Lewis Structures

Report
Bonding Review
Covalent Bonds (2 nonmetals)
…atoms share e– to get a full valence shell
C
1s2 2s2 2p2 4 valence e1s2 2s2 2p5
F
7 valence e-
*Both need 8 v.e – for a full outer shell (octet rule)!*
xx
o
o
C
o
o
x
x
Fx
xx
Draw the Lewis dot structure for the
following elements (write e- config first):
Si
1s2 2s2 2p6 3s2 3p2
4 valence e-
O
1s2 2s2 2p4
6 valence e-
P
1s2 2s2 2p6 3s2 3p3
5 valence e-
B
1s2 2s2 2p1
3 valence e-
Ar
1s2 2s2 2p6 3s2 3p6
8 valence e-
Br
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
7 valence e-
Notice any trends…?
1
2
3
4
5
6
7
H
8
He
Be
Na
Mg
K
Ca
Rb
Sr
Cs
Ba
TRANSITION METALS
Li
B
C
N
O
F
Ne
Al
Si
P
S
Cl
Ar
Se
Br
Kr
Te
I
Xe
The group # corresponds to the # of valence e–
Let’s bond two F atoms together…
Each F has 7 v. e– and each needs 1 more e–
F
F
F F
F2
Now let’s bond C and F atoms together…
carbon tetrafluoride (CF4)
F
F
C
F
F
F C F
F
F
Lewis Structures: 2D Structures
NH3
CH2O
H
N
H
H
CO2
H
H
C
H
H
CH4
O
C
O
(0)
(0)
(0)
O
S
SO2
O
Drawing Lewis Structures
1.
Sum the # of valence electrons from all atoms
Anions: add e– (CO32- : add 2 e– )
2.
Cation: subtract e– (NH4+: minus 1 e– )
Predict the arrangement of the atoms
• Usually the first element is in the center (often C, never H)
3.
4.
Make a single bond (2 e–) between each pair of atoms
Arrange remaining e– to satisfy octets (8 e– around each)
• Place electrons in pairs (lone pairs)
• Too few? Form multiple bonds between atoms:
double bond (4 e– ) and triple bond (6 e– )
5.
Check your structure!

All electrons have been used

All atoms have 8e-
Exceptions: Remember that H only needs 2e– !
Lewis Structure Practice
Draw a Lewis Structure for the following
compounds:
• CH4
• H2O
• NF3
• OF2
O
H
H
F
Br
H C N
O
N
F
F
F
• HCN
H
F
• HBr
O
• NO3-
N
O
• CO3
O
O
2-
C
O
O
Lewis Structure Trends
Here are some useful trends…
C group
O
• Forms a combo of 4 bonds and no LP (Lone Pairs)
• i.e. CO2
N group
• Forms a combo of 3 bonds and 1 LP
• i.e. NH3
O group
H
N
H
• Forms a combo of 2 bonds and 2 LP
• i.e. CH2O
F group (halogens)
• Forms 1 bond and 3 LP
• i.e. OF2
O
F
)
0
(
O
(0
H
F
Note that these are NOT always true!
)
0
(
C
)
Carbonite CO22Carbonate?
CO32-
Resonance Structures
Resonance structures differ only in the position of the electrons
Show resonance
O
C
O
O
Show movement of e-
C
O
O
O
C
O
O
O
• The actual structure is a hybrid (average) of the resonance
structures
• Technically NOT two single bonds and one double bond
• All 3 Oxygen atoms share the double bond
• 3 equal bonds (somewhere between a double and single)
Arrow formalism: curved arrows show electron movement
Predicting Molecular Shape: VSEPR
(Valence Shell Electron Pair Repulsion)
• Electrons repel each other
• The molecule adopts a 3-D shape to keep the
electrons (lone pairs and bonded e-) as far apart as
possible
• Different arrangements of bonds/lone pairs result in
different shapes
• Shapes depend on # of bonds/lone pairs (“things”)
and LP around the central atom
Selected Shapes and Geometries using VSEPR
“Things”
Carbon Dioxide: CO2
Lewis Structure
O
C
O
(0 )
(0 )
(0 )
O
C
O
 Two “things” (bonds or lone pairs)
 Linear geometry
 0 LP → Linear Shape
 180o Bond angle
Formaldehyde: CH2O
Lewis Structure
O
O
C
H
C
H
H
 Three “things”
 Trigonal planar geometry
 0 LP → Trigonal planar shape
 120° bond angles
H
Sulfur Dioxide: SO2
Lewis Structure
O
S
A
S
O
B
O
A
 Three “things”
 Trigonal planar geometry
 1 LP → Bent shape
 120° bond angles
O
A
Methane: CH4
Lewis Structure
H
H
C
H
H
 Four “things” (bonds/LP)
 Tetrahedral geometry
 0 LP → Tetrahedral shape
 109.5o bond angles
Ammonia: NH3
Lewis Structure
H
N
H
H
 Four “things” (bonds/LP)
 Tetrahedral geometry
 1 LP → Trigonal pyramid shape
 107o bond angles
Water: H2O
Lewis Structure
 4 “things” (bonds/LP)
O
H
H
 Tetrahedral Geometry
 2 LP → Bent Shape
 104.5o bond angle
Hydrogen Chloride: HCl
Lewis Structure
H Cl
 Four “things” (bonds/LP)
 Tetrahedral geometry
 3 LP → Linear Shape
H
Cl
Cl
 No Bond angle
A special note…
For any molecule having only two atoms…
 e.g. N2, CO, O2, Cl2, HBr, etc.
N N




