The Science of Everything Podcast - Episode 130: Transition Metal Chemistry

Episode Date: July 31, 2022

I introduce the unique and diverse chemistry of the transition metal elements, also known as the d-block metals. I begin with an overview of transition metal properties and ores, and then discuss coordination complexes, ligands, denticity, chelation, coordination geometries, isomerism, and the difference between strong and weak field ligands. We then examine how crystal field theory can explain many properties of transition metals, including their unique colouration and magnetic properties. The episode concludes with an overview of organometallic compounds and ligand field theory, including how pi-bonding can explain the difference between strong and weak field ligands.   If you enjoyed the podcast please consider supporting the show by making a PayPal donation or becoming a Patreon supporter. https://www.patreon.com/jamesfodor https://www.paypal.me/ScienceofEverything

You're listening to The Science of Everything podcast, episode 130, Transition Metal Chemistry. I'm your host, James Fodor. In this episode, we're going to talk about the chemistry of transition metals, which are the metals that exist along the middle of the periodic table, which actually encompasses kind of the majority of the elements in the periodic table, which often get neglected in popular discussions and in introductory courses. So hopefully we're going to remedy a bit of that today.

In particular, I'm going to talk about, transition metals, what they are, what's special about them. And then I'm going to talk about two of the major types of structures that they form, so coordination complexes and organometallic compounds. I'm going to talk about concepts such as coordination number and the geometry of ligands, isomerism, chelation, strong and weak field, ligands as well as the different reaction mechanisms. I'll also talk about the differences between crystal field theory and ligand field theory and how those structural theories are able to help us understand some of the various properties of transition metals. And I'll apply some of these concepts to understand certain more well-known

properties of transition metals, including their magnetic properties and their reactivity and their distinctive coloration. Recommended pre-listing for this episode is episode 15, chemical bonding. You might also gain some benefit from listening to episodes 119 and 120 on computational chemistry, although that's not essential. It might provide a little bit of extra background. All right, so let's jump in and start talking about transition metals. So the name transition metal is relatively recent. It's only about 100 years old. It refers to metals that exist along kind of the middle of the periodic table, the idea being that it's a sort of transition zone between the metals on the far left of the periodic table

in groups 1 and 2, and the non-metals or metalloids on the far right of the periodic table in groups 13 to 17. So transition metals occupy groups 3 to 12 of the periodic table, so those are the bits the bit sort of in the middle from rows 3, 4, 5, and 6. Another name, which is probably a more useful name for transition metals, are the D-block elements, and they're often coloured differently on a periodic table to distinguish them from other types of metals. They're called the D-block elements

because their chemistry is mostly characterized by the D-Orbitals, which is where the valence electrons are contained in these metals, at least most of the time. technically the two series that are often shown at the very bottom of the periodic table, the lanthanide and actinide series, technically these are also transition metals, although they are rarely seen in many chemical reactions. They're pretty unusual, and some of them are only found artificially. So I'm not really going to talk about those. Those, if I do cover them, will be the subject of another episode. So we're just going to talk about the D-block elements,

which are those sort of found in the middle. So scandium to zinc on the top and so on down from that. Just to explain the name D-block elements, recall from episode 15 or otherwise we talked about the different orbitals that are found around an atom. So an orbital is essentially, you can think of it as like a region of space surrounding the nucleus of an atom where electrons live or electrons can exist stably. There are different types of orbitals which are characterized essentially by their shape and orientation in space. So the simplest orbitals is the S-orbital, which is just spherical. Then there are the P orbitals. There are three of those, and they're shaped kind of like an hourglass, so they're sort of bulged at the ends and pinched in the middle, and there are three of those.

Then there are the D orbitals. Now, those have more complicated shapes. It will be easier to just look those up. They kind of have multiple lobes and stick out in different directions, and there are five of those. The next major type of orbitals are the F orbitals, of which there are seven, and they characterize the valence electron activity in the F orbitals characterize the F-block elements, which are, as I said before, those that are located below the main part of the periodic table, the lanthanides and the actinites. Technically those are transition metals, but we're not going to talk about those here. We're just going to talk about the D-block elements that are characterized by valence electrons located in those D-orbitales. So let's talk about some

important general properties of these transition metals or D-block elements. I'll use those interchangeably. So one is that, as I already said, that chemistry is characterized by valence electrons in the D orbitals. So valence electrons, meaning electrons that are in the outermost shell and are responsible for chemical activity. They're either lost or gained or shared with other atoms as the element in question forms bonds, breaks bonds, and so forth. So those are the valence electrons. They're the things that contribute to the chemistry. The inner electrons in inner shells typically don't contribute to chemistry. So that's one important aspect. Transition metals also typically form structures called coordination compounds. So I'll talk about those

a little bit later, but that's another characteristic of the transition metals. Many transition metals can also be found in many different oxidation states in nature. The oxidation state of a metal essentially refers to how many electrons it's either lost or gained. So because metals typically lose electrons, we'll talk a bit more about that later, they tend to acquire positive oxidation states, which means that they've lost electrons, so they're sort of positively charged, right? Now, Now, they might not literally have a charge on them because often they're in a compound. They're bonded to other atoms or other compounds, right? So they don't literally have a charge.

But you can think of an oxidation state, kind of like the charge that it would have if it was just by itself, loosely speaking. So oxidation states are important for understanding transition metals. And the fact that many transition metals can have a wide range of different oxidation states from plus one to plus seven or so, depending on the compound they're in, is quite important because it means they're kind of quite flexible. There's lots of different compounds that they can be in. There are lots of different things that they can bind to or react with. So they can be found in many different states.

Transition metals also often show strong or at least noticeable magnetic properties. This is also due to the presence of unpaired D electrons, electrons in those D orbitals. I'll talk more about that later. And finally, many transition metals or compounds containing transition metals show very exotic or distinctive coloration, often dependent on the oxidation state. And once again, that is traceable or explainable in terms of the electrons that are present in those valence d-orbitales. So we'll explain all of that more in due course. But these are some of the important characteristics

that distinguish transition metals. So color, oxidation states being highly variable, magnetic properties, ability to form many different compounds, and tendency to form coordination complexes. Transition metals are widely used in industry and also in biology, so they're very important for producing metal ores. Well, we extract metal from the metal oars, I should say, and use many transition metals in all sorts of industrial applications. And they're also widely used in biology. They have uses as the catalytic centers for many enzymes. And I probably won't have time to talk about that in this episode, but we'll discuss, we've talked about enzymes before in some of the biochemistry episodes, and we'll talk about some of the details of that. in future episodes. So many of the transition metals you would have heard of, zinc, copper, nickel,

cobalt, iron, chromium, titanium, silver, gold, platinum, tugston, mercury. These are all transition metals or de-block elements. There's others as there as well that are less well known, such as hafnium and zoconium and itrium and so forth. But many of these are widely known metals and widely used in industry, and they have been throughout history. The transition metals that are furthest to the right of the periodic table are those that are easiest to reduce. So reducing a transition metal or the metal in a transition metal compound involves giving it back the electrons that it lost when it reacted with some other element. So typically we don't find transition metals in their metallic form in nature. There are a few exceptions like gold. But in general,

we don't find them in those forms, we find them in all form, which means that they are forming a compound with more highly electronegative elements. Electronativity, if you recall, is a measure of how easy it is for a particular element to grab electrons from other elements. Elements to the top and the right of the periodic table tend to have very high electronegativities, essentially because they have relatively high number of protons and a relatively small size from the proton to the outer shell of the electrons.

