BONDING
A. IONIC BONDS
Ionic bond – electrostatic attraction between two oppositely charged ions
+
_
cation + anion
+
or
_
metal + nonmetal
or

Very strong bond

Involves a transfer of electrons

Results in the formation of salts (or group 1A hydroxides)
+
_
metal + polyatomic ion
(All salts are ionic compounds but not all ionic compounds are salts.)
Ionic compound – compound composed of positive and negative ions
Ex:
Na + Cl  Na+Cl-
Na +
Cl
+
 Na Cl
+
Na Cl
Characteristics of Ionic Compounds:

Crystalline solids at room temperature

Have high melting points (MP indicates bond strength)

Tend to dissolve in water

Conduct electricity in the molten state
Look at the reaction of Na + Cl:
1.
Nao + ionization energy  Na+1 + 1 e-
ENDO (bonds breaking requires energy)
2.
Clo + 1e-  Cl- + electron affinity energy
EXO
(bonds forming releases energy)
3.
Na+1 + Cl-  NaCl + lattice energy
EXO
(bonds forming releases energy)
Lattice energy – the energy given off when two oppositely charged ions bond to form
a crystal lattice or solid salt.
Octet Rule – atoms react to achieve the electron configuration of a noble gas
One endothermic step and two exothermic steps results in a NET EXOTHERMIC process.
Compounds form to achieve the lowest possible energy.
Ionic Crystal Structure of Compounds
NaCl crystal:
In the 3 D structure of a NaCl
crystal, each ion is surrounded
by 6 oppositely charged ions.*
There are, therefore, many strong
ionic bonds between all adjacent
oppositely charged ions.
These large attractive forces result
in a stable structure with a very
high melting point.
A crystal’s arrangement is a regular repeating pattern called a lattice.

the shape of a crystal depends on the internal arrangement of the ions (or atoms)

the many possible crystal arrangements of ions depends on their sizes
Coordination number – the number of ions of opposite charge that surround each ion in a crystal
*the coordination number of NaCl is 6
B. METALLIC BONDS
Metallic bonds – attractions of free-floating valence electrons for positively charged metal ions
- may form between atoms of the same metal or atoms of different metals.
Alloys – solutions of different metals made by melting metals together and cooling them
Atoms that form metallic bonds must have:

vacant valence orbitals that can hold additional electrons

low ionization energies so that loosely held electrons are available for bonding
Atoms in metals are packed in compact crystal forms that are most commonly cubic in shape.
Characteristics of Metals:

malleable and ductile

good conductors of electricity
C. COVALENT BONDS
Covalent bond – mutual attraction of different nuclei to the same electron’s orbital
- occurs between nonmetal atoms

Very strong bond

Involves a sharing of electrons

Results in the formation of molecular compounds
Molecular compound – compound composed of nonmetals held together by covalent bonds
Ex:
Cl + F  Cl F
Cl + F  ClF
Cl
F
Characteristics of Molecular Compounds:

Tend to be liquids or gases at room temperature (though some are solids)

