Chemical Bonding:
Basic Concepts
Chapter 6
Why do atoms
form bonds?
To become more stable
What is a bond?
• A mutual electrical attraction between the
nuclei and valence electrons of different
atoms that binds the atoms together.
Bonding Forces
Electron – electron
repulsive forces
Proton – proton
repulsive forces
Electron – proton
attractive forces
Three types of bonds
What type of bond will form?
• Type of bond is dependent on differences in
electronegativity
Remember- Electronegativity is …
A measure
of the
ability of
an atom in
a chemical
compound
to attract
electrons
Rule of thumb
• Metal and non-metal usually form an ionic
bond
• Non-metal and non-metal from a covalent
bond
• Metals come together to form metallic bonds
What type of bond?
Elements
Electronegativity
difference
Bond
Type
Morenegative
atom
Chlorine
& Calcium
3.0 – 1.0
= 2.0
Ionic
Chlorine
Chlorine
& Oxygen
3.5 – 3.0
= 0.5
Covalent
Bromine &
Chlorine
3.0 – 2.8
= 0.2
Covalent
Oxygen
Chlorine
Pause for a Cause
Create a “PFC” page in your notebook to use for this unit.
PFC #1
Determine what type of
bond is present in each
of the following
1. Magnesium chloride
2. Water (H2O)
3. Chlorine gas (Cl2)
4. Tungsten shavings
Check your answers:
Properties of each type of bond
Covalent Bond
A chemical bond that results in the
sharing of electrons between two atoms
Characteristics of Covalent Bonds
a) Usually occur between atoms of nonmetals
b) Result in a particle called a molecule
c) Have low melting points (therefore are not
solids at room temperature)
d) Do not usually conduct electricity.
Ionic Bond
• The chemical bond resulting from the
electrostatic attraction between
positive and negative ions.
Electrons are transferred
• Opposite charges attract (electrostatic force)
Ionic substances have predictable characteristics.
a) Have extremely high melting points
b) Tend to be soluble in water.
c) Usually form a crystalline lattice structure as a solid.
d) Are good conductors of electricity in the molten state.
Metallic Bonding
Bonding that occurs between atoms of a metal due to
delocalized electrons forming an electron sea
Characteristics of Metallic Bonding
•The highly mobile “sea of electrons” in metals
account for its high conductivity.
•Metals are better conductors of heat and
electricity than ionic and molecular compounds.
•Metals are also malleable, ductile and shiny.
Pause for a Cause #2
Tell what kind of bonds I have…
1. Heat me up and I can
conduct electricity.
2. My electrons flow like fish
in the sea.
3. My bonding partner and I
are very dull (not shiny).
4. Write your own riddle
about one type of bond.
Check your answers
1. Ionic
2. Metallic
3. Covalent
4. Let’s hear some of yours!
How strong are bonds?
Ionic (Lattice energy)
Dependent on two factors
Covalent (bond strength)
Dependent on two factors
1. Charge
1. Bond type
(single/double/triple)
2. Size
2. Size (Must remember your
size trends)
Shorter bonds are stronger
Longer bonds are weaker
Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely
separate one mole of a solid ionic compound into gaseous
ions.
cmpd
lattice energy
MgF2
2957
MgO
3938
LiF
1036
LiCl
853
Q= +2,-1
Q= +2,-2
r F < r Cl
Shorter bonds are stronger
bonds
Lengths of Covalent Bonds
Bond
Type
Bond
Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond
9.4
Pause for a Cause #3
Label each set as covelent or
ionic and circle the substance
with the greatest lattice or bond
energy
1. MgCl2 or MgO
2. H2O or H2S (both have
single bonds)
3. N2 (triple bond) or H2
(single bond)
4. NaCl or KCl
Check your answers
Polar covalent bond or polar bond is a covalent bond with
greater electron density around one of the two atoms
electron poor
region
H
electron rich
region
F
e- poor
e- rich
H
F
d+
d-
9.5
Chemical
compounds tend
to form so that
each atom, by
gaining, losing, or
sharing electrons,
has an octet of
electrons in its
highest occupied
energy level.
