Honors Chemistry
Unit 6
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Lewis Dot Structures
VSPER Structures
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Name_____________________________________________________
Family Name
Electron configuration
Valance Shell
Electrons
Electron Dot
Notation
Oxidation #
Alkali Metals
Alkaline Earth
Metals
Boron Family
Carbon Family
Nitrogen Family
Oxygen Family
Halogens
Noble Gas
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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
Examples :
Na2O
Ionic or Covalent?
CH4 Ionic or Covalent?
AlCl3
Ionic or Covalent?
NBr3
Ionic or Covalent?
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4
5
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Lewis Structures
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:
O
C
O
Cl
C
O
Cl
Cl
C
C
Cl
1. SOCl2 (Sulfur in middle)
2. NCCN
3. CH3C(O)CH3
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Lewis Structures for
Polyatomic Species
Examples: NH4+
and
PO43-
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: NH4+
Atom
How many
x
# of e-s Total for each element
N
1
5
=
5
H
4
1
=
4
+
1
losing 1
=
-1
Total valence electrons = 8
8
Example: PO43-
Atom
How many
x
# of e-s Total # of electrons
P
1
5
=
5
O
4
6
=
24
-3
1
gaining 3
=
3
Total valence electrons = 32
Example: CO32-
Atom
How many
x
# of e-s total # of electrons
C
1
4
=
4
O
3
6
=
18
-2
1
gaining 2
=
2
Total valence electrons = 24
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Show math calculations and then draw the Lewis structure
1.
H3O+
2.
NO3-
3.
CN-
4.
SO32-
5.
OH-
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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.
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 ion.
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)
O CO
A)
B)
OOC
C) C O O
The structure with all formal charges equal to zero is the better structure.
Which of the following is the best Lewis structure for formaldehyde, CH2O?
A)
H C O
H
B) C O H
C) H C O H
H
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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 C N
S N C
Why?
Nitrate, NO3-
O N O
O N O
O
O
O N O
O
Acetate, C2H3O2-
H
H C C O
H O
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Summary of Writing Covalent Lewis Structures
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 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 3rd period and beyond can have more than eight electrons around
them but elements in the 2nd 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 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?
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VSEPR Theory
(Valence Shell Electron Pair Repulsion)
 Predicts the molecular shape of the resulting molecule
 Electrons on central atom arrange themselves as far apart as possible
 Unshared pairs on the central atom repel the most
 Shared pairs on the central atom repel the least
 To get the shape ONLY LOOK AT WHAT IS CONNECTED TO THE
CENTRAL ATOM!!!!!
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VSEPR Theory
Type
Overall Geometry
AX2
Molecular Geometry
Bond Angle (o)
Examples
Linear
180
BeF2, CO2
AX3
Trigonal Planar
120
BCl3, H2CO
AX2E
Trigonal Planar
<120
SO2
AX4
Tetrahedral
109.5
CH4, (SO4)-2
AX3E
Tetrahedral
<109.5
NH3
Bent
Trigonal Pyramidal
(107)
AX2E2
Tetrahedral
Bent
<<109.5
H2O
(104.5)
AX5
AX4E
Trigonal
Bipyramidal
Trigonal Bipyramidal
180
PCl5
120
See-Saw
177
SF4
104
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AX3E2
Trigonal Bipyramidal
T-Shape
87.5
ClF3
AX2E3
Trigonal Bipyramidal
Linear
180
XeF2, I3-1
AX6
Octahedral
90
SF6
AX5E
Octahedral
<90
IF5
Square Pyramidal
<90
AX4E2
Octahedral
Square Planar
90
(BrF4)-1, XeF4
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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 NF3 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.
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VSEPR Practice (molecular geometry)
1) Linear
a. linear= 0 lone pair
Draw:
BeF2
CO2
2) Trigonal Planar
a. trigonal planar= O lone pair
Draw:
BCl3
H2CO
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b. bent= 1 lone pair
Draw:
SO2
3) Tetrahedral
a. Tetrahedral= 0 lone pair
Draw:
CH4
(SO4)2-
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b. trigonal pyramidal= 1 lone pair
Draw:
NH3
c. bent= 2 lone pairs
Draw:
H2O
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4) Trigonal bipyramidal
a. Trigonal bipyramidal= 0 lone pairs
Draw:
PCl5
b. See-saw= 1 lone pair
Draw:
SF4
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c. T-shaped= 2 lone pairs
Draw:
ClF3
d. linear= 3 lone pairs
Draw:
XeF2
I3-
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5) Octahedral
a. octahedral= 0
lone pair
Draw:
SF6
b. square pyramidal= 1
lone pair
Draw:
IF5
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c. square planar= 2
lone pairs
Draw:
(BrF4)-
XeF4
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32
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XeF2
34
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