Ligands - Online Tutor, Practice Problems & Exam Prep

Ligands act as Lewis bases, donating lone pairs to metal cations, forming complex ions. They can be neutral (L ligands) or negatively charged (X ligands). Monodentate ligands donate one lone pair, while bidentate ligands donate two, often creating cyclic structures. Common examples include water, ammonia, and ethylenediamine. Understanding the classification of ligands is crucial for grasping complex ion formation and their geometries in coordination chemistry.

Ligands

In this video, we're going to take a look at ligands. Now, a ligand represents a Lewis base because it bonds to a metal cation in a complex ion by using its lone pair. Remember a Lewis base is an electron pair donor. The transition metal catalyst accepts that lone pair, so it represents our Lewis acid. Now, since ligands use their lone pair or pairs to grab onto metal cations, they are referred to as chelating agents. We've seen that ligands can either be neutral in the form of like ammonia or they can be negative like in the form of cyanide ions or a halogen. Now, we're going to say ligands are typically either neutral and termed L ligands or negatively charged and termed X ligands. Now the number of elements in a molecule that can donate a lone pair characterizes a ligand. We're going to say ligands that possess the ability to donate 1 lone pair at a time are referred to as monodentate ligands. And then ligands that possess the ability to donate 2 lone pairs at a time are referred to as bidentate ligands. So we're going to be looking at monodentate and bidentate ligands.

Now bidentate ligands give rise to rings in the complex ion form. So basically when you have a bidentate ligand connecting to a transition metal cation, it helps to make a cyclic complex ion. So, in the next video, we'll take a look at the chart provided below where we look at what exactly makes a ligand monodentate, what makes it bidentate, and then how is it an X or L ligand? So click on to the next video and let's explore the chart together.

Ligands

So we've separated our chart here in terms of monodentate ligands and bidentate ligands. Let's first take a look at our monodentate ligands. Now, what makes them a monodentate ligand is that in each case, there's only one element that's able to donate a lone pair. So we have our ligands. We have names for those ligands. Now the naming of ligands and the naming of complex ions is really just reserved for general chemistry. Within organic chemistry, we're not going to have to worry about naming complex transition metals, how they help to make certain geometries, and later on the types of reactions that are possible with these complex ion structures. So don't worry about naming. And then we'll see that certain ligands have abbreviations. And once we've covered those three topics, we'll take a look and see is the ligand I'm looking at an x ligand or an l ligand or maybe even a combination of both. So if we take a look, what makes the monodentate ligands is that in each case, there's only one element that will donate a lone pair.

In water, we have just the oxygen being able to donate a lone pair. Here we have ammonia and its name would be amine or amino. Again, don't worry about naming too much. We have just the nitrogen. Here we have phosphorus with its one lone pair, \(PR_3\). Now here, that \(R\) could represent a benzene ring. And in that case, it would be a triaryl phosphino and it would be written like \(PPh_3\), where the \(PH\) stands for our benzene ring. It could also stand in for an alkyl group like maybe a methyl. So you could have a phosphorus connected to three methyl groups. Next, we have carbon monoxide which is carbonyl. Here it's only the carbon that will donate its lone pair and act as the Lewis base. Acetonitrile, here it's only the nitrogen. For ethylene, what's being donated is a lone pair and it's coming from our pi bond here. Benzene has three pi bonds so any one of the three could donate a lone pair. \(X\) is our halogen so we have fluoride, chloride, bromide, or iodide. \(H^-\) is our hydride ion. Then we have a cyanide molecule or cyanide ion. So here it's just the carbon that can donate a lone pair. Hydroxide is just the oxygen. Here we have cyclopentadienyl. So technically in this one, we can get a lone pair from this lone pair here or from one of these pi bonds. But it's only one at a time. That's what makes it monodentate. We're only donating one lone pair at a time. And then here we have an allyl group. So here it is this carbon with its lone pair that is donating to the transition metal cation.

Now, some of these have abbreviations. So our carbon monoxide, termed carbonyl, its abbreviation is just \(CO\). Here we have Acetonitrile. Its name would be Acetonitrile. It's written as \(MeCN\), where \(Me\) represents methyl. Benzene, its abbreviation is \(Ph\). Some of these abbreviations you should be quite familiar with. Halogens are just \(X\). Cyanide ion is \(CN\). Cyclopentadienyl is \(Cp\). The others don't have abbreviations.

