Chapter 11
Chemical Bonding
Forces that hold atoms together
The Nature of Bonding
• There are several major types of bonds.
Ionic, covalent and metallic bonds are the
three most common types of bonds.
• Covalent bonds – electrons are shared
between atoms.
• Ionic bonds – electrons are transferred
between atoms, creating cations and anions.
• Metallic bonds – two or more metals bonded
together.
The Nature of Covalent Bonding
• There are two different types of covalent
bonds, polar covalent and nonpolar
covalent.
– polar covalent – electrons are not shared
equally between the two bonded atoms. The
electrons are pulled toward the more
electronegative of the elements.
– nonpolar covalent – electrons are shared
equally between the two bonded atoms.
Electronegativities
9_12
IA
IIA
Li
1.0
Be
1.5
Na
0.9
Mg
1.2
K
0.8
H
2.1
VIIIB
IIIB
IVB
VB
VIB
VIIB
Ca
1.0
Sc
1.3
Ti
1.5
V
1.6
Cr
1.6
Mn
1.5
Fe
1.8
Co
1.8
Rb
0.8
Sr
1.0
Y
1.2
Zr
1.4
Nb
1.6
Mo
1.8
Tc
1.9
Ru
2.2
Cs
0.7
Ba
0.9
La–Lu
1.1–1.2
Hf
1.3
Ta
1.5
W
1.7
Re
1.9
Os
2.2
Fr
0.7
Ra
0.9
Ac–No
1.1–1.7
IIIA
IVA
VA
VIA
VIIA
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
IB
IIB
Ni
1.8
Cu
1.9
Zn
1.6
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Rh
2.2
Pd
2.2
Ag
1.9
Cd
1.7
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Ir
2.2
Pt
2.2
Au
2.4
Hg
1.9
Tl
1.8
Pb
1.8
Bi
1.9
Po
2.0
At
2.2
The formation of a bond
between two hydrogen atoms.
Source: Andrey K. Geim/High Field Magnet Laboratory/University of Nijmegen
Probability representations of the electron sharing
in HF. (a) What the probability map would look like
if the two electrons in the H–F bond were shared
equally. (b) The actual situation, where the shared
pair spends more time close to the fluorine atom
than to the hydrogen atom.
The Nature of Covalent Bonding
• Ionic bonds are formed when there is an
electronegativity difference (DEN) greater
than 2.0.
• Polar covalent bonds form when there is a
DEN between 0.5 and 1.7.
• Nonpolar covalent bonds form when there
is a DEN between 0 and 0.49.
The Nature of Covalent Bonding
• If the DEN is between 1.7 and 2.0, an ionic
bond will form if a metal is one of the
elements, and a polar covalent bond will
form if only nonmetals or metalloids are
present.
The Nature of Covalent Bonding
• What type of bond is formed between the
following elements?
• N and O
K and F
• Mg and Cl
P and F
• C and H
Bond Polarity
• Covalent bonding between unlike atoms
results in unequal sharing of the
electrons
– One end of the bond has larger electron
density than the other
•• F 
H

• The result is bond polarity
– The end with the larger electron density
gets a partial negative charge
– The end that is electron deficient gets a
partial positive charge
The three possible types of bonds:
(a) a covalent bond formed between identical atoms;
(b) a polar covalent bond, with both ionic and
covalent components; and
(c) an ionic bond, with no electron sharing.
Dipole Moment
• Bond polarity results in an unequal electron
distribution, resulting in areas of partial
positive and partial negative charge
• Any molecule that has a center of positive
charge and a center of negative charge in
different points is said to have a dipole
moment (two different poles of charge).
Dipole Moment
• If a molecule has more than one polar
covalent bond, the areas of partial negative
and positive charge for each bond will
partially add to or cancel out each other
• The end result will be a molecule with one
center of positive charge and one center of
negative charge
• The dipole moment effects the attractive
forces between molecules and therefore the
physical properties of the substance
(a) The charge
distribution in
the water
molecule.