Have low melting points

Contain 2 types of forces:
1. intramolecular forces – strong covalent bonds within the molecules
2. intermolecular forces – weaker attractive forces between the molecules
Ex: Water molecules in an ice crystal:
+
INTRAMOLECULAR:
Strong covalent bonds
in which electrons are
shared by the O and H
_
+
+
_
+
INTERMOLECULAR:
The positive end of one H2O mc
is attracted to the negative end
of an adjacent H2O molecule.
_
+
+
Diatomic Molecule – molecule composed of 2 atoms (this includes all HONClBrIF’s)
Both nuclei attract the same
Look at H2 :
electrons
Both atoms strive
H
1s1
x
H
1s1
to fill their 1s orbital
so both H’s attract
the pair of bonding
electrons
+
+
H–H
Shape of H2 is linear
+
+
atomic orbitals overlap
to form new molecular orbital
s to s bonding is non-directional because s orbitals can approach each other from any direction
and overlap to form a bond
Look at F2 :
F
2pz is where bonding occurs
py
py
x
1s2 2s2 2px2y2z1
x
F
x
F
+
pz
py
new molecular
orbital forms
+
pz
x
F
1s2 2s2 2px2y2z1
2pz orbitals overlap
so now both atoms
achieve a stable octet
py
F–F
Shape of F2 is linear
directional bonds - bonding in which atom’s orbitals must approach at a specific direction in order for the
orbitals to overlap to form a bond (s to p, p to p, s to d, & p to d)
non-polar covalent bond – covalent bond in which there is an equal attraction for the shared electrons
All bonds in HONClBrIFs are nonpolar (purely) covalent bonds where there is an equal sharing of electrons.
electronegativity – a number that indicates the relative attraction an atom has for a shared pair of
electrons in a bond; the higher the number, the greater the attraction
electronegativity difference ( EN) is an indication of the type of bond formed:
electronegativity difference 0 – .3 = nonpolar covalent bond (equal sharing of electrons)
electronegativity difference .4 – 1.6 = polar covalent bond
(unequal sharing of electrons)
electronegativity difference
(transfer of electrons)
> 1.7 = ionic bond
polar covalent bond – a covalent bond between 2 atoms in which one atom has a greater pull (attraction)
on the shared electrons (unequal sharing)
Look at H2O :
O
_
..
x
O 1s2 2s2 2px2y1z1
+
H
x
O
x
H
O
+
:
+
H
H
H
1s1
1 H’s 1s overlaps the
O’s 2px lobe and the
H2O mc’s shape
is bent b/c the
+ H
bonds are polar b/c O has
H 1s1
other H’s 1s overlaps
the O’s 2pz lobe
y and z axes are
perpendicular
a greater pull on e- than H
O = 3.5; H = 2.1 EN = 1.4)
H
MOLECULAR POLARITY
polar molecule – a molecule in which one region of the molecule has a higher electron density causing
the molecule to have positive and negative regions or ends.
The polarity of a molecule is determined by the following:

the polarity of the bonds

the overall distribution of the valence electrons

the symmetry/shape of the molecule
Compare the following molecules:
Methane
CH4
H
H C H
H
Ammonia
Water
NH3
H2O
_
--
H N H
H
H O
H
Bonds: C - H
C = 2.5
H = 2.1
EN = .4
Bonds: N - H
N = 3.0
H = 2.1
EN = .9
Bonds: O - H
O = 3.5
H = 2.1
EN = 1.4
bonds are slightly polar
bonds are moderately polar
bonds are very polar
shape is tetrahedral
shape is pyramidal
shape is bent
molecule is non-polar
molecule is moderately polar
molecule is very polar
b/c mc is symmetrical
b/c mc is asymmetrical
b/c mc is asymmetrical
and has 1 exposed e- pair
and has 2 exposed e- pairs
gas at room temperature
liquid that evaporates easily
liquid at room temp
(no attractive forces b/t mc)
(weak intermolecular forces)
(strong intermolecular forces)
low boiling point
high boiling point
MOLECULAR GEOMETRY
VSEPR (valence shell electron pair repulsion) Theory – because electron pairs repel, molecules adjust
their shape so that the valance (outer shell) electron pairs are as far apart as possible.
Shape
Bond Angle
Linear
180o
Example
.. ..
: F: F:
..
Bent
..
H:O :
105o
F2
..
H2O
..
H
(The asymmetric water molecule has 2 exposed (unshared) electron pairs on one side of the
molecule. Due to their repulsion of like charges, they require more space and consequently
decrease the angle between the two hydrogen atoms.)
Trigonal Planar
120o
F
..
B
F
Pyramidal
107o
BF3
F
..
H : N : H
NH3
..
H
Tetrahedral
109.5o
H
..
H : C : H
..
H
CH4
POLYATOMIC IONS
Polyatomic ion – a group of atoms that act as a unit and carry a charge

