CH225.6
Today:
•Octet rule and formal charges
• Lewis structure
• Resonance
For next class read Lecture 7
The first sentence in The Books of Bokonon is this:
“All of the true things I am about to tell you are shameless lies.”
Anyone unable to understand how a useful religion can be founded on lies will
not understand this book either.
(From Cat’s Cradle by one of the best American writers, Kurt Vonnegut)
To paraphrase Vonnegut: Anyone unable to understand how a useful theory can
have little truth in it will not understand Chemistry.
Chemical bonding according to Lewis
Chemistry begins when atoms and ions are joined together to form molecules and
molecular ions. First, we shall look at some simplistic ideas of chemical bonding.
In 1916, Lewis defined covalent bond as the interaction between two atoms held
together by two electrons. The electrons are shared by the atoms.
Example: In H2, proton-electron attractive interactions offset the repulsions between
the electrons and nuclei (protons). The net interaction is the H-H bond.
This is represented as H:H, or more simply H–H. This representation tells us that
there are two electrons located between the two hydrogen nuclei, and that these two
electrons form the bond holding the atoms.
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Octet rule and the number of bonds in a molecule
The octet rule is an expectation that the valence electrons in a molecule can be
arranged to give every p-block atom 8 e- in bonds & lone pairs. Superficially, this
arrangement makes all valence orbitals (s & p) fully occupied (termed a closed shell
electronic configuration).
The octet rule works most of the time for compounds of carbon, nitrogen, oxygen,
and fluorine. All other elements rarely obey the octet rule, unless the electrons are
deliberately artificially arranged, e.g. consider PF5 (20 valence electron pairs):
Trigonal bipyramidal
PF5 molecule with five
P-F bonds:
Unreasonable representation
of PF5 as a tetrahedral (sp3
hybridized) ionic compound:
Based on the octet rule, one can predict the number of bonds x in a main-group
compound :
where h is the number of hydrogen atoms, and p is the number of p-block atoms.
These atoms can accommodate as many as 2h + 8p electrons in the valence shells.
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Lewis structure
Definition from the IUPAC “Gold book”
Lewis structures are based on the octet rule and should be augmented by formal
charges when necessary. The formal charge on an atom is calculated as follows,
based on the number of bonds to the atom, it’s group number and the number of
non-bonding (i.e. lone-pair) electrons:
These ideas often make inorganic Lewis structures rather artificial constructs.
Selected examples of Lewis structures are considered on the following slides.
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Examples of Lewis structures
The following generalizations are helpful in constructing Lewis structures:
1. Hydrogen atoms are often (but not always) terminally located.
2. The least electronegative atom in a molecule often occupies the central position
and the more electronegative atoms occupy the outer (terminal) positions.
3. The best structure exhibits the least number of formal charges. Inorganic chemists
prefer drawing structures without formal charges, when possible.
Example: F2
x = (2 x 8 - 14)/2 = 1 bond
Formal Charge = 7 - 6 - 1 = 0
The resulting molecule, written F-F, demonstrates an important point. It has six pairs
of electrons which are not shared by the atoms. Such pairs of electrons are called
non-bonding electrons or lone pairs.
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Examples of Lewis structures
Example: O2
x = (2 x 8 - 12)/2 = 2 bonds; Formal Charge = 6 - 4 - 2 = 0
Oxygen needs two electrons to attain the octet. This places two shared pairs
between the oxygen atoms. Two pairs of electrons between two atoms constitute a
double bond. The bond order in the molecule is 2, and O2 is written as O=O. Note
that the Lewis structure of O2 is actually wrong; the molecule is a diradical with two
unpaired electrons.
Example: N2
x = (2 x 8 - 10)/2 = 3 bonds; Formal Charge = 5 - 2 - 3 = 0
Dinitrogen is written as NN to show the triple bond, and the bond order is 3.
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Examples of Lewis structures
Example: NF3, nitrogen trifluoride
x = (4 x 8 - 26)/2 = 3 bonds
Nitrogen Formal Charge = 5 - 2 - 3 = 0
Example: NO2–, nitrite ion
x = (3 x 8 - 18)/2 = 3 bonds
Nitrogen Formal Charge = 5 - 2 - 3 = 0
Example: NO2, nitrogen dioxide
This molecule violates the octet rule
Nitrogen Formal Charge = 5 - 1 - 3 = +1
Note how the structure appears in
Wikipedia, without the formal
charges which do not make any
practical sense:
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Problematic Lewis structures
Example: Carbon monoxide, CO
x = (2 x 8 - 10)/2 = 3 bonds;
Carbon Formal Charge = 4 - 2 - 3 = -1
Oxygen Formal Charge = 6 - 2 - 3 = +1
The bond order in the molecule is 3. This is in full agreement with the experimental
properties of carbon monoxide.
