[habbababba]

I was surprised to read that the hydrated sodium ion is not classified as a complex ion. This poses the following question: what determines whether a metal ion can form a complex ion with ligands through dative covalent bonds whereas other metal ions interact with the negative or slightly negative chemical species through electrostatic ion-dipole attractions?

Thanks.

[magician4]

let's take a look at IUPAC , shall we?
http://goldbook.iupac.org/C01203.html

now then, is Na⁺_(aq.)  (quote)

A molecular entity formed by loose association involving two or more component molecular entities (ionic or uncharged) (...)

(unquote)
or is it not?

do we have (at least) two different sub-entities ?  yes , we do ( i.e. water, the sodium-ion )

are they at least "loosely associated" ?   yes, they are ( as we could learn from,  for example, ion-mobility experiments: the what is moving there is far bigger than the diameter usually assigned to the "naked" sodium ion), as the hull of water surrounding the ion is at least loosely bound to the ion itself, hence will have to move about with it.


  following IUPAC's definition ( and this organization is the highest authority there is in chemistry), hydrated sodium is a complex.
(besides, it makes sense . but then again,"making sense " sometimes is irrelevant when you're talking definitions. so , better let's state: all elements required to fullfill the definition of a complex are being fullfilled by [Na(H₂O)_{(approx. 6)}]⁺ , and therefore it IS a complex.


however, there might be some elder or individual definitions around still, like having to do with "empty d (or higher) metal orbitals, electron density of the ligand is donated to" and thatlike.

yeah, in this sense sodium ions, hydrated, are no complexes, simply as their empty d-orbitals are far far faaaaar away, and hence won't become occupied whilst solvation.


however , talking for me personally, I don't find those elder definitions very usefull for the problem at hand, i.e. explaining whats going on when sodium ions are being hydrated.

   insofar, I agree with IUPAC


regards

Ingo

[habbababba]

The IUPAC definition of course puts all doubts to sleep, not because the IUPAC gets the last word but because the definition seems reasonable to me as well.

The source I'm using however is a 'well-established' international high school curriculum and I'd like to think that they're using the 'elder' definition as you put it.

Now IF they are, they make a big confusion when they say the following: "Suggest how the aluminium ion is able to form dative covalent bonds in its complexes such as [AlF₆]^3-".

Now if we look carefully, Na⁺ and Al³⁺ are isoelectronic. Wouldn't this mean that the empty d orbitals of Al³⁺ are also far away as in the case with Na⁺ and hence should not become occupied during solvation? Or is there something else that makes Al³⁺ different than Na⁺ despite both being isoelectronic?

Thanks.

[magician4]

I agree: from a first glance, this seems to be inconsequential to the max.

however, taking a closer look we might wish to concede that, though being isoelectronic, Al³⁺ and Na⁺ are by no means identical in their chemical behaviour: their charge density is hugely different.

as a consequence, aluminium abhorrs the "pure" ionic state (which, as you know, is a borderline case anyway, almost never fullfilled, not even close to, not even in Al₂O₃ , if memory serves): it will prefer bonds with a high covalent characteristic, leading in turn to ( as there would be 6 electrons around the aluminium with 3 covalent bonds: not so good) at times "exotic" molecules like Al₂Cl₆ , for example

same is true for fluorine, with known "exotic" results like H⁺[FHF]^- as a result.

bring the two of them together, and there u are: unusual, covalent results, with quite strong bonds holding the structure together at that.
  the stability of the result does make the notation of "complex formation" inescapable, it seems to me, doesn't it?

... though i would love to here an explanation involving "d orbitals at aluminium" here 


"chemistry is a patchwork of explanations which at times will crash, esp. when it comes to the border between two of them"


regards

Ingo

[habbababba]

however, taking a closer look we might wish to concede that, though being isoelectronic, Al³⁺ and Na⁺ are by no means identical in their chemical behaviour: their charge density is hugely different.

So the higher charge density 'forces' the aluminium ion into forming covalent bonds. Simple and reasonable.

... though i would love to here an explanation involving "d orbitals at aluminium" here 

And this is where I'm trying to get. Often in chemistry textbooks, chemists introduce the concept of coordination complexes when they discuss the d-block elements. I am aware of coordination complexes formed by the representative metals, but chemists often do not talk about these complexes as nearly as they talk about the d-block coordination complexes, as if there is something 'special' about the electronic configuration of the d-block elements that allows them to form coordination complexes... Or isn't there?

Thanks.

[magician4]

... Or isn't there?

more often than not, they are brightly coloured, which raises curiosity?

but then again, so is Al₂Cl₆ , a very nice yellow in my opinion...
... and let's not forget about N₂O and thatlike, elementary sulfur...and those higher halogenes.. when talking 'bout maingroup elements and their ability to reach out for  the LUMO with only very little energy required.

well, yes, HOMO-LUMO distances might be the key here, which makes them d-group elements somewhat special

but asides from that...


regards

Ingo

[habbababba]

well, yes, HOMO-LUMO distances might be the key here, which makes them d-group elements somewhat special

Tell me more please.

[magician4]

as a rule of thumb, the LUMO of a given central atom shouldn't be too far away from the HOMO, else the incorporation of this orbital in the general scheme of hybridization would make it too expensive to be worthwhile:
that's why sulfur doesn't include d-orbitals when making SO₄^2-

Now, let's consider an early  maingroup element-ion: they (almost) all prefer noble gas configuration, don't they? So, the LUMO for these ions is skyhigh, and if it wasn't for charge distribution (this of course outweights the effort) , the resulting bonds should be quite weak.
... and that in fact is what we observe, esp. for group I and II elements: those bonds are weak and easily broken (though that doesn't mean that there are no such bonds!)