So they tend to pull on those electrons very tightly and grab new ones. So fluorine, oxygen, nitrogen, these have very high electron. negativities. The other extreme is on the bottom left of the periodic table, which are much larger and have sort of fewer protons relative to the number of electrons in the outer shell, and therefore they're not as able to pull on those outer electrons as effectively, and so they have very low electronegativities. So because transition metals are kind of in the middle, they're transitional, right? They have kind of moderate to low electronegativities, and so their electrons, or at least their valence electrons, tend to get pulled away by more highly low-negative elements, especially elements like

oxygen and nitrogen and carbon as well. They tend to be found in compounds where they have reacted with these more highly electronegative elements. In order to separate them out from these elements from all form to get them back into pure form, we need to reduce them. Essentially give them back the electrons that were stolen away by oxygen and carbon and nitrogen and so forth. Now to do that requires a lot of heats because you basically have to give back the energy. It's energetically favorable for the highly electronegative elements to grab those electrons. So you need to add the energy back to kind of undo that energetically favorable step. So you need to use a source of energy. And historically, the first elements that were utilized for making metal tools were those on the

relatively far right of the transition metals like copper and zinc and nickel because they're easier to reduce having a higher electronegativity. So they're kind of closer to the carbons and nitrogen elements, right, which have those high-electric negativity. So they're kind of easier to give back those electrons to. Iron and cobalt are further to the left, so further away from carbon- nitrogen, so they have lower electron-negativities, and therefore they're harder to reduce, which is why, in a sense, the Iron Age comes after the Bronze Age, right? Because, I mean, this is not universally true, but that that happened in all societies, but in general, societies found out how to work bronze before they found out how to work iron, precisely because

of this greater difficulty in smelting iron, that is in purifying iron compared to bronze or copper and such, because it requires much higher temperatures. The earliest transition metals, meaning those to the furthest on the left of the periodic table, require the very highest temperatures, because they're the hardest to reduce. You have to add the most energy in, and those are really only being used widely in modern industries, such as titanium, vanadium, and chromium. Much harder to reduce much higher temperatures again required. So that's an interesting kind of preview of some of the more common transition metals. And bear in mind that all the ones I just mentioned there are on the first row of the periodic table that contains deblock elements. So that's the fourth

row from the top. The ones further down from that become generally less well known, because they're less why they use, but still important in various applications. All right. So that's a bit of an introduction about transition metals. Now let's talk a bit about coordination complexes. Now I mentioned that transition metals tend to form coordination complexes. I didn't really explain what they are. So let's talk about them because they're very important for understanding how transition metals work. So a coordination complex is kind of a special type of chemical structure that we haven't talked about before, and you tend not to hear about too much in like intro to chemistry courses. You know, they talk about your metallic bonding between different metals and ionic bonding,

when you have a non-metal and a metal that one gives up the electron to the other, and then they form an ionic compound. Or they'll talk about covalent compounds, which is when multiple non-metals share electrical, between each other, like organic compounds, for example. Those are the sort of three main types of compounds that you talk about. But there's another type of, well, it's not literally a compound, but it's another type of chemical structure, which is called a co-ordination complex. And this is what happens when you react a transition metal with typically a non-metal. So, again, one of the more electronegative elements. And what happens is that because transition metals

are kind of larger, and they have a larger number of valence electrons, because remember we're talking d-orbitales now, there's five of those. Remember, usually when we just talked about non-transition metals, like on the far left of the periodic table or when we talked about non-metals, there was the rule of eight, right? You needed eight electrons to fill up a valent shell. So that's not true anymore. That rule of eight goes out the window for transition metals. Usually they require 18 electrons in the valence shell to fill them, although that's not always sure either. Sometimes it's 16 and they're occasional exceptions to that. But as a rule of thumb, you can think of them requiring 18 electrons in the Atomos shell to fill it up. instead of only eight, which, again, we would be used to for other elements. So the fact that there are so many more valence electrons now essentially means that there is a lot more room for kind of chemistry to happen around the edges, so to speak, at the valence shell, because there's many more electrons involved, twice as many.

So the fact that there's kind of more electrons to go around, there's more room for stuff to happen, what we tend to form instead of like a regular kind of covalent compound or something, is we form what's called a coordination complex. And this is where you have the transition metal atom at the center, and then it's surrounded by other species. They could be elements or other molecules, which are called ligands. So a ligand is just a chemical species that binds to something else.

In this case, it's binding to or interacting with the central metal ion. The ligand will donate electrons to the central metal ion, because remember metals are always deficient in electrons. They're on the left or center to left side of the periodic table. So they're looking for extra electrons to fill up their outermost shell. In this case, they need to get 18 typically, not always, but let's say usually 18 electrons. Again, not eight as before. So they need to have a bunch of ligands come along, which have extra electrons to kind of donate those electrons to the central metal ion, thereby filling up its outer shell.

So ligands are effectively any atom or compound that can donate electrons to the central metal ion. And as I said, typically ligands consist of nonmetals of some form, which have more highly electronegative elements. So they tend to grab the electrons of the central metal ion and donating their own electrons, kind of filling up the outer shells of both elements. So it's kind of like a coordination complex and the ligand interactions with the central metal ion. They're kind of like a covalent bond, but they're different. The chemistry is different, so we shouldn't really think of it that way. It's not a regular covalent bond. It's its own thing, right?

So we call it a coordination complex. As I mentioned before, the central metal ion typically has multiple ligands associated with it. And each of those ligands may donate electrons to the metal ion via one atom or multiple atoms from the same molecule. Each atom that directly binds to the transition metal ion that gives up electrons to it is called a donor atom because it's donating electrons, right? Makes sense. The total number of donor atoms that are bound to the central metal ion is called the coordination number of that metal complex. So the coordination number is the total number of atoms from ligands that donating electrons to the central metal ion.

You could have all of those donor atoms on a single ligand, that can happen, or they could be spread across multiple ligands in really any combination you like. It just entirely depends on the ligand and the metal ion. To give us a little bit more vocabulary for understanding that, there's another concept called denticity. The denticity of a ligand refers to the number of donor atoms on that ligand or in that ligand. So denticity is related to the word dentist, right? So it relates to bite. So the idea is how much bite does that ligand have with respect to the metal ion? If it has a dentistry, of one, then that means it only

donates electrons to the central metal ion via one donor atom. So it has relatively little bite, and we call that a monodentate ligand. It only has one bytes worth, if you like, of strength with respect to the central metal ion. A bi-dentate

ligand donates electrons to the central metal ion via two different atoms on the same molecule. Tri-dentate has three different atoms on the same ligand molecule, which all donate electrons to the central metal. metal ion. So the way that you can calculate the coordination number of a given transition metal in a particular complex is just to add up the denticities of all the ligands that are associated with it.

Often there'll all be monodentate and then there's just one each, right? But then sometimes they're bi-dentate or tridentate ligands or there can be even more. The idea there is that the higher the density of the ligand, the tighter it tends to bind to the central metal ion. And you can kind of understand that, right, because in order to then become disassociated with the central metal ion, all of those two or three or how many bonds with that metal ion have to be broken at once. And that's kind of difficult, right, to break them all at once. And so they tend to bind very tightly when you have high densities. This is closely related to a phenomenon called chelation, which is just what we've been talking about. It's the additional bonding

strength of a ligand to a metal ion when that ligand is able to form multiple bonds with a particular metal ion, so-called keylating agents or keyletors such compounds, usually organic compounds, that are able to form a bunch of these bonds with the same metal ion, often six, and thereby have a very strong, form a very strong bond with it, and they're able to extract metals from, you know, a solution or from a biological substrate. So keylating agents are sometimes used to treat heavy metal poisoning in humans, and they have other applications as well when you need to extract metals from an environment. So let's now talk about sort of the structure of coordination complexes a bit more and introduce the notion of coordination geometries.

