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It is written here that chelation increases overall entropy?

I thought it should decrease entropy because the ligands are now bonded to the central metal atom forming a ring like structure. According to me entropy must decrease because the ligands can no longer move around so there is a decrease in randomness. Why this logic is not correct?

  • 3
    $\begingroup$ Because you don't chelate bare cation. It's actually a substitution. $\endgroup$
    – Mithoron
    Commented Dec 16, 2016 at 14:24
  • $\begingroup$ Can you please elaborate?? $\endgroup$
    – Arishta
    Commented Dec 16, 2016 at 14:25
  • 2
    $\begingroup$ Read up on Martin's answer here (chemistry.stackexchange.com/questions/31005/…) $\endgroup$ Commented Dec 16, 2016 at 14:27

2 Answers 2


The chelate effect is very simple. Consider the addition of $\ce{Na2H2edta}$ to a random transition metal (e.g. manganese(II)). Your starting complex is a hydrate complex; we shall assume it is $\ce{[Mn(H2O)6]^2+}$. The reaction equation is given below:

$$\ce{[Mn(H2O)6]^2+ + edta^4- <=>> [Mn(edta)]^2- + 6 H2O}\tag{1}$$

You started with two molecular entities, you end up with seven. Since $2<7$, the product side has a higher entropy (more independent particles). The same is true for any chelating ligand that displaces more than one monodentate ligand.

Yes, upon chelation the chelating ligand does lose degrees of freedom concerning bond rotation. However, the number of particles greatly outweighs that.


I think Mithoron has the right idea. Say before chelation your cation is in water, surrounded by its typical hydration sphere of solvent molecules coordinating to it. Then the chelating agent comes along, displacer the water. I'm guessing the driving force for the substitution is enthalpic in origin, the cation-chelator coordinate bonds are more energetically stable/preferable.

Intuitively, which system can be arranged more ways: 1) Bulky chelator molecules strewn about a sea of solvent molecules, and metal cations coordinated by some number of solvent molecules, or 2) that number of solvent molecules freely dispersed in the solution, and the chelators coordinating the metals...?

Since the free chelators themselves require an ordered solvent structure around them, my intuition suggests that performing the chelation and in the bulk solution swapping chelators for solvent molecules would increase entropy, because these newly freed solvent molecules can occupy nearly any position and orientation (whereas previously the free chelators required solvation spheres).

tl;dr Perhaps the entropy increase of dissembling the ordered solvation spheres around the free chelator molecules compensates for your entropy decrease of forming the ordered chelator-cation complex, leading to a net increase in entropy.

I think your situation may be interesting to compare to the rationalization of the hydrophobic effect.


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