Chemistry for Engineers Question Bank - 2075 Verified Questions

Course Introduction

Chemistry for Engineers Question

Bank

Chemistry for Engineers provides a foundational understanding of chemical principles and their practical applications within engineering contexts. The course explores topics such as atomic and molecular structure, chemical bonding, thermodynamics, reaction kinetics, equilibria, and material properties. Emphasis is placed on how these concepts underpin processes in areas including materials science, environmental engineering, and energy systems. Through lectures, problem-solving, and laboratory experiments, students develop analytical skills and gain insight into the chemical processes essential for innovation and design in various engineering disciplines.

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 \(\underline{\text{ correct? }}\)

A) A pure substance may be separated by filtration or distillation into two or more components.

B) A heterogeneous mixture is also known as a solution.

C) A heterogeneous mixture is composed of two or more substances in the same phase.

D) The composition is uniform throughout a homogeneous mixture.

E) The combination of a liquid and a solid always results in a heterogeneous mixture.

Answer: D

Q2) An electrically charged atom or group of atoms is a(n)________.

A) element

B) ion

C) molecule

D) heterogeneous mixture

E) solution

Answer: B

Q3) Density is an example of a(n)________ property,which is one that does not depend on the amount of a substance.

Answer: intensive

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3

Chapter 2: Mathlets Review: The Tools of Quantitative Chemistry

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Q1) How many significant figures are there in the number 8.500?

A) 1

B) 5

C) 3

D) 4

E) 2

Answer: D

Q2) What is the volume of a cube that has an edge length of 0.013 m?

A) 2.2 \(\times\) 10^-3 m³

B) 2.2\(\times\) 10^-3 km³

C) 2.2 \(\times\) 10^-3 cm³

D) 2.2 \(\times\) 10^-3 mm³

E) 2.2 cm³

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

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Page 4

Chapter 3: Atoms, molecules, and Ions

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Sample Questions

Q1) Silver has two stable isotopes with masses of 106.90509 u and 108.9047 u.The atomic weight of silver is 107.868 u.What is the percent abundance of each isotope?

A) 50.0% Ag-107 and 50.0% Ag-109

B) 51.8% Ag-107 and 48.2% Ag-109

C) 55.4% Ag-107 and 44.6% Ag-109

D) 48.2% Ag-107 and 51.8% Ag-109

E) 44.6% Ag-107 and 55.4% Ag-109

Answer: B

Q2) What are the names of four metalloids?

Answer: boron,silicon,germanium,arsenic,(antimony,and tellurium)

Q3) What is the molar mass of sodium sulfate?

A) 55.06 g/mol

B) 119.1 g/mol

C) 78.05 g/mol

D) 142.0 g/mol

E) 110.0 g/mol

Answer: D

Q4) Millikan's oil drop experiment determined the charge of the ________. Answer: electron

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Chapter 4: Chemical Reactions

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Q1) What is the oxidation number of each O in BaFeO₄?

A) +6

B) +2

C) -2

D) +3

E) 0

Q2) The reaction of elemental chlorine with potassium iodide yields elemental iodine and potassium chloride.Write a balanced chemical equation for this reaction.

A) Cl₂(g)+ KI(s)\(\to\) I(s)+ KCl₂(s)

B) Cl₂(g)+ 2 KI(s)\(\to\)I₂(s)+ 2 KCl(s)

C) Cl₂(g)+ KI₂(s)\(\to\)I₂(s)+ KCl₂(s)

D) Cl(g)+ KI(s)\(\to\) I(s)+ KCl(s)

E) Cl₂(g)+ 2 K₂I(s)\(\to\) I₂(s)+ 2 K₂Cl(s)

Q3) A(n)________ agent gains electrons in an oxidation-reduction reaction.

Q4) A solution is a homogeneous mixture composed of one or more ____________ dissolved in a solvent.

Q5) _____-oxides produce acids when reacted with water.

