Bonding General Concepts

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Chapter 8 Bonding General Concepts
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Valence Electrons
Main Group Elements
1A 2A 3A 4A 5A 6A 7A 8
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Ionic Bonding Gaining (anion) and Losing
(cation) Electrons to attain Inert gas
electronic structure
More on that later . . .
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Covalent Bonding Sharing electrons to attain
inert gas electron configurations.
E
d
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Best bond distance Lowest energy
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In order for a bond to occur, the
Electron-proton interaction must be Greater than
proton-proton and electron -electron repulsion.
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Electronegativity
The ability of an atom to attract Electrons
towards itself.
There is a measureable Bond energy
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A
B
C
Expected A Bond Energy B energy C energy
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D Actual A - Expected A
If D is not zero, one atom must Pull electrons
more than the other
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d -
d
A polarized bond
A dipole moment
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Electronegativity
2.1 H
4.0 F
1.0 Li
3.5 O
3.0 Cl
2.5 C
0.7 Cs
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A generalization is that a D electronegarivity of
1.7 or greater is classified as ionic.
Ionic
Covalent Polar Covalent
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Cl
K
3.16
0.82
D 3.16 - 0.82 2.34
IONIC
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Dipole
S
Br
2.58
2.96
D 2.96 - 2.58 0.38
Covalent
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3.44
2.22
DE 1.22
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No net dipole
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Ions Electron Configuration and sizes
1. Predict Ions from position on periodic Table.
Rb Br S Ca Al
P
Zn Cu Ag Hg
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2. Predict ions sizes
Na Mg Al O F
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Use of thermodynamics for prediction Of lattice
energies NaCl
IP
DHVap
Na(s) Na (g) Na (g)
EA
DHDis
1/2 Cl2(g) Cl (g) Cl - (g)
NaCl (s)
More later!
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Covalent Bonding Sharing electrons to attain
inert gas electron configurations.
Inner
Lewis Electron Dot Structures
Mostly between groups 4 - 7 and H
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Two atoms with unfilled outer Orbitals can share
electrons.
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Rules Hydrogen gets 1 bond to look like
helium. Other atoms try to have filled s and
p-orbitals.
H
O
H
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Rules Hydrogen gets 1 bond to look like
helium. Other atoms try to have filled s and
p-orbitals.
H
O
H
First assign the valence electrons
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Rules Hydrogen gets 1 bond to look like
helium. Other atoms try to have filled s and
p-orbitals.
H
O
H
First assign the valence electrons Now start
sharing
O
H
H
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A more general approach Simply count up (and
hold) the valence electrons. Place them
within bonds (2 at a time) until all bonding
situations seem filled.
4 6 10 electrons
C O
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A more general approach Simply count up (and
hold) the valence electrons. Place them
within bonds (2 at a time) until all bonding
situations seem filled.
4 6 10 electrons
C O
2
C O
8 electrons
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2
C O
8 electrons
by electron pair counting, neither has 8
C O
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Example
O H O S O H O
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Examples
H O C H O
O H O Cl O O
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Resonance
The ability to draw more than one structure for
a molecule by only moving electrons.
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Resonance
The ability to draw more than one structure for
a molecule by only moving electrons.
Consider the carbonate ion
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- 2
O C O O
Lets consider the electrons available for bonding
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- 2
6
4
6
O C O O
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Lets consider the electrons available for bonding
Total 66642 24
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24 electrons . . . lets put them in.
O C O O
24-6 18 left
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24 electrons . . . lets put them in.
O C O O
24-6 18 left
Lets put these around the outside atoms
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18 to distribute
O C O O
Are all of the atoms OK?
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O C O O
Now all the atoms are satisfied . . . but why did
we use that particular lone pair?
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Resonance Structures
O C O O
O C O O
O C O O
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Exceptions to the Octet Rule
Sometimes less Sometimes more
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Group II BeCl2
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Cl Be Cl
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7 2 7 16
Cl Be Cl
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Cl Be Cl
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Cl Be Cl
Not happy
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Cl Be Cl
One might suggest a stabilization Where the
central atom shares additional electrons, but
evidence Still shows only two bonds.
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F B F F
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7 3 7
F B F F
e 24
7
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F B F F
e 24
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F B F F
e 24
Again an unfulfilled central atom.
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F B F F
e 24
How about resonance?
No matter how you show it, however, There are
only three bonds.
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F
e 34
F S F
F
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F
e 34
F S F
F
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F
e 34
F S F
F
Still have 2 electrons to deal with!
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F
e 34
F S F
F
The central atom is associated with 10 electrons
. .an expanded shell.
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-
I I I
This has 21 1 electrons 22
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I I I
This has 21 1 electrons 22
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I I I
This has 21 1 electrons 22
After we have placed electrons around the
terminal atoms, we go back to the central atom .
. .with 6 more electrons.
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I I I
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What are we finding for shapes?
bond or lone pair
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What are we finding for shapes?
bond or lone pair
H
F
O
Be
B
F
Cl
Cl
F
H
F
F
S
F
F
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F F Xe F F
36 electrons
4 electrons
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Electron Dot Structures and Molecular Structure
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Tetrahedral Orientation of Bonds and Lone Pairs
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Tetrahedral Orientation of Bonds and Lone Pairs.
What you see
Bent
What might we predict about the bond angle?
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electron pair repulsions are greater than
bond pair repulsions.
Thus, the bond angle for the bonds gets smaller
as the non-bonding angle gets larger.
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Tetrahedral Orientation of Bonds and Lone Pairs
Three bonds and one lone pair.
What you see is pyramidal.
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Tetrahedral Orientation of Bonds and Lone Pairs.
Three bonds and one lone pair.
You dont see
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Tetrahedral Orientation of Bonds and Lone Pairs.
Four bonds.
You see it all!
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Planar orientation of three bonds and or lone
pairs.
BF3
With three bonds
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Planar orientation of three bonds and or lone
pairs.
Planar, Trigonal
BH3
With three bonds
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Planar orientation of two bonds and one lone
pair.
You see bent
O
S
O
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Linear orientation of two bonds
CO
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Linear orientation of two bonds
I3-
You dont see
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Four Bonds and One Lone Pair
SF4
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Four Bonds and Two Lone Pairs
XeF4
You cant see
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Formal Charges
Bonding Pair Electrons
Group
FG GN - LPe - 1/2 BPe
Lone Pair Electrons
Formal Charge
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CO2 Which structure is better?
O C O
O C O
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CO2 Which structure is better?
O C O
FC
6-4-20
O C O
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CO2 Which structure is better?
O C O
FC
4-0-40
0
O C O
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CO2 Which structure is better?
O C O
FC
0
0
0
O C O
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CO2 Which structure is better?
O C O
0
0
0
O C O
6-6-1
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CO2 Which structure is better?
O C O
0
0
0
O C O
4-0-4
1
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CO2 Which structure is better?
O C O
0
0
0
O C O
6-2-3 1
O
-1
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How about sulfuric acid?
O H O S O H O
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One of the Important take-home messages. If we
think of bonds between atoms as sticks, we
might imagine
Bond strength
Bond length