Made up of nonmetals

Bonds WITHIN polyatomic ions are covalent bonds

Cannot exist independently (can only exist as part of a compound)
Examples:
Hydroxide
OH-1
O H
Hydronium
Ammonium
H3O+1
NH4+1
H
O
H
H
+1
H
H
N
+1
H
H
+1
[ O - H ] -1
H–O–H
H
(linear)
H
H-N-H
H
+1
(pyramidal)
(tetrahedral)
Sulfite
Sulfate
Phosphate
SO3-2
SO4-2
PO4-3
O
O
S
O
-2
O
O
S O
O
O
O
O S O
-2
(pyramidal)
O
-2
O
O S O
O
(tetrahedral)
P
-3
O
O
-2
O
O P O
O
-3
(tetrahedral)
Coordinate covalent bond – a covalent bond in which one atom contributes both electrons in the
shared pair
MOLECULAR ORBITALS (orbitals of molecules)
Molecular orbitals are the result of the overlap of atomic orbitals when 2 atoms form a bond.
Types of molecular orbitals:
1. sigma bond – bond formed when 2 atomic orbitals combine to form a molecular orbital
that is symmetrical along the axis connecting 2 nuclei

END-TO-END orbital overlap is extensive therefore bonds are very strong
Two “s” atomic orbitals can overlap to form a sigma-bond molecular orbital:
+
s atomic orbital
+
s atomic orbital
sigma-bonding molecular orbital
Two “p” atomic orbitals can overlap head to head to form a sigma-bonding molecular orbital:
+
p atomic orbital
+
p atomic orbital
sigma-bonding molecular orbital
2. pi bond - bond formed when 2 atomic orbitals combine to form a molecular orbital in
which the bonding electrons are located in sausage shaped regions above and
below the bond axis

SIDE-TO-SIDE orbital overlap is not as extensive (as in a sigma bond) therefore
the bonds tend to be weaker
Two “p” atomic orbitals can overlap side-to-side to form a pi-bonding molecular orbital:
+
p atomic orbital
+
p atomic orbital
pi-bonding molecular orbital
ORBITAL HYBRIDIZATION
Hybridization – the combining of 2 or more orbitals of nearly the same energy into orbitals of equal energy

occurs when atoms promote electrons into nearby orbitals to increase their bonding capacity
(occurs in atoms of Gr. 2, 3, 4 or 6A)

occurs when several atomic orbitals mix to form the same number of hybrid orbitals

hybrid orbitals are equivalent because they have the same size, shape and energy

all sigma bonds and lone pair electrons (exposed or nonbonding e-) require hybrid orbitals
Look at BeH2:
Beryllium promotes one of its2s electrons to its empty 2p orbital
Be’s one 2s orbital and one 2p orbitals then mix to form two sp hybrid orbitals
2s
2p
hybrid sp orbitals
Be
mix to form
unhybridized 2p orbitals
and
Now the 1s orbitals of the Hydrogen atoms can overlap with the sp hybrid orbitals of the Be:
sp
The resulting molecule is linear:
H – Be – H
Be
because the sp hybrid orbitals form 180o angles
H
H
s
H
atomic
orbital
and one hydrogen overlaps on each end.
s
Be
hybrid
orbitals
H
atomic
orbital
Bonding
Beryllium Hydride
molecule
The H – Be bonds are sigma bonds because they are the result of the end-to-end orbital overlap.
Hybridization of Carbon - Carbon promotes one of its 2s electrons to its empty 2p orbital in order to
s
p
p
p
increase its bonding capacity.
Look at Methane: When methane forms – Carbon’s one 2s orbital and three 2p orbitals mix to
( CH4 )
s
form four sp3 hybrid orbitals to provide four bonding sites for the four H atoms.
p
p
p
C:
s
s
s
s
sp3
H’s:
H
+
Hydrogen
atomic orbitals
H
H
H
+
Carbon
hybrid orbitals
Hydrogen
atomic orbitals
Methane molecule
All four bonds formed are sigma bonds because they are the result of end-to-end orbital overlap.
The shape of the methane molecule is tetrahedral because the sp3 bond angles are 109.5o
Look at Ammonia (NH3) :
s
p
p
p
N
When H bonds with N, four sp3 molecular orbitals must form:
Three of the orbitals provide
bonding sites for the H’s
H’s
H
H
H
One of the orbitals provides a
site for the lone pair of e..

