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I'm still pretty confused on when to do double and triple bonds on a Lewis structure. For example in the compound AsF3, I'm not sure if it's supposed to have a double bond on two of the fluorines. I calculated formal charge and with the double bonds on two fluorines, I get a formal charge of 0 on arsenic and all of the fluorines. If I do single bonds I get a formal charge of 2+ on arsenic and 0 on fluorine. Also, the Lewis is supposed to have a total of 26 electrons and I only got 26 when I did two double bonds. So, doesn't it make more sense to have two double bonds on two fluorines?
AsF3 would not have any double bonds. Instead, there should be a single bond from As to F, with each F having three lone pairs and As having one lone pair. This adds up to 26 electrons and they all have a formal charge of 0, as As has 5 valence electrons, which the 3 bonds and 2 lone electrons satisfy.
Savannah Mance 3B wrote:Why does arsenic have lone pairs? And how would you know it would have lone pairs being the central atom?
From practice what helps me is abiding b y the rules of electronegativity. Though if it helps play with the problem by filling up all bound atoms valence shells and if electrons are left over it is an indicator that there may be LP on central atom (these electrons have to go somewhere).
Make sure that your formal charges equal the charge of the molecule, however they should still be minimized. Electrons should be counted before drawing the molecule to help determine how many bonds/lone pairs you will have
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