Now what type of ligands are they? Well, remember, \(L\) ligands are neutral. \(X\) ligands are negatively charged. So here, water has no charge, so it's an \(L\) ligand. Ammonia is an \(L\) ligand. \(PR_3\) is an \(L\) ligand. Here, \(CO\) is also an \(L\) ligand. \(L\) ligand, \(L\) ligand. Now when it comes to benzene, we have three pi bonds. Any one of the three pi bonds can donate a lone pair, one at a time. But the fact that we have three pi bonds that any of them could donate, we're gonna say that this is an \(L3\) ligand. Okay? That means it has three pi bonds that can donate a lone pair. Here we have a halogen. We finally see a negative charge so these would be \(X\) ligands. They all have negative charges. Now, for our cyclopentadienyl, it has a negative charge because of the lone pair on this carbon, so that's \(X\) plus the lone pairs here. Either one could be donated. So technically, this one would be an \(X\) \(L2\) ligand. This one, we could have this lone pair being donated to make a pi bond, or we could have this pi bond here being donated. So this would be an \(XL\) ligand. Remember, an \(L\) ligand is neutral and \(X\) ligand is negative, and these last two are a combination of both.

Now, here we have our bidentate ligands. What makes them bidentate is that we have two places, two locations where lone pairs can be donated. In the first one, we have ethylenediamine, and it's both nitrogen with their lone pairs that can be donated. And then we have here, this is acetylacetonato as its name. And it's both oxygens here that can be donated. Just like both oxygens here that can be donated. They each have their own abbreviations as well. So when it comes to ethylenediamine, its abbreviation would be \(en\). Now we're gonna say here that these two lone pairs, there's no negative charge involved. So we'd say that this would represent an \(L2\) ligand. Here, oxalate, it would be \(ox^{-2}\) as its abbreviation. You have two negative oxygens that can donate a lone pair at the same time. So that's an \(X2\) ligand. And then here for acetylacetonato, we have two oxygens as the site they both donate a lone pair at the same time. So here, its abbreviation is \(acac\). Both are neutral and they can donate lone pairs, so they will be \(L2\) ligands. These are the most common types of ligands, that are possible when it comes to complex ion formation. Remember, when it comes to neutral ligands, they are called \(L\) ligands. Negatively charged ligands are \(X\) ligands. If we have a negative charge, then that can represent an \(X\) ligand. If we look at pi bonds being used to donate a lone pair, that's what is characteristic of an \(L\) ligand. So that's what you look out for. Just keep in mind some of these common types of ligands as we explore more and more about complex ions. Now that we've talked about the differences between them, we'll take a look in the next video on the following examples where we're just being asked to give the complex ion structure. So click on to the next video and let's tackle the first example together.

Ligands Example 1

Alright. So, let's take a look at this example. It says, Provide the complex ion structure when a titanium 3 ion combines with 4 ammonia molecules and 2 chloroanions. Alright, so we have a titanium 3 ion. It is combining with 4 ammonia molecules. So remember, ammonia is NH3 and there are 4 of them, plus 2 chloroanions. Chlorine is in group 7 so its charge is minus 1. To get the correct complex ion structure, we need to see what the overall charge will be. So, titanium contributes a plus 3 charge. Our 4 ammonias are neutral, so they contribute 0. Then, we have our 2 chloride anions. So each one is minus 1 and there are 2 of them, so that’s 2 minus 1 which gives us negative 2. Add everything together so we have plus 3 minus 2, which gives us plus 1 overall for the charge of our complex ion. Because our complex ion has a charge, we're going to use brackets. Here goes our titanium. It is connected to 4 ammonia molecules. So, we put those in parenthesis. And then our 2 chloroanions, close parenthesis and our overall charge is plus 1, which we calculated. So that represents the structure of our complex ion once we have our titanium 3 ion combining with the 6 ligands involved. Now that you've seen this example, take a look at the next example and see if you can actually do it. Once you do that, come back and see if your answer matches up with mine.