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Chapter 5: Stoichiometry: Quantitative Information About Chemical Reactions

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Q1) How many grams of dioxygen are required to completely burn 4.3 g of C₂H₅OH?

A) 19 g

B) 15 g

C) 3.0 g

D) 42 g

E) 58 g

Q2) A compound consists of only C and F.It contains 38.71% C by mass.What is the empirical formula of the compound?

A) CF

B) CF₂

C) CF₃

D) C₂F

E) CF₄

Q3) What volume of 2.52 M HCl is required to prepare 176.5 mL of 0.449 M HCl?

A) 9.91 \(\times\) 10² mL

B) 1.56 \(\times\) 10² mL

C) 31.4 mL

D) 0.0318 mL

E) 2.00 \(\times\) 10² mL

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

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Sample Questions

Q1) When 10.0 g KOH is dissolved in 100.0 g of water in a coffee-cup calorimeter,the temperature rises from 25.18 S1U1P1\(\circ\)S1S1P0C to 47.53 S1U1P1\(\circ\)S1S1P0C.What is the enthalpy change per gram of KOH dissolved in the water? Assume that the solution has a specific heat capacity of 4.18 J/g.K.

A) -116 J/g

B) -934 J/g

C) -1.03 \(\times\) 10³ J/g

D) -2.19 \(\times\) 10³ J/g

E) -1.03 \(\times\) 10⁴ J/g

Q2) What is the change in internal energy of the system (\(\Delta\)U)if 7 kJ of heat energy is evolved by the system and 99 kJ of work is done on the system for a certain process?

A) 92 kJ

B) -106 kJ

C) -7 kJ

D) -92 kJ

E) 106 kJ

Q3) Internal energy and enthalpy are state functions.What is meant by this statement?

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Chapter 7: The Structure of Atoms

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

1)A paramagnetic substance is attracted to a magnetic field.

2)An atom with no unpaired electrons is ferromagnetic.

3)Atoms with one or more unpaired electrons are paramagnetic.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 3

E) 1,2,and 3

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

1)A diamagnetic substance is strongly attracted to a magnetic field.

2)Substances that retain their magnetism after they are withdrawn from a magnetic field are called ferromagnetic.

3)Most transition metals and all lanthanide metals are ferromagnetic.

A) 1 only

B) 2 only

C) 3 only

D) 1 and 3

E) 1,2,and 3

Q3) A point in a standing wave that has zero amplitude is called a(n)________.

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Chapter 8: The Structure of Atoms and Periodic Trends

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Q1) The element ________ has the following electron configuration: [Rn]5f¹²7s².

Q2) What is the ground state electron configuration for Cr³⁺?

A) [Ar]

B) [Ar]3d⁷4s²

C) [Ar]3d¹4s²

D) [Ar]3d²4s¹

E) [Ar]3d³

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

A) 2s²

B) 3s²

C) 1s²

D) 4d³

E) 1s⁶

Q4) ________ rule states that the most stable arrangement of electrons is that which contains the maximum number of unpaired electrons,all with the same spin direction.

Q5) Electron ________ is defined as the energy change for a process in which a gas phase atom acquires an electron.

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Chapter 9: Bonding and Molecular Structure

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Q1) What are the approximate H-N-H bond angles in NH₄⁺?

A) 109.5S1U1P1\(\circ\)S1S1P0

B) 120S1U1P1\(\circ\)S1S1P0

C) 109.5S1U1P1\(\circ\)S1S1P0 and 120S1U1P1\(\circ\)S1S1P0

D) 90S1U1P1\(\circ\)S1S1P0

E) 90S1U1P1\(\circ\)S1S1P0 and 120S1U1P1\(\circ\)S1S1P0

Q2) Linus Pauling noticed that the energy of a polar bond is often greater than expected.He attributed the greater bond energy to

A) a coulombic attraction between atoms with partially positive and negative charges.

B) the greater bond lengths of the heteronuclear bonds.