(b) The water
molecule
behaves as if it
had a positive
end and a
negative end, as
indicated by the
arrow.
(a) Polar water molecules are strongly attracted to
positive ions by their negative ends. (b) They are
also strongly attracted to negative ions by their
positive ends.
Polar water molecules are strongly
attracted to each other.
Electron Configuration in Ionic Bonding
• Metals tend to lose their valence electrons,
leaving a complete octet in their nextlowest energy level.
• Sodium – (1 valence electron) loses 1
electron and becomes Na+1.
• Na ([Ne]3s1)  1e- + Na+1([Ne])
• Calcium – (2 valence electrons) loses 2
electrons and becomes Ca+2.
• Ca ([Ar]4s2)  2e- + Ca+2([Ar])
Electron Configuration in Ionic Bonding
• Nonmetals tend to gain or share valence
electrons to complete an octet in their
highest energy level.
• Oxygen – (6 valence electrons) gains two
electrons to become O-2 .
• O ([He]2s22p4) + 2e-  O-2 ([He] 2s22p6)
• Phosphorus – (5 valence electrons) gains
three electrons to become P-3.
• P ([Ne]3s23p3) + 3e-  P-3 ([Ne] 3s23p6)
Formation and Properties of Ionic
Compounds
• Ionic bonds – forces of attraction that bind
cations and anions together.
• Ionic compound – consists of electrically
neutral group of ions joined by
electrostatic forces.
• Example: Sodium chloride
Formation and Properties of Ionic
Compounds
• At room temperature, most ionic
compounds are crystalline solids, where
ions are arranged in various 3-D patterns.
• Because of the large attractive forces of
the ions to each other the compounds
become very stable and have high melting
points.
The shape of snowflakes results
from bonding.
Source: ©Clyde
H. Smith/Peter
Arnold, Inc.
Sodium
Chloride
Crystals
The structure of lithium fluoride.
Electron Configuration in Ionic Bonding
• Scientists have learned that all of the elements
within each group behave similarly because they
have the same number of valence electrons.
• Valence electrons - # of electrons in the highest
occupied energy level of an atom.
• The number of valence electrons is related to
the group numbers on the periodic table.
Electron Configuration in Ionic Bonding
•
•
•
•
•
•
•
•
•
Group 1 elements = 1 valence electron.
Group 2 elements = 2 valence electrons.
Groups 3-12 elements = 2 valence electrons.
Group 13 elements = 3 valence electrons.
Group 14 elements = 4 valence electrons.
Group 15 elements = 5 valence electrons.
Group 16 elements = 6 valence electrons.
Group 17 elements = 7 valence electrons.
Group 18 elements = 8 valence electrons.
Determining Valence Electrons for
an Ion or a Compound
• 1. Multiply the number of valence electrons by
the number of moles of each element.
• 2. Add up all the electrons for each of the
elements.
• 3. If there is a charge and it is negative, add
that number of electrons to the total.
• 4. If there is a charge and it is positive, subtract
that number of electrons from the total.
• Total # of electrons should always be an even
number!
Determining Valence Electrons Examples
• Determine the number of valence electrons in
each of the following compounds and ions:
• NH4+1
• CH2ClBr
• PO4-3
Electron Configuration in Ionic Bonding
• Valence electrons are the only electrons
involved in bonding, and are the only ones
written when drawing electron dot structures.
• In forming compounds, atoms tend to achieve
the electron configuration of a noble gas, having
8 valence electrons which as known as having a
stable octet (octet for 8 valence electrons).
