How to represent molecules using Lewis dot structures? (With Examples)

The Lewis dot structures are used to show the shared electron pairs between the bonded atoms in the molecule and the lone pairs of electrons if any. The representative structures follow the octet rule wherein the atoms combine by either transfer of electrons (loss or gain) or by sharing of valence electrons in a way that the valence shell of the atoms attains the octet configuration.

The electrons involved are shown as dots. A single bond is made up of two electrons, a double four and a triple bond six. Similarly, an atom can have one, two or three lone pairs of electrons.

The rules employed to draw a correct Lewis Structure are:

1) The least electronegative atom is the central atom.

Few Examples:

Molecule

Central Atom

CO2

BCl3

SO42-

C

B

S

 

2) The Hydrogen tries to attain the duplet and not the octet configuration. Therefore, it occupies the end or the terminal positions as shown in the images. 

Hydrogen Position in Lewis Acid Structures

3) The valence electrons of all the atoms in the neutral molecule are added up. If the molecule is an anion, the charge carried by the anion is added up as electrons. Similarly, for the cations, depending on the charge carried by the molecule, the equivalent number of the electrons are deducted from the final answer.

Example,

Molecules

Valence Electrons

Addition of the valence electrons

Neutral

SiCl4

SOCl2

H2O2

 

 

Si -4, Cl- 7

S - 6, O - 6, Cl - 7

H - 1, O - 6

 

4 x1+ 7 x 4 (for 4 Cl atoms) = 32

6x1 + 6x1 + 2 x 7= 26

2 x 1 + 6 x 2 = 14

 

Anion

CO32-

 

PO43-

 

HCO3-

 

 

 

C - 4, O – 6

 

P - 5, O – 6

 

H - 1, C - 4, O – 6

 

 

4 x 1+ 3 x 6 + 2 (for -2 charge) = 24

 

5 x 1 + 4 x 6 + 3 (for -3 charge)= 32

 

1 + 4 + 3 x 6 + 1 (for -1 charge) = 24

 

Cations

NH4+

 

 

 

N – 5, H - 1

 

 

5 x 1 + 4 x 1 – 1 (for +1 charge)= 8

4) Divide the number of electrons obtained from step 3 by 2 to get the electron pairs. It is easier to assign Lewis dots when the electrons are paired. Note that odd electrons disobey the octet rule.

For Example,

Molecules

Valence Electrons

Addition of the valence electrons

Valence Electrons Divided by 2 = Electron Pairs

Neutral

SiCl4

 

SOCl2

 

H2O2

 

 

Si -4, Cl- 7

 

S- 6, O - 6, Cl - 7

 

H - 1, O - 6

 

4+ 7 x 4 (4 Cl atoms) = 32

 

6 + 6 + 2 x 7= 26

 

2 x 1 + 6 x 2 = 14

 

32/2 = 16

 

26/2 = 13

 

14/2= 7

Anion

CO32-

 

PO43-

 

HCO3-

 

 

C - 4, O – 6

 

P - 5, O – 6

 

H - 1, C - 4, O – 6

 

4+ 3 x 6 + 2 (for 2 electrons) = 24

 

5 + 4 x 6 + 3 = 32

 

1 + 4 + 3 x 6 + 1 = 24

 

24/2= 12

 

32/2= 16

 

24/2= 12

Cations

NH4+

 

 

N – 5, H - 1

 

5 + 4 x 1 – 1 = 8

 

8/2 = 4

 

5) Keeping the least electronegative atom as the central atom, attach the atoms with each other using a single bond. One bond is one electron pair. Deduct the number of single bonds from the electron pairs.

The remaining unshared pair of the electrons are used as lone pairs (first satisfying the terminal atoms and then the central atom in a way that every atom has an octet configuration) or in multiple bonding.

Example, Neutral Molecules (Uncharged)

Lewis Acid Examples

Example, Anions

Example, Cations

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