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# Formal charge and dot structures

Definition of formal charge, and how minimization of formal charge can help choose the more stable dot structure.  Created by Jay.

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• It's ok for Sulfur to have 6 bonds although it isn't in the 4th orbital or higher?
• The octet rule can be broken by elements starting in the 3rd period and below. It is attributed to the not so large energy difference between 3p and 3d orbitals, which allow for additional shared pairs beyond the octet. Examples of this include SF6, PCl5, SO4(2-), etc.
• Can't we use the formula of Formal Charge as :

FC = ""Total Number of Valence Electrons in free atom - Total Number of electrons in Lone Pairs - 1/2 Total Number of Bonding Electron""
• Yes this formula works out to give you the correct answer.
Total number of electrons in lone pair + 1/2 Total number of bonding electrons = Number of valence electrons in bonded atom
• 1) at , how come the formal charge be +? does that mean the molecule is cation?
(#V.E. in free atom - # V.E in bonded atom) does that mean the molecule has one more electron that is not bonded? then doesn't it make the NH4+ anion? as it has one more electron?

2) at , is there any reason why S is attached to 4 Os first and Hs are attached at each end? why can't Hs be attached to S directly? Thank you
• NH3 is a neutral atom, N has 3- charge and H +1, in NH4 the N forms a dative covalent bond with the H (since H will probs lose it´s electron, +1), nay, the overall charge will be+1
(1 vote)
• I dont get how to draw a dot structure when calculating a formal charge....As it is a bit different with respect to the usual dot structure! Please help me out!
• We draw the dot structure in the exact same manner, and then calculate the formal charges for the atoms in the molecule.

Remember that formal charge is calculated by taking the # of valence electrons, minus the lone electrons and the bonds, and we show that charge next to the molecule.

Take ::O=C=O:: for example. Each O's formal charge would be calculated by: 6 (valence) - 4 (lone electrons) - 2 (bonds) = 0. C's formal charge is 4 (valence) - 4 (bonds) = 0.

Hopefully this helps you.
• Why is the Nitrogen in the center if you mention that the least electronegative atom in the center (wouldn't it be Hydrogen in this case)?
• H can't go in the centre because it can bond to only one other atom. H must always be a terminal atom.
• Why is a lower formal charge preferred?
• It takes energy to remove electrons and create a positive charge, so a lower formal charge usually indicates a lower energy level.
• At
Q: Is there a different method to calculate the formal charges?
I think my teacher has a different method.
( V electrons) - (the bonds)- (the number of free electrons).
Is it correct or did I make a mistake ?
• Both methods are correct.
The method your teacher used is best when you know the electron-dot structure of the molecule ( e.g, H-O-SO₂-OH).
The methods used in the video are best for when you don't know the electron dot structure (e.g. H₂SO₄). They are just summaries of the results obtained from your teacher's formula.
• What is the difference between formal charge and partial charge?
• The formal charge is the charge that an atom appears to have when we count the electrons according to certain arbitrary rules.
The partial charge is the charge that the atom really has.