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Chapter 8
Bonding: General Concepts
8.1
8.2
8.3
8.4
8.5
8.6
8.7
8.8
8.9
8.10
8.11
8.12
8.13
Types of Chemical Bonds
Electronegativity
Bond Polarity and Dipole Moments
Ions: Electron Configurations and Sizes
Energy Effects in Binary Ionic Compounds
Partial Ionic Character of Covalent Bonds
The Covalent Chemical Bond: A Model
Covalent Bond Energies and Chemical Reactions
The Localized Electron Bonding Model
Lewis Structures
Exceptions to the Octet Rule
Resonance
Molecular Structure: The VSEPR Model
A Chemical Bond
• Forces that hold groups of atoms together and make them
function as a unit.
• A bond will form if the energy of the aggregate is lower than
that of the separated atoms.
Types of Chemical Bonds:
• Ionic Bonding – electrons are transferred
• Covalent Bonding – electrons are shared equally
• Intermediate cases
Copyright © Cengage Learning. All
rights reserved
3
Ionic Compound
Cation is always smaller than atom from which
it is formed.
Anion is always larger than atom from which it
is formed.
Bond energy:
It is the energy required to break the bond.
Bond length:
It is the equilibrium distance where the energy of the system is
minimum
How does a bonding force develop between two identical
atoms ?
The Interaction of Two Hydrogen Atoms
When H atoms are brought
close together, there are two
unfavorable potential energy
terms, proton-proton repulsion
& electron-electron repulsion ,
and one favorable term, protonelectron attraction.
A bond will form (that is, the
two H atoms will exist as a
molecular unit)if the system
can lower its total energy in the
process .
Energy profile as a function of the distance between the
nuclei of the hydrogen atoms.
The zero point of energy is :” the
atoms at infinite separation .
At very short distances , when
the atoms are very close together
Bond length is the distance at
which the system has minimum
energy .
The potential energy of each
electron is lowered because of
the increased attractive forces in
this area
In the hydrogen molecule and in many other molecules in which
electrons are shared by nuclei is called
Covalent bonding.
Covalent Bond
 No electron transfer
Electrons are shared between two atoms,
positioned between the two nuclei
electrons are shared equally between identical
atoms
Example: H2, O2, etc.
Polar Covalent Bond
•
•
•
Unequal sharing of electrons between atoms in a
molecule.
Results in a charge separation in the bond (partial
positive and partial negative charge).
When a sample of hydrogen fluoride HF gas is
placed in an electric field , the molecules tend to
orient themselves, with the fluoride end closest to
the positive pole and the hydrogen end closest to
the negative pole.
+δ
-δ
Polar covalent bond or polar bond is a covalent bond with
greater electron density around one of the two atoms. The
molecule is called “Dipolar”.
H
F
e- poor e- rich
F
H
The Effect of an Electric Field on Hydrogen Fluoride
Molecules
(a) When no electric
field is present, the
molecules
are
randomly oriented.
(b) When the field is
turned on, the
molecules tend to line
up with their
negative ends toward
the positive pole and
their
positive ends toward
to determine how polar a bond will be, we introduce
Electronegativity
• Electronegativity
– the ability of an atom in a molecule to attract shared
electrons to itself
– I.e., …how much an atom “wants” electrons within a
bond
– determined by Linus Pauling (1901 - 1995)
•The general trend is electronegativity increase as we go right
and up in the periodic table
The Relationship Between Electronegativity and Bond Type
•Electronegativity difference increase  polarity increases
Example 8.1:
:
• Arrange
(Order)the following bonds in order of increasing polarity:
•
– H-H, O-H, Cl-H, S-H, and F-H
– Solution:
– The polarity of the bond increases as the difference in the electronegativity
increases.
H-H < S-H < Cl-H < O-H < F-H
(2.1) (2.1) (2.5) (2.1) (3.0) (2.1) (3.5) (2.1)
–0
0.4
Electronegativity
difference
Covalent bond
0.9
(4.0) (2.1)
1.4
Polarity increases
1.9
polar covalent bond
Bond Polarity and Dipole Moments
A molecule with a center of negative charge and a
center of positive charge is said to be dipolar or has a
dipole moment.
For example hydrogen fluoride
HF behaves in electric field as if it had two centers of charge, H
positive and F negative.
• Place a polar molecule in an electric field
– the molecule will line up so that its “negative” end will
line up with the positive pole and the “positive end” will
line up
with the negative pole
• Dipole
Moment
– a polar molecule has a dipole moment
– a polar molecule has a center of positive charge and a center of negative
charge
– Ex: H-F
– The dipolar character of a molecule is represented by an arrow pointing to the negative
charge center with the tail of the arrow indicating the positive center of charge.
