Low melting & boiling points
Low heats of fusion & vaporization
High vapor pressure
May be soft, as wax
May be crystalline, as sugar
(weak lattice, based on dipole-dipole or
H bonding)
Molecules are neutral. NEVER conduct.
Properties of Molecular
Substances
Breaking a bond is
Breaking a
endothermic.
bond is
Making
a bond is
endothermic.
exothermic.
Making a bond
is exothermic.
Describe changes in chemical
potential energy that accompany
bond formation or bond breaking.
As PE , stability .
As PE , stability .
Describe the relationship
between stability & potential
energy
H + H  H2 + energy
Exothermic
- energy term is on product side
- bond formation releases energy
Does the above equation
represent an endothermic or
exothermic process?
How do you know?
H2 + energy  H + H
Endothermic
- energy term is on reactant side
- bond breaking absorbs energy
Does the above equation
represent an endothermic or
exothermic process?
How do you know?
To achieve the electron
configuration of the nearest
noble gas!
Why do atoms form bonds?
Covalent Bonds
Electrons are shared
All Nonmetals in
formula
How do you identify a
covalent formula?
The simultaneous attraction of
electrons and two different
nuclei.
Covalent bonds result from …
•Molecular, which give exact
C6H6
composition of molecule
•Empirical, which give lowest whole
CH
number ratio of atoms in molecule
•Sometimes they are the same. H O
2
•Otherwise, molecular is a whole
number multiple of empirical
What kinds of formulas do
molecular substances have?
•Shows which elements & how
many atoms of each.
•Shows connectivity or how
atoms are linked
•Shows type of bond – single,
double or triple
H-O-H
O=C=O
O-N=O
Structural Formula
•Shows type & number of atoms.
•Shows connectivity
•Shows type of bonds
•Shows all nonbonding valence
electrons, in addition to the bonding
valence electrons
Bonding electrons
are in between
two atoms!
Lewis Diagram
One electron pair
or
Two electrons
Shared between atoms
Represented by 2 dots or 1 dash
between atoms
Single Bond
Two electron pairs
or
Four electrons
Shared between atoms
Represented by 4 dots or 2 dashes
between atoms
Double Bond
Three electron pairs
or
Six electrons
Shared between atoms
Represented by 6 dots or 3 dashes
between atoms
Triple Bond
Energy change that occurs when a
bond is formed between two atoms.
Symbol = D0
Bond Energy
As the number of electrons shared
between 2 atoms increases, the
attractive interactions increase, & the
bond energy increases.
Triple > Double > Single
Bond Energy
Distance between two bonded nuclei.
The more shared electrons between 2
nuclei, the greater the attractive
interactions, the shorter the bond
length.
Bond Length
Overlap of two orbitals occurs on the
line directly connecting the two nuclei.
All single bonds are sigma bonds.
 Bond
Overlap of two orbitals occurs above
and below the line that directly
connects the two nuclei.
Double bonds consist of one  & one  bond.
Triple bonds consist of one & two  bonds.
 Bond
1)
2)
3)
Total up the valence electrons from all the atoms in
the molecule.
Draw the skeleton, including a single bond between
every atom.
Compare numbers:
1)
2)
4)
5)
# of electrons needed for each atom in the skeleton to have an
octet/duet
# of electrons available after drawing skeleton
Distribute electrons
Verify by performing two validity checks
Procedure for writing a Lewis
structure for a molecular
substance
1. Total # of valence electrons = 2
2. Skeleton: H – H
3. Compare: 0 electrons needed & 0
electrons available
4. No electrons to distribute
5. Verify.
Lewis Diagram of H2
1.
2.
3.
4.
5.
# of valence electrons = 14
Skeleton: Cl – Cl
Compare: Need 12, Have 12
Distribute
Verify
Lewis Diagram of Cl2
Bonding electrons are located
BETWEEN two atoms.
Nonbonding electrons are
located on one atom only.
Bonding vs. Nonbonding
electron pairs
1. # of valence electrons = 12
2. Skeleton: O - O
3. A) Compare: Need 12, Have
10, deficient by 2, add 1 bond
B) O  O. Compare: Need 8,
have 8.
4. Distribute
5. Verify
Lewis Diagram of O2
1. # of valence electrons = 10
2. Skeleton: N – N
3. A) Compare: Need 12, Have 8,
deficient by 4, add 2 bonds
B) N  N. Compare: Need 4, have
4.
4. Distribute
5. Verify
Lewis Diagram of N2
1.
2.
3.
4.
5.
