Lewis Dot Structure Valence Calculator
Select Element
H – Hydrogen
C – Carbon
N – Nitrogen
O – Oxygen
F – Fluorine
P – Phosphorus
S – Sulfur
Cl – Chlorine
Br – Bromine
I – Iodine
Xe – Xenon
B – Boron
Be – Beryllium
Select Element
H – Hydrogen
C – Carbon
N – Nitrogen
O – Oxygen
F – Fluorine
P – Phosphorus
S – Sulfur
Cl – Chlorine
Br – Bromine
I – Iodine
Xe – Xenon
B – Boron
Select Element
H – Hydrogen
C – Carbon
N – Nitrogen
O – Oxygen
F – Fluorine
P – Phosphorus
S – Sulfur
Cl – Chlorine
Total Valence Electrons: 0
Total Electron Pairs: 0
Status: –
How to Determine a Lewis Dot Structure
Lewis dot structures (also known as electron dot structures) are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. They are essential for understanding chemical bonding and molecular geometry.
Step 1: Calculate Total Valence Electrons
Sum the valence electrons for every atom in the molecule. If the species is a polyatomic ion, add one electron for every negative charge and subtract one electron for every positive charge.
Example: CO2
Carbon (C) is in Group 14: 4 valence electrons.
Oxygen (O) is in Group 16: 6 valence electrons (x 2 atoms).
Total: 4 + (6 * 2) = 16 valence electrons.
Carbon (C) is in Group 14: 4 valence electrons.
Oxygen (O) is in Group 16: 6 valence electrons (x 2 atoms).
Total: 4 + (6 * 2) = 16 valence electrons.
Steps for Drawing the Structure
- Identify the Central Atom: Usually the least electronegative atom (never Hydrogen).
- Connect Atoms: Draw single bonds between the central atom and surrounding atoms. Each bond represents 2 electrons.
- Distribute Remaining Electrons: Complete the octets of the outer atoms first.
- Check the Octet Rule: If the central atom does not have an octet, move lone pairs from outer atoms to form double or triple bonds.
Example: Water (H2O)
1. Total Electrons: (1 * 2) + 6 = 8 electrons.
2. Central atom: O (Hydrogen cannot be central).
3. Single bonds to H: Uses 4 electrons (2 pairs).
4. Remaining 4 electrons are placed as lone pairs on Oxygen.
1. Total Electrons: (1 * 2) + 6 = 8 electrons.
2. Central atom: O (Hydrogen cannot be central).
3. Single bonds to H: Uses 4 electrons (2 pairs).
4. Remaining 4 electrons are placed as lone pairs on Oxygen.
Formal Charge and Exceptions
While the octet rule is a general guideline, there are exceptions. Elements in Period 3 and below (like Sulfur or Phosphorus) can have an expanded octet, holding more than 8 electrons. Boron and Beryllium often have incomplete octets.
The Formal Charge is used to determine the most stable Lewis structure when multiple possibilities exist. It is calculated as: [Valence Electrons] – [Non-bonding Electrons] – [1/2 * Bonding Electrons].