Lewis Structure

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Theory

Molecules: Group or cluster of atoms of same or different elements which together behave as single unit and have characteristic properties.

Chemical Bond: The attractive force which holds various atoms or ions together in different species.

Questions arises:

Why do some atoms combine while some do not?
Why hydrogen molecule exists as H₂ and not H₃ or H₄?
Why do molecules have definite shapes?
Nature of force existing between the combining atoms?

These questions were tried to be answered by various theories. One was Lewis approach to chemical bonding.

Lewis Approach to Chemical Bonding:

In 1916, Kossel and Lewis gave some explanation for the chemical bonding. According to them,

Outer shell of an atom has maximum accommodation capacity of 8 electrons.
These electrons occupy the corners of a cube. For example, magnesium has 2 valence electrons, so it means that two corners of the cube are occupied.
When chemical bonds are formed, they try to achieve the stable octet by having 8 electrons.
That is why, noble gases are inert, because they already have 8 electrons in the corners of the cube.
Chemical combination occurs either by transfer of electrons or mutual sharing of electrons.

Lewis Symbols:

Valence electrons: The outershell electrons of the atoms which are involved in the chemical bond formation in a molecule.

Lewis introduced the simple notations to represent the valence electrons in an atom called as Lewis symbols or electron dot symbols.

Significance: Common valency of the element is either equal to the number of dots or 8 minus number of dots.

According to Kossel,

Electropositive atom loses electron to form positively charged ion (Cation) to attain noble gas configuration.

$Na â†’ {Na}^{+} + e^{-}$

Electronegative atom gains electron to form negatively charged ion (Anion) to attain noble gas configuration.

$Cl + e^{-} â†’ {Cl}^{-}$

The negative and positive ions are stabilised by electrostatic attraction which results in formation of electrovalent bond.

${Na}^{+} + {Cl}^{-} â†’ NaCl$

Lewis Structures:

The following basic steps are generally followed for writing Lewis dot structures of molecules and ions:

1. Calculate the total number of electrons (T) by assuming the octet (duplet for H atom) of each combining atom.

For example: for CO, T = 8+8 = 16

for ${CO}_{3}^{2 -}$ , T = 8 + (3x8) = 32

for ${NH}_{4}^{+}$ , T = 8 + (4x2) = 16

2. Calculate the total number of valence electrons (V) of all the combining atoms.

for CO, V = 4 + 6 = 10

While calculating V for polyatomic anions, add the negative charge to V.

for ${CO}_{3}^{2 -}$ , V = 4 + (3x6) + 2 = 24

While calculating V for polyatomic cations, subtract the positive charge from V.

for ${NH}_{4}^{+}$ , V = 5 + (4x1) - 1 = 8

3. Calculate number of shared electrons (S) by using following formula.

S = T - V

for CO, S = 16 - 10 = 6

for ${CO}_{3}^{2 -}$ , S = 32 - 24 = 8

for ${NH}_{4}^{+}$ , S = 16 - 8 = 8

4. Calculate number of unshared electrons (U) by using following formula.

U = V - S

for CO, U = 10 - 6 = 4

for ${CO}_{3}^{2 -}$ , U = 24 - 8 = 16

for ${NH}_{4}^{+}$ , U = 8 - 8 = 0

5. Write skeleton structure by placing the least electronegative atom in the centre and more electronegative atoms on the terminal positions.

6. Basic requirement is that each bonded atom gets an octet of electrons.

Formal Charge:

In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom.

Formal charge of an atom is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

In simple words, it can be calculated as:

$F o r m a l c h a r g e = V - U - \frac{S}{2}$

For example,

FC_P = 5 - 0 â€“ (8/2) = +1

FC_O = 6 - 4 â€“ (4/2) = 0

FC_O = 6 - 6 â€“ (2/2) = -1

Advantages of Formal charge:

Formal charges help in selection of the lowest energy structure from a number of Lewis structures for a given molecule/ion.
The structure having the smallest formal charge on its atoms has the lowest energy and thus most stable.

2nd structure is more stable.

Limitations of Octet Rule:

Formation of compounds with electron deficient atoms: The atoms having less than 4 valence electrons should try to lose electrons and form ionic bonds. But there are some examples available where these type of elements do form covalent compounds and do not even complete their octets.
Formation of super octet molecules: In some compounds, there are more than eight shared number of electrons around the central atom. This means that, they have more than 8 electrons around them. This is called as super octet or expanded octet.
Formation of compounds of xenon: According to octet rule, noble gases are inert, but Kr and Xe form compounds with F and O.
Inability to predict energy changes during bond formation.
This theory does not tell about the shape of molecules.