Why Do Some Elements Have Multiple Oxyanions While Others Do Not?

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    Electronegativity
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Discussion Overview

The discussion revolves around the reasons some elements exhibit multiple oxyanions while others do not, exploring the underlying theories related to electronegativity, oxidation states, and electron configurations. The scope includes theoretical considerations and chemical behavior of various elements and their oxyanions.

Discussion Character

  • Exploratory
  • Technical explanation
  • Debate/contested

Main Points Raised

  • One participant questions the existence of multiple oxyanions for elements like chlorine (ClO-, ClO2-, ClO3-, ClO4-) compared to iodine (IO3-) and bromine (BrO3-), suggesting a possible link to electronegativity.
  • Another participant notes that periodate (IO4-) and perbromate (BrO4-) do exist, and mentions additional compounds like H5IO6, HBrO, and IO-, indicating a broader context for iodine and bromine.
  • A claim is made that fluorine cannot form an oxyanion like FO3- due to its high electronegativity and small size, making it resistant to oxidation by oxygen.
  • Discussion includes the different behaviors of pnictides (nitrogen group) regarding oxidation states, with phosphorus being able to form H3PO4 in the +5 oxidation state due to having more electrons available compared to nitrogen.
  • Another participant suggests that the ability to form compounds is not solely determined by the number of electrons in the outer shell, referencing noble gas compounds like XeF6 as an example.

Areas of Agreement / Disagreement

Participants express differing views on the factors influencing the formation of multiple oxyanions, with no consensus reached on the primary reasons behind these chemical behaviors.

Contextual Notes

Limitations include the potential influence of electronegativity, oxidation states, and electron configurations, which are not fully resolved in the discussion.

Who May Find This Useful

Chemistry students, educators, and enthusiasts interested in the behavior of elements and their oxyanions, as well as those exploring the implications of electronegativity and oxidation states in chemical compounds.

so-crates
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Is there any theory as to why you have ClO-, ClO2-, ClO3-, and ClO4-, but only IO3- and BrO3- ? And why is there no FO3- ? I suspect this has something to do with electronegativety but I'm uncertain. And how come you have NO3-, but PO4 3-?
 
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Periodate and perbromate exist (IO4- and BrO4-). There is also H5IO6, HBrO, and IO-.

You are correct. Fluorine is too small and too electronegative to be oxidized by oxygen. Fluorine is the only element more electronegative than oxygen.

All the pnictides (nitrogen group) have different behaviors when it comes to oxidation states. But basically, the reason you have H3PO4 for P in the +5 O.S. and not HPO3 is because phosphorous has more electrons to take away than nitrogen. Nitrogen only has 7 electrons total, and only 5 valence electrons. Once there are three oxygens taking electrons away from nitrogen it's over, no more oxygens can fit. Phosphorous has 15 electrons total and there are more than the 5 considered to be valence electrons that are available to be taken by oxygen. Another thing that contributes to this electron grabbing is the greater difference in electronegativity between phosphorous and oxygen.
 
heh

so-crates

Bill and Ted
 
Well, it's not completely determined by number of electrons in the outer shell, because you can relatively easily make noble gas compounds like XeF6, but for all practicle purposes its just looking at what can make a full shell, or half full shell.
 

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