Drawing Lewis
Structures
Some issues about Lewis Structures to be discussed:
(1) Drawing “valid” Lewis structures which follow the “octet” rule
(holds almost without exception for first full row)
(2) Drawing structures with single, double and triple bonds
(3) Dealing with isomers (same composition, different constitution)
(4) Dealing with resonance structures (same constitution, different
bonding between atoms)
(5) Dealing with “formal” charges on atoms in Lewis structures
(6) Dealing with violations of the octet rule:
Molecules which possess an odd number of electrons
Molecules which are electron deficient
Molecules which are capable of making more than four covalent
bonds
The Lewis Model of Chemical Bonding
• In 1916 G. N. Lewis proposed that atoms
combine in order to achieve a more stable
electron configuration.
• Maximum stability results when an atom
is isoelectronic with a noble gas.
• An electron pair that is shared between
two atoms constitutes a covalent bond.
Covalent Bonding and Lewis Structures
(1) Lewis “dot” (electron) structures of valence electrons
for atoms
(2) Use of Periodic Table to determine the number of
“dots”
(3) Use of Lewis structures to describe the electronic
structures of atoms and molecules
(4) Works best for covalent bonds and for elements in the
first full row of the Periodic Table: H, He, Li, Be, B, C, N,
O, F, Ne
(5) Works with restrictions for second full row of the
Periodic Table and beyond: Na, Mg, Al, Si, P, S, Cl, Ar
Valence electrons for Elements
• Represent the number of valence electrons as dots
• Valence number is the same as the Periodic Table Group
Number
Lewis “dot-line” representations of atoms and
molecules
• Electrons of an atom are of two types: core electrons
and valence electrons
• The number of valence electrons is equal to the group
number of the element for the representative elements.
• For atoms the first four dots are displayed around the
four “sides” of the symbol for the atom.
• If there are more than four electrons, the dots are paired
with those already present until an octet is achieved.
• Ionic compounds are produced by complete transfer of
an electron from one atom to another.
• Covalent compounds are produced by sharing of one or
more pairs of electrons by two atoms.
Covalent bonding and Lewis structures
(1) Covalent bonds are formed from sharing of electrons by
two atoms.
(2) Molecules possess only covalent bonds.
(3) The bedrock rule for writing Lewis structures for the
first full row of the periodic table is the octet rule for C,
N, O and F: C, N, O and F atoms are always
surrounded by eight valence electrons.
(4) For hydrogen atoms, the doublet rule is applied: H
atoms are surrounded by two valence electrons.
Writing Lewis structures
• The skeletal structure of a polyatomic ion /
molecule indicates the order in which the atoms
are attached to one another
• It consists of one or more central atom(s) and at
least 2 terminal atoms
• A central atom is bonded to two / more atoms in
the structure
• A terminal atom is bonded to only one other atom
• In writing a skeletal structure the idea that every
atom must be connected to the rest of the
structure by at least one bond is applied.
Valence electrons and number of bonds
Number of bonds elements prefers depending on the number of valence electrons.
In general -

F am i l y
H al o g en s
F , B r, C l , I
C al co g en s
O,S
N i tro g en
N,P
C arb o n
C ,Si
# C o v al en t B o n d s*
O


N

3 bond often
C

4 bond
X
1 bond
often
2 bond often
always
The above chart is a guide on the number of bonds formed by these atoms.
Lewis Structure, Octet Rule Guidelines
When compounds are formed they tend to follow
the Octet Rule.
Octet Rule: Atoms will share electrons (e-) until it is
surrounded by eight valence electrons.
Rules of the (VSEPR) gamei) O.R. works mostly for second period elements.
Many exceptions especially with 3rd period elements (d-orbitals)
ii) H prefers 2 e- (electron deficient)
.
. N:
.
.
:O:
.
..
:F:
.
iii) :C:
4 unpaired 3unpaired 2unpaired
1unpaired up = unpaired e4 bonds
3 bonds
2 bonds
1 bond
O=C=O
NN
O=O
F-F
iv) H & F are terminal in the structural formula (Never central)
10
Lewis Structure Tutorial
10.7.00 6:16 PM
Atomic Connectivity
The atomic arrangement for a molecule is usually given.
HNO3
CH2ClF
Cl
H
C
O
F
N
CH3COOH
O
O
H
H
H C
H
H2Se
O
H
Se
C
H
H2SO4
O3
O
H
H O
S
O
O
O H
O H
O
In general when there is a single central atom in the molecule, CH2ClF, SeCl2, O3 (CO2,
NH3, PO43-), the central atom is the first atom in the chemical formula.
Except when the first atom in the chemical formula is Hydrogen (H) or fluorine (F). In
which case the central atom is the second atom in the chemical formula.
Find the central atom for the following:
1) H2O
a) H
b) O
2) PCl3
a) P
b) Cl
3) SO3
a) S
b) O
4) CO32-
a) C
b) O
5) BeH2
a) Be
b) H
6) IO3-
a) I
b) O
O
Note the following
• Hydrogen atoms are nearly always terminal atoms,
they form only one bond. (H has 2 electrons in
valence shell)
• In polyatomic molecules and / ions, the central
atom(s) usually have the lowest electronegativity
except for hydrogen (that is always terminal even
when bonded to a more electronegative atom)
• In oxo acids, hydrogen atoms are usually bonded to
oxygen atoms.
• With the major exception of carbon compounds in
which long chains of carbon atoms are common,
polyatomic molecules and ions usually form
compact structures
Steps for writing Lewis structures
1.Determine the total number of valence electrons.
– The total number of valence electrons for a molecule
is the sum of the valence electrons for each atom.
• N2O4 ----- (2 x 5) + (4 x 6) = 34 valence electrons
– For a polyatomic anion, which has one / more extra
electrons, add one electron for each unit of negative
charge
• NO3- ----- 5 + (3 x 6) + 1 = 24 valence electrons
– For a polyatomic anion, which is missing one / more
electrons, subtract one electron for each unit of the
positive charge
• NH4+ ----- 5 + (4 x 1) – 1 = 8 valence electrons
2. From the chemical formula, determine the
atom connectivity for the structure.
• Given a chemical formula, ABn, A is the central
atom and B flanks the A atom. i.e., NH3, NCl3,
NO2. In these examples, N is central in the
structure.
• H and F are never central atoms.
3. Write the skeletal structure and connect bonded
atoms with an electron – pair bond (dash)
4. Place electron pair around terminal atoms so
that each atom (except Hydrogen) has an octet)
5. Assign any remaining electrons as lone pairs
around the central atom(s).
6. If at this point a central atom has fewer than 8
electrons, a multiple bond(s) is likely. Move one
or more lone pairs from a terminal atom(s) to a
region between it and the central atom to form
a double or triple bond.