2E 11 b)

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Labiba Sardar 2A
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2E 11 b)

Postby Labiba Sardar 2A » Sat Nov 16, 2019 9:48 pm

When drawing the Lewis structure for ICl3, how do you know that iodine will have 2 lone pairs around it? Why can't you add multiple bonds, making the atoms have an expanded octet?

Jessica Castellanos
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Re: 2E 11 b)

Postby Jessica Castellanos » Sat Nov 16, 2019 9:52 pm

By adding double bonds, the Cl will gain charges which is not favorable because it's not stable, so instead of adding double bonds, I should have lone pairs. Hope this helps!

Anish Patel 4B
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Re: 2E 11 b)

Postby Anish Patel 4B » Sat Nov 16, 2019 9:54 pm

When you complete the octets of the chlorine atoms, you have 24/28 electrons filled for the molecule. Adding double or triple bonds would decrease this number of overall electrons used by taking one from each atom. The only way to fill the quota of 28 electrons is to add two lone pairs, which is possible since I can form expanded octets.

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Re: 2E 11 b)

Postby KnarGeghamyan1B » Sat Nov 16, 2019 9:56 pm

Also, if we added double bonds, the Iodine atom would have a +2 charge since it would have 5 bonds around it, and we try to make the central atom have as close to a formal charge of 0 as possible.

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Re: 2E 11 b)

Postby stephaniekim2K » Sat Nov 16, 2019 10:00 pm

Chlorine cannot have more than a single bond so we know that there are 3 single bonds on iodine, meaning there will be two left over lone pairs on iodine.

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