Chemical Bonds
The Formation of
Compounds From Atoms
Chapter 11
Hein and Arena
Version 2.0
12th Edition
Eugene Passer
Chemistry Department
Bronx Community
1 College
© John Wiley and Sons, Inc
Chapter Outline
11.1 Periodic Trends in
Atomic Properties
11.2 Lewis Structures of Atoms
11.6 Electronegativity
11.7 Lewis Structures of
Compounds
11.3 The Ionic Bond: Transfer of 11.8 Complex Lewis Structures
Electrons From One Atom
11.9 Compounds Containing
to Another
Polyatomic Ions
11.4 Predicting Formulas of
11.10 Molecular Shape
Ionic Compounds
11.5 The Covalent Bond:
Sharing Electrons
11.11 The Valence Shell
Electron Pair (VSEPR)
2
Model
11.1
Periodic Trends in
Atomic Properties
3
Characteristic properties and trends of the
elements are the basis of the periodic
table’s design.
4
These trends allow us to use the periodic
table to accurately predict properties and
reactions of a wide variety of substances.
5
Metals and Nonmetals
6
Chemical Properties
of Metals
• metals tend to lose
electrons and form
positive ions.
Chemical Properties
of Nonmetals
• nonmetals tend to
gain electrons and
form negative ions.
When metals react with nonmetals, electrons
are usually transferred from the metal to the
nonmetal.
7
Physical Properties Physical Properties
of Metals
of Nonmetals
• lustrous
• malleable
• good conductors of
heat
• good conductors of
electricity
• nonlustrous
• brittle
• poor conductors of
heat
• poor conductors of
electricity
8
Metalloids have properties that
are intermediate between metals
and nonmetals.
9
The Metalloids
1. boron
2. silicon
3. germanium
4. arsenic
5. antimony
6. tellurium
7. polonium
10
Nonmetals
arefound
foundtotothe
theleft
right
metalloids.
Metals are
ofofthethemetalloids
11
11.1
Atomic Radius
12
Atomic radii
increase down a
group.
For each step down a group, electrons enter
the next higher energy level.
13
11.2
Radii of atoms tend to decrease
from left to right across a period.
Eachincrease
For
This
time an in
representative
positive
electron
nuclear
is added, a
elements
charge
proton
is
pulls
also
within
all
the same
electrons
added
to the
period,
closer to
the energy
nucleus.
nucleus.level
remains constant
as electrons are
added.
14
11.2
Ionization Energy
15
The ionization energy of an atom is the
energy required to remove an electron from
an atom.
Na + ionization energy → Na+ + e-
16
• The first ionization energy is the
amount of energy required to remove
the first electron from an atom.
He + first ionization energy → He+ + eHe + 2,372 kJ/mol → He+ + e-
• The second ionization energy is the
amount of energy required to remove
the second electron from an atom.
He+ + second ionization energy → He2+ + eHe+ + 5,247 kJ/mol → He2+ + e-
17
• The first ionization energy is the
amount of energy required to remove
the first electron from an atom.
He + first ionization energy → He+ + eHe + 2,372 kJ/mol → He+ + e-
• The second ionization energy is the
amount of energy required to remove
the second electron from an atom.
He+ + second ionization energy → He++ + eHe+ + 5,247 kJ/mol → He++ + e-
18
As each succeeding electron is removed from
an atom, ever higher energies are required.
19
Ionization energies gradually increase from left to
right across a period.
Noble
Gases
1
2
VIIA
VA
IA
IVA
IIA
11.3
VIA
3
4
IIIA
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
20
Gases
nonmetals have higher ionization
potentials than metals
VIIA
VA
IA
VIA
IVA
IIA
IIIA
Distance of Outer Shell Electrons From Nucleus
Ionization energies of Group A elements decrease
from top to bottom in a group.
Noble
nonmetals
metals
11.3
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
21
11.2
Lewis Structures of Atoms
22
Metals form cations and nonmetals form
anions to attain a stable valence electron
structure.
