Bonding – General Concepts
What is a Bond?
• A force that holds atoms together.
• 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.
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.
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.
• Most Noble gases do not have values.
• Difference in electronegativity between
atoms tells us how polar.
Electronegativity: The ability of an
atom in a molecule to attract shared
electrons to itself.
Zero
Covalent
Intermediate
Polar
Covalent
Large
Ionic
Covalent Character
decreases
Ionic Character increases
Electronegativity Bond
difference
Type
Ionic Bonds
 Electrons are transferred
 Electronegativity differences are
generally greater than 1.7
 The formation of ionic bonds is
always exothermic!
Determination of
Ionic Character
Electronegativity
difference is not
the final
determination of
ionic character
Compounds are ionic
if they conduct
electricity in their
molten state
Coulomb’s Law
 Q1Q2 
E  (2.31 x 10 J  nm) 

 r 
 Q1Q2 
E

 r 
19
• Q is the charge.
• r is the distance between the centers.
• If charges are opposite, E is negative
– exothermic
• Same charge, positive E, requires energy
to bring them together.
– endothermic
Size of ions
• Ion size increases down a group.
• Cations are smaller than the atoms they
came from.
• Anions are larger.
• across a row they get smaller, and then
suddenly larger.
• First half are cations.
• Second half are anions.
Periodic Trends
• Across the period nuclear charge
increases so they get smaller.
• Energy level changes between anions and
cations.
Li+1
B+3
Be+2
C+4
N-3
O-2
F-1
Table of
Ion Sizes
Size of Isoelectronic ions
• Iso - same
• Iso electronic ions have the same # of
electrons
• Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
• All have 10 electrons.
• All have the configuration 1s22s22p6
Size of Isoelectronic ions
• Positive ions have more protons so they
are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Ionic Compounds
• We mean the solid crystal.
• Ions align themselves to maximize
attractions between opposite charges,
• and to minimize repulsion between like
ions.
• Can stabilize ions that would be unstable
as a gas.
• React to achieve noble gas configuration
Li < Na < C < O < F
Na < Li < C < O < F
F < O < C < Li < Na
Na < Li < F < O < C
Na < Li < C < F < O
82%
12%
6%
O
C
0%
i<
Na
<L
i<
C
F<
<
O
F<
<
Na
<L
Na
O
F<
Na
<L
i<
<C
C
<
<
Li
O
<
<
F
F
0%
<
O
<
<C
Na
Li
<
1.
2.
3.
4.
5.
List the following atoms in order of
increasing ionization energy: Li, Na, C,
O, F.
Sodium losing an electron is an ________
process and fluorine losing an electron is an
_______ process.
40%
20%
20%
15%
d
ic
n
m
in
fo
r
or
e
ne
ed
e
m
he
r
ex
ot
at
io
rm
ic
th
e
ic,
rm
m
ex
ot
he
er
m
ic ,
en
do
do
th
e
en
ic,
en
do
th
rm
er
m
ic ,
ex
ot
he
rm
rm
ic
ic
5%
ex
ot
he
endothermic, exothermic
exothermic, endothermic
endothermic, endothermic
exothermic, exothermic
more information needed
en
do
th
1.
2.
3.
4.
5.
Which of the following statements
is true about the ionization energy
of Mg+?
27%
18%
bo
ve
ea
o.
..
d
o.
..
no
ne
of
th
an
o
lt
eq
ua
lt
be
ill
It
w
It
w
ill
be
eq
ua
lt
o
an
d
an
d
io
th
e
o
o
lt
eq
ua
eq
ua
be
be
ill
ill
It
w
9%
o.
..
...
9%
It
w
1. It will be equal to the ionization
energy of Li.
2. It will be equal to and opposite in
sign to the electron affinity of Mg.
3. It will be equal to and opposite in
sign to the electron affinity of Mg+.
4. It will be equal to and opposite in
sign to the electron affinity of Mg2+.
5. none of the above
36%
74%
Choose the compound
with the most ionic bond.
