Physical Science Exam Review - 3207 Verified Questions

Physical Science

Exam Review

Course Introduction

Physical Science is an interdisciplinary course that introduces students to the fundamental principles of physics and chemistry. Topics include the structure of matter, forms of energy, force and motion, properties of substances, and the interactions of matter and energy in various physical environments. Emphasis is placed on developing scientific inquiry skills through laboratory experiments, real-world applications, and problem-solving activities, preparing students to understand and appreciate the scientific processes that govern everyday phenomena.

Recommended Textbook Chemistry A Molecular Approach 2nd Canadian Edition by Nivaldo J. Tro

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25 Chapters

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Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) How many square kilometres (km²)are in 246 square miles? (1 km = 0.621371 miles)

A)153 km²

B)481 km²

C)396 km²

D)95.0 km²

E)637 km²

Answer: E

Q2) What answer should be reported,with the correct number of significant figures,for the following calculation?

6.83192458 + log(6.7831 × 10⁴)

A)11.66335

B)11.663353

C)11.6634

D)11.6633528

E)11.663

Answer: A

Q3) Define matter.

Answer: Matter is anything that occupies space and has mass.

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) If 102.7 g of tin is mixed with 10.9 g of copper to create the alloy pewter,how many metal atoms total are present in the pewter?

A)5.81 × 10²³ metal atoms

B)5.22 × 10²³ metal atoms

C)3.75 × 10²³ metal atoms

D)9.66 × 10²³ metal atoms

E)6.24 × 10²³ metal atoms

Answer: E

Q2) Dalton's atomic theory states

A)that all elements have several isotopes.

B)that matter is composed of small indestructible particles.

C)that the properties of matter are determined by the properties of atoms.

D)that energy is neither created nor destroyed during a chemical reaction.

E)that an atom is predominantly empty space.

Answer: B

Q3) Are anions typically larger or smaller than their corresponding atom? Why?

Answer: Anions are larger than their corresponding atom because the anion contains more electrons than the atom.Since electrons repel one another and determine the size of the atom or ion,adding electrons to the atom to form an anion makes it larger.

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Page 4

Chapter 3: Molecules,compounds,and Nomenclature

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Sample Questions

Q1) Identify the cyanide ion.

A) \(\mathrm { MnO } _ { 4 }^-\)

B)H \( \mathrm{CO}_{3}^{-} \)

C) \(\mathrm { CO } _ { 3 } ^ { 2 - }\)

D)HCN

E) \(\mathrm { CN } ^ { - }\)

Answer: E

Q2) The compound Cu(NO₂ )₂,is named

A)copper nitrite(II).

B)copper(I)nitrite.

C)copper(I)nitrite(II).

D)copper(II)nitrite.

Answer: D

Q3) In which set do all elements tend to form anions in binary ionic compounds?

A)C,S,Pb

B)K,Fe,Br

C)Li,Na,K

D)N,O,I

Answer: D

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Chapter 4: Chemical Reactions and Stoichiometry

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Q1) Lithium and nitrogen react to produce lithium nitride: 6Li(s)+ \(\mathrm { N } _ { 2 }\) (g)? \(2 \mathrm { Li } _ { 3 }\) N(s)

How many moles of \(\mathrm { N } _ { 2 }\) are needed to react with 0.400 mol of lithium?

A)2.40

B)0.400

C)0.133

D)1.2

E)0.0667

Q2) HBr,HI,HNO₃,NaBr,and KNO₃ are all classified as A)acids.

B)nonelectrolytes.

C)strong electrolytes.

D)weak electrolytes.

Q3) Which one of the following compounds is insoluble in water?

A)Na₂SO₄

B)KNO₃

C)BaSO₄

D)Li₂CO₃

Page 6

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Chapter 5: Gases

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Sample Questions

Q1) How many molecules of N₂ are in a 400.0 mL container at 1.0399 bar and 135 °C?

A)7.01 × 10^{21 molecules}

B)7.38 × 10^{21 molecules}

C)2.12 × 10^{22 molecules}

D)2.23 × 10^{22 molecules}

Q2) Define hypoxia.

A)oxygen starvation

B)increased oxygen concentration in body tissues

C)increased nitrogen concentration in body tissues and fluids

D)nitrogen starvation

Q3) At what temperature does argon have a root mean square velocity of 492 m s^-1?

A)321 K

B)291 K

C)340 K

D)388 K

E)409 K

Q4) Define effusion.

Q5) Define pressure.