# of valence electrons = 8
Skeleton: H – O – H
Compare: Need 4, Have 4
Distribute
Verify
Lewis Diagram of H2O
-
1. # of valence electrons = 8
H
2. Skeleton:
H–N–H
3. Compare: Need 2, Have 2
4. Distribute
5. Verify
Lewis Diagram of NH3
- -
1. # of valence electrons = 8
H
2. Skeleton: H – C – H
H
3. Compare: Need 0, Have 0
4. Distribute (nothing)
5. Verify
Lewis Diagram of CH4
- -
1. # of valence electrons = 32
Cl
2. Skeleton: Cl – C – Cl
Cl
3. Compare: Need 24, Have 24
4. Distribute
5. Verify
Lewis Diagram of CCl4
-
1. # of valence electrons = 12
H H
-
2. Skeleton: H – C – C - H
3. A) Compare: Need 4, Have 2. Add 1
bond.
H H
B) Compare:
H – C  C - H Need 0, Have 0
4. Distribute (nothing)
5. Verify
Lewis Diagram of C2H4
1. # of valence electrons = 10
2. Skeleton: H – C – C - H
3. A) Compare: Need 8, Have 4. Add 2
bonds.
H–CC-H
B) Compare: Need 0, Have 0
4. Distribute (nothing)
5. Verify
Lewis Diagram of C2H2
1. # of valence electrons = 16
2. Skeleton: O – C - O
3. A) Compare: Need 16, Have 12,
deficient by 4, add 2 bonds
B) O  C  O. Compare: Need 8,
have 8.
4. Distribute
5. Verify
Lewis Diagram of CO2
Sometimes, more than one valid
Lewis structure can be written
for a molecule. For CO2:
<-->
<-->
What are Resonance
Structures?
The atoms are in the same
location.
The electrons are distributed
differently.
Resonance Structures
1. Molecules with an odd # of electrons can
never satisfy the octet rule for all their
atoms. (NO, NO2, ClO2)
2. Some molecules have an atom with less
than an octet. (BF3, BeH2)
3. Some molecules have an atom with more
than an octet. (PCl5, SF6)
What are the three general ways
the octet rule breaks down?
Molecules with odd # of electrons
1. Consider NO
5 + 6 = 11 valence
electrons
2. Skeleton: N – O
3. A) Compare: the N needs 6 & the O
needs 6 for a total of 12. Have only 9
available. Add one bond.
B) N = O now each atom needs 4 for a
total of 8. Have
only
7
available.
.
..
4. Distribute: :N = O:
The N atom has
only
7
valence
e
5. Verify
Molecules with odd # of electrons
1. Consider NO2. 5 + 2(6) = 17 electrons
2. Skeleton: O – N – O
3. A) Compare: each O needs 6 & N
needs 4. Need 16 total. Have 13
available. Add 1 bond.
B) O = N – O Compare: Need 4 + 2 +
6 or 12 electrons.
.. .Have
.. 11 available.
4. Distribute: :O = N – O:
..
The N atom
5. Verify
has only 7 e-
Molecules with odd # of electrons
1. Consider ClO2 7 + 2(6) = 19 valence e2. Skeleton is O – Cl – O
3. Compare: Need 6 + 4 + 6 = 16 eHave: 19 – 4 = 15. Deficient by 1 e-.
Can’t fix.
..
. ..
The Cl atom
4. Distribute :O
–
Cl
–
O:
.. .. ..
has only 7
5. Verify
valence e- !
Molecules with atoms that have
less than an octet
Occurs in molecules with Be and B.
Be likes to have 4 valence electrons in
molecules.
B likes to have 6 valence electrons in
molecules.
Molecules with atoms that contain
more than an octet
• Only the central atom can have more than an
octet.
• And only if it belongs in rows 3-7 of the PT.
• Consider PF5. 5 + 5(7) = 40 valence e-.
• Since P is in row 3 it has empty d orbitals
available which can be used for bonding.
• Skeleton:
The P has >
than an octet.
F
F P
F
F
F
Distribute the
remaining 30 e- by
placing 6 e- on each
of the 5 F atoms.
A group of covalently bonded
atoms that has gained or lost
electrons and hence acquired a
charge
Polyatomic Ions
Lewis Diagrams of Polyatomic
Ions
• When calculating the total # of valence
electrons, you must adjust for the
charge of the ion.
– Total up the electrons contributed from
each atom.
– Add 1 electron for each negative charge.
Or
– Subtract 1 electron for each positive
charge.
2. Skeleton:
- -
1. # of valence electrons = 9 – 1 = 8
+1
H
H–N–H
H
3. Compare: Need 0, Have 0
4. Distribute (nothing)
5. Verify
Lewis Diagram of NH4+1
# of valence electrons = 6 + 1 + 1 = 8
Skeleton: [O – H]-1
Compare: Need 6, Have 6
Distribute: [:O - H]-1
Verify
: :
1.
2.
3.
4.
5.
Lewis Diagram of OH-1
Determining Molecular Shape
from the Lewis Structure
1. Count up the number of electron
domains on the central atom.
