1. Figure: 08-01
Title:
Chemical bonds.
Caption:
Examples of substances in which (a) ionic, (b) covalent, and (c) metallic bonds are found.

2. Figure: UN
Title:
Give It Some Thought.
Caption:
Three Lewis structures.

3. Figure: 08-02
Title:
Formation of sodium chloride.
Caption:
Sodium metal and chlorine gas react vigorously to form the ionic substance sodium chloride.

4. Figure: 08-03
Title:
The crystal structure of sodium chloride.
Caption:
In this three-dimensional array of ions, each Na+ ion is surrounded by six Cl– ions, and each Cl– ion is surrounded by six Na+ ions.

5. Figure: UN
Title:
Equation 8.2
Caption:
Formation of NaCl.

6. Figure: 08-04
Title:
The Born–Haber cycle.
Caption:
This representation shows the energetic relationships in the formation of ionic solids from the elements. By Hess's law, the enthalpy of formation of NaCl(s) from elemental sodium and chlorine (Equation 8.5) is equal to the sum of the energies of several individual steps (Equations 8.6 through 8.10).

7. Figure: 08-05
Title:
The covalent bond in H2.
Caption:
(a) The attractions and repulsions among electrons and nuclei in the hydrogen molecule. (b) Electron distribution in the H2 molecule. The concentration of electron density between the nuclei leads to a net attractive force that constitutes the covalent bond holding the molecule together.

8. Figure: UN
Title:
Formation of H2.
Caption:
Lewis structures.

9. Figure: UN
Title:
Formation of Cl2.
Caption:
Lewis structures.

10. Figure: UN
Title:
Lewis structures.
Caption:
H2 and Cl2.

11. Figure: UN
Title:
Lewis structures.
Caption:
Some simple hydrogen compounds.

12. Figure: UN
Title:
Sample Exercise 8.3
Caption:
Lewis structures.

13. Figure: UN
Title:
Lewis structure of CO2.
Caption:
Formation of double bonds.

14. Figure: UN
Title:
Lewis structure of N2.
Caption:
Formation of a triple bond.

15. Figure: 08-06
Title:
Electronegativities of the elements.
Caption:
Electronegativity generally increases from left to right across a period and decreases from top to bottom down a group.

16. Figure: 08-07
Title:
Electron density distribution.
Caption:
This computer-generated rendering shows the calculated electron-density distribution on the surface of the F2, HF, and LiF molecules. The regions of relatively low electron density (net positive charge) appear blue, those of relatively high electron density (net negative charge) appear red, and regions that are close to electrically neutral appear green.

17. Figure: 08-08
Title:
Dipole and dipole moment.
Caption:
When charges of equal magnitude and opposite sign Q+ and Q– are separated by a distance r, a dipole is produced. The size of the dipole is given by the dipole moment, , which is the product of the charge separated and the distance of separation between their centers:  = Qr.

18. Figure: 08-09
Title:
Charge separation in the hydrogen halides.
Caption:
Blue represents regions of lowest electron density, red regions of highest electron density. In HF the strongly electronegative F pulls much of the electron density away from H. In HI the I, being much less electronegative than F, does not attract the shared electrons as strongly, and consequently there is far less polarization of the bond.

19. Figure: UNEx8.6
Title:
Sample Exercise 8.6
Caption:
Skeleton structure.

20. Figure: UNEx8.6
Title:
Sample Exercise 8.6
Caption:
Partial Lewis structure of PCl3.

21. Figure: UNEx8.6
Title:
Sample Exercise 8.6
Caption:
Completed Lewis structure.

22. Figure: UNEx8.6
Title:
Practice Exercise
Caption:
Lewis structure.

23. Figure: UNEx8.7
Title:
Sample Exercise 8.7
Caption:
Partial Lewis structure.

24. Figure: UNEx8.7
Title:
Sample Exercise 8.7
Caption:
Lewis structure of HCN.

25. Figure: UNEx8.7
Title:
Practice Exercise
Caption:
Lewis structure.

26. Figure: UNEx8.8
Title:
Sample Exercise 8.8
Caption:
Lewis structure of a polyatomic ion.

27. Figure: UNEx8.8
Title:
Practice Exercise
Caption:
Lewis structures of ions.

28. Figure: UN
Title:
Formal charges.
Caption:
The formal charge for each atom in CO2.

29. Figure: 08-10
Title:
Oxidation number and formal charge
Caption:
(a) The oxidation number for any atom in a molecule is determined by assigning all shared electrons to the more electronegative atom (in this case Cl). (b) Formal charges are derived by dividing all shared electrons equally between the bonded atoms. (c) The calculated distribution of electron density on an HCl molecule. Regions of relatively more negative charge are red; those of more positive charge are blue. Negative charge is clearly localized on chlorine.

30. Figure: UNEx8.9
Title:
Sample Exercise 8.9
Caption:
Three possible Lewis structures

31. Figure: UNEx8.9
Title:
Sample Exercise 8.9
Caption:
Lewis structures and formal charges.

32. Figure: UNEx8.9
Title:
Practice Exercise
Caption:
Lewis structures and formal charges.

33. Figure: 08-11
Title:
Ozone.
Caption:
Molecular structure (top) and electron-distribution diagram (bottom) for the ozone molecule, O3.

34. Figure: 08-12
Title:
Resonance.
Caption:
Describing a molecule as a blend of difference resonance structures is similar to describing a paint color as a blend of primary colors. (a) Green paint is a blend of blue and yellow. We cannot describe green as a single primary color. (b) The ozone molecule is a blend of two resonance structures. We cannot describe the ozone molecule in terms of a single Lewis structure.