Representation of dipole moment
•Molecules with polar bonds and have net dipole moment:
Water molecule is polar molecule
The water molecule has a dipole moment .
The same type of behavior is observed for the NH3
The structure and charge
distribution of the ammonia
molecule
 The dipole moment of the ammonia
molecule oriented in an electric
field
 Polar molecule
 It has dipole moment
• Some molecules have polar bonds, but are nonpolar
– these molecules have no net dipole moment
– due to the overall geometry of the molecule, the bond
polarities cancel out, so the molecule has no net dipole
moment
• Ex: A- Linear molecule: e.g: CO2 (O=C=O)
c. Tetrahedral molecules:
e.g: CCl4 , and CH4
Tetrahedral molecules with four identical bonds 109.5 degrees apart
Example 8.2:
For each of the following molecules, show the direction of the bond
polarities and indicate which ones have a dipole moment:
•HCl
•Cl2
•CH4 ( tetrahedral molecule)
•H2S (V-shaped molecule)
Answer:
•HCl
The electronegativity of chlorine is greater than that of
hydrogen (3.02.1). Thus the chlorine will be partially negative,
and the hydrogen will be partially positive. The HCl molecule
has a dipole moment.
•Cl2
The two chlorine atoms share the electrons equally. No bond polarity
occurs, and the Cl2 molecule has no dipole moment
•CH4 ( tetrahedral molecule)
Carbon has a slightly higher electronegativity
(2.5) than does hydrogen (2.1). this lead to
small partial positive charges on the hydrogen
atoms and small partial negative charge on the
carbon. Bond polarities cancel. The molecule
has no dipole moment.
•H2S (V-shaped molecule)
The H2S molecule has a dipole
moment.
Ions: Electron Configurations and sizes:
•Atoms in stable compounds usually have a noble gas electron
configuration.
• In ionic compounds ;The nonmetals form anions, and the
metals form cations to achieve a noble gas electron
configuration ( Na+ - Cl-)
•When two nonmetals react to form a covalent bond, they share
electrons to complete the valence electron configuration of
both atoms. That is, both nonmetals have noble gas electron
configuration.
•When metal react with nonmetal to form ionic compound,
the valence electron configuration of the nonmetal achieves
the electron configuration of the next noble gas atom and
the valence orbital of the metal are emptied.
In this way both ions achieve noble gas electron
configuration.
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Ca [Ar]4s2
Al [Ne]3s23p1
Na+ [Ne]
Ca2+ [Ar]
Al3+ [Ne]
Atoms gain electrons so
that anion has a noble-gas
outer electron
configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
Predicting formulas of ionic compounds
When we speak in this text of the stability of an ionic
compound, we are referring to the solid state.
• Group I forms +1 ions
• Group II forms +2 ions
• Group III forms +3 ions
• Group VI forms -2 ions
• Group VII forms -1 ions
Example of Ionic Compounds
 MgO magnesium oxide is formed of Mg2+ and O2Mg :[Ne]4s2
Mg2+ :[Ne]
O: 1s22s22p4
 CaO
O2- 1s22s22p6 or [Ne]
formed from Ca2+ and O2-.
Ca :[Ar]4s2
Ca2+ :[Ar]
 Al2O3 is formed of 2Al3+ and 3O2-.
Al:[Ne]3s23p1 Al3+ :[Ne]
• Size of Ions
– Positive ion (Cation)
• formed by the loss of an electron from the parent atom
• cation is smaller than the parent due to the loss of electron pair
repulsion and the increased nuclear charge felt by each electron
– Negative ion (Anion)
• formed by the parent atom gaining an electron
• anion is larger than the parent atom due to increased electron
pair repulsion
Isoelectronic ions
Ions containing the same number of electrons.
The number of electrons and the number or protons are affect on
the size of ions. But in case of isoelectronic ions, ions have the
same number of electrons.
ions with the same number of electrons
Na+ , Mg+2, Al+3, F-, O-2, N-3
all contain 10 electrons
all have the same electron configuration as Ne
in terms of size, N-3>O-2>F->Na+>Mg+2>Al+3
the ion with the greater number of protons in an
isoelectronic series will be the smallest due to the greater
nuclear charge pulling the electrons in closer
8
10 electrons
1s2 2s2 2p6
F
9
10 electrons
1s2 2s2 2p6
+
Na
11
10 electrons
1s2 2s2 2p6
+2
Mg
12
10 electrons
1s2 2s2 2p6
+3
Al
13
2 2s
2 2p6 the 10 electrons and the positive charge
The
attraction1s
force
between
10
electrons
on the nucleus increase with increase the nuclear charge ‘Z’ increases.
