Coordination Complexes: Videos & Practice Problems

Coordination complexes are formed by transition metals bonded to ligands, with a counter ion ensuring overall neutrality. The coordination number indicates the number of ligands attached, commonly 2, 4, or 6, leading to specific geometries: linear for 2, tetrahedral for d10 configurations, and octahedral for 6. For d8 configurations, square planar geometry is observed. Understanding these structures and their electron configurations is crucial for predicting complex ion behavior in chemical reactions.

Coordination Complexes

In this video, we're going to take a look at the concept of coordination complexes. Now, we're going to say that the most prevalent feature of transition metal chemistry is the formation of coordination complexes or compounds. These structures are composed of a complex ion that is bonded or connected to ions or molecules called ligands. To maintain the overall neutrality of the compound, a counter ion is used. Alright. If we take a look here at this example, we have nickel bonded to 4 ammonia molecules and then we have 2 chloride ions outside of the brackets. This thing overall together is our coordination complex. Our coordination complex could be thought of as an anion. So taking this logic, let's highlight this part in yellow. Since it's in a bracket, it's all together. And outside the bracket, we have these two chlorines. So, what's in the bracket represents our complex ion, and what's outside the bracket is our counter ion. Outside the bracket, we have two chlorines, which are two chloride ions. They are our counter ions. Now, where did this 2 come from? That 2 came from my complex ion portion. So it's still together as Ni4NH3+2. This is my complex ion. Now, we can go a little bit further and break the complex ion further into its components. We're going to say that this complex ion is composed of 4 ammonia molecules. We use the term molecules because ammonia is neutral, and they represent our ligands. So remember, ligands could either be neutral or they could have a charge. Examples of a neutral ligand are ammonia, but you could also have water as another example. Ligands can also be negative. So examples of negative ligands could be the cyanide ion, the azide ion, or even another halogen as a ligand. So those would be negative. And then if we think about it, the ligands here are 4 ammonia molecules. They're neutral. They don't provide any charge, but yet our complex ion overall has a plus 2 charge. If it has an overall charge of plus 2 and the ligands aren't contributing anything, then that plus 2 charge has to be coming from the nickel. Nickel itself has to be plus 2. So this would be our transition metal. Again, our coordination complex is just a more complex ionic compound that can be broken down into its counter ion and then its complex ion, counterions. Then you can further break down the complex ion into its ligand portions and its transition metal portion. Connected to complex ion or complex ions, coordination complexes, all that stuff, we have what are called coordination numbers. Now, we're going to say the coordination number is the number of ligands bonded to the central metal cation. So your coordination number is how many of these ligands are connected to my transition metal. In this case, we'd say that there were 4 ligands connected to nickel 2+ ion. Now, we're going to say the most common coordination numbers are 2, 4, 6 ligands attached to my transition metal cation. There are other types, there are other numbers as well, but these are just the 3 most common ones that we see in organic chemistry. Alright. So now that we've gone over what a coordination complex is, how it's composed of a complex ion, counterions, ligands, and transition metal. We'll take a look at the examples below. So I want you to click on to the next video and see how I approach the first example question dealing with our coordination complex given.

Coordination Complexes Exercise 1

So here it says, correctly label all the components of the coordination complex. Now, remember we say that traditionally when we have brackets, what's inside the brackets represents our complex ion, and what's outside our brackets represents our counterions. In this example, we have 2 sodiums outside of the brackets, so they represent our counterions. Na is in group 1, so its charge is plus 1. So we'd say we have 2 sodium ions that represent our counterions. Then we have tin with 6 chlorines. Alright. So where did this 2 come from? Well, that 2 must have come from the complex ion. Sodium is positive. In this case, our complex ion now is negative. Negative. Remember, our complex ion we said could be positive or negative. So here it would be SnCl6 still in brackets, and the overall charge would be -2. So this represents our complex ion.

Now, we're going to say what is that complex ion further broken down into? We're going to say here that it is composed So these are our ligands. Remember, ligands are what are directly connected to my transition metal. And if we think about it, we have 6 chloride ions, so their overall contribution is a minus 6 in terms of charge. Right? So 6 times negative one, that's negative 6. And the overall charge here is minus 2, though. How is that possible? Well, that would mean that tin has to be plus 4 in terms of charge. So, plus 4 coming from the tin, and the negative 6 coming from the 6 chlorines, overall that gives us a negative 2 charge left over. So we're dealing with Sn4+ion. So this represents our transition metal, cation. So those would be all the components that we have here for our coordination complex. What's in the brackets can be seen as our complex ion. It can be either positive or negative. Then you can go further into the complex ion and see what's the transition metal ion used and what are the types of ligands used.

Now that we've seen this example, click on to the next video and see if you can figure out the number of ligands within the provided complex ion. Come back and see, does your answer match up with mine?

Coordination Complexes Exercise 2

For this one, it says determine the number of ligands in the complex ion. Now, remember, the complex ion portion is the part that is within the brackets. That fluoride is a counter ion because it's outside the brackets. So fluoride is in group 7, so its charge is minus 1. So to have an overall neutral coordination complex, that means it has to be plus 1. So that plus 1 and that negative one cancel out, and that's why the coordination complex overall is neutral.

Now, we're going to say within this complex ion portion, we see that chromium is our transition metal and it's connected to what appears to be 4 water molecules and 2 bromide ions. So that is a total of 6 ligands that are attached to my chromium ion. So that would be our answer.

Now, I didn't ask for the charge of the chromium, but let's take a look at it and see if we can figure it out. So here, the bromines each have a charge of minus 1. Together, that would mean that they're minus 2. But our complex ion's charge overall is plus 1. This could only happen if the chromium itself was plus 3. Because we have plus 3 minus 2 gives me plus 1 overall. Remember, water molecules are neutral, so they don't contribute to the charge at all.

Alright. So just remember, when it comes to a coordination complex, you have to be able to spot the complex ion portion and the counter ion portion. You may also be asked to look at the complex ion and break it down a little bit further into its ligand portions and its transition metal portion. And when it comes to our ligands, it's normal to see 2, 4, or 6 of them connected to our transition metal. Now that we've seen this, click on to the next video where we dive a little bit deeper into the different types of molecular geometries that are possible based on the number of ligands attached to my transition metal.

Coordination Complexes

Now we can say that coordination complexes form predictable geometries based on their coordination and sometimes their electron configuration. So remember, the most common types of coordination numbers are 2, 4, and 6, where we have 2, 4, or 6 ligands attaching to our transition metal. Now, when it comes to our coordination of 2, we have 2 ligands attached to our transition metal. In this case, we don't have to worry about the electron configuration of our transition metal, so we'll ignore this portion. If our transition metal is connected to just 2 ligands, then its geometry would just be linear. A good example of this is we could have copper connected to 2 bromines and we can say the overall charge is minus 1. If we were to draw this out, it would just be copper connected to the 2 bromines and brackets with a charge on the outside. When it comes to a coordination of 4, 2 different geometries are possible but we're gonna come back to that. Let's go on to a coordination of 6. So if we have a coordination of 6, because there's only one type of geometry associated with it, we don't have to worry about the electron configuration of the transition metal. We'll skip that portion. If your transition metal is connected to 6 ligands, then we're gonna say that its geometry would be octahedral. Now, a good example of this, we could use, cobalt connected to 6 ammonia. And let's say that the overall complex ion is plus 