However, the formal charges placed on this molecule by following the rules of writing
Lewis structures are very misleading. To avoid a confusion, they are usually omitted
in the literature. E.g., this is how the molecule appears in Wikipedia:
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Problematic Lewis structures
Example: ClO4–, perchlorate ion
chlorine = 7 - 4 = +3
x = (5 x 8 - 32)/2 = 4 bonds; Formal Charge on
d(Cl-O) = 1.44 Å (experiment)
Predicted single Cl-O bond:
0.99 + 0.66 = 1.65 Å
A Lewis structure cannot reliably predict the bond order in compounds of the maingroup elements outside Period 2. There is a significant double-bond character in the
Cl-O bonds of ClO4−. Common representations of the structure:
From Wikipedia,
note the depiction
of the Cl-O bonds
as “partial double”:
The nature of bonding in Inorganic compounds is often based on experimental
information; a non-professional (e.g., a student) may find it non-trivial to predict.
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Problematic Lewis structures
Example: Ozone, O3
x = (3 x 8 - 18)/2 = 3 bonds;
Central Oxygen Formal Charge = 6 - 2 - 3 = +1
Singly-bonded Oxygen Formal Charge = 6 - 6 - 1 = -1
The formal charges placed on this molecule by following the rules of writing Lewis
structures are completely misleading. To avoid a confusion, this is how the molecule
appears in a popular Inorganic Chemistry textbook:
Note, however, the wrong Lewis structure of ozone in Wikipedia:
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Resonance
A single Lewis structure is often an inadequate description of a molecule. The Lewis
structure of ozone, for instance, suggests (incorrectly) that the two oxygen-oxygen
bonds are different. In fact, they are identical.
It is possible to draw two Lewis structures to represent ozone. These representations
are called resonance structures or resonance forms.
1.28 Ǻ
1.28 Ǻ
116º
It is technically impossible to describe ozone by a single Lewis structure. Resonance
Lewis structures are drawings (representations) of no real existence.
The above diagram must be read in the following way. Ozone has two identical
oxygen-oxygen bonds. They have properties intermediate between those of typical
single O-O (1.48 Å) and double O=O (1.21 Å) bonds. The experimentally determined
oxygen-oxygen bond length in ozone is 1.28 Å.
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Resonance
The two resonance forms of ozone are clearly degenerate. In other cases the
resonance structures may be non-degenerate. The cyanate ion, [CNO]–, is one such
example:
1.19 Å
Computed APT atomic charges:
-0.89
1.22 Å
+0.84
-0.95
The relative contributions (weight) of the resonance forms can be qualitatively
established based on electronegativity. In the case of [CNO]–, the structure has a
greater contribution of (it resembles more) the resonance form with a negative
charge on oxygen.
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Resonance
The use of resonance structures is widespread in Organic Chemistry. However, even
the classical examples of usage are probably not obvious or make 100% sense.
Example: Dimethylformamide, DMF
Lewis structure
Experimental structure
d(C-N) = 1.45 Å
Note: sp3 hybridized nitrogen is
expected to be pyramidal.
Planar (flat) structure
d(C-N) = 1.34 Å
a partial double bond
There is said to be a 2nd resonance structure for DMF
that explains the partial C-N double bond. Yet, it is not
obvious (a) why this should make a significant
contribution, and (b) it does not explain why the
nitrogen fragment is planar (perfectly flat) in DMF, i.e.
the real structure is not a ‘resonance hybrid’.
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Resonance
Some molecules cannot be satisfactorily described even with the help of resonance
structures.
1.35 Å (a double bond)
Example: Butadiene
1.46 Å (a partial double bond)
Note that a single C-C bond is 1.53 – 1.54 Å
No reasonable second resonance structure can be proposed for butadiene.
Lewis structure
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Hypervalence
When the total number of electrons around a main-group atom exceeds 8, the
species is called hypervalent.
The term 'hypervalence' may suggest that compounds which have more than 8
electrons are unusual and are adopting an unconventional bonding mode. In fact,
there are so many examples of 'hypervalent' compounds of the main-group elements
from Periods 3 to 6 that this terminology is simply misleading.
Examples: PCl5 and SF6
Predicted bonds: (6x8-40)/2 = 4
10 e- around P
(7x8-48)/2 = 4
12 e- around S
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Conclusions
At the end, we are probably left with mixed feelings about Lewis structures.
We shall acknowledge that Lewis structures form the foundation of structural Organic
Chemistry and are widely used in this area, in combination with some additional
techniques, e.g. the ‘electron-pushing’. This is largely for the reason that organic
molecules are often made of the atoms that obey the octet rule (H, C, N, O).
Inorganic Chemistry is a bigger and considerably more complex subject.
Unfortunately, Lewis structures are not inorganic chemist’s ‘trusted friend’.
Figuratively speaking, they cheat on us, they lie to us, and make false promises. One
may ask, why we do not abandon the use of Lewis structures in Inorganic
Chemistry? There are two reasons. One the one hand, there is simply no better
alternative when you start ‘from scratch’. On the other hand, between all the lies,
Lewis structures too often tell us something valuable, give us useful practical hints.
Inorganic chemists use Lewis structures in combination with experimental structure
determination and quantum-chemical calculations.
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