16%
5%
5%
KC
l
LiF
l
Na
C
KF
LiCl
KF
NaCl
LiF
KCl
LiC
1.
2.
3.
4.
5.
l
0%
In which pair do both compounds
exhibit predominantly ionic
bonding?
100%
Ca
O
d
an
Rb
Cl
an
d
Na
F
nd
0%
H2
O
O3
0%
KI
a
BH
3a
nd
nd
5a
0%
3
HF
0%
Na
2S
O
PCl5 and HF
Na2SO3 and BH3
KI and O3
NaF and H2O
RbCl and CaO
PC
l
1.
2.
3.
4.
5.
Which of the following arrangements
is in order of increasing size?
.
S2
–
>K
+
>C
l–
. ..
Ca
2+
>
>
Ca
2+
Ga
3+
>
0%
>.
..
0%
S2
–
>C
Ca
2+
Ga
3+
>
Ga
3+
>
S2
–>
>
K+
>
l–
>C
l–
>.
..
..
>.
l–
Ca
2+
>C
K+
>
Ca
2+
S2
–
Ga3+ > Ca2+ > K+ > Cl– > S2–
S2– > Cl– > K+ > Ca2+ > Ga3+
Ga3+ > S2– > Ca2+ > Cl– > K+
Ga3+ > Ca2+ > S2– > Cl– > K+
Ga3+ > Ca2+ > S2– > K+ > Cl–
33%
0%
Ga
3+
>
1.
2.
3.
4.
5.
67%
Which of the following species
would be expected to have the
lowest ionization energy?
0%
M
g2
+
0%
O2
-
0%
50%
Na
+
50%
Ne
FNe
O2Mg2+
Na+
F-
1.
2.
3.
4.
5.
Sodium
Chloride
Crystal
Lattice
Ionic compounds form solids at ordinary temperatures.
Ionic compounds organize in a characteristic crystal
lattice of alternating positive and negative ions.
Forming Ionic Compounds
• Lattice energy - the energy associated
with making a solid ionic compound from
its gaseous ions.
• M+(g) + X-(g)  MX(s)
• This is the energy that “pays” for making
ionic compounds.
• Energy is a state function so we can get
from reactants to products in a round
about way.
Calculating Lattice Energy
• Lattice Energy = k(Q1Q2 / r)
• k is a constant that depends on the
structure of the crystal.
• Q’s are charges.
• r is internuclear distance.
• Lattice energy is greater with more highly
charged ions.
• This bigger lattice energy “pays” for the
extra ionization energy.
• Also “pays” for unfavorable electron
affinity.
Estimate DHf for Sodium Chloride
Na(s) + ½ Cl2(g)  NaCl(s)
Lattice Energy
Ionization Energy for Na
Electron Affinity for Cl
-786 kJ/mol
495 kJ/mol
-349 kJ/mol
Bond energy of Cl2
239 kJ/mol
Enthalpy of sublimation for Na
109 kJ/mol
Na+(g) + Cl-(g)  NaCl(s)
Na(g)  Na+(g) + e½ Cl2(g)  Cl(g)
Cl(g) + e-  Cl-(g)
Na(s)  Na(g)
Na(s) + ½ Cl2(g)  NaCl(s)
-786 kJ
+ 495 kJ
+ ½(239 kJ)
- 349 kJ
+ 109 kJ
-412 kJ/mol
Lattice Energies of Alkali Metals Halides (kJ/mol)
F-
Cl-
Br-
I-
Li+
1036
853
807
757
Na+
923
787
747
704
K+
821
715
682
649
Rb+
785
689
660
630
Cs+
740
659
631
604
Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)
OH-
O2-
Na+
900
2481
Mg2+
3006
3791
Al3+
5627
15,916
What about 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.
• They reach a distance with the lowest
possible energy.
• The distance between is the bond length.