So in all coordination complexes, you've got a central metal ion. It's called the central because it's kind of, well, it's central to the whole structure. Also, it's literally at the center. So imagine that's always there, the big central metal ion. And then surrounding it, there will be some number of donor atoms. Donor atoms may be part of the same ligand, they may not. let's put that to the side for the moment, and just think about them as individual atoms. So it's often helpful to just start with thinking of them as monodentate, so just donating a single

electron pair. There are many different geometries that you can have, depending on how many of these donor atoms there are, surrounding the central metal ion. The most common type of coordination geometry is called an octahedral geometry, and that consists of, you know, the sort of simplest way you can imagine surrounding a central metal ion, which is sort of, they're all roughly spherical, right? With ligands, right? There's one above, one below, one to the left and one to the right, and one to the back and one to the front, right? So it's one either side in all three dimensions, so that gives you six in total. It's a little confusing with the name, but an octahedral geometry or an octahedral coordination complex has six donor atoms surrounding

the central metal ion. Often there'll be six ligands as well, but if they're all monodentatial ligands, so they all donate a single electron pair with a single donor atom, but they may not be, right? So you may have, if you had, for example, an octahedral geometry, that could be because there are six mono-dentate ligands, or it could be three bi-dentate ligands, or two tridentate ligands, or it could even be, and this isn't very common, but it could happen, you could have one hexadentate ligand. So hopefully that gets the point across, right? So the number of donor atoms that you have is dependent on the coordination geometry of the metal iron complex in question.

So octahedral is the most common with a total coordination number of six. Another very common one is square planer. So a square planer structure is so called because it forms a square on a plane. So unlike an octahedral structure, it's not kind of three-dimensional. The bonds all exist on a single plane

and basically the coordination complex forms a square with the central metal atom at the center and then at the corners of the square, each corner has a donor atom there. So it's quite different to an octahedral structure where it's kind of like it's at the center of a cube and you have donor atom at the center of the face of each cube. A square planar is very different. You have the central metal ion at the center of a square and then a donor atom at each of the corners of that square. So very different geometry.

And in that case, then, you obviously have, it's a four coordinate structure or the coordination number is four for a square planar. complex. Another four-coordinate geometry is tetrahedral, where you have a single central metal ion, as always, and it's surrounded on all sides by equally spaced donor atoms. So there are still four of them, but instead of being spread across a plane like they are in the square planer case, they're equally spaced around it to form a tetrahedral structure. That's a very common chemical structure, this tetrahedral. So you've got the four coordinate geometries, tetrahedral and square planar, and the six coordinate geometry of octahedral. Those are by far the most common.

There are some others as well. There are two coordinate complexes. They can be linear or bent and other higher coordination geometries that can sometimes occur, but we're not going to worry too much about those. We will focus mainly on the octahedral geometry and occasionally discuss square planar as well. But obviously it is important to bear in mind that the properties of the transition metal complex, the coordination complex are going to be dependent on the coordination geometry, which in turn depends on the central metal ion, its oxidation state, and the ligation.

that it's associated with. Now, just to make things a little bit more complicated, it's not enough to specify a coordination complex to just say, well, here's the metal ion, here's its oxidation state, and here are the ligands that are involved, and then we're done. Well, we're not done, right, because there's also this phenomenon of isomerism. Now, I've talked about this before. I don't exactly remember in which episodes, but isomerism is a phenomenon in chemistry whereby we have the same chemical formula, but different chemical properties, because of different arrangement of the same ingredients. So you think of a chemical formula is telling you what the ingredients are for that molecule or compound, like a simple example, H2O, two hydrogens, one oxygen.

However, that's not always sufficient by itself to fully specify the compound or the chemical structure, because the same ingredients can be combined or arranged in different ways. and that occurs very commonly for organic compounds, but it also occurs for coordination compounds. It also occurs for coordination complexes. So I'll just briefly talk through the four main different types of isomerism that occur for coordination complexes. Don't worry too much if you don't quite understand the differences between them, but it is important to have, again, in the back of your mind when you're thinking about coordination complexes, that it is important to think about the three-dimensional arrangement of the ligands as well,

not just which ligands are present. So let's go through them. The first type is called coordination isomerism, and this is pretty easy to understand, right? Although I do have to take a step back and introduce the notion of a counter ion. So coordination complexes are often charged, meaning that you get a central metal ion, a bunch of ligands associate with it, sharing electrons, like donating electrons, and therefore sharing them with the central metal ion. But that whole complex often still has a charge. Often it's still positively charged. Sometimes they can be negative, but usually they're positively charged, meaning that there's still a deficiency of electrons.

So the complex as a whole has a charge. Therefore, in most context, it won't exist by itself. It will be associated with a counter-ion. Just like when you have ions in a solution, they will almost always balance each other out so that you'll have just as many positive as negative ions, because in nature, large accumulations of charge don't exist. They're not stable, right?

They tend to equalize each other out, so other ions will be, sort from somewhere else. The counter ion for a coordination complex is some ion, like any ion, really, theoretically at least, with the right charge to counterbalance the total charge on the coordination complex. So for example, if the coordination complex has a charge of plus one, then maybe it will associate with the counter ion of chlorine or bromine, which have charges of minus one. Critically, the counter ion isn't part of the coordination complex, is just associated with it, either in a lattice or in solution.

if it's dissolved, but either way, there'll be a one-to-one relation between the number of these counter ions and the complex, but the counter-iron isn't part of the complex itself. Now, coordination isomerism is when you can have a swap occurring with a ligand that swaps positions with the counter-ion. So, for example, if you have a central metal ion with bromine as a ligand and chlorine as a counter-ion, they can interchange and chlorine becomes a ligand and bromine becomes the counter-iron. They can do that because they have the same charge. obviously they'll have to have the same charge, otherwise the whole structure will be disrupted, right? So that's important, right? Because in one case, the bromine is part of the coordination compound,

the coordination complex, and the chlorine isn't. But in the other case, the chlorine is and the bromine isn't. So it makes a difference. And although the overall chemical structure is the same, I should say the overall chemical composition is the same, the structure is different. So that's coordination isomerism, swapping around counter ions with ligands. Then there's linkage isomerism. This occurs with certain ligands which can actually have different atoms serving as the donor atom. So one example is nitrogen dioxide. That's nitrogen and two oxygens. For that compound, because nitrogen and oxygen are both highly electronegative,

either of those can actually serve as electron donors to the central metal ion. So you can have a situation where the nitrogen is the donor atom or one of the oxygens is the donor atom. It's the same ligand, but the actual donor atom, is different in the two cases. So that's called linkage isomerism because you've changed the link, the connection between the central metal ion and the ligand. You haven't changed the ligand, but you've changed the way that the ligand is connected to the central metal ion. Then we have geometric isomorism. So this is when you have the same set of ligands, but they're arranged differently surrounding the centre. Now this is quite difficult to explain without a diagram,

but for both square planar and octahedral geometries, there are non-equivalent ways of arranging two different ligands surrounding the centre. If you want to think of the octahedral case, that's remember when you have one either side, one front to back and one up and down. Suppose I have two different monodentate ligands, right? So one adentate, it only donates one electron pair to the central metal ion. So for an octahedral structure, if I have two different monodentate ligands, four of one type and two of another type, just to take a simple example, there's two different ways of arranging those around the central metal ion.

One way to do it would be to put all of one type of ligand in a plane, right? So we can think of that as side to side to front to back, right? And then I put the different ligand two, let's say, I put that top and bottom. So in this arrangement, ligand one is in a plane, and then ligand two is top and bottom. But a different way to do that, I could put the two ligand twos next to each other and then just spread the ligand ones over the rest of it. See, the difference is in the first case, which is called the trans configuration, the two different ligands are actually not next to each other.

