Modified Lewis Structures
Modified Lewis structures show both the total valence electrons in the molecule,
or ion, and its three-dimensional structure. Structures are derived by considering each
valence shell electron pair (or non-bonding unpaired electron) to be stereochemically
significant in that their arrangement will be such as to minimize repulsions between
them. Minimization of electron-electron pair repulsions dictates the following
arrangements for 2-6 electron pairs about a given atom: 2, linear; 3, trigonal planar; 4,
tetrahedral; 5, trigonal bipyramidal; 6, octahedral. Note that molecular geometries are
described by consideration of the arrangement of nuclei since the locations of electron
pairs cannot be experimentally determined.
In determining the number of valence electron pairs for central atoms, and thus
their geometrical arrangement, the following rules will be helpful.
1. Assign each terminal atom enough electrons such that it will have a filled
valence shell as a result of sharing an electron from the "central" atom. In the
case where hydrogen or halogen serve as terminal atoms no adjustments are
necessary. However, in the case of atoms such as oxygen, sulfur and nitrogen
this will require that they be considered as anions in combination with a cationic
central atom. That is, O1-, S1- and N2- can achieve an octet by sharing one
electron with another atom.
2. After making any adjustments of terminal atoms as described above, assign the
central atom a number of electrons and a formal charge such that the sum of
the charges on the central and terminal atoms is consistent with the overall
charge on the species whose structure is under consideration.
3. The total number of electrons around the central atom is determined by adding
the number of electrons assigned to it in step 2 together with the number of
electrons contributed by the terminal atoms (one electron each). The total
number of electrons around the central atom divided by 2 is the number of
electrons pairs that are stereochemically active and whose arrangement about
the central atom determine the geometry of the species.
Examples:
BF4 1
B1
4
4F @ 1 each
4
Total
F
8 e or 4 pairs
Conclusion: tetrahedral arrangement
F
B
F
1F
CO32
1-
C1+
3
3O @ 1 each
3
Total
6 e or 3 pairs
O
O
11+
1C O
1-
O
C O
O
1-
Conclusion: trigonal planar arrangement
After such structures have been drawn, if formal positive and negative charges
reside on adjacent atoms then these may be combined to create an additional (π) bond
between the two atoms, as shown above for carbonate. This may be done repeatedly
provided the total number of bonds about a second row element does not exceed
four, and the total number of bonds about a third (or higher) row element does
not exceed six. In a some cases the "best" structure still has charge separation, i.e.,
opposite charges on adjacent atoms; for example, nitrate ion.
1-
O
O
1+
N O
1-
Refinements of structures with respect to deviations from ideal angles can be made by
consideration of the spatial requirements of non-bonding electron pairs and the relative
energies of each type of electron pair repulsion. Since non-bonding electron pairs are
restricted to move in the vicinity of a single atom they are expected to be closer to the
nucleus and to effectively subtend a greater angle than a bonding electron pair which is
by necessity shared with another atom. Repulsions between non-bonding pairs (np) will
be greater than between non-bonding pairs and bonding pairs (bp), which in turn are
greater than interactions between bonding pairs, i.e., np-np > np-bp > bp-bp. Note that
π-bonding electron pairs are not generally considered as being stereochemically active
in the determination of gross geometry but they may cause specific distortions from
ideal geometries because of the much greater electron density between the multiply
bonded nuclei. In general, any structure having non-bonding electron pairs will have
bp-bp angles that are less than the ideal angles and where more than one non-bonding
pair is present they will be located in such a way to minimize the number of 90b np-np
and np-bp interactions.
In localized orbital bonding models it is frequently necessary to resort to
resonance structures to adequately describe the electronic structure of an ion or
molecule. The following rules are helpful in determining what resonance forms are
likely to be significant.
1. The relative positions of nuclei do not change.
2. All structures must have the same number of electron pairs.
3. Terminal atoms will generally adhere to the octet rule.
4. Structures with zero formal charge on all atoms will be most important. If,
however, formal charges are required, negative charge should be localized on
the most electronegative elements. Structures with large separation between
opposite charges are not likely to be important, nor are structures with like
charges on adjacent atoms.
1-
O
1+
O
N O
O
1-
O
11+
1N O
1-
O
O
1+ 1N O