Lewis
Structures
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Lewis Structures
•Are models
•The representations of the electron arrangements in atoms,
ions, or molecules by showing the valence electrons as dots
placed around the symbols for the elements
•Also called Lewis Dot Diagrams or Electron Dot Diagrams
•Can be drawn for atoms, molecules, anions, cations or ionic
compounds
•Useful when determining the geometry or shape of a
molecule
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Lewis Structures
Why are they important?
•You can visualize the electrons involved in
chemical bonds
•You can gain a greater understanding of how
chemical bonds form
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Lewis Structures
1) Count up total number of valence electrons
2) Decide on arrangement of atoms and connect all atoms with single bonds
- “least electronegative atoms usually in the middle
- “single” atoms usually in center;
C always in center,
H always on outside.
3) Complete octets on exterior atoms, lone pair electrons (not H, though)
4) Check:
- valence electrons math with Step 1
- all atoms (except H) have an octet; if not, try multiple bonds
- any extra electrons? Put on central atom
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Lewis Structures
How Are They Drawn?
1) Count up all of the valence electrons for each atom in the formula
for anions, add the charge of the ion to the number of valence electrons
for cations, subtract the charge of the ion from the number of electrons
2) Determine the number of octet electrons the formula should have
3) Determine the Number of Bonding Electrons
Subtract valence electrons from octet electrons
4) Determine the Number of Bonds
Divide # Bonding electrons by 2
5) Draw the Structure with the correct Number of Bonds, least electronegative
element is usually the central atom
Bond all atoms together by single bonds, then add in the multiple bonds until
the rules in the notes are followed
6) Determine number of Lone Pair Electrons and arrange them around the
atoms until the octet rule is satisfied for all atoms (except Hydrogen)
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Some Examples
CO2
1) Valence Electrons = 16



Carbon is in group 4, 4 valence
Oxygen is in group 6, 6 valence X 2 atoms = 12 total
4 + 12 = 16
2) Octet Electrons = 8 ea. X 3 = 24
3) Bonding Electrons = 24 – 16 = 8
4) Number of Bonds = (8/2) = 4
5) O = C = O
6) Non bonding electrons = 16 – 8 = 8
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Carbon Dioxide
This Lewis Structure uses the correct
number of electrons, but does not obey
the octet rule.
O=C=O
This Lewis Structure uses
the correct number of electrons
and obeys the octet rule.
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Hydrogen Cyanide
 HCN
1) Valence electrons = 10
2) Octet electrons = 18
3) Bonding electrons = 8
4) Number of Bonds = 4
5) H – C
N
6) Number of Nonbonding electrons = 2
H–C
N:
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Exceptions to the Octet Rule
(there are always exceptions)
BH3
 Each hydrogen accommodates 2 electrons, or one bond.
 The boron atom in BH3, on the other hand, has only 6
total electrons.
 Because boron is a smaller atom, it does not have
enough space to accommodate a full octet of 8 electrons.
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BF3
 BF3 has two potential Lewis structures shown below.
 Left structure (Structure I) is better because it minimizes
interaction between molecules.
 Structure II shows B and F with formals charges.
 F is a more electronegative atom (attracts more electrons)
and will have more 3 lone pairs, shown in Structure I.
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Formal Charge
 Determines which Lewis structure is correct if many
structures are possible
 Compares number of electrons around a bonded
atom to the number of electrons a lone atom
possesses
Formal charge = #valence electrons – (#nonbonding + ½
bonding)
 The best Lewis Structure is the one where the formal
charges are as close as possible to zero
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Figure I: Formal charge on top Fluorine is 0.
Formal charge on right Fluorine is 0.
Formal charge on left fluorine is 0.
Formal charge on Boron is 0.
Figure II: Formal charge on top Fluorine is 0.
Formal charge on right Fluorine is 1.
Formal charge on left fluorine is 0.
Formal charge on Boron is -1.
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NO – Nitric Oxide
 Exception: odd number of electrons
 Consider NO – 11 valence electrons
 Best course of action:
 Maximize number of bond
 Make sure neither atom in the 2nd period exceeds an octet
 One atom will have an odd electron count
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1st 2nd or 3rd period elements as central atoms
 Expanded octets: an exception
 Atoms in the 3rd period or higher can old more than 8
electrons
 They can hold 8, 10, or 12 electrons around the central
atom
 Examples: XeF4, SF6, PCl5
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Bonding and Shapes of Molecules
Number
of Bonds
Number of
Unshared Pairs
on Central Atom
0
3
0
4
0
3
1
2
2
-Be=C=
B
C
Shape
Examples
Linear
BeCl2
CO2
Trigonal planar
BF3
Tetrahedral
CH4, SiCl4, CCl4
Trigonal Pyramidal
NH3, PCl3
Bent
H2O,
:
2
Covalent
Structure
:
N
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O:
Linear
BeCl2
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Linear
Carbon dioxide
Linear
geometry
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Trigonal Planar
BF3
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Tetrahedral
Methane
Methane –The first member of the paraffin (alkane) hydrocarbons series.
a.k.a. (marsh gas, CH4).
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Tetrahedral
SiCl4 Silicon tetrachloride
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Tetrahedral
Carbon tetrachloride – CCl4
Carbon tetrachloride – “carbon tet” had been used as dry cleaning solvent
because it is extremely non-polar.
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Pyramidal
..
N
NH3
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H
Ammonia
H
H
Trigonal Pyramidal
Phosphorus trichloride
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Bent
d(-)
SO2
Water
O
H
H
d(+)
Polar molecule
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