C) one of the many unexplainable phenomena that scientists encounter.

D) the ability of heteronuclear species to form double and triple bonds.

E) the greater number of valence electrons found in heteronuclear molecules.

Q3) If a molecule has a positive and negative end,the molecule is said to have a(n)________ moment.

Q4) The molecular geometry of a molecule whose central atom has four single bonds and one lone pair of electrons is ________.

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11

Chapter 10: Bonding and Molecular Structure: Orbital

Hybridization and Molecular Orbitals

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Q1) Refer to Diagram 9-1.Consider the molecules B₂,C₂,N₂ and F₂.Which two molecules have the same bond order?

A) B₂ and C₂

B) B₂ and F₂

C) C₂ and N₂

D) C₂ and F₂

E) N₂ and F₂

Q2) Refer to diagram 9-1.Identify the molecule with the shortest bond length.

A) O₂

B) C₂

C) B₂

D) F₂

E) N₂

Q3) Draw a Lewis structure of xenon trioxide.What is the hybridization of the xenon atom in this molecule?

Q4) In molecular orbital theory,the bond order is defined as 1/2 \(\times\) (the number of electrons in ________ orbitals minus the number of electrons in antibonding orbitals).

12

Q5) The hybridization of the carbon atom in CF₃⁺ is

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Chapter 11: Gases and Their Properties

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Q1) A statement of Avogadro's hypothesis is that equal volumes of gases under the same conditions of pressure and temperature will contain equal numbers of

Q2) When 0.5000 grams of an unknown hydrocarbon,C_xH_y,is completely combusted with excess oxygen,1.037 L CO₂ gas and is produced at 98.3 S1U1P1\(\circ\)S1S1P0C and 1.000 atm.What is the empirical formula of the hydrocarbon? (R = 0.08206 L.atm/mol.K)

A) CH

B) CH₂

C) C₂H₃

D) C₃H₅

E) C₃H₈

Q3) A gaseous mixture containing 1.5 mol Ar and 3.5 mol CO₂ has a total pressure of 7.3 atm.What is the partial pressure of CO₂?

A) 2.2 atm

B) 1.4 atm

C) 17 atm

D) 5.1 atm

E) 7.3 atm

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Page 14

Chapter 12: Intermolecular Forces and Liquids

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Q1) Above its critical temperature and pressure,a substance becomes a(n)________ fluid.

Q2) The molecules in solid PH₃ are attracted to each other by ___.

A) London dispersion forces

B) hydrogen-bonding

C) dipole-dipole interactions

D) London dispersion forces and dipole-dipole interactions

E) London dispersion forces and hydrogen-bonding

Q3) Which one of the following molecules will exhibit dipole-dipole intermolecular forces as a pure liquid or solid?

A) CS₂

B) C₂H₂

C) CCl₄

D) F₂

E) PCl₃

Q4) The ________ equation relates the equilibrium vapor pressure of a volatile liquid to the molar enthalpy of vaporization at a given temperature.

Q5) Which ion,K⁺ or Ca²⁺,is expected to have the more negative enthalpy of hydration? Why?

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Chapter 13: The Chemistry of Solids

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Q1) Elements that have their highest energy electrons in a \(\textbf{ filled }\) band of molecular orbitals that is separated from the lowest empty band by an energy difference much too large for electrons to jump between bands are called ____.

A) semiconductors

B) metals

C) conductors

D) insulators

E) isomorphs

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) A salt with a 1:1 ratio of anions to cations may pack in a face-centered cubic unit cell with the anions at the lattice points and the cations occupying one-half of the ________ holes.Zinc sulfide is an example of this structure.

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Chapter 14: Solutions and Their Behavior

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Q1) A solution in which there is more dissolved solute than in a saturated solution is known as a(n)________ solution.

Q2) What is the mole fraction of urea,CH₄N₂O,in an aqueous solution that is 54% urea by mass?

