Chemistry: Demo: Shifting Equilibrium of FeSCN2+

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We've spent a lot of time looking at various applications of Le Chatelier's Principle. Now, let's look at a really simple illustration of what it really looks like not on paper, but in real life.
What I have here is a solution of ferric ion, Fe^3+, and thiocyanate in water. And in equilibrium is the presence of a complex ion, iron thiocyanate two plus. And the color of this iron thiocyanate two plus is red-brown when it's concentrated, so it looks more like this when it's very highly concentrated, but, when it's relatively dilute, it looks like this, sort of orangish-red solution. And what I'm going to do is I'm going to try some things and make some predictions based on Le Chatelier's Principle. So, for instance, if we add ferric ion - Le Chatelier's Principle says that we should make more of this red-brown thiocyanate - so let's take a look at that. I have here a solution of ferric nitrate and I'm going to shoot some of this in. And you can see that the solution is turning more red, meaning making more of the iron thiocyanate at the expense of the thiocyanate that's in solution. So the iron three plus concentration goes up, thiocyanate goes down, but, most importantly, the iron thiocyanate goes up, and so we get more red-brown color in our solution.
Now, let's look at the related reaction, where we add the other reagent. This is a solution of potassium thiocyanate. And I'm going to add some potassium thiocyanate to this solution. Again, we'll give it a couple of shots. And you can see that we're developing more red color again, indicative of a greater concentration of iron thiocyanate.
Now, let's go the other way. Suppose we could decrease the concentration of iron. If we decrease the concentration of iron in our solution, what's going to happen? Well, Le Chatelier's Principle says that this reaction ought to shift back towards reactants, towards the things that are on the top here, and so we ought to decrease the concentration of ferric thiocyanate, and that means that the color ought to go away, or at least decrease. And the way we're going to do that is we're going to show it some hydroxide. Hydroxide reacts with ferric iron to form ferric hydroxide, which is insoluble, and so we're going to decrease the concentration of ferric ion, because it's going to precipitate out of solution. And I'm going to give this a couple of shots. And you can see that the color is dissipating, that it's gone almost entirely colorless. You can't really see the precipitates. It's so finely divided, it's just sort of floating here. But, take my word for it, there is a little bit of a precipitate. And a lot of the color that you see is the presence of the precipitate, not the presence of the iron thiocyanate complex ion.
And let's do now the final reaction, and the final reaction would be to decrease the concentration of thiocyanate in solution. And again, by Le Chatelier's Principle, we expect that we should decrease the color of the solution. And what happens is that the silver complexes with the thiocyanate to form silver thiocyanate. And the silver thiocyanate, again, is insoluble, so it forms a precipitate. And look at that. The color disappears, or at least decreases substantially. Again, it's the color of the iron thiocyanate, the red-brown, that has decreased it's concentration, so the color decreases, exactly as how we would predict based on Le Chatelier's Principle.
So to summarize, Le Chatelier's Principle allows you to predict how an equilibrium should shift. When we first set it up, before we did anything, it as an equilibrium. We had iron three plus and thiocyanate ion in equilibrium with this complex ion. And then we perturbed the equilibrium, and Le Chatelier's Principle says that the equilibrium will shift to minimize the change, and that means either making more iron thiocyanate, when we add more reactants, or making less of the complex ion iron thiocyanate, when we take away some of the reactants.
Chemical Equilibrium
Shifting Chemical Equilibrium
CIA Demonstration: Shifting the Equilibrium of FeSCN^2+ Page [1 of 1]