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A more stable Lewis structure will have a formal charge of 0. That is why we try to draw the Lewis Structure in different ways to see if we can get the formal charge to 0. In the example provided of (S04^2-), the first Lewis Structure we drew had a final formal charge of (-1); whereas, the second one had a final formal charge of 0, making it more stable.
The arrangement that allows for the lowest formal charge is the one that matches closest to the true structure. Also, it is important to remember that the atom with the lower electron affinity should not get the most negative formal charge.
The goal of manipulating Lewis Structures is to minimize the formal charge of the overall molecule and thus find the most stable form of the molecule. During the lecture example with SO4^2-, we calculated each individual atom's formal charge and then added them all up to see the overall charge, which we used to compare to the overall formal charges of other forms of the same molecule. The more stable form of SO4^2- had two oxygen atoms (had double bonds with S) with formal charges of O, two oxygen atoms (had single bonds with S) with formal charges of -1, and a sulfur atom with a formal charge of +2. The overall formal charge of the molecule was thus 0 because (+2) + (-1) + (-1) = 0, and this form of the molecule is clearly more stable than the other form, which had all oxygens bound to sulfur by single bonds and an overall formal charge of -2.
The goal is to minimize formal charge overall but if it comes down to minimizing the formal charge on one of the outer atoms at the expense of the center atom's formal charge increasing than it's better to have the center atom with the lowest possible formal charge.
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