A) 0.26

B) 0.74

C) 0.54

D) 0.37

E) 0.80

Q3) What concentration unit is used in the calculation of osmotic pressure for a dilute solution?

A) molality

B) weight percent

C) mass fraction

D) mole fraction

E) molarity

Q4) If an egg's shell is carefully dissolved using an acid,the egg white and yolk will remain intact inside a membrane.If the egg is then placed in distilled water,it will slowly expand until it bursts.Why?

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Chapter 15: Chemical Kinetics: the Rates of Chemical Reactions

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Q1) For a certain overall third-order reaction with the general form aA \(\to\) products,the initial rate of reaction is 0.42 M·s^-1 when the initial concentration of the reactant is 0.25 M.What is the rate constant for this reaction?

A) 0.0093 M^-2.s^-1

B) 27 M^-2.s^-1

C) 0.42 M^-2.s^-1

D) 0.15 M^-2.s^-1

E) 110 M^-2.s^-1

Q2) For the first-order decomposition of N₂O₅ at a high temperature,determine the rate constant if the N₂O₅ concentration decreases from 1.04 M to 0.62 M in 375 seconds.

A) 5.99 \(\times\) 10^-4 s^-1

B) 1.59 \(\times\) 10^-3 s^-1

C) 1.74 \(\times\) 10^-3 s^-1

D) 1.38 \(\times\) 10^-3 s^-1

E) 1.94 \(\times\) 10² s^-1

Q3) The pre-exponential,A,in the Arrhenius equation is called the ________ factor.

Q4) Termolecular elementary steps are rare.Why?

Q5) Radioactive isotopes decay by ________-order kinetics.

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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)Product concentrations appear in the numerator of an equilibrium constant expression.

2)A reaction favors the formation of products if K >> 1.

3)Stoichiometric coefficients are used as exponents in an equilibrium constant expression.

A) 1 only

B) 2 only

C) 3 only

D) 2 and 3

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) The symbol Q is called the ________.

Q4) In 1913,the Haber-Bosch process was patented.The product of the Haber-Bosch process is ________.

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19

Chapter 17: The Chemistry of Acids and Bases

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Q1) What volume of water must be added to 14.8 mL of a pH 2.0 solution of HNO₃ in order to change the pH to 4.0?

A) 14.8 mL

B) 147 mL

C) 1.47 \(\times\) 10³mL

D) 37 mL

E) 85 mL

Q2) When a Lewis acid combines with a Lewis base,the base supplies both electrons to the bond.This type of chemical bond is called a(n)________ covalent bond.

Q3) At 25 S1U1P1\(\circ\)S1S1P0C,all of the following ions produce an acidic solution,except ____.

A) NH₄⁺

B) HSO₃^-

C) HPO₄^2-

D) [Fe(H₂O)₆]³⁺

E) [Al(H₂O)₆]³⁺

Q4) A molecule that can behave as either a Brønsted-Lowry acid or base is termed

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

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Q1) Consider the titration of 300.0 mL of 0.569 M NH₃ (K_b = 1.8 \(\times\) 10^-5)with 0.500 M HNO₃.After 150.0 mL HNO₃ has been added,what is the pH of the solution?

A) 4.64

B) 9.36

C) 6.36

D) 11.36

E) 7.00

Q2) A buffer is composed of 0.400 mol H₂PO₄^- and 0.400 mol HPO₄^2- diluted with water to a volume of 1.00

L.The pH of the buffer is 7.210.How many moles of HCl must be added to decrease the pH to 6.210?

A) 0.200 mol

B) 0.327 mol

C) 0.360 mol

D) 0.400 mol

E) 3.60 mol

Q3) For a monoprotic acid titration,the ________ point in a titration is the point where the moles of strong base added equals the moles of acid initially present.

Page 21

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Chapter 19: Entropy and Free Energy

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Q1) At what temperature (in kelvin units)is the entropy of a pure crystal 0.0 J/K.

