Bonding: General Concepts
Chapter 8
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.
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
 Q1Q 2 
J nm 

 r 
This is a statement of Coulomb’s Law where:
Q1 and Q2 = numerical ion charges
r = distance between ion centers (in nm)
When E is positive (+), repulsion is indicated.
When E is negative (-), attraction is indicated.
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
H
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)
Interaction of two hydrogen atoms.
Covalent Bonding
- covalent bonds are formed by sharing electrons
between nuclei.
H    H  H-H
- coordinate covalent bonds are bonds where both
shared electrons originate on the same atom
H3 N  + H +  H3 N-H +
Types of Covalent Bonds
Polar covalent bond -- covalent bond in which
the electrons are not shared equally
because one atom attracts them more
strongly than the other. A dipole moment
exists.
Nonpolar covalent bond -- covalent bond in
which the electrons are shared equally
between both atoms. No dipole moment
exists.
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
B
Be
2.0
1.5
Al
Mg
1.2
Ca
Sc
Ti
V
Cr
Mn
Co
Ni
Cu
1.8
1.9
1.9
1.9
1.0
1.3
1.5
1.6
1.6
0.8
Y
Zr
Nb
Mo
Tc
Rh
Pd
Sr
Ru
Ag
Rb
1.6
1.8
1.9
2.2
2.2
2.2
1.9
W
Re
Os
Ir
Pt
1.7
1.9
2.2
2.2
2.2
0.8
Cs
1.0
Ba
1.2
1.4
La-Lu
Hf
Ta
1.5
1.5
Fe
0.7
0.9
1.0-1.2
1.3
Fr
Ra
Ac
Th
Pa
U
Np-No
1.1
1.3
1.4
1.4
1.4-1.3
0.7
0.9
Au
2.4
Zn
Si
P
1.5
1.8
2.1
Ga
Ge
As
4.0
3.5
3.0
2.5
F
O
N
C
S
2.5
Se
2.4
Cl
3.0
Br
2.8
1.6
1.8
2.0
Cd
In
Sn
Sb
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
1.6
Te
(b)
Pauling Electronegativity Values
I
2.5
Percent Ionic Character
xA  xB
100% 
% Ionic Character (IC) 
xA
where xA is the larger electronegativity and xB
is the smaller value.
Watch significant figures!!!
Ionic Bond
Polar Covalent
Nonpolar Covalent
% IC > 50 %
% IC 5 - 50 %
% IC < 5 %
Three Possible Types of Bonds
Nonpolar Covalent
(Electrons equally
shared.)
Polar Covalent
(Electrons shared
unequally.)
Ionic
(Electrons are
transferred.)
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

F

H
F
 H
H


F
H



 H
F
F

 H

(a)
F
(b)
The Effect of an electric field on hydrogen fluoride molecules.
08_133


+
H
O



H

(a)
(b)
Dipole Moment for the water molecule.
08_134
3
+

N
H
H

H




(a)
(b)
Dipole moment for the ammonia molecule.
Nitrogen Trichloride
Does nitrogen trichloride exhibit a dipole
moment?
Yes. It has three nonpolar bonds but, also,
has a lone pair of electrons which makes
it assymetrical and therefore, polar.
Bond Polarity vs Molecular Polarity
Bond polarity depends upon electronegativity
difference (% Ionic Character)
Molecular polarity depends up the symmetry
of the molecule.
08_151
Nonpolar molecule--zero dipole moment.
Table 8.2 on page 357 in Zumdahl.
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.
See Table 8.3 on page 361 in Zumdahl.
08_136
O
F
(1.40)
140
(1.36)
136
S
Cl
(1.84)
184
(1.81)
181
Ga
Se
Br
(0.62)
62
(1.98)
198
(1.95)
195
Li
Be
(0.60)
60
(0.31)
31
Na
Mg
Al
(0.95)
95
(0.65)
65
(0.50)
50
K
Ca
(1.33)
133
(0.99)
99
Rb
Sr
In Sn Sb
(1.48)
148
(1.13)
113
(0.81) (0.71) (0.62)
81
71
62
Te
(2.21)
221
I
(2.16)
216
Sizes of ion related to position on the periodic table.
Isoelectronic Ions
Ions containing 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]
 Q1Q 2 
Lattice Energy = k 

