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Formal Charge
Start the process by drawing possible Lewis
Structures.
- Choose the central atom. Usually the least electronegative.
- Count total valence electrons for the molecule.
- Start by placing one pair of electrons for each bond.
- Satisfy the octet rule. This may require making double or triple bonds (and removing lone pairs)
- If you can draw multiple structures, check the formal charge. See below.
Formal Charge
Formal charge is used to decide if a Lewis
Structure is plausible.
Formal Charge
Rules for formal charges:
- Formal charges in must add up to the charge of the molecule or ion.
- We generally choose the Lewis structure in which the atoms bear formal charges closest to zero
- We generally choose the Lewis structure in which any negative charges are on the more electronegative atoms. Note: formal charges do not represent actual charge. It is merely a tool that helps us determine plausible Lewis structures.
Resonance
Ozone: (O 3 )
Resonance Cont.
2. Resonance: All elements want an
octet, and we can do that in multiple
ways by moving the terminal atom's
electrons around (bonds too).
Assign formal charges
Find the most ideal resonance
structure.
It is the one with the least formal charges that adds up to zero or to the molecule's overall charge.) The most electronegative atom should have a negative charge and least electronegative should have positive charge
3 Exceptions to the Octet Rule
3. Molecules where an atom has more than
an octet of electrons (i.e. ClF 3 , PCl 5 ,
XeF 2 ).
SULFUR is most common exception for
more than an octet.
This is fairly common for elements in the 3rd period (row) and below. However, elements in the first two periods, H – Ne, cannot violate the octet rule in this way.
Determining Molecular Shape
Examples:
1. H 2 O
- Oxygen will be central. Assign it a number of 6
- |charge x quantity| of hydrogens = 2
- Not an ion so don’t need to add or subtract anything
- Total = 8
- divide by 2 = 4 (so for water there are 2 bonds and 2 lone pairs)
Determining Molecular Shape
Examples: NO 3 -
- Nitrogen will be central. Assign it a number of 5
- Ignore other atoms because they are oxygen
- A -1 ion (add 1 electron)
- Total = 6
- Divide 6 by 2 = 3 (so 3 bonds and no lone pairs – note that one of the bonds is a double bond)
Determining Molecular Shape
A trick to determine the number of bonds + lone pairs in a molecule (from Sherry Berman Robinson)
- Determine the central atom. Assign it a number the same as its Group number (number of valence electrons). (ex. Oxygen would be 6, phosphorus would be 5, etc.)
- If the other atoms are Sulfur or Oxygen, ignore them. If not, then assign them a number equal to the absolute value of |charge x quantity|
- If a negative ion, add number of electrons equal to charge
- If a positive ion, subtract number of electrons equal to charge
- Add up numbers from steps 1-3. This is equal to the total number of electrons available for bonding
- Divide the total by 2. This is your number of bonds + lone pairs.
Formal Charge References
Examples: see the following websites:
http://en.wikipedia.org/wiki/Formal_charge#Form al_Charge
http://wine1.sb.fsu.edu/chm1045/notes/Bonding/ Drawing/Bond06.htm
http://www.chem.ucalgary.ca/courses/351/Carey 5th/Ch01/ch1-3depth.html
http://www.science.uwaterloo.ca/~cchieh/cact/c 20/dotstruc.html References: Chemistry 7th ed., Chang; http://en.wikipedia.org/wiki/Formal_charge#Formal_Charge
