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Chemistry

Exercises

ChemistryExercises
Table of contents
  1. Preface
  2. 1 Essential Ideas
    1. Introduction
    2. 1.1 Chemistry in Context
    3. 1.2 Phases and Classification of Matter
    4. 1.3 Physical and Chemical Properties
    5. 1.4 Measurements
    6. 1.5 Measurement Uncertainty, Accuracy, and Precision
    7. 1.6 Mathematical Treatment of Measurement Results
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  3. 2 Atoms, Molecules, and Ions
    1. Introduction
    2. 2.1 Early Ideas in Atomic Theory
    3. 2.2 Evolution of Atomic Theory
    4. 2.3 Atomic Structure and Symbolism
    5. 2.4 Chemical Formulas
    6. 2.5 The Periodic Table
    7. 2.6 Molecular and Ionic Compounds
    8. 2.7 Chemical Nomenclature
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  4. 3 Composition of Substances and Solutions
    1. Introduction
    2. 3.1 Formula Mass and the Mole Concept
    3. 3.2 Determining Empirical and Molecular Formulas
    4. 3.3 Molarity
    5. 3.4 Other Units for Solution Concentrations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  5. 4 Stoichiometry of Chemical Reactions
    1. Introduction
    2. 4.1 Writing and Balancing Chemical Equations
    3. 4.2 Classifying Chemical Reactions
    4. 4.3 Reaction Stoichiometry
    5. 4.4 Reaction Yields
    6. 4.5 Quantitative Chemical Analysis
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  6. 5 Thermochemistry
    1. Introduction
    2. 5.1 Energy Basics
    3. 5.2 Calorimetry
    4. 5.3 Enthalpy
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  7. 6 Electronic Structure and Periodic Properties of Elements
    1. Introduction
    2. 6.1 Electromagnetic Energy
    3. 6.2 The Bohr Model
    4. 6.3 Development of Quantum Theory
    5. 6.4 Electronic Structure of Atoms (Electron Configurations)
    6. 6.5 Periodic Variations in Element Properties
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  8. 7 Chemical Bonding and Molecular Geometry
    1. Introduction
    2. 7.1 Ionic Bonding
    3. 7.2 Covalent Bonding
    4. 7.3 Lewis Symbols and Structures
    5. 7.4 Formal Charges and Resonance
    6. 7.5 Strengths of Ionic and Covalent Bonds
    7. 7.6 Molecular Structure and Polarity
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  9. 8 Advanced Theories of Covalent Bonding
    1. Introduction
    2. 8.1 Valence Bond Theory
    3. 8.2 Hybrid Atomic Orbitals
    4. 8.3 Multiple Bonds
    5. 8.4 Molecular Orbital Theory
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  10. 9 Gases
    1. Introduction
    2. 9.1 Gas Pressure
    3. 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
    4. 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
    5. 9.4 Effusion and Diffusion of Gases
    6. 9.5 The Kinetic-Molecular Theory
    7. 9.6 Non-Ideal Gas Behavior
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  11. 10 Liquids and Solids
    1. Introduction
    2. 10.1 Intermolecular Forces
    3. 10.2 Properties of Liquids
    4. 10.3 Phase Transitions
    5. 10.4 Phase Diagrams
    6. 10.5 The Solid State of Matter
    7. 10.6 Lattice Structures in Crystalline Solids
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  12. 11 Solutions and Colloids
    1. Introduction
    2. 11.1 The Dissolution Process
    3. 11.2 Electrolytes
    4. 11.3 Solubility
    5. 11.4 Colligative Properties
    6. 11.5 Colloids
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  13. 12 Kinetics
    1. Introduction
    2. 12.1 Chemical Reaction Rates
    3. 12.2 Factors Affecting Reaction Rates
    4. 12.3 Rate Laws
    5. 12.4 Integrated Rate Laws
    6. 12.5 Collision Theory
    7. 12.6 Reaction Mechanisms
    8. 12.7 Catalysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  14. 13 Fundamental Equilibrium Concepts
    1. Introduction
    2. 13.1 Chemical Equilibria
    3. 13.2 Equilibrium Constants
    4. 13.3 Shifting Equilibria: Le Châtelier’s Principle