23
Thesestable
This
rearrangements
structure often
occur
consists
by losing,
of twogaining,
s and
sixsharing
or
p electrons.
electrons.
24
The Lewis structure of an atom is a
representation that shows the valence
electrons for that atom.
• Na with the electron structure 1s22s22p63s1
has 1 valence electron.
• Fluorine with the electron structure 1s22s22p5
has 7 valence electrons
25
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.
Paired
electrons
B
Unpaired
electron
Symbol of
the element
2
1
2s 2p
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
27
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.
S
2
4
2s 2p
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
28
Lewis Structures of the first 20 elements.
29
11.4
11.3
The Ionic Bond
Transfer of Electrons From
One Atom to Another
30
The chemistry of many elements,
especially the representative ones, is to
attain the same outer electron structure
as one of the noble gases.
31
With the exception of helium, this
structure consists of eight electrons in
the outermost energy level.
32
After sodium loses its 3s electron, it has attained the
same electronic structure as neon.
33
After chlorine gains a 3p electron, it has attained the
same electronic structure as argon.
34
Formation of NaCl
35
The
3s electron
of sodium
transfersion
to (Cl
the-) 3p
of
A
sodium
ion (Na+)
and a chloride
areorbital
formed.
chlorine.
The force holding Na+ and Cl- together is an ionic bond.
Lewis representation of sodium chloride formation.
36
Formation of MgCl2
37
2+
2+) and two
Two
A
The
magnesium
3s
forces
electrons
holding
ionof(Mg
Mg
magnesium
two
transfer
chloride
Cl- together
toions
the are
half-filled
(Cl-ionic
) are
3p
formed.
bonds.
orbitals of two chlorine atoms.
38
In NaCl
the crystal
is made
each
upsodium
of cubicion
crystals.
is surrounded by six
chloride ions.
39
In the crystal each chloride ion is surrounded by six
sodium ions.
40
11.5
The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl
41
11.5
Relative Size of
Sodium Ion to Chloride Ion
42
A sodium ion is smaller than a sodium atom
because:
(1) the sodium atom has lost its outermost
electron.
(2) The 10 remaining electrons are now
attracted by 11 protons and are drawn closer
to the nucleus.
43
11.6
A chloride ion is larger than a chlorine atom
because:
(1) the chlorine atom has gained an electron
and now has 18 electrons and 17 protons.
(2) The nuclear attraction on each electron has
decreased, allowing the chlorine to expand.
44
11.6
• Metals usually have one, two, or three
electrons in their outer shells.
• When a metal reacts it:
– usually loses one, two, or three electrons
– attains the electron structure of a noble
gas
– becomes a positive ion.
• The positive ion formed by the loss of
electrons is much smaller than the
metal atom.
45
• Nonmetals usually have one, two or
three electrons in their outer shells.
• When a nonmetal reacts it:
– usually gains one, two, or three electrons
– attains the electron structure of a noble
gas
– becomes a negative ion.
• The negative ion formed by the gain of
electrons is much larger than the
nonmetal atom.
46
47
11.4
Predicting Formulas of
Ionic Compounds
48
In almost all stable chemical compounds of
representative elements, each atom attains a
noble gas electron configuration.
49
• Metals will lose electrons to attain a
noble gas configuration.
• Nonmetals will gain electrons to attain a
noble gas configuration.
Barium and Sulfur Combine.
– barium loses two electrons to sulfur and
attains a xenon configuration.
– sulfur gains two electrons from barium and
attains an argon configuration.
Ba →[Xe]6s
[Xe] +2 2e-
S +[Ne]3s
2e- →2[Ar]
3p4
Ba + S → BaS
50
Because of similar electron structures,
the elements of a family generally form
compounds with the same atomic ratios.
