Introduction to Coordination Chemistry Concepts

An introduction to coordination chemistry is going to be the topic in this first lesson on a whole chapter on the subject. My name is Chad, and welcome to Chad's Prep, where my goal is to take the stress out of learning science. Now, in addition to high school and college science prep, we also do MCAT, DAT, and OAT prep as well. I'll leave a link in the description below for where you can find those courses. Now, this lesson's part of my new general chemistry playlist I'm still releasing for at least a couple more weeks, several lessons a week, and I'll start on a new playlist here soon, so if you want to be notified every time I post a new lesson or when I get started on the new playlist, And subscribe to the channel, click the bell notification. Alright, so coordination chemistry. This all kind of is centered around what we call a coordinate covalent bond. And we learned back in the day that a covalent bond is just the sharing of electrons, and typically between two non-metals is the way we looked at it. And one non-metal kicked in an unpaired electron, and the other non-metal kicked in an unpaired electron, and then they shared those two electrons as a single covalent bond. Well, gonna have a little bit different... players here, we're still going to have some sharing of electrons going on, but where those electrons come from, so and who shares what is going to be a little bit different. So instead of having two different atoms each sharing one electron, we're going to have one atom sharing both the electrons and the other atom sharing nothing, just saying thank you very much. So when we got this unequal sharing of electrons like this, we're dealing with what are called Lewis acids and Lewis bases, which we got introduced to back in the acid and base chapters. We got Lewis acids, Lewis bases. So you might recall that the hallmark of a Lewis base is that it had to have a lone pair of electrons to share, and it's going to share that lone pair as a bond with the Lewis acid. Well, in these coordination complex things, it's a little strange here. So you're typically going to see a pattern kind of arise, though. In the center of it, you're going to have a metal ion right there, and so that is the Lewis acid, it turns out, so we'll call him the central metal ion. Now technically every once in a while you might actually see a neutral metal in there, but most of the time it's going to be a metal ion. So, and then it's going to be bonded to one or more molecules on the outside, and I say molecules, they could be neutral molecules, they could be anions as well. You'll find out that most of these are going to be anions, but some of them will actually be neutral, which is going to be a little unusual. So, but they're all going to have to have a lone pair of electrons to donate to make that bond, that's what makes them the Lewis base, and we call them ligands. or ligands, depending on who you talk to. I'm gonna go with ligands, that's how I've heard it. So I'm gonna run with that. So what you're gonna find is you have a single central metal ion bonded to multiple ligands is normally the way it's gonna work. So, and we're gonna deal with three different sets of numbers of bonds coming off that central metal ion. We're gonna deal with two bonds coming off it, in which case they'd be 180 degrees apart and we'd call it linear. So the central metal ion can be making four bonds in which there are two different geometries associated with that. Square planar. where they're all in a single plane 90 degrees apart, or tetrahedral, where it's a three-dimensional structure, where they're 109.5 degrees apart. And then finally, like this one here, where the central metal ion is making six bonds, and that's going to be octahedral, same shape we've seen before, where all the adjacent ones are 90 degrees apart, or opposite ones are 180 degrees apart, however you want to look at it, but same octahedral shape we've dealt with back in molecular geometry. So those are the common geometries we're going to deal with in this chapter. It turns out there are... you know, coordination, well, let's not get there. There are geometries that are going to involve more than just six bonds around that central metal ion that we're just not going to discuss in this chapter. These are the only ones we're going to talk about. So linear, square planar, tetrahedral, octahedral, and these are going to correspond to what we call coordination numbers. So of two, four, four, and six. And the coordination number, again, is just the number of bonds coming off that central metal ion. Now, if we take a look at this guy right here, so it turns out that... Your central metal ion and all the ligands it is bonded to is what we call the coordination sphere. So, and then everything outside the sphere is not part of the coordination sphere. And it turns out in this case, this guy is a complex cation. So it's a complex ion. So, and when your central metal ion bonded to ligands has an overall positive charge, we call it a complex cation. When it's overall got a negative charge, we call it a complex anion. And when it's neutral, we'll just call it a neutral complex. So in this case, we've got a complex cation. And then we've just got simple anions, and those simple anions might be monatomic ions or polyatomic ions, but they're not going to be a complex in this case, because they're just simple chloride ions. Now, there's nothing that says you can't have both a complex cation and a complex anion. It's not normally what you're going to run into, but there's nothing that says that you can't have both. But in this case, we had a complex cation with just a simple counter ion, a simple anion in response here. Alright, so this is the coordination sphere, the metal ion and everything in between. And the reason it gets kind of singled out and gets a special name is it turns out what's in that coordination sphere may not be available to undergo normal chemical reactions. So let's write this out real quick. If we actually wrote this out, the way that we write the formula here is you start with the central metal ion. Then you list all the ligands it's bonded to. And in this case, they're all ammonia and there are six of them. And then you put what's ever in the coordination sphere in brackets. And you go cation first, anion second. And so since the complex cation, or the complex is the cation, we put it first. Had it been the anion, we'd put it second. So then the counter ions will come last. And so it's always cation before anion, just like with normal ionic compounds that don't involve complex ions. So and there's our lovely formula. Now let's say we had something a little bit different here. Let's say we gave you the formula with water as a ligand. So this is going to be a little unusual here, and you might have seen it with the ammonia here. So most ligands are going to have a negative charge, it turns out, but not all of them. And the two most prominent ones that don't are the ammonia here and the water here. We call this amine, we call this aqua in this context. And notice those are neutral molecules, and a lot of students get a little, you know, confused here because we're used to ionic bonding, we've got plus and minus, cations and anions. Well, in this case, so you're definitely going to have a cation at the central metal ion, but the ligands around it oftentimes will be anions, but they don't have to be. So with water and ammonia, those are neutral molecules, and the hallmark again is to be a Lewis base, they just needed to have a lone