1. Chapter 2 Alkanes and Cycloalkanes: Introduction to Hydrocarbons1

2. Classes of Hydrocarbons1

3. Classes of Hydrocarbons
Hydrocarbons only contain carbon and hydrogen atoms.
Hydrocarbons are either classed as aliphatic or aromatic.
Aliphatic hydrocarbons contain three main groups: alkanes which only have carbon-carbon single bonds, alkenes which have a carbon-carbon double bond, or alkynes which have a carbon-carbon triple bond. 4

4. Classes of Hydrocarbons
Aromatic hydrocarbons are more complex but the simplest aromatic hydrocarbon is benzene. Aromatic hydrocarbons are called arenes. 4

5. Electron Waves and Chemical Bonds

6. Models for Chemical Bonding
The Lewis model of chemical bonding predates the idea that electrons have wave properties.
Two widely used theories of bonding based on the wave nature of an electron are:
Valence Bond Theory, and
Molecular Orbital Theory 2

7. Formation of H2 from Two Hydrogen Atoms+ e– + e–
Which electrostatic forces are involved as two hydrogen
atoms approach each other and form a H-H bond.
These electrostatic forces are:
attractions between the electrons and the nuclei
repulsions between the two nuclei
repulsions between the two electrons 6

8. Potential Energy vs Distance Between Two Hydrogen Atoms
weak net attraction at long distances
Potential energy
H• + H• H
Internuclear distance 6

9. Potential Energy vs Distance Between Two Hydrogen Atoms
attractive forces increase faster than repulsive forces as atoms approach each other
Potential energy
H• + H• H H H H
Internuclear distance 6

10. Potential Energy vs Distance Between Two Hydrogen Atoms
maximum net attraction (minimum potential energy) at 74 pm internuclear distance
Potential energy
74 pm
H• + H• H H H H
-436 kJ/mol H2
Internuclear distance 6

11. Potential Energy vs Distance Between Two Hydrogen Atoms
repulsive forces increase faster than attractive forces at distances closer than 74 pm
Potential energy
74 pm
H• + H• H H H H
-436 kJ/mol H2
Internuclear distance 6

12. Models for Chemical Bonding
Valence Bond Theory constructive interference between two half-filled atomic
orbitals is basis of shared-electron bond
Molecular Orbital Theory derive wave functions of molecules by combining wave functions of atoms 2

13. Behavior of Waves
Waves interactions include: Constructive interference when the waves are in phase and reinforce each other Destructive interference when the waves are out of phase and oppose each other 2

14. Valence Bond Model for Bonding in Hydrogen
Electron pair can be shared when half-filled orbital of one atom overlaps in phase with half-filled orbital of another. For example with overlap of two 1s orbitals of two hydrogen atoms shown below:

15. Valence Bond Model
The approach of the two hydrogen atoms can be modeled showing electrostatic potential maps. The high electron density between the nuclei is apparent.
Electrons feel the attractive force of the protons
Orbitals begin to overlap
Optimal distance between nuclei
High electron density between the nuclei

16. The Sigma (s) Bond
A bond in which the orbitals overlap along a line connecting the atoms is called a sigma (s) bond. Two perpendicular views are shown below.

17. Bonding in H2: The Molecular Orbital Model
Electrons in molecules occupy molecular orbitals (MOs) just as electrons in an atom occupy atomic orbitals (AOs).
MOs are combinations of AOs.
Two electrons per MO.
The additive combination of two atomic orbitals generates one bonding orbital.
The subtractive combination of the two atomic orbitals generates an antibonding orbital.

18. Molecular Orbital Model for H2
Addition of the AOs to form the bonding MO (s)
Subtraction of the AOs to form the antibonding MO (s*)

19. Molecular Orbital Digrams
Format is AOs on the sides and MOs in the middle.
Combination of n AOs results in n MOs.
Bonding MOs lower in energy than antibonding MOs.
Fill electrons in MOs the same as for AOs – lowest first.