51
52
The elements
group numbers
of a family
for thehave
representative
the same
elements are
outermost
electron
equalconfiguration
to the total except
numberthat
of
outermost
the
electrons
electrons
are in different
in the atoms
energy
of the
levels.
group.
53
10.17
• The atomic ratio of the alkali metal
sodium to chlorine is 1:1 in NaCl.
• The atomic ratios of the other alkali
metal chlorides can be predicted to also
be 1:1.
• LiCl, KCl, CsCl, FrCl
54
• The atomic ratio of hydrogen to
nitrogen is 3:1 in ammonia (NH3).
Nitrogen is the first member of group
5A.
• The atomic ratio of hydrogen when
combined with other group 5A
elements can be predicted to also be
3:1.
• PH3, AsH3, SbH3, BiH3
55
11.5
The Covalent Bond:
Sharing Electrons
56
A covalent bond consists of a pair of
electrons shared between two atoms.
In the millions of chemical compounds
that exist, the covalent bond is the
predominant chemical bond.
57
Substances which covalently bond exist
as molecules.
Carbon dioxide bonds covalently.
It exists as individually bonded
covalent molecules containing
one carbon and two oxygen
atoms.
58
11.7
The term molecule is not used when
referring to ionic substances.
Sodium chloride bonds ionically.
It consists of a large aggregate of
positive and negative ions. No
molecules of NaCl exist.
59
11.7
Covalent bonding in the hydrogen molecule
Two 1s orbitals from each of
two hydrogen atoms overlap.
Each 1s orbital contains 1
The two nuclei are
electron.
shielded from each
other by the electron
pair. This allows the
two nuclei to draw
close together.
11.8
The most likely
The
orbital
of the
region
to find
the
electrons
includes
two electrons
is
both
hydrogen
between
the two
nuclei.
nuclei.
60
of the molecule
Covalent bonding in The
theorbital
chlorine
The two nuclei are
electrons includes
Two 3p orbitals from
each
of
both
chlorine
two chlorine atomsnuclei.
overlap.
shielded from each
other by the
electron
pair.
This
The most
likely
allows
regionthe
to two
find the
nuclei
to draw is
two electrons
close
together.
between
the two
nuclei.
11.9
Each
Each unpaired
3pchlorine
orbital now has 8
electrons
on each chlorine
atomin its outermost
energy level.
contains 1 electron.
61
Covalent bonding with equal sharing of
electrons occurs in diatomic molecules
formed from one element.
hydrogen
chlorine
iodine
nitrogen
A dash may replace a pair of dots.
62
11.6
Electronegativity
63
electronegativity: The relative attraction
that an atom has for a pair of shared
electrons in a covalent bond.
64
• If the two atoms that constitute a
covalent bond are identical, then there
is equal sharing of electrons.
• This is called nonpolar covalent
bonding.
• Ionic bonding and nonpolar covalent
bonding represent two extremes.
65
• If the two atoms that constitute a
covalent bond are not identical, then
there is unequal sharing of electrons.
• This is called polar covalent bonding.
• One atom assumes a partial positive
charge and the other atom assumes a
partial negative charge.
– This charge difference is a result of the
unequal attractions the atoms have for
their shared electron pair.
66
Partial positivePartial
charge
negative charge
on hydrogen. on chlorine.
Polar Covalent Bonding in HCl
+
:
:
H Cl
-
Chlorine
hasthat
a greater
attraction
forelement
the
Shared
Theof
shared
electron
electron
pair. pair
The attractive
force
an atom
an
has
shared electron pair than is
hydrogen.
to chlorine than
for shared electrons in a molecule closer
or a polyatomic
ion
to hydrogen.
67
is known as its electronegativity.
A scale of relative electronegativities
was developed by Linus Pauling.
68
Electronegativity generally
decreases increases
down a left
group
to right
for
representative
across
a periodelements
.
.
69
The
electronegativities
metals are
The
electronegativities
of of
thethe
nonmetals
arelow.
high.
