Covalent Bonding
VSEPR theory, Lewis structure and
resonance
Predicting the arrangement of
bonding and nonbonding
electrons, and molecules
Why?
 Knowing
the shape of molecules helps
predict their behavior, solubility and how
biological molecules interact and much
more.
 We will look at bond and molecule
polarity and 3-D structure.
Type of Bond and Bond
Polarity
Metallic: delocalized valence e-, (sea of
electrons), piece of metal, good electrical
conductors, very strong bond, very high mp/bp.
Ex. Gold ring, copper wire, steel basket, etc.
b)
Ionic: usually metals with non-metals, absolute
Electronegativity difference greater than 1.7,
brittle, high mp/bp, strong bond, hydrophillic.
Ex. Ca Cl2 |1.0-3.16| = 2.16
Hint: look at each bond individually (don’t multiply
by 2)
So e- have been transferred from Ca to Cl, will form 3
separate ions in solution, Ca2+ Cl1- Cl1a)
c) Pure covalent bond: generally non-metals,
electrons shared equally between the 2 atoms,
weak bonds, low mp/bp, hydrophobic,
lipophilic
Electronegativity difference < 0.4
Ex. Methane (CH4) |2.55-2.20| = 0.35
d) Polar covalent bond: generally non-metals,
electrons shared unequally between the 2
atoms, weak bonds, low mp/bp, hydrophilic,
lipophobic
Electronegativity difference between 0.4 and
1.7
Ex. Water H2O
|3.44-2.20| = 1.27
Greek letter delta (lowercase) is used to indicate a partial charge
Determine the type of bond
A.
B.
C.
D.
E.
F.
Silicon tetraiodide
Ammonia, NH3
Silver spoon
Sodium fluoride
Laughing gas, N2O
Alloy of gold and nickel
Bellringer
1.
2.
3.
4.
5.
How many valence electrons does N
have?
How many does C have?
And hydrogen?
How many valence electrons does H
need total to have a complete energy
level?
How many does carbon need?
Writing a Lewis Structure
1.
2.
3.
4.
Sum the valence electrons for all the atoms. Do
not worry about which electrons came from
which atom.
Use a pair of electrons to form a covalent bond
between each pair of bonded atoms. Use ―
Arrange electron pairs around the attached
atoms to satisfy A) hydrogen having two electrons
and B) the octet rule for nonmetal elements of
Period 2 (except Boron)
Remaining electrons (pairs or single) are placed
around the central atom. The Octet rule is more
like a suggestion…it is not always followed. C, N,
O, F have to have octets, others would like to.
VSEPR



For methane the Lewis
structure appears to have
hydrogen atoms attached to
the carbon resulting in 900
angles between the hydrogen
atoms.
Due to the repulsion of the
electrons in the bonds, the
molecule attains a 3D
structure where the angles
between the bonds are 109.50
The repulsion of the electron
pairs gives the molecule a
shape or geometry.
Bonding and nonbonding
electrons
 Bonding
electron
pairs (n) are in a
covalent bond
around a central
atom.
 Nonbonding
electron pair (m)
are in lone pairs
around the central
atom.