Ligands Example 2

So, for example 2, we have to provide the complex ion structure when copper(I) combines with ethylene diamine and 2 cyanide anions. So we have copper(I), and then we have ethylene diamine. Here we can use the abbreviated form of ethylene diamine, so we just say that's en. And then we have 2 cyanide anions. So 2 CN^-. The overall charge of our complex ion would be +1 from the copper(I). Ethylene diamine is an L-two ligand so it doesn't contribute any charge. And then we have 2 cyanide ions so that they're contributing -2 overall. So when we add all that up, our complex ion should have a negative one charge. Now we write it out. So we have copper. Remember brackets because we're going to have a charge. Copper is connected to our ethylene diamine. So we just do parenthesis en. Then it's connected to 2 cyanide ions. So parenthesis 2, close brackets, and then the -1 on the outside. So this would represent our complex ion structure:

[ Cu+ ( en ) ( 2 CN- ) ] -1

Remember, it's important to know the types of ligands that are attaching to your transition metal because that'll have an effect on the overall charge of your complex ion. Remember that ligands that have no charge that are neutral are typed L ligands whereas ligands with a negative charge are typed as X ligands. These are important terms that you need to keep in mind when we delve deeper and deeper into complex ions and the eventual organic chemical reactions that they will undertake.

Provide the complex ion structure and geometry when a Ni atom combines with 2 chloro anions and 2 water molecules.

So here we need to provide the complex ion structure and geometry when a nickel atom combines with 2 chloroanions and 2 water molecules. Alright. So let's do the complex ion structure first. A nickel atom means that it's in its neutral form because, remember, if it had a charge, we wouldn't use the term atom; we'd say ion. Also, if it had a charge, they would show it to us. Remember, transition metals have more than one charge. So, if they're not giving you a charge, we're assuming that it's neutral. Sometimes you may see them write a 0 here to illustrate that it has no charge. We have 2 chloroanions, so that's 2 Cl-. And then we have 2 water molecules. Water is neutral. So, our overall charge, we have 0 here, minus 2 from the 2 chlorines, plus 0 from the waters. The complex ion should have a minus 2 charge overall. The structure would be nickel. We write it in the order that's given to us. So, we have nickel, we have 2 chlorides, and then we have 2 waters. We put that in parentheses and then the 2 on the outside. Then the overall charge is minus 2, so this would be Ni(Cl2(H2O)2)2- for our complex ion structure. Now, we need to figure out its geometry. Going back to coordination numbers, we have 4 ligands attached to our nickel atom: 2 chloroanions and 2 waters. So remember, if your coordination number is 4, it kind of sucks because that means you could be either tetrahedral or square planar. Now, do you remember the type of electron configurations that determine which one we have? Well, remember, if you have a tetrahedral geometry, that's because you have a d10 electron configuration. And if you're square planar, that means you have a d8 electron configuration. So we have to look at the transition metal and its electron configuration. Nickel here, if you look at your periodic table, remember nickel would be [Ar] 4s23d8. And it's in its neutral form so we don't have to worry about removing any electrons. Here, it is a d8 electron configuration. Therefore, we should draw it as square planar. So we're going to have our nickel in the center. So if we have nickel in the center, it's connected to 2 chloros and 2 waters. Now, there are 2 different ways you could draw this. So remember, with square planar, we have these 2 as being into the page of the paper and then these two bonds that point out of the plane of the paper. There are 2 different ways I could draw this square planar. So let's say I decided to draw the 2 chlorines being on the same side with each other. So they’re both going into the paper. And then I drew the 2 water molecules as also being on the same side. Here we put brackets and the outside charge. So, remember, what do we say if things are on the same side? Well, if you're on the same side, then you would be cis. So, this would be a cis configuration for this square planar structure. But let's say they're not on the same side. Let's say they're on opposite sides. So I put chlorine on opposite sides where one is going into the plane of the paper and one is pointing out of the plane of the paper. And the same thing with the water molecules. We put brackets and the charge on the outside. So, what do we say if you're on opposite sides? Well, if you're on the same side, you're cis. If you’re on opposite sides, you are trans. So here, both are potential answers for the way to draw the geometry of this particular complex ion. So remember, when it has a coordination number of 4, it gets tricky because it could be either tetrahedral or square planar. You have to look at the electron configuration of the transition metal to determine which one is favored. A d10 electron configuration gives us tetrahedral whereas a d8 electron configuration gives us square planar. Alright. Now, I want you guys to attempt to do this practice question left at the bottom of the page. I'll give you guys a hint. We're going to say here that we have a cobalt(III) ion connecting to 3 ethylenediamines. Ethylenediamine is a bidentate ligand. It has two sites where lone pairs can attach to our transition metals. Bidentate ligands, remember, they make cyclic, complex ions. Also, because they have 2 sites of lone pair donation, it's kind of like counting it as 2 ligands. So, what I'm basically saying is, 3 ethylenediamine is equivalent to 6 ligands. Because each ethylenediamine is donating 2 lone pairs. So, 2 ligands per ethylenediamine. I'm telling you this hint because what kind of geometry do we have when we have 6 ligands involved? Okay. So that's all I'm going to tell you. Attempt to see if you can draw this out. If you get stuck, don't worry. Click on to the next video and see how I approach this practice question. Good luck, guys.

