Chemical Bonding & Lewis Dot
Structures
I. Covalent & Ionic bonding
A. chemical bonding –
1. octet rule – chem compounds
tend to
form so that each
atom (gain e-, share
e-, lose e-)
has an octet (8) of electrons
in
its highest occupied E-level.( or
have
a noble gas
configurations)
[why??? because nature favors a
lower PE & atoms have lower PE when
bonded---the bond length between 2
bonded atoms is the distance between
the atoms when the PE is the least---this PE is the bond energy or how
much E is required to break the chem
bond chart p.168 pm = picometer =
10-12m]
2. attractions between nuclei &
valence e- of
dif. atoms that hold
atoms together
3. atoms bond & redistribute their
valence eto make the atoms
more stable
4. types of bonds
a. ionic bonds
b. covalent bonds —polar covalent
& nonpolar
covalent
c. metallic bonds
d. hydrogen bonds
5. How to determine if a bond will be
ionic,
polar covalent, or nonpolar
covalent
a. look at the electronegativity
chart!!!
p. 162 & p.151
1) Bonding between atoms of
dif.
elements is
rarely purely ionic or
purely covalent. It usually falls
somewhere in between
the 2
extremes
depending on how strongly
the atoms of each attract
electrons
….their
electronegativity.
2) look at differences in
electroneg. to
determine
type of bond (p. 151)
a) difference from 0 to
0.3 =
nonpolar
covalent bond (p.162)
b) difference from 0.3 to
1.7 = polar
covalent bond
c) difference from 1.7 to
3.3 = ionic
bond
d) the atom with the
larger
electroneg. will be the more
negative atom
B. Covalent bonding – sharing electron
pairs
between 2 atoms—a
molecule is produced -1. molecules – produced by covalent
bonds
between
atoms
a) diatomic molecules – contain
only 2 atoms
ex: O2, N2, H2
b) chemical formula – shows #’s
& kinds of
atoms in a
compound using element
symbols & subscripts
2. nonpolar-covalent bond –
“pure”covalent
bond where the
electrons are shared
equally so
there is an even distribution of e3.polar covalent –
covalent bond
where
there is an unequal
attraction for the shared
electrons….which
means
one end of the
molecule is
slightly more neg. δ- & the
other end is
slightly more
positive δ+
4. octet rule applies to most there are
exceptions
a) less than an octet – H (2), Be
(4), B (6)
b) more than an octet – SF6, PCl5
c) odd # of electrons – NO (5 + 6 =
11)
5. characteristics of covalent cmpds
a) low melting & boiling pts
b) most are rubbery, powdery,
liquids or gases
c) does not conduct electricity
d) polar molecules dissolve in
other polar
molecules
e) nonpolar molecules dissolve in
nonpolar
molecules
6. single bond – sharing 1 pair
of
electrons (short
bond length)
7. double bond – sharing 2 pr of
ebetween 2 atoms
(shorter bond
length)
8. triple bond – sharing 3 pr of
ebetween 2 atoms
(shortest bond
length)
C. Ionic bonding—transfer of
electrons from
one atom to another
1. ionic compound -- made of pos &
neg ions that
are combined so # of
pos & neg charges are
equal
2. formula unit – shows the types & #
of
ions in an ionic cmpd ( pos.
ion =
cation- always goes
first)
3. ions reduce PE by combining in
crystal
formations
4. to compare bond strengths – lattice
E–
E released when 1 mole of
gaseous ions
form into a solid crystal
structure
5. characteristics of ionic cmpds
a) high melting & boiling pts
b) exist as crystals that are brittle,
shifting the structure
(striking it) will
cause cleavage
(break along
patterns)
c) many are soluble in water
d) conduct electricity when
dissolved or
melted
6. polyatomic ion – charged group of
covalently bonded atoms – has
both ionic & molecular char.--- they
combine with other ions to form
ionic cmpds
D. Metallic bonding—chem bonding
between
metals that results from
the attraction
between metal
atoms & the SEA of esurrounding them – the bonding is the
same in all directions
1. in metals the highest E level has
very few e2. the vacant p (& d) orbitals overlap
3. this allows the valence e- to roam
freely
throughout the
entire metal
a) this allows for great
conductivity of heat
& electricity
b) small distances between orbitals
allow
for a wide range of photons
to be absorbed
easily (& re-emitted
easily in the form of
light which
makes metals shiny)
c) since bonding is same in all
directions
(metals don’t
break easily) metals
are
malleable (able to be hammered
into sheets) & ductile (able to be
pulled
into thin wires)
d) very high melting & boiling pts.
