# Lecture 14: Lewis Structures and Formal Changes

• Lewis Structure - Octate form-stable, inert because the shell is full combining capacities of atoms

• Exceptions
• ﻿$H$﻿ seeks a duet rather than an octet
• Electron deficient compounds do not have enough ﻿$e$﻿'s to satisfy the octet of all actions
• For resonance hybrids, more than one Lewis structure can be drawn to represent the molecule or complex ion
• Elements of the third row of the Periodic Table or beyond may be surrounded by more than 8 ﻿$e$﻿'s

### Steps for Drawing Lewis Structures

1. Draw the species with atoms connected by single 2﻿$e^-$﻿ bonds
2. All atoms connected to the central atom by single bonds that uses a pair of electron to form the bond
3. Count the total number of valence shell ﻿$e^-$﻿
4. It includes all the atoms, the central atom too
5. Subtract ﻿$e^-$﻿ in 1 from 2
6. You get total number of valence electrons and from this, subtract the electrons used up to form single bonds
7. Then use the rest to complete the octets, beginning with the more electronegative atoms
8. Remaining electrons are placed on the atoms beginning with the more electronegative atoms so that they get the octet first
9. Assign formal charges to all the atoms
10. Share lone pairs of ﻿$e^-$﻿ to offset the formal charge and complete the octets
11. Transform lone pairs into bond pairs in such a way that the formal charges are removed

### How to Assign Formal Charges

1. The least electronegative atom is the central atom
2. Count valence electrons, minus total number of non-bonding electrons - total number of bonding electrons
3. Calculate this for each atom and then conclude compounds formal charge (total)
4. Draw 2 (or more) possible resonance structures in brackets with an arrow between showing back and forth
• Negative ﻿$FC$﻿ 's should be placed on electronegative elements
• ﻿$FC$﻿ of the same sign should not be next to each other
• Number of ﻿$FC$﻿ should be minimized