20. Energy-Level Diagram for H2 MOs

21. Introduction to Alkanes: Methane, Ethane, and Propane21

22. Small Alkanes General formula for alkanes is CnH2n+2.
Smallest alkane is methane CH4 - also the most abundant.
Ethane (C2H6) and propane (C3H8) are the next alkanes.
Natural gas is 75% methane 10% ethane and 5% propane. These alkanes have the lowest boiling points. 22

23. Structures of Alkanes
All carbons in methane, ethane and propane have four bonds.
Bond angles (which are close to 109.5o) and bond lengths are: 22

24. sp3 Hybridization and Bonding in Methane

25. Structure and Bonding Theory
The dilemma: Methane has tetrahedral geometry.
This is inconsistent with electron configuration of carbon of 1s2, 2s2, 2px1,2py1 with only two unfilled orbitals.

26. sp3 Hybrid Orbitals
Linus Pauling proposed a mixing or hybridization of the s and three p orbitals to create 4 equal unfilled orbitals called sp3 orbitals.

27. Properties of sp3 Hybrid Orbitals
All four sp3 orbitals are of equal energy.
The axes of the sp3 orbitals point toward the corners of a tetrahedron.
σ Bonds involving sp3 hybrid orbitals of carbon are stronger than those involving unhybridized
2s or 2p orbitals.

28. Bonding with sp3 Hybrid Orbitals
Bonding in methane involves orbital overlap between each partially filled carbon sp3 orbital and a partially filled s orbital of the hydrogen atom.

29. Bonding and Structure of Ethane
Ethane also has tetrahedral geometry about the carbon atoms.
Hybridization can be used to rationalize the bonding.
The C-H bonds are formed as described for methane.
The C-C bond is formed by overlap of sp3 orbitals on each of the carbon atoms.

30. C-C Bond Formation in Ethane
Two half-filled sp3 orbitals on each C
Electrons with opposite spin
Overlap of orbitals to form a bonding orbital.

31. Structure of Ethylene and sp2 Hybridization
Ethylene is planar with bond angles close to 120o.
sp3 Hybridization cannot be used to explain this bonding.
Three atoms are bonded to each carbon so three hybrid orbitals are formed. Called sp2 orbitals.
One p orbital is not hybridized.

32. sp2 Hybrid Orbitals
The 2s and two of the 2p orbitals are mixed to form three sp2 orbitals with a trigonal planar arrangement. The 2pz orbital remains half filled.

33. Sigma (s) Bonding in Ethylene
Form C-H bonds by overlap of sp2 and s orbitals
Form C-C bond by overlap of sp2 orbitals on each carbon
These are all sigma (s) bonds. An unfilled p orbital remains on each carbon atom.

34. Pi (p) Bonding in Ethylene
Form second C-C bond by overlap of p orbitals on each carbon
This called a pi (p) bond and the electrons in the bond are called p electrons.

35. Structure of Acetylene and sp Hybridization
Acetylene is linear with bond angles of 180o.
sp3 and sp2 Hybridization cannot explain this bonding.
sp Hybridization explains this. There are two half filled p orbitals no hybridized.

36. sp Hybrid Orbitals
The 2s and one of the 2p orbitals are mixed to form two sp orbitals with a linear arrangement. The 2py and 2pz orbitals remain half filled.

37. Sigma (s) Bonding in Acetylene
Form C-H bonds by overlap of sp and s orbitals
Form C-C bond by overlap of sp orbitals on each carbon
These are all sigma (s) bonds. Two unfilled p orbitals remain on each carbon atom.

38. Pi (p) Bonding in Acetylene
Form one p bond by overlap of py orbitals on each carbon
Form second p bond by overlap of pz orbitals on each carbon
There are two pi (p) bonds and a total of 4 p electrons.

39. Hybridization of Carbon
Carbons bonded to four atoms are sp3 hybridized with bond angles of approximately 109.5o.
Carbons bonded to three atoms are sp2 hybridized with bond angles of approximately 1200 and one C-C p-bond.
Carbons bonded to two atoms are sp hybridized with bond angles of approximately 1800 and two C-C p-bonds.

40. Which Theory of Chemical Bonding Is Best?21

41. Theories of Chemical Bonding
Approaches to chemical bonding:
Lewis model;
Orbital hybridization model;
Molecular orbital model.