 r 
Q1, Q2 = charges on the ions
r = shortest distance between centers of the
cations and anions
Mg2+(g) + O2-(g)
08_139
Electron affinity
737
Mg2+(g) + O(g)
Mg2+(g) +
-3916
1
2
O2(g)
247
Lattice
energy
2180
Ionization energy
Na(g) + F(g)
Na+(g) +
Mg(g) +
Mg(s) +
-602
MgO(s)
1
2
1
2
495
O2(g)
O2(g)
150
Overall
energy
change
109
1
2
F2(g)
77
Ionization
energy
Na(g) +
Na(s) +
1
2
1
2
-328
Electron
affinity
-923
Lattice
energy
Na+(g) + F-(g)
F2(g)
F2(g)
-570
NaF(s)
Comparison
.
of the energy changes in the formation of sodium
fluoride and magnesium oxide.
Bond Energies
Bond breaking requires energy (endothermic).
Bond formation releases energy (exothermic).
H = (bonds broken)  (bonds formed)
energy required
energy released
Draw the Lewis Structure for each reactant and
product before doing any calculations!
Single, Double, & Triple Bonds
Single bonds -- one shared pair of
electrons.
Double bonds -- two shared pairs of
electrons.
Triple bonds -- three shared pairs of
electrons.
See bond energy Tables 8.4 & 8.5 on pages 373-374
in Zumdahl.
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.
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
Covalent Compounds
Ionic Compounds
1



  
K
 F 
  
 F  F 

 
In ionic compounds, electrons are transferred
and ions are formed. In covalent compounds,
electrons are shared to form a molecule.
Potassium gains the stability of argon,
bromine of krypton, and fluorine of neon.
Lone Pairs & Bonding Pairs
 
 F  F 

 
Electrons shared between atoms are
bonding pairs. Electrons that are not
involved in bonding are called lone pairs.
Each fluorine has three lone pairs and
one bonding pair shared between them.
Electron Deficient Molecules
Beryllium chloride -- BeCl2 -- is electron deficient
with four electrons. It forms a linear molecule.
Boron trifluoride -- BF3 -- is electron deficient
with six electrons. It forms a trigonal planar
molecule.
See page 381 for the reaction between boron
trifluoride and ammonia.
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.
Rules for Writing Lewis
Structures
•
Sum the valence electrons from all the
atoms.
•
Use a pair of electrons to from a bond
between each pair of bound atoms.
•
Arrange remaining electrons to satisfy the
duet rule for hydrogen and the octet rule
for the second-row elements.
Lewis Structures
NO+
•5 e- + 6 e- - 1 e- = 10 e•Each atom has an octet
and is satisfied.
 N  O  



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 called a resonance hybrid.
See the resonance structures for the nitrate
ion on page 385 in Zumdahl.
Odd-Electron Molecules
NO2
• contains 17 electrons.
• cannot satisfy the octet rule.
• a more sophisticated model is neededthe molecular orbital model.
Stereochemistry
The study of the threedimensional arrangement of
atoms or groups of atoms within
molecules and the properties
which follow such arrangement.
VSEPR Model
Valence Shell Electron Pair
Repulsion -- 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.(Parent
Geometry)
4. Determine the name of molecular structure
from positions of the atoms.(Actual
Geometry)
Molecular Geometry
Parent Geometry is
Actual Geometry is the
electron pair arrangement
about the central atom.
arrangement of atoms about
the central atom.
•linear
•linear
•trigonal planar
•bent
•tetrahedral
•trigonal pyramid
•trigonal bipyramidal
•seesaw
•octahedral
•T-shaped
•square pyramid
•square planar
08_142
Lone
pair
N
N
H
H
H
(a)
(b)
Lone pair of electrons on the ammonia molecule.
08_143
Lone pair
Bonding
pair
O
Bonding
pair
O
H
(a)
H
Lone pair
(b)
O
H
(c)
H
Lone pairs on the water molecule.
08_144

Cl
Cl
P
P
Cl
Cl
Cl
Cl
Octahedral structure for phosphorus hexachloride.
08_145
Xe
Octahedral structure for xenon.
08_150
F
F
F
90 °
Xe
F
Xe
leads to the
structure
F
F
F
F
(a)
F
F
F
Xe
F
F
180°
F
leads to the
structure
Xe
F
F
(b)
Parent and actual geometry for xenon tetrafluoride.
08_152
I
I
I
I
I
I
I
I
I
(a)
(b)
(c)
Three possible arrangements of the electron pairs in triiodide ion.
08_06T
Table
8.6 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion
Number of
Electron Pairs
Arrangement of Electron Pairs
2
Linear
A
3
Trigonal
planar
A
4
Tetrahedral
A
90°
5
Trigonal
bipyramidal
6
Octahedral
120° A
A
Example
VSEPR Model Summary
•
Determine the Lewis structure(s) for the
molecule.
•
For molecules with resonance structures, use any
of the structures to predict the molecular
structure.
•
Sum the electron pairs around the central atom to
determine the parent geometry.
•
The arrangement of the pairs is determined by
minimizing electron-pair repulsions.(Actual
Geometry)
VSEPR Model Summary
(Continued)
•
Lone pairs require more
space than bonding pairs
since they are tightly
attracted to only one
nucleus. Lone pairs
produce slight distortions
of bond angles less than
120o.