A bond is the sharing of 2 electrons.

Covalent bonds share electrons in order to form a stable octet around each atom in the molecules. Hydrogen is the exception it only requires 2 electrons (a duet) to be stable.

How do we draw a covalent Lewis Dot Structure?

Level 1 (basic)

1. Add up all the valance electrons of the atoms involved. ex CF₄

So C has 4 and F has 7 (x4 we have 4Fs) = 32 valence electrons

2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others. Never H.

So C will be surrounded by F's.

3. Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.

We have now placed 8 electrons as 4 bonds. We have 32-8= 24 more to place.

4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.

All 32 electrons are now in place, count the dots around each F. 6 dots and a bond (2 electrons) is 8. We have our octet.

The carbon has 4 bonds (2electrons) for its 8.

DONE

Level 2 (Double and Triple bonds)

Same rules apply until #4

1. Add up all the valance electrons of the atoms involved. ex CO₂

So C has 4 and O has 6 (x2 ) = 16 valence electrons

2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others. Never H.

So C will be surrounded by O's.

3. Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.

We have now placed 4 electrons as 2 bonds. We have 16-4=12 more to place.

4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.

All 16 electrons are now in place, count the dots around each O. 6 dots and a bond (2 electrons) is 8. We have our octet.

The carbon has 2 bonds (2electrons) for its 4....?

We need 8, so move a pair of electrons from the O to between the C and O. It will share 2 pairs of electrons instead of 1. It now has a double bond instead of a single bond.

now they all have an octet, it cleans up like this

Make it symmetrical.

Level 3-Lewis Dots of Polyatomic Ions

Same rules apply, at the end they get brackets and a charge

AP Chemistry and or College Level Rules

1. Determine whether the compound is covalent or ionic. If covalent, treat the entire molecule. If ionic, treat each ion separately. Compounds of low electronegativity metals with high electronegativity nonmetals (DE_N > 1.7) are ionic as are compounds of metals with polyatomic anions. For a monoatomic ion, the electronic configuration of the ion represents the correct Lewis structure. For compounds containing complex ions, you must learn to recognize the formulas of cations and anions.

2. Determine the total number of valence electrons available to the molecule or ion by:

(a) summing the valence electrons of all the atoms in the unit and
(b) adding one electron for each net negative charge or subtracting one electron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs (E.P.) available.

3. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom.

4. Determine a provisional electron distribution by arranging the electron pairs (E.P.) in the following manner until all available pairs have been distributed:

a) One pair between the central atom and each ligand atom.
b) Three more pairs on each outer atom (except hydrogen, which has no additional pairs), yielding 4 E.P. (i.e., an octet) around each ligand atom when the bonding pair is included in the count.
c) Remaining electron pairs (if any) on the central atom.

5. Calculate the formal charge (F) on the central atom.

a) Count the electrons shared as bonds. Total = b
b) Count the electrons owned as lone pairs. Total = n
c) F = V - (n + b/2), where V = number of valence electrons for the atom.

6. If the central atom formal charge is zero or is equal to the charge on the species, the provisional electron distribution from (4) is correct. Calculate the formal charge of the ligand atoms to complete the Lewis structure.

7. If the structure is not correct, calculate the formal charge on each of the ligand atoms. Then to obtain the correct structure, form a multiple bond by sharing an electron pair from the ligand atom that has the most negative formal charge.

a) For a central atom from the second (n = 2) row of the periodic table continue this process sequentially until the central atom has 4 E.P. (an octet).
b) For all other elements, continue this process sequentially until the formal charge on the central atom is reduced to zero or two double bonds are formed.

8. Recalculate the formal charge of each atom to complete the Lewis structure.

on to Formal Charge

Chemical Demonstration Videos

Click the Chemical Formula to see the Lewis Structure

Steps for Writing Lewis Structures

1. Find the total valence electrons for the molecule. Explain How Examples: H₂S, NCl₃, OH^-


2. Put the least electronegative atom in the center.
   Note: H always goes outside.
   Examples: NOCl, CF₂Cl₂, HCN


3. Put two electrons between atoms to form a chemical bond. Examples: CH₄, NH₃, I₂


4. Complete octets on outside atoms.
   Note: H only needs two valence electrons.


5. If central atom does not have an octet, move electrons from outer atoms to form double or triple bonds.
   Examples: O₂, N₂, C₂H₄



Advanced Steps

6. If you have extra electrons after the above steps add them to the central atom. Note: elements in the Period Three (usually S, P, or Xe) can have more than eight valence electrons.
   Examples: ClF₃, SF₄,XeH₄


7. Check the Formal Charges to make sure you have the best Lewis Structure. Explain How
   Examples: SO₄^2-, N₂O, XeO₃

Notable Exceptions to the Octet Rule

• H only needs 2 valence electrons.
• Be and B don’t need 8 valence electrons.
• S and P sometimes have more than 8 val. Electrons.
• Elements in Period Three, Four, etc (on the periodic table) can hold more than 8 valence electrons.