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Chapter 14 Slide 1 of 50 Chapter 14 Covalent Bonding: Orbitals
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Transcript of Chapter 14ocw.aca.ntu.edu.tw/ocw_files/099S125/ch14.pdf · Chapter 14 Slide 6 of 50 Several...
Chapter 14 Slide 1 of 50
Chapter 14
Covalent Bonding: Orbitals
Chapter 14 Slide 2 of 50
Two bonding theories
Two bonding theories will be discussed in this chapter:
• Valence Bond Theory• Molecular Orbital Theory
Chapter 14 Slide 3 of 50
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 4 of 50
Bonding In H2
Chapter 14 Slide 5 of 50
Bonding In H2S
Chapter 14 Slide 6 of 50
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 7 of 50
sp3 Hybridization Scheme
Hybridization
109.50
Chapter 14 Slide 8 of 50
Bonding in Methane
Chapter 14 Slide 9 of 50
Bonding in Ammonia
Chapter 14 Slide 10 of 50
The sp2 Hybridization Scheme
Chapter 14 Slide 11 of 50
The sp Hybridization Scheme
Chapter 14 Slide 12 of 50
Hybrid Orbitals Involvingd Subshells
• This hybridization allows for expanded valence shellcompounds.
• A 3s electron can be promoted to a 3d subshell which gives rise to a set of five sp3d 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 sp3d2
hybrid orbitals. These molecules have an octahedralmolecular geometry.
Chapter 14 Slide 13 of 50
The sp3d Hybrid Orbitals
Chapter 14 Slide 14 of 50
The sp3d2 Hybrid Orbitals
Chapter 14 Slide 15 of 50
Hybrid Orbitals and TheirGeometric Orientations
Chapter 14 Slide 16 of 50
Hybrid Orbitals andMultiple 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
sp2-hybrid
Chapter 14 Slide 17 of 50
Descriptions of Ethylene
Rotational barrier for double bond
Chapter 14 Slide 18 of 50
Valence Bond Theory of the Bonding in Acetylene
sp-hybrid
Chapter 14 Slide 19 of 50
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 andsquare planar compounds.
Chapter 14 Slide 20 of 50
Geometric IsomerismIn 2-Butene
Chapter 14 Slide 21 of 50
Characteristics ofMolecular 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 orbitalsand places a high electron probability away from the region between the bonded atoms.
Chapter 14 Slide 22 of 50
The 1s Orbital
Ψ (r,θ,φ) = R(r) Υ0,0(θ,φ)
Υ0,0(θ,φ) = 1/2π1/2
Chapter 14 Slide 23 of 50
Constructive Interference
Chapter 14 Slide 24 of 50
Destructive Interference
Chapter 14 Slide 25 of 50
Molecular Orbitals and Bondingin the H2 Molecule
++
+ -
Chapter 14 Slide 26 of 50
H2 M.O. energy level diagram
Bond order = ½ ( # of bonding electrons - # of anti-bonding electrons )
Electron configuration of H2 : (σ1s)2
B.O. of H2 = ½ (2 - 0) = 1
Bond energy = 435 kJ/mol
Bond length = 74 pm
H2
Chapter 14 Slide 27 of 50
M.O. energy level diagram
H2-H2
+
He2
Chapter 14 Slide 28 of 50
H2, H2, +H2,
-He2
--0(σ1s)2(σ1s*)2He2
108238½(σ1s)2(σ1s*)1H2-
106269½(σ1s)1H2+
744351(σ1s)2H2
Bond length (pm)
Bond energy (kJ/mol)
B.O.Electron configuration
Species
Chapter 14 Slide 29 of 50
Hetero-nuclear Diatomic Molecule
Lewis Structure
Chapter 14 Slide 30 of 50
Electron configuration of Li2 : KK(σ2s)2
B.O. of Li2 = ½ (2 - 0) = 1Bond length = 267 pm
2nd Period Homo-nuclear Diatomic Molecules
Chapter 14 Slide 31 of 50
Molecular Orbitals Formed byCombining 2p Atomic Orbitals
++ +
____
_ + +__
1 node ⊥ inter-nuclear axis
+ +_ _
+_
1 node along inter-nuclear axis
_ +
+__
Chapter 14 Slide 32 of 50
B2 M.O. energy level diagram
Diamagnetic Paramagnetic
Chapter 14 Slide 33 of 50
O2 M.O. energy level diagram
Paramagnetic
Chapter 14 Slide 34 of 50
Paramagnetism of Oxygen
Chapter 14 Slide 35 of 50
Molecular Orbitals of Homo-nuclearDiatomic Molecules of 2nd Period
Chapter 14 Slide 36 of 50
Bonding in Benzene
• The structure of benzene (C6H6), 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 37 of 50
The σ-Bonding Framework
Chapter 14 Slide 38 of 50
The π-Molecular Orbitals of Benzene
E
+
++ _
_
node
node
π-M.O. of benzene
Chapter 14 Slide 39 of 50
The Molecular Orbitals of Benzene
3 nodes
2 nodes
Chapter 14 Slide 40 of 50
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 compoundsimply 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 41 of 50
Some RepresentativeAromatic Compounds
Chapter 14 Slide 42 of 50
p-Aminobenzoic acid
Chapter 14 Slide 43 of 50
The absorption spectrum
Chapter 14 Slide 44 of 50
Conjugated Double Bonds
E
π-M.O.
Bonding
Anti-bonding
Chapter 14 Slide 45 of 50
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 46 of 50
Energy vs N
bonding
Anti-bonding
Chapter 14 Slide 47 of 50
The 2s Band in Lithium Metal
Bonding
Anti-bonding
e- e-Valence band
Conduction band
Chapter 14 Slide 48 of 50
Band Overlap in Magnesium
Valence band
Conduction band
Chapter 14 Slide 49 of 50
Band Structure of Insulatorsand Semiconductors
Chapter 14 Slide 50 of 50
Temperature vs Resistance