O O
Cl Cl
H Br
Geometry = Linear
Shape = Linear
Bond Angle(s)? = None
It is much like geometry…
what is formed by connecting two points?
…a line.
You will need to commit these to memory!
“Things”
VSEPR Practice
(w/o aid of yellow sheet)
• CO2
• CH3COO-
G:
S:
Angle:
G:
S:
Angle:
• ClO2-
• PBr3
G:
S:
Angle:
G:
S:
Angle:
• NO2-
• AsO43-
G:
S:
Angle:
G:
S:
Angle:
Electronegativity and Bond Type
The electronegativity difference between two elements helps
predict what kind of bond they will form.
Electronegativity
Bond type
difference
≤ 0.4

0.5 – 1.8

> 1.8
Definition
Covalent
e- are evenly shared

Polar covalent e- are unevenly shared

Ionic
e- are exchanged (gained or lost)
Practice with Bond Types
Sample Bonds Electronegativity Difference Bond Type?
Ionic
NaCl
3.0 – 0.9 = 2.1
Covalent
Cl-Cl
3.0 – 3.0 = 0
Polar covalent
C-O
3.5 – 2.5 = 1.0
Covalent
C-H
2.5 – 2.1 = 0.4
H
2.1
Li
1.0
Na
0.9
K
0.8
Be
1.5
Mg
1.2
Ca
1.0
B
2.0
Al
1.5
C
2.5
Si
1.8
N
3.0
P
2.1
O
3.5
S
2.5
F
4.0
Cl
3.0
Br
2.8
I
2.5
Electronegativity
difference
≤ 0.4
0.5 – 1.8

> 1.8
Bond type
Covalent

Polar covalent

Ionic
Dipole Moments and Polarity
• Occurs in polar covalent bonds
• Uneven distribution of e• Atoms become partially charged
Partially
“+”
charged
end
δ+
H Cl
δ-
Arrow points
toward partially
“-” end
Polarity Examples
1. Check molecule for dipole moments (polar bonds)
2. When determining overall polarity, an imbalanced structure
will likely be polar (at least partially)
3. Even with polar bonds, a balanced structure is non-polar
overall
4. Any structure with lone pairs on the central atom is
automatically polar!
Try these with your neighbors…
•
•
•
•
•
HCN
CO2
CO32CH2O
SO2
Polar
Non-polar
Non-polar
Polar
Polar
•
•
•
•
•
CH4
CH3F
C3H8
CO
NH3
Non-polar
Polar
Non-polar
Polar
Polar

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