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Chapter 6: Thermochemistry

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Q1) Calculate the amount (mass)of acetic acid (C₂H₄O₂,molar mass = 60.052 g mol^-1)that causes a bomb calorimeter with a heat capacity of 8.43 kJ °C^-1 to have a temperature increase from 24.5 °C to 36.8 °C.The D_rU for the combustion of acetic acid is -874.2 kJ mol^-1.

A)6.18 g

B)7.12 g

C)2.18 g

D)9.66 g

E)8.68 g

Q2) How much heat is absorbed/released when 40.00 g of NH₃(g)reacts in the presence of excess \(\mathrm { O } _ { 2 }\) (g)to produce NO(g)and H₂O(l)according to the following chemical equation? 4NH₃(g)+ 5O₂(g)? 4NO(g)+ 6H₂O(l) \(\Delta _ { \mathrm { r } }\) H° = 1168 kJ

A)685.8 kJ of heat are absorbed.

B)685.8 kJ of heat are released.

C)2743 kJ of heat are absorbed.

D)2743 kJ of heat are released.

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Page 8

Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Food can be cooked by ________ radiation.

A)ultraviolet

B)gamma

C)microwave

D)X-ray

E)radio

Q2) Identify the ground-state electron configuration for Se²^-.

A)[Ar]4s²3d¹⁰4p⁴

B)[Ar]4s²3d¹⁰4p²

C)[Ar]4s²4p⁶

D)[Ar]4s²3d¹⁰4p⁶

E)[Ar]4s²3d⁸4p⁶

Q3) Determine the longest wavelength of light required to remove an electron from a sample of potassium metal if the binding energy for an electron in K is 1.76 × 10³ kJ mol^-1.

A)147 nm

B)68.0 nm

C)113 nm

D)885 nm

E)387 nm

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Chapter 8: Periodic Properties of the Elements

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Q1) Use the periodic table to determine the ground-state electron configuration for the following element: Zr

A)[Kr]5s²4d¹⁰5p²

B)[Kr]5s²5p²

C)[Kr]5s²5d²

D)[Kr]5s²4d²

E)[Kr]5s²4d¹⁰5p⁴

Q2) Using the periodic table,identify the element with the following exception electron configuration: [Kr]5s¹4d⁷

A)Mo

B)Tc

C)Rh

D)Ru

E)Pd

Q3) Choose the statement that is TRUE.

A)Outer electrons efficiently shield one another from nuclear charge.

B)Core electrons effectively shield outer electrons from nuclear charge.

C)Valence electrons are the most difficult of all electrons to remove.

D)Core electrons are the easiest of all electrons to remove.

Q4) Define ionization energy.

Page 10

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Give the number of valence electrons for SF₄.

A)28

B)30

C)32

D)34

Q2) Choose the bond below that is most polar.

A)C-N

B)C-F

C)C-O

D)C-C

E)F-F

Q3) Describe a covalent bond.

Q4) Place the following in order of decreasing magnitude of lattice energy. NaF \(\quad\)RbBr \(\quad\)KCl

A)RbBr > NaF > KCl

B)NaF > KCl > RbBr

C)KCl > NaF > RbBr

D)NaF > RbBr > KCl

E)RbBr > KCl > NaF

Q5) Define formal charge.

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Chapter 10: Chemical Bonding II: Molecular Shapes, valence

Bond Theory, and Molecular Orbital Theory

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Sample Questions

Q1) Determine the electron geometry (eg)and molecular geometry (mg)of SiF₄.

A)eg = tetrahedral,mg = trigonal pyramidal

B)eg = octahedral,mg = square planar

C)eg = trigonal bipyramidal,mg = trigonal pyramidal

D)eg = tetrahedral,mg = bent

E)eg = tetrahedral,mg = tetrahedral

Q2) Determine the electron geometry (eg),molecular geometry (mg),and polarity of TeCl₆.

A)eg = octahedral,mg = octahedral,nonpolar

B)eg = trigonal bipyramidal,mg = trigonal bipyramidal,nonpolar

C)eg = octahedral,mg = square planar,polar

D)eg = trigonal bipyramidal,mg = seesaw,polar

E)eg = tetrahedral,mg = trigonal pyramidal,polar

Q3) List the number of sigma bonds and pi bonds in a single bond.

A)1 sigma,0 pi

B)0 sigma,1 pi

C)1 sigma,1 pi

D)1 sigma,2 pi

Q4) Explain why oil and water do not mix.

