CH2Notes - Chapter
Two
 Chemical
Bonds
...

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Unformatted text preview: Chapter
Two
 Chemical
Bonds
 CHEM
1211K
 Fall
2010
 1
 Objec7ves
 2
 Key
Ideas
 •  Bond
forma4on
is
accompanied
by
a
lowering
 of
energy
due
to
the
aArac4ons
between
 oppositely
charged
ions
or
between
nuclei
and
 shared
electron
pairs.
 •  The
electron
configura4ons
of
individual
 atoms
control
how
the
atoms
combine
with
 one
another.
 3
 Why
is
it
important?
 •  The
existence
of
compounds
is
central
to
the
science
 of
chemistry.
 •  By
seeing
how
bonds
form
between
atoms,
we
come
 to
see
how
chemists
design
new
materials.
 •  Research
into
ar4ficial
blood,
new
pharmaceu4cals,
 agricultural
chemicals,
and
the
polymers
used
in
 materials
such
as
compact
discs,
cell
phones,
and
 synthe4c
fibers
is
based
on
an
understanding
of
how
 atoms
link
together.
 4
 Chemical
Bonds
 •  A
chemical bond
is
the
link
between
atoms.
 –  When
a
chemical
bond
forms
between
two
atoms
 the
resul4ng
arrangement
of
the
two
nuclei
and
 their
electrons
has
a
lower
energy
than
the
total
 energy
of
the
separate
atoms.
 5
 Chemical
Bonds
 •  Ionic bond •  Covalent bond
 6
 Ionic
Bonds
 •  An
ionic solid
is
an
assembly
of
ca4ons
and
 anions
stacked
together
in
a
regular
array.
 –  Ionic
solids
are
examples
of
crystalline solids
 Figure
2.1
 7
 The
Ions
that
Elements
Form
(2.1)
 •  Atoms
gain
or
lose
electrons
in
such
a
way
as
 to
achieve
noble
gas
configura4on.
 –  The
noble
gas
core
configura4on
is
called
an
octet.
 ns2np6
 •  s‐block
metals
and
p‐block
metals
in
periods
2
 and
3
 •  p‐block
nonmetals
and
metals
in
periods
4
and
 lower 8
 The
Ions
that
Elements
Form
(2.1)
 •  When
the
d‐block
is
occupied
by
electrons,
 (n‐1)d‐orbitals
are
lower
in
energy
than
the
ns
 orbitals.

 •  Example:
What
is
the
electron
configura4on
of
 the
Fe2+
ion?
 9
 The
Ions
that
Elements
Form
(2.1)
 Figure
C.6
 Figure
C.7
 10
 The
Ions
that
Elements
Form
(2.1)
 •  Sec4on
summary:
 –  To predict the electron configura9on of a monatomic ca9on, remove outermost electrons in the order: np, ns, and (n‐1)d. –  To form a monatomic anion, add electrons un9l the next noble gas configura9on has been reached. –  The transfer of electrons results in the forma9on of an octet (or duplet) of electrons in the valence shell on each of the atoms. –  Metals achieve an octet (or duplet) by electron loss and nonmetals achieve it by electron gain. 11
 Lewis
Symbols
(2.2)
 •  Lewis symbols
are
a
simple
way
to
keep
track
 of
valence
electrons
when
atoms
form
bonds.
 –  The
element
is
represented
by
the
chemical
 symbol.
 –  A
single
dot
represents
an
electron
alone
in
an
 orbital.
 –  A
pair
of
dots
represents
two
paired
electrons
 sharing
an
orbital.
 12
 Lewis
Symbols
(2.2)
 13
 Lewis
Symbols
(2.2)
 •  To
work
out
the
formula
of
an
ionic
compound
 using
Lewis
symbols:
 –  Represent
the
ca4on
by
removing
the
dots
from
the
 symbol
for
the
metal
atom.
 –  Represent
the
anion
by
transferring
those
dots
to
the
 Lewis
symbol
for
the
nonmetal
atom
to
complete
the
 valence
shell.
 –  It
might
be
necessary
to
adjust
the
numbers
of
atoms
 of
each
kind
to
that
the
dots
removed
from
the
metal
 atom
symbols
are
accommodated
by
the
nonmetal
 atom
symbols.
 –  Write
the
charge
of
each
ion
as
a
superscript
in
the
 normal
way,
using
brackets
to
indicate
that
the
charge
 is
the
overall
charge
of
the
ion.
 14
 Lewis
Symbols
(2.2)
 •  Example:

Draw
a
Lewis
structure
to
represent
 lithium
chloride.
 •  Example:

Draw
a
Lewis
structure
to
represent
 magnesium
bromide.
 15
 Lewis
Symbols
(2.2)
 •  Sec4on
summary:
 –  Formulas of compounds consis9ng of the monatomic ions of main‐group elements can be predicted by assuming that ca9ons have lost all their valence electrons and anions have gained electrons in their valence shells un9l each ion has an octet of electrons (or a duplet in the case of H, Li, and Be). 16
 The
Energe7cs
of
Ionic
Bond
 Forma7on
(2.3)
 •  Why
do
ionic
solids
form?

What
is
the
driving
 force
that
makes
the
solid
more
favorable
than
 separated
sodium
and
chlorine
atoms?
 17
 The
Energe7cs
of
Ionic
Bond
 Forma7on
(2.3)
 Figure
2.4
 •  Why
the
crystal
of
sodium
 chloride
has
lower
energy
 than
the
separate
atoms?
 –  Sodium
atoms
release
 electrons
 –  Electrons
released
by
sodium
 aAach
to
chlorine
 –  Resul4ng
sodium
ca4ons
and
 chlorine
anions
clump
 together
in
a
crystal
 18
 The
Energe7cs
of
Ionic
Bond
 Forma7on
(2.3)
 •  Sec4on
summary:
 –  The energy required for the forma9on of ionic bonds is supplied largely by the aJrac9on between oppositely charged ions. •  It must be strong enough to overcome the energy investment needed to make the ions. 19
 Interac7ons
Between
Ions
(2.4)
 •  An ionic solids is NOT held together by bonds between specific pairs of ions •  An
ionic
bond
is
a
“global”
characteris4c
of
the
 en4re
crystal—a
net
lowering
of
energy
of
the
 en4re
crystal.
 20
 Interac7ons
Between
Ions
(2.4)
 •  To
determine
strength
of
interac4on:
 –  Look
at
charge
first.
 –  Then
look
at
ionic
radii.
 •  AJrac9on is greatest when ions are small and highly charged. 21
 Interac7ons
Between
Ions
(2.4)
 Figure
2.8
 22
 Interac7ons
Between
Ions
(2.4)
 •  Sec4on
summary:
 –  Ionic solids typically have high mel9ng points and are briJle. –  The Coulombic interac9on between ions in a solid is large when the ions are small and highly charged. 23
 Covalent
Bonds
 •  A
covalent bond
is
a
pair
of
electrons
shared
 between
two
atoms.
 –  Covalent
bonds
arise
because
they
contain
 nonmetals
which
do
not
form
ca4ons!
 Page
63
 24
 Lewis
Structures
(2.5
and
2.6)
 •  Remember
that
in
electrons
are
shared
in
 covalent
bonds.
 –  When
possible,
the
valence
electrons
in
a

 compound
are
distributed
in
such
a
way
that
each
 main‐group
element
in
a
molecule
(except
H
or
 He)
is
surrounded
by
eight
electrons
(an
octet
of
 electrons).


 –  Hydrogen
and
helium
should
have
two
electrons
 in
such
a
structure.
 25
 Lewis
Structures
(2.5
and
2.6)
 •  Lewis
structures
are
a
good
way
to
show
 covalent
bond
forma4on.
 –  A
pair
of
dots
between
the
two
atom
symbols
or
a
 dashed
line
represents
the
covalent
bond.
 –  This
is
called
a
bonding pair.
 –  Look
atH2
as
an
example:
 H