pair of electrons. And if you recall that ammonia's Lewis structure shows that lone pair of electrons on the nitrogen. It is that lone pair on ammonia that was used to make this covalent bond here, here, here, here, here, and here. All right, so if we take a look at this formula down here now, you're supposed to realize that it's iron, five waters and a chloride that form the coordination sphere, and then there are two chloride counter ions, and that's what you're supposed to get from this formula. And so the question we're going to deal with here is if we react this with excess AgNO3, and let's read that question verbatim. So if excess AgNO3 is reacted with one mole of this lovely coordination compound, how many moles of AgCl will be produced? And so the idea is the Ag plus here is going to be reacting with some of the chlorides here And what you're supposed to realize is that the chloride that's within the coordination sphere is not available For this chemical reaction only these two chlorides out here and so with one mole of this lovely coordination Compound even though it total has three moles of chloride ions one two three Only these two are available for the chemical reaction So only has two moles of what we say free chloride ions to react with the Ag plus and so we're only going to get two moles of AgCl, not three. Alright, so what's inside the coordination sphere versus what's outside the coordination sphere, very important. Alright, it turns out we've also got some ligands that can make more than one bond. And so, in this case over here, we've only actually got three ligands. So this is ethylenediamine. It's got two nitrogens separated in space, and both of those nitrogens have a lone pair just like ammonia's nitrogen has a lone pair. And so as a result, One molecule can actually make two bonds to the iron. So, and when this is true, we call this a polydentate ligand. And if it's two, it's specifically bidentate. It had it been three bonds to the central metal ion, it would have been tridentate, tetradentate, pentadentate, hexadentate, so on and so forth. So, but polydentate just generically here. So, but ethylenediamine is a well-known bidentate ligand. It can make two bonds. And it's not just enough to have two atoms that each have a lone pair, they also have to be separated by enough space that they can bind like say 90 degrees apart in an octahedral complex like this. So if we write the formula for this guy, we've got Fe and it turns out the abbreviation for ethylenediamine is En, and we've got one here, one here, and one here for a total of three, and then two plus. And so in this case, I don't have an entire coordination compound, so cation and anion, I've just got the cation. So we can only write the formula of the cation, I don't, I never showed you what the counter ions were in this case, so we'll just stick with just the cation. But the big thing is a lot of students will get tripped up here, and if I ask you, what's the coordination number of the iron in this compound, a lot of students get tripped up and they would say three, because they see three ligands. But again, the coordination number is not the number of ligands, it's the number of bonds being made to the central metal ion. And since each of these three ethylene diamines can make two bonds each, well, three times two is still one, two, three, four, five, six bonds being made of that central metal ion. It's still gonna have a coordination number of six, which is gonna correspond to that octahedral geometry. And so you're gonna have to memorize, you know, which of these ligands are bidentate. And there's only a handful. Most of the ligands are gonna be monodentate. And I put on the next page of the handout here, the study guide, I put a big list of them. I'll put it up on the board here as well. In this case, you're supposed to know that ethylenediamine is bidentate, so is o-phenanthroline and oxalate and carbonate. And then diethylentriamine is tridentate. And then EDTA, for short, is mostly what we call it. But ethylenediamine tetraacetic acid can make up to six bonds to a central metal ion. All right, you should also know one other term here. And that term is going to be called a chelating agent. So, and it turns out that any polydentate ligand can act as what we call a chelating agent. So, and these chelating agents, when they're bonded to the metal, we call that a metal chelate. So, and these are kind of important. They're important for a couple of reasons. So, but they sequester metal ions by binding them and surrounding them. So, in this case, this would be an example of a metal chelate right here. So, and again, this is kind of important for a couple different reasons. So, one of them might be if you get heavy metal poisoning. If you get heavy metal poisoning, like let's say you get lead poisoning or something like this, Well, lead tends to bind very strongly to your proteins in an irreversible fashion, and so once you get poisoned by lead, usually you're probably going to just stay poisoned by lead, and it's going to cause you some serious irreversible harm. However, if shortly thereafter getting that lead poisoning, maybe they're going to give you some sort of chelating agent in your system, and that's going to start binding to those lead ions before they have a chance to bind to your proteins, and as a result, because they're not going to be strongly bound to your body, they might actually get passed through your system. And so chelating agents can... kind of help with say heavy metal poisoning and things of this sort. Also, sometimes, you know, we run chemical reactions, it's pretty common in biology, like say some of these DNA reactions we do, we often have to add magnesium into the reaction to get these reactions to go. It turns out a lot of DNA reactions, in the biological sense, are dependent on having Mg2 plus around. And so you might be running a biochemical reaction in this context, and when you want it to stop, well, what we often do is just add in some sort of chelating agent. to kind of sequester off all the magnesium ions. And if there's no magnesium ions free, you know, to help out with these biochemical reactions with the DNA, well then the reaction stops. And so it's a good way to kind of start and stop a reaction in this case with the presence or absence of a chelating agent. So, but big thing here is you should just realize that any polydentate ligand could be called a chelating agent and when they're bound to that metal that's called a metal chelate. All right, on that same lovely table full of ligands, you're going to find out that it also lists their names. So, and a couple different names. Their actual name, like H2O, is water. So, but also how they're named when we name them as part of a complex ion or a coordination compound. And so, it turns out water is going to be named as aqua. Ammonia is going to be named as amine with two Ms and things of this sort. And that's going to be important because that's actually the focus of the very next lesson. We're going to learn how to name complex ions and coordination compounds in general. Now if you found this introduction to coordination chemistry helpful, then a like and a comment let me know are pretty much the best things you can do to support the channel. And if you're looking for practice and if you're looking for say final exam reviews, this usually comes at the end of the semester, or practice final exams, then take a look at my general chemistry master course. I'll be sure to leave a link in the description below. A free trial is available. Happy studying.

