Equilibrium system given: Fe3+(aq) + SCN- (aq) (colorless) <- -> FeSCN2+(aq) (re

Question

Equilibrium system given: Fe3+(aq) + SCN- (aq) (colorless) <- -> FeSCN2+(aq) (red) Why do the following change colors as followed? Please be specific. 1. 0.1M AgNO3 turned milky white with precipitate 2.+ conc. HCl turned lighter orange 3.+ 0.1M Hg(NO3)2 turned almost completely clear 4.+1M Na2PO4 yellowish with precipitate 5.+0.1M Na2C2O4 very light yellow 6.+ Solid NaF yellow-orange with a seemingly higher viscosity The only one I think I have is the Silver Nitrate one, but I'm unsure of it as well. I think that it turned white because the ferric thiocyanate and silver nitrate together formed silver thiocyanate, which is a white solid and that overpowered the reaction and it shifted that way.

Explanation / Answer

FOLLOW THIS The Iron (III) chloride is going to dissociate in your solution into Fe3+ and 3Cl-. Your elemental iron III has an oxidation number of +3 and so does the Iron atom in [FeSCN]2+. But when you add your Iron (III) chloride and it dissociates in solution, you are effectively adding in Fe3+ ions. In order for the reaction to re-establish equilibrium, it will try to use up the Fe3+ ions you just added and the reaction will shift to the right. Therefore, you will be making more [FeSCN]2+ and the reaction will become a darker red. I think the reason you got stuck is because the iron in [FeSCN]2+ also has an oxidation number of 3+ so it seems like you would be adding iron on both sides of the reaction. But, this would be incorrect. The only way you would add iron o

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