The Octet Rule
The Ionic Bond
Li +
1s22s1
Li+
F
1s22s22p5
-
F
Li+ +
F
-
22s22p6
[He]
1s21s[Ne]
Li
e +
F
Li+ + e
F
-
Li+
-
F
-
A covalent bond is a chemical bond in which two or more electrons are shared by two
atoms.
Why should two atoms share electrons?
F
+
7e-
F
F
7e-
8e-
F
8e-
Lewis structure of F2
single covalent bond
lone pairs
F
F
lone pairs
F
F
single covalent bond
lone pairs
lone pairs
Writing Lewis Structures
1. Draw skeletal structure of compound showing what
atoms are bonded to each other. Put least
electronegative element in the center.
2. Count total number of valence e-. Add 1 for each
negative charge. Subtract 1 for each positive charge.
3. Complete octets for as many atoms as possible,
starting with the outer atoms.
4. If all octets are not filled, create double or triple
bonds within the structure (with C, S, P, O or N)
9.6
Write the Lewis structure of nitrogen trifluoride (NF3).
F
N
F
F
9.6
Write the Lewis structure of the carbonate ion (CO32-).
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
9.6
A resonance structure is one of two or more Lewis structures for a single molecule
that cannot be represented accurately by only one Lewis structure.
-
+
O
O
-
O
+
O
O
O
What are the resonance structures of the
carbonate (CO32-) ion?
-
O
C
O
O
-
O
C
O
O
-
-
-
O
C
O
O
-
9.8
Pause for a Cause #5
Draw Lewis structures for the following
molecules
1. HBr
2. CF4
3. NH2Cl
4. SiCl3Br
5. NH3
6. H2S
Violations of the Octet Rule
B and Sn tends to form bonds
with 6 electrons.
3rd row and heavier elements
CAN exceed the octet rule using
empty valence d orbitals.
Be tends to form bonds with 4
electrons.
SF4
BF3
Two possible skeletal structures of formaldehyde (CH2O)
H
H
C
O
H
C
O
H
An atom’s formal charge is the difference between the number of valence
electrons in an isolated atom and the number of electrons assigned to that atom
in a Lewis structure.
formal charge
on an atom in
a Lewis
structure
=
total number
total number
of valence
of nonbonding
electrons in electrons
the free atom
-
1
2
(
total number
of bonding
electrons
)
The sum of the formal charges of the atoms in a molecule or ion must equal the
charge on the molecule or ion.
9.7
H
-1
+1
C
O
formal charge
on an atom in
a Lewis
structure
H
=
total number
of bonding
pairs
+
total number
of nonbonding
electrons
formal charge =
on C
formal charge =
on O
9.7
H
0
0
C
O
H
formal charge
on an atom in
a Lewis
structure
=
C – 4 eO – 6 e2H – 2x1 e12 etotal number
of
-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
total number
of nonbonding
electrons
-
1
2
(
total number
of bonding
electrons
)
formal charge = 4 - 0 - ½ x 8 = 0
on C
formal charge = 6 - 4 - ½ x 4 = 0
on O
9.7
Formal Charge and Lewis Structures
1.
For neutral molecules, a Lewis structure in which there are no formal
charges is preferable to one in which formal charges are present.
2.
Lewis structures with large formal charges are less plausible than those with
small formal charges.
3.
Among Lewis structures having similar distributions of formal charges, the
most plausible structure is the one in which negative formal charges are
placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O?
H
-1
+1
C
O
H
H
0
0
C
O
H
9.7
What shape do molecules have?
Determined by Valence Shell Electron Pair
Repulsion Theory
VSEPR Theory- electrostatic repulsion between
the valence-level pairs surrounding an atom
causes these pairs to be oriented as far
apart as possible.
What does it mean?
Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic repulsions between the
electron (bonding and nonbonding) pairs.