Note:we are only concerned with the electrons around the central atom.
Lewis Structure Example
Write a Lewis Structure for the cyanide
anion (CN1-).
Step 1: Carbon has 4 valence electrons
Nitrogen has 5 valence electrons
1- adds 1 valence electron
Total valence electrons: 4+5+1 = 10
Lewis Structure Example
Step 2: Bond the carbon and nitrogen atoms
with a pair of electrons:
C-N
10 electrons – 2 electrons in the bond = 8 electrons remaining.
Lewis Structure Example
Step 3: Place the remaining electrons around the
atoms.
C-N
Since carbon and nitrogen are in Period 2, they each must have an
octet. So, rearrange the electrons…
Step 4: make any double or triple bonds so that all have an octet
(remember, H only needs 2 electrons to have a complete shell).
Since this is a polyatomic ion, place brackets and the charge
[C ≡ N]1-
Lewis Structure Example
C-N
This structure shows that the carbon
and nitrogen are held by a triple
covalent bond with each atom
retaining a lone pair of electrons.
Final structure:
C
N
-
Examples
 Water
 Bromine
gas
 Methane
 Carbon
dioxide
Practice on whiteboards
G.
Carbon tetrachloride
Nitrogen (check to see if diatomic)
Hydrochloric acid
Ammonium ion
Fluorine (diatomic?)
Silicon dioxide
CH3 C N (hydrogens on 1st carbon, carbon bonded
H.
Nitrate ion
A.
B.
C.
D.
E.
F.
to carbon and N at the end)
Practice Lewis structures
Write the formula, total valence electrons, AND draw
the correct Lewis structure.
1.
Ammonia
2.
Carbon bromine trihydride
3.
Sulfur dioxide
4.
CH3 OH (methanol)
5.
Sulfate ion (SO4)26.
Hydrogen peroxide
7.
Nitrate ion
8.
PBr3
9.
BeI2
10. Carbon dioxide
Practice Lewis structures
Draw the best Lewis Dot Structure for each of the
following species. (some are exceptions to the octet rule)
a) S F2
b)
B Cl3 (no double bonds)
c) C H2 C H2 (carbons bonded to each other)
d)
P Br5
e) S I2
f)
B H2–
g)
N I3
h)
Cl F4+
i)
sulfate ion
j) chlorate ion
Bellringer
1.
2.
3.
How many total valence electrons
does a molecule of methanol have?
CH3 OH
Draw the Lewis structure for it. (hint
carbon is bonded to oxygen also)
Draw the Lewis structure for As Cl5
Videos Lewis structures,
resonance, etc
 Crash
course Lewis structures + Models
https://www.youtube.com/watch?v=a8LF
7JEb0IA
 Lewis
structures and resonance 5 steps
https://www.youtube.com/watch?v=1Zlnzy
Hahvo
Resonance
 Resonance
occurs when more than one
Lewis structure may be drawn for a given
compound.
Resonance
 Resonance
appears to be like the
electrons are “flipping” or “moving” to
different positions within the molecule
while the atoms remain in a fixed location.
Combined Resonance structure
The dashed line represents the spread of the
electron energy around the molecule. Since 2
electron energies are spread around 3 areas each
dashed line has a value of 2/3 electron energy.
Resonance
 Instead
of the electrons moving,
resonance shows that the electron energy
becomes spread throughout the
molecule.
On your own
a)
Resonance structures for the thiocyanate
ion and label formal charges
b)
Resonance structures for carbon disulfide
c) Resonance structures for NCO1-
Evaluating Lewis structures for
Molecular Compounds with the
Formal charge
Molecular compounds form in such a way that the
formal charge on each atom is minimize; the closer
to zero the better.
Formal Charge: a measure of the distribution of
electron energy around a given nucleus.
The best structure will have least formal charges or
better spread out. If 2 or more structures are equal
then are probably resonance forms.
We will use formal
• Higher Eneg atom = zero or neg preferred charge when
evaluating Lewis
• Lower Eneg atom = zero or pos preferred

Structures

Calculation of Formal Charge:
FC  # valence e   (# nonbonded e   1 # bonded e  )
2
Formal Charge



Formal charge is a way to evaluate the distribution
of electron energy throughout a molecule.
The most stable Lewis structure for a particular
molecule is one that gives the lowest charge on
each atom in the molecule.
Lets try the nitrate ion….
The + and – seen in this
structure is showing the
formal charge on each
atom. They should be
circled.
Since all three structures have
the same formal charge, then
all 3 are equally plausible.
Formal Charge
 Using
formal charge
to evaluate a
structure:
Carbon dioxide CO2
There are 3 possible
Lewis structures for
this, draw all 3…..
Which structure is right….all three follow the rules for drawing Lewis
structures….all are resonance structures….but are they equal
resonance structures? Lets find out.
Best Lewis Structure
 Since
the formal charges are zero on all atoms
(which is as good as you can get), the correct
structure of CO2 is
Practice Formal charges
to turn in individually