Provide the complex ion structure and geometry when a Co³⁺ ion combines with 3 ethylenediamines.

Alright. So from the little bit of hints that I gave to you, let's approach this question. So we know that we have a cobalt(III) ion and it's connecting to 3 ethylenediamines. Remember, in terms of abbreviation that's just en. Now, we're going to say that the structure would be cobalt connected to 3 ethylenediamines. The overall charge would be plus 3 coming from the cobalt(III) and no charge contribution from the ethylenediamines because they are neutral ligands. So overall, the complex ion should have a plus 3 charge overall. So brackets and then 3+ on the outside. Now, we have to draw this. Remember, ethylenediamine is a bidentate ligand and they form rings when they combine with the transition metal to make the complex ion structure. We have our cobalt in the center. As these 3 ethylenediamines are equivalent to 6 ligands, 6 monodentate ligands, this gives us an octahedral shape. In an octahedral, we'd have two bonds on the plane of the paper, two bonds that go into the plane of the paper, and two bonds that point up out of the plane of the paper. All we have to do now is draw these 3 ethylenediamines. Remember, ethylenediamine has the structure of NH2 connected to a CH2. So that's one part of it done. Same thing here, NH2 connected to a CH2, CH2 connected to an NH2. Then NH2 connected to a CH2. Connected to another CH2. Connected to an NH2. Okay, I'm drawing it as best as I can. So we'd get something like this. As we can see, each bidentate ligand helps to make a cyclic structure for a complex ion. And then brackets and then the charge on the outside. So you get something like this as the geometry for a complex ion. So, just keep in mind when it comes to complex ion structures, the charge of your ligands are important and the charge of your transition metal is important. They all contributing together gives us the overall charge of our complex ion. Also, remember the rules that we established when it comes to the geometry of these complex ion structures. If you have a coordination number of 2, which equals 2 ligands, then you're linear. If you have 6 ligands, then you're octahedral. When you have a coordination number of 4, then you can be either tetrahedral or square planar. It's all based on the electron configuration of your transition metal. Keep these in mind when you're writing either the complex ion structure or its geometry.

Here’s what students ask on this topic:

What is a ligand in coordination chemistry?

A ligand in coordination chemistry is a molecule or ion that donates a pair of electrons to a central metal atom or ion to form a coordination complex. Ligands act as Lewis bases, meaning they are electron pair donors. They can be neutral molecules like water (H₂O) or ammonia (NH₃), or negatively charged ions like chloride (Cl^-) or cyanide (CN^-). The nature of the ligand, including its charge and the number of lone pairs it can donate, influences the geometry and reactivity of the resulting complex ion.

What is the difference between monodentate and bidentate ligands?

Monodentate ligands are ligands that donate only one pair of electrons to a metal center, forming a single bond. Examples include water (H₂O) and ammonia (NH₃). Bidentate ligands, on the other hand, can donate two pairs of electrons from two different atoms within the same molecule, forming two bonds with the metal center. This often results in the formation of a ring structure within the complex. An example of a bidentate ligand is ethylenediamine (en), which has two nitrogen atoms each donating a lone pair.

What are L ligands and X ligands?

L ligands are neutral ligands that do not carry a charge. Examples include water (H₂O) and ammonia (NH₃). X ligands are negatively charged ligands, such as chloride (Cl^-) and cyanide (CN^-). The classification into L and X ligands helps in understanding the electronic properties and the type of bonding interactions they have with the metal center in coordination complexes.

How do ligands affect the geometry of coordination complexes?

Ligands influence the geometry of coordination complexes by the number of lone pairs they donate and their spatial arrangement around the metal center. Monodentate ligands typically lead to simpler geometries like linear, tetrahedral, or octahedral shapes. Bidentate and polydentate ligands can create more complex geometries, often forming ring structures within the complex. The steric and electronic properties of the ligands also play a crucial role in determining the final geometry of the coordination complex.

What are some common examples of ligands in coordination chemistry?

Common examples of ligands in coordination chemistry include water (H₂O), ammonia (NH₃), chloride (Cl^-), cyanide (CN^-), ethylenediamine (en), and acetylacetonate (acac). These ligands vary in their charge, the number of lone pairs they can donate, and their ability to form different types of bonds with metal centers, influencing the properties and reactivity of the resulting coordination complexes.