Therefore the size of ions decreases as the nuclear charge ‘Z’
increases. Also we can see of isoelectronic ions decreases with
increasing atomic number
O-2> F- > Na+ > Mg+2 > Al+3 size of isoelectronic ions
• Example :Arrange the ions Se-2, Br-, Rb+, and Sr+2 in order of
decreasing size
• Answer: This is an isoelectric series of ions with the krypton electron
configuration. Since these ions all have the same number of electrons,
• Z (ion)
36 36 36 36 (no. of electrons)
• Their sizes will depend on the nuclear charge .
• Z (atom)
34 (Se2-) 35 (Br-) 37(Rb+) 38 (Sr2+) (no. of protons)
Since the nuclear charge is greatest for (Sr2+) , it is the smallest of these
ions.
• The (Se2-) ion is largest. (Z increases and size decreases)
•
•
Ion Se2- > Br- > Rb+ > Sr2+
Example 8.4
Choose the largest ion in each of the following groups:
•Li+, Na+, K+, Rb+, Cs+
(Group 1A)
2+, Cs+, I-, Te2•Ba
(isoelectronic series)
Answer:
•Since size increases down a group Cs+ is largest.
•This is an isoelectronic series of ions with xenon electron
configuration (size decreases with increasing Z):
Ion
Te2- > I- > Cs+ > Ba2+
Z (atom)
Z (ion)
52
54
53
54
55
54
56 (no. of protons)
5 (no. of electrons)
The Covalent Chemical Bond: A model:
The Localized Electron Bonding Model
• The model assumes that A molecule is composed of atoms that are
bound together by sharing pairs of electrons using the atomic orbitals of
the bound atoms.
• electron pairs are localized on a particular atom or in the space between
two atoms.
Localized Electron Bonding Model
• Lone pairs
– electrons localized on a particular atom
• Bonding pairs (or shared pairs)
– electrons localized in the space between two atoms
• On to Lewis structures!
Lewis Structure
Lewis structure of a molecule Shows how
valence electrons are arranged among atoms in a
molecule.
 Octet rule, atoms tend to combine in such a way
that they each have eight electrons in their valence
shells, giving them the same electronic
configuration as a noble gas.
 The
Hydrogen forms stable molecules where it shares 2 electrons.
When 2 hydrogen atoms ,each with 1 electron , combine to form
the H2 molecule .
H.
.H
H:H
By sharing electrons , each H in H2 , in effect, has 2 electrons
,that is , each H has a filled valence shell .
• In covalent bond formation, atoms go as far as possible
toward completing their octets (duplets) by sharing
electron pairs.
• Each fluorine atom also has three pairs of electrons not
involved in bonding. These are the lone pairs .
H+H  H:H or H-H
..
:
..
F F
.
..
:
.
..
:
..
..
F F
:
..
..
..
:
or
:
..
F F
..
..
:
Steps for Writing Lewis Structures
1. Sum the valence electrons from all the atoms.
2. Use a pair of electrons to form a bond between each pair of bound
atoms.
3. Atoms usually have noble gas configurations. Arrange the remaining
electrons to satisfy the octet rule (or duet rule for hydrogen).
•The central atom is usually written first in the formula [the
central atom is usually the least electronegative atom,
exceptions in H2O, and NH3 where O and N are the central
atoms]
•Hydrogen is never the central atom
Examples:
Lewis structure of water
1- H2O (water): 1H and 8O
..
H  O H
Sum of valence electrons = 1 + 1 + 6 = 8
single covalent bonds
..
H
+ O
+ H
2- CO2 (carbon dioxide):
H O H
8e
2e
2e6C
or
H
O
H
and 8O
Sum of valence electrons = 4 + 6 + 6 = 16
Double bond – two atoms share two pairs of electrons
double bonds
or
O
O
C
O C O
double bonds
8e- 8e- 8e-
3- CN- (Cyanide ion): 6C and 7N
Sum of valence electrons = 4 + 5 + 1 = 10
Triple bond – two atoms share three pairs of electrons
N
c
or
8e- 8etriple bond
4. NO+ (Nitrogen oxide ion) 7N and 8O
Sum of valence electrons = 5 + 6 – 1 = 10
C
N
triple bond
[NO]+
5. N2 (Nitrogen)
7N and 7N
Sum of valence electrons = 5 + 5 = 10
NN
6. CO2
..
