I've been studying coordination chemistry, but I still have a fundamental question about the area: why do transition metals even form coordination compounds, while main group metals do not? Why can't they just form simple ionic compounds, treating their d electrons as valence electrons?

On that note about d-orbitals as valence electrons: when do they act as valence electrons and form "simple" ionic (such as $\ce{CrCl_3}$) or covalent ($\ce{MnO_4^-}$) bonding, and when does coordination happen instead? Lately, I've began to think that transition metals don't even form simple ionic bonds or covalent bonds. It used to seem to me that in the chromate ion, $\ce{CrO_4^{2-}}$ (by drawing the Lewis structure) the d-orbitals in chromium participate as valence electrons in covalent bonding, but then how is it colored without d–d transitions? Sulfate, with $\ce{S}$ instead of $\ce{Cr}$, is colorless $\ce{SO4^2-}$. Thus, there must be coordination chemistry involved in $\ce{CrO4^2-}$ (is $\ce{O^2-}$ even a good ligand)? And with $\ce{CrCl_3}$, is it indeeed a simple ionic bond like I thought, or is there coordination going on with that, too?

Basically, why can transition metals exhibit coordination in addition to (at least what seems like) ionic and covalent bonds?

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    $\begingroup$ I wouldn't say main group metal do not. I would actually say they almost certainly do, and you're not likely to see an unsolvated metal in solution. They do tend to be more labile bonds, however, due to the more metallic character of the alkali and alkaline earth metals. Look at the mechanism for the Weinreb Ketone Synthesis for an example of a synthetically useful of main group metal coordination. $\endgroup$ – SendersReagent Apr 20 '16 at 1:28
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    $\begingroup$ For the colour of chromate thing, please search - it's been answered before more than once. If chromate doesn't work search for permanganate. Furthermore I would strongly advise against this black-and-white approach of classifying things so strictly into one category or another. Some approaches to bonding are more instructive for certain compounds and less useful for others. For example I hardly expect CrCl3 to be completely ionic, either in solution or even in the solid state. There are numerous examples of solid state structures being driven by coordination preferences. $\endgroup$ – orthocresol Apr 20 '16 at 1:43
  • $\begingroup$ Check out this question about a main group metal complex. They certainly do exist and some of them are stable enough to be crystallised as salts with all sorts of weird anions like this. $\endgroup$ – bon Apr 20 '16 at 9:15

Your approach is fundamentally wrong. Main group elements form complexes all the time. It is just that these complexes are typically kinetically labile and quickly decompose, rearrange and so on.

Main group elements’ complexes are also typically colourless due to the lack of d-electrons or due to the fully populated d-subshell meaning that all electron transitions are in the UV range, so you cannot easily study them via colour changes.

The state of the d-subshell (full or empty, nothing in-between) also means that no additional stability of certain geometries exist; the s and p orbitals can more or less rearrange to whatever is required.

And finally note that ‘simple’ ionic compounds are, in fact, huge coordinative compounds. A single sodium ion in a $\ce{NaCl}$ crystal could be described as a $\ce{[Na(Cl)_{\frac{6}{6}}]}$ complex.


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