Lewis Symbols of Atoms and Ions
• Also known as electron dot symbols
• Use symbol of element to represent nucleus and inner
electrons
• Use dots around the symbol to represent valence electrons
– put one electron on each side first, then pair
• Elements in the same group have the same Lewis symbol
– Because they have the same number of valence
electrons
• Cations have Lewis symbols without valence electrons
• Anions have Lewis symbols with 8 valence electrons
••
••
•
••
••
Li•
Be•
•
Li•
•B•
•
•C•
Li+1
•
•N•
•
••
:F:
•
•O:
•
:F:
•• -1
[:F:]
••
•
:Ne:
••
The Nature of Covalent Bonding
• Structural formula – chemical formulas
that show the arrangement of atoms in
molecules and polyatomic ions.
• Octet rule – atoms gain or lose electrons
to acquire the stable electron configuration
of a noble gas, usually having 8 valence
electrons.
Lewis Structures
• You can represent the formation of the
covalent bond in H2 as follows:
H
. + .H
:
H H
– This uses the Lewis dot symbols for the hydrogen
atom and represents the covalent bond by a pair
of dots.
Lewis Structures
• The shared electrons in H2 spend part of
the time in the region around each atom.
:
H H
– In this sense, each atom in H2 has a helium
configuration.
Lewis Structures
• The formation of a bond between H and Cl
to give an HCl molecule can be
represented in a similar way.
: :
. + .Cl:
: :
H
: :
H Cl
– Thus, hydrogen has two valence electrons about
it (as in He) and Cl has eight valence electrons
about it (as in Ar).
Lewis Structures
• Formulas such as these are referred to as
Lewis electron-dot formulas or Lewis
structures.
bonding pair
: :
: :
H Cl
lone pair
– An electron pair is either a bonding pair (shared
between two atoms) or a lone pair (an electron
pair that is not shared).
The Nature of Covalent Bonding
• Exceptions to the octet rule:
– H needs 2 electrons to be stable
– Be needs 4 electrons to be stable
– B needs 6 electrons to be stable
The Nature of Covalent Bonding
• Steps for Drawing Lewis-dot structures
1. Determine the number of valence
electrons in the molecule.
- When drawing determining valence electrons
for an ion, add electrons if it an anion, and
subtract electrons if it is a cation.
2. The first element in the compound will be
the central atom. Exception: hydrogen
will never be the central atom.
The Nature of Covalent Bonding
Steps for Drawing Lewis-dot Structures
3. Use one pair of electrons to bond each
outer or terminal atom to the central
atom.
4. Make all outer or terminal atoms stable
using the valence electrons.
5. Put any remaining electrons around the
central atom as lone pairs.
The Nature of Covalent Bonding
• Draw the Lewis structure for:
• NH3
• PO43• CHFClBr
• PF5-2
The Nature of Covalent Bonding
• Single covalent bond – a bond in which
two atoms share a pair of electrons.
• Double covalent bond – a bond in which
two atoms share two pairs of electrons.
• Triple covalent bond – a bond in which two
atoms share three pairs of electrons.
The Nature of Covalent Bonding
• If you have used up all of the valence electrons
and you still need two more electrons to make
the central atom stable, you must have one
double bond.
• If you still need four more electrons to make the
central atom stable, you must have either one
triple bond or two double bonds.
• Double and triple bonds exist most commonly
between C, N, O, and S atoms.
The Nature of Covalent Bonding
• Draw Lewis structures for:
• NOCl
• CO2
• N2
• SiO3-2
The Nature of Covalent Bonding
• Resonance structures – molecules or ions
that can have two or more different Lewis
structures. They must contain a double
bond to have any resonance structures.
• Resonance structures don’t truly have a
single bonds or a double bond, but a
hybrid mixture of bonds where the extra
bond is spread equally among the other
single bonds.
Resonance Structures
• Draw Lewis structures for:
• NOCl
• CO2
• N2
• SiO3-2
The Nature of Covalent Bonding
• Diamagnetic – substance where all of the
electrons of the central atom are paired or
bonded with other atoms.
• Paramagnetic – substance where all of the
electrons of the central atom are not
paired or bonded with other atoms.