Covalent Bonds
Polar-Covalent bonds
 Electrons are unequally shared
 Electronegativity difference between .3 and 1.7
Nonpolar-Covalent bonds
 Electrons are equally shared
 Electronegativity difference of 0 to 0.3
Covalent Bonding Forces
 Electron – electron
repulsive forces
 Proton – proton
repulsive forces
 Electron – proton
attractive forces
Bond Length Diagram
The Octet Rule
Combinations of elements tend to form so that
each atom, by gaining, losing, or sharing
electrons, has an octet of electrons in its
highest occupied energy level.
Diatomic Fluorine
Formation of Water by the Octet Rule
Comments About the Octet Rule
 2nd row elements C, N, O, F observe the octet
rule.
 2nd row elements B and Be often have fewer
than 8 electrons around themselves - they are
very reactive.
 3rd row and heavier elements CAN exceed the
octet rule using empty valence d orbitals.
 When writing Lewis structures, satisfy octets
first, then place electrons around elements
having available d orbitals.
Lewis Structures
Shows how valence electrons are arranged
among atoms in a molecule.
Reflects central idea that stability of a
compound relates to noble gas electron
configuration.
Completing a Lewis Structure -CH3Cl
Make carbon the central atom
Add up available valence electrons:
Join peripheral atoms
to the central atom
with electron pairs.
H
..
..
Complete octets on
H
atoms other than
hydrogen with remaining
electrons
C
..
H
Total = 14
..
Cl
..
..
C = 4, H = (3)(1), Cl = 7
Multiple Covalent Bonds:
Double bonds
H
H
C
H
H
C
H
C
H
H
Ethene
Two pairs of shared electrons
C
H
Multiple Covalent Bonds:
Triple bonds
H
C
C
H
H
C
C
H
Ethyne
Three pairs of shared electrons
Resonance
Resonance is invoked when more than one valid
Lewis structure can be written for a particular
molecule.
H
H
Benzene, C6H6
H
H
H
H
H
H
H
H
H
H
The actual structure is an average of the resonance
structures.
The bond lengths in the ring are identical, and
between those of single and double bonds.
Resonance Bond Length and Bond Energy
Resonance bonds are shorter and stronger
than single bonds.
H
H
H
H
H
H
H
H
H
H
H
H
Resonance bonds are longer and weaker than double
bonds.
Resonance in Ozone, O3
O
O
O
O
O
O
Neither structure is correct.
Oxygen bond lengths are identical, and
intermediate to single and double bonds
Resonance in Polyatomic Ions
Resonance in a carbonate ion:
Resonance in an acetate ion:
Localized Electron Model
Lewis structures are an application of the
“Localized Electron Model”
L.E.M. says: Electron pairs can be
thought of as “belonging” to pairs of
atoms when bonding
Resonance points out a weakness in the
Localized Electron Model.
Models
Models are attempts to explain how
nature operates on the microscopic level
based on experiences in the macroscopic
world.
Models can be physical
as with this DNA model
Models can be mathematical
Models can be theoretical
or philosophical
Fundamental Properties of Models
A model does not equal reality.
Models are oversimplifications, and are
therefore often wrong.
Models become more complicated as they
age.
We must understand the underlying
assumptions in a model so that we don’t
misuse it.
VSEPR – Valence Shell Electron Pair
Repulsion
X+E
Overall Structure
Forms
2
3
4
Linear
AX2
Trigonal Planar
AX3, AX2E
Tetrahedral
AX4, AX3E, AX2E2
Trigonal bipyramidal
AX5, AX4E, AX3E2, AX2E3
Octahedral
AX6, AX5E, AX4E2
5
6
A = central atom
X = atoms bonded to A
E = nonbonding electron pairs on A
VSEPR: Linear
AX2
CO2
VSEPR: Trigonal Planar
AX3
AX2E
BF3
SnCl2
VSEPR: Tetrahedral
AX4
CCl4
AX3E
PCl3
AX2E2
Cl2O
VSEPR: Trigonal Bi-pyramidal
AX5
PCl5
AX4E
SF4
AX3E2
ClF3
AX2E3
I 3-
VSEPR: Octahedral
AX6
SF6
AX5E
BrF5
AX4E2
ICl4-