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How are perchlorate or chlorate or chlorite ions and their respective acids or compounds formed. $\ce{Cl}$ can't form more than one bond but still...

Perchlorate ion from Wikipedia $\rightarrow$'Perchlorate ion'

Chlorous acid containing Chlorite ion from wikipedia $\rightarrow$'Chlorous acid containing Chlorite ion'

Chlorate ion from Wikipedia $\rightarrow$'Chlorate ion'

Chlorine is making more than 1 bond...How is it possible???

Am I having some confusion or misunderstanding? If so, please help to clarify

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In terms of valence bond theory Chlorine can have an expanded valence because it has an empty 3d subshell. –  1110101001 Aug 12 at 19:04

3 Answers 3

up vote 6 down vote accepted

The order of bonding, and so the valence state of Cl in $\ce{ClO_{x}-}, x>1$ compounds is very debatable. Generally, two models exist.

  • $\ce{Cl}$ atom, just like $\ce{S, P}$ and some others has unoccupied $d$-orbitals in the valence shell. It is possible to move some electrons from $p$-orbitals to $d$-orbitals, producing half-occupied orbitals that can participate in covalent bond formation

  • $\ce{O}$ atom has 6 electrons at the valence shell with four orbitals. It is possible to move one electron from chlorine atom to oxygen atom to form $\ce{O^{\dot{-} }}$ ion, that can than form a bond with newly formed half-filled orbital of chlorine. However, being electronegative and having high positive charge, chlorine partially pulls back the electron pair of newly formed bond.

Currently the second option is in favor of theoreticians, both by electron population analysis and relatively high charges in the oxochloric compounds, however, in school-level chemistry people traditionally resort to the first option. On the other hand, the second approach is the only option to adequately describe nitric acid and nitrates as nitrogen does not have $d$-orbitals on the valence shell.

Now, the stability of oxochloric compounds is a completely different matter. All of them are not very stable compounds, easily freeing their oxygen. Still, the process has significant activation energy, the state in between is even less favorable then oxochloric molecule. both oxygen and chlorine have electron-rich shells that repel potential reactants. Perchlorate ion is especially stable, at its surface is completely covered by outer negative shells of oxygen, so electrons cannot reach positively charged chlorine.

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I like this answer. –  Dissenter Aug 12 at 20:14

The Cl-O bond has a dissociation enthalpy of 269.1 kJ/mole in the gas phase. That's not very low. This suggests that yes, the Cl-O bond is stable. Similarly, the Cl-Cl bond has a dissociation enthalpy of 242.580 kJ/mole. Again, not very low, but not very high. For comparison, consider triply bound diatomic nitrogen - the third strongest covalent bond known. This species has a dissociation enthalpy of 945 kJ/mole - about three times higher than the aforementioned energies.


Chlorine can have multiple bonds, as you've shown in your Lewis structures. Try looking into molecular orbital theory.

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I don't understand what is the big deal with bonds in acids or esters in highier oxidation states - I always thought of them as of single dipolar bonds. They shouldn't be written as double bonds - it an archaism for me - now chlorine oxoanions aren't considered hypervalent. Even hypervalency of molecules like iodine heptafluoride has been criticised.

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