Q2) For a chemical system,\(\Delta\)_rGS1U1P1\(\circ\)S1S1P0 and \(\Delta\)_rG are equal when

A) the system is in equilibrium.

B) the reactants and products are in standard state conditions.

C) the equilibrium constant,K,equals 0.

D) the reaction quotient,Q,is less than 1.

E) the reactants and products are in the gas phase.

Q3) What is the sign of \(\Delta\)Hº(system)and \(\Delta\)Sº(system)if a chemical reaction is nonspontaneous at all temperatures under standard conditions?

A) (\(\Delta\)Hº)(system)is negative,and \(\Delta\)Sº(system)is positive.

B) (\(\Delta\)Hº)(system)is negative,and \(\Delta\)Sº(system)is negative.

C) (\(\Delta\)Hº)(system)is positive,and \(\Delta\)Sº(system)is positive.

D) (\(\Delta\)Hº)(system)is positive,and \(\Delta\)Sº(system)is negative.

E) none of these

Q4) ________ changes only occur in the direction that leads toward chemical equilibrium.

Q5) In any chemical process,energy must be conserved.This is a statement of the ________ law of thermodynamics.

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Chapter 20: Principles of Reactivity: Electron Transfer Reactions

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Q1) Al³⁺ is reduced to Al(s)at an electrode.If a current of 2.75 ampere is passed for 36 hours,what mass of aluminum is deposited at the electrode? Assume 100% current efficiency.

A) 9.2 \(\times\) 10^-3 g

B) 3.3 \(\times\) 10¹ g

C) 9.9 \(\times\) 10¹ g

D) 1.0 \(\times\) 10² g

E) 3.0 \(\times\) 10² g

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

Q3) Explain the function of a salt bridge in a voltaic cell.

Q4) The standard cell potential of the following electrochemical cell is 0.19 V.Pt | Sn⁴⁺(aq,1.0 M),Sn²⁺(aq,1.0 M)|| Cu²⁺(aq,0.200 M)| Cu

Which factor will increase the measured cell potential of the galvanic cell?

A) switching from a platinum to a graphite anode

B) increasing the size of the anode

C) decreasing the concentration of Cu²⁺

D) increasing the concentration of Sn⁴⁺

E) decreasing the temperature of the cell

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Chapter 21: Environmental Chemistry Earths

Environment,energy,and Sustainability

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Q1) What form of radiation is trapped by greenhouse gases?

A) microwaves

B) UV radiation

C) visible radiation

D) infrared radiation

E) radio waves

Q2) The pH of the oceans in the 18th century was 8. 2.What would the pH be if the hydronium concentration of the oceans was doubled?

A) 16.4

B) 9.28.0

C) 8.0

D) 7.9

E) 4.1

Q3) What is the correct name for NO₃^-?

A) nitric oxide

B) nitrite

C) nitrogen trioxide

D) nitrous oxide

E) nitrate

Q4) 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) Dinitrogen monoxide,or laughing gas,can be made by decomposing ammonium nitrate at 250 S1U1P1\(\circ\)S1S1P0C.Write a balanced equation for this reaction.

A) NH₄NO₃(s)\(\to\) N₂O(g)+ 2 H₂(g)+ O₂(g)

B) NH₄NO₃(s)+ 2 N₂(g)\(\to\) 3 N₂O(g)+ 2 H₂(g)

C) NH₄NO₃(s)\(\to\) N₂O(g)+ 2 H₂O(g)

D) 2

NH₄NO₃(s)\(\to\) N₂O(g)+ 2 HNO₂(g)+ H₂O(g)+ 2 H₂(g)

E) 3

NH₄NO₃(s)\(\to\) N₂O(g)+ 2 N₂(g)+ 8 H₂O(g)

Q2) Give two examples of the allotropes of elemental oxygen.

Q3) The primary product of the reaction of potassium and oxygen is ________.This compound is used in oxygen generators.

Q4) A(n)________ reaction is one where an element or compound is simultaneously oxidized and reduced.

