AP Notes Chapter 8
Bonding and Molecular Structure:
Fundamental Concepts
Valence e- and Bonding
Covalent
Ionic
Resonance & Exceptions to Octet Rule
Bond Energy & Length
Structure, Shape & Polarity of Compounds
What is a Bond?






A force that holds atoms together.
Why?
We will look at it in terms of energy.
Bond energy the energy required to break
a bond.
Why are compounds formed?
Because it gives the system the lowest
energy.
Covalent compounds?





The electrons in each atom are attracted to
the nucleus of the other.
The electrons repel each other,
The nuclei repel each other.
The reach a distance with the lowest
possible energy.
The distance between is the bond length.
Thus Hydrogen is Diatomic!
Bond Formation
Covalent Character
e
Why Isn’t Helium Diatomic?
. .
He2
E
.
He + He
Inter-nuclear Distance
.
F+F
F2
2p ____ ____ ___
2s ____
F
___ ____ ____ 2p
____ 2s
F
Ionic Bonding



An atom with a low ionization energy
reacts with an atom with high electron
affinity.
The electron moves.
Opposite charges hold the atoms together.
Li
+
Cl
1s22s1
[Ne] 3s23p5
2s ___
3p _____ _____ ___
1s _____
3s _____
[Ne]
Li + Cl
2s ___
3P _____ _____ _____
1s _____
3s _____
[Ne]
LiCl
2s ___
3P _____ _____ _____
1s _____
3s _____
[Ne]
Electronegativity
Describes the relative
ability of an atom
within a molecule to
attract a shared pair
of electrons to itself.
Electronegativity
Pauling electronegativity
values, which are unitless, are the norm.
Electronegativity
Range from 0.7 to 4.0
Figure 9.9 – Kotz & Treichel
Bond: A - B
DEN = | ENA - ENB |
Bond Character
“Ionic Bond” - Principally Ionic
Character
“Covalent Bond” - Principally
Covalent
Character
Determining Principal
Character of Bond
DEN
~0
covalent
ionic
1.7
~4
F-F
Non-polar
EN = 0
DEN = |3.0 - 3.5|
= 0.5
N-O
N
O
Slightly polar
Ca - O
Ca
D EN = |1.0 - 3.5|
= 2.5
O
Ionic Bond with some
covalent character
Electronegativity





The ability of an electron to attract shared
electrons to itself.
Pauling method
Imaginary molecule HX
Expected H-X energy =
H-H energy + X-X energy
2
D = (H-X) actual - (H-X)expected
Electronegativity






D is known for almost every element
Gives us relative electronegativities of all
elements.
Tends to increase left to right.
decreases as you go down a group.
Noble gases aren’t discussed.
Difference in electronegativity between
atoms tells us how polar.
Zero
Covalent
Intermediate
Polar
Covalent
Large
Ionic
Covalent Character
decreases
Ionic Character increases
Electronegativity Bond
difference
Type
Dipole Moments




A molecule with a center of negative
charge and a center of positive charge is
dipolar (two poles),
or has a dipole moment.
Center of charge doesn’t have to be on an
atom.
Will line up in the presence of an electric
field.
How It is drawn
d+ d-
H-F
Which Molecules Have Them?




Any two atom molecule with a polar bond.
With three or more atoms there are two
considerations.
There must be a polar bond.
Geometry can’t cancel it out.
Ionic Radii -- Cations
Ionic Radii -- Anions
Molecular Polarity
Vector Sum of Bond
Polarities
MgBr2
Mg - Br
Br
EN = |1.2 - 2.8| = 1.6
Mg
Br
Covalent BOND w/much ionic
character, BUT NON-POLAR molecule
Lewis
Structures
The most important
requirement for the
formation of a stable
compound is that the
atoms achieve noble
gas e configuration
Valence Shell Electron
Pair Repulsion Model
(VSEPR)
The structure around a
given atom is determined
principally by minimizing
electron-pair repulsions
VSEPR
Electron Bond
pairs Angles
2
180°
Underlying
Shape
Linear
3
120°
4
109.5°
Tetrahedral
5
90° &
120°
6
90°
Trigonal
Bipyramidal
Octagonal
Trigonal Planar
LEWIS STRUCTURES
1 : draw skeleton of species
2 : count e- in species
3 : subtract 2
e
for each bond
in skeleton
4 : distribute remaining e-
Distinguish Between
ELECTRONIC
Geometry
&
MOLECULAR
Geometry
CH4
Bond angle = 109.50
Electronic geometry: tetrahedral
Molecular geometry: tetrahedral
+
H3O
Bond angle ~ 1070
Electronic geometry: tetrahedral
Molecular geometry: trigonal pyramidal
H2O
Bond angle ~ 104.50
Electronic geometry: tetrahedral
Molecular geometry: bent
NH2
Bond angle ~ 104.50
Electronic geometry: tetrahedral
Molecular geometry: bent
“Octet Rule” holds for
connecting atoms, but
may not for the central
atom.
BaI2
Bond angle =1800
Electronic geometry: linear
Molecular geometry: linear
BF3
Bond angle =1200
Electronic geometry: trigonal planar
Molecular geometry: trigonal planar
PF5
Bond angle = 1200 & 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: trigonal bipyramidal
SF4
Bond angle = 1200 & 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: see-saw
ICl3
Bond angle <= 900
Electronic geometry: trigonal bipyramidal
Molecular geometry: T-shape
I3
Bond angle = 1800
Electronic geometry: trigonal bipyramid
Molecular geometry: linear
PCl6
Bond angle = 900
Electronic geometry: octahedral
Molecular geometry: octahedral
BrF5
Bond angle ~ 900
Electronic geometry: octahedral
Molecular geometry: square pyramidal
ICl4
Bond angle = 900
Electronic geometry: octahedral
Molecular geometry: square planar
Actual shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
2
3
3
4
4
4
2
3
2
4
3
2
0
0
1
0
1
2
linear
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
5
5
5
5
5
4
3
2
0
1
2
3
trigonal bipyrimidal
See-saw
T-shaped
linear
Actual Shape
NonElectron Bonding Bonding
Pairs
Pairs
Pairs Shape
6
6
6
6
6
6
5
4
3
2
0
1
2
3
1
Octahedral
Square Pyramidal
Square Planar
T-shaped
linear
What happens
when there are not
enough electrons
to “satisfy” the
central atom?
EXAMPLES
Ethene
Acetic Acid
Oxygen
Nitrogen
RESONANCE
&
FORMAL CHARGE
 Sometimes
Resonance
there is more than one valid
structure for an molecule or ion.
 NO3
 Use double arrows to indicate it is the
“average” of the structures.
 It doesn’t switch between them.
 NO2
 Localized electron model is based on
pairs of electrons, doesn’t deal with odd
numbers.
EXAMPLES
Nitrate ion
Ozone
FORMAL CHARGE
the charge assigned to an atom
in a molecule or polyatomic ion
FC atom = Family# - [LPE + ½(BE)]
Sum FC’s atoms = ion charge
Closer sum FC’s is to zero more stable
Formal Charge
 For
molecules and polyatomic ions
that exceed the octet there are
several different structures.
 Use charges on atoms to help
decide which.
 Trying to use the oxidation numbers
to put charges on atoms in
molecules doesn’t work.
Formal Charge
The difference between the number of
valence electrons on the free atom and
that assigned in the molecule.
 We count half the electrons in each bond
as “belonging” to the atom.

