1. Figure 1. 1: The position of carbon in the periodic table
Figure 1.1: The position of carbon in the periodic table. Other elements commonly found in organic compounds are shown in the colors typically used to represent them.
Fig. 1-1, p. 3

2. Figure 1. 2: A schematic view of an atom
Figure 1.2: A schematic view of an atom. The dense, positively charged nucleus contains most of the atom’s mass and is surrounded by negatively charged electrons. The three-dimensional view on the right shows calculated electron-density surfaces. Electron density increases steadily toward the nucleus and is 40 times greater at the blue solid surface than at the gray mesh surface.
Fig. 1-2, p. 4

3. Figure 1. 3: Representations of s, p, and d orbitals
Figure 1.3: Representations of s, p, and d orbitals. The s orbitals are spherical, the p orbitals are dumbbell-shaped, and four of the five d orbitals are cloverleaf-shaped. Different lobes of p orbitals are often drawn for convenience as teardrops, but their true shape is more like that of a doorknob, as indicated.
Fig. 1-3, p. 5

4. Figure 1. 5: Shapes of the 2p orbitals
Figure 1.5: Shapes of the 2p orbitals. Each of the three mutually perpendicular, dumbbell-shaped orbitals has two lobes separated by a node. The two lobes have different algebraic signs in the corresponding wave function, as indicated by the different colors.
Fig. 1-5, p. 6

5. Figure 1. 4: The energy levels of electrons in an atom
Figure 1.4: The energy levels of electrons in an atom. The first shell holds a maximum of 2 electrons in one 1s orbital; the second shell holds a maximum of 8 electrons in one 2s and three 2p orbitals; the third shell holds a maximum of 18 electrons in one 3s, three 3p, and five 3d orbitals; and so on. The two electrons in each orbital are represented by up and down arrows, ↑↓. Although not shown, the energy level of the 4s orbital falls between 3p and 3d.
Fig. 1-4, p. 5

6. p. 9

7. Figure 1.7: The cylindrical symmetry of the H–H σ bond in an H2 molecule. The intersection of a plane cutting through the σ bond is a circle.
Fig. 1-7, p. 11

8. Figure 1. 8: Relative energy levels of H atoms and the H2 molecule
Figure 1.8: Relative energy levels of H atoms and the H2 molecule. The H2 molecule has 436 kJ/mol (104 kcal/mol) less energy than the two H atoms, so 436 kJ/mol of energy is released when the H–H bond forms. Conversely, 436 kJ/mol must be added to the H2 molecule to break the H–H bond.
Fig. 1-8, p. 11

9. Figure 1.9: A plot of energy versus internuclear distance for two hydrogen atoms. The distance between nuclei at the minimum energy point is the bond length.
Fig. 1-9, p. 12

10. Figure 1.10: Four sp3 hybrid orbitals (green), oriented to the corners of a regular tetrahedron, are formed by combination of an s orbital (red) and three p orbitals (red/blue). The sp3 hybrids have two lobes and are unsymmetrical about the nucleus, giving them a directionality and allowing them to form strong bonds to other atoms.
Fig. 1-10, p. 13

11. Figure 1.11: The structure of methane, showing its 109.5° bond angles.
Fig. 1-11, p. 13

12. p. 14

13. Figure 1. 12: The structure of ethane
Figure 1.12: The structure of ethane. The carbon–carbon bond is formed by σ overlap of two carbon sp3 hybrid orbitals. For clarity, the smaller lobes of the sp3 hybrid orbitals are not shown.
Fig. 1-12, p. 14

14. Figure 1. 12: The structure of ethane
Figure 1.12: The structure of ethane. The carbon–carbon bond is formed by σ overlap of two carbon sp3 hybrid orbitals. For clarity, the smaller lobes of the sp3 hybrid orbitals are not shown.
Fig. 1-12a, p. 14

15. Figure 1. 12: The structure of ethane
Figure 1.12: The structure of ethane. The carbon–carbon bond is formed by σ overlap of two carbon sp3 hybrid orbitals. For clarity, the smaller lobes of the sp3 hybrid orbitals are not shown.
Fig. 1-12b, p. 14

16. p. 15

17. Figure 1. 13: An sp2-hybridized carbon
Figure 1.13: An sp2-hybridized carbon. The three equivalent sp2 hybrid orbitals (green) lie in a plane at angles of 120° to one another, and a single unhybridized p orbital (red/blue) is perpendicular to the sp2 plane.
Fig. 1-13, p. 15

