Prediction of Complex Ion Dissociation Question

[betterfuture]

So I was wondering how it's possible to predict the species that form after a complex ion reacts with, say, a strong acid or something. Is there a way to do so? I know that transition metals can take on many oxidation states so how do we figure out the how to write out the molecular formula of a complex ion forming as well as how to write out the reaction taking place?

Somehow, this question required just that and I am not sure how to figure it out, quite honestly.

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[aldol16]

I don't think you have to predict the complex ion that forms upon addition of a strong electrolyte. I think what's important for this question that the answer does not emphasize is that there is enough information to arrive at the correct answer without predicting the reaction itself. That is, you know that copper usually only takes a few oxidation states - +1 is the most common because that allows Cu to achieve a d10 configuration when it's in complex (it uses the s-orbitals to bond since these have the greatest penetration to the nucleus and thus "promotes" its 2 s electrons to the d subshell). That's a little complicated, so you might also know that copper does not like high valent states and prefers to be in the +1 oxidation state. Either way, CuCl6 2- implies Cu(IV), which you should never have seen before and essentially does not exist. So that rules out B and D immediately. You can then rule out C because this is likely an equilibrium and not a quantitative reaction because all the chloride ions are doing is competing for ligation sites with aqua.

 

[betterfuture]

I'm going to have to read that a couple of times! So basically your saying transition metals have preferred oxidation states?

 

[aldol16]

I'm going to have to read that a couple of times! So basically your saying transition metals have preferred oxidation states?

Some do! For instance, Zn(II), Cu(I), Ni(0), Pd(0), etc. Just look up "noble metals" and a lot of the inert ones should be there. The preferred oxidation states have to do with the unusual energetic stabilization these metals get from having d10 or d5 shells.

 

[betterfuture]

So if the orbitals are half full or have full complete d orbital they are most stable? As in, it's easier that the lower energy 4s lose its 1 electron and have a full d orbital because it's most stable that way.(copper)

Like for iron, it's common oxidation might be 3+, since it can lose its 2 electrons in 4s orbital and 1 from the d orbital to attain a half d orbital making it stable?

 

[aldol16]

Yeah, half-filled and fully filled d-shells make metals more stable. That's why there's the so-called "18 electron rule" for metals where you can't do much chemistry with metals that have 18 around it - for tetrakis palladium, for instance, you have to blow off a phosphine before you can do any chemistry in most cases.

+3 is indeed one of the most common oxidation states of iron - if you remember from general chemistry, you usually see ferrous and ferric.

 

[betterfuture]

I think you meant iron

 

[aldol16]

Thanks! I corrected it now.