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Chapter 23: The Chemistry of the Transition Elements

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Q1) Which of the following d-block ion electron configurations may be high spin \(\underline{\text{ or }}\) low spin when placed in an \(\underline{\text{ octahedral ligand}}\) field?

A) [Ar]3d⁸

B) [Ar]3d³

C) [Ar]3d⁹

D) [Ar]3d⁴

E) [Ar]3d²

Q2) In which of the following complexes does the transition metal have a d⁸ configuration?

A) Cu(H₂O)₆²⁺

B) Fe(CN)₆^3-

C) Ni(CO)₄

D) PtCl₄^2-

E) Zn(NH₃)₄²⁺

Q3) Phenanthroline (C₁₂H₈N₂)can bind to metal ions using both its ammine groups.It is referred to as a(n)________ ligand,meaning it is "two-toothed."

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Chapter 24: Carbon: Not Just Another Element

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Q1) A peptide bond is the amide linkage that is formed in a condensation reaction involving the __________ group of one amino acid with the carboxylic acid group of a second amino acid.

Q2) The addition of Br₂ is used as the reaction to distinguish between alkanes and alkenes.What is the observation that accompanies this test?

A) A bright red color is produced when Br₂reacts with an alkene.

B) Bromine dissolves in alkanes but not in alkenes.

C) Bromine is a dark red liquid.When it adds to the double bond of an alkene to make the dibromide,it becomes colorless.

D) The red color of bromine disappears when it dissolves in alkanes.

E) Bromine reacts with alkanes to form a precipitate.

Q3) What is the product of the addition of excess F₂ to acetylene?

A) 1,2-difluoroethane

B) ethylene

C) 1,1,2,2-tetrafluoroethane

D) ethane

E) 1,1-difluoroethylene

Q4) The process by which long chain hydrocarbons in petroleum are shortened is called

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Chapter 25: Biochemistry

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54 Verified Questions

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Sample Questions

Q1) Which of the following concerning protein structure is not correct?

A) The primary structure of a protein is the sequence of amino acids linked together by peptide bonds.

B) Hydrogen bonds between amide linkages in the protein backbone define protein secondary structure.

C) Helices and sheets are common tertiary structures in proteins.

D) How a protein folds into its overall 3-dimensional shape is referred to as the protein tertiary structure.

E) In cases where more than one polypeptide chain makes up the protein,the interaction of the different chains is referred to as the quaternary structure.

Q2) Which of the following is not a polysaccharide?

A) glycogen

B) lactose

C) cellulose

D) amylose starch

E) amylopectin starch

Q3) What is the role of tRNA in protein synthesis?

Q4) The oxygen that bridges two sugars in a disaccharide is known as a ___ bond.

Q5) The amide linkages in a protein and also called ___ bonds.

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Chapter 26: Nuclear Chemistry

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Sample Questions

Q1) All of the following statements concerning nuclei are true EXCEPT

A) only hydrogen-1 and helium-3 have more protons than neutrons.

B) from He to Ca,stable nuclei have roughly equal numbers of protons and neutrons.

C) elements with odd atomic numbers have more stable isotopes than do those with even atomic numbers.

D) the neutron to proton ratio in stable nuclei increases as mass increases.

E) more stable isotopes have an even number of neutrons than an odd number.

Q2) Explain the treatment named boron neutron capture therapy (BNCT).

Q3) If a nucleus undergoes positron particle emission

A) its atomic number decreases by four and its mass number decreases by two.

B) its atomic number decreases by two and its mass number decreases by four.

C) its atomic number increases by one and its mass number is unchanged.

D) its atomic number decreases by one and its mass number is unchanged.

E) its atomic number is unchanged and its mass number decreases by one.

Q4) Technetium-99m is routinely used in medical imaging.The italics m means the nucleus is ________.

Q5) A unit used to quantify biological damage is called the ________ or rem.

Q6) All isotopes of atomic number greater than ________ are unstable and radioactive.

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