70
11.1
The polarity of a bond is determined by the
difference in electronegativity values of the
atoms forming the bond.
71
• If the electronegativity difference
between two bonded atoms is greater
than 1.7-1.9, the bond will be more
ionic than covalent.
• If the electronegativity difference is
greater than 2, the bond is strongly
ionic.
• If the electronegativity difference is
less than 1.5, the bond is strongly
covalent.
72
If the electronegativities are the same, the bond
is nonpolar covalent and the electrons are shared
equally.
The molecule is
nonpolar covalent.
Electronegativity
Difference = 0.0
Electronegativity
2.1
H
H
Electronegativity
2.1
Hydrogen Molecule
73
11.10
If the electronegativities are the same, the bond
is nonpolar covalent and the electrons are shared
equally.
The molecule is
nonpolar covalent.
Cl
Electronegativity
3.0
Cl
Electronegativity
Difference = 0.0
Electronegativity
3.0
Chlorine Molecule
11.10
74
If the electronegativities are not the same, the
bond is polar covalent and the electrons are
shared unequally.
The molecule is
polar covalent.
+
H
Electronegativity
2.1
Cl
Electronegativity
Difference = 0.9
Electronegativity
3.0
Hydrogen Chloride Molecule
11.10
75
If the electronegativities are very different, the
bond is ionic and the electrons are transferred to
the more electronegative atom.
No molecule exists.
The bond is ionic.
Electronegativity
Difference = 2.1
Na+
Electronegativity
0.9
ClElectronegativity
3.0
Sodium Chloride
11.10
76
A dipole is a molecule that is
electrically asymmetrical, causing it to
be oppositely charged at two points.
A dipole can be written as
+
77
An arrow can be used to indicate a dipole.
The arrow points to the negative end of the
dipole.
Molecules of HCl, HBr and H2O are polar .
O
H
Cl
H
Br
H
H
78
A molecule containing different kinds of
atoms may or may not be polar depending
on its shape.
The carbon dioxide molecule is nonpolar
because its carbon-oxygen dipoles cancel
each other by acting in opposite directions.
79
Relating Bond Type to
Electronegativity Difference.
80
11.11
11.7
Lewis Structures
of Compounds
81
In writing Lewis structures, the most
important consideration for forming a
stable compound is that the atoms attain
a noble gas configuration.
82
• The most difficult part of writing
Lewis structures is determining the
arrangement of the atoms in a molecule
or an ion.
• In simple molecules with more than
two atoms, one atom will be the central
atom surrounded by the other atoms.
83
Cl2O has two possible arrangements.
The two chlorines can be bonded to each other.
Cl-Cl-O
The two chlorines can be bonded to oxygen.
Cl-O-Cl
Usually the single atom will be the central atom.
84
Procedures for Writing
Lewis Structures
85
Valence Electrons of Group A Elements
Atom
Group
Valence Electrons
Cl
7A
7
H
1A
1
C
4A
4
N
5A
5
S
6A
6
P
5A
5
I
7A
7
86
Step 1. Obtain the total number of
valence electrons to be used in the
structure by adding the number of valence
electrons in all the atoms in the molecule
or ion.
–If you are writing the structure of an ion,
add one electron for each negative charge
or subtract one electron for each positive
charge on the ion.
87
Write the Lewis structure for H2O.
Step 1. The total number of valence electrons
is eight, two from the two hydrogen atoms and
six from the oxygen atom.
88
Step 2. Write the skeletal arrangement of
the atoms and connect them with a single
covalent bond (two dots or one dash).
– Hydrogen, which contains only one bonding
electron, can form only one covalent bond.
– Oxygen atoms normally have a maximum of
two covalent bonds (two single bonds, or
one double bond).
89
Write the Lewis structure for H2O.
Step 2. The two hydrogen atoms are connected
to the oxygen atom. Write the skeletal
structure:
:
H:O
H
or
H:O:H
Place two dots between the hydrogen and
oxygen atoms to form the covalent bonds.