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Chapter 11: Liquids, solids, and Intermolecular Forces

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Sample Questions

Q1) How much energy is required to heat 87.1 g acetone (molar mass = 58.08 g mol^-1)from a solid at -154.0 °C to a liquid at -42.0°C? The following physical data may be useful: _fusH = 7.27 kJ mol^-1

C_liq = 2.16 J g^-1 °C^-1

C_gas= 1.29 J g^-1 °C^-1

C_sol = 1.65 J g^-1 °C^-1

T_melting= -95.0 °C

A)8.48 kJ

B)18.5 kJ

C)32.2 kJ

D)29.4 kJ

E)9.97 kJ

Q2) Choose the substance with the highest vapour pressure at a given temperature.

A)SiS₂

B)RbCl

C)CH₃SCH₃

D)BF₃

E)SbH₃

Q3) Define volatile.

Q4) Define the boiling point of a liquid.

Page 13

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Chapter 12: Solutions

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Sample Questions

Q1) A solid can be purified through which technique?

A)recrystallization

B)dilution

C)dissolution

D)boiling

E)melting

Q2) To make a 2.00 mol kg^-1 solution,one could take 2.00 moles of solute and add

A)1.00 L of solvent.

B)1.00 kg of solvent.

C)enough solvent to make 1.00 L of solution.

D)enough solvent to make 1.00 kg of solution.

Q3) A compound is found to have a molar mass of 598 g mol^-1.If 35.8 mg of the compound is dissolved in enough water to make 175 mL of solution at 25 °C,what is the osmotic pressure of the resulting solution?

A)3.42 × 10^-3 bar

B)8.48 × 10^-3 bar

C)5.01 × 10^-3 bar

D)5.99 × 10^-3 bar

E)8.06 × 10^-3 bar

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) Given the following balanced equation,determine the rate of reaction with respect to [SO₃].If the rate of SO₂ loss is 1.19 × 10^-3 mol L^-1 s^-¹,what is the rate of formation of SO₃?

2SO₂(g)+ O₂(g) 2SO₃(g)

A)3.56 × 10^-3 mol L^-1 s^-1

B)1.19 × 10^-3 mol L^-1 s^-1

C)1.78 × 10^-3 mol L^-1 s^-1

D)1.42 × 10^-2 mol L^-1 s^-1

E)7.12 × 10^-3 mol L^-1 s^-1

Q2) The rate constant for the first-order decomposition of N₂O is 3.40 s^-1.What is the half-life of the decomposition?

A)0.491 s

B)0.204 s

C)0.236 s

D)0.424 s

E)0.294 s

Q3) Explain what the exponential factor in the Arrhenius equation represents.

Q4) What is the difference between average reaction rate and instantaneous reaction rate?

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction at equilibrium.What effect will increasing the temperature have on the system?

Fe₃O₄(s)+ CO(g) 3FeO(s)+ CO₂(g) _rH°= +35.9 kJ

A)The reaction will shift in the direction of reactants.

B)The equilibrium constant will increase.

C)The equilibrium constant will decrease.

D)No effect will be observed.

E)The reaction will shift in the direction of products.

Q2) Consider the following reaction: NO(g)+ SO₃(g) NO₂(g)+ SO₂(g)

A reaction mixture initially contains 2.0 bar NO₂ and 1.0 bar

SO₂.Determine the equilibrium pressure of SO₂ if K_p for the reaction at this temperature is 0.0118.

A)0.011 bar

B)0.037 bar

C)0.10 bar

D)0.0038 bar

E)0.067 bar

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Page 16

Chapter 15: Acids and Bases

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Sample Questions

Q1) Calculate the pH for an aqueous solution of acetic acid that contains \(2.15 \times 10 - 3 \mathrm {~mol} \mathrm {~L} ^ { - 1 }\) hydronium ion.

A)4.65 × 10^-12

B)2.15 × 10^-3

C)2.67

D)11.33

Q2) Calculate the pH of a 0.60 mol L^-1 H₂SO₃ solution that has the stepwise dissociation constants K_a1 = 1.5 × 10^-2 and \(K _ { \mathrm { a } 2 } = 6.3 \times 10 ^{- 8}\)

A)1.02

B)1.06

C)1.82

D)2.04

Q3) Describe a molecule that can be a Lewis acid.

Q4) Describe the relationship between molecular structure and acid strength.

Q5) What is the autoionization of water?

Q6) What does the term amphoteric mean?

Q7) What is the difference between a strong and weak acid?

Page 17

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Which of the following solutions is a good buffer system?