Alright, so coordination chemistry. This all kind of is centered around what we call a coordinate covalent bond. And we learned back in the day that a covalent bond is just the sharing of electrons, and typically between two non-metals is the way we looked at it.

And one non-metal kicked in an unpaired electron, and the other non-metal kicked in an unpaired electron, and then they shared those two electrons as a single covalent bond. Well, gonna have a little bit different... players here, we're still going to have some sharing of electrons going on, but where those electrons come from, so and who shares what is going to be a little bit different. So instead of having two different atoms each sharing one electron, we're going to have one atom sharing both the electrons and the other atom sharing nothing, just saying thank you very much.

So when we got this unequal sharing of electrons like this, we're dealing with what are called Lewis acids and Lewis bases, which we got introduced to back in the acid and base chapters. We got Lewis acids, Lewis bases. So you might recall that the hallmark of a Lewis base is that it had to have a lone pair of electrons to share, and it's going to share that lone pair as a bond with the Lewis acid.

Well, in these coordination complex things, it's a little strange here. So you're typically going to see a pattern kind of arise, though. In the center of it, you're going to have a metal ion right there, and so that is the Lewis acid, it turns out, so we'll call him the central metal ion.

Now technically every once in a while you might actually see a neutral metal in there, but most of the time it's going to be a metal ion. So, and then it's going to be bonded to one or more molecules on the outside, and I say molecules, they could be neutral molecules, they could be anions as well. You'll find out that most of these are going to be anions, but some of them will actually be neutral, which is going to be a little unusual.

So, but they're all going to have to have a lone pair of electrons to donate to make that bond, that's what makes them the Lewis base, and we call them ligands. or ligands, depending on who you talk to. I'm gonna go with ligands, that's how I've heard it.

So I'm gonna run with that. So what you're gonna find is you have a single central metal ion bonded to multiple ligands is normally the way it's gonna work. So, and we're gonna deal with three different sets of numbers of bonds coming off that central metal ion.

