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Equilibrium concentration question

  1. Oct 28, 2014 #1
    It has been a bit since I took chemistry, but I have been wondering something, and cant seem to remember. Lets say that I have the following:

    Fe3+ + SCN- <-> FeSCN2+

    How can I calculate the point where Fe3+ concentrations become effectively zero? That would say, if I know that I have 1M Fe3+ in a solution, and I plan to add SCN- in large amounts to put pressure on the left side, forcing the reaction to the right, how can I calculate the concentration of SCN- needed to make the Fe3+ run out?

    Thanks in advance.
     
  2. jcsd
  3. Oct 28, 2014 #2

    Borek

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    Staff: Mentor

    Define "effectively zero".

    Fe3+/SCN- is a particularly difficult case, as there are six consecutive equilibria involved.

    Even for huge excess of SCN- you can expect concentrations of Fe3+ (actually some hydrated forms) to be present. In 0.01M FeCl3 with 1 M KSCN (so a hundredfold excess of the ligand), concentration of Fe3+ is in the 10-9 M range.
     
  4. Oct 28, 2014 #3
    Basically, the reason I am asking is due to a question that was asked of me by someone. They are trying to determine the concentration of Cl- in a solution by using Hg(SCN)2, with Fe(NO3)3 added into the mix. He told me that the professor stated that Fe was present in "large amounts."

    So I pointed out that since FeSCN will be formed regardless of Cl- being present, with a high enough Fe3+ concentration, all that SCN is going to get used up with no help of Cl-. (The Hg binding to Cl drives the reaction further to the right, but if high concentrations of Fe3+ is there, eventually the Cl- wont drive the equation anymore)

    He then asked me what sort of concentration would be needed for that. And so here I am.
     
  5. Oct 28, 2014 #4

    Borek

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    Not enough data for a reasonable answer. Without seeing what the initial concentrations are it is not clear what can happen.
     
  6. Oct 28, 2014 #5
    Ok, well if we have the two following:
    Cl-(aq) + Hg(SCN)2(aq) <-> HgCl2(aq) +2SCN-(aq)

    SCN- (aq) + Fe3+ <-> Fe(SCN)2+

    This is how it was presented to me. It caught my eye that the end reaction doesnt use Cl at all. Being that Hg(SCN)2 is listed as aq, and he confirmed that there is no solid present, this tells me that A) the top reaction shouldnt even be in an equilibrium, as all of the Hg(SCN)2 should be broken up into ions to begin with. (if memory served correctly) And B) "large amounts" of Fe3+ seems to me like it would drive the right side of the top equation to exhaustion no matter what in an equilibrium. And since they are measuring the absorbance of the Fe(SCN)2+ complex, it seemed like an odd way of testing for Cl.
     
  7. Oct 28, 2014 #6

    Borek

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    Any reason why it should?

    It definitely is.

    No, SCN- is a ligand and it complexes dissolved mercury, this happens in parallel to precipitation. Which process dominates depends on the concentrations. Not knowing them we won't get top any conclusions.

    That's not how the equilibrium works.

    Can't say I understand the idea behind the procedure you described, and not knowing expected/used concentrations of all ions involved it is impossible to say what is happening in the solution, but your analysis doesn't look correct to me.
     
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