Chemistry: Complexes and Ligands

This lesson has not been reviewed.

Please purchase the lesson to review.

I have here a solution of iron-3 nitrate. Now, we would just normally write on paper Fe[3] plus aqueous and we'd be happy with that. But remember, at a molecular level, just like when sodium is in water, that iron is going to be surrounded by water molecules. But the interesting difference between sodium when it's in water and iron-3 plus when it's in water is that the binding between the iron and the water molecules is much stronger than that between the sodium - so strong, in fact, that when we dry it - we get rid of all the water; we put it in an oven - we still have the water molecules stuck on the metal. These are, in fact, very strong bonds. And in fact, there's a very definite number - so a much more ordered system than we had for the case of sodium. And, in this case, we've got 6 water molecules strongly bound to the iron-3 plus. And this one of thousands and thousands of examples of what we call coordination compounds.
So in fact, this is an entire area of chemistry characterized by a transition metal at the center of a set of ligands - ligands being either neutral molecules or, very often, anions, stuck on the metal. So, in this particular case, they're all water, but ligands can come in a variety of different flavors. Very common ligands for this type of a compound would include neutral molecules like ammonia or water but also anions, as I mentioned, like hydroxides. Cyanides are very common. Acetate, chemical - these are all examples of ligands that we refer to as monodentate - literally, "single-toothed," or they have only one place within the molecule or ion that they can attach to the transition metal.
Well, as you might guess, there's a variety of molecules also that are polydentate - multiple-tooth - that can attach to the metal at more than one point. So let's look at a couple examples of that. This guy is called ethylenediamene. And you can see the name comes from, essentially, the ethylene framework - two-carbon framework - and diamine - 2 amine groups. And this is a very common bidentate ligand for transition metals in that it has 2 places - the 2 nitrogens here - that can attach to the metal. So for instance, this is a molecule of ethylenediamene and these two nitrogens here are attached to a central metal. And this is the framework that holds them together. And can have, in fact, up to three of these on a metal at one time. I'll show you a picture of that in a moment. Other examples of bidentate ligands would include acetoacetonate, coming from a historical organic nomenclature. So don't worry about memorizing it. But I'll point out to you here that we can draw two other resonance structures of this molecule. That would be good practice for you if you haven't done that this week. And we have equivalent oxygens then, because it's a hybrid of these different resonance structures. These oxygens both can attach to a metal in the same way that these nitrogens could both attach to the metal.
Oxalate is another example of a molecule. We've looked at a dicarboxylic acid that's deprotonated. Both of these oxygens highlighted in green. Remember, these are equivalent to these oxygens also, through resonance structure arguments. But so that one oxygen from each of the two carbons can act as a bidentate ligand.
So what does this look like? Here's an example of cobalt surrounded by 3 ethylenediamenes, the ligand we just saw a moment ago. In fact, this model, just three of these guys stuck on the metal rather than just one. So this would be an example of 3 bidentate ligands and in this case, we have an example of a metal surrounded by one polydentate, or hexadentate, in this case, ligand.
So what do we have here? We've got 1 molecule - in this case, a tetra-anion - wrapped around the iron core. And you'll notice that this entire thing - that whole molecule - completely encapsulates that transition metal. Now that's a really interesting notion. I mean, It's as if the analogy would be if this is the iron metal, this is the molecule - the polydentate or often called chelating - in fact, all of this is called chelating oftentimes. This polydentate ligand wraps itself completely around the iron and sequesters it, in a sense. In fact, these are, in effect, also called sequestering agents - or sederafores - lots of different names for this. But the basic notion is that you have one molecule that can completely surround a transition metal. This has got all kinds of interesting uses. It's in mayonnaise, for instance. It's used as a preservative. It might even show up in your grandpa's hidden flask. The role of EDTA is simply to surround trace impurities of metals and stop them from acting as enzymes to decompose the mayonnaise or the alcohol or something in the alcohol or whatever you have that you're trying to preserve. Nature uses its own variation of these polydentate, or chelating agents, to sequester iron - to get iron that's normally at real low concentrations. In fact, in your intestine is a protein called transferin which is a polydentate ligand, essentially, that can wrap around iron and pull it in an make use of it. And then, bacteria have their own version of transferin. They emit a sederafore, which comes out of the cell and grabs the iron from you and pulls it back into the cell so that they can enjoy the iron. So there's this continual battle going on in your body for these small amounts of iron. And the weapons of choice are these sequestering agents, or these polydentate ligands that can wrap themselves around these transition metals.
