Chapter 14. Covalent Bonding: Orbitals. Chapter 14 Slide 1 of 50

Transcription

1 Chapter 14 Covalent Bonding: Orbitals Chapter 14 Slide 1 of 50

2 Two bonding theories Two bonding theories will be discussed in this chapter: Valence Bond Theory Molecular Orbital Theory Chapter 14 Slide 2 of 50

3 Valence Bond Theory Valence Bond (VB) Theory states that a covalent bond is formed by the pairing of two electrons with opposing spins in the region of overlap of atomic orbitals between two atoms. This overlap region has a high electron charge density. In general, the more extensive the overlap between two atomic orbitals, the stronger is the bond between two atoms. The valence bond theory attempts to find the best approximation of optimal orbital overlap for all the bonds in a molecule. Chapter 14 Slide 3 of 50

4 Bonding In H 2 Chapter 14 Slide 4 of 50

5 Bonding In H 2 S Chapter 14 Slide 5 of 50

6 Several Important Points Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms. Bonding electrons are localized in the region of atomic orbital overlap. For orbitals with directional lobes, maximum overlap occurs when atomic orbitals overlap end to end; that is, a hypothetical line joining the nuclei of the bonded atoms passes through the region of maximum overlap. The molecular geometry depends on the geometric relationships among the atomic orbitals of the central atom that participate in bonding. Chapter 14 Slide 6 of 50

7 Hybridization sp 3 Hybridization Scheme Chapter 14 Slide 7 of 50

8 Bonding in Methane Chapter 14 Slide 8 of 50

9 Bonding in Ammonia Chapter 14 Slide 9 of 50

10 The sp 2 Hybridization Scheme Chapter 14 Slide 10 of 50

11 The sp Hybridization Scheme Chapter 14 Slide 11 of 50

12 Hybrid Orbitals Involving d Subshells This hybridization allows for expanded valence shell compounds. A 3s electron can be promoted to a 3d subshell which gives rise to a set of five sp 3 d hybrid orbitals. These molecules have a trigonal bipyramidal molecular geometry. One 3s electron and one 3p electron can be promoted to two 3d subshells which gives rise to a set of six sp 3 d 2 hybrid orbitals. These molecules have an octahedral molecular geometry. Chapter 14 Slide 12 of 50

13 The sp 3 d Hybrid Orbitals Chapter 14 Slide 13 of 50

14 The sp 3 d 2 Hybrid Orbitals Chapter 14 Slide 14 of 50

15 Hybrid Orbitals and Their Geometric Orientations Chapter 14 Slide 15 of 50

16 Hybrid Orbitals and Multiple Covalent Bonds Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (σ) bonds. All single bonds are sigma bonds. No nodal plane along inter-nuclear axis A bond formed by parallel, or side-by-side, orbital overlap is called a pi (π) bond. One nodal plane along inter-nuclear axis sp 2 -hybrid Chapter 14 Slide 16 of 50

17 Descriptions of Ethylene Rotational barrier for double bond Chapter 14 Slide 17 of 50

18 Valence Bond Theory of the Bonding in Acetylene sp-hybrid Chapter 14 Slide 18 of 50

19 Geometric Isomerism Geometric isomers are isomers that differ only in the geometric arrangement of certain substituent groups. Two main types of geometric isomers: cis: substituent groups are on the same side trans: substituent groups are on opposite sides Each compound is distinctly different in both physical and chemical properties. Usually formed across double bonds, in cyclic and square planar compounds. Chapter 14 Slide 19 of 50

20 Geometric Isomerism In 2-Butene Chapter 14 Slide 20 of 50

21 Characteristics of Molecular Orbitals Molecular orbitals (MOs) are mathematical equations that describe the regions in a molecule where there is a high probability of finding electrons. Bonding molecular orbitals (σ, π) are at a lower energy level than the separate atomic orbitals and have a high electron probability, or electron charge density. Antibonding molecular orbitals (σ*, π*) are at a higher energy level than the separate atomic orbitals and places a high electron probability away from the region between the bonded atoms. Chapter 14 Slide 21 of 50

