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
• HBr
• OF2
H
F
O
• HCN
H
N
F
H Br
F
O F
H C N
O
F
• NO3• CO3
O
N
O
O
2-
O
C
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 1 bond and 3 LP
• i.e. OF2
O
F
)
0
(
O
(0
H
• Forms a combo of 2 bonds and 2 LP
• i.e. CH2O
F group (halogens)
C
F
Note that these are NOT always true!
)
)
0
(
Carbonite CO22Carbonate?
CO32-
Resonance Structures
Resonance structures differ only in the position of the electrons
O
Show resonance
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
Lewis Structures and Formal Charge
For each atom in a molecule…
• Formal charge = # valence e– – # lone e– – # bonds
or = # valence e– – dots – lines
• Sum of formal charges = overall charge on molecule/ion
Formal Charges
Water
H2O
Hydronium
H3O+
H
O
H
H
O
H
H
H: 1 - 1 = 0
O: 6 - 4 - 2 = 0
Total charge = 0
H: 1 - 1 = 0
O: 6 - 2 - 3 = +1
Total charge = +1
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
H
 4 “things” (bonds/LP)
O
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
3.0 – 0.9 = 2.1
3.0 – 3.0 = 0
3.5 – 2.5 = 1.0
2.5 – 2.1 = 0.4
NaCl
Cl-Cl
C-O
C-H
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?
Ionic
Covalent
Polar covalent
Covalent
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
Intermolecular Forces (IMF’s)
• Intramolecular Forces = bonding within a
molecule
e.g. ionic, covalent, polar covalent bonds
• Intermolecular Forces = interactions between
two molecules
…Intercity v. Intracity v. Innercity
• Intermolecular Forces are ALL
weaker than Intramolecular
bonds
IMF’s: Ion-Ion Force
+
+
-
+
Opposite Charges Attract
Similar Charges Repel
Cl–
Na+
Na+
Na+
Cl–
Attractive and
repulsive forces
between two separate
ions.
Cl–
IMF’s: Ion-Dipole Force
The interaction
between an ion
and another
molecule that has
a dipole moment.
(polar covalent)
δ+
Cl
Na+
H
δCl
δ-
+
H
δ+
Lewis Structure
H
δ-
H
δ+
O
O
Na+
H
H
δ+
IMF’s: Dipole-Dipole Force
The interaction
between two
separate molecules,
each having a
dipole moment.
(polar covalent)
δ+
δH
Cl
H
Cl
δ-
δ+
H
HCl = Stomach Acid
Cl
H
Cl
IMF’s: Hydrogen Bonding
H
O
H
H
O
A specific type of
dipole-dipole
interaction
between an H
bond donor and
an H bond
acceptor.
H
O
H
O
H
H
H
H
H bond donor: an H bonded to N, O, or F
H bond acceptor: any lone pair of e–
H
O
O
H
H
IMF’s: London Dispersion Forces
Involves an
instantaneous
dipole. This
dipole will
induce dipoles
in other
molecules.
All molecules
will exhibit LDF
↑ mass,↑ LDF
H
NO!
Probable? Yes
H
δ+
δ-
YES!
Possible? Yes
Why instantaneous?
δ-
δ+
This dipole will only remain
for an instant! The electrons
will quickly move to another
part of the molecule!
H
H
H
H
H
H
H
H
H
H
H
H
Instantaneous = WEAKEST!
H
H
IMF Review
STRONGEST
Ion-Ion
Involves an ion (+ and – charged)
Ion-Dipole
Hydrogen Bonding
Involves a dipole
(polar molecule)
Dipole-Dipole
London Dispersion Forces (LDFs)
 a.k.a. van der Waals Forces
Weakest
Involves a non-polar molecule
*Remember: These are all weaker than actual bonds
(ionic, covalent, etc.). These are just attractions.