So how do they work? Why are these so much more stable than monodentate ligands? Because that seems to be what I'm implying. Well, let me tell you about my daughter. My daughter, who I love dearly, she brings entropy a whole new meaning. She is very, very disorganized - has been since the day she was born. And we send her out with a full set of clothes and she'd come back wearing a sock or something like that. Clothes would be everywhere. And we never could find half of them. Well, this wonderful invention came about, or rediscovered, which was a pair of mittens attached through a string that went through her jacket. Now the wonderful thing about this was, as long as she was wearing any one of these pieces, the other two were right with her at all times - wonderful idea. And usually, she'd just be wearing this one mitten, for instance, and the jacket would be dragging behind.
This is a great idea and this uses the same exact principle as chelating agents. There are lots of things you can do with this idea. Let's suppose you're at a picnic. So we've got our dog here and we're ready to enjoy our food and I've got ketchup, mustard, vinegar, everything, even the dessert, and it's all tied together. And so, this handy dandy low-entropy picnic pack - this is taking advantage of the same idea, that if we tie these things all together, I don't have to worry about finding them all. And again, that's the whole principle behind a polydentate ligand, that once you've tied everything together, there's a very low cost in entropy to actually get them together to use, because they're already together.
So imagine a thought experiment now where instead of EDTA, you've got the individual pieces of EDTA stuck to a metal. You can imagine equilibrium, we could describe, where this complex comes together with EDTA to give the EDTA complex plus the 6 free ligands that have been liberated here. And think about the huge increase in entropy now once this has happened. All of these ligands that were stuck to the metal are now free to go wherever they want, and our cost was just this one EDTA molecule that came in to wrap around the metal. So there's a huge increase in entropy. There is very little change in the enthalpy, because the same bonds are being made, essentially - so no change in enthalpy, huge change in entropy, favoring this forward reaction. The equilibrium constant we know is going to be much greater than 1. So that is the stability of a chelating ligand, and this is called the chelate effect, very often, which is, for the most part, an entropy-driven process.
So finally, the last thing we have to think about is simply, okay, we're familiar then with this notion of a coordination complex. How do we actually assign oxidation states for a complex-looking thing like this? The secret is simply in your mind to pull the ligands off, bringing with them the electrons that made the bond. That's going to be real important to remember that. In the bonding that's going on between iron and water, for instance, we're going to take water off as a neutral water molecule. So it's a heterolytic cleavage of that bond. So, for instance, if we have silver dichloride, that anion, we would pull off the chlorines as chlorides. And so we'd end up with an overall charge on the silver of plus 1 in order for that plus 2 minus to equal the overall charge of the complex ion. Here's another example here. We start with an overall neutral compound, in this case. Well, we have 2 ammonias that we pull off at neutral ammonias, 2 chlorines that we pull off as chlorides. And as a result, we have platinum having to be a 2 plus ion in order for that charge to balance the 2 minus charges we pulled of for this compound to be neutral.
Let's do another example. In this case, we have a complex ion which we pull off cyanide ions, CN^- ions, and as a result of that, we have 6 minus charges. Therefore, in order to have an overall charge of 3 minus, we must have a 3 plus charge on the ion. So that would be our oxidation state of the iron, in this case. So we can go on and on. In fact, let me leave it up to you to work out these other two examples. I'll just point out that there are other neutral molecules that can act as ligands as well, like ammonia, like water. In this case nitrogen or dinitrogen is acting as a ligand. Down here, carbon monoxide is acting as a ligand. Phospheme is acting as a ligand in the same way ammonia acts as a ligand.
So the key to understanding oxidation states in coordination complexes is to disassemble them into their components, to take off the ligands, at least mentally, look at the charges of the ligands and then make sure that your oxidation state of the metal is appropriate to balance the overall charge of either the neutral molecule or anion or cation, just depending on what you have.
Once we understand how to determine the oxidation state, we know the electron configuration and we know that that is so essential to predicting chemical reactivity. So that's what we're going to do as soon as we have a little lesson in naming these compounds.
Transition Elements
Coordination Compounds
Complexes and Ligands Page [2 of 2]