A)a solution that is 0.10 mol L^-1 NaCl and 0.10 mol L^-1 HCl

B)a solution that is 0.10 mol L^-1 HCN and 0.10 mol L^-1 LiCN

C)a solution that is 0.10 mol L^-1 NaOH and 0.10 mol L^-1 HNO₃

D)a solution that is 0.10 mol L^-1 HNO₃ and 0.10 mol L^-1 KNO₃

E)a solution that is 0.10 mol L^-1 HCN and 0.10 mol L^-1 NaCl

Q2) Which of the following is the correct equation relating Q to K_sp for an unsaturated solution?

A)Q > K_sp

B)Q < K_sp

C)Q = K_sp

D)Q K_sp

E)Q K_sp

Q3) Name the chemical compound that forms stalactites and stalagmites.

Q4) Explain the common ion effect with respect to molar solubility.

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Chapter 17: Gibbs Energy and Thermodynamics

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Q1) Consider the following reaction at constant pressure.Use the information here to determine the value of S_surrat 298 K.Predict whether or not this reaction will be spontaneous at this temperature. \(\quad\)\(\quad\)N₂(g)+

2O₂(g) 2NO₂(g) _rH = +66.4 kJ

A) S_surr = +223 J K^-1 mol^-1,reaction is spontaneous

B) S_surr = -223J K^-1 mol^-1,reaction is not spontaneous

C) S_surr = -66.4 J K^-1 mol^-1,reaction is spontaneous

D) S_surr = +66.4 kJ K^-1 mol^-1,reaction is not spontaneous

E) S_surr = -66.4 J K^-1 mol^-1,reaction is not spontaneous

Q2) Why is heating your home with gas more efficient than heating it with electricity?

Q3) Give the standard states for a gas,liquid,solid,and solution.

Q4) Why can endothermic reactions be spontaneous?

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Chapter 18: Electrochemistry

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Sample Questions

Q1) What is the balanced equation for the galvanic cell reaction expressed using shorthand notation below?

\(\quad\)\(\quad\)\(\quad\)Ni(s)| Ni²⁺(aq)|| Cl₂(g)| Cl^-(aq)| C(s)

A)Ni(s)+ 2Cl^-(aq) Ni²⁺(aq)+ Cl₂(g)

B)Ni(s)+ Cl₂(g) Ni²⁺(aq)+ 2Cl^-(aq)

C)Ni²⁺(aq)+ 2Cl^-(aq) Ni(s)+ Cl₂(g)

D)Ni²⁺(aq)+ 2Cl^-(aq) NiCl₂(s)

Q2) For the galvanic cell reaction,expressed below using shorthand notation,what half-reaction occurs at the cathode?

\(\quad\)\(\quad\)Zn(s)| Zn²⁺(aq)|| Ni²⁺(aq)| Ni(s)

A)Zn(s) Zn²⁺(aq)+ 2 e^-

B)Zn²⁺(aq)+ 2 e^- Zn(s)

C)Ni(s) Ni²⁺(aq)+ 2 e^-

D)Ni²⁺(aq)+ 2 e^- Ni(s)

Q3) Why,if we multiply a reaction by 2,don't we multiply its E°_red by 2?

Q4) What is electrolysis?

Q5) Why are iron nails coated with zinc?

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Complete the following equation of transmutation.

\(\begin{array} { l }

27 \\

13

\end{array}\) Al + 4

2 He ? \(\begin{array} { l }

30 \\

15

\end{array}\) P + ________

A) \({ } _ { + 1 } ^ { 0 } \mathrm { e }\)

B) \({ } _ { 0 } ^ { 0 } \mathrm {~g}\)

C) \({ } _ { - 1 } ^ { 0 } \mathrm { e }\)

D) \({ } _ { 1 } ^ { 1 } \mathrm { H }\)

E) \({ } _ { 0 } ^ { 1 } \mathrm { n }\)

Q2) Explain the concept of "magic numbers."

Q3) Describe what is meant by the "valley of stability."

Q4) How does a dosimeter measure exposure to radioactivity?

Q5) Define chain reaction in terms of the fission of uranium nucleus.

Q6) What is the mass defect?

Q7) Define radioactivity.

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Chapter 20: Organic Chemistry I: Structures

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Q1) What is the difference between the R/ S designation and d/l designation of chiral compounds?

Q2) Which compound is a saturated hydrocarbon?

A)1-butyne

B)acetylene

C)3-methylheptane

D)2-methylheptene

E)ethene

Q3) The general formula for alkanes is C_nH₂_n₊₂.What is the value of n in 3-ethyl-2-methylpentane ? Is this molecule chiral?