So yes, there are "for real bonds" even for these ions when they become, for example,  hydrated - those are not just electrostatic aggregates here - but they are weak: somewhere in the ballpark of hydrogen bridges, if memory serves.
However, though being weak, they allow for the existence of "defined" structures like [Ca(H₂O)₄(OH)₂]_(aq.) and [Ca(H₂O)₅(OH)]⁺_(aq.) ,  with their individual pKb-values of 1.37 and 2.43, respectively.

... and if something has an individual pKb-value, we'd better call it a species, right? And if those were species, then what else instead of "complexes" should we call them?

regards

Ingo

[habbababba]

as a rule of thumb, the LUMO of a given central atom shouldn't be too far away from the HOMO

Not quite sure if I'm catching what you're saying here, but are you implying that the LUMO of the d-block elements is 'closer' to the HOMO of the ligands than the LUMO of the representative metals is to the HOMO of the same ligands? (Hope I made myself clear...)

So yes, there are "for real bonds" even for these ions when they become, for example,  hydrated - those are not just electrostatic aggregates here - but they are weak

I've always learned that these interactions can co-exist simultaneously at many instances like you said.

Nevertheless, if we look at the metal and the ligand as a Lewis acid and a Lewis base, respectively, we can use the following rule of thumb: hard acid cations form complexes in which simple Coulombic interactions are dominant, and soft acid cations form complexes in which covalent bonding is important.

Now I would like to go back to the same ions, Na⁺ and Al³⁺. Taking into account the higher charge density of Al³⁺, it follows that Al³⁺ can be classified as a harder Lewis acid than Na⁺. Supposedly it is fair to make such comparison, then isn't Al³⁺ actually the species forming complexes in which Coulombic interactions are more dominant whereas Na⁺ is the species forming complexes with still dominant Coulombic interactions (because it's still a hard Lewis acid) but to a lesser extent than Al³⁺?

The reason why I'm trying to find out whether there is a contrast line between these two types of interactions is because often when it is given that ligands form dative covalent bonds with the metal ion, the color and magnetism of the resultant complex ion arise from the splitting of d orbital energies. But then the same model (The Crystal Field Model) that explains the splitting of the d orbital energies approximates that the metal-ligand bonding is entirely ionic, contradictory to the existence of dative covalent bonding in the same complex.

More than one model is on many occasions advantageous, but mind boggling in this particular case.


Once again, thanks.

[magician4]

Not quite sure if I'm catching what you're saying here, but are you implying that the LUMO of the d-block elements is 'closer' to the HOMO of the ligands than the LUMO of the representative metals is to the HOMO of the same ligands? (Hope I made myself clear...)

no, this wasn't the idea I wanted to sell you: I was talking the "naked" ion, and how it would need to hybridize to result in "good" A-orbitals for the latter MO's
(i.e.: this is a bit "the organic chemist's point of view")

I've always learned (...)
More than one model is on many occasions advantageous, but mind boggling in this particular case.

One of the things I've learned in chemistry is that there is only one theory which gives the for real answer: quantummechanics.( and there are no contradictions in Q.M., just some at times hard-to-solve equations. *⁾)
However, being humans, we need something more easy to grasp, some "rule of thumb" type foreshorthenings of QM: this gives us our patchwork of "chemical laws". And when it comes to a clash of those concepts...

"mind boggling" you said, and I'd agree.
Try to take a step sideways, and look for a fresh perspective: more often than not you'll succeed


regards

Ingo



*⁾
unfortunately , those answers more often than not remind me of Douglas Adam's Hitchhiker: "42"

================
"Maybee we couldn't solve your problem, but we hope we at least helped to loosen it up some"

[habbababba]

no, this wasn't the idea I wanted to sell you: I was talking the "naked" ion, and how it would need to hybridize to result in "good" A-orbitals for the latter MO's
(i.e.: this is a bit "the organic chemist's point of view")

But so do all the ions that don't belong to the d-block. So I'll jump to the question: how does the hybridization of the d-block metals differ from that of the other s & p-block metals?

"Maybee we couldn't solve your problem, but we hope we at least helped to loosen it up some"

You sure have and for that I thank you.

[magician4]

So I'll jump to the question: how does the hybridization of the d-block metals differ from that of the other s & p-block metals?

whenever you wish to create a hybrid orbital, you'll have to "pay" for the new orbital(s) included. Hence, your "new" total state will be energetically elevated with respect to what it would be without hybridization.
(sometimes named "hybridization energy")

the amount of this elevation depends on the difference between the former ground state and the energy of the former unoccupied orbital:*⁾  the bigger this difference, the higher the energy of the new hybrid will be
  your "new" hybridized ion will be energetically elevated

now, if you already are in the d-block occupation with your ion, even without hybridization, those empty "other" d-orbitals are nearby: HOMO-LUMO difference is small

... and if you're not, then there's hell to pay

regards

Ingo


*⁾
however, even with the "next empty" orbital not being that far away, there's a price to pay which at time might be too high if it doesn't get properly compensated for ( by better bonds, more favourable geometries and thatlike)
... as we could learn from , for example, PH₃ ( which is next to not hybridized at all) or water ( which is not, contrary to urban legend , sp³ with two equal lonepairs, but something near sp²-p instead [see for example  Levine I.N. “Quantum chemistry” (4th edn, Prentice-Hall) p. 470–2, 475) ]

[habbababba]

Your efforts for clarification were of great assistance. Thank you indeed.