42. Considerations of Chemical Bonding
Lewis and Orbital hybridization models work together and success in organic depends on writing correct Lewis structures.
Molecular orbital theory provides insights into structure and reactivity lacking in the other models. This model requires higher level theory which will not be presented. The results of MO theory will be used – for example electrostatic potential maps.

43. Isomers of Butane
There is only one isomer for each of the molecular formulas CH4, C2H6 and C3H8.
For C4H10 there are two distinct connectivities of the carbon atoms. They are constitutional isomers.
Bondline formulas

44. Isomers of Butane The isomers have different physical properties.
All carbon atoms are sp3 hybridized.

45. Higher n-Alkanes
n-Alkanes are straight-chain alkanes with general formula CH3(CH2)nCH3. n-Pentane is CH3CH2CH2CH2CH3 and n-hexane is CH3CH2CH2CH2CH2CH3. These formulas can be abbreviated as CH3(CH2)3CH3 or CH3(CH2)4CH3.

46. Isomers of C5H12 There are three isomers C5H12.
It is important to realize that these are all representations of isopentane.

47. Isomers of higher n-alkanes
For higher n-alkanes there are many isomers and it is not possible to easily predict how many isomers can be formed.

48. IUPAC Nomenclature of Unbranched Alkanes21

49. IUPAC Naming
Alkane names are the basis of the IUPAC system of nomenclature. The –ane suffix is specific to alkanes.

50. The IUPAC Rules for Branched Alkanes
Rules for naming branched alkanes:
Find the longest continuous carbon chain and its IUPAC name. This is the parent alkane.
Identify the substituents on this chain.
substituent
longest chain (5 carbons)

51. The IUPAC Rules for Branched Alkanes
Rules for naming branched alkanes:
3. Number the longest continuous chain in the direction that gives the lowest number to the first substituent.
4. Write the name of the compound. The parent alkane is the last part of the name and is preceded by the names of the substituents and their numerical locations (locants). Hyphens separate the locants from the words.
2-methylpentane

52. The IUPAC Rules for Branched Alkanes
Rules for naming branched alkanes:
When the same substituent appears more than once,
use the multiplying prefixes di-, tri-, tetra-, and so on. A separate locant for each substituent. Locants are separated from each other by commas and from the words by hyphens.
2,2-dimethylbutane
2,3-dimethylbutane

53. Alkyl Groups
Alkyl groups are substituents derived from alkanes. They lack one hydrogen at the point of attachment.
The alkyl group is named from the alkane by replacing the -ane suffix with –yl. For example a CH3CH2CH2CH2- substituent is a butyl group. 2

54. Classification of Carbon Atoms
Carbon atoms are defined as primary, secondary, tertiary or quaternary. A primary carbon is directly attached to one other carbon. A secondary carbon is directly attached to two other carbons. A tertiary carbon to 3 and a quaternary carbon to 4. 2

55. Complex Alkyl Groups (Substituents)
Secondary and tertiary groups may have common names and IUPAC names. The base name of these groups is the longest chain including the attachment carbon form and the substituents are located on this chain. 2

56. Naming Highly Branched Alkanes
When two or more different substituents are present number from the end closest to the first point of difference.
When two or more different substituents are present, they are listed in alphabetical order in the name. Prefixes such as di-, tri-, and tetra- are used but ignored when alphabetizing. tert-Butyl precedes isobutyl. sec-Butyl precedes tert-butyl.
4-ethyl-3,5-dimethyloctane 2

57. Naming Highly Branched Alkanes
When two or more different substituents are present number from the end closest to the first point of difference.
If the first substituent is located an equal distance from each end then the second substituent becomes the first potential point of difference and so on. 2

58. Naming Cycloalkanes
Cycloalkanes contain a ring of carbons and have general formula CnH2n.
Add the prefix cyclo- to the name of the corresponding alkane. 2

59. Naming Cycloalkanes Identify and name substituents as before.
For one substituent no numbers are used. 2

60. Naming Cycloalkanes
For multiple substituents the locations must be specified.
Number the carbon atoms of the ring in the direction that gives the lowest number to the substituents at the first point of difference.
First substituent is on C1 by default. 2