We're gonna deal with two bonds coming off it, in which case they'd be 180 degrees apart and we'd call it linear. So the central metal ion can be making four bonds in which there are two different geometries associated with that. Square planar.

where they're all in a single plane 90 degrees apart, or tetrahedral, where it's a three-dimensional structure, where they're 109.5 degrees apart. And then finally, like this one here, where the central metal ion is making six bonds, and that's going to be octahedral, same shape we've seen before, where all the adjacent ones are 90 degrees apart, or opposite ones are 180 degrees apart, however you want to look at it, but same octahedral shape we've dealt with back in molecular geometry. So those are the common geometries we're going to deal with in this chapter. It turns out there are...

you know, coordination, well, let's not get there. There are geometries that are going to involve more than just six bonds around that central metal ion that we're just not going to discuss in this chapter. These are the only ones we're going to talk about. So linear, square planar, tetrahedral, octahedral, and these are going to correspond to what we call coordination numbers.

So of two, four, four, and six. And the coordination number, again, is just the number of bonds coming off that central metal ion. Now, if we take a look at this guy right here, so it turns out that... Your central metal ion and all the ligands it is bonded to is what we call the coordination sphere.

So, and then everything outside the sphere is not part of the coordination sphere. And it turns out in this case, this guy is a complex cation. So it's a complex ion.

So, and when your central metal ion bonded to ligands has an overall positive charge, we call it a complex cation. When it's overall got a negative charge, we call it a complex anion. And when it's neutral, we'll just call it a neutral complex. So in this case, we've got a complex cation.

And then we've just got simple anions, and those simple anions might be monatomic ions or polyatomic ions, but they're not going to be a complex in this case, because they're just simple chloride ions. Now, there's nothing that says you can't have both a complex cation and a complex anion. It's not normally what you're going to run into, but there's nothing that says that you can't have both. But in this case, we had a complex cation with just a simple counter ion, a simple anion in response here.

Alright, so this is the coordination sphere, the metal ion and everything in between. And the reason it gets kind of singled out and gets a special name is it turns out what's in that coordination sphere may not be available to undergo normal chemical reactions. So let's write this out real quick. If we actually wrote this out, the way that we write the formula here is you start with the central metal ion.

Then you list all the ligands it's bonded to. And in this case, they're all ammonia and there are six of them. And then you put what's ever in the coordination sphere in brackets.

And you go cation first, anion second. And so since the complex cation, or the complex is the cation, we put it first. Had it been the anion, we'd put it second.

So then the counter ions will come last. And so it's always cation before anion, just like with normal ionic compounds that don't involve complex ions. So and there's our lovely formula.

Now let's say we had something a little bit different here. Let's say we gave you the formula with water as a ligand. So this is going to be a little unusual here, and you might have seen it with the ammonia here. So most ligands are going to have a negative charge, it turns out, but not all of them.

And the two most prominent ones that don't are the ammonia here and the water here. We call this amine, we call this aqua in this context. And notice those are neutral molecules, and a lot of students get a little, you know, confused here because we're used to ionic bonding, we've got plus and minus, cations and anions.

Well, in this case, so you're definitely going to have a cation at the central metal ion, but the ligands around it oftentimes will be anions, but they don't have to be. So with water and ammonia, those are neutral molecules, and the hallmark again is to be a Lewis base, they just needed to have a lone pair of electrons. And if you recall that ammonia's Lewis structure shows that lone pair of electrons on the nitrogen. It is that lone pair on ammonia that was used to make this covalent bond here, here, here, here, here, and here. All right, so if we take a look at this formula down here now, you're supposed to realize that it's iron, five waters and a chloride that form the coordination sphere, and then there are two chloride counter ions, and that's what you're supposed to get from this formula.

And so the question we're going to deal with here is if we react this with excess AgNO3, and let's read that question verbatim. So if excess AgNO3 is reacted with one mole of this lovely coordination compound, how many moles of AgCl will be produced? And so the idea is the Ag plus here is going to be reacting with some of the chlorides here And what you're supposed to realize is that the chloride that's within the coordination sphere is not available For this chemical reaction only these two chlorides out here and so with one mole of this lovely coordination Compound even though it total has three moles of chloride ions one two three Only these two are available for the chemical reaction So only has two moles of what we say free chloride ions to react with the Ag plus and so we're only going to get two moles of AgCl, not three. Alright, so what's inside the coordination sphere versus what's outside the coordination sphere, very important. Alright, it turns out we've also got some ligands that can make more than one bond.

And so, in this case over here, we've only actually got three ligands. So this is ethylenediamine. It's got two nitrogens separated in space, and both of those nitrogens have a lone pair just like ammonia's nitrogen has a lone pair.