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
Arrangement of
electron pairs
Molecular
Geometry
linear
linear
B
B
10.1
Cl
Be
Cl
lone pairs
on to
central
atom
20
atoms
bonded
central
atom
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
linear
linear
AB3
3
0
trigonal
planar
trigonal
planar
Arrangement of
electron pairs
Molecular
Geometry
10.1
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
linear
linear
AB3
3
0
trigonal
planar
trigonal
planar
AB4
4
0
Arrangement of
electron pairs
tetrahedral
Molecular
Geometry
tetrahedral
10.1
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
linear
linear
AB3
3
0
trigonal
planar
trigonal
planar
AB4
4
0
tetrahedral
tetrahedral
0
trigonal
bipyramidal
trigonal
bipyramidal
AB5
5
Arrangement of
electron pairs
Molecular
Geometry
10.1
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB2
2
0
linear
linear
AB3
3
0
trigonal
planar
trigonal
planar
AB4
4
0
tetrahedral
tetrahedral
trigonal
bipyramidal
octahedral
Arrangement of
electron pairs
AB5
5
0
trigonal
bipyramidal
AB6
6
0
octahedral
10.1
Molecular
Geometry
10.1
10.1
lone-pair vs. lone pair >
repulsion
lone-pair vs. bonding >
pair repulsion
bonding-pair vs. bonding
pair repulsion
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
trigonal
planar
bent
AB3
3
0
trigonal
planar
AB2E
2
1
trigonal
planar
10.1
VSEPR
Class
AB4
AB3E
# of atoms
bonded to
central atom
# lone
pairs on
central atom
4
0
3
1
Arrangement of
electron pairs
Molecular
Geometry
tetrahedral
tetrahedral
tetrahedral
trigonal
pyramidal
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
tetrahedral
tetrahedral
AB3E
3
1
tetrahedral
trigonal
pyramidal
AB2E2
2
2
tetrahedral
bent
Class
Arrangement of
electron pairs
Molecular
Geometry
O
H
H
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB5
5
0
trigonal
bipyramidal
trigonal
bipyramidal
AB4E
4
1
trigonal
bipyramidal
distorted
tetrahedron
Class
Arrangement of
electron pairs
Molecular
Geometry
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB5
5
0
trigonal
bipyramidal
trigonal
bipyramidal
AB4E
4
1
trigonal
bipyramidal
distorted
tetrahedron
AB3E2
3
2
trigonal
bipyramidal
T-shaped
Class
Arrangement of
electron pairs
Molecular
Geometry
F
F
Cl
F
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB5
5
0
trigonal
bipyramidal
trigonal
bipyramidal
AB4E
4
1
trigonal
bipyramidal
distorted
tetrahedron
AB3E2
3
2
trigonal
bipyramidal
T-shaped
AB2E3
2
3
trigonal
bipyramidal
linear
Class
Arrangement of
electron pairs
Molecular
Geometry
I
I
I
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB6
6
0
octahedral
octahedral
AB5E
5
1
octahedral
square
pyramidal
Class
Arrangement of
electron pairs
Molecular
Geometry
F
F
F
Br
F
F
10.1
VSEPR
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB6
6
0
octahedral
octahedral
AB5E
5
1
octahedral
square
pyramidal
AB4E2
4
2
octahedral
Class
Arrangement of
electron pairs
Molecular
Geometry
square planar
F
F
Xe
F
F
10.1
10.1
Predicting Molecular Geometry
1.
Draw Lewis structure for molecule.
2.
Count number of lone pairs on the central atom and number of atoms
bonded to the central atom.
3.
Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
O
S
AB2E
F
O
F
S
AB4E
F
distorted
tetrahedron
bent
F
10.1
Pause for a Cause
Pg. 211 #48
48. Draw the Lewis Structure and
predict the molecular geometry
for each of the following:
a. SCl2
b. PI3
c. Cl2O
d. NH2Cl
e. SiCl3Br
f. ONCl
Dipole Moments and Polar Molecules
electron poor
region
electron rich
region
F
H
d+
d-
10.2
10.2
Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
10.2
a.
b.
c.
d.
e.
f.
g.