90
Step 3. Subtract two electrons for each
single bond you used in Step 2 from the
total number of electrons calculated in
Step 1.
– This gives you the net number of electrons
available for completing the structure.
91
Write the Lewis structure for H2O.
Step 3. Subtract the four electrons used in Step
2 from eight to obtain four electrons yet to be
used.
H:O:H
92
Step 4. Distribute pairs of electrons (pairs
of dots) around each atom (except
hydrogen) to give each atom a noble gas
configuration.
93
H:O:
H
or
: :
: :
Write the Lewis structure for H2O.
Step 4. Distribute the four remaining electrons
in pairs around the oxygen atom. Hydrogen
atoms cannot accommodate any more
electrons.
H:O:H
These
arrangements
are
Lewis
structures
The shape of the molecule is not shown by the
because
each
atom
has
a
noble
gas
electron
Lewis structure.
structure.
94
Write a Lewis structure for CO2.
Step 1. The total number of valence electrons
is 16, four from the C atom and six from each
O atom.
95
Write a Lewis structure for CO2.
Step 2. The two O atoms are bonded to a
central C atom. Write the skeletal structure and
place two electrons between the C and each
oxygen.
O:C:O
96
Write a Lewis structure for CO2.
Step 3. Subtract the four electrons used in Step
2 from 16 (the total number of valence
electrons) to obtain 12 electrons yet to be used.
O:C:O
97
Write a Lewis structure for CO2.
4
electrons
:O:C:O:
:
: :
:
:
: :
:O:C:O:
:
: :
: :
Step 4. Distribute the 12 electrons (6 pairs)
around the carbon and oxygen atoms. Three
possibilities exist.
I
II
III
:O:C:O:
6 6
6
6
electrons
electrons
electronselectrons
Many of the atoms in these structures do not
have eight electrons around them.
98
: :
:: ::
Write a Lewis structure for CO2.
Step 5. Remove one pair of unbonded
electrons from each O atom in structure I and
place one pair between each O and the C atom
forming two double bonds.
::OO:::C:O
:O::
double bondEach atom now has 8 double bond
Carbon is sharing 4
electrons around it. electron pairs.
99
11.8
Complex Lewis Structures
100
There are some molecules and
polyatomic ions for which no single
Lewis structure consistent with all
characteristics and bonding information
can be written.
101
32
Write a Lewis structure for NO .
Step 1. The total number of valence electrons
is 24, 5 from the nitrogen atom and 6 from
each O atom, and 1 from the –1 charge.
102
32
Write a Lewis structure for NO .
Step 2. The three O atoms are bonded to a
central N atom. Write the skeletal structure
and place two electrons between each pair of
atoms.
:
O
O:N:O
103
32
Write a Lewis structure for NO .
Step 3. Subtract the 6 electrons used in Step 2
from 24, the total number of valence electrons,
to obtain 18 electrons yet to be placed.
:
O
O:N:O
104
32
: :
: : :
: :
Write a Lewis structure for NO .
Step 4. Distribute the 18 electrons around the
N and O atoms.
:O
:O:N:O:
electron deficient
105
32
: :
: : :
: :
Write a Lewis structure for NO .
Step 4. Since the extra electron present results
in nitrate having a –1 charge, the ion is
enclosed in brackets with a – charge.
:O
:O:N:O:
-
106
32
: :
:
:
: :
Write a Lewis structure for NO .
Step 5. One of the oxygen atoms has only 6
electrons. It does not have a noble gas
structure. Move the unbonded pair of electrons
from the N atom and place it between the N
and the electron-deficient O atom, making a
double bond.
:O
:O
-
electron deficient
N O:
:
107
32
Write a Lewis structure for NO .
Step 5. There are three possible Lewis
structures.
A molecule or ion that shows multiple
correct Lewis structures exhibits resonance.
Each Lewis structure is called a resonance
structure.
