Chemical Education Journal (CEJ), Vol. 13, No. 2/Registration No. 13-19 /Received March 2, 2010.
URL = http://chem.sci.utsunomiya-u.ac.jp/cejrnlE.html

Bonding and Structure - Empirical Rules
for Writing Lewis Structure

T. Parthasarathy

E-mail: parthasarathytrediff.com

Introduction

Formal Charge

Hyper and Hypo Valency Structure

Guideline for identifying the Central Atom

Monoplayer Molecules

Multiplayer Molecules

Empirical Rules

References

Introduction

Chemistry is an experimental science. Correlation of a range of experimental observation is possible with a number of qualitative concepts that are supported by appropriate theories. Structural formulae showing the arrangement of chemical bonds play a central role in such an approach and it is useful to establish the concepts involved.

The earliest description of the formation of a covalent bond is that of a shared electron pair. Lewis structures depict schematically how such pair are shared and give a topological picture of bonding in a molecule. Lewis concept [1] is the very direct interpretation of a bond in term of electrons, and even in modern discussion, it is quite often helpful to make use of Lewis's model.

Lewis structure is a preliminary step to write a most plausible molecular structure for a given molecular formula, before confirming through experimental means. Lewis theory of covalent bond formation is used to tackle this problem. The essence of this theory is to picturise covalent bond as (i) two-centre two electron bond and (ii) atoms with octet electronic structure.

The most plausible Lewis structure should be able to explain its actual shape and geometry when hybridization and valence shell electron pair repulsive model (VSEPR) concepts are applied. It should also explain bond parameters in a qualitative way from the resonance concept. Then it can be said that the method adopted in Lewis structures is acceptable. Although quite a few articles on Lewis structures have appeared in literature [2-9], a comprehensive method is yet cited. Writing Lewis structures becomes increasingly difficult as the number of atoms in a molecule or molecule ion increases. The topic dealing with Lewis theory is not vividly and elaborately presented in most of the general chemistry text-books and too much emphasis is often laid on the original concept of octet rule. Also modern Lewis concept of hyper and hypo valency is not emphasized. A comprehensive treatise for dealing with Lewis structure has not been dealt with in textbooks. As a result, many students of chemistry resort to a hit and trial method to write Lewis structure. Hence, the need for a student - friendly approach has arisen.

This paper attempts to overcome this difficulty of writing a Lewis structure firstly to select a central atom and then apply the given empirical rules. The modern theories of Lewis structure uses the concept of formal charge (FC).

Formal Charge is the residual charge on an atom in a Lewis structure. It may be zero, positive or negative. FC depends on number of valence electrons of an atom present in the Lewis structure.

FC = n_VE - n_LP - n_LP

Where n_LP = number of valence electron; n_LP = number of lone pair electrons;
nBP = number of bonded pair electrons.
FC is helpful in choosing between more than one possible Lewis structures. It is generally the case that the lowest energy structure is the one with the smallest formal charge on the atoms (and which usually lies between-1 and +1). For example, for Formaldehyde, CH2O the following two possible structures can be written as:

1(a)	1(b)

In both structures, each H atom has a FC of zero. In structure (a), the O atom has a FC of 6 - 2 - 1/2 (6) = + 1 and C atom has a FC of 4 - 2 -1/2 (6) = - 1

In structure (b), the FCs on O and C are 6 - 4 - 1/2 (4) = 0, and 4 - 0 -1/2 (8) = 0, respectively. Therefore, structure (b) is most plausible.
If FCs are inevitable, in all of the plausible Lewis structures, those with minimum FCs are to be considered. Generally negative FC should reside on more electro-negative atom and positive FC should reside on less electronegative atom. Sum of all FCs is equal to zero for neutral molecule. The charge of a molecular ion is equal to the sum of FCs.

Hyper and Hypo Valency Structure

Second period element like C, N, O and F obey the octet rule quite well. However, third period and subsequent elements show deviations from it. For example PCl5, SF6, IF7 in which, the central atom has 10, 12, 14 electrons respectively. These are called hyper valent compounds.
The occurrence of hyper valence is widespread for the elements of periods 3 to 6.

Hyper valency is due to availability of low lying unfilled "d" orbitals, which can accommodate the additional electrons.
Cases of hypo valency central atom (<8 electrons) are also known. For example BeX2, BX3, AlX3, GaX3, InX3, (X = halogen) in which the central atom has 4, 6, 6, 6, 6, respectively. Hypo valency is the feature of Be, B and 13 and 14 (Sn and Pb) group elements.

Guideline for identifying the Central Atom

Molecules may be identified as mono- and multiplayer molecules. Monoplayer molecules are those in which all other atoms around the central atom are bonded. Multiplayer molecules are those in which more than one central atom is present. The central atom of monoplayer molecules can be identified based on electro negativity, size, higher oxidation state, presence of "d"-orbitals and period it belongs to [10, 11]. It is difficult to identify the central atom for complex (multiplayer) molecules. This task of identifying the central atom can be overcome from the following guidelines.

