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1

We are learning to:
1.
Represent compounds with Lewis structures.
2.
Apply the VSEPR theory to determine the molecular geometry of a compound.
We are looking for:
1. Draw Lewis structures for compounds and polyatomics based upon octet, formal charge, and resonance.
2. Determine the molecular geometry of a Lewis structure using VSEPR.
2

Name_____________________________________________________
Family Name Electron configuration Valance Shell
Electrons
Electron Dot
Notation
Alkali Metals
Alkaline Earth
Metals
Boron Family
Carbon Family
Nitrogen Family
Oxygen Family
Oxidation #
Halogens
Noble Gas
3

Introduction to Lewis Structures
1.
Draw electron dot for each element
2.
Determine if bond will be ionic or covalent a.
Metal + nonmetal  Ionic b.
Nonmetal + nonmetal  Covalent
3.
If Ionic: Draw arrows showing the electrons transferring from cation to anion a.
Metals will show a positive charge b.
Nonmetals will show a negative charge
4.
If Covalent: Circle the two electrons that make a shared bond. C, N, O, F must obey the octet rule and achieve 8 valence electrons by covalent bonding (H must always only have 2).
Examples : (also name each compound)
Na
2
O Ionic or Covalent?
CH
4
Ionic or Covalent?
AlCl
3
Ionic or Covalent?
NBr
3
Ionic or Covalent?
CO Ionic or Covalent?
4

Formal Charges & Lewis Structures
Used when one or more atoms have fewer or more bonds than usual; for example, sometime, C only forms 3 bonds instead of its usual 4. Also for elements that do NOT have to obey the octet rule (8 valence electrons). C, N, O, F must
obey the octet rule.
On a molecule, formal charges should add up to zero:
: C O:
On a polyatomic ion, formal charges must add up to the charge on the ion.
[ .. ..
:Cl O:
]
..
-
When more than one Lewis structure is possible, use the following rules:
1.
A Lewis structure with formal charges of zero is preferable to one with non-zero formal charges. Small formal charges are preferable to large formal charges.
2.
Lewis structures with negative formal charges on the more electronegative atom are more likely than Lewis structures with negative formal charges on the less electronegative atom.
3.
Lewis structures with unlike charges close together are more likely than Lewis structures with opposite charges widely separated.
4.
Lewis structures with like charges on adjacent atoms are very unlikely.
..
Ex) A) :
Ö C O:
B)
Ö C Ö
.. .. ..
The structure with all formal charges equal to zero is the better structure.
Which of the following is the best Lewis structure for formaldehyde, CH
2
O?
A) H C Ö
H
..
B) C O H
H
..
C) H C Ö H
5

Resonance Structures:
Even after using the rules for formal charges, sometimes more than one Lewis structure can be written. This often occurs when using unshared electron pairs to form multiple bonds. In this case both (or more) must be written. These are called resonance structures. In reality, the resulting structure is a hybrid of all the resonance structures. Use a double headed arrow between them.
*
Very important*: when writing resonance structures, all the atoms must be shown in the
same positions; only the positions of the electrons are different.
We can know the position of the atoms by experiment, but the position of the electrons cannot be determined; only the probability of finding an electron in a certain region can be known.
Ex) thiocyanate ion, SCN -
.. ..
S C N
But not
..
:S N C:
Nitrate, NO
3
-
..
:O N O:
..
..
..
:O:
..
..
..
..
Why?
O N O:
.. ..
:O:
..
:O N O
..
:O:
..
..
..
Acetate, C
2
H
3
O
2
-
H
..
H C C=O
..
H :O:
6

If single bonds are used and not all atoms have an octet (Except IA, IIA, and IIIA) – Try unshared pairs then double and triple bonds:
CO
2
O C O
CCl
2
O
C
2
Cl
2
Cl C O
Cl
Cl C C Cl
1.
O
2
4.
HClO
2
2. NCCN
3.
O
3
5.
N
2
H
4
7