And so as a result, One molecule can actually make two bonds to the iron. So, and when this is true, we call this a polydentate ligand. And if it's two, it's specifically bidentate. It had it been three bonds to the central metal ion, it would have been tridentate, tetradentate, pentadentate, hexadentate, so on and so forth.

So, but polydentate just generically here. So, but ethylenediamine is a well-known bidentate ligand. It can make two bonds. And it's not just enough to have two atoms that each have a lone pair, they also have to be separated by enough space that they can bind like say 90 degrees apart in an octahedral complex like this.

So if we write the formula for this guy, we've got Fe and it turns out the abbreviation for ethylenediamine is En, and we've got one here, one here, and one here for a total of three, and then two plus. And so in this case, I don't have an entire coordination compound, so cation and anion, I've just got the cation. So we can only write the formula of the cation, I don't, I never showed you what the counter ions were in this case, so we'll just stick with just the cation. But the big thing is a lot of students will get tripped up here, and if I ask you, what's the coordination number of the iron in this compound, a lot of students get tripped up and they would say three, because they see three ligands. But again, the coordination number is not the number of ligands, it's the number of bonds being made to the central metal ion.

And since each of these three ethylene diamines can make two bonds each, well, three times two is still one, two, three, four, five, six bonds being made of that central metal ion. It's still gonna have a coordination number of six, which is gonna correspond to that octahedral geometry. And so you're gonna have to memorize, you know, which of these ligands are bidentate.

And there's only a handful. Most of the ligands are gonna be monodentate. And I put on the next page of the handout here, the study guide, I put a big list of them. I'll put it up on the board here as well.

In this case, you're supposed to know that ethylenediamine is bidentate, so is o-phenanthroline and oxalate and carbonate. And then diethylentriamine is tridentate. And then EDTA, for short, is mostly what we call it.

But ethylenediamine tetraacetic acid can make up to six bonds to a central metal ion. All right, you should also know one other term here. And that term is going to be called a chelating agent. So, and it turns out that any polydentate ligand can act as what we call a chelating agent.

So, and these chelating agents, when they're bonded to the metal, we call that a metal chelate. So, and these are kind of important. They're important for a couple of reasons. So, but they sequester metal ions by binding them and surrounding them.

So, in this case, this would be an example of a metal chelate right here. So, and again, this is kind of important for a couple different reasons. So, one of them might be if you get heavy metal poisoning. If you get heavy metal poisoning, like let's say you get lead poisoning or something like this, Well, lead tends to bind very strongly to your proteins in an irreversible fashion, and so once you get poisoned by lead, usually you're probably going to just stay poisoned by lead, and it's going to cause you some serious irreversible harm. However, if shortly thereafter getting that lead poisoning, maybe they're going to give you some sort of chelating agent in your system, and that's going to start binding to those lead ions before they have a chance to bind to your proteins, and as a result, because they're not going to be strongly bound to your body, they might actually get passed through your system.

And so chelating agents can... kind of help with say heavy metal poisoning and things of this sort. Also, sometimes, you know, we run chemical reactions, it's pretty common in biology, like say some of these DNA reactions we do, we often have to add magnesium into the reaction to get these reactions to go.

It turns out a lot of DNA reactions, in the biological sense, are dependent on having Mg2 plus around. And so you might be running a biochemical reaction in this context, and when you want it to stop, well, what we often do is just add in some sort of chelating agent. to kind of sequester off all the magnesium ions.

And if there's no magnesium ions free, you know, to help out with these biochemical reactions with the DNA, well then the reaction stops. And so it's a good way to kind of start and stop a reaction in this case with the presence or absence of a chelating agent. So, but big thing here is you should just realize that any polydentate ligand could be called a chelating agent and when they're bound to that metal that's called a metal chelate. All right, on that same lovely table full of ligands, you're going to find out that it also lists their names. So, and a couple different names.

Their actual name, like H2O, is water. So, but also how they're named when we name them as part of a complex ion or a coordination compound. And so, it turns out water is going to be named as aqua.

Ammonia is going to be named as amine with two Ms and things of this sort. And that's going to be important because that's actually the focus of the very next lesson. We're going to learn how to name complex ions and coordination compounds in general. Now if you found this introduction to coordination chemistry helpful, then a like and a comment let me know are pretty much the best things you can do to support the channel.

And if you're looking for practice and if you're looking for say final exam reviews, this usually comes at the end of the semester, or practice final exams, then take a look at my general chemistry master course. I'll be sure to leave a link in the description below. A free trial is available. Happy studying.

Transcript for:Introduction to Coordination Chemistry Concepts

Transcript for:
Introduction to Coordination Chemistry Concepts