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Transcript of Chapter 14ocw.aca.ntu.edu.tw/ocw_files/099S125/ch14.pdf · Chapter 14 Slide 6 of 50 Several...
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Chapter 14 Slide 1 of 50
Chapter 14
Covalent Bonding: Orbitals
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Chapter 14 Slide 2 of 50
Two bonding theories
Two bonding theories will be discussed in this chapter:
• Valence Bond Theory• Molecular Orbital Theory
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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.
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Chapter 14 Slide 4 of 50
Bonding In H2
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Chapter 14 Slide 5 of 50
Bonding In H2S
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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.
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Chapter 14 Slide 7 of 50
sp3 Hybridization Scheme
Hybridization
109.50
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Chapter 14 Slide 8 of 50
Bonding in Methane
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Chapter 14 Slide 9 of 50
Bonding in Ammonia
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Chapter 14 Slide 10 of 50
The sp2 Hybridization Scheme
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Chapter 14 Slide 11 of 50
The sp Hybridization Scheme
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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.
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Chapter 14 Slide 13 of 50
The sp3d Hybrid Orbitals
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Chapter 14 Slide 14 of 50
The sp3d2 Hybrid Orbitals
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Chapter 14 Slide 15 of 50
Hybrid Orbitals and TheirGeometric Orientations
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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
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Chapter 14 Slide 17 of 50
Descriptions of Ethylene
Rotational barrier for double bond
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Chapter 14 Slide 18 of 50
Valence Bond Theory of the Bonding in Acetylene
sp-hybrid
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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.
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Chapter 14 Slide 20 of 50
Geometric IsomerismIn 2-Butene
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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.
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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
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Chapter 14 Slide 23 of 50
Constructive Interference
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Chapter 14 Slide 24 of 50
Destructive Interference
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Chapter 14 Slide 25 of 50
Molecular Orbitals and Bondingin the H2 Molecule
++
+ -
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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
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Chapter 14 Slide 27 of 50
M.O. energy level diagram
H2-H2
+
He2
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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
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Chapter 14 Slide 29 of 50
Hetero-nuclear Diatomic Molecule
Lewis Structure
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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
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Chapter 14 Slide 31 of 50
Molecular Orbitals Formed byCombining 2p Atomic Orbitals
++ +
____
_ + +__
1 node ⊥ inter-nuclear axis
+ +_ _
+_
1 node along inter-nuclear axis
_ +
+__
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Chapter 14 Slide 32 of 50
B2 M.O. energy level diagram
Diamagnetic Paramagnetic
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Chapter 14 Slide 33 of 50
O2 M.O. energy level diagram
Paramagnetic
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Chapter 14 Slide 34 of 50
Paramagnetism of Oxygen
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Chapter 14 Slide 35 of 50
Molecular Orbitals of Homo-nuclearDiatomic Molecules of 2nd Period
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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.
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Chapter 14 Slide 37 of 50
The σ-Bonding Framework
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Chapter 14 Slide 38 of 50
The π-Molecular Orbitals of Benzene
E
+
++ _
_
node
node
π-M.O. of benzene
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Chapter 14 Slide 39 of 50
The Molecular Orbitals of Benzene
3 nodes
2 nodes
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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.
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Chapter 14 Slide 41 of 50
Some RepresentativeAromatic Compounds
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Chapter 14 Slide 42 of 50
p-Aminobenzoic acid
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Chapter 14 Slide 43 of 50
The absorption spectrum
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Chapter 14 Slide 44 of 50
Conjugated Double Bonds
E
π-M.O.
Bonding
Anti-bonding
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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.
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Chapter 14 Slide 46 of 50
Energy vs N
bonding
Anti-bonding
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Chapter 14 Slide 47 of 50
The 2s Band in Lithium Metal
Bonding
Anti-bonding
e- e-Valence band
Conduction band
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Chapter 14 Slide 48 of 50
Band Overlap in Magnesium
Valence band
Conduction band
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Chapter 14 Slide 49 of 50
Band Structure of Insulatorsand Semiconductors
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Chapter 14 Slide 50 of 50
Temperature vs Resistance