactive isotopes,such as C-14,are found in nature because they are formed continuously by nuclear reactions.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 2

E) 1,2,and 3

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