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CaptainOfSmug
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[*] posted on 7-3-2013 at 10:49
Help with an equilibrium reaction


Hi all! I recently have been trying to learn and master equilibrium reactions. However, I recently came across a problem that has seemingly made my brain hurt. Essentially the equation is (net ionic) Fe+3 + SCN- -->Fe(SCN)+2. I intially increased the concentration by adding more Fe(NO3)3 to the solution driving the reaction to the product because the dark red color became more dark red. I then added KSCN which made the solution even darker red. Here's where my question comes into play. I then added AgNO3 (in excess) to my stock solution of Fe(NO3)3 which formed a cloudy white solution. Is this considered a precipitate which is mainly AgSCN? I then added KSCN into the tube and I could tell immediately the silver precipitate floated to the bottom but an orange supernatant liquid was also present. I can't tell which way the reaction went? It seems to be like it went both ways? Any help here would be much appreciated!
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[*] posted on 7-3-2013 at 11:44


Quote: Originally posted by CaptainOfSmug  
Hi all! I recently have been trying to learn and master equilibrium reactions. However, I recently came across a problem that has seemingly made my brain hurt. Essentially the equation is (net ionic) Fe+3 + SCN- -->Fe(SCN)+2. I intially increased the concentration by adding more Fe(NO3)3 to the solution driving the reaction to the product because the dark red color became more dark red. I then added KSCN which made the solution even darker red. Here's where my question comes into play. I then added AgNO3 (in excess) to my stock solution of Fe(NO3)3 which formed a cloudy white solution. Is this considered a precipitate which is mainly AgSCN? I then added KSCN into the tube and I could tell immediately the silver precipitate floated to the bottom but an orange supernatant liquid was also present. I can't tell which way the reaction went? It seems to be like it went both ways? Any help here would be much appreciated!


Adding silver will ppt AgSCN, removing thiocyanate from the solution, causing your complex to break up (shift to reactants, loss of red colour). Adding more thiocyanate will cause your reaction to shift to products (orange colour forms back), but also redissolves some of the precipitate as [Ag(SCN)4]3- (Kf = 5e+9 according to my old Porile "Modern University Chemistry" textbook).




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CaptainOfSmug
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[*] posted on 7-3-2013 at 12:30


Okay thanks! So when AgNO3 was added the equilibrium shifted to the reactants Fe3+ and SCN-? That being said the concentration of Fe(SCN)2+ decreased while the Fe3+ concentration increased? I guess I'm still a little confused why the addition of AgNO3 which was only involved in the last reaction affected the equilibrium of Fe3+ + SCN- -><- Fe(SCN)2+ even though it wasn't involved in that equation. Is it because the the thicyanate is removied which increases the concentration of Fe+3?
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[*] posted on 7-3-2013 at 12:55


Quote: Originally posted by CaptainOfSmug  
Okay thanks! So when AgNO3 was added the equilibrium shifted to the reactants Fe3+ and SCN-? That being said the concentration of Fe(SCN)2+ decreased while the Fe3+ concentration increased? I guess I'm still a little confused why the addition of AgNO3 which was only involved in the last reaction affected the equilibrium of Fe3+ + SCN- -><- Fe(SCN)2+ even though it wasn't involved in that equation. Is it because the the thiocyanate is removied which increases the concentration of Fe+3?

Yes- removing the SCN- makes the reaction shift to the left to generate more SCN- to replace it.




Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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CaptainOfSmug
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[*] posted on 7-3-2013 at 14:25


Okay thanks for clarifying this for me! I was having trouble sort of wrapping my head around this :P
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