–
–
Single, double, & triple bonds each count
as ONE domain.
Lone electron pairs (or even a lone
singleton) counts as ONE domain.
2. Count up the number of atoms that are
bonded to the central atom.
3. Compare these two numbers to get the
shape.
Shapes on the Regents Exam
# of electron
domains on
central atom
# of atoms
bonded TO the
central atom
Shape
Example
2
2
Linear
CO2
4
4
Tetrahedral,
109
CH4
4
3
Trigonal Pyramid,
107
NH3
4
2
Bent, 105
H2O
Additional Shapes: Know *reds
# of electron
domains on
central atom
# of atoms
bonded TO the
central atom
Shape
Example
3
3
*Trigonal Planar,
bond angle = 120
BF3, BH3, SO3
5
5
*Trigonal
Bipyramid
PF5
5
4
See-Saw
5
3
T-Shape
5
2
Linear
6
6
*Octahedral
6
5
Square Pyramid
6
4
Square Planar
SF6
Tetrahedral Molecule
Trigonal Pyramidal Molecule
Bent Molecule
Shape of CH4
1. Inspect central atom in Lewis diagram:
2. Central atom has
- 4 electron domains
- 4 bonded atoms
3. Shape is tetrahedral
with 109 bond angles
Shape of NH3
1. Inspect central atom in Lewis diagram:
2. Central atom has
- 4 electron domains
- 3 bonded atoms
3. Shape is trigonal pyramid with 107
bond angle
Why isn’t the bond angle 109?
Because the lone pair on the N atom spreads out
& squeezes the bonding pairs together.
Shape of H2O
1. Inspect central atom in Lewis Diagram:
2. Central atom has
- 4 electron domains
- 2 atoms bonded to it
3. Shape is Bent with a 105 angle
Why isn’t the bond angle 109?
Because the two lone pairs on the O spread
out more than the 2 bonding pairs and
squeeze the bonding pairs together.
Shape of BF3
Recall: B is an exception to the octet rule!
1. Inspect central atom in Lewis Diagram:
2. Central atom has
or
- 3 electron domains
- 3 atoms bonded to it
3. Shape is Trigonal Planar with a 120
angle
Shape of BeF2
Recall: Be is an exception to the octet rule!
1. Inspect central atom in Lewis Diagram:
2. Central atom has
- 2 electron domains
- 2 atoms bonded to it
3. Shape is Linear with a 180 angle
Shape of SO2
1. Inspect central atom in Lewis Diagram:
2. Central atom has
- 3 electron domains
- 2 atoms bonded to it
3. Shape is Bent with a 120 angle
Shape of PF5
1. Inspect central atom in Lewis Diagram:
2. Central atom has
F
F
- 5 electron domains
F P
F
- 5 atoms bonded to it
F
3. Shape is Trigonal Bipyramid
The P is allowed to have an
“expanded octet” because it
has empty 3d orbitals that can
hold valence electrons.
PF5
Two kinds of F atoms in PF5 :
Axial – set of two
Equatorial – set of three
The axial F atoms have 3
nearest neighbors at 90.
The equatorial F atoms have
2 nearest neighbors at 90 &
2 nearest neighbors at 120.
The equatorial F
atoms are less
crowded than the
axial F atoms!
Shape of SF6
• Inspect central atom in Lewis diagram.
• Central S atom has
– 6 electron domains
– 6 atoms bonded to it
• Shape is octahedral!
• The S is allowed to have an “expanded
octet” because it has empty 3d orbitals
that can hold valence electrons.
Linear
2 points make a line!
Shape of all diatomics?
Trigonal Pyramid
Shape of NH3, NF3, PH3,
etc?
Bent, with bond angle of
105
Shape of H2O, H2S, H2Se,
etc.?
Tetrahedral
Shape of CH4, CCl4, etc.?
Bond has poles – the ends are
different!
Bond has a permanent partial
separation of charge
Polar Bond
Electronegativity
Difference is 0.5 to 1.7
Polar Bond
No poles.
Ends are the same.
Symmetric electron cloud.
Nonpolar Bond
Electronegativity difference = 0
to 0.5
NO separation of charge in bond.
Nonpolar Bond
Electronegativity
difference  1.7
Ionic Bond
Full Separation of Charge
At least +1 and -1.
Ionic Bond
Formula has a metal and
a nonmetal
Ionic
Molecular Polarity
• Depends on how the atoms are arranged
in the molecule.
• Some molecules which contain polar
bonds are nonpolar overall. Common.
• Some molecules which contain nonpolar
bonds are polar. Less common.
Polar.
Water is bent.
The O end is a bit negative
& the H end is a bit
positive.
WATER, H2O
Results from attractions
between nucleus on 1 atom &
electrons on another atom.
Bonding
releases energy.