O
..
C
..
O
..
Exceptions to the Octet Rule
The Incomplete Octet (Be &B)
Some atoms are satisfied with less than an octet
-Be (1s22s2) is stable with only four valence electrons
- Boron (1s2 2s2 2p1) also tends to form compounds
with less than eight electrons
Odd-Electron Molecules:
- Some molecules have an odd number of electrons
can't satisfy octet rule; usually N has the odd
number
The Expanded Octet:
– Atoms in and beyond the 3rd period can
have more than eight electrons when in a
compound
– Thus, when drawing Lewis electron-dot
formulas, extra electrons go on the central
atom.
- : 3(7) + 1 = 22 valence electrons
I
3
Example:
SF6 [6 + 6 (7) = 48 valence electrons ]
[
..
:I
..
PCl5 : 5 + 5(7) = 40 electrons
..
: Cl
..
..
I:
..
....
I
..
..
: Cl :
P
: Cl :
: Cl:
..
..
..
:
Cl
..
]
-
Resonance: Blending of Structures
Consider the Lewis structure for the nitrate ion NO-3:
: O:
[
.
.. O
. ..
N
...
.O
..
]
-
Based on Lewis structure it should be two types of N…O
bonds
.
Experiments show that NO-3 has only one type of
N…O bond with length and strength between those
expected for a single and double bonds.
.
This can be explained on the basis of the resonance
structures of NO-3:
All valid. We cannot find two bond
lengths (hypothetical N-O vs N=O
NO3-
..
:O:
[ ..
..O..
N
-
]
...
.O
..
..
:O:
[.
.O.
.
N
-
...
.O
..
]
:O:
[.
...O..
N
-
..
.O
.
]
Physical evidence shows that NO3- has three equivalent bonds
The correct description of NO3- is not one of the three Lewis structures,
but an “average” of the three Lewis structures
resonance structure: one of two or more Lewis structures
representing a single molecule that cannot be described fully with
only one Lewis structure
Example: Given the Lewis structure for ozone, O3
Formal Charge
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.
1. For neutral molecules, sum of formal charges must equal zero.
2. For ions, the sum of the formal charges must equal charge.
Consider the Lewis structure for POCl3. Assign the formal
charge for each atom in the molecule.
Cl
P: 5 – (4+0) = +1
O: 6 – (6+1) = –1
Cl: 7 – (6+1) = 0
Cl
P
Cl
O
VSEPR Model
(Valence-Shell Electron-Pair Repulsion model)
This model is useful in predicting the geometries of molecules formed
from nonmetals.
The structures of molecules play a very important role in determining their
chemical properties .
The main postulate of this model is that “ The structure around a given
atom is determined principally by minimizing electron-pair
repulsions.”
The bonding and nonbonding pairs around a given atom will be positioned
as far apart as possible.
Number of electron pairs
Structure
Example
o
180
2
Linear
3
Trigonal planar
Cl
Cl
Be
F
o
120
B
F
F
4
H
Tetrahedral
o
109.5
C
H
H
H
5
Trigonal bipyrimidal
Cl
o
90
Cl
o
120
Cl
P
Cl
Cl
6
Octahedral
F 90
F
o
F
S
F
F
F
Predicting a VSEPR Structure
• 1. Draw Lewis structure for the molecule.
• 2. Put the electron pairs as far apart as possible.
• 3. Determine positions of atoms fro the way electron
pairs are shared.
• 4. Determine the name of molecular structure from
positions of the atoms.
The Bond Angles In the CH4, NH3, and H2O
Molecules
bonding-pair vs. bonding
pair repulsion
<
lone-pair vs. bonding
pair repulsion
<
lone-pair vs. lone pair
repulsion
Example:
Predict the molecular structure and bond angles for the
following molecule or ion:
1.HCN
2.PH3
3.CHCl3
4.NH4+
5.H2CO2
6.CO2
Answer:
a) HCN
•Draw the Lewis structure
H-CN:
2- count the electron pairs
•2 electron pairs as single bonds
•2 effective electrons as triple bond
3-determine the position of atoms
H-CN:
4- determine the name of molecular structure
•Linear , 180o
b) PH3 trigonal pyramid  109.5o
C)CHCl3 tetrahedral, 109.5o
d) NH4+ tetrahedral, 109.5o
e) H2CO trigonal planar, 120 o
f)CO2 linear, 180 o