    5. 13.4 Equilibrium Calculations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  15. 14 Acid-Base Equilibria
    1. Introduction
    2. 14.1 Brønsted-Lowry Acids and Bases
    3. 14.2 pH and pOH
    4. 14.3 Relative Strengths of Acids and Bases
    5. 14.4 Hydrolysis of Salt Solutions
    6. 14.5 Polyprotic Acids
    7. 14.6 Buffers
    8. 14.7 Acid-Base Titrations
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  16. 15 Equilibria of Other Reaction Classes
    1. Introduction
    2. 15.1 Precipitation and Dissolution
    3. 15.2 Lewis Acids and Bases
    4. 15.3 Multiple Equilibria
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  17. 16 Thermodynamics
    1. Introduction
    2. 16.1 Spontaneity
    3. 16.2 Entropy
    4. 16.3 The Second and Third Laws of Thermodynamics
    5. 16.4 Free Energy
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  18. 17 Electrochemistry
    1. Introduction
    2. 17.1 Balancing Oxidation-Reduction Reactions
    3. 17.2 Galvanic Cells
    4. 17.3 Standard Reduction Potentials
    5. 17.4 The Nernst Equation
    6. 17.5 Batteries and Fuel Cells
    7. 17.6 Corrosion
    8. 17.7 Electrolysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  19. 18 Representative Metals, Metalloids, and Nonmetals
    1. Introduction
    2. 18.1 Periodicity
    3. 18.2 Occurrence and Preparation of the Representative Metals
    4. 18.3 Structure and General Properties of the Metalloids
    5. 18.4 Structure and General Properties of the Nonmetals
    6. 18.5 Occurrence, Preparation, and Compounds of Hydrogen
    7. 18.6 Occurrence, Preparation, and Properties of Carbonates
    8. 18.7 Occurrence, Preparation, and Properties of Nitrogen
    9. 18.8 Occurrence, Preparation, and Properties of Phosphorus
    10. 18.9 Occurrence, Preparation, and Compounds of Oxygen
    11. 18.10 Occurrence, Preparation, and Properties of Sulfur
    12. 18.11 Occurrence, Preparation, and Properties of Halogens
    13. 18.12 Occurrence, Preparation, and Properties of the Noble Gases
    14. Key Terms
    15. Summary
    16. Exercises
  20. 19 Transition Metals and Coordination Chemistry
    1. Introduction
    2. 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
    3. 19.2 Coordination Chemistry of Transition Metals
    4. 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
    5. Key Terms
    6. Summary
    7. Exercises
  21. 20 Organic Chemistry
    1. Introduction
    2. 20.1 Hydrocarbons
    3. 20.2 Alcohols and Ethers
    4. 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
    5. 20.4 Amines and Amides
    6. Key Terms
    7. Summary
    8. Exercises
  22. 21 Nuclear Chemistry
    1. Introduction
    2. 21.1 Nuclear Structure and Stability
    3. 21.2 Nuclear Equations
    4. 21.3 Radioactive Decay
    5. 21.4 Transmutation and Nuclear Energy
    6. 21.5 Uses of Radioisotopes
    7. 21.6 Biological Effects of Radiation
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  23. A | The Periodic Table
  24. B | Essential Mathematics
  25. C | Units and Conversion Factors
  26. D | Fundamental Physical Constants
  27. E | Water Properties
  28. F | Composition of Commercial Acids and Bases
  29. G | Standard Thermodynamic Properties for Selected Substances
  30. H | Ionization Constants of Weak Acids
  31. I | Ionization Constants of Weak Bases
  32. J | Solubility Products
  33. K | Formation Constants for Complex Ions
  34. L | Standard Electrode (Half-Cell) Potentials
  35. M | Half-Lives for Several Radioactive Isotopes
  36. Answer Key
    1. Chapter 1
    2. Chapter 2
    3. Chapter 3
    4. Chapter 4
    5. Chapter 5
    6. Chapter 6
    7. Chapter 7
    8. Chapter 8
    9. Chapter 9
    10. Chapter 10
    11. Chapter 11
    12. Chapter 12
    13. Chapter 13
    14. Chapter 14
    15. Chapter 15
    16. Chapter 16
    17. Chapter 17
    18. Chapter 18
    19. Chapter 19
    20. Chapter 20
    21. Chapter 21
  37. Index