They're far apart. They're on opposite sides of the complex. Whereas in the second configuration, which is called the cis configuration, the two odd ligands out, so to speak, are actually next to each other as close as they can get. Again, that's a little hard to visualize without a diagram, but hopefully the basic idea makes sense, right, that there's actually two different configurations which are not equivalent to each other. Same ligands, but different ways of arranging them that are not equivalent. You have to break a bond and rearrange it to actually get the same thing. You can't just rotate between

them. There's a similar type of isomerism called the FAC and MIR isomers from the facial arrangement and the marital arrangement, which are the technical terms. So ciss and trans and fack and mure, they're often referred to. So don't worry about the details of that. The important point is that the arrangement of the ligands around the central metal ion can make a difference as well to the to the particular isomer. And the final type of isomerism that we'll talk about is the most difficult to explain, and it's called optical isomerism. And optical isomerism is effectively the same thing as chirality, or it's an instance of chirality, which we've talked about before when we talked about organic chemistry. Two optical isomers are identical, except that one is a non-superimposable

mirror image of the other. One effect of this is that they rotate plain-polarized light in opposite directions. I won't go into the details of explaining that, but this is a very important, but also kind of subtle phenomenon. In a lot of reactions, optical isomers behave the same as each other. But there are certain chemical interactions and also certain physical processes, such as polarization of plane polarized light, where you can distinguish between the two different isomers. So that's all I'll say about that here. Look back on to the episodes we've done on organic chemistry to hear more about chirality. That then brings us to coming back to a summary of the different types of isomerism that I've talked about. So

coordination isomerism is when you swap around a ligand with a counter ion. Linkage isomerism is when you have the same ligand, but a different donor atom is actually connected to the central metal ion. Geometric isomerism is when you have the same ligands, but they're connected to the central metal ion in a different arrangement that's not equivalent. And then optical isomerism is when you have two non-superimposable mirror image versions of the same thing. So those types of isomerism are all important for characterizing a coordination complex, in addition to the identity of the metal ion, in addition to oxidation state,

and in addition to the identity of the different ligands, bearing in mind their denticity, so the number of donor atoms that they have. So you can see why there are such a wide range of different coordination complexes, because there's so many variables at play here, that can be different depending on the specifics of the case. Now, there's another very important property of coordination complexes, and actually it extends beyond coordination complexes to organometallic compounds as well, which I'll talk about later.

And this refers to whether a ligand is a strong-filled ligand or a weak-filled ligand. So remember, the ligand is just a compound or an atom of some kind that buys to or forms a bond with the central metal ion, and ligands will donate electrons to the central metal. ion. The central metal ion is trying to get extra electrons to fill up its outer shell, and the electronegative elements also want to grab up the metal ions, extra electrons, so they're highly electronegative, so they're trying to grab those up. So they kind of both grab up the electrons, sort of sharing them between them, and therefore form a bond, arranged in those sort of different

coordination geometries that we talked about. Again, typically the octahedral arrangement, which will be the main focus, where there's six ligands, or at least six donor atoms surrounding the central metal ion. So all of these ligands are trying to, a sense they're trying to donate electrons to the central metal ion, but different ligands have different effects on the orbitals of that central metal ion. Now to explain this, we need to think a little bit about the nature of the valence electrons surrounding the central metal ion. Remember that I said that transition metals are also are called D-block elements, and they're called D-block elements because their valence electrons are found in D-orbitales. So there are five D-orbitales, which are distinguished from each other,

essentially by their orientation and the shape of where the electrons are found, the electron clouds. Also, recall that there are up to two electrons in each of the orbitals. So five orbitals, five-d-orbital means a total of 10 electrons. If you have an atom all by itself, not bound to anything, these five D-orbitales will all have the same energy. The technical term for that is degenerate. So they're all the same, energetically speaking. So that means that an electron doesn't care which of these orbitals it goes into. They're all energetically equivalent, right?

So you can imagine five lines next to each other written out on a page, which is how you write the five orbitals. And you can write electrons as arrows. They're written as arrows because one will be spin up and one is spin down. I've talked about that in previous episodes. Don't worry if you're not quite sure what that is. It's just a property that electrons have. In each orbital, you can have one that spin up and one that spin down.

So up to 10 in total in these five d-orbitals, right? However, as soon as you bring ligands into the picture, there's a complication because no longer are all of these de-orbitales degenerate. They're not energetically equivalent anymore. Some of them increase in energy, so they become higher energy, and some of them decrease in energy, so they become lower in energy. Now, the average energy of all of the orbitals has to be the same, but they can split. You can think of this as a splitting of the orbitals, at least that's

kind of how I think about it, right? Splitting in the sense that they're no longer energetically equivalent. Some of them have higher energy, some lower energy. And the way it typically happens in an octahedral coordination geometry is that two of the orbitals get higher energy and three of them have lower energy than they would be if the atom was just by itself. So remember, this is an effect of the coordination complex being formed. Once the metal ion comes into association with the ligands, those ligands change the energy levels of the orbitals. So instead of all being the same, now there are two that are going to be higher energy and three that are going to be at a lower energy level than before. The average will be the same, but the energy of each orbital specifically

is now different. When I say that two are promoted and two sort of go down to, when I say that two have higher energy and three have lower energy, that's true for octahedral geometry. It's different for different types of geometries, but let's just focus on octahedral for the moment, because it's the most common and the easiest to talk about. Now, this happens, this splitting of the energy levels in the D-orbitales occurs for all coordination geometries and all ligands, right? But different ligands have a different size effect on the size of the split between the high and the low orbitals. So they all cause the split, but some cause a bigger split than others. And this is what we call strong versus weak-filled ligands.

So strong field ligands cause a big split between the energy levels, whereas weak field ligands cause only a small split between the energy levels, and of course there's some that are in between, that are kind of moderate. Now remember that they all cause this splitting of the energy levels, some orbitals to have high energies than otherwise, and some to have lower energies than they otherwise would, but the strong ones do it more and the weak ones do it less. And this turns out to be critical for many of the most interesting

and important properties of coordination complexes, because you don't have this sort of behavior for non-transition metals or non-metals because their valence electrons are not in the de-orbitales, right? So you don't have this splitting that emerges. It's just different behavior. The identity of the metal ion matters as well. So certain metal ions in certain oxidation states have stronger fields than others, but often you mostly focus on the ligand.

So the metal ion matters and its oxidation state matters, but the ligand is typically more important, at least for a lot of purposes that we look at, and so that will be mostly what we discuss. So that forms a good transition point to talking about the electronic structure of D-block metals. I guess we've already been talking about this to some extent, but the difference between strong-field and weak-field ligands is kind of a property of the ligands. Now we're really going to focus on what effect this has on the structure of the electrons surrounding the central metal ion.

And to understand this, we introduce a third. theory that's called crystal field theory. Now this is an older theory, but it still checks out. It's still useful to understand the structure of de-block metals. We'll introduce ligand-field theory later, which is a more kind of updated version. But let's start with the simpler version first, crystal field theory. Now, the basic idea of crystal field theory is that we take the de-orbital of a transition metal, and we sort of slot them in the context of the particular coordination geometry of, you know, whatever the geometry the transition metal complex adopts. Again, normally we're going to talk about octahedral, so let's go with the octahedral geometry.

The analysis is going to be different if it's a different geometry, but let's go with octahedral. Now, it will be too difficult for me to explain this because you just sort of have to see the diagrams to appreciate this, but I'll give you the result. And the result is this, when we kind of superimpose the geometry of where the ligands are situated in the octahedral complex over the physical three-dimensional shape of the D-Orbitals, what we find is that three of these orbitals are arranged such that they point in between where the ligands are, whereas two of the orbitals are arranged such that the orbitals point directly to where some of the ligands are located. So let's think about what consequences as.

When I say the orbitals point in that direction, I mean the clouds of electron density, electrons spend most of their time or where the electron density is highest, are sort of clustered either near to or away from the side of the ligands. Crystal field theory makes a simplifying assumption which kind of assumes that each ligand site is sort of just a bunch of electron density located at a point where that ligand is surrounding, you know, in the octahedral geometry of the central metal ion. So that means that you don't actually want the de-orbital, the electron density of those de-orbitals pointing near the ligands, right? Because that's going to be electron density near other electron density. That's energetically unfavorable because electrons repel.

