Chapter 2 Hydrocarbon Frameworks: Alkanes

1. Chapter 2Hydrocarbon Frameworks:Alkanes

2. Classes of Hydrocarbons

3. Hydrocarbons Aliphatic Aromatic

4. Hydrocarbons Aliphatic Aromatic Alkanes Alkenes Alkynes

5. H H H H C C H H Hydrocarbons • Alkanes are hydrocarbons in which all of the bonds are single bonds. Aliphatic Alkanes

6. H H C C H H Hydrocarbons • Alkenes are hydrocarbons that contain a carbon-carbon double bond. Aliphatic Alkenes

7. HC CH Hydrocarbons • Alkynes are hydrocarbons that contain a carbon-carbon triple bond. Aliphatic Alkynes

8. H H H H H H Hydrocarbons • The most common aromatic hydrocarbons are those that contain a benzene ring. Aromatic

9. Electron Waves andChemical Bonds

10. Models for Chemical Bonding The Lewis model of chemical bonding predatesthe idea that electrons have wave properties. There are two other widely used theories ofbonding that are based on the wave nature of anelectron. • Valence Bond Theory • Molecular Orbital Theory

11. Formation of H2 from Two Hydrogen Atoms e– e– + + • Examine how the electrostatic forces change as two hydrogen atoms are brought together. • These electrostatic forces are:attractions between the electrons and the nucleirepulsions between the two nucleirepulsions between the two electrons

12. H H Internuclear distance Figure 2.1 weak net attraction atlong distances Potentialenergy H• + H•

13. H H H H Internuclear distance Figure 2.1 attractive forces increasefaster than repulsive forcesas atoms approach each other Potentialenergy H• + H• H H

14. H H H H Internuclear distance Figure 2.1 maximum net attraction (minimum potential energy)at 74 pm internuclear distance 74 pm Potentialenergy H• + H• H H -436 kJ/mol H2

15. Figure 2.1 1s 1s H H 2 H atoms: each electron "feels" attractive force of one proton H H H2 molecule: each electron "feels" attractive force of both protons

16. H H H H Internuclear distance Figure 2.1 repulsive forces increasefaster than attractive forcesat distances closer than 74 pm 74 pm Potentialenergy H• + H• H H -436 kJ/mol H2

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

18. Bonding in H2:The Valence Bond Model

19. Valence Bond Model • Electron pair can be shared when half-filled orbital of one atom overlaps in phase withhalf-filled orbital of another.

20. Valence Bond Model 1s 1s H H in-phase overlap of two half-filledhydrogen 1s orbitals H H  bond of H2

21. Valence Bond Model •  Bond: orbitals overlap along internuclear axis • Cross section of orbital perpendicular to internuclear axis is a circle. H H

22. Valence Bond Model of H2 The 1s orbitals of two separated hydrogen atoms are far apart. Essentially no interaction. Each electron is associated with a single proton.

23. Valence Bond Model of H2 As the hydrogen atoms approach each other, their 1s orbitals begin to overlap and each electron begins to feel the attractive force of both protons.

24. Valence Bond Model of H2 The hydrogen atoms are close enough so that appreciable overlap of the the two 1s orbitals occurs. The concentration of electron density in the region between the two protons is more readily apparent.

25. Valence Bond Model of H2 A molecule of H2. The two hydrogen 1s orbitals have been replaced by a new orbital that encompasses both hydrogens and contains both electrons.

26. Bonding in H2:The Molecular Orbital Model

27. Main Ideas • Electrons in a molecule occupy molecular orbitals (MOs) just as electrons in an atom occupy atomic orbitals (AOs). • Two electrons per MO, just as two electrons per AO. • Express MOs as combinations of AOs.

28. MO Picture of Bonding in H2 Linear combination of atomic orbitals method expresses wave functions of molecular orbitalsas sums and differences of wave functionsof atomic orbitals. • Two AOs yield two MOs • Bonding combinationAntibonding combination • MO = (H)1s + (H')1s 'MO = (H)1s - (H')1s

29. Fig. 2.6: Energy-Level Diagram for H2 MOs 1s 1s AO AO

30. Fig. 2.6: Energy-Level Diagram for H2 MOs MO antibonding * 1s 1s AO AO bonding  MO

31. Introduction to Alkanes:Methane, Ethane, and Propane CnH2n+2

32. The Simplest Alkanes • Methane (CH4) CH4 • Ethane (C2H6) CH3CH3 • Propane (C3H8) CH3CH2CH3 bp -160°C bp -89°C bp -42°C

33. sp3 HybridizationandBonding in Methane

34. Structure of Methane • tetrahedral • bond angles = 109.5° • bond distances = 110 pm • but structure seems inconsistent withelectron configuration of carbon

35. Electron configuration of carbon • only two unpaired electrons • should form bonds to only two hydrogen atoms • bonds should be at right angles to one another 2p 2s

36. sp3 Orbital Hybridization • Promote an electron from the 2s to the 2p orbital 2p 2s

37. sp3 Orbital Hybridization 2p 2p 2s 2s

38. sp3 Orbital Hybridization • Mix together (hybridize) the 2s orbital and the three 2p orbitals 2p 2s

39. sp3 Orbital Hybridization • 4 equivalent half-filled orbitals are consistent with four bonds and tetrahedral geometry 2p 2 sp3 2s

40. sp3 Orbital Hybridization

41. Nodal properties of orbitals p + – + s

42. Shape of sp3 hybrid orbitals • take the s orbital and place it on top of the p orbital p + – + s

43. + Shape of sp3 hybrid orbitals • reinforcement of electron wave in regions where sign is the same • destructive interference in regions of opposite sign s + p + –

44. – Shape of sp3 hybrid orbitals • orbital shown is sp hybrid • analogous procedure using three p orbitals and one s orbital gives sp3 hybrid • shape of sp3 hybrid is similar sp hybrid +

45. + – Shape of sp3 hybrid orbitals • hybrid orbital is not symmetrical • higher probability of finding an electron on one side of the nucleus than the other • leads to stronger bonds sp hybrid

46. + – – The C—H  Bond in Methane In-phase overlap of a half-filled 1s orbital of hydrogen with a half-filled sp3 hybrid orbital of carbon: + sp3 s H C gives a  bond. + H—C  C H

47. Justification for Orbital Hybridization • consistent with structure of methane • allows for formation of 4 bonds rather than 2 • bonds involving sp3 hybrid orbitals are stronger than those involving s-s overlap or p-p overlap

48. Bonding in Ethane

49. Structure of Ethane • tetrahedral geometry at each carbon • C—H bond distance = 110 pm • C—C bond distance = 153 pm C2H6 CH3CH3

50. The C—C  Bond in Ethane • In-phase overlap of half-filled sp3 hybridorbital of one carbon with half-filled sp3hybrid orbital of another. • Overlap is along internuclear axis to give a  bond.