Determine the best Lewis structure for each of
the following (give all possible resonance
structures and formal charges on each atom):
1.
2.
3.
4.
5.
6.
Ge O32Br O41Se Cl4
Sb F5
N Si O1H C N or C N H or C H N
Practice Lewis
 Determine
the best Lewis structure for
each of the following (give all possible
resonance structures)
1.
2.
3.
4.
5.
6.
Ge O32Br O41Se Cl4
Sb F5
N Si O1HCN
Homework Formal charge
For each: name it, draw the Lewis structure (all
possible resonance), tell how many ve-, label formal
charges and if several resonance structures, tell
which is favored.
1.
CO322.
PO32-
3.
ClO41-
4.
H C N or N C H or N H C
Bellringer
1.
2.
Draw all resonance structures for
SO2
Label formal charges and
determine which structure is the
most favored.
Quiz Friday
Over Lewis,
formal
charges,
resonance.
Not grade
cam
VSEPR
Valence
Shell
Electron
Pair
Repulsion
Theory
VSEPR videos
 Dr
McCord VSERP
https://www.youtube.com/watch?v=keH
S-CASZfc
VSEPR theory

Electron pairs (or groups of pairs) try to avoid one another
because of repulsions between like-charged particles

Electron pairs (bonded or lone) tend to spread out as far
as possible, lone pairs need most space.

Regions where electrons are likely to be found will be
called electron domains.

A double or triple bond comprises a single electron
domain.
Lewis and Model (give
shape name See page
200 blue book)
 Ammonia!
 Methane!
 Water!
 Acetylene
HCCH
 BF3 (equal bonds)
Molecular shapes and
electron domains
 Pg
200 shape names
 Ex Draw Lewis structure, determine
polarity of each bond, label polar bonds
(with δ+ /δ- or →), give electron
arrangement and molecular geometry.
 Ammonia
 Carbon tetrabromide
 Water
Whiteboards
For each: Draw Lewis structure, determine
polarity of each bond, label polar bonds (with
δ+ /δ- or →), give electron arrangement and
molecular geometry. Then build the model
and have it checked.
1.
2.
3.
4.
5.
Carbon dioxide
Phosphorous triiodide
Silicon tetrabromide
HCN
Sulfur dioxide
Bellringer
1.
2.
3.
4.
5.
Draw the Lewis structure for the
hydronium ion.
Does it have resonance structures? If so
what are they?
Label formal charges.
What molecular shape does it have?
What is the name of the electron
arrangement?
Molecular geometry

Molecular 3-D shape names
See page 265 for some basic shape names
Electron arrangement looks at how many regions of
electron density are around a central atom
Molecular geometry looks at only the bonded atoms
around the central atom
Electron pairs (bonded or lone) tend to spread out as far as
possible, lone pairs need most space.
Ex ammonia

Carbon tetrabromide

Water





Getting the shape

1.
2.
3.
4.



Construct a Lewis structure from the formula.
Sum the valence electrons for all the atoms. Do not worry
about which electrons came from which atom.
Use a pair of electrons to form a covalent bond between
each pair of bonded atoms.
Arrange electron pairs around the attached atoms to satisfy
A) hydrogen having two electrons and B) the octet rule for
most elements.
Remaining electrons (pairs or single) are placed around the
central atom. The Octet rule is more like a suggestion…it is not
always followed.
Determine the SN of the central atom to determine the electron
domain geometry.
Determine the AXnEm for the structure to determine the
attachment domain geometry.
Give the shape of the molecule.
Drawing in 3-D