Monoplayer Molecules

1. A single atom of a polyatomic molecule or ion is considered as the central atom.
   Example: F2O, Cl2O, Br2O, BeCl2, BCl3, BH3, BH4^-, CO3^2-, NO3^-, PO4^3-, SO3^2-, SF6, SiF6^2-, SO2, CCl4, SO3, IF7, NH3, PCl5, BrF5, XeF6, H2O, SO4^2-, ClF3, CrO3^2-.
2. A) However, in some exceptional cases like N2O, S2O, oxygen is not the central atom. In these cases, FC condition is not satisfied with 'O' as the central atom.
   B) Hydrogen and fluorine never act as central atoms.
   Example: N3H, NSF
3. If two or three odd-elements atoms are present in a molecule, atom with larger size, low electronegativity, higher oxidation state, presence of "d"-orbital, element third and higher period take precedence over other atoms.
   Example: SOCl2, POCl3, PNCl3, XeOF4, HCN, NOCl, NH2Cl, SOF4, CNO^-, CNS^-.
4. The central atoms of certain molecules can be identified only based on their root names. In fact, root manes were given based on experimental evidences.
   Example:

   Thiazylfluoride	NSF
   NitrosylChloride	NOCl
   Nitrylchloride	NO2Cl
   Cyanate	(OCN)-
   Isocyanate	(NCO)-
   Thiocyanate	(SCN)-
   Thiosulphate	(S2O32- = SSO32-)

5. If the compound is an acid, O, H atoms can be written in the form of OH group.
   Example:
   

In selecting the central atom, the minimal FC rule is to be kept in mind mandatorily.

Multiplayer Molecules

The problem in identifying central atom is easy if multiplayer molecule is written in terms of monolayer fragments.

(1) Split the molecule into equal halves; identify the central atom in each fragment from the rules mentioned above.

H2O2	HO - OH		N2H4	H2N - NH2
S2Cl2	ClS - SCl		S2O42-	-O2S - SO2-
S2O62-	-O3S - SO3-		Se2F6	F3Se - SeF3
N2O4	O2N - NO2		P2O64-	2-O3P - PO32-
Cl2O4	O2Cl - ClO2		I2Cl6	Cl3I - ICl3
H2S2O6	HO3S - SO3H		C2H4	H2C = CH2
C2H6	H3C - CH3		C2H2

(2) If the molecule cannot be split into equal halves, it can be split into equal halves by placing an/all extra atom/s between two equal fragments.

N2O5	O2N - O - NO2		P2O5	O2P - O - PO2
I2O5	O2I - O- IO2		Cl2O7	O3Cl -O-ClO3
P2O74-	2-O3P - O - PO32-		Cr2O72-	-O3Cr - O - CrO3-
H2S2O7	HO3S - O - SO3H		S3O62-	-O3S - S - SO3-

(3) If it is an acid and its basicity is known, O, H atoms can be written in the form of - OH group

H3PO4	(HO)3 PO		H3PO3	(HO)2 H PO
H3PO2	(HO)2 PH		H2SO4	(HO)2 SO2
H3BO3	(HO)3 B		H4SiO4	(HO)4 Si
HClO4	HO ClO3		H2CO3	(HO)2 CO
H2S2O7	HO-SO2 - O - SO2-OH

(4) In peroxy and disulfide compound O - O and S - S single bonds are present

Hydrogen peroxide	H2O2	HO - OH
Hydrogen disulfide	H2S2	HS - SH
Peroxysulphuric acid	H2SO5	(HO) (HO-O) SO2
Peroxydisulphuric acid	H2S2O8	O - O (SO2-OH)2
	HO- SO2 - O - O - SO2 -OH
Peroxydisulphuryl difluoride	S2O6F2	FO2S - O - O - SO2F
Peroxyphosphoric acid	H3PO5	(HO)2 PO (O-OH)
Peroxydiphosphoric acid	H4P2O8	O - O (PO (OH)2)2
	(HO)2 (O) P - O - O - P (O) (OH)2
Bis(pentafluorosulphur)	S2F10	F5S - SF5

(5) If the root nomenclature of the molecule is mentioned, its central atom can be easily identified based on its root structure.

Aminosulphuric acid	NH2SO3H	(HO) SO2 (NH2)
Chlorine perchlorate	Cl2O4	Cl ClO4
Thiosulphate	S2O32-	S SO32-
Chlorosulphuric acid	HSO3Cl	(HO) SO2 (Cl)
Dithionate	S2O62-	-O3S - SO3-
Tetrathionate	S4O62-	-O3S - S2 -SO3-
Cyanamide	CNNH2	NC - NH2
Methylsulphuric acid	CH3SO3H	(CH3) SO2-OH

Empirical Rules

1. Count the valence electrons of all the atoms in given molecular formula.
2. Identify the central atom as per the guidelines.
3. Connect the other atoms to the central atom through single bonds.
4. Electrons are placed on the peripheral atom in the form of pairs until they attain octet (hydrogen case duplet) state.
5. Remaining electrons are to be placed on the central atom in the form of pair of electrons or lone electron (if one is left).
6. In order to satisfy the octet state of central atom convert lone pairs present on peripheral into bond pairs.
7. The elements that belong to third and higher period when designated as central atom may get more than eight electrons to satisfy minimal FC conditions.
8. The structure of molecules involving second period element cannot show more than eight electrons.
9. The structure with either zero or minimal FC, will be the most plausible structure.

Lewis structures for few case examples arrived at, using above guidelines and empirical rules are presented below.

The objective of this article to help the student overcome the difficulty in writing the Lewis structure for octet, hyper and hypo valency compounds is hopefully achieved.

References

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5. Carroll, J. A., J. Chem. Educ.,1986, 63, 28 - 30.
6. DeKock, R. L., J. Chem. Educ., 1987, 64, 934 - 941.
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8. Malerich, C. J., J. Chem. Educ., 1987, 64, 403.
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11. Ahmad, W. Y. and Zakaria, M. B., J. Chem. Educ., 2000, 77, 329 - 331.