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Chapter 9B: Advanced Bonding Theory
CHEM 101 Fall 2020
Dr. Lauren Genova
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Chapter 9
By the end of this chapter, you will be able to:
• Differentiate between a sigma (σ) bond and pi (𝜋) bond
• Determine the hybrid orbitals (hybridization) associated with various molecular geometries
Chapter 9B: Learning Objectives
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Fig. 10-1, p.439
• Valence bond theory states that covalent bonds are formed by overlapping and pairing electrons in atomic orbitals:
H: 1s
H: 1s
Valence Bond Theory
• We say that the orbitals on 2 different atoms overlap when a portion of one orbital and a portion of a second orbital occupy the same region of space.
• If the overlap is constructive, electron density will increase in the region between the two nuclei, resulting in a bonding orbital (system is stabilized)
(max stability)
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• Two types of bonds are permitted in valence bond theory: – Sigma (σ) bonds – formed by end-to-end overlap of two
atomic orbitals (which may be s, p, or hybrid)
– Pi (𝛑) bonds – formed by side-to-side overlap of two atomic p orbitals
1.) Sigma and Pi Bonding
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Sigma Bonding • Sigma (σ) bond: Covalent bond in which the highest
electron density lies between the two atoms along the bond axis. – All single bonds are sigma bonds – Free rotation is allowed along the axis of a sigma bond
bond axis
bond axis
bond axis 5
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• Pi (𝛑) bond: Covalent bond in which electron density is greatest around—not along—the bonding axis. – A double bond contains 1 sigma and 1 pi bond – A triple bond contains 1 sigma and 2 pi bonds – Free rotation is not allowed with a pi bond (thus adding
rigidity to molecules which contain them) • This rigidity imparted by pi bonds has a major influence on
the conformations of complex molecules in solution
Pi Bonding
Double bond Triple bond
𝛑 σσ
σ 𝛑 σσ σ 𝛑
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Free Rotation Differences between σ and 𝛑
Image courtesy of: https://chemistry.stackexchange.com/questions/68480/why-can-a-sigma-bond-rotate
(σ)
(σ)
(𝛑)
bond axis
bond axis
bond axis
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Practice: Identifying σ and 𝛑 bonds 2.) How many sigma (σ) and pi (𝛑) bonds are in this molecule?
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• The overlap of atomic orbitals from valence bond theory results in predicted bond angles that differ from the observed bond angles (from VSEPR theory)
• Example: H2O
• We require a more detailed model to explain this discrepancy
H: 1s
O:H: 1s
Expected bond angle from p orbitals (valence bond theory) = 90o
Problem with Valence Bond Theory
Actual bond angle (VSEPR) = 104.5o104.5
o
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Observed reality: Tetrahedral (109.5o)
Hybridization • Hybridization is the concept of mixing atomic orbitals into new
hybrid orbitals (with different energies/shapes from the original atomic orbitals) suitable for the pairing of e– – Atomic orbitals can hybridize to maximize
distances between e–
– This helps to account for the problem with valence bond theory: namely, that the straight overlap of atomic orbitals doesn’t account for all observed molecular shapes.
Valence bond theory prediction:
Bond angles = 90o apart
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A Closer Look at Hybridization
Hybridization – the mixing of atomic orbitals to generate new sets of orbitals that form covalent bonds with other atoms.
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Scientist Spotlight Dr. Herdeline “Digs” Ardoña Assistant Professor, Department of Chemical and Biomolecular Engineering, University of California, Irvine (Irvine, CA)
Topic: Hybridization (with a focus on pi-stacking)
Favorite quote: “Winners never quit on something that makes and/or will make them happy.”
Email: hardona@uci.edu
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Hybridization and Geometry • Goal for this unit:
to determine the hybridization of a molecule
• How? By identifying its electron-pair geometry!
Recall: electron-pair geometry includes bonding pairs (electrons shared between 2 atoms) AND nonbonding pairs (lone pairs)
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Number of connections
Type of hybridization
Molecular geometry
2 sp Linear
3 sp2 Trigonal planar
4 sp3 Tetrahedral
5 sp3d Trigonal bipyramidal
6 sp3d2 Octahedral
Dr. G’s method for determining hybridization: 1.) Count the number of connections (number of orbitals) around the central atom whose hybridization you have been asked to identify
• Order when counting number of connections: s, p, p, p, d, d • INCLUDE THE LONE PAIRS while you’re counting!
2.) Number of connections = hybridization!
3.) Hybridization and Geometry
• Goal for this unit: to determine the hybridization of a molecule
• How? By identifying its electron-pair geometry!
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These two unhybridized p orbitals ultimately lead to a triple bond (one σ and two 𝛑 bonds)
Linear (2 connections): sp Hybridization
• Formed by mixing one s and one p orbital
• Example: each C in H–C≡C–H
s p
s p
s, p, p, p, d, d
sp hybridized
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Trigonal Planar (3 connections): sp2 Hybridization This unhybridized p orbital ultimately leads to a double bond (one σ and one 𝛑 bond)
• Formed by mixing one s and two p orbitals 16
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Trigonal Planar (3 connections): sp2 Hybridization
• Formed by mixing one s and two p orbitals • Examples: the C in H2CO, the O in H2CO
s
p p
s p
p
s, p, p, p, d, d sp2 hybridized
This unhybridized p orbital ultimately leads to a double bond (one σ and one 𝛑 bond)
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Tetrahedral (4 connections): sp3 Hybridization
s
p
p
p
• Formed by mixing one s and three p orbitals • Example: CH4 – four sigma bonds
s, p, p, p, d, d
sp3 hybridized
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Other sp3 Hybrid Examples
• Formed by mixing one s and three p orbitals • Example: NH3 (3 sigma bonds, 1 lone pair)
Remember to count the lone pairs (non-bonding e–)!s
p
p
p
s, p, p, p, d, d
sp3 hybridized
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Other sp3 Hybrid Examples
• Formed by mixing one s and three p orbitals • Example: H2O (2 sigma bonds, 2 lone pairs)
Remember to count the lone pairs (non-bonding e–)!s
pp
p
s, p, p, p, d, d
sp3 hybridized
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Trigonal Bipyramidal (5 connections): sp3d Hybridization
• Formed by mixing one s, three p, and one d orbital
• Example: PF5 – six sigma bonds
s
p
p p
d s, p, p, p, d, d
sp3d hybridized
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• Formed by mixing one s, three p, and two d orbitals
• Example: SF6 – six sigma bonds
Octahedral (6 connections): sp3d2 Hybridization
s
p
p p
d
d s, p, p, p, d, d
sp3d2 hybridized
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4.) Describe the hybridization around each central atom in the amino acid glycine.
O
C O H
H
H NH H
sp3
sp3
sp3
sp2 ••
•• ••C
1 2
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Hybridization Practice
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