1s1
 H∙

+

H∙


→


H:H
 26
 Lewis
Structures
(2.5
and
2.6)
 •  Valence
Electron
Count
Method
for
Drawing
Lewis
 Structures
 –  Determine
which
is
the
central
atom
 •  It’s usually the least electronegative element. •  It’s also easily identified if there is only one of one type of atom and several of another type. The lonely guy is usually the central atom. –  Connect
each
atom
to
the
central
atom
using
a
 single
bond
(single
bonds
represent
2
electrons).
 27
 Lewis
Structures
(2.5
and
2.6)
 •  Valence
Electron
Count
Method
for
Drawing
Lewis
Structures
 –  Distribute
the
remaining
electrons,
in
pairs.
 •  Start with one atom and fill its octet. •  Proceed to the next atom, etc. until you run out of electrons. •  These are called lone pairs. –  COUNT
 •  Does everyone have a complete octet? •  Does the total number of electrons match the count you did at the beginning? 28
 Lewis
Structures
(2.5
and
2.6)
 •  Valence
Electron
Count
Method
for
Drawing
Lewis
 Structures
 –  Incomplete
octets?
 –  Move
pairs
of
electrons
to
turn
single
bonds
into
 double
bonds
or
double
bonds
into
triple
bonds.
 –  You
don’t
have
enough
electrons
in
your
 structure?
 –  Add
lone
pairs
to
the
central
atom.

This
may
cause
you
 to
violate
the
octet
rule,
but
that’s
okay
in
some
cases.
 29
 Lewis
Structures
(2.5
and
2.6)
 •  Draw
the
Lewis
structure
for
CCl2F2.
 30
 Lewis
Structures
(2.5
and
2.6)
 •  Draw
the
Lewis
structure
for
NF3.
 31
 Lewis
Structures
(2.5
and
2.6)
 •  Draw
the
Lewis
structure
for
CO2.
 32
 Lewis
Structures
(2.5
and
2.6)
 •  Draw
the
Lewis
structure
for
the
ammonium
 ion.
 33
 IV.
Lewis
Structures
(6.9—6.12)
 •  Things
get
a
liAle
more
complicated
if
there
is
 more than one central atom
in
the
molecule.
 –  We
have
to
consider
the
maximum
number
of
 bonds
a
given
atom
can
form.
 –  Write
the
Lewis
structure
for
methanol
(molecular
 formula
CH4O).
 •  What’s
the
central
atom?
 34
 Lewis
Structures
(2.5
and
2.6)
 •  Sec4on
summaries:
 –  Nonmetal atoms share electrons un9l each has completed its octet (or duplet). –  A Lewis structure shows the arrangement of electrons as lines (bonding pairs) and dots (lone pairs). –  The Lewis structure of a polyatomic species is obtained when all the valence electrons are used to complete the octets (or duplets) of the atoms present by forming single or mul9ple bonds and leaving some electrons as lone pairs. 35
 Resonance
(2.7)
 O3
can
be
drawn
in
2
ways
‐
 O O O O O O •  Molecules
like
O3
are
not
adequately
represented
by
a
 single
Lewis
structure.
 –  Both
O3
structures
are
valid.
 –  Both
have
the
same
energy.
 –  The experimental bond lengths match neither structure. 36
 Resonance
(2.7)
 •  Because
the
bond
lengths
in
O3
are
iden4cal,
a
 beAer
model
is
a
blend
of
the
two
Lewis
 structures
with
eah
bond
intermediate
in
 proper4es
between
a
single
and
a
double
 bond.
 –  The
blending
of
structures
is
called
 –  The
blended
structures
is
a

 37
 Resonance
(2.7)
 •  The
electrons
that
are
shown
in
different
 posi4ons
in
a
set
of
resonance
structures
are
 said
to
be
delocalized.
 38
 Resonance
(2.7)
 •  To
exhibit
resonance,
each
contribu4ng
 structure
must
have
nuclei
in
the
same
 posi4on.
 –  Only
the
loca4ons
of
lone
pairs
and
bonding
pairs
 are
changed.
 39
 Resonance
(2.7)
 •  Sec4on
summary:
 –  Resonance is a blending of structures with the same arrangement of atoms but different arrangements of electrons. –  It spreads mul9ple bond character over a molecule and results in a lower energy. 40
 Formal
Charge
(2.8)
 •  Experimental
determina4ons
of
structures
are
 the
only
way
to
unequivocally
iden4fy
correct
 structures.
 •  When
we
can’t
do
an
experiment,
we
look
at
 formal charge.
 Formal
charge
=
V
–
(L
+
½
B)
 Formal
charge
=
valence
–
(lone
+
½
bonding)
 41
 Formal
Charge
(2.8)
 •  Example:

Assign
formal
charges
to
each
atom
 in
both
possible
structures
for
NCO‐.
 N C A O ‐
 NC B
 O ‐
 ‐
 N C O C
 42
 Formal
Charge
(2.8)
 •  There
are
three
primary
criteria
for
selec4ng
 the
dominant
or
“best”
resonance
structure.
 –  Smaller
formal
charges
(either
posi4ve
or
 nega4ve)
are
preferable
to
larger
charges.
 –  Avoid
like
charges
(+
+
or
‐
‐
)
on
adjacent
atoms.
 –  A
more
nega4ve
formal
charge
should
exist
on
an
 atom
with
a
larger
electronega4vity
value.
 43
 Formal
Charge
(2.8)
 •  Formal
charge
and
oxida4on
number
both
 give
us
informa4on
about
the
number
of
 electrons
around
an
atom
in
a
compound,
but
 they
are
different
methods
and
onen
have
 different
values.
 –  Formal
charge
exaggerates
covalent
character
of
 bonds.
 –  Oxida4on
number
exaggerates
ionic
character
of
 bonds.
 44
 Resonance
(2.7)
 •  Draw
the
resonance
hybrid
for
the
perchlorate
 ion.
 45
 Formal
Charge
(2.8)
 •  Sec4on
summary:
 –  The formal charge gives an indica9on of the extent to which atoms have gained or lost electrons in the process of covalent bond forma9on. –  Atom arrangements and Lewis structures with the lowest formal charges are likely to have the lowest energy. 46
 Excep7ons
to
the
Octet
Rule
 •  C,
N,
O,
and
F
obey
the
octet
rule
rigorously.
 •  Other
atoms
may
not
always
obey
the
octet
 rule,
depending
on
the
situa4on


 –  Some
molecules
have
an
odd
number
of
electrons
 –  Some
atoms
can
accommodate
more
than
eight
 valence
electrons

 –  Some
atoms
are
happy
to
have
less
than
an
octet
 47
 Expanded
Valence
Shells
(2.10)
 •  When
the
central
atom
in
a
molecule
has
 empty
d‐orbitals,
it
may
be
able
to
 accommodate
more
than
eight
valence
 electrons.

This
is
called
an
expanded octet.
 –  These
electrons
may
be
present
either
as
lone
or
 bonding
pairs
of
electrons.
 48
 Lewis
Structures
(2.5
and
2.6)
 •  Draw
the
Lewis
structure
for
I3‐.
 •  Draw
the
Lewis
structure
for
PCl5.
 49
 The
Unusual
Structures
of
Some
 Group
III
Compounds
(2.11)
 •  A
compound
that
contains
an
atom
with
more
atoms
 aAached
to
it
than
is
permiAed
by
the
octet
rule
is
called
 a
hypervalent compound.
 •  Group
III
atoms
onen
have
an
incomplete octet—they
 exist
in
molecules
with
only
six
electrons
in
their
valence
 shell.
 –  Look
at
the
formal
charges
on
two
possible
Lewis
 structures
for
BF3.
 50
 The
Unusual
Structures
of
Some
 Group
III
Compounds
(2.11)
 •  Sec4on
summary:
 –  Compound of boron and aluminum may have unusual Lewis structures in which they have incomplete octets. 51
 Ionic
Versus
Covalent
Bonds
 •  Ionic
and
covalent
bonding
are
two
extreme
 models
of
the
chemical
bond.
 –  Most actual bonds lie somewhere between purely ionic and purely covalent.
 52
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  All
bonds
can
be
viewed
as
resonance
hybrids
 of
purely
covalent
and
purely
ionic
structures.
 :
Cl
 Cl
:
 ‐
 :
Cl
 Cl
:
 +
 ‐
 :
+
 :
Cl
 Cl
 •  In

molecule
composed
of
different
elements,
 the
resonance
has
unequal
contribu4ons
from
 the
two
ionic
structures.
 : H
 Cl
 