35. Figure: UN
Title:
Resonance structures for the nitrate ion.
Caption:
Three Lewis structures.

36. Figure: UNEx8.10
Title:
Sample Exercise 8.10
Caption:
Resonance structures.

37. Figure: UNEx8.10
Title:
Sample Exercise 8.10
Caption:
Lewis structure.

38. Figure: UNEx8.10
Title:
Practice Exercise
Caption:
Formate ion.

39. Figure: 08-13
Title:
Benzene, an "aromatic" compound.
Caption:
(a) Benzene is obtained from the distillation of fossil fuels. More than 4 billion pounds of benzene is produced annually in the United States. Because benzene is a carcinogen, its use is closely regulated. (b) The benzene molecule is a regular hexagon of carbon atoms with a hydrogen atom bonded to each one.

40. Figure: UN
Title:
Benzene.
Caption:
Resonance structures of benzene.

41. Figure: UN
Title:
Shorthand structure of benzene.
Caption:
The carbon–carbon framework of benzene.

42. Figure: UN
Title:
Give it Some Thought.
Caption:
Lewis structure of hexatriene.

43. Figure: UN
Title:
BF3.
Caption:
Equivalent resonance structures.

44. Figure: UN
Title:
Resonance structures of BF3.
Caption:
Resonance structures showing the most and less important structures.

45. Figure: UN
Title:
Reaction between NH3 and BF3.
Caption:
Lewis structures of the reactants and products.

46. Figure: UN
Title:
Structure of PCl5.
Caption:
Lewis structure.

47. Figure: UN
Title:
Structure of PCl5.
Caption:
Ball-and-stick model.

48. Figure: UNEx8.11
Title:
Sample Exercise 8.11
Caption:
Lewis structure of ICl4–.

49. Figure: UNEx8.11
Title:
Practice Exercise
Caption:
Lewis structure for XeF2.

50. Figure: UN
Title:
PO43–.
Caption:
Two Lewis structures.

51. Figure: UN
Title:
Dissociation of Cl2.
Caption:
Bond enthalpy.

52. Figure: UN
Title:
Atomization of CH4.
Caption:
Reaction used to define the average bond enthalpy for the C—H bond.

53. Figure: 08-14
Title:
Using bond enthalpies to calculate Hrxn.
Caption:
Average bond enthalpies were used to estimate Hrxn for the reaction in Equation Breaking the C—H and Cl—Cl bonds produces a positive enthalpy change (H1), whereas making the C—Cl and H—Cl bonds causes a negative enthalpy change, H2. The values of H1 and H2 are estimated from the values in Table From Hess's law, Hrxn = H1 + H2.

54. Figure: UNEx8.12
Title:
Sample Exercise 8.12
Caption:
Oxidation of ethane.

55. Figure: UNEx8.12
Title:
Practice Exercise
Caption:
Decomposition of N2H2.

56. Figure: UN
Title:
Chemistry at Work
Caption:
Lewis structures of nitroglycerine and TNT.

57. Figure: UN
Title:
Sample Integrative Exercise
Caption:
Lewis structures of phosgene.

58. Figure: UN
Title:
Sample Integrative Exercise
Caption:
Formal charges.

59. Figure: UN
Title:
Sample Integrative Exercise
Caption:
Lewis structures for a chemical reaction.

60. Figure: UNE1
Title:
Exercise 8.1
Caption:
Lewis structure of X.

61. Figure: UNE1
Title:
Exercise 8.1
Caption:
Lewis structure of X.

62. Figure: UNE1
Title:
Exercise 8.1
Caption:
Lewis structure of X.

63. Figure: UNE2
Title:
Exercise 8.2
Caption:
Molecular models of four ions.

64. Figure: UNE3
Title:
Exercise 8.3
Caption:
Orbital diagram.

65. Figure: UNE4
Title:
Exercise 8.4
Caption:
Lewis structure.

66. Figure: UNE5
Title:
Exercise 8.5
Caption:
Partial Lewis structure.

67. Figure: UNE6
Title:
Exercise 8.6
Caption:
One Lewis structure for XeO3.

68. Figure: UNE56
Title:
Exercise 8.56
Caption:
Incomplete Lewis structure for naphthalene.

69. Figure: UNE65
Title:
Exercise 8.65
Caption:
Three chemical reactions.

70. Figure: UNE66
Title:
Exercise 8.66
Caption:
Three chemical reactions.

71. Figure: UNE86
Title:
Exercise, 8.86
Caption:
Lewis structures of N2O.

72. Figure: UNE88
Title:
Exercise 8.88
Caption:
Gas-phase isomerization reaction.

73. Figure: UNE88
Title:
Exercise 8.88
Caption:
Gas-phase isomerization reaction.

74. Figure: UNE88
Title:
Exercise 8.88
Caption:
Gas-phase isomerization reaction.

75. Figure: UNE88
Title:
Exercise 8.88
Caption:
Gas-phase isomerization reaction.

76. Figure: UNE90
Title:
Exercise 8.90
Caption:
Structure of cyclotrimethylenetrinitramine (RDX).

77. Figure: UNE91
Title:
Exercise 8.91
Caption:
Reaction of A=A.

78. Figure: UNE103
Title:
Exercise 8.103
Caption:
Ball-and-stick model of P4.

79. Figure: 08-T01
Title:
Table 8.1
Caption:
Lewis Symbols.

80. Figure: 08-T02
Title:
Table 8.2
Caption:
Lattice Energies for Some Ionic Compounds.

81. Figure: 08-T03
Title:
Table 8.3
Caption:
Bond Lengths, Electronegativity Differences, and Dipole Moments of the Hydrogen Halides.

82. Figure: 08-T04
Title:
Table 8.4
Caption:
Average Bond Enthalpies (kJ/mol).

83. Figure: 08-T05
Title:
Table 8.5
Caption:
Average Bond Lengths for Some Single, Double, and Triple Bonds.