A)n = 7; yes,it is chiral

B)n = 7; no,it is not chiral

C)n = 8; yes,it is chiral

D)n = 8; no,it is not chiral

E)n = 5; yes,it is chiral

Q4) Why are there so many more carbon compounds than the number of compounds made up of all the rest of the elements combined?

Q5) What is meant by the term,"structural isomer"?

Q6) What are constitutional isomers?

Page 22

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Chapter 21: Organic Chemistry II: Reactions

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Q1) Which of the following is the balanced chemical equation for the catalytic hydrogenation (addition of H₂)to CH₃C CH?

A)CH₃C CH + H₂ CH₃CH=CH₂

B)CH₃C CH + 2H₂ CH₃CH=CH₂

C)CH₃C CH + H₂

CH₃CH₂CH₃

D)CH₃C CH + 2H₂

CH₃CH₂CH₃

E)CH₃C CH + H₂ CH₂=C=CH₂

Q2) What is the product obtained from the reaction between a ketone and a primary alcohol?

A)a secondary alcohol

B)a tertiary alcohol

C)a hemiacetal

D)a hemiketal

E)a carboxylic acid

Q3) Draw a clear mechanism for S_N2 substitution reactions.Make sure to indicate each step of the mechanism (if there is more than one)as well as to use arrows to show the electron flow.

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Chapter 22: Biochemistry

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Q1) How many chromosomes are in the nuclei of human cells?

A)42

B)44

C)46

D)48

E)52

Q2) Which of the following shows all of the tripeptides that can be formed from one molecule each of glycine (Gly),valine (Val),and leucine (Leu)?

A)GlyValLeu,GlyLeuVal,ValLeuGly,ValGlyLeu,LeuGlyVal,LeuValGly

B)GlyValLeu,GlyLeuVal,LeuGlyVal

C)ValGlyLeu,GlyValLeu,GlyLeuVal,LeuGlyVal

D)ValGlyLeu and GlyLeuVal

E)GlyValLeu

Q3) Which one of the following amino acids contains a hydrophobic side chain?

A)valine

B)histidine

C)glutamine

D)glutamic acid

Q4) What is a codon?

Q5) How do cis- and trans-fats differ?

Page 24

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Chapter 23: Chemistry of the Nonmetals

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Q1) What is the name of powdered form of carbon,which is a component of soot?

A)diamond

B)graphite

C)fullerenes

D)bucky balls

E)carbon black

Q2) Which of the following statements is TRUE?

A)All silicates are composed of SiO₄ tetrahedra that bond together through an oxygen to form single chains.

B)The fibrous nature of asbestos comes from its silicate network structure.

C)In some silicates,ionic bonding between a metal ion and an Si_xO_y^z^- ion is what holds the structure together.

D)The flakiness of mica comes from its single silicate chain structure.

E)Silicates are easily soluble in water because they contain silicate anions in their structure.

Q3) Describe the major production method for obtaining oxygen.

Q4) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q5) Why are phosphate compounds added to detergents?

Page 25

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Chapter 24: Metals and Metallurgy

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Q1) Hydrometallurgy is

A)refining of metal ores using oxidation-reduction reactions.

B)refining of metal ores using heat.

C)refining of metal ores using reactions with aqueous solutions.

D)forming metal parts using heat and small crystals of metal.

E)forming tiny metal crystals using heat and water spray.

Q2) What is an advantage of hydrometallurgical processes over pyrometallurgical processes?

Q3) An interstitial alloy contains hydrogen in half of the tetrahedral holes of a closest-packed metal,M.What is the formula of this alloy?

A)MH

B)MH₄

C)MH₂

D)M₂H

E)M₄H

Q4) Describe the difference between a substitutional alloy and an interstitial alloy.

Q5) Why is zinc used to coat steel objects?

Q6) What is the difference between ferromagnetism and paramagnetism?

Q7) Why does a "two-phase" structure occur in a substitutional alloy?

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Choose the electron configuration for Cr³⁺.

A)[Ne]3s²3p⁶

B)[Ar]4s¹3d²

C)[Ar]4s²3d¹

D)[Ar]4s²3d⁷

E)[Ar]3d³

Q2) How many moles of aqueous ions will be produced from the dissolution of 1.0 mole of K₃[FeCl₆] in water?

A)4

B)10

Q3) How many unpaired electrons would you expect for the complex ion [MnF₆]^4-? A)1 B)0 C)5 D)2 E)3

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