So what you want to have, in terms of being energetically favorable, is for the lobes of your electron density of the de-orbital to point as far away as possible from the sites of the ligands. So you want them to point in between the ligands, in the kind of the gaps between where the ligands aren't scented. And it turns out, looking at the geometries, that that happens for three of the de-orbitales, but not for two of them. And so those two de-orbitales form the de-orbitales that are higher in energy. Remember I said two are higher in energy than the normal? And then three, the three where the de-orbital point away from the ligands, those are lower energy at the normal. So you can understand that from purely the geometry of the de-orbital and the location

of the ligands. It turns out for a tetrahedral arrangement, that was for octahedral. For a tetrahedral is actually the reverse. You've got three that go up and two that down. But let's not confuse things. Let's focus on the octahedral. Two have higher energy and three have lower energy. Average is the same, but two are higher and three are lower. So crystal field theory allows us to understand why this is the case. It's actually because of the three-dimensional arrangement of the electrons in the metal valence de-orbitales relative to the electrons donated by the ligands. And you can describe quantitatively how much effect this has on the valence electron energy of adding extra electrons and get fairly accurate results, which is pretty cool.

But there are many other effects of this as well, including on colour and magnetic properties. So let's talk about those in turn. First of all, colour. Remember that I mentioned that transition metals and coordination complexes in particular often have characteristic colours, and that can change based on the oxidation state of the metal. Now, the colour of a material is determined by the wavelength of photons, that it is capable of absorbing, or photons of visible light. So if it can absorb a particular wavelength, that means that it doesn't reflect that wavelength of color,

and so the color of the photons that hit it and bounce off will not be of that color, right? So you won't see that color in reflected light. A black substance is one that absorbs all, or nearly all, of the incoming light. So it doesn't reflect much of any color, and so you see it as pretty dark. A white object is one that reflects all of the colors,

or most of the wavelengths that are incident on it. And then objects that are particularly colored will absorb some wavelengths but not others. And so the result you'll see is a mixture of the colors corresponding to the wavelengths that it reflects. Now, coordination complexes have valence energy levels of those de-orbitales, which are at the right level to absorb

a number of different wavelengths of visible light. And exactly which wavelengths of visible light, absorbs depends on the precise energy gap, which is the magnitude of the field strength, right? So remember strong field ligands separate the gap more, weak field ligands, less of a gap, always a gap, but it's the size of it that varies. And also the element of the central metal ion and its oxidation state matter too. So all of these things and the ligands determine the exact size of this gap. And that in turn determines which wavelengths of light are absorbed by the complex and which are reflected. And so the colour.

is determined then by whatever's left over. Whatever's not absorbed is reflected, and then so we see the color of the complex as a combination of those reflected colors. And so that's essentially why you have so many different colors of like different oxidation states of metals in coordination complexes. It's because slight changes in the ligands that they're bound to

or the oxidation state can slightly change the energy gap between the less energetically favorable and the more organically favorable of the valence deorbital thereby causing it to absorb different lights, causing it to reflect different wavelengths of light, and thereby appear a different colour. And so that's really cool, right? We can understand the different colours of these elements based on one sort of unifying principle from crystal field theory, based on the physical geometrical arrangement of the d-orbital's relative to the ligands. Another sort of payoff of this theory is that

we can understand the magnetic properties of transition metals as well. I mentioned before that many transition metals are magnetic, or at least can be magnetic. Particularly, many are paramagnet. which means that they interact with an external magnetic field, but relatively weakly. And the reason for that is because paramagnetic materials, individual complexes in this case, individual complexes interact with the magnetic field, but they don't kind of all line up with each other. So each complex you can think has a magnetic dipole. The whole complex itself is like a little magnet, I guess you can think of it that way.

But they're all pointing in slightly different directions because of the different. orientations of the complexes in the material and therefore they don't line up properly. So at an individual complex level you see a magnetic effect but the material as a whole doesn't show a magnetic effect or at least it shows a relatively weak one because they don't line up with each other. Paramagnets are different from ferro magnets. Ferro magnets like iron magnets is where you have those individual like complex level magnetic dipoles but they all line up across larger regions of the metal and then you see like macroscopic magnetic behavior. So most transition metal complexes aren't ferromagnets, but many of them

are paramagnetic. The extent to which they exhibit magnetic behavior is dependent as is their color on the gap, the energy gap between those deorbital, right? So strong field versus weak field ligands again. And it works like this. Remember I said that typically transition metals will like to have 18 electrons in their valence shell, of which 10 will be up to 10, if you get all 18, then it will be 10, will be found in the de-orbitales, because there's five of those, and they can hold two each, right? But that doesn't mean that they always have that number in their outer shell, because they can often form ions, for example, which are electron deficient. So it really depends on the specific case. They like to fill up with 18, but they certainly

don't always do that. anyway so the number of electrons that we have in our five valence deorbitales depends on the oxidation state of the metal and the ligands that it's bound to so it could have up to 10 but often they have less than that and this is the fact that i mentioned before that coordination complexes can often have different oxidation states and form a wide range of different compounds that's it a bit different to organometallic chemistry which i'll talk about later but here we're talking about coordination complexes So number of electrons actually in these de-orbitales can vary quite a bit. As such, often they're not all full.

Because they're not all full, like all the orbitals aren't full, we have to decide, well, where do the electrons go? If I start adding them, where do they go first? Well, we know that they're going to go to the lowest energy ones first. And remember I said that in the octahedral complex, there are always three that are lower energy and two that are high energy. So you'd think, well, surely the electrons are going to go to the bottom three first, and then only when those are, filled, will they start filling the top two? So you'll have up to six in the bottom, and then they'll start filling the top. So you might think that's how it works, but actually that's not always how it works. And the reason for that is essentially a quantum mechanical reason, which we won't really get into here,

but basically, electrons don't like to share their orbital with another electron. Let's put it that way. Now, they can and will, if there's nowhere else for their pair electron to go, but they prefer not to. it's energetically disfavorable. So what tends to happen is that one electron will go in one orbital and then another will go in a different orbital and so on. They'll fill up one in each orbital, and only when each orbital already has an electron in it will then the second one go into each of those orbitals.

But there's a limit to that as well, right? Because suppose there's a huge energy gap between one orbital and the other. Well, then the second electron, well, I was going to say it has a choice. Obviously it doesn't really have a choice, right? but there's two countervailing factors that are going on here. One is it's energetically disfavorable to pair up with another electron in the same orbital. You can have up to two. You can't have any more than that, but you can have up to two, but it's energetically disfavorable.

On the other hand, so you can then go into the free orbital, the one that has no electrons in it, but that's at a higher energy level. So there's a trade-off here, which is bigger? Is it the disfavorable energy effect of pairing up with a second electron in the same orbital, or go into a higher energy orbital that has no electrons in it, but it is a higher energy. And obviously, it's just the matter of which of those effects is bigger. You measure the energy difference and which is larger, it goes into the other one, right? So whichever is energetically favorable, that will happen.

And here's the key point of all of that. The disfavorable effect of pairing up two electrons in one orbital is pretty much the same across the board. But the energetic disfavorability effect of moving from the, lower energy, you know, the three lower energy deorbitals, are moving from that to one of the two higher ones. That depends on the field strength, right? Is it a weak field ligand or a strong field ligand? Because if it's only a weak field ligand, that means that the energy gap between the three lower ones and the two higher deorbital is pretty small. And so the electrons actually will fill up those higher energy orbitals first before they pair up with any of the electrons in the lower

energy orbitals. Because, yeah, it's a little bit energetically disfavorable to go to the higher ones, but it's even worse to pair up with an existing electron. So we'll go to the higher ones first. That's what happens in a weak-filled ligand complex. In the strong-field ligand case, the energy gap is much larger, and it's too disfavorable for those extra electrons to go from the low-energy orbitals to the higher-energy orbitals. So they don't. They pair up instead. Weak-field case, electrons don't pair up. They go singles in each of the orbitals, whereas strong field, the gap is too big and they do pair up. Okay, but why do we care about this?

