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Lewis Dot Structures

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Alex Martiniouk

on 13 January 2014

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Transcript of Lewis Dot Structures

AP Chemistry
Lewis Dot Structures
Octet Rule
Odd Number of Valence Electrons
Exceptions to the Octet Rule
Example: Benzene
Charged Ions have either
more or less valence electrons
than the elements that constitute them (+1, +2...=less e-, -1, -2...=more e-)
They follow the octet rule
- the added/removed electron usually helps the ion achieve octets.
Don't forget to draw a
bracket with the charge
of the ion around the structure!
Charged Ions
Example №1: Cyanide
More than one valid structure
for a molecule exists.
Not one structure is a correct representation of the molecule
, but rather, the molecule is the average of its contributing structures.
For example, NO3- (nitrate ion) can be depicted with
one double bond and two single bonds
. But the actual molecule contains only
one type of bond
, a
of the single and double bonds (the pi-electron(s) appear to be delocalized)
Alex Martiniouk & Sophey Ho
Duet Rule
But that's just the tip of the iceberg!
Lewis structures show how the
valence electrons
are arranged among the atoms in a molecule.
To create a
molecule, elements in the
2nd period
share electrons to have
eight total
valence electrons. For elements in the 3rd period and beyond, the octet rule should be satisfied before atoms are placed on elements with available
The pairs of electrons not involved in bonding are
lone pairs
forms stable molecules when it shares two electrons because it has a
filled valence shell
does not form bonds
because it already has a filled valence shell.
Example №2: Nitrosonium (NO+)
Odd molecules (coined by G.N. Lewis) are molecules that contain an odd number of valence electrons, and thus,
all of the atoms will not be able to form octets
So what do we do???
To draw the structure of an odd molecule, the
formal charge
must be minimized
for each atom in the molecule (the formal charge can be negative on the most electronegative atoms).
But what is
Formal Charge
is the
assigned charge of an atom in a molecule
which can be used to to determine the
efficiency of electron distribution
(useful for drawing structures of odd molecules) *
FC disregards relative electronegativity
* (slight charges are not accounted for). FC is calculated for each atom separately.

you are comparing the the valence electrons of the non-bonded atom to the valence electrons assigned to the atom in the molecule.
The sum of all formal charges in a molecule will also equal its overall charge.
The equation:
FC = V - (N + B/2)
, where
is the number of valence electrons in the isolated atom (the non-bonded element),
is the number of valance electrons the atom does not share, and
is the number of electrons present in covalent bonds (divide by two because each bond has two electrons).
^(Nitric Oxide)
Which one is correct?
(N has octet)
(O has octet)
(N has octet)
Fewer than 8 valence electrons
Exceptions to the Octet Rule
Example: Boron Trifluoride
Boron and Beryllium
(second period) tend to form molecules where they are electron-deficient (having less than 8e- in valence shell), which makes them
very reactive
, especially with molecules with lone pairs.
But, BF3 is also
violently reactive
with molecules containing lone pairs, behaving
More than 8 valence electrons
Exceptions to the Octet Rule
Example: Triiodide (I3-)
This is present in molecules with elements of
period 3 and beyond
, as their valence shells are not limited to eight electrons (the s and p orbitals).
Sulfur Hexaflouride
Triiodide has 3(7)-1=
This means that even with the least amount of shared electrons (single bonds),
the octet rule will be violated
This leaves 22-4=18 electrons, meaning 1 atom will have to exceed the octet. But I3- has
3 atoms with empty 3d orbitals
. In this case, we
assume the central atom exceeds the octet.
Sulfur atom:
proposed by Kekulé
resonance hybrid
also has
lower energy
than any of the
contributing structures
Can also be denoted with the
resonance hybrid
Thiele used a dashed-line circle
Partial pi-bonds can also be indicated with dashed or curved lines.
Example: CO2
Example: H2O
Step 1: Sum up the valence electrons in a molecule. Don't worry about keeping track of which electrons come from which atoms! Only the

is important.
Step 2: Using a pair of electrons per bond, draw the bond between
each pair
of bound atoms.
Step 3: Distribute remaining electrons to achieve noble gas configuration, satisfying the octet and duet rules.
Example: HF
Example: NH3
The most important requirement for the formation of a stable compound is that the atoms achieve
noble gas configuration.
In this case, neon is a
noble gas
, and does not form bonds. Note that
only its valence electrons
(2s22p6) are represented by the Lewis structure. The 1s2 electrons are
core electrons
and are not shown.
*Note that the bonds between atoms is indicated by a line, rather than by the separate electrons. This is the standard notation.
(Few are actually stable, becomes more common with atoms of larger periods.)
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