Chapter 2 Alkanes and Cycloalkanes: Introduction to Hydrocarbons

Transcription

1 Chapter 2 Alkanes and Cycloalkanes: Introduction to Hydrocarbons

2 2.1 Classes of Hydrocarbons

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

5 2.2 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

7 Formation of H 2 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

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

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

10 Potential Energy vs Distance Between Two Hydrogen Atoms Potential energy 74 pm maximum net attraction (minimum potential energy) at 74 pm internuclear distance H + H H H H H H H -436 kj/mol H 2 Internuclear distance

11 Potential Energy vs Distance Between Two Hydrogen Atoms Potential energy 74 pm repulsive forces increase faster than attractive forces at distances closer than 74 pm H + H H H H H H H -436 kj/mol H 2 Internuclear distance

12 Models for Chemical Bonding Valence Bond Theory constructive interference between electron waves of 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

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

14 2.3 Bonding in H 2 : The Valence Bond Model

15 Valence Bond Model 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:

16 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. Orbitals begin to overlap Electrons feel the attractive force of the protons Optimal distance between nuclei High electron density between the nuclei

17 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.

18 2.4 Bonding in H 2 : The Molecular Orbital Model

19 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.

20 Molecular Orbital Model for H 2 Addition of the AOs to form the bonding MO (s) Subtraction of the AOs to form the antibonding MO (s*)

21 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.

22 Energy-Level Diagram for H 2 MOs

23 2.5 Introduction to Alkanes: Methane, Ethane, and Propane

24 Small Alkanes General formula for alkanes is C n H 2n+2. Smallest alkane is methane CH 4 - also the most abundant. Ethane (C 2 H 6 ) and propane (C 3 H 8 ) are the next alkanes. Natural gas is 75% methane 10% ethane and 5% propane. These alkanes have the lowest boiling points.

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

26 2.6 sp 3 Hybridization and Bonding in Methane

27 Structure and Bonding Theory The dilemma: Methane has tetrahedral geometry. This is inconsistent with electron configuration of carbon of 1s 2, 2s 2, 2p x1,2p y 1 with only two unfilled orbitals.

28 sp 3 Hybrid Orbitals Linus Pauling proposed a mixing or hybridization of the s and three p orbitals to create 4 equal unfilled orbitals called sp 3 orbitals.

29 Properties of sp 3 Hybrid Orbitals All four sp 3 orbitals are of equal energy. The axes of the sp 3 orbitals point toward the corners of a tetrahedron. σ Bonds involving sp 3 hybrid orbitals of carbon are stronger than those involving unhybridized 2s or 2p orbitals.

30 Bonding with sp 3 Hybrid Orbitals Bonding in methane involves orbital overlap between each partially filled carbon sp 3 orbital and a partially filled s orbital of the hydrogen atom.

31 2.7 Bonding in Ethane

32 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 sp 3 orbitals on each of the carbon atoms.

33 C-C Bond Formation in Ethane Two half-filled sp 3 orbitals on each C Electrons with opposite spin Overlap of orbitals to form a bonding orbital.

34 2.8 sp2 Hybridization and Bonding in Ethane

35 Structure of Ethylene Ethylene is planar with bond angles close to 120 o. sp 3 Hybridization cannot be used to explain this bonding. Three atoms are bonded to each carbon so three hybrid orbitals are formed. Called sp 2 orbitals. One p orbital is not hybridized.

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

37 Sigma (s) Bonding in Ethylene Form C-H bonds by overlap of sp 2 and s orbitals Form C-C bond by overlap of sp 2 orbitals on each carbon These are all sigma (s) bonds. An unfilled p orbital remains on each carbon atom.

38 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.

39 2.9 sp Hybridization and Bonding in Acetylene

40 Structure of Acetyene Acetylene is linear with bond angles of 180 o. sp 3 and sp 2 Hybridization cannot explain this bonding. sp Hybridization explains this. There are two half filled p orbitals no hybridized.

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

42 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.

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

44 Hybridization of Carbon Carbons bonded to four atoms are sp 3 hybridized with bond angles of approximately o. Carbons bonded to three atoms are sp 2 hybridized with bond angles of approximately and one C-C p-bond. Carbons bonded to two atoms are sp hybridized with bond angles of approximately and two C-C p-bonds.

45 2.10 Which Theory of Chemical Bonding Is Best?

46 Theories of Chemical Bonding Approaches to chemical bonding: 1.Lewis model; 2.Orbital hybridization model; 3.Molecular orbital model.

47 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.

48 2.11 Isomeric Alkanes: The Butanes

49 Isomers of Butane There is only one isomer for each of the molecular formulas CH 4, C 2 H 6 and C 3 H 8. For C 4 H 10 there are two distinct connectivities of the carbon atoms. They are constitutional isomers. H H C H C H C H C H H H H H H H C H C H C H C H H H H H Bondline formulas

50 Isomers of Butane The isomers have different physical properties. All carbon atoms are sp 3 hybridized.

51 2.12 Higher n-alkanes

52 Higher n-alkanes n-alkanes are straight-chain alkanes with general formula CH 3 (CH 2 ) n CH 3. n-pentane is CH 3 CH 2 CH 2 CH 2 CH 3 and n-hexane is CH 3 CH 2 CH 2 CH 2 CH 2 CH 3. These formulas can be abbreviated as CH 3 (CH 2 ) 3 CH 3 or CH 3 (CH 2 ) 4 CH 3.

53 2.13 The C 5 H 12 Isomers

54 Isomers of C 5 H 12 There are three isomers C 5 H 12. It is important to realize that these are all representations of isopentane.

55 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.

56 2.13 IUPAC Nomenclature of Unbranched Alkanes

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

58 2.15 Applying the IUPAC Rules: The Names of the C 6 H 14 Isomers

59 The IUPAC Rules for Branched Alkanes Rules for naming branched alkanes: 1.Find the longest continuous carbon chain and its IUPAC name. This is the parent alkane. 2.Identify the substituents on this chain. substituent longest chain (5 carbons)

60 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

61 The IUPAC Rules for Branched Alkanes Rules for naming branched alkanes: 5. 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

62 2.16 Alkyl Groups

63 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 CH 3 CH 2 CH 2 CH 2 - substituent is a butyl group.

64 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.

65 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.

66 2.17 IUPAC Names of Highly Branched Alkanes

67 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

68 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.

69 2.18 Cycloalkane Nomenclature

70 Naming Cycloalkanes Cycloalkanes contain a ring of carbons and have general formula C n H 2n. Add the prefix cyclo- to the name of the corresponding alkane.

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

72 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.

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

74 2.19 Sources of Alkanes and Cycloalkanes

75 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.

76 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.

77 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 (CH 3 (CH 2 ) 23 CH 3 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.

78 2.20 Physical Properties of Alkanes and Cycloalkanes

79 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.

80 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.

81 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.

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

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

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

85 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.

86 2.22 Physical Properties of Alkanes and Cycloalkanes

87 Acidity of Hydrocarbons Hydrocarbons are very weak acids. Alkynes have the lowest pka.

88 Combustion of Hydrocarbons Combustion of hydrocarbons is exothermic generating CO 2 and water.

89 Combustion of Relative Stability All isomers of C 8 H 18 generate 8 molecules of CO 2 and 9 of H 2 O 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

90 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.

91 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.

92 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 CH 3 Cl and CH 3 OH. 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 CH 4 and CH 3 Li.