 ‐
 : H
:
 Cl
 
 +
 ‐
 :
+
 : H
 
 Cl
 53
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  The
charges
on
the
atoms
in
HCl
are
called
par9al charges
and
are
represented
by
δ+


and
δ‐.
 •  A
bond
in
which
ionic
contribu4ons
to
the
 resonance
result
in
par4al
charges
is
called
a
 polar covalent bond
 –  Any
bond
between
two
different
atoms
(EXCEPT
C
and
 H)
is
considered
a
polar
covalent
bond.
 54
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  The
two
atoms
in
a
polar
covalent
bond
form
 an
electric dipole
(or
just
dipole).
 –  A
dipole
is
a
par4al
posi4ve
charge
next
to
an
 equal
but
opposite
par4al
nega4ve
charge.
 Figure
page
77
 55
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  A
covalent
bond
is
polar
if
one
atom
has
 greater
aArac4on
for
electrons
than
the
other
 atom.
 –  The
ability
of
an
atom
to
aAract
electrons
to
itself
 when
in
a
molecule
is
called
electronega9vity.
 56
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  Electronega4vity
increases
up
a
group
and
 from
len
to
right
across
a
period.
 57
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 Figure
2.13
 58
 Correc7ng
the
Covalent
Model:

 Electronega7vity
(2.12)
 •  Sec4on
summary:
 –  Electronega9vity is a measure of the pulling power of an atom on the electrons in a bond. –  A polar covalent bond is a bond between two atoms with par9al electric charges arising from their differences in electronega9vity. –  The presence of par9al charges gives rise to an electric dipole moment. 59
 Correc7on
the
Ionic
Model:

 Polarizability
(2.13)
 •  All
ionic
bonds
have
some
covalent
character,
 but
why?
 Figure
2.14
 60
 Correc7on
the
Ionic
Model:

 Polarizability
(2.13)
 •  Atoms
and
ions
that
readily
undergo
a
large
 distor4on
are
said
to
be
highly
 polarizable.
 –  The
more
electrons
an
atom
posses,
the
more
 polarizable
it
is.
 •  How
4ghtly
the
electrons
are
held
by
the
 nucleus
also
affects
polarizablity.
 61
 Correc7on
the
Ionic
Model:

 Polarizability
(2.13)
 •  Sec4on
summary:
 –  Compounds composed of highly polarizing ca9ons and highly polarizable anions have significant covalent character in their bonding. 62
 Bond
Strengths
(2.14)
 •  The
strength
of
a
chemical
bond
is
measured
 by
the
energy
required
to
separate
bonded
 atoms.

This
is
called
the
dissocia9on energy
 Figure
2.15
 63
 Varia7on
in
Bond
Strength
(2.15)
 64
 Varia7on
in
Bond
Strength
(2.15)
 •  Trends
in
bond
strength
also
correlate
to
 atomic
radii.
 –  Smaller
atoms
can
get
closer
 


together
and
form
stronger
bonds.
 –  Larger
atoms
cannot
get
close
 


enough
to
form
strong
bonds.
 65
 Figure
2.18
 Varia7on
in
Bond
Strength
(2.15)
 •  The
presence
of
lone
pairs
may
influence
the
 strength
of
bonds.
 66
 Varia7on
in
Bond
Strength
(2.15)
 Figures
page
81
of
text
 67
 Varia7on
in
Bond
Strength
(2.15)
 •  Sec4on
summary:
 –  The bond strength increases as the mul9plicity of a bond increases. –  Bond strength decreases as the number of lone pairs on neighboring atoms increases. –  Bond strength decreases as the atomic radii increase. –  Bonds are strengthened by resonance. 68
 Bond
Lengths
(2.16)
 •  A
bond length is
the
distance
between
the
 centers
of
two
atoms
joined
by
a
covalent
 Figure
page
82
 bond.
 Figure
2.15
 69
 Bond
Lengths
(2.16)
 Figure
2.21
 70
 Bond
Lengths
(2.16)
 •  Sec4on
summary:
 –  The stronger the bond, the shorter it is. –  The weaker the bond, the longer it is. •  Example:

Which
molecule
has
the
longest
C
 to
O
bond?

Which
has
the
strongest?
 CO2,
CO,
CH3OH
 71
 ...
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This note was uploaded on 01/16/2012 for the course CHEM 1211 taught by Professor Ford during the Fall '09 term at Georgia Tech.

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