Chapter 8 - Chemistry
Chapter
8
Study
Guide
 


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Unformatted text preview: Chemistry
Chapter
8
Study
Guide
 
 Chemical
bond:
a
force
that
holds
atoms
together
in
a
molecule
or
compound
 • Electrostatic
forces:
positive
and
negative
charges
holding
atoms
together
 Ionic
bond:
atoms
held
together
by
charges;
cation
(+)
and
anion
(‐)
 • metal
+
nonmetal
(or
metal
+
polyatomic
ion)
 • Electrons
transferred
from
metal
to
nonmetal

 • strong
bond
 • properties:
 
 
 High
melting
point
 
 
 Hard,
brittle
crystalline
solid
(at
room
temperature)

 
 
 Often
soluble
in
water
 • ionic
compounds
arrange
in
a
crystal
lattice
pattern
which
makes
an
ionic
crystal

 Covalent
bond:
electrons
are
shared,
not
transferred
 • two
nonmetals
 • Weak
bond

 • Properties:
 
 
 Low
melting
point
 
 
 Gas,
liquid,
soft
solid
at
(room
temperature)
 Nonpolar
covalent
bond:

 o electrons
shared
equally
 Polar
covalent
bond:
 o Electrons
shared
unequally
 o Leads
to
partial
charges
(δ+,
δˉ)
 Electronegatvity:
ability
of
an
element
to
attract
electrons
in
a
bond
 • electronegatvity
trend:
 Increase
from
bottom
to
top
and
increase
from
left
to
right

 




 
 
(points
toward
Flourine)
 • • • • • 
 
 



 
 the
greater
the
difference
in
electronegativities,
the
greater
the
partial
ionic
charge
on
 the
atoms
and
the
more
polar
the
bond

 Use
electronegativity
to
find
if
a
bond
is
polar:
 Ex.








H‐F








4.0‐2.2=1.8
so
bond
is
polar
covalent
 nonmetals
have
higher
electronegativities,
so
tend
to
gain
electrons
(make
negative
 charges)
 metals
have
lower
electronegativities,
so
lose
electrons
(make
positive
charges)
 δ:
partial
charge
δ+
and
δ‐

 
 put
on
the
side
of
Lewis
structure
which
is
more
partially
negative
or
positive
 can
also
use













to
point
to
more
partially
negative
atom
 Bonds:
using
electronegatvities,
subtract
the
highest
from
the
lowest
(doesn’t
matter
which)
 • nonpolar
covalent:
0‐0.4
 • polar
covalent:
0.5‐1.7
 • ionic:
1.8
<
 
Octet
Rule:
maximum
of
8
valence
electrons
for
most
elements
 • Atoms
trying
to
reach
noble
gas
configurations
which
have
8
valence
electrons
(stable)
 • Exceptions:
H,
He
(doesn’t
make
bonds),
Li,
Be,
B
 
 Some
atoms
can
have
expanded
octet
i.e.:
S,
Cl
 Lewis
Symbols:
dots
placed
around
an
element’s
symbol
to
represent
its
valence
electrons
 • Place
a
dot
on
each
side
and
if
more
electrons
go
around
again
;
Each
side
has
no
more
 than
2
dots
and
no
more
than
8
dots
all
together
 • Also
used
to
figure
out
how
many
electrons
needed
to
reach
octet
 Ex.
 N
needs
3
more
electrons
to
reach
octet
 Lewis
Formulas:
 • Single
bond:
one
shared
pair
of
electrons
 • Double
bond:
two
shared
pairs
of
electrons
 • Triple
bond:
three
shared
pairs
of
electrons
 • Atoms
combine
in
different
ways
to
reach
octets
 o ex:
oxygen
can
form
either
double
bonds
(O₂)
or
2
single
bonds
(H₂O)
 O=O
 