IMF Practice
Formaldehyde: CH2O
Lewis Structure
O
O
H
C
C
H
H
H
Trigonal planar geometry
120° bond angles
Polar C=O bond = Net dipole moment
IMF= Dipole-Dipole
Methane: CH4
Lewis Structure
H
H
C
H
H
Tetrahedral geometry
109o bond angles
Covalent bonds = No net dipole
IMF = London dispersion forces
Ammonia: NH3
Lewis Structure
H
N
H
H
Trigonal pyramid
107o bond angles
Polar Bonds, Lone pairs = Dipole
1 H bond acceptor (LP), 3 H bond donors (N-H)
IMF = Hydrogen Bonding
Carbon Dioxide: CO2
Lewis Structure
O
C
O
(0)
(0)
(0)
O
C
O
Linear geometry
180o Bond angle
C=O bond is polar, but…
Dipoles cancel = No net dipole
IMF = London Dispersion Forces!
Water: H2O
Lewis Structure
H
O
H
Bent
Polar bonds, lone pairs = Net dipole
2 H bond acceptor (LP), 2 H bond donors (O-H)
IMF = Hydrogen bonding
Orbital Review
• Each atom contains electrons in atomic orbitals
Each orbital can hold a max of 2 electrons
2p
E
Valence e-
2s
1s
 Core e-
Orbitals are in energy levels
Inner energy levels = core e- (kernel e-)
Outermost energy level = valence e-
Valence e- are used for bonding
Ionic bonding: transfer of those eCovalent bonding: sharing of those e-
Orbitals and Bonding
• Covalent bonds are formed by sharing two e• Sharing occurs in overlapping atomic orbitals
• Example: H2 (H-H)
E
1s
1s : 1s
Use 1s orbitals for bonding
• Example: H2O
From VSEPR: bent, 104.5°
angle between H atoms
Use two 2p orbitals for bonding?
1s
2p
E
2s
90°
1s
2p
2p
How can we explain this
bonding and match VSEPR?
Bonding
• Single bonds (one sigma bond) σ
Overlap of s orbitals on bond axis
• Double bonds (one sigma + one pi bond)
Sharing of electrons between p orbitals
perpendicular to the bonding atoms
Termed “pi” or π bonds
Bond Axis of σ bond
2p
2p
O=O or O2
One π bond
Hybrid Orbitals
+
=
H2
2H: 1s
+
=
C: 2s, 2px, 2py, 2pz
TRIGONAL
PLANAR
=
sp2
C: three sp2
C: 2s, 2px, 2py
+
sp
C: two sp
C: 2s, 2p
+
LINEAR
=
TETRAHEDRAL
C: four sp3
sp3
Hybrid Orbitals: Tetrahedral
Mix starting atomic orbitals together to form hybrid orbitals
Lone pairs Bonds
with H
H
O
H
1 s and 3 p orbitals on O mix to form 4 identical sp3 hybrid
orbitals (tetrahedron)
• Two sp3 orbitals contain O lone pairs
• Two form  bonds with H 1s orbitals
Hybrid Orbitals: Trigonal Planar
 bond
 bonds
O
H
C
H
• 1 s and 2 p orbitals on C mix to form 3 identical sp2 hybrid
orbitals (trigonal planar).
• One p orbital remains unhybridized.
• sp2 orbitals form  bonds with O and H atoms
• Leftover p orbital forms  bond with O
Hybrid Orbitals: Linear Bonding
 bonds
 bonds
O
C
O
(0)
(0)
(0)
• 1 s and 1 p orbital on C mix to form 2 identical sp
hybrid orbitals (linear)
• Two p orbitals remain unhybridized.
• Each sp orbital forms a  bond with O
• Each leftover p orbital forms a  bond with O
Determining Hybridization
• To determine an atom’s hybridization, count “things”
around the atom
“Things” = bonded atoms and lone pairs
• Two things: sp hybrid orbitals
• Three things: sp2 hybrid orbitals
• Four things: sp3 hybrid orbitals
Methane
Acetonitrile
H
H
C
H
H
4 “things”
sp3 hybridization
sp
H
H
C
C
H
sp3
N