4. heat of vaporization – measure of
bond
strength – the amt of heat
required to
vaporize metal into
individual metal atoms
in the
gaseous (vaporized) state
E. Other types of bonds (weaker)
1. Hydrogen bonding –
intermolecular
attractions
that occur between polar
covalent molecules that have a
hydrogen
atom bonded to an
extemely electroneg.
atom (N,
O, F)—the H atom will be δ+
and be attracted to the δ(unshared pair
of e-) of another like molecule
a) this accounts for the high
boiling pt of
H2O, HF, NH3
2. London Forces – due to the
constant
motion of emomentary polarity occurs in
all
atoms and molecules (distribution of
euneven) this can weakly hold
atoms/molecules
together
a) the larger the atomic/molar
mass the
greater the attraction
& the greater
London
forces can increase the bp
II. Electron Dot Notation
A. electron configuration that
shows how
the valence e- are
arranged around an atom ……shown
usually as dots
1. ex: p. 170
III. Lewis Structures
A. electron dot notation can be used to
show electron config for atoms &
molecules (& ionic cmpds) *fill in dots
in counterclockwise & in
singles
p.170
B. Lewis structures keep the dots to
show
unshared e- & use a line
to show a shared
pair of electrons
1. structural formula - shows only
the shared
pair of e- (p.171)
Lewis Structures -- atoms:
Sodium: Na.
Sodium ion: Na+
Phosphorus:
Bromine:
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Ionic Compound Lewis Structure
Examples:

Bromide ion: We need the brackets to
show that the bromide ion "owns" all
of the electrons rather than sharing
them. This is particularly important
when we make an ionic compound
such as sodium chloride.
Sodium chloride: NaCl
bromide: KBr
Potassium
Aluminum chloride: AlCl3
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Covalent Compound Lewis Structure
Examples:
Water H2O
NH3
Ammonia
Ammonium ion NH4+
CH4 Methane
Hydrogen sulfide
H2S
CO2 Carbon dioxide
Carbon monoxide
Choosing a Bond Type - Lewis
Structure Examples:
Barium iodide check
electronegativity difference
….formula?
According to electronegativity
dif. It should be polar
covalent…..it’s not, it’s ionic!
ionic: pos cation goes first
BaI2
nonmetal to nonmetal =
covalent!
nonmetal to metal = ionic!
Carbon disulfide check
electronegativity
difference….formula?
covalent: less electronegative
element goes first CS2
Arsenic triiodide check
electronegativity
difference….formula?
covalent: AsI3
Hydrogen selenide check
electronegativity
difference….formula?
covalent: H2Se
Oxygen difluoride check
electronegativity
difference….formula?
covalent: OF2
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Multiple Bonds - Examples
Carbonate ion : CO32-
However, other equally symmetrical
structures are possible,
so:
This is called resonance
IV. Molecular Geometry
A. VSEPR theory– “vesper”stands for:
Valence –
Shell—Electron –
Pair – Repulsion Theory
1. model for determining molecular
geometry which
relies on one simple
principle: electrons repel one
another.
2. in every molecular structure the
electron pairs on the
central
atom, whether shared in bonds or
unshared
("lone pairs"), will
distribute around the central
atom so as to be as far apart as
possible.
It's important to keep in mind that it is
the electron pairs on the central atom
that are doing the repelling and
generating the shape, not the electrons
on the atoms attached to the center.
3. How to find geometry:
a) Draw the Lewis structure
for the
molecule.
b) Count the number of
atoms that are
“stuck”
to it. (don’t count the number of
bonds!)
c) Count lone pairs of
electrons that are
“stuck” to it. (don’t count the
number of
bonds!)
IMPORTANT: This does
NOT mean to count the number of
lone
pairs on all of the
atoms in the molecule.
Lone
pairs on other atoms aren't important
what's important is only
what's directly
stuck to the
atom you're interested in.