I mean, why does it matter if they pair up or don't pair up? Because when electrons pair up, there's always one that spin up and one that spin down, so the spins cancel each other out. And that means that the magnetic moment cancels each other out as well. Basically, the magnetic properties of the complex depend on the spins of the electrons. Because each electron effectively has an intrinsic internal magnetic pole to it. but if it's paired up then it's always paired with one that has the exact opposite magnetic pole and they cancel out so paired electrons have no net magnetic pole magnetic dipole but unpaired ones do and so when you're trying to work out what are the magnetic properties of a compound as a whole or like a coordination complex as a whole we need to

look at how many unpaired electrons are there that's really what counts the paired ones don't do anything, magnetically speaking, the unpaired ones do. That's what connects us to the weakfield versus strongfield ligands. Because in the weakfield ligand case, remember, the electrons don't pair up. I mean, they will pair up if, you know, if all of the orbitals get full. But I'm talking about the case where they're not all full, right? Let's say there are five electrons. In the weakfield case, there'll be one electron in each of the five orbitals. Because, you know, it has, sure, some of them have to go to a slightly higher energy level, but that's still favorable instead of pairing up, right? So in that case, you actually have five unpersersers.

electrons. And that gives rise to a sizable magnetic moment. That's called a high spin complex, because the total unpaired spin is high. Whereas in the strong field case, the gap is too high. It's too energetically disfrable to jump up to the higher energy orbitals, and so the electrons pair up instead. So if you had five electrons in these de-orbitales, two of them would pair up, two would pair up, and then there'd be one that was left over. In the strongfield case, you've only got one unpaired electron instead of five in the Wigfield case. So five versus one, that's a big difference in magnetic properties, right? And that's why we have the high spin complex, we call a high spin case when there's five unpaired, and the low spin case when there's only one

unpaired. So what is the significance of all that? Well, that's effectively what determines the magnetic properties of the complex as a whole, because if it has a high spin, that means that it's going to have the whole complex has a magnetic moment, which means it can interact. with an external magnetic field. Whereas low spin complex has a much lower, I mean, it still has a magnetic dipole, but much smaller ones, so it reacts much less with an external magnetic field.

And if there are no unpaired electrons in it, it's not paramagnetic at all, so it won't react in that same way with an external magnetic field. Whether a macroscopic sample of that compound is, exhibits magnetic behaviors, depends on additional properties like whether neighboring complexes

kind of align with each other. That's what differentiates. pharomagnets from paramagnets, but nevertheless, the magnetic properties of transition metal complexes still depend critically on the number of unpaired electrons in those d-orbitales, which in turn depends on whether the complex has weak field versus strong-field ligands. So we can explain the magnetic properties, the color properties, and many of the other chemical properties of transition-mental coordination complexes, purely based on crystal field theory, which basically just says compare the orientation in space of the D-orbital, of the 5D orbitals,

with the spatial orientation of the ligands surrounding it. Again, we've talked about the octahedral case, but you can do the same thing for the tetrahedral or the square planar case or other rarer cases as well, and a similar sort of analysis will follow. So crystal field theories are really powerful way of understanding the electronic structure, and therefore many of the chemical and physical properties of coordination complexes. So in the final part of today's episode, I'm going to talk about organometallic chemistry. And this is a bit of a step up in terms of complexity that there's some additional elements here

that we'll have to explain. But just to distinguish it from what we've been talking about so far. So organometallic chemistry is still about the chemistry of transition metal elements. So we're still talking about those same set of elements. The difference is that coordination complexes are really what happens when you get a transition metal and you associate it with ligands that are mostly inorganic non-metallic chemistry. We're now taking those transition metal ions and associating them with organic compounds. Now there's not necessarily a hard and fast distinction here because, you know, some,

in some cases it's a little unclear exactly which category you should go in, but there are general distinctions, right? So that's the big difference. Organom metallic means that we're combining an organic compound or ligand of some sort with the central metal ion. And organic means it's carbon containing. So a typical definition is that an organic metallic compound contains at least one carbon to metal bond. So there's bonding of carbons straight to the metal ions. And there may be other elements involved as well in the organic ligands, but it involves at least one carbon to ligand bond. There are a number of

important differences between organomotallic compounds and coordination complexes. So as I mentioned, And the crucial difference, which is sort of what's used to define them, is that organometallic compounds involve a bond between at least one carbon atom and the metal ion, whereas coordination complexes don't. Other properties that differ include the fact that coordination complexes are usually charged. So coordination complexes are usually ions, which then have a counter ion, whereas organomotallic compounds are typically neutral compounds. Coordination complexes also have a highly variable de-orbital count.

So I mentioned that there's lots of different possibilities, depending on the oxidation state of the central metal ion, for the number of electrons in the deorbitales. Whereas in organometallic compounds, it's much more common to have a fixed count of electrons in the deorbitales. And indeed, it's more common for them to be full in the compound. So to obey the 18 electron rule that I mentioned before. That can be useful for some coordination complexes, but it's a lot more useful for organometallic complexes. So it's more likely that they sort of fill up those valence orbitals and have the full complement of 18 or sometimes 16 electrons. And also, coordination complexes are usually soluble in water.

That's related to the fact that they're charged, whereas organometallic compounds being neutral are soluble in organic solvents. So organometallics behave a bit more like organic compounds, and the chemistry is sort of closer to that, whereas coordination complexes are kind of more like inorganic, pure inorganic chemistry, if you like. All right, so those are some of the differences, and that's why they kind of studied a little bit separately, because again, the chemistry is kind of different. There are some similarities, of course. They both involve transition metals, and as we'll see, that you can also apply crystal field theory and ligand field theory to both of them. The concept of ligands and denticity and coordination geometries, all of those things are still relevant, but now there's some additional twists, which I'll talk about here. So one of the important things that we need to do when we're talking about organometallic chemistry is to, introduce ligand field theory, which I've mentioned a couple of times, but I haven't really explained what it is yet.

Ligand field theory is kind of a modern extension or an update of crystal field theory. Crystal field theory is in a sense quite simple. It only considers the geometry of the ligands. It treats each ligand as really a point of electron mass, which is not realistic. It's good enough for many purposes, but it's not sufficient and it does break down at a certain point. Ligand field theory is more detailed because it really goes from instead of just talking about ligands as kind of point particles or point atoms, it actually fully incorporates the orbitals of the ligands and how they form new hybrid molecular orbitals in the coordination compound, or in the organometallic compound in this case. So basically, ligand field theory is an application of molecular

orbital theory, where we move away from talking about the orbitals of individual atoms, and we actually start thinking in terms of molecular orbitals. This is something I've talked about before in some previous episodes on electronic structure, which is those that I did computational chemistry, because in a compound, you don't actually have like S or P or D orbitals. I've talked a long time about these D orbitals on the metal ion, but actually, in reality, those don't exist. D orbitals only truly exist in an isosal.

At a Noticulated atom. What you have ligands with that or other types of bonds like a carbon metal bond, you don't have D-Orbitals anymore, certainly not at the valence shell. What you have now are molecular orbitals, which represent some sort of hybrid of those initial atomic orbitals, combined together in a complicated way, depending on the exact geometrical arrangement. And so it's really the molecular orbitals that we should be thinking about, not the atomic orbitals, because the atomic orbitals are really just a fiction. to fiction. They're a way of thinking about it as simplification. Often the molecular orbitals will resemble the atomic orbitals in some way, but they're not the same as them. And so Ligod-fil theory takes that seriously and then extends the crystal-fill theory analysis for a proper molecular orbital

treatment. Now, the details are quite complicated, and I'll just try to explain some of the basic points here, but let me kind of give a hint or an idea of where we're going with this. And the basic idea is as follows. In the simple case of pie bonding, and if you recall, again, I've covered this before, pie bonding is basically when you have two circular orbitals, or close-ish-to-circular orbitals, overlapping each other kind of face-on. So this can be S-and-sorbitals, kind of like two spheres just overlapping each other, or it could be the head of a P-orbital, remember that's kind of like the hourglass shape, overlapping with an S-orbital. So that's also kind of like two spheres overlapping, although the P-orbital's not exactly a sphere,

but it's kind of similar ratio in shape. Both of those are cases of sigma bonds. They're the simplest type of bond that you can form between two different orbitals, and they basically represent spheres or near spheres overlapping, right? So what crystal field theory does, effectively, it only considers sigma-type bonds and the molecular orbitals that kind of form from that.