Making a bond …
Absorbs energy.
Breaking a bond …
(almost all)
Molecular compounds.
Covalent Compounds are
Triple Bond
N2
Double Bond
O2
Single Bond
F2, Cl2, Br2, I2
released.
When bonds are made,
energy is …
absorbed.
When bonds are broken,
energy is …
increases.
As the energy of a system
, the stability generally
…
System releases
energy. Its energy
level goes down.
Exothermic
1)
2)
3)
4)
5)
Soft
Low melting point & low boiling point
Does not conduct electricity in any phase
Does not dissolve in water
React slowly
Properties of Molecular
Substances
Electrons are shared
equally between the
two atoms
Nonpolar Bonds
Molecule must contain polar
bonds and they must be
arranged asymmetrically.
Molecular Polarity
Depends on the shape
(Bent & Pyramidal are polar
Linear & tetrahedral, polarity depends
on composition.)
Molecular Polarity
1) Noble gas atoms (kickballs)
2) 7 Diatomic Elements (footballs)
3) CXHY or pure hydrocarbons.
4) Larger molecules that have high
symmetry
Nonpolar Molecules
•
If it’s not one of the 4 easy
categories of nonpolar molecules, it
is a polar molecule!
Polar Molecules
Weak Intermolecular Forces
(Dispersion or Van der Waals)
Low boiling points & melting points
Tend to be gases
Nonpolar Molecules
Intermolecular Forces are Dipoledipole forces
Stick together better than
nonpolar molecules
Tend to have higher melting points,
boiling points, Hf, & Hv than
nonpolar substances.
Polar Molecules
Covalent bond where both
electrons in the bond are
donated by 1 atom.
Coordinate Covalent Bond
2 electrons to
contribute.
..
H:N:H
..
H
+ H+ 
No electrons
to contribute.
H
..
H:N:H
..
H
4 identical N-H
bonds!
Coordinate Covalent Bond
+
Both electrons in the
bond are donated by the
same atom.
Coordinate Covalent Bond
Compound that contains a
polyatomic ion.
Compound that exhibits both
covalent & ionic bonding
Metal
cation
: :
[Na]+1
[:O -
-1
H]
Polyatomic
anion
Lewis Diagram for NaOH
- -
H–N–H
+1
: :
H
:Cl:-1
H
Polyatomic
cation
Nonmetal
anion
Lewis Diagram for NH4Cl
- -
H–N–H
H
Polyatomic
cation
+1
: :
H
[:O - H]-1
Polyatomic
anion
Lewis Diagram for NH4OH
Subtract the
electronegativities of the
2 atoms.
How to calculate the polarity
of a bond
Use the symmetry &
conmposition of the molecule to
help you.
Don’t get a number for it.
Just polar or nonpolar overall.
To determine molecular polarity
Bent molecules
Pyramids, either trigonal or
square base
See-Saw
T-shape
Low symmetry shapes:
POLAR
Linear
Trigonal PLANAR
Tetrahedrons
Trigonal BIpyramids
Octahedrons
Square Planar
High symmetry shapes
All the ends or corners have to
match.
The corners or ends have to be
the same element.
For a high symmetry shape
to be NONPOLAR
CX4 is nonpolar. Corners match!
CXY3 and CX2Y2 are polar.
Corners don’t match!
Example of Symmetry &
Composition
Cdia, Cgraph, SiO2, & SiC
Four network covalent
substances are …
Crystal Lattices!
Network Covalent
Substances form …
High melting & boiling points
Hard
Brittle
Nonconductors
Properties of network
covalent substances
Strong directional covalent
bonds
Network Covalent
substances have …
Combinations of atomic orbitals
on an atom. Used to describe
bonding in molecules.
Hybridization
One-to-one correspondence
between number of domains and
type of hybridization.
# of domains
Type of
Geometry of
hybridization
hybrid
orbitals
# of p
orbitals left
over
2
sp
Linear
2
3
sp2
Trigonal
planar
1
4
sp3
Tetrahedral
0
5
dsp3
Trigonal
bipyramid
0
6
d2sp3
Octahedral
0
Figuring out hybridization
# of orbitals in = # of orbitals out.
Within a set of hybrid orbitals: orbitals have
same energy & same shape but point in
different directions.
Hybrid orbitals are different from the
atomic orbitals used to construct them.
“Rules” for Hybridization
Result from combination of one
s and three p atomic orbitals.
Produce 4 new orbitals pointing
to the corners of a tetrahedron.
sp3 hybrid orbitals
Formed by “head-on” overlap of
atomic orbitals.
Electron density is on the line
connecting the two nuclei
involved in the bond.
Sigma or  bonds
Result from side-to-side overlap
of p orbitals.
Two regions of overlap, above
and below the line connecting
the two nuclei.
Pi or  bonds