Bonding between main group elements is obviously almost entirely covalent and it is often easy to see these covalent bonds as 2c-2e bonds localised between two atoms. The bonding in borane, boron clusters and hypervalent compounds like SF6 introduce a more delocalised view of covalency as the bonding can be rationalised in an MO diagram. Nevertheless, covalent bonding can still be thought of as electrons occupying a molecular orbital which forms when orbitals from two or more atoms overlap (this orbital mixing requires good energy, size and symmetry match).
However, I find it hard to rationalise covalency and trends in it when considering other compounds and I was wondering if there is a good, simple(ish) way of rationalising (even if not predicting) the presence of covalency in other compounds.
People say if you have a polarising cation and a polarisable anion like in AgCl for example, you get covalency. I suppose this also makes sense because that's essentially saying that the atomic orbitals need to be of a similar energy to mix. However I am going to run through three examples that I struggle to rationalise or have questions regarding it.
MeLi Li is clearly quite polarising (high charge density) and must therefore polarise C and there is some overlap of orbitals = COVALENCY. However, it is tetrameric in solution. Now compare it to MeMgBr; this is monomeric in solution (and in organic chemistry is softer) indicating more covalency - makes sense because of the increased polarisability of Mg relative to Li. However, is the tetramer vs monomer as a result of the difference in covalency, if so why? Is it fair to say that moving towards a more monomeric molecular structure is indicative of covalency (eg BeCl2 - chain vs MgCl2 - layered lattice vs CaCl2 - rutile)?
Suboxide Group 1 suboxides are possible (eg Rb2O9) when Rb is exposed to limited oxygen. This forms a cluster with delocalised electrons. It seems bizarre to me because this is one of the least polarising cations going - there are no suboxides higher up the group despite it being easier to oxidise these metals lower down the group. The only vaguely analogous example is the FeS clusters seen in biology with highly delocalised electrons but the electronegativities make more sense here.
Zn Zn is a soft Lewis acid. Why? It easily loses two electrons and then is left with a filled 3d shell. Surely that would mean it is disinclined to covalently bond and should be relatively similar to Ca?