Chemistry Lewis Structures

Study Notes

Determine the total number of valence electrons by adding the valence electrons of each atom in the molecule and accounting for any charge.
Draw a skeletal structure with the central atom (usually the least electronegative) surrounded by terminal atoms (e.g., hydrogen and fluorine).
Fill the octets of the terminal atoms (or add two electrons to hydrogen) using the remaining valence electrons.
Place any remaining electrons on the central atom, potentially leading to an expanded octet.
Calculate the formal charge (FC) of each atom using the formula: FC = (Valence electrons - ½ Bonding electrons - Nonbonding electrons).
Minimize formal charges by creating multiple bonds using nonbonding electrons, especially when neighbouring atoms have opposite charges.
Ensure all atoms have an allowed electron count, following the octet rule for C, N, O, and F, while considering the exceptions for elements like hydrogen, beryllium, and boron, as well as expanded octets for elements in periods 3 and beyond.

Electronegativity (EN) is an atom's ability to attract electrons in a bond.
EN increases across a period and up a group in the periodic table.
Fluorine has the highest EN (4.0) according to the Pauling scale.
The difference in EN (ΔEN) determines bond polarity:
- Large ΔEN (> 1.9) indicates ionic bonding, like in NaCl (ΔEN ≈ 2.23).
- Intermediate ΔEN (0.5 - 1.9) indicates polar covalent bonding, like in PCl5 (ΔEN ≈ 0.97).
- Small ΔEN (< 0.5) indicates pure covalent bonding, like in Cl2 (ΔEN ≈ 0).

Lewis structures show all bonding (b) and non-bonding (nb) electrons, as well as formal charges.
A complete shell is typically an octet, except for:
- Hydrogen: satisfied with 1 electron pair.
- Beryllium: satisfied with 2 electron pairs.
- Boron/Aluminum: satisfied with 3 electron pairs.
- Elements in periods 3 and beyond: can have expanded octets if they are the central atom.
Bonding electrons can be involved in single, double, or triple bonds.

BF3:
- Three fluorine atoms surround a central boron atom.
- All atoms have a formal charge of 0.
NO3−:
- Three oxygen atoms surround a central nitrogen atom.
- The formal charge is distributed as follows: one oxygen atom has a formal charge of -1, and the nitrogen atom has a formal charge of +1.
BrOF2+:
- Two fluorine atoms and one oxygen atom surround a central bromine atom.
- The formal charge is distributed as follows: the bromine atom has a formal charge of +1, and the oxygen atom has a formal charge of -1.

iClicker #1: The correct Lewis structure for chlorate anion (ClO3−) shows the chlorine atom with one double bond to an oxygen atom and two single bonds to the other two oxygen atoms. All oxygen atoms have a formal charge of -1, and the chlorine atom has a formal charge of +1.
iClicker #2: The correct Lewis structure for CO places a triple bond between the carbon and oxygen atoms, with one lone pair on carbon and two lone pairs on oxygen. This arrangement minimizes formal charges and satisfies the octet rule for both atoms.