CaptainOfSmug
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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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DraconicAcid
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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).
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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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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DraconicAcid
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| 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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Okay thanks for clarifying this for me! I was having trouble sort of wrapping my head around this 
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