Lewis Lab
Introduction
In this exercise you will construct Lewis structures.
The existence of chemical compounds with fixed composition implies that the atoms in
compounds must be connected in characteristic patterns. Early models showed the atoms hooked
together like links on a chain. Modern representations are a good deal more abstract and often
mathematical in nature. Nevertheless, it is possible to represent molecular structures by using
Lewis structures.
The chemical bonds that hold atoms together in molecules generally consist of pairs of electrons
shared between two atoms. Atoms tend to share outer electrons in such a way that each atom in
the union has a share in an octet of electrons in its outermost shell. This generalization has come
to be known as the octet rule. The location of each element in the Periodic Table provides
information about the number of electrons in the outermost level of the atoms. Carbon, for
example, is in Group 4A and has four outer electrons; thus, it must share four additional
electrons from other atoms in order to achieve a share in eight outer electrons (an octet). This is
summarized in the table below. Oxygen, in Group 6A, has six outer electrons and shares two
electrons from other atoms in order to achieve an octet. Hydrogen is a special case, needing to
share its one electron with only one electron from anther atom in order to achieve the stable outer
electron configuration of the non-reactive element (He).
Atom
Outer Electrons
Carbon
Nitrogen
Fluorine & Chlorine
Hydrogen
4
5
7
1
Electrons needed
from another atom
4
3
1
1
Electrons shared
(no. of bonds)
4
3
1
1
A single bond consists of one shared pair of electrons; a double bond is two shared pairs (i.e. 4
electrons total), and a triple bond is three shared pairs (6 electrons). On paper, the bonds are
represented by single, double, or triple lines. Electrons not involved in bonding are termed
unshared electrons.
Writing Lewis Structures
The Lewis structure of a diatomic molecule is obtained by writing the chemical symbols for the
two atoms and deciding (from their locations in the periodic table) how many valence electrons
each supplies. The sum of these numbers in the number of dots that must appear in the Lewis
structure. In the case of carbon monoxide, for instance, the C atom has four valence electrons
and the O atom has six, so ten dots must appear. Moreover, the dots must be added to the
structure in pairs so that each atom (other than hydrogen) ends up with an octet and all the dots
are used.
The same ideas apply to polyatomic molecules. Each atom of a polyatomic molecule completes
its octet by sharing pairs of electrons with its immediate neighbors. Each pair of electrons shared
by two neighbors is a covalent bond, just as in a diatomic molecule. Two simple examples are as
follows:
Example 1: CH4
H
|
H-C-H
|
H
Example 2: NH3
H
:N - H
H
Procedure for Writing Lewis Structures
The procedure for writing Lewis structures for polyatomic molecules is the same as for diatomic
molecules. We use all the dots representing the valence electrons, and we arrange them so that
each atom has an octet of electrons. The only complication is that we now need to know which
atoms are linked to which other atoms. For instance, we need to know that the arrangement of
atoms in carbon dioxide is OCO and not COO. As in this case, one clue is that the less
electronegative atom is often the central atom. Another clue is that chemical formulas,
especially simple ones, are often written with the central atom first, followed by the terminal
atoms (the atoms attached to the central atom). An exception is the convention of writing acid
formulas with hydrogen at the front. If the species is an oxoacid, those hydrogen atoms would
be attached to oxygen atoms, which in turn are attach to the central atom. This is the case in
H2SO4, which has the structure (HO)2SO2.
Expanded Octets
The octet rule is based on the idea that eight electons fill a valence shell consisting of one s
orbital and three p orbitals. However, if an atom has empty d orbitals available, it may be able to
use them to accommodate more that eight electrons and hence to “expand its octet” to 10, 12, or
even more electrons. That may allow the central atom of a molecule to form additional multiple
bonds to the atoms attached to it, or to form bonds to more atoms. This is the case for the Period
3 elements of the p block, for which the 3d orbitals are only slightly higher in energy than the 3s
and 3p orbitals; it is also true of the elements in later periods. However, because 2d orbitals do
not exist, Period 2 elements cannot expand their octets. Another factor – and possibly the main
factor – in determining whether more atoms can bond to a central atom than are allowed by the
octet rule is the size of the central atom. A P atom is big enough for five Cl atoms to fit around
it. An N atom is too small, so NCl5 is unknown.
Incomplete Octets
Some elements at the left of the p block, must notably boron, form compounds in which their
atoms have incomplete octets. One example is the colorless gas boron trifluoride. The central
boron atom has only six electrons.
Odd-electron molecules
Some molecules end up with an odd number of valence electrons. In such molecules, at least
one atom cannot have an octet. This is the case with nitric oxide. Since the N atom supplies 5
valence electrons and the O atom supplies 6, the NO molecule contains 11 valence electrons. In
this example, the odd electron (unpaired electron) lies on the nitrogen atom.
The procedure for writing Lewis structures for molecules and ions can be summarized as
follows:
Step1. Calculate the total number of electron dots to be used by adding the numbers of valence
electrons of all the atoms. Each hydrogen atom supplies on electron. Each representative-group
element supplies its group number of electrons. For a cation, subtract on dot for each positive
charge. For an anion, add one dot for each negative charge. Divide this number by 2 to obtain
the number of electron pairs.
Step2. Arrange the chemical symbols for each atom in the formula so that the terminal atoms
surround the central atom. Initial hydrogen atoms should be attached to any oxygen atoms
present, or if no oxygen is present, to the central atom.
Step3. Use electron pairs to form single bonds linking each atom to its neighbor. Then try to
place any remaining electron pairs around the atoms so as to complete the necessary octet. If
there are not enough electrons, use on or more of the lone pairs to form double or triple bonds to
the central atom. Terminal halogen atoms always form single bonds.
Step 4. Use the expanded octet, incomplete octets, and odd-numbered octets where necessary.
Resonance
In many cases it is possible to write several different Lewis structures for the same molecule.
For example, three possible Lewis structures are possible for the nitrate ion (NO3-). Each Lewis
diagram is as equally correct as the others, however the correct structure is a blend of all three
structures – the resonance hybrid.
Lewis Lab
Procedure
Draw the Lewis structures (in Observations) for each of the following compounds:
O2, N2, H2O, CO2, O3, CF2Cl2, CHF2Cl, SO2, CO, NO, NO2, SF4, SF6.
Molecule
O2
N2
H2O
CO2
Total
outer
electrons
Lewis structure (s)
O3
CF2Cl2
CHF2Cl
SO2
CO
NO
NO2
SF4
SF6