Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: jcais on July 10, 2006, 04:59:53 PM
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Hello,
We performed the following: Fe3+ + SCN- <> FeNCS2+
Fe3+ and SCN- are lightly colored. The product is dark.
The <> means equilibrium.
We added a couple drops of HCl and the solution became lighter. So, I am thinking that the equilibrium shifted left.
Is this because the HCl enhanced the products side and so to lessen the products, the equil. shifted to the left to reduce the products?
This equation was given to us in addition:
Fe3+ + 6Cl- <> FeCl(6)3-
Do you think the HCl created more products upon addition, so it had to shift left?
Thank you for your time and help. :)
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I'd argue that Chloride ions added to your equilibrated system react with Iron(III) aquo ions, thus reducing the concentration of Iron(III) aquo ions, pushing the equilibrium towards the left.
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We performed the following: Fe3+ + SCN- <> FeNCS2+
Is this because the HCl enhanced the products side and so to lessen the products, the equil. shifted to the left to reduce the products?
This equation was given to us in addition:
Fe3+ + 6Cl- <> FeCl(6)3-
Do you think the HCl created more products upon addition, so it had to shift left?
The simplified equations written correctly are:
[Fe(H2O)6]3+(aq) + SCN-(aq) <-> [Fe(H2O)5(SCN)]2+(aq)
[Fe(H2O)6]3+(aq) + 4Cl-(aq) <-> [FeCl4]-(aq) + 6H2O(l) (doesn't matter if 3, 4 or 6(!) chlorides actually take part in the reaction, it's the concept that matters)
Therefore when you add the chloride ions the second equilibrium will shift to the right and remove aqueous Fe3+ ions from the equilibriums, forcing the first equilibrium to shift to the left as you and Dan correctly said. This accounts for the lighter colour.
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Hint: HSCN is not a strong acid. Although its pKa is sometimes given as 4.0 and sometimes as 0.9, no idea which one is correct - if 4.0 effect will be much more visible. Iron SCN- complexes have substantially higher stability (creation) constants - but you better check it, as that's what I remember, no time to dig in my books now :)
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Does the addition of HCl remove more of the reactant Fe ion or the product Fe ion? I am thinking it has more of an effect on the reactant Fe ion to cause a shift left, right? Thank you :-)
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Hint: HSCN is not a strong acid. Although its pKa is sometimes given as 4.0 and sometimes as 0.9, no idea which one is correct - if 4.0 effect will be much more visible.
A thiocyanate salt like KSCN would be a better source for thiocyanate in this reaction.
Does the addition of HCl remove more of the reactant Fe ion or the product Fe ion? I am thinking it has more of an effect on the reactant Fe ion to cause a shift left, right? Thank you :-)
The HCl removes the reactant [Fe(H2O)6]3+(aq) from the equilibrium, so the equilibrium with thiocyanate shifts to the left to replace the lost [Fe(H2O)6]3+(aq).
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HCl presence both protonates SCN- (removing it from the reaction) and complexes Fe3+ (removing it...). IMHO first process is much more important, but I don't have time to check details right now.
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Thank you all! :D
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HCl presence both protonates SCN- (removing it from the reaction) and complexes Fe3+ (removing it...). IMHO first process is much more important, but I don't have time to check details right now.
I just realised that after I posted that :-[
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A thiocyanate salt like KSCN would be a better source for thiocyanate in this reaction.
It doesn't matter in this case - if you start with KSCN but you ad strong acid to teh solution, you will protonate SCN- ions. SCN-/HSCN ratio will be given by the Henderson-Hasselbalch equation - assuming Ka is 4.0 at pH=2 (0.01M HCl - not that concentrated) only about 1% of thiocyanate is available as SCN-. Everything else is HSCN.
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I just realised that after I posted that :-[
Don't worry, you are organic type, ionic equilibria is my domain :)