 How
to
draw
Lewis
structures:
 • Add
up
all
valence
electrons
of
atoms
(picture
must
have
same
number)
 • Draw
skeleton
of
structure
 o Central
atom
is
usually
the
first
in
the
formula
 o Arrangement
of
atoms
is
usually
symmetrical
 o Hydrogen
atoms
form
only
one
bond,
so
only
go
on
outside
of
structure
 • Distribute
remaining
electrons
 • When
don’t
have
enough
electrons,
make
multiple
bonds
(usually
on
the
central
atom)
 o Ex:
SO₂
 Resonance
hybrids:
2
or
more
Lewis
formulas
that
represent
the
bonding
a
molecule
 • The
real
structure
of
the
molecule
is
somewhere
in
between
the
pictures
 • All
structures
follow
octet
rule
 • Don’t
change
positions
of
atoms,
only
electrons
 • Connect
pictures
with
double
headed
arrow

 • resonance
often
just
switches
which
atom
has
the
single
bond,
which
has
double
bond
 Carbon
forms
4
covalent
bonds:
variations
of
single,
double,
triple
bonds
 Molecular
Shapes:
 • Valence
Shell
Electron
Repulsion
Theory
(VSEPR):
tendency
of
electron
pairs
to
adjust
 orientation
of
orbitals
to
maximize
distance
between
them
 • Bonded
atoms
and
unshared
electron
pairs
are
as
far
apart
as
possible
 • bond
angle
from
central
atom
and
atoms
bonded
to
it
 Parent
structure:
a
3D
shape
concerning
the
central
atom
and
bonded
atoms/unshared
 electron
pairs
 • Linear:
2
atoms/
electron
pairs
 
 HCN,
F₂
 • Trigonal
planar:
3
atoms/electron
pairs
 NO₃ˉ,
SO₂
 • Tetrahedral:
4
atoms/electron
pairs

 H₂O,
CH₄,
NH₃
 Molecular
Structure:
 • Linear:
2
groups
of
shared
electron
pairs,
usually
double
bonds;
no
unshared
pairs
 • Bent:
2
groups
of
shared
electron
(single
or
double
bonds);
1
or
2
unshared
electron
 pairs
around
central
atom
 o Parent
structure
can
be
trigonal
planar
or
tetrahedral
 • Trigonal
planar
(molecular
structure):
3
groups
shared
electrons
around
central
atom
(2
 single,
1
double);
no
unshared
electrons
 • Trigonal
pyramidal:
3
pairs
shared
electrons;
1
unshared
pair
around
central
atom
 • Tetrahedral:
4
pairs
shared
electrons,
no
unshared
around
central
atom
 Bond
Angle:
 • Tetrahedral
(molecular
structure):
109.5˚
 • Trigonal
planar
(molecular):
120˚
 • Linear
(molecular):
180˚
 • Trigonal
pyramidal
(Molecular):
107˚
 • Bent
(tetrahedral
parent
structure):
105˚
 • Bent
(trigonal
planar
parent
structure):
119˚
 
To
determine
Molecular
Shape:
 1) Draw
Lewis
structure
 2) Determine
regions
of
electrons
around
central
atom
(shared
and
unshared)
 3) Determine
electron
(parent)
shape
(by
adding
atoms
&
unshared
e
pairs)
 4) Determine
molecular
shape
by
considering
only
positions
of
bonded
atoms
 Ex:
CO₂


16
valence
electrons



1) 

2)
2
regions
 
 
 
 
 3)
Linear
 4)Linear
 Polarity:
partially
positive
or
negative
molecule
 • Polyatomic
molecule
that
has
nonpolar
bonds
cannot
be
polar
 • Some
polyatomic
bonds
with
polar
bonds
may
be
dipoles
depending
on
geometry
 • To
be
nonpolar
molecule
with
polar
bonds
 o The
magnitudes
of
the
the
atoms
have
to
be
the
same

 o Directions
of
bonded
atoms
have
to
cancel
out
 • Polar
molecule
(dipole)
with
polar
bonds:
 o Bonds
are
polar
and
not
symmetrical
 o The
atoms
around
central
atom
aren’t
identical
(magnitudes
don’t
cancel)
 Ex:
 
 
 
 






 
 
 ...
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