chapter 10 chemical bonding ii
DESCRIPTION
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 PresentationTRANSCRIPT
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Chapter 10
Chemical Bonding II
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Lewis Structure Molecular Structure
Structure determines chemical properties
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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
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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
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Know how to spellthe names!
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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
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BF3, NO3−, H2CO
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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
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SO2
119°
Electron domain arrangement is not necessarily the same as the molecular structure.
Bent or V-shaped
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CH4
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NH3
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H2O
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PCl5
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SF4
To minimize repulsion, electron lone pairs are always placed in equatorial positions for trigonal bipyramidal geometry.
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BrF3
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XeF2
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SF6
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BrF5
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XeF4
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Polarity of a molecule
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How to quantify the polarity of a bond?
Dipole moment
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+ −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
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H — Fdipole
dipole moment of a bond ≠ 0 ↔ polar bond
dipole moment of a bond = 0 ↔ nonpolar bond
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The Pauling Electronegativity Values
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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
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v1
v2
v = v1 + v2
v3 = v2
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SO3
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120°
=
=
Net dipole moment = 0, nonpolar molecule
SO3
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109.5°
CCl4
=
=Net dipole moment = 0, nonpolar molecule
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Polarity of a molecule depends on the polarity of its bonds AND the geometry of the molecule.
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NH3
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How are electrons shared in covalent bonds?
Valence Bond Theory
Molecular Orbital Theory
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Valence Bond Theory:
Orbital Overlap as a Chemical Bond
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CH4
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CH4
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4 electron domains
sp3 hybridization
tetrahedral arrangement
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NH3
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sp3 on OH2O
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C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
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C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
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The σ Bonds in Ethylene
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C = C
Ethylene
H
H
H
H
~120°
All atoms are in the same plane
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A Carbon-Carbon Double Bond Consists of a σ and a π Bond
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3 electron domains
sp2 hybridization
trigonal planar arrangement
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H−C ≡ C−H
Acetylene
Linear molecule
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N2 :N≡N:
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2 electron domains
sp hybridization
Linear arrangement
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Single bond: σ
double bond: one σ, one π
triple bond: one σ, two π
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C
H
H
CH
=O
O C ≡ N:
::
¨¨
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PCl5
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The Orbitals Used to Form the Bonds in PCl5
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5 electron domains
dsp3 hybridization
trigonal bipyramidal arrangement
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SF6
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An Octahedral Set of d2sp3 Orbitals on Sulfur Atom
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6 electron domains
d2sp3 hybridization
octahedral arrangement
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paramagnetismdiamagnetism
paired electrons unpaired electrons
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Liquid O2 is paramagnetic
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How are electrons shared in covalent bonds?
Valence Bond Theory
Molecular Orbital Theory
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EH
Atoms → atomic orbitals
Molecules → molecular orbitals
Molecular Orbital (MO)≈ Linear Combination of Atomic Orbitals (LCAO)
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H2
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Electron configuration
(σ1s)2
Pauli principleand Hund’s ruleapply
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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
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Electron configuration
(σ1s)2 bond order = 1
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(σ1s)2(σ1s*)1 bond order = 0.5
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(σ1s)2(σ1s*)2 bond order = 0
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(σ1s)2(σ1s*)1 bond order = 0.5
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Li2
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(σ2s)2 bond order = 1
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(σ2s)2(σ2s*)2 bond order = 0
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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
+
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Homonuclear diatomic molecules/ions
Heteronuclear diatomic molecules/ions
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Problems Chapter 10
1-7 ,33,35,39,42,47,51,53,59,61,63,69,
71,85,86