I mean, crystal field theory doesn't explicitly think in molecular orbital terms, but essentially that's what it does. It's like molecular orbital theory, but only with the simplest type of molecular orbitals, or bonds between the atoms. Only sigma-type bonding. It turns out that for some complexes, that's fine, because they only have sigma-type bonding,

and so you actually don't need really anything more than that. That's why Crystal Field Theory is so successful, because in a bunch of cases, actually, you don't need anything more. The simple sigma bonds is enough. But, and this is the key point here, for a lot of compounds, especially organometallic compounds, pie bonding is important. pie bonds between ligands, particularly carbon atoms, and the metal ion become very important.

And pie bonding is not considered at all by a crystal field theory. It doesn't fit into the framework. Now, you may be wondering what's a pie bond. Again, I have talked about this, so just go over quickly. A pie bond is like the next level up in complexity from a sigma bond. A sigma bond is kind of like spheres overlapping. A pie bond is basically, remember those kind of hourglass shapes from the P orbitals? those could meet head-on, right, with like lobe overlapping with lobe, that would be a sigma bond. But now imagine the meeting side-on, right? Kind of like putting two hourglasses next to each other. So the top lobe bit will touch side-on with the top-lob bit and the bottom-low bit will touch side-on with the bottom-low bit.

That's what a pi-bond looks like. It's kind of side-on p-orbital overlap instead of end-to-end p-orbital overlap. So S-orbital cannot form pi bonds. They don't have the right geometry, but p-orbital can, and critically D-orbital is also can form pie bonds. So in order to incorporate pie bonding, however, you need ligand field theory, because ligand field theory essentially takes the position of the ligands and the orbitals, critically the orbitals that are contributed that are donated by those ligands, and then overlaps them with the orbitals contributed by the metal, and then essentially constructs a hybrid orbital,

or more correctly, a molecular orbital out of that. Now, you might wonder, well, how do we know which orbitals of the ligands interact with which orbitals of the metal ion. The answer to that is actually that we use, we use symmetry. It turns out that there is a very complicated language for explaining the symmetry properties of the different orbitals. Just to give you an idea of that, a sphere has extremely high symmetry. I can rotate it any which way and it looks the same. So S orbitals have very high symmetry. P orbitals still have symmetry, but not as much, right? Because obviously if I take an hourglass shape and I rotate it kind of 90 degrees towards me, it looks completely different. It kind of just looks like a sphere now. But then I rotated another 90 degrees and it actually

looks the same again. Because if you'd like you tip the hourglass upside down, it looks sort of the same. Also, if you rotate it around its central axis, an hourglass looks the same. D-orbitals have different symmetry again because their shapes are more complicated. So the symmetry that an orbital has is actually critical for determining which other orbitals it interacts with. And the basic idea is that two orbitals will only interact with each other. By interact, I mean essentially form a new molecular orbital if they are of the same symmetry. Now, I will probably at some point devote a full episode to talking about symmetry in chemistry and group theory and how it's used to explain that. It's quite complicated, actually. So I'm not going to try to explain that now, and it doesn't really matter for our

purposes. All we need to understand is that different orbitals have different symmetry properties, and two different orbitals, say from the metal and from the ligand, have to have the same symmetry in order to interact and form a bond between each other or equivalently form a molecular orbital. A Molecular orbital is essentially the same thing as a bond. So what we do in ligand field theory, again, the update from crystal field theory, is that we take our ligands, position them appropriately. It depends on the coordination geometry of the complex. We take our metal and then we position the deorbital appropriately. And then we work out which orbitals of the ligands are going to interact with those deorbital.

And that depends both on the nature of the ligands and also their positioning, so their geometry. and using symmetry, we work out which orbitals are going to interact with which, and then from that we can determine what the molecular orbitals will look like. We can determine what symmetry properties they have, but also what energy level they will have. And there's complicated calculations that go into that, which we don't have to worry about here. For more on that, you could look at episodes 190 and 120 computational chemistry,

which kind of talk a bit about that. From here, all we need to understand is that we can use this theory to predict what the energy levels will be of the resulting complex, As I said, the results of this more complicated approach are pretty much the same as that you get from the simple crystal field theory approach as long as there's no pie bonding. But when pie bonding comes into mix, then things change. And so that's really what the payoff is here from Ligginfield theory. It's understanding the effects of pie bonding. You might be wondering, well, like, what does that actually matter in terms of practical effects, right?

In terms of like predicting chemical behavior. Well, there's one thing that hasn't been explained, right, that I talked about before. Remember I talked about the difference between strong and weak-field ligands. I said some ligands split the energy between the de-orbitales a lot and some split it only a little. Crystal field theory can explain why the splitting occurs, but it can't explain why some ligands split more and some split less. Crystal field theory actually just says that all it should matter is the geometry of the complex. It shouldn't actually matter what the ligands are. They should split it the same, right?

Because there's no notion of like the orbitals donated by the ligands that's just treated as a point. And so that's the limitation of crystal field theory that it can't explain why some ligands are strongfield and why some are weak field. And that's the beauty of ligand field theory that it can explain this. Liggen field theory is able to give an explanation for why some ligands, again, given a particularly geometry, here we're talking mostly octahedral, so why some ligands are strongfield and why some are weak field. It's essentially because of the role of pie bonding, which we can understand now through the lens of ligand field theory. So that leads us then to a discussion of how exactly this pie bonding effect from ligand field theory is able to explain

the difference between strong field and weak field ligands. And the basic idea is actually fairly simple. It's just a repetition or another example of the same phenomena that we've already seen in the case of crystal field theory. So what happens is that the orbitals of the ligand that engage in pie bonding, so these p-orbitales, interact with some of the de-orbitales of the metal ion. So remember, the interaction can only occur if they're both of the same symmetry class. So it's actually these lower-energy de-orbitals. Remember, we have five de-orbitales, but then they're split into the two upper-energy ones

and the three lower-energy ones because of the sigma bonding. Well, now those three lower-energy ones are actually split again as a result of the interaction with the P orbitals as a result of this pie bonding with the ligand. Only some ligands are able to do this, but those that can cause an additional splitting of these lower 3D orbitals that have the right symmetry structure. And so when this happens, as I said,

that there's a splitting, although in this case it's a little bit more complicated, right, because the ligands, the pi orbitals that interact with these lower energy D orbitals are also kind of added to the mix of total orbitals that we're dealing with. So we sort of start with three lower energy d-orbitales, and then we add in three new orbitals from the ligand. So we actually get six orbitals. So it's those six that are split.

The difference between the strong-field ligand and the weak-field ligand is really just a difference of how these p-orbital from the ligand interact with and split the lower-energy three d-orbital from the metal ion. So there's effectively two cases. In one case, the ligand has lower energy orbitals that it's, that it's sort of interacting with, the P orbitals are lower energy than the existing low energy deorbitals from the metal ion. So the new B orbitals coming from the ligands are lower energy than the existing low energy ones. And so what that's going to do is it's actually going to split the low energy deorbitals into two

and push the current low-energy orbitals higher. So it actually pushes the energy levels together. Remember, there's an initial splitting of the 5D orbitals into the two higher and three lower. So this extra layer of splitting actually causes the lower-energy three orbitals to be pushed closer in energy to the two upper-energy levels. So that actually causes weak-field ligands. So ligands that donate lower-energy-level P orbitals are actually responsible for weak-field effects. So they're weak-filled ligands, right? Because they actually push the three low-energy de-orbitales

higher or closer in energy to the two upper ones, closing the gap between them. Whereas the flip-side case, the other case is when the ligands actually have higher-energy P-orbital that then interact. And what they actually do is they push the low-energy de-orbital up above the energy level of the two previously higher ones. They actually push the three low-energy ones up so high

that they actually overtake the two previously high-energy ones, which are actually no longer the high-energy ones, they actually push it all the way above. The result of that, it may not be entirely obvious, but the effect of that is actually that the overall gap between the new high-energy orbitals and the new low-energy orbitals is actually bigger than before. So it actually splits them such that the gap is bigger, and this results in a strong-field ligand. So just to recap, in both of these cases, the weak-field or the strong-field ligands, you have ligands where there's an addition of new bonding orbitals from the P orbitals of the ligand. In one case, the weak field case, the electrons that are added by the ligand, are lower in energy

than the de-orbitales of the metal. That causes a splitting of the lower energy three orbitals from the metal ion, compressing the gap between the high and low energy of the de-orbitales, and thereby causing a weak-field effects, so reducing the energy gap. The other case is when the energy of the orbitals contributed by the ligands is actually higher than the energy of the D-orbitals. And that actually pushes the level of the formerly low energy 3D orbitals up above the formerly high-energy two orbitals, causing an increase in the gap between the low and the high energy levels. And that results in a strong field effect.