1.
Using whatever information available, write the symbols for the elements in the correct arrangement.
2.
Calculate the total number of valence electrons. Remember to adjust amount for polyatomic ions!
3.
Place one pair of electrons (single line) between each pair of bonded atoms.
4.
Beginning at the outside of the formula, place the remaining electrons in pairs until there are eight electrons around each atom (two for hydrogen) or all electrons have been used. a.
If there are extra electrons, place them on the central atom. b.
Atoms in the 3 rd period and beyond can have more than eight electrons around
them and can form more than 4 bonds, but elements in the 2 nd period cannot.
5.
If not enough electrons are available to give all atoms (except H) an octet, move unshared pairs to form double or triple bonds. a.
However, Be, B, and other Group IIA elements may have fewer than eight
electrons.
6.
Examine your Lewis structure to see if resonance structures are needed. (Is more than one position possible for multiple bonds)
7.
Check the formal charges of the atoms. It should be at the lowest possible, even if the octet rule is not fulfilled (exception: C, N, O, F must have an octet and H two electrons).
Want to try to have a formal charge of zero for the central atom.
8.
Check your answer. Does the Lewis structure show: a.
the correct number of atoms? b.
the correct number of electrons? c.
the right number of electrons around each atom? d.
the minimum number of formal charges?
e.
The brackets around and the charge on it if it is a polyatomic ion?
8

Lewis Structures
Formula
( Give the name of the compound under the formula)
1.
Rb
3
N
2.
NF
3
3.
SCl
2
4.
InBr
3
5.
PCl
5
Ionic
Or
Covalent?
Number
Of
Valence
Electrons
6.
SF
6
7.
SeCl
4
Lewis
Structure
9

Formula
**( You do NOT have to name these)**
Ionic
Or
Covalent?
Number
Of
Valence
Electrons
8.
C
2
H
6
O
(Carbon-
Carbon-Oxygen bonding pattern)
9.
CCl
2
F
2
(Carbon in middle)
10.
NH
2
Cl
11.
SOCl
2
(Sulfur in the middle)
12.
CH
3
COCH
3
Lewis
Structure
10

Lewis Structures for
Polyatomic Species
Examples: NH
4
+ and PO
4
3-
When the charge is +
Subtract the charge value from the total number of valence electrons
When losing electrons – they usually come off the central atom!!
When the charge is –
Add the charge value to the total number of valence electrons
When gaining electrons - usually don’t add to the central atom!!
Example: NH
4
+ _____________________________
# of e s Total for each element Atom
N
How many x
1 5 = 5
H
1+
4 1 = 4
losing 1 = -1
Total valence electrons = 8
11

Example: PO
4
3 ____________________________________
Atom How many x # of e s Total # of electrons
P
O
1
4
5 =
6 =
5
24
3- gaining 3 = 3
Total valence electrons = 32
C
O
2-
O
1-
Example: BrO
3
-
______________________________________________________________
Atom How many x # of e s total # of electrons
Br 1 7 = 7
3 6 = gaining 1 =
18
1
Total valence electrons = 26
CO
3
2 _________________________________________
Atom How many x # of e s total # of electrons
1
3
4 =
6 = gaining 2 =
4
18
2
Total valence electrons = 24
12

Show math calculations and then draw the Lewis structure
1. H
3
O + ____________________________________
2. NO
3
_________________________________________
3. CN - _cyanide____
4. SO
3
2 ________________________________________
5. OH __________________________________________
13

14

V
S
E
P
R



Unshared pairs


CENTRAL ATOM
15

AX
5
AX
4
E
AX
3
E
AX
2
E
2
AX
2
E
AX
4
Type Overall Geometry
AX
2
AX
3
Linear
Trigonal Planar
Tetrahedral
Trigonal
Bipyramidal
Molecular Geometry Bond Angle ( o )
180
Examples
SiS
2
, CO
2
Bent
Tetrahedral
120
<120
109.5
Trigonal Pyramidal
Bent
Trigonal
Bipyramidal
See-Saw
<109.5
(107)
<<109.5
(104.5)
180
120
<180 (177)
<120 (104)
BCl
3
, H
2
CO
SO
2
CH
4
, (SO
4
) -2
NH
3
H
2
O
PCl
5
SF
4
16

