Bonds
Forces that hold groups of
atoms together and make
them function as a unit.
Bond Energy
- It is the energy required to break a
bond.
- It gives us information about the
strength of a bonding interaction.
Bond Length
The distance where the system
energy is a minimum.
+
+
H atom
H atom
Sufficiently far apart
to have no interaction
+
+
H atom
H atom
The atoms begin to interact
as they move closer together.
+
+
Energy (kJ/mol)
08_130
0
H
HH
H
HH
-458
0
H2molecule
Optimum distance to achieve
lowest overall energy of system
(a)
H
0.074
Internuclear distance (nm)
(HH bond length)
(b)
H
Ionic Bonds
-
Formed from electrostatic attractions of
closely packed, oppositely charged ions.
-
Formed when an atom that easily loses
electrons reacts with one that has a high
electron affinity.
Ionic Bonds
E = 2.31  10
19
J nm (Q1Q2 / r )
Q1 and Q2 = numerical ion charges
r = distance between ion centers (in nm)
Periodic Trends
Electronegativity
The ability of an atom in a molecule
to attract shared electrons to itself.
 = (H  X)actual  (H  X)expected
Increasing electronegativity
08_132
H
Decreasing electronegativity
2.1
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
0.8
1.0
1.3
1.5
1.6
1.6
1.5
1.8
1.9
1.9
1.9
1.6
1.6
1.8
2.0
2.4
2.8
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
0.8
1.0
1.2
1.4
1.6
1.8
1.9
2.2
2.2
2.2
1.9
1.7
1.7
1.8
1.9
2.1
2.5
Cs
Ba
La-Lu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
0.7
0.9
1.0-1.2
1.3
1.5
1.7
1.9
2.2
2.2
2.2
2.4
1.9
1.8
1.9
1.9
2.0
2.2
Fr
Ra
Ac
Th
Pa
U
Np-No
0.7
0.9
1.1
1.3
1.4
1.4
1.4-1.3
(a)
Increasing electronegativity
H
Decreasing electronegativity
2.1
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
2.0
1.5
Co
Ni
Cu
Sc
Ti
V
Cr
Mn
1.5
1.6
1.6
1.5
1.8
1.9
1.9
1.9
1.0
1.3
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Sr
Y
Ag
1.9
1.4
1.9
2.2
1.2
1.8
2.2
1.0
1.6
2.2
Hf
Ta
W
Re
Os
Ir
Pt
Ba
La-Lu
1.7
1.9
2.2
2.2
2.2
Ca
0.7
1.0-1.2
1.3
1.5
Fr
Ra
Ac
Th
Pa
U
Np-No
1.1
1.3
1.4
1.4
1.4-1.3
0.9
Au
2.4
Si
P
S
Cl
2.5
3.0
1.5
2.1
Ga
Ge
As
Se
Br
Zn
1.6
1.8
2.0
2.4
1.6
Cd
In
Sn
Sb
Te
1.7
1.7
1.8
1.9
2.1
Hg
Tl
Pb
Bi
Po
At
1.9
1.8
1.9
1.9
2.0
2.2
1.2
Fe
3.0
2.5
4.0
3.5
1.8
Al
Mg
0.9
0.7
(b)
B
Be
F
O
N
C
2.8
I
2.5
Polarity
A molecule, such as HF, that has a center
of positive charge and a center of negative
charge is said to be polar, or to have a
dipole moment.
H F
+

08_131
H
F

H




 H F
F



 H F

H
F
 H F

H

F
H


 H F
F

 H F

(a)

(b)

08_133


+
H
O


H

(a)
(b)

08_134
3
+

N
H
H

H




(a)
(b)
Achieving Noble Gas Electron
Configurations (NGEC)
Two nonmetals react: They share
electrons to achieve NGEC.
A nonmetal and a representative group
metal react (ionic compound): The
valence orbitals of the metal are emptied
to achieve NGEC. The valence electron
configuration of the nonmetal achieves
NGEC.
Isoelectronic Ions
Ions containing the the same number of
electrons
(O2, F, Na+, Mg2+, Al3+)
O2> F > Na+ > Mg2+ > Al3+
largest
smallest
Lattice Energy
The change in energy when separated
gaseous ions are packed together to form
an ionic solid.
M+(g) + X(g)  MX(s)
Lattice energy is negative (exothermic)
from the point of view of the system.
Formation of an Ionic Solid
1. Sublimation of the solid metal
M(s)  M(g) [endothermic]
2. Ionization of the metal atoms
M(g)  M+(g) + e [endothermic]
3. Dissociation of the nonmetal
1/2X (g)  X(g)
[endothermic]
2
Formation of an Ionic Solid
(continued)
4. Formation of X ions in the gas phase:
X(g) + e  X(g) [exothermic]
5. Formation of the solid MX
M+(g) + X(g)  MX(s)
[quite
exothermic]
Lattice Energy = k(Q1Q2 / r )
Q1, Q2 = charges on the ions
r = shortest distance between centers of the
cations and anions
Models
Models are attempts to explain
how nature operates on the
microscopic level based on
experiences in the macroscopic
world.
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.
Bond Energies
Bond breaking requires energy (endothermic).
Bond formation releases energy (exothermic).
H = D(bonds broken)  D(bonds formed)
energy required
energy released
Localized Electron Model
A molecule is composed of atoms
that are bound together by sharing
pairs of electrons using the atomic
orbitals of the bound atoms.
Localized Electron Model
1. Description of valence electron
arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to
share electrons or hold lone pairs.
Lewis Structure
-
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.
Lewis Structures
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
1.
Least electronegative is central atom
2.
Connect others with single bond
3.
Fill octet of outer atoms
4.
Share additional pairs for octet on central atom if
necessary
5.
Extra electrons to central atom as pairs
6.
Use formal charge to help make decisions
Formal Charge
The difference between the number of
valence electrons (VE) on the free atom
and the number assigned to the atom in the
molecule.
We need:
1. # VE on free neutral atom
2. # VE “belonging” to the atom in the
molecule
Formal Charge
O C O
(-1)
(0)
(+1)
Not as good
O C O
(0)
(0)
Better
(0)
Resonance
Occurs when more than one valid Lewis
structure can be written for a particular
molecule.
These are resonance structures. The actual
structure is an average of the resonance
structures.
VESPR
VSEPR Model
The structure around a given
atom is determined principally
by minimizing electron pair
repulsions.
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the
way electron pairs are shared.
4. Determine the name of molecular structure
from positions of the atoms.
Electronic
Geometry
END