(Polar molecules, Non-polar molecules, etc.)

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Evonne Hsu 1J
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Postby Evonne Hsu 1J » Mon Nov 16, 2020 3:24 pm

I'm sure Dr.Lavelle has gone over this in one of his lectures, but I'm still confused about when to use a single/double/triple bond and when to just leave it as a lone pair. I realize that there are thousands of compounds with a lot of exceptions but what is the general rule to look out for when drawing out compounds?

Crystal Hsueh 2L
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Re: Bonding

Postby Crystal Hsueh 2L » Mon Nov 16, 2020 3:26 pm

There are a lot of exceptions and ways to draw Lewis structures (hence resonance structures) but I think generally you should try to make bonds over leaving electrons as lone pairs. I believe this tends to make the structure more stable depending on the element and its formal charge. Most of the time you should be able to figure it out from the number of valence electrons.

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Re: Bonding

Postby Chem_Mod » Mon Nov 16, 2020 3:27 pm

A good guiding principle is to find formal charge. For example, if nitrogen is double bonded with two lone pairs (satisfying the octet rule), then the formal charge is -1. Whereas, if nitrogen is triple bonded with one lone pair, the formal charge would be zero.

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Re: Bonding

Postby OwenSumter_2F » Wed Nov 18, 2020 9:28 am

Also, as far as molecular shape goes, bonds don't change the shape whether they are single or double or triple, yet lone pairs can.

Shrey Pawar 2A
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Re: Bonding

Postby Shrey Pawar 2A » Wed Nov 18, 2020 9:34 am

The first thing I would look at when deciding is the formal charge of the individual elements. If the charge is closer to neutral with more bonds that would be the solution as opposed to using lone pair electrons.

Madisen Brown -1C
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Re: Bonding

Postby Madisen Brown -1C » Wed Nov 18, 2020 6:57 pm

I usually take formal charge into consideration when determining whether a single, double, or triple bond would yield the most stable structure.

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