iluvcats
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The iron(II) ion forms complexes with monodentate ethanoate ions and bidentate ethanedioate ions. The complexes with ethanedioate ions are more stable. What is the best explanation for this?

A. Ethanedioate ions form stronger bonds than ethanoate ions with iron(II) ions
B. Ethanedioic acid is a stronger acid than ethnoic acid
C. The formation of the ethanedivoate complex produces more particles in solution
D. Ethanedioic acid forms stronger hydrogen bonds than ethnic acid


The question is from the unit 5 paper of Jan 2015. The correct answer is C but I don't get why it is.
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Wolfram Alpha
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Because consider this:
[Fe(H2O)6]2+ + EDTA4- ---> [Fe(EDTA)]2- + 6H20
There is an increase in entropy as the number of moles has increased from 2 on LHS to 7 on RHS. This means Gibbs Free energy is more negative and thus the complex is more stable.

I only used aqeuous Fe complex to demonstrate a point. If you use any Iron(II) complex that's bonded to monodentate or bidentate ligands and react this with a multidentate ligand suxh as EDTA4-, there will always be an increase in number of particles as EDTA4- can form many co-ordinate bonds with the central metal ion so only 1 of it is needed...I hope that makes sense
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charco
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(Original post by iluvcats)
The iron(II) ion forms complexes with monodentate ethanoate ions and bidentate ethanedioate ions. The complexes with ethanedioate ions are more stable. What is the best explanation for this?

A. Ethanedioate ions form stronger bonds than ethanoate ions with iron(II) ions
B. Ethanedioic acid is a stronger acid than ethnic acid
C. The formation of the ethanedivoate complex produces more particles in solution
D. Ethanedioic acid forms stronger hydrogen bonds than ethnic acid


The question is from the unit 5 paper of Jan 2015. The correct answer is C but I don't get why it is.
ethnic acid?
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iluvcats
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why is the complex more stable when Gibbs Free energy is more negative? I get that the entropy increases and systems tend towards chaos, and that this reaction is spontaneous but how does that relate to stability?

(Original post by Wolfram Alpha)
Because consider this:
[Fe(H2O)6]2+ + EDTA4- ---> [Fe(EDTA)]2- + 6H20
There is an increase in entropy as the number of moles has increased from 2 on LHS to 7 on RHS. This means Gibbs Free energy is more negative and thus the complex is more stable.

I only used aqeuous Fe complex to demonstrate a point. If you use any Iron(II) complex that's bonded to monodentate or bidentate ligands and react this with a multidentate ligand suxh as EDTA4-, there will always be an increase in number of particles as EDTA4- can form many co-ordinate bonds with the central metal ion so only 1 of it is needed...I hope that makes sense
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charco
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(Original post by iluvcats)
why is the complex more stable when Gibbs Free energy is more negative? I get that the entropy increases and systems tend towards chaos, and that this reaction is spontaneous but how does that relate to stability?
Gibbs free energy change tells us that the complex is more likely to form.

Ligands in complexes have a tendency to break off and reattach. This is called lability and is measured by the stability constant. When the stability constant is high the complex is stable.

However, once a complex is formed a polydentate ligand is attached to the transition metal ion by more than one bond. This means that whenever a bond breaks it is still held in place by the other bond and is more likely to reattach back to the transition metal ion.
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