:O
N O
:O
:O
N O:
-
: :
:
:
: :
:O:
-
: :
:
:
: :
: :
:
:
: :
-
:O:
O N O:
108
11.9
Compounds Containing
Polyatomic Ions
109
A polyatomic ion is a stable group of
atoms that has either a positive or
negative charge and behaves as a single
unit in many chemical reactions.
110
Sodium nitrate, NaNO3, contains one
sodium ion and one nitrate ion.
Na
+
nitrate ion NO
3
-
: :
:
:
: :
sodium ion
Na+
:O
:O
N O:
111
• The nitrate ion is a polyatomic ion
composed of one nitrogen atom and
three oxygen atoms.
• It has a charge of –1
• One nitrogen and three oxygen atoms
have a total of 23 valence electrons.
: :
:
:
: :
Na
+
-
:O
:O
N O:
112
• The –1 charge on nitrate adds an
additional valence electron for a total
of 24.
• The additional valence electron comes
from a sodium atom which becomes a
sodium ion.
: :
:
:
: :
Na
+
-
:O
:O
N O:
113
• Sodium nitrate has both ionic and
covalent bonds.
• Covalent
Ionic bonds
bonds
exist are
between
present
thebetween
sodium
ions carbon
the
and the and
carbonate
oxygen
ions.
atoms within
the carbonate ion.
Na
+
covalent
bond
-
: :
:
:
: :
ionic
bond
:O
:O
N O:
covalent
bond
covalent
bond114
• When sodium nitrate is dissolved in
water the ionic bond breaks.
• The sodium
nitrate ion,
ionswhich
and nitrate
is heldions
together
separate
by
from eachbonds,
covalent
otherremains
formingasseparate
a unit. sodium
and nitrate ions.
Na
+
-
: :
:
:
: :
:
:
: :
Na
+
-
:O
:O
:O
N :O: N O:
115
11.10
Molecular Shape
116
The 3-dimensional arrangement of the
atoms within a molecule is a significant
feature in understanding molecular
interactions.
117
118
11.12
11.11
The Valence Shell
Electron Pair (VSEPR) Model
119
The VSEPR model is based on the idea
To
accomplish
this
minimization,
the
that electron pairs will repel each other
electron
pairs
will
be
arranged
as
far
electrically and will seek to minimize
apart
as
possible
around
a
central
atom.
this repulsion.
120
BeCl2 is a molecule with only two pairs of
electrons around beryllium, its central
atom.
Its electrons are arranged 180o apart for
maximum separation.
121
• BF3 is a molecule with three pairs of electrons
around boron, its central atom.
o apart
• This
Its electrons
arrangement
are ofarranged
atoms is 120
called
trigonal
for
maximum separation.
planar.
122
• CH4 is a molecule with four pairs of electrons
around carbon, its central atom.
• However,
An obvious
since the
choice
molecule
foris 3-dimensional,
its atomic
arrangement
the
molecularisstructure
a 90o angle
is tetrahedral
between its
with
atoms
a
with all
bond
angle
of its
of atoms
109.5oin
. a single plane.
123
Ball and stick models of methane, CH4, and carbon
tetrachloride, CCl4.
124
11.13
• Ammonia, NH3, has four electron pairs
around nitrogen.
The arrangement of
the electron pairs is
tetrahedral.
125
One
of
its
electron pairs is a
nonbonded
The shape (lone)
of the
pair.
NH3 molecule is
pyramidal.
126
• Water has four electron pairs around
oxygen.
The arrangement
of electron pairs
around oxygen is
tetrahedral.
127
Two
of
its
electron pairs are
The H2O molecule
nonbonded
(lone)
is bent.
pairs.
128
Structure Determination Using
VSEPR
1. Draw the Lewis structure for the
molecule.
2. Count the electron pairs and arrange
them to minimize repulsions.
3. Determine the positions of the atoms.
4. Name the molecular structure from
the position of the atoms.
129
130