Not all of the hybrid orbitals of the central atom must be used for bonding,
N
H
That is, lone pair electrons can also be accommodated in hybrid orbitals.
H
H

The 3 bonds formed are sigma bonds because they are the result of end-to-end orbital overlap.

The sp3 bond angles here are 107o, because the bonding electrons are repeled by the unshared
(lone) pair of electrons and require more room, therefore the molecular shape is pyramidal.
MULTIPLE BONDS
Look at Ethene: ( C2H4 ) : Each carbon will form two C – H single bonds and one C = C double bond
H
C
H
single bond - 1 shared pair of electrons
H
double bond - 2 shared pairs of electrons
C
H
In order for each Carbon to bond to three atoms: One s and two p orbitals of each C form
three sp2 hybrid bonding orbitals
s
p
p
p
Each Carbon’s additional p orbital remains unhybridized
H
O
H
O
C
O
(C)
T

Two of the sp2 orbitals of each C overlap end-to-end with an s orbital of H to form a sigma bond.

The other sp2 orbital of each Carbon overlaps end-to-end to form a C – C sigma bond.*

The unhybridized p atomic orbital of each Carbon overlap side-to-side to form a C – C pi bond.*
*The C = C double bond is comprised of 1 sigma bond and 1 pi bond.

As a result of the sp2 orbital bonding, the H – C – H bond angles are 120o.
Hydrogen
atomic orbitals
Carbon atomic orbitals
and hybrid orbitals
Hydrogen
atomic orbitals
Ethene molecule
Look at Ethyne ( C2H2 ) : Each carbon will form one C – H single bond and one C
H
C
C
H
=
C triple bond
triple bond - 3 shared pairs of electrons
In order for each Carbon to bond to two atoms: One s and one p orbital of each C form
two sp hybrid bonding orbitals
s
p
p
p
Each Carbon’s additional two p orbitals remain unhybridized
H
O
C
O
(C) (C)
T T

One of the sp orbitals of each C overlap end-to-end with an s orbital of H to form a sigma bond.

The other sp orbital of each Carbon overlaps end-to-end to form a C – C sigma bond.*

The unhybridized p atomic orbitals of each Carbon overlap side-to-side to form two C – C pi bonds.*
*The C = C triple bond is comprised of 1 sigma bond and 2 pi bond.

As a result of the sp orbital bonding, the H – C – H bond angles are 180o and the shape of the
molecule is linear.

Hydrogen
atomic orbitals
Carbon atomic orbitals
and hybrid orbitals
There are 3 types of hybrid orbitals:
Hydrogen
atomic orbitals
Ethyne molecule
To determine the type of hybrid orbitals formed:

sp – comprised of 2 orbitals
1. Count the # of atoms bonded (to the central atom)
or count the number of sigma bonds

sp2 – comprised of 3 orbitals
2. Count the number of non-bonding electron pairs

sp3 – comprised of 4 orbitals
3. Sum 1 + 2 above to find # of hybrid orbitals
4. Match # of hybrid orbitals to type (see bullets at left)
RESONANCE

the extremely rapid shifting of the bonding of a molecule between 2 or more stable arrangements

2 or more equally valid electron dot structures can be drawn for a molecule

the bonds are considered a mix or hybrid of the extremes

makes a molecule more stable
Ozone O3
..
O
O
O
O
Sulfur Trioxide
O
O
O
O
O
O
O
SO3
O
..
S
O
O
O
O
O
O
O
S
S
S
O
O
O
Benzene C6H6
Or
O
O