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Chapter 10 Chemical Bonding II. Lewis Structure Molecular Structure. Structure determines chemical properties. Electron domain/group: area where electrons appear in Lewis structures. It can be electron lone pairs, single bonds, double bonds, triple bonds, or single electrons. - PowerPoint PPT Presentation
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Chapter 10
Chemical Bonding II
Lewis Structure Molecular Structure
Structure determines chemical properties
Electron domain/group: area where electrons appear in Lewis structures.
It can be electron lone pairs, single bonds,
double bonds, triple bonds, or single electrons.
H2O, NH3, CH4, O2, N2, SCl2, CCl4, PCl3, NO+, NH4
+, CO, CO2
Valence Shell Electron Pair Repulsion (VSEPR) model
The lowest energy arrangement of a given number of electron
domains is the one that minimizes the repulsions among them.
The shape of ABn molecules or ions depend on the number of
electron domains surrounding the central A atom.
number of electron domains: 2 to 6
Know how to spellthe names!
How to predict geometry of a molecule?
1) Draw the Lewis structure of the molecule or ion, and countthe number of electron domains around the central atom.
2) Determine the electron domain arrangement by arranging theelectron domains about the central atom so that the repulsionsamong them are minimized.
3) Use the arrangement of the bonded atoms to determine the molecular geometry.
CO2
BF3, NO3−, H2CO
Electron domains for multiple bonds exert a greaterrepulsion force on adjacent electron domains than do electron domains for single bonds.
lone pair-lone pair > lone pair-bonding pair > bonding pair- bonding pair
SO2
119°
Electron domain arrangement is not necessarily the same as the molecular structure.
Bent or V-shaped
CH4
NH3
H2O
PCl5
SF4
To minimize repulsion, electron lone pairs are always placed in equatorial positions for trigonal bipyramidal geometry.
BrF3
XeF2
SF6
BrF5
XeF4
Polarity of a molecule
How to quantify the polarity of a bond?
Dipole moment
+ −Dipole
Dipole has a magnitude and a direction — vector
Magnitude (length) of a dipole — dipole moment
μ = qr
q — charge, r — distance between + and − charge
H — Fdipole
dipole moment of a bond ≠ 0 ↔ polar bond
dipole moment of a bond = 0 ↔ nonpolar bond
The Pauling Electronegativity Values
Polarity of a molecule
dipole moment of a molecule ≠ 0 ↔ polar molecule
dipole moment of a molecule = 0 ↔ nonpolar molecule
dipole of a molecule = sum of all the bond dipoles
v1
v2
v = v1 + v2
v3 = v2
SO3
120°
=
=
Net dipole moment = 0, nonpolar molecule
SO3
109.5°
CCl4
=
=Net dipole moment = 0, nonpolar molecule
Polarity of a molecule depends on the polarity of its bonds AND the geometry of the molecule.
NH3
How are electrons shared in covalent bonds?
Valence Bond Theory
Molecular Orbital Theory
Valence Bond Theory:
Orbital Overlap as a Chemical Bond
CH4
CH4
4 electron domains
sp3 hybridization
tetrahedral arrangement
NH3
sp3 on OH2O
C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
The σ Bonds in Ethylene
C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
A Carbon-Carbon Double Bond Consists of a σ and a π Bond
3 electron domains
sp2 hybridization
trigonal planar arrangement
H−C ≡ C−H
Acetylene
Linear molecule
N2 :N≡N:
2 electron domains
sp hybridization
Linear arrangement
Single bond: σ
double bond: one σ, one π
triple bond: one σ, two π
C
H
H
CH
=O
O C ≡ N:
::
¨¨
PCl5
The Orbitals Used to Form the Bonds in PCl5
5 electron domains
dsp3 hybridization
trigonal bipyramidal arrangement
SF6
An Octahedral Set of d2sp3 Orbitals on Sulfur Atom
6 electron domains
d2sp3 hybridization
octahedral arrangement
paramagnetismdiamagnetism
paired electrons unpaired electrons
Liquid O2 is paramagnetic
How are electrons shared in covalent bonds?
Valence Bond Theory
Molecular Orbital Theory
EH
Atoms → atomic orbitals
Molecules → molecular orbitals
Molecular Orbital (MO)≈ Linear Combination of Atomic Orbitals (LCAO)
H2
Electron configuration
(σ1s)2
Pauli principleand Hund’s ruleapply
Bond order = ½ (number of bonding electrons − number of antibonding electrons)
If bond order > 0, the molecule is stable
If bond order = 0, the molecule is not stable
bond order = 1 → single bond
bond order = 2 → double bond
bond order = 3 → triple bond
Electron configuration
(σ1s)2 bond order = 1
(σ1s)2(σ1s*)1 bond order = 0.5
(σ1s)2(σ1s*)2 bond order = 0
(σ1s)2(σ1s*)1 bond order = 0.5
Li2
(σ2s)2 bond order = 1
(σ2s)2(σ2s*)2 bond order = 0
O2, F2, Ne2: (σ2s) (σ2s*) (σ2p) (π2p) (π2p
*) (σ2p*)
B2, C2, N2: (σ2s) (σ2s*) (π2p) (σ2p) (π2p
*) (σ2p*)
1)Electron configuration.
2)Bond order → stable molecule/ion?
3)Paramagnetic or Diamagnetic?
O2, F2, N2, N2−, N2
+
Homonuclear diatomic molecules/ions
Heteronuclear diatomic molecules/ions
Problems Chapter 10
1-7 ,33,35,39,42,47,51,53,59,61,63,69,
71,85,86