Chemical Bonds
The Formation of
Compounds From Atoms
Chapter 11
Hein and Arena
1
11.2
Lewis Structures of Atoms
2
Metals form cations and nonmetals form
anions to attain a stable valence electron
structure.
3
Thesestable
This
rearrangements
structure often
occur
consists
by losing,
of twogaining,
s and
sixsharing
or
p electrons.
electrons.
4
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
5
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.
7
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.
8
Lewis Structures of the first 20 elements.
9
11.4
11.3
The Ionic Bond
Transfer of Electrons From
One Atom to Another
10
The chemistry of many elements,
especially the representative ones, is to
attain the same outer electron structure
as one of the noble gases.
11
With the exception of helium, this
structure consists of eight electrons in
the outermost energy level.
12
After sodium loses its 3s electron, it has attained the
same electronic structure as neon.
13
After chlorine gains a 3p electron, it has attained the
same electronic structure as argon.
14
Formation of NaCl
15
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.
16
Formation of MgCl2
17
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.
18
In NaCl
the crystal
is made
each
upsodium
of cubicion
crystals.
is surrounded by six
chloride ions.
19
In the crystal each chloride ion is surrounded by six
sodium ions.
20
11.5
The ratio of Na+ to Cl- is 1:1
There is no molecule of NaCl
21
11.5
11.6
Electronegativity
22
electronegativity: The relative attraction
that an atom has for a pair of shared
electrons in a covalent bond.
23
• 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.
24
• 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.
25
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.
26
is known as its electronegativity.
A scale of relative electronegativities
was developed by Linus Pauling.
27
Electronegativity generally
decreases increases
down a left
group
to right
for
representative
across
a periodelements
.
.
28
The
electronegativities
metals are
The
electronegativities
of of
thethe
nonmetals
arelow.
high.
29
11.1
The polarity of a bond is determined by the
difference in electronegativity values of the
atoms forming the bond.
30
• 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.
31
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
32
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
33
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
34
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
35
A dipole is a molecule that is
electrically asymmetrical, causing it to
be oppositely charged at two points.
A dipole can be written as
+
36
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
37
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.
38
Relating Bond Type to
Electronegativity Difference.
39
11.11
11.7
Lewis Structures
of Compounds
40
In writing Lewis structures, the most
important consideration for forming a
stable compound is that the atoms attain
a noble gas configuration.
41
• 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.
42
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.
43
Procedures for Writing
Lewis Structures
44
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
45
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.
46
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.
47
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).
48
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.
49
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.
50
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
51
Step 4. Distribute pairs of electrons (pairs
of dots) around each atom (except
hydrogen) to give each atom a noble gas
configuration.
52
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.
53
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.
54
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
55
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
56
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.
57
: :
:: ::
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.
58
11.8
Complex Lewis Structures
59
There are some molecules and
polyatomic ions for which no single
Lewis structure consistent with all
characteristics and bonding information
can be written.
60
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.
61
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
62
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
63
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
64
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:
-
65
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:
:
66
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:
67
11.9
Compounds Containing
Polyatomic Ions
68
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.
69
Sodium nitrate, NaNO3, contains one
sodium ion and one nitrate ion.
Na
+
nitrate ion NO
3
-
: :
:
:
: :
sodium ion
Na+
:O
:O
N O:
70
• 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:
71
• 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:
72
• 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
bond73
• 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:
74
11.10
Molecular Shape
75
The 3-dimensional arrangement of the
atoms within a molecule is a significant
feature in understanding molecular
interactions.
76
77
11.12
11.11
The Valence Shell
Electron Pair (VSEPR) Model
78
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.
79
BeCl2 is a molecule with only two pairs of
electrons around beryllium, its central
atom.
Its electrons are arranged 180o apart for
maximum separation.
80
• 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.
81
• 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.
82
Ball and stick models of methane, CH4, and carbon
tetrachloride, CCl4.
83
11.13
• Ammonia, NH3, has four electron pairs
around nitrogen.
The arrangement of
the electron pairs is
tetrahedral.
84
One
of
its
electron pairs is a
nonbonded
The shape (lone)
of the
pair.
NH3 molecule is
pyramidal.
85
• Water has four electron pairs around
oxygen.
The arrangement
of electron pairs
around oxygen is
tetrahedral.
86
Two
of
its
electron pairs are
The H2O molecule
nonbonded
(lone)
is bent.
pairs.
87
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.
88
89