Polar or non-polar?
Let’s draw the Lewis
Structures
HF And find out!
H2O
NH3
CO2
SO3
H2S
CCl4
Valence Bond Theory and NH3
N – 1s22s22p3
3 H – 1s1
If the bonds form from overlap of 3 2p orbitals on nitrogen
with the 1s orbital on each hydrogen atom, what would
the molecular geometry of NH3 be?
If use the
3 2p orbitals
predict 900
Actual H-N-H
bond angle is
107.30
10.4
Hybridization – mixing of two or more atomic
orbitals to form a new set of hybrid orbitals.
1. Mix at least 2 nonequivalent atomic orbitals (e.g. s
and p). Hybrid orbitals have very different shape
from original atomic orbitals.
2. Number of hybrid orbitals is equal to number of
pure atomic orbitals used in the hybridization
process.
3. Covalent bonds are formed by:
a. Overlap of hybrid orbitals with atomic orbitals
b. Overlap of hybrid orbitals with other hybrid
orbitals
10.4
10.4
10.4
Predict correct
bond angle
10.4
Formation of sp Hybrid Orbitals
10.4
Formation of sp2 Hybrid Orbitals
10.4
How do I predict the hybridization of the central atom?
Count the number of lone pairs AND the number
of atoms bonded to the central atom
# of Lone Pairs
+
# of Bonded Atoms
Hybridization
Examples
2
sp
BeCl2
3
sp2
BF3
4
sp3
5
sp3d
PCl5
6
sp3d2
SF6
CH4, NH3, H2O
10.4
10.4
10.5
10.5
10.5
10.5
10.5
Sigma (s) and Pi Bonds (p)
1 sigma bond
Single bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid
(vinegar) molecule CH3COOH?
O
H
s bonds = 6
H
C
C
O
H
+1=7
p bonds = 1
H
10.5
Dipole Moments and Polar Molecules
electron poor
region
electron rich
region
H
F
d+
d-
m=Qxr
Q is the charge
r is the distance between charges
1 D = 3.36 x 10-30 C m
10.2
10.2
10.2
Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
10.2
Molecular Forces
• Intramolecular Forces
– The forces of attraction between atoms in
a molecule.
– Covalent, Ionic, Metallic bonds
• Intermolecular Forces
– The forces of attraction between
molecules.
– Dipole-dipole, Hydrogen bonding, and
London dispersion forces.
Dipole-Dipole Forcesforce of attraction found
when the positive end of
one polar molecule is
attracted to the negative
end of another polar
molecule.
Dipole- equal but opposite charge
separated by a short distance.
Polarity is dependent upon
differences in electronegativity and
the symmetry of the molecule
- NH3, HF, H2O, CH3Cl are all polar
- CO2, CCl4 are nonpolar
Hydrogen Bonding- a
particularly strong dipoledipole attractive force
formed when
hydrogen bonds with a
strongly electronegative
nonmetal (F, O, or N)
1. Polar Bonds cause
hydrogen to leave its
electrons almost exposed.
ie. water
.
A nonpolar molecule is transformed into a
dipole due to a distortion of its’ electron cloud
in response to an approaching dipole.
-Negative end of the dipole molecule
causes an imbalance of the electron
cloud in the nonpolar molecule and
therefore a bulging away.
* ie. water solutions of iodine. Water is
polar but iodine molecules are nonpolar
London Dispersion
Forces
Intermolecular
attractive force that
exist from the creation
of instananeous or
temporary dipoles .
Fritz London
1900-1954
-They are the only force
of attraction between
nonpolar molecules.
- They increase in
significance as
molecularsize increases.
London Forces in Hydrocarbons
E
X
A
P
L
E
1
E
X
2
E
X
3
Relative magnitudes of
forces
The types of bonding forces vary in
their strength as measured by average
bond energy.
Strongest
Ionic bonds
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
London forces (less than 1 kcal)
Weakest