61. Naming Cycloalkanes
If the ring has fewer carbons than the alkyl group attached to it then the ring is the substituent. 2

62. Sources of Alkanes and Cycloalkanes
Natural is mainly methane with ethane and propane.
Petroleum is a liquid mixture containing approximately 150 hydrocarbons. Half of these are alkanes or cycloalkanes.
Distillation of crude oil gives fractions based on boiling point. 2

63. Petroleum Refining
The yield of the more useful petroleum fraction used as automotive fuel is increased by two processes:
Cracking. Cracking is the cleavage of carbon–carbon bonds in high molecular weight alkanes induced by heat (thermal cracking) or with catalysts (catalytic cracking).
Reforming. Reforming converts the hydrocarbons in petroleum to aromatic hydrocarbons and highly branched alkanes, both of which are better automotive fuels than unbranched alkanes and cycloalkanes. 2

64. Other Natural Sources of Alkanes
Solid n-alkanes are waxy and coat the outer surface of many living things to prevent loss of water. Examples include: Pentacosane (CH3(CH2)23CH3 is found in the waxy outer layer of many insects.
Hentriacontane is a component of beeswax and the outer layer of leaves of tobacco, peach trees and others.
Hopanes are found in petroleum and geologic sediments. 2

65. Physical Properties of Alkanes and Cycloalkanes4

66. Boiling Point
Boiling points of n-alkanes increase with increasing molecular weight (number of carbons). Branched alkanes generally have lower boiling points than unbranched alkanes with the same number of carbons. 2

67. Intermolecular Forces and Boiling Point
Attractive forces between molecules in the liquid phase affect the boiling point of the liquid. These Intermolecular forces are van der Waals forces and may be divided into three types: Dipole-dipole (including hydrogen bonding); Induced dipole-dipole; or Induced dipole-induced dipole. 2

68. Intermolecular Forces and Alkanes
Alkanes have no dipole so the van der Waals forces are the temporary induced dipole-induced dipole. This interaction is dynamic and fluctuates. 2

69. Intermolecular Forces and Alkanes
Long chain alkanes have more induced dipole-induced dipole interactions so the boiling point increases with increasing chain length. 2

70. Intermolecular Forces and Alkanes
Branched alkanes have lower surface area than isomeric n-alkanes and therefore have lower boiling points. 2

71. Melting Point
Solid n-alkanes are soft low melting solids. The same intermolecular forces hold the molecules together in the solid state. 2

72. Solubility of Alkanes in Water
Alkanes (and all hydrocarbons) are virtually insoluble in water and are said to be hydrophobic.
The densities of most alkanes are in the range g/mL therefore alkanes float on the surface of water. 2

73. Chemical Properties of Alkanes and Cycloalkanes4

74. Acidity of Hydrocarbons
Hydrocarbons are very weak acids. Alkynes have the lowest pKa. 2

75. Combustion of Hydrocarbons
Combustion of hydrocarbons is exothermic generating CO2 and water. 2

76. Combustion of Relative Stability
All isomers of C8H18 generate 8 molecules of CO2 and 9 of H2O yet different amounts of energy. This energy difference must be directly related to the relative energies of the isomers.
Least stable isomer
Most stable isomer
Least energy released 2

77. Oxidation and Reduction in Organic Chemistry
Assuming the oxidation state of H is +1 and O is -2 it is possible to calculate the oxidation state of C in compounds containing C, H and O. 2

78. Oxidation and Reduction in Organic Chemistry
Oxidation of carbon corresponds to an increase in the number of bonds between carbon and oxygen or to a decrease in the number of carbon–hydrogen bonds.
Reduction corresponds to an increase in the number of carbon–hydrogen bonds or to a decrease in the number of carbon–oxygen bonds. 2

79. Oxidation and Reduction in Organic Chemistry
Any element more electronegative than C has the same effect as O on the oxidation state of C. Oxidation state of C is +2 in CH3Cl and CH3OH.
Any element less electronegative than C has the same effect as H on the oxidation state of C. Oxidation state of C is -4 in CH4 and CH3Li. 2