Chapter 2 Alkanes and Cycloalkanes: Introduction to Hydrocarbons

1. Chapter 2Alkanes and Cycloalkanes: Introduction to Hydrocarbons

2. 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. Electron Waves andChemical 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 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

8. H H Internuclear distance Potential Energy vs Distance Between Two Hydrogen Atoms weak net attraction atlong distances Potentialenergy H• + H•

9. H H H H Internuclear distance Potential Energy vs Distance Between Two Hydrogen Atoms attractive forces increasefaster than repulsive forcesas atoms approach each other Potentialenergy H• + H• H H

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

11. H H H H Internuclear distance Potential Energy vs Distance Between Two Hydrogen Atoms repulsive forces increasefaster than attractive forcesat distances closer than 74 pm 74 pm Potentialenergy H• + H• H H -436 kJ/mol H2

12. Models for Chemical Bonding • Valence Bond Theoryconstructive interference between two half-filled atomic • orbitals is basis of shared-electron bond • Molecular Orbital Theoryderive wave functions of moleculesby combining wave functions of atoms

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

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. The approach of the two hydrogen atoms can be modeled showing electrostatic potential maps. The high electrondensity between the nuclei is apparent. Valence Bond Model 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 orbitalsgenerates an antibonding orbital.

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

19. 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. Molecular Orbital Digrams

20. Energy-Level Diagram for H2 MOs

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

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.

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:

24. sp3 HybridizationandBonding in Methane

25. Structure and Bonding Theory The dilemma:Methane has tetrahedral geometry. This is inconsistent with electron configuration of carbon of1s2, 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 bondingorbital.

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 hybridorbitals 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 sp2and s orbitals Form C-C bond by overlap of sp2orbitals 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 and2pz orbitals remain half filled.

37. Sigma (s) Bonding in Acetylene Form C-H bonds by overlap of spand s orbitals Form C-C bond by overlap of sporbitals 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 withbond angles of approximately 1200 and one C-C p-bond. Carbons bonded to two atoms are sp hybridized withbond angles of approximately 1800 and two C-C p-bonds.

40. Which Theory of Chemical Bonding Is Best?

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 structureand 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. Bondlineformulas

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 asCH3(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 ofUnbranched Alkanes

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)