Drawing Lewis Structures of Ionic and Covalent Compounds

 

1.     Determine whether the compound is molecular or ionic. If the species to be depicted is an ionic compound, the cation and the anion should be drawn separately with the appropriate charges.

 

2.     Calculate the total number of electrons to be shown in the structure by

a.      adding the number of valence shell electrons for each atom.

b.     adding or subtracting electrons to account for the charge of the species if it is an ion.

 

3.     Calculating the probable number of bonds. A helpful (?) Rule of Thumb.

 

a.      Determine the number of electrons each atom would need to have around it in a “good” Lewis Structure (see the Table I below) and determine the sum for the whole molecule or ion.

 

b.     Subtract the number of electrons calculated above from the number calculated in Step 1 to determine the number of electrons to be shared in the structure. Divide this by 2 to calculate the number of bonds in the structure.

 

c.      Note that when this Rule of Thumb fails, the predicted number of bonds will be fewer than the number required to join the atoms! If this happens, connect the atoms with single bonds and go to Step 4.

 

4.     Connect the atoms with single or multiple bonds as lines or dot-pairs. (See Table I below.)

 

5.     Put in the remaining electrons as lines or dot-pairs around the atoms from most to least electronegative ,usually starting with the outside atoms, to fill their valence shells. Put any remaining electrons around the central atom.

 

6.     Calculate the formal charge of each atom from

 

            Formal Charge = (# of valence electrons in the neutral atom)

                                                             - (# of unshared electrons + 1/2 # shared electrons)

 

7.     Move lone pair electrons into multiple bonds with adjacent atoms to reduce formal charges and to make formal charge agree with electronegativity, if possible. Be sure not to exceed the limits of the valence shells. Note: Hydrogen can accommodate only two electrons in its valence shell (comprised only of a 1s orbital). The second period elements can accommodate eight (thus the “octet rule”) since the second shell is comprised of one 2s and three 2p orbitals. Elements beyond the second period may accommodate more than eight electrons since they do have d orbitals available for bonding. Boron and Aluminum often have only six electrons in a Lewis Structure.)

Table I : Rules of Thumb for Drawing Lewis Structures

 

Rule / Group

IA

IIA

IIIA

IVA

VA

VIA

VIIA

1)  Number of Electrons to Satisfy

2

4

6

8

8

8

8

2)  Number of Bonds Usually Formed

1

2

3

4

3

2

1

Notes:

1.     IA, IIA, and IIIA elements (other than hydrogen and boron) are usually involved in ionic or metallic bonding.

2.     IIIA elements can accommodate 8 electrons by forming 4 bonds.

3.     p-Block elements of the third period or higher can accommodate more than 8 electrons by forming more than 4 bonds.

4.     Hydrogen can accommodate only two electrons in its valence shell (comprised only of a 1s orbital). The second period elements can accommodate eight (thus the “octet rule”) since the second shell is comprised of one 2s and three 2p orbitals. Elements beyond the second period may accommodate more than eight electrons since they do have d orbitals available for bonding. Boron and Aluminum often have only six electrons in a Lewis Structure, but can accommodate eight.


Resonance Structures

 

Rules of Resonance:

1)         Move electrons only, not atoms. Except for “hyperconjugation” only p and non-bonding electrons should be moved.

2)         Move electrons (generally in pairs)

                        -from an atom into a bond with an adjacent atom or vice versa,

                        -from a bond to an adjacent bond,

                        -or (rarely invoked) from an atom to an adjacent atom.

3)         Do not exceed the valence capacity of an atom.

 

Judging the Contribution of Resonance Structures:

            A “best” or “most important” resonance structure should have

1)         The maximum number of paired electrons (i.e. the fewest unpaired)

2)         The maximum number of shared electrons (i.e. the greatest number of bonds)

3)         The fewest number of formal charges with the lowest absolute values.

4)         Formal charges of the same sign far apart and/or formal of opposite sign close together.

5)         Distribution of positive and negative charges should agree with electronegativity.

 

Consider OCN-

 


YT = a qA + b qB + c qC + d qD

 

Importance?