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Postby Guzman_1J » Sun Dec 01, 2019 8:49 pm

I was going through my notes and there's something I didn't fully understand: "oxoacids more readily lose H+ if resulting anion is stabilized by electron withdrawing atoms which delocalize and stabilized the negative charge". Could someone explain how the electron withdrawing atoms influence the negative charge or exactly what this is trying to say?

Sydney Pell 2E
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Re: notes

Postby Sydney Pell 2E » Sun Dec 01, 2019 10:03 pm

This pretty much means that oxoacids can dissociate better and thus are stronger acids when, after giving off a proton, they are able to delocalize the negative charge. This means that oxoacids composed of atoms with similar electronegativities (and therefore would delocalize the charge) will be more likely to give away their proton.

For example, Fe(III)-O- would be a stronger oxoacid than Fe(II)-O- since the +3 charge in the first one would pull on the negative charge of the Oxygen more than the +2 charge of the second one would. So, the first oxoacid is stronger since the negative charge is not localized just on the oxygen, but is more spread out and stable.

Sarah Nichols 4C
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Joined: Sat Aug 24, 2019 12:18 am

Re: notes

Postby Sarah Nichols 4C » Sun Dec 01, 2019 10:14 pm

I think this is saying that oxoacids with other electronegative atoms are stronger because these atoms delocalize the negative charge and make it less likely for a proton to bind/stay bound. For example, in OCl- (hypochlorite ion), the negative charge is largely localized on the single oxygen atom; as a result, hypochlorous acid is weak because the bond between the oxygen and hydrogen is strong and hard to break. However, in ClO4- (perchlorate ion), the negative charge is delocalized over all four oxygen atoms, weakening the hydrogen bond. Therefore, perchloric acid is very strong because it will readily dissociate (lose H+)

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