AX
6
AX
5
E
AX
3
E
2
AX
2
E
3
AX
4
E
2
Octahedral
T-Shape
Linear
Octahedral
Square Pyramidal
Square Planar
90
<90
90
<90
(87.5)
180
ClF
3
XeF
2
, I
3
-1
SF
6
IF
5
(BrF
4
) -1 , XeF
4
17

Here are some rules about using VSEPR to describe molecular geometries:
1. Electron pair repulsions decrease in the following order: a. (nonbonding-nonbonding)> (nonbonding-bonding)> (bonding-bonding) b. triple bond > double bond > single bond
2. Nonbonding pairs occupy equatorial positions in the trigonal bipyramid.
3. Two nonbonding pairs in an overall octahedral structure occupy positions across from one another.
4. Bond angles decrease with increasing electronegativity of the non-central atoms, provided that there is at least one nonbonding pair in the species. For example, the F-N-F bond angle in NF
3
is smaller than the H-N-H bond angle in ammonia. Picture the electrons being closer to the more electronegative element fluorine than to the nitrogen. This allows the nonbonding pair to squeeze the bonding pairs closer together, resulting in a smaller bond angle. In ammonia, the bonding electrons are closer to the more electronegative element N so the electrons are already close to each other.
The bond angle doesn't get any smaller.
18

19

1) Linear a.
linear= 0 lone pair
Draw:
SiS
2
CO
2
2) Trigonal Planar a.
trigonal planar= O lone pair
Draw:
BCl
3
H
2
CO b.
bent= 1 lone pair
Draw:
SO
2
O
3
20

3) Tetrahedral a.
Tetrahedral= 0 lone pair
Draw:
CH
4 b.
trigonal pyramidal= 1 lone pair
Draw:
NH
3 c.
bent= 2 lone pairs
Draw:
H
2
O
(SO
<<109
4
) 2-
21

4) Trigonal bipyramidal a.
Trigonal bipyramidal= 0 lone pairs
Draw:
PCl
5 b.
See-saw= 1 lone pair
Draw:
SF
4 c.
T-shaped= 2 lone pairs
Draw:
ClF
3
22

d.
linear= 3 lone pairs
Draw:
XeF
2
5) Octahedral a.
octahedral= 0 lone pair
Draw:
SF
6 b.
square pyramidal= 1 lone pair
Draw:
IF
5
I
3
-
23

c.
square planar= 2 lone pairs
Draw:
(BrF
4
) XeF
4
24

Formula
(write the name under the formula)
1) SiF
4
Lewis
Structure
Lewis Structure/VSEPR Problems
Number of
Bonded Atoms
Number of
Lone Electron
Pairs
Molecular
Geometry
(name & sketch)
2) BBr
3
3) NF
3
4) H
2
O
5) AsCl
5
6) SF
6
25

Formula
(write the name under the formula)
1) KCl
Lewis Structures and VSEPR Practice
**Only do the VSEPR Shape on Covalent Structures**
Ionic or
Covalent
Lewis Structure
VSEPR Shape
Drawing
Shape Name
2) BH
3
3) CCl
4
4) CO
3
2-
5) SF
2
6) Na
2
O
7) AsBr
3
8) H
3
O +
26

Name:_________________________________
Lewis Structure/VSEPR Practice WS
Formula Compound Name Lewis
Structure
Resonance?
Yes or No
VSEPR Structure (include at least 1 angle)
1) Si I
4
2) A l
N
3) NO
3
-
4) BeC l
2
5) PF
3
VSEPR
Name
27

Formula
6) TeH
6
7) Rb
2
O
8) P I
5
9) PO
4
3-
10) XeF
2
Compound Name Lewis
Structure
Resonance?
Yes or No
VSEPR Structure (include at least 1 angle)
VSEPR
Name
28