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The most stable compounds typically have an overall total charge of zero (the exception would be if the compound or element is an anion or cation, that is, if they have an overall charge other than zero). And just a quick note, formal charge = # of valence electrons an atom has - (# bonds + lone pair electrons), or # of valence electrons an atom has - (# of lone pair electrons + # of shared electrons/2)
You know you are dealing with the most stable version when the Formal Charges are closest to zero. Try to find the different structure for the naturally occurring molecule, and whichever is closest to zero is the most stable.
In cases where it is impossible to entirely eliminate the formal charge on each atom, the most stable structure will have the negative formal charges on the more electronegative atoms and the positive formal charges on the less electronegative atoms.
Using the equation given in class, you can find the formal charge of a molecule. When the FC equation equals zero, you have reached the most stable form of the compound. If it does not equal zero, you know some modifications can be made to the electrons or bonds in order to reach a FC of zero and have a more stable compound.
Clara Cho wrote:How do you know which atom can hold more of the negative charge?
To maximize the stability of your structure, any remaining negative formal charge should be preferably assigned to the more electronegative elements first. I'm sure Lavelle will not expect us to memorize the electronegativities for all the various elements, so while there are several exceptions, just remember that electronegativity generally increases going left to right across a period and decreases going down a group. As an example, in the sulfate ion (SO4 2-), one would ideally place the negative formal charges on the oxygen atoms before the sulfur atom, as oxygen has a higher electronegativity.
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