Organic Chemistry 351

11. FORMAL CHARGES   4  Ch 1.6

 In some covalent bonds, one atom "provides" both electrons. This is referred to as  a "coordinate polar covalent" bond and can be found in both charged organic "species" and neutral organic compounds . Formally, in such bonds the donor atom will bear a  = charge and the acceptor a  Bcharge).  These charges are not really full-blown charges but partial charges i.e. d+ & d-- since the electrons in such bonds are somewhat delocalised. Some familiar and not (yet!) so familiar examples of this type of bonds are found in:

In organic compounds,  formal charges are a method of keeping track of electrons and they may or may not correspond to real charges to ensure the dreaded "octet". In most cases, if the Lewis structure shows that an atom has a formal charge, it actually bears at least part of that charge.
 The concept of formal charges demonstrates which atoms bear most of the charge in a charged species and it illustrates the charged atoms in an overall neutral molecule.  Obviously, they point out nucleophilic and electrophilic sites in these species.
 In nitromethane, CH3NO2, one O in the N=O in neutral whereas the other one has a "formal" negative charge (balanced in this overall neutral cpd by the=N).
The equation for calculating formal charges (see also Bp14) is [e= electrons]
Formal Charge = #val.electrons - ( Lonepair electron  + bonding electrons/2 )
or
Formal Charge = Periodic Table Group# - # of dots - #lines

 where v= # of valence e (also Per. Table #), Lp= total # of lone pair e and b= total # of bonding e
Try using this equation with the Ns and Os in CH3NO2, hydronium ion, hydroxide ion, water, ammonia, ammonium ion, etc.  The general rule is ;
A four-bonded N is +, a three-bonded O is + a single-bonded O is -ve, a two-bonded N are -ve and, a triply bonded C can be either +  or - dependent whether it has 6 or 8 electrons (see below). [ C: means 8 electrons and thus a carbanionC- ]
 

Element
+ (cation) neutral - (anion)
C
three bonds only
CH3+ (and CH2=CH+)
four bonds
CH4 ( & CH2=CH2 & HCtripleCH)
three bond WITH  1 lone pair
CH3:-
N
four bonds NH4
(& >CH=NH2+,
three bonds + 1 lone pair NH3
(& >CH=NH)
two bond + 2 lone pairs NH2-
O
three bonds H3O+ [one lone pair] (& >C=OH+) two bonds H2O+ 2 lone pairs one bond + 3 lone pairs HO- 
(and CH3CO2- :i.e. acetate)

VERY REACTIVE CARBON i.e.REACTION INTERMEDIATES

 CARBOCATIONS, CARBON RADICALS & CARBANIONS are known to exist for a short time (called "intermediates") during the course of many OC rxns. They are usually too unstable to be isolated [though some especially stable ones have been.

 Carbocations are 6 electron species ("electrophiles") hybridised sp2 and are thus trigonal planar in shape with an empty 2p orbital; they react very rapidly with electron-rich ("nucleophilic") sites in other molecules in order to achieve a filled shell (or "octet"), i.e. CH3+  (Note: in some older organic chemistry books (Sykes!) , these species were called "carbonium" ions.)

 Carbon radicals are 7 electron species hybridised sp2 (also "electrophiles" cf. carbocations), 1 electron in the 2p orbital: they react with electron rich species or other radicals --> 8 shared electrons, i.e. CH3.

 Carbanions are 8 electron species ("nucleophiles") again usually hybridised sp2  and are trigonal planar with the 2p orbital containing a lone pair of electrons; these anions react very rapidly with electron-poor  ("electrophilic") species i.e. CH3-.
N.B. Some textbooks may claim that the methyl carbanion is sp3 hybridised. For the purposes of this course, it is best to consider this unlikely and that ALL CARBANIONS SHOULD BE CONSIDERED hybridised sp2.