18. Figure 1. 13: An sp2-hybridized carbon
Figure 1.13: An sp2-hybridized carbon. The three equivalent sp2 hybrid orbitals (green) lie in a plane at angles of 120° to one another, and a single unhybridized p orbital (red/blue) is perpendicular to the sp2 plane.
Fig. 1-13a, p. 15

19. Figure 1. 13: An sp2-hybridized carbon
Figure 1.13: An sp2-hybridized carbon. The three equivalent sp2 hybrid orbitals (green) lie in a plane at angles of 120° to one another, and a single unhybridized p orbital (red/blue) is perpendicular to the sp2 plane.
Fig. 1-13b, p. 15

20. Figure 1. 14: The structure of ethylene
Figure 1.14: The structure of ethylene. Orbital overlap of two sp2-hybridized carbons forms a carbon–carbon double bond. One part of the double bond results from σ (head-on) overlap of sp2 orbitals (green), and the other part results from π (sideways) overlap of unhybridized p orbitals (red/blue). The π bond has regions of electron density on either side of a line drawn between nuclei.
Fig. 1-14, p. 16

21. Figure 1. 14: The structure of ethylene
Figure 1.14: The structure of ethylene. Orbital overlap of two sp2-hybridized carbons forms a carbon–carbon double bond. One part of the double bond results from σ (head-on) overlap of sp2 orbitals (green), and the other part results from π (sideways) overlap of unhybridized p orbitals (red/blue). The π bond has regions of electron density on either side of a line drawn between nuclei.
Fig. 1-14a, p. 16

22. Figure 1. 14: The structure of ethylene
Figure 1.14: The structure of ethylene. Orbital overlap of two sp2-hybridized carbons forms a carbon–carbon double bond. One part of the double bond results from σ (head-on) overlap of sp2 orbitals (green), and the other part results from π (sideways) overlap of unhybridized p orbitals (red/blue). The π bond has regions of electron density on either side of a line drawn between nuclei.
Fig. 1-14b, p. 16

23. Figure 1. 15: An sp-hybridized carbon atom
Figure 1.15: An sp-hybridized carbon atom. The two sp hybrid orbitals (green) are oriented 180° away from each other, perpendicular to the two remaining p orbitals (red/blue).
Fig. 1-15, p. 18

24. Figure 1. 16: The structure of acetylene
Figure 1.16: The structure of acetylene. The two sp-hybridized carbon atoms are joined by one sp–sp σ bond and two p–p π bonds.
Fig. 1-16, p. 18

25. Figure 1. 16: The structure of acetylene
Figure 1.16: The structure of acetylene. The two sp-hybridized carbon atoms are joined by one sp–sp σ bond and two p–p π bonds.
Fig. 1-16a, p. 18

26. Figure 1. 16: The structure of acetylene
Figure 1.16: The structure of acetylene. The two sp-hybridized carbon atoms are joined by one sp–sp σ bond and two p–p π bonds.
Fig. 1-16b, p. 18

27. Atoms +
lone pairs*
# of hybrid
orbitals
Unhybridized
p-orbitals
Example
H3C-CH sp3 0 0
H2C=CH sp2 1 1
H-CºC-H sp 2 2
s bonds +
lone pairs*
hybrid
p-bonds
* Exception: lone pairs on atoms adjacent to other atoms with p-orbitals will also be in p-orbitals for resonance

28. Table 1-2, p. 18

29. Figure 1. 17: Molecular orbitals of H2
Figure 1.17: Molecular orbitals of H2. Combination of two hydrogen 1s atomic orbitals leads to two H2 molecular orbitals. The lower-energy, bonding MO is filled, and the higher-energy, antibonding MO is unfilled.
Fig. 1-17, p. 21

30. Figure 1.18: A molecular orbital description of the C=C π bond in ethylene. The lower-energy, π bonding MO results from a combination of p orbital lobes with the same algebraic sign and is filled. The higher-energy, π antibonding MO results from a combination of p orbital lobes with the opposite algebraic signs and is unfilled.
Fig. 1-18, p. 22

31. p. 23

32. Table 1-3, p. 23

33. p. 24

34. p. 30

35. p. 30

36. p. 31

37. p. 31

38. p. 31

39. p. 31

40. p. 31

41. p. 31

42. p. 31

43. p. 31

44. p. 32

45. p. 32

46. p. 32

47. p. 33

48. p. 33

49. p. 33

50. p. 34