22 The 1s Orbital Ψ (r,θ,φ) = R(r) Υ 0,0 (θ,φ) Υ 0,0 (θ,φ) = 1/2π 1/2 Chapter 14 Slide 22 of 50

23 Constructive Interference Chapter 14 Slide 23 of 50

24 Destructive Interference Chapter 14 Slide 24 of 50

25 Molecular Orbitals and Bonding in the H 2 Molecule Chapter 14 Slide 25 of 50

26 H 2 M.O. energy level diagram H 2 Bond order = ½ ( # of bonding electrons - # of antibonding electrons ) Electron configuration of H 2 : (σ 1s ) 2 B.O. of H 2 = ½ (2-0) = 1 Bond energy = 435 kj/mol Bond length = 74 pm Chapter 14 Slide 26 of 50

27 M.O. energy level diagram H 2 + H 2 - He 2 Chapter 14 Slide 27 of 50

28 H 2, H 2, + H 2, - He 2 Species Electron configuration B.O. Bond energy (kj/mol) Bond length (pm) H 2 (σ 1s ) H 2 + (σ 1s ) 1 ½ H 2 - (σ 1s ) 2 (σ 1s *) 1 ½ He 2 (σ 1s ) 2 (σ 1s *) Chapter 14 Slide 28 of 50

29 Hetero-nuclear Diatomic Molecule Lewis Structure Chapter 14 Slide 29 of 50

30 2nd Period Homo-nuclear Diatomic Molecules Electron configuration of Li 2 : KK(σ 2s ) 2 B.O. of Li 2 = ½ (2-0) = 1 Bond length = 267 pm Chapter 14 Slide 30 of 50

31 _ Molecular Orbitals Formed by Combining 2p Atomic Orbitals + _ _ + _ + + _ + _ 1 node inter-nuclear axis _ + 1 node along inter-nuclear axis Chapter 14 Slide 31 of 50

32 B 2 M.O. energy level diagram Diamagnetic Paramagnetic Chapter 14 Slide 32 of 50

33 O 2 M.O. energy level diagram Paramagnetic Chapter 14 Slide 33 of 50

34 Paramagnetism of Oxygen Chapter 14 Slide 34 of 50

35 Molecular Orbitals of Homo-nuclear Diatomic Molecules of 2 nd Period Chapter 14 Slide 35 of 50

36 Bonding in Benzene The structure of benzene (C 6 H 6 ), discovered by Michael Faraday in 1825, was not figured out until 1865 by F.A. Kekulé. Kekulé discovered that benzene has a cyclic structure and he proposed that a hydrogen atom was attached to each carbon atom and that alternating single and double bonds joined the carbon atoms together. This kind of structure gives rise to two important resonance hybrids and leads to the idea that all three double bonds are delocalized across all six carbon atoms. Chapter 14 Slide 36 of 50

37 The σ-bonding Framework Chapter 14 Slide 37 of 50

38 The π-molecular Orbitals of Benzene _ E node + + _ node π-m.o. of benzene Chapter 14 Slide 38 of 50 +

39 The Molecular Orbitals of Benzene 3 nodes 2 nodes Chapter 14 Slide 39 of 50

40 Aromatic Compounds Many of the first benzene-like compounds discovered had pleasant odors and hence acquired the name aromatic. In modern chemistry, the term aromatic compound simply refers to a substance with a ring structure and with bonding characteristics and properties related to those of benzene. All organic compounds that are not aromatic are called aliphatic compounds. Chapter 14 Slide 40 of 50

41 Some Representative Aromatic Compounds Chapter 14 Slide 41 of 50

42 p-aminobenzoic acid Chapter 14 Slide 42 of 50

43 The absorption spectrum Chapter 14 Slide 43 of 50

44 Conjugated Double Bonds E Antibonding Bonding π-m.o. Chapter 14 Slide 44 of 50

45 Band Theory This is a quantum-mechanical treatment of bonding in metals. The spacing between energy levels is so minute in metals that the levels essentially merge into a band. When the band is occupied by valence electrons, it is called a valence band. A partially filled or low lying empty band of energy levels, which is required for electrical conductivity, is a conduction band. Band theory provides a good explanation of metallic luster and metallic colors. Chapter 14 Slide 45 of 50

46 Energy vs N Antibonding bonding Chapter 14 Slide 46 of 50

47 The 2s Band in Lithium Metal Anti-bonding Conduction band e- e- Bonding Valence band Chapter 14 Slide 47 of 50

48 Band Overlap in Magnesium Conduction band Valence band Chapter 14 Slide 48 of 50

49 Band Structure of Insulators and Semiconductors Chapter 14 Slide 49 of 50

50 Temperature vs Resistance Chapter 14 Slide 50 of 50