The Nature of Covalent Bonding
• Single bonds are longer (length between
the atoms) than double and triple bonds.
• Double bonds are longer than triple bonds.
• Single bonds are not as strong as double
bonds, and can be broken much easier
than double bonds.
• Triple bonds are stronger than double
bonds.
Bonding Theory
• The valence-shell electron pair repulsion
(VSEPR) model predicts the shapes of
molecules and ions by assuming that the
valence shell electron pairs are arranged as far
from one another as possible.
– To predict the relative positions of atoms around a
given atom using the VSEPR model, you first note
the arrangement of the electron pairs around that
central atom.
Predicting Molecular Geometry
• The following rules and figures will help
discern electron pair arrangements.
1. Draw the Lewis structure
2. Determine how many bonding pairs are around
the central atom. Count a multiple bond as one
pair.
3. Determine how many lone pairs, if any, are around
the central atom.
All diatomic molecules have a linear shape.
Arrangement of Electron
Pairs About an Atom
2 pairs
Linear
3 pairs
Trigonal planar
5 pairs
Trigonal bipyramidal
4 pairs
Tetrahedral
6 pairs
Octahedral
Molecular Geometry Examples
• NH3
• NOCl
• PO43-
• CO2
• CHFClBr
• SF2
• PF5
• N2
• SeF6
• SiO3-2
Polar Bonds and Molecules
• Nonpolar covalent bond – equal sharing of
electrons between two atoms.
• Polar covalent bond – unequal sharing of
electrons between two atoms.
• In polar covalent bonds the electrons are
pulled closer to the atom with the larger
electronegativity value.
• This creates a partial positive and a partial
negative pole within the bond.
Polar Bonds and Molecules
• Polar bonds can create polar or nonpolar
molecules and ions.
• If the centers of partial positive and partial
negative charge are in the same location,
the molecule or ion is nonpolar.
• If the centers of partial positive and partial
negative charge are in different locations,
the molecule or ion is polar.
Polar Bonds and Molecules
• The easiest way to determine if a molecule
or ion is polar or nonpolar is to look at the
central atom.
• If the central atom has lone pairs of
electrons, the molecule or ion is polar.
• If the central atom does not have any lone
pairs of electrons, the molecule or ion is
nonpolar.
Examples: Polar or Nonpolar?
• Determine whether
each of the following
molecules or ions are
polar or nonpolar:
• NO2-1
• N2
• CN-1
• CH4
• SO3-2
Polar Bonds and Molecules
Attractions Between Molecules
• Molecules are attracted to one another by
a variety of forces.
• These intermolecular forces are weaker
than ionic or covalent bonds.
• These forces are responsible for whether
or not a molecular compound is a solid,
liquid, or a gas.
Polar Bonds and Molecules
• van der Waals forces – consist of
dispersion forces and dipole interactions
(dipole-dipole moments).
• Dispersion forces – weakest of all
intermolecular forces. They are caused by
the motion of electrons. The strength of
dispersion forces increases with the
increasing number of electrons in a
molecule.
Polar Bonds and Molecules
• All molecules contain dispersion forces.
• As molar mass and the number of
electrons increase, dispersion forces
increase.
• Halogens are the most common molecules
to have dispersion forces. Fluorine is a
gas, Bromine is a liquid and Iodine is a
solid.
Polar Bonds and Molecules
• Dipole interactions – occur when polar
molecules or ions are attracted to one
another. This occurs when a partial
positive charge and a partial negative
charge come close to each other.
• Dipole interactions are very similar to, but
much weaker than ionic bonds.
Polar Bonds and Molecules
• Hydrogen bonds – force exerted between
a hydrogen atom bonded to an F, O, or N
atom in one molecule and an unshared
pair on another F, O, or N atom in a
nearby molecule.
• Hydrogen bonds can have a great effect
on the boiling point of a substance.
Intermolecular Forces Examples