19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds

1.

Write the electron configurations for each of the following elements:

(a) Sc

(b) Ti

(c) Cr

(d) Fe

(e) Ru

2.

Write the electron configurations for each of the following elements and its ions:

(a) Ti

(b) Ti2+

(c) Ti3+

(d) Ti4+

3.

Write the electron configurations for each of the following elements and its 3+ ions:

(a) La

(b) Sm

(c) Lu

4.

Why are the lanthanoid elements not found in nature in their elemental forms?

5.

Which of the following elements is most likely to be used to prepare La by the reduction of La2O3: Al, C, or Fe? Why?

6.

Which of the following is the strongest oxidizing agent: VO43,VO43, CrO42−,CrO42−, or MnO4−?MnO4−?

7.

Which of the following elements is most likely to form an oxide with the formula MO3: Zr, Nb, or Mo?

8.

The following reactions all occur in a blast furnace. Which of these are redox reactions?

(a) 3Fe2O3(s)+CO(g)⟶2Fe3O4(s)+CO2(g)3Fe2O3(s)+CO(g)⟶2Fe3O4(s)+CO2(g)

(b) Fe3O4(s)+CO(g)⟶3FeO(s)+CO2(g)Fe3O4(s)+CO(g)⟶3FeO(s)+CO2(g)

(c) FeO(s)+CO(g)⟶Fe(l)+CO2(g)FeO(s)+CO(g)⟶Fe(l)+CO2(g)

(d) C(s)+O2(g)⟶CO2(g)C(s)+O2(g)⟶CO2(g)

(e) C(s)+CO2(g)⟶2CO(g)C(s)+CO2(g)⟶2CO(g)

(f) CaCO3(s)⟶CaO(s)+CO2(g)CaCO3(s)⟶CaO(s)+CO2(g)

(g) CaO(s)+SiO2(s)⟶CaSiO3(l)CaO(s)+SiO2(s)⟶CaSiO3(l)

9.

Why is the formation of slag useful during the smelting of iron?

10.

Would you expect an aqueous manganese(VII) oxide solution to have a pH greater or less than 7.0? Justify your answer.

11.

Iron(II) can be oxidized to iron(III) by dichromate ion, which is reduced to chromium(III) in acid solution. A 2.5000-g sample of iron ore is dissolved and the iron converted into iron(II). Exactly 19.17 mL of 0.0100 M Na2Cr2O7 is required in the titration. What percentage of the ore sample was iron?

12.

How many cubic feet of air at a pressure of 760 torr and 0 °C is required per ton of Fe2O3 to convert that Fe2O3 into iron in a blast furnace? For this exercise, assume air is 19% oxygen by volume.

13.

Find the potentials of the following electrochemical cell:

Cd | Cd2+, M = 0.10 ‖ Ni2+, M = 0.50 | Ni

14.

A 2.5624-g sample of a pure solid alkali metal chloride is dissolved in water and treated with excess silver nitrate. The resulting precipitate, filtered and dried, weighs 3.03707 g. What was the percent by mass of chloride ion in the original compound? What is the identity of the salt?

15.

The standard reduction potential for the reaction [Co(H2O)6]3+(aq)+e−⟶[Co(H2O)6]2+(aq)[Co(H2O)6]3+(aq)+e−⟶[Co(H2O)6]2+(aq) is about 1.8 V. The reduction potential for the reaction [Co(NH3)6]3+(aq)+e−⟶[Co(NH3)6]2+(aq)[Co(NH3)6]3+(aq)+e−⟶[Co(NH3)6]2+(aq) is +0.1 V. Calculate the cell potentials to show whether the complex ions, [Co(H2O)6]2+ and/or [Co(NH3)6]2+, can be oxidized to the corresponding cobalt(III) complex by oxygen.

16.

Predict the products of each of the following reactions. (Note: In addition to using the information in this chapter, also use the knowledge you have accumulated at this stage of your study, including information on the prediction of reaction products.)

(a) MnCO3(s)+HI(aq)⟶MnCO3(s)+HI(aq)⟶

(b) CoO(s)+O2(g)⟶CoO(s)+O2(g)⟶

(c) La(s)+O2(g)⟶La(s)+O2(g)⟶

(d) V(s)+VCl4(s)⟶V(s)+VCl4(s)⟶

(e) Co(s)+xsF2(g)⟶Co(s)+xsF2(g)⟶

(f) CrO3(s)+CsOH(aq)⟶CrO3(s)+CsOH(aq)⟶

17.

Predict the products of each of the following reactions. (Note: In addition to using the information in this chapter, also use the knowledge you have accumulated at this stage of your study, including information on the prediction of reaction products.)