SO4-2
 Molecules try to achieve as low a formal
charge as possible.
 Negative formal charges should be on
electronegative elements.

Assignment of
e
1. Lone pairs belong
entirely to atom in
question
2. Shared e- are divided
equally between the
two sharing atoms
The sum of the formal
charges of all atoms in
a species must equal
the overall charge on
the species.
A useful equation
 (happy-have)
POCl3
-2
 SO4
-2
 SO3
-2
 PO4
 SCl2

/ 2 = bonds
P is central atom
S is central atom
S is central atom
P is central atom
S is central atom
Exceptions to the octet
BH3
 Be and B often do not achieve octet
 Have less than and octet, for electron
deficient molecules.
 SF6
 Third row and larger elements can
exceed the octet
 Use 3d orbitals?
 I3

Exceptions to the octet

When we must exceed the octet, extra
electrons go on central atom.
ClF3
 XeO3
 ICl4
 BeCl2

If nonequivalent Lewis
structures exist, the
one(s) that best describe
the bonding in the
species has...
FAVORED LEWIS
STRUCTURES
1. formal charges closest to
zero
2. negative formal charge
is on the most
electronegative atom
EXAMPLES
Carbon dioxide
Thiocyanate ion
Sulfate ion
BOND
ENERGY &
LENGTH
Bond Energies
E = (Bonds Broken) – (Bonds Made)
Bonds form between atoms
because bonded atoms
exhibit a lower energy.
Thus, energy is required to
break bonds and energy is
released when bonds are
formed.
Bond Order = # bonds to a specific set
of elements
C-C the BO=1
C=C the BO=2
C
C the BO=3
Fractions are possible
COVALENT BONDS
Bond Dissociation
Energy
Table 9.9 (text)
Bond Energy
(kJ/mol)
H-F 565
H-Cl 432
H-Br 366
H-I
299
Bond Energy
(kJ/mol)
Cl-Cl 242
Br-Br 193
I-I
151
Bond Energy
(kJ/mol)
CC
CC
CC
346
610
835
Bond Energy
(kJ/mol)
N  N 163
N  N 418
N  N 945
Use bond energies to
predict DHc for acetylene
(C2H2).
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Internuclear Distance
Energy
0
Bond Length
Internuclear Distance
Energy
Bond Energy
0
Internuclear Distance
Bond:
CC
CC
CC
Energy Length
(kJ/mol) (pm)
346
610
835
154
134
121
Bond:
CO
CO
Energy Length
(kJ/mol) (pm)
358
745
143
122
Binary Ionic Compounds
metal(s) + non-metal (g) ---> salt(s)
Lattice Energy
Energy change occurring
when separated
gaseous ions are
packed together to form
an ionic solid
+
M (g) + NM (g) --> M-NM
What is the lattice
energy of NaCl(s)?
+
Na (g)
+
Cl (g)
---> NaCl(s)
Lattice Energies
LiCl 834
NaCl 769
KCl 701
Li2O 2799
Na2O 2481
K2O 2238
:
:
:
:
:
:
BeCl2 3004
MgCl2 2326
CaCl2 2223
BeO 4293
MgO 3795
CaO 3414
LE = Lattice Energy
 Q 1Q 2 
LE  k

 r 
Where: k = proportionality constant
dependent on structure of solid and on
electron configuration of the ions
Where: Q1 & Q2 = charges on the ions
Where: r = the shortest distance
between the centers of the cation and
anion