So high-energy orbitals from the ligand results in a big gap between energy levels in the complex, whereas low-energy orbitals from the ligand results in a small gap between the energy levels in the resulting complex. So don't worry if some of the details are a little bit obscure there. The main point there is that it's the energy levels of the p-orbitals contributed by the ligands that determine whether or not that ligand is a weak field or a strong-field ligand, and the way that they interact with the low-energy de-orbitales from the metal ion. So that's how you get the series of ligands with different effects on the field, which again, is pretty cool how much of an understanding we can get just based on the different shapes of

the metal ion orbitals and the ligand orbitals and how they interact with each other to form molecular orbitals. So we're nearly done for today. There's just a couple of last things that I wanted to cover. I wanted to say a little bit about the types of ligands that typically bind in organometallic complexes, because I've sort of been talking about them in fairly general terms. There are many hundreds of different ligands, and so I obviously won't even try to cover them all, but I will just say a little bit about some of the types of molecules that can be involved in this kind of binding.

So carbon monoxide is a fairly common ligand. Obviously, that's a fairly small one. Hydrogen or dihydrogen, so two hydrogen atoms bonded together is also a ligand. Many of the other ligands are various alkali compounds. So basically this means some number of carbons bonded together with single bonds. You can also have double bonds and triple bond carbons involved. And so the denticity, remember that's the number of bonds formed between the ligand and the metal, is affected by the number of carbons.

So if there's one carbon, then it's generally going to form one carbon bond with the metal ion. If there are two in the ligand, then it might form two carbon to metal bonds, or three or four and so forth. And it can go all the way up to six for an octahedral complex. So there is a vast array of different compounds depending on exactly how many carbons there are, whether there are single double or triple carbon bonds, depending on what other atoms are present like oxygen or nitrogen or phosphate or others. There's also ligands which bind to the metal ion in a ring structure.

So benzene, for example, can be a ligand. That's six carbon atoms with conjugate bonds between them. Don't worry if you don't know what that is. I won't explain that here. So they can actually bind directly to the metal and form a very strong hexidentate binding there with six bonding sites, one for each of the carbon. so that's a particularly interesting ligand.

So there's all sorts of these very complicated different arrangements possible, and you can hopefully begin to see how much diversity there is in organometallic chemistry, because it's like a combination of all of the range of possibilities with transition metals, in addition to a huge range of possibilities of bonding types and numbers of carbons and non-carbon atoms and so forth in organic chemistry combined together. So there is a vast array of chemistry possible here. All right. And the very last thing that I wanted to talk about is just to give you an idea of some of the types of reactions that are possible with organometallic compounds. Again, there's many, many reactions that are possible. So I won't even try to go through very many of them. I'll just mention a few of the different types to give you an idea of how the chemistry can work here.

Like in many cases when it comes to categorizing chemical reactions, there's a few sort of key types that come up again and again. So one of the most simple ones is a substitution reaction. This is when one electron donor, so essentially one orbital that's donated by ligand, is replaced with another. And there's many, many possibilities of these, when you can have one ligand that substitutes out for another, or one ligand that is like a tridentate ligand, that one of those donor pairs is removed, and then a different ligand binds instead. There's many possibilities here, particularly when you have high densities of carbon compounds that have three or four carbons, or as well as benzene rings, many possibilities there for slippages and movements where you can

have the carbon compound that's binding multiple times, but then it sort of moves over and opens a new binding site for a new ligand to come in. Lots of complex chemistry possible were there with these sort of substitutions. Then there are also additional reactions. So an addition reaction is when you, in this case it's going to be ligands that come in and form an additional bond that wasn't there previously. So addition reactions typically change the coordination geometry. If you had, for example, a square planar compound and then you have new ligands that come in from sort of the top and the bottom, and then they bind that could now change into an octahedral complex. This changes the oxidation state as well as the coordination number, obviously, because more ligands have come in.

The reverse of an addition reaction is an elimination reaction. This is when you reduce the coordination number and lower the oxidation state in this case, where you maybe go from a octahedral geometric arrangement to say a square planar arrangement by removing two of those ligands. There are yet more types of reactions as well. One type that's important for organometallic chemistry is called an alkali migratory insertion. This changes the coordination number by one instead of changing it by two in the as for the addition and elimination cases that we just talked about. A migratory insertion effectively happens where there's some kind of alkyl compound that's bound as a ligand to the metal and it kind of leaves right, it disassociates.

but then it stays attached to the metal indirectly via another ligand. So instead of binding directly to the metal, it moves off and binds to a different ligand of that metal. So instead of having two different groups leave or come at the same time, which was the case, say, in the elimination reaction, there's only one that leaves, but it does still stay bound to the overall complex. So you can see that just from this very brief overview, that there are a wide range of different types of reactions possible based on whether we're substituting out ligands, adding ligands, or removing ligands. and based on the change in the coordination number, the change in the number of ligands, and obviously what those ligands are. I haven't talked about the types of reactions of coordination complexes, but you can think of them as sort of broadly similar in that there are going to be substitution,

addition and elimination reactions, although the naming nomenclature is slightly different there. All right, well, there's plenty more to talk about with regard to transition metal chemistry, but I'm going to leave it there for now. That's more than enough already that we've talked about in this episode. So let's just do a quick recap. So we talked about what transition metals are and what makes them special. and particularly interesting, particularly that the valence electrons are in the de-orbitales, which means that they have quite different chemistry to, say, non-transition metals and also non-metals.

We talked about coordination complexes, which is when you have a central metal ion that forms these special bonds with the surrounding, typically inorganic ligands. We talked about coordination number, denticity, coordination geometries, chelation, isomerism, and then we introduced crystal field theory to explain the difference between strong and weak-field ligands, and the effect that the splitting of the energy levels has on the color of these complexes, as well as the magnetic properties and other chemical properties of these complexes. I then introduced organometalic complexes or organomatelic compounds, which is similar to coordination compounds,

except in this case that you have organic ligands or carbon to metal binding directly, and they tend to have different chemical properties to coordination complexes as well, so we study them separately. We talked about ligand field theory and introduced that as a extension of crystal field theory and use that to explain why some ligands are strong-filled and why some are weak-field based on the pie bonding interactions as opposed to crystal field theory, which just dealt with sigma bond pairing interactions. We also talked about some of the different ligands that occur in organometallic chemistry and a brief overview of some of the different types

of reactions that can occur. So hopefully you found this interesting. If so, please consider giving the podcast a favorable review on the aggregator of your choice. If you'd like to get in touch with me, you can send me an email on Fods12 at gmail.com. That's FODDS12 at gmail.com. I've been getting quite a lot of feedback recently, which is really great. I'd love to hear from all your listeners. Also, if you'd like to support the podcast financially, you can become a Patreon supporter on Patreon. Thanks also to all of those who make one-off donations via PayPal, which you can also do if you would like. I appreciate all of you. And hopefully, I have been able to thank all of you individually when you make the donation. If not, please can you touch with me and I'll rectify that.

So thanks once again for listening and I'll talk to you next time.