(a) Fe(s)+H2SO4(aq)⟶Fe(s)+H2SO4(aq)⟶

(b) FeCl3(aq)+NaOH(aq)⟶FeCl3(aq)+NaOH(aq)⟶

(c) Mn(OH)2(s)+HBr(aq)⟶Mn(OH)2(s)+HBr(aq)⟶

(d) Cr(s)+O2(g)⟶Cr(s)+O2(g)⟶

(e) Mn2O3(s)+HCl(aq)⟶Mn2O3(s)+HCl(aq)⟶

(f) Ti(s)+xsF2(g)⟶Ti(s)+xsF2(g)⟶

18.

Describe the electrolytic process for refining copper.

19.

Predict the products of the following reactions and balance the equations.

(a) Zn is added to a solution of Cr2(SO4)3 in acid.

(b) FeCl2 is added to a solution containing an excess of Cr2O72−Cr2O72− in hydrochloric acid.

(c) Cr2+ is added to Cr2O72−Cr2O72− in acid solution.

(d) Mn is heated with CrO3.

(e) CrO is added to 2HNO3 in water.

(f) FeCl3 is added to an aqueous solution of NaOH.

20.

What is the gas produced when iron(II) sulfide is treated with a nonoxidizing acid?

21.

Predict the products of each of the following reactions and then balance the chemical equations.

(a) Fe is heated in an atmosphere of steam.

(b) NaOH is added to a solution of Fe(NO3)3.

(c) FeSO4 is added to an acidic solution of KMnO4.

(d) Fe is added to a dilute solution of H2SO4.

(e) A solution of Fe(NO3)2 and HNO3 is allowed to stand in air.

(f) FeCO3 is added to a solution of HClO4.

(g) Fe is heated in air.

22.

Balance the following equations by oxidation-reduction methods; note that three elements change oxidation state.
 Co(NO3)2(s)⟶Co2O3(s)+NO2(g)+O2(g)Co(NO3)2(s)⟶Co2O3(s)+NO2(g)+O2(g)

23.

Dilute sodium cyanide solution is slowly dripped into a slowly stirred silver nitrate solution. A white precipitate forms temporarily but dissolves as the addition of sodium cyanide continues. Use chemical equations to explain this observation. Silver cyanide is similar to silver chloride in its solubility.

24.

Predict which will be more stable, [CrO4]2− or [WO4]2−, and explain.

25.

Give the oxidation state of the metal for each of the following oxides of the first transition series. (Hint: Oxides of formula M3O4 are examples of mixed valence compounds in which the metal ion is present in more than one oxidation state. It is possible to write these compound formulas in the equivalent format MO·M2O3, to permit estimation of the metal’s two oxidation states.)

(a) Sc2O3

(b) TiO2

(c) V2O5

(d) CrO3

(e) MnO2

(f) Fe3O4

(g) Co3O4

(h) NiO

(i) Cu2O

19.2 Coordination Chemistry of Transition Metals

26.

Indicate the coordination number for the central metal atom in each of the following coordination compounds:

(a) [Pt(H2O)2Br2]

(b) [Pt(NH3)(py)(Cl)(Br)] (py = pyridine, C5H5N)

(c) [Zn(NH3)2Cl2]

(d) [Zn(NH3)(py)(Cl)(Br)]

(e) [Ni(H2O)4Cl2]

(f) [Fe(en)2(CN)2]+ (en = ethylenediamine, C2H8N2)

27.

Give the coordination numbers and write the formulas for each of the following, including all isomers where appropriate:

(a) tetrahydroxozincate(II) ion (tetrahedral)

(b) hexacyanopalladate(IV) ion

(c) dichloroaurate(I) ion (note that aurum is Latin for "gold")

(d) diamminedichloroplatinum(II)

(e) potassium diamminetetrachlorochromate(III)

(f) hexaamminecobalt(III) hexacyanochromate(III)

(g) dibromobis(ethylenediamine) cobalt(III) nitrate

28.

Give the coordination number for each metal ion in the following compounds:

(a) [Co(CO3)3]3− (note that CO32− is bidentate in this complex)

(b) [Cu(NH3)4]2+

(c) [Co(NH3)4Br2]2(SO4)3

(d) [Pt(NH3)4][PtCl4]

(e) [Cr(en)3](NO3)3

(f) [Pd(NH3)2Br2] (square planar)

(g) K3[Cu(Cl)5]

(h) [Zn(NH3)2Cl2]

29.

Sketch the structures of the following complexes. Indicate any cis, trans, and optical isomers.

(a) [Pt(H2O)2Br2] (square planar)

(b) [Pt(NH3)(py)(Cl)(Br)] (square planar, py = pyridine, C5H5N)

(c) [Zn(NH3)3Cl]+ (tetrahedral)

(d) [Pt(NH3)3Cl]+ (square planar)

(e) [Ni(H2O)4Cl2]

(f) [Co(C2O4)2Cl2]3− (note that C2O42−C2O42− is the bidentate oxalate ion, −O2CCO2−)−O2CCO2−)

30.

Draw diagrams for any cis, trans, and optical isomers that could exist for the following (en is ethylenediamine):

(a) [Co(en)2(NO2)Cl]+

(b) [Co(en)2Cl2]+

(c) [Pt(NH3)2Cl4]

(d) [Cr(en)3]3+

(e) [Pt(NH3)2Cl2]

31.

Name each of the compounds or ions given in Exercise 19.28, including the oxidation state of the metal.

32.

Name each of the compounds or ions given in Exercise 19.30.

33.

Specify whether the following complexes have isomers.

(a) tetrahedral [Ni(CO)2(Cl)2]

(b) trigonal bipyramidal [Mn(CO)4NO]

(c) [Pt(en)2Cl2]Cl2

34.

Predict whether the carbonate ligand CO32−CO32−will coordinate to a metal center as a monodentate, bidentate, or tridentate ligand.

35.

Draw the geometric, linkage, and ionization isomers for [CoCl5CN][CN].

19.3 Spectroscopic and Magnetic Properties of Coordination Compounds

36.

Determine the number of unpaired electrons expected for [Fe(NO2)6]3−and for [FeF6]3− in terms of crystal field theory.

37.

Draw the crystal field diagrams for [Fe(NO2)6]4− and [FeF6]3−. State whether each complex is high spin or low spin, paramagnetic or diamagnetic, and compare Δoct to P for each complex.

38.

Give the oxidation state of the metal, number of d electrons, and the number of unpaired electrons predicted for [Co(NH3)6]Cl3.

39.

The solid anhydrous solid CoCl2 is blue in color. Because it readily absorbs water from the air, it is used as a humidity indicator to monitor if equipment (such as a cell phone) has been exposed to excessive levels of moisture. Predict what product is formed by this reaction, and how many unpaired electrons this complex will have.

40.

Is it possible for a complex of a metal in the transition series to have six unpaired electrons? Explain.

41.

How many unpaired electrons are present in each of the following?

(a) [CoF6]3− (high spin)

(b) [Mn(CN)6]3− (low spin)

(c) [Mn(CN)6]4− (low spin)

(d) [MnCl6]4− (high spin)

(e) [RhCl6]3− (low spin)

42.

Explain how the diphosphate ion, [O3P−O−PO3]4−, can function as a water softener that prevents the precipitation of Fe2+ as an insoluble iron salt.

43.

For complexes of the same metal ion with no change in oxidation number, the stability increases as the number of electrons in the t2g orbitals increases. Which complex in each of the following pairs of complexes is more stable?

(a) [Fe(H2O)6]2+ or [Fe(CN)6]4−

(b) [Co(NH3)6]3+ or [CoF6]3−

(c) [Mn(CN)6]4− or [MnCl6]4−

44.

Trimethylphosphine, P(CH3)3, can act as a ligand by donating the lone pair of electrons on the phosphorus atom. If trimethylphosphine is added to a solution of nickel(II) chloride in acetone, a blue compound that has a molecular mass of approximately 270 g and contains 21.5% Ni, 26.0% Cl, and 52.5% P(CH3)3 can be isolated. This blue compound does not have any isomeric forms. What are the geometry and molecular formula of the blue compound?

45.

Would you expect the complex [Co(en)3]Cl3 to have any unpaired electrons? Any isomers?

46.

Would you expect the Mg3[Cr(CN)6]2 to be diamagnetic or paramagnetic? Explain your reasoning.

47.

Would you expect salts of the gold(I) ion, Au+, to be colored? Explain.

48.

[CuCl4]2− is green. [Cu(H2O)6]2+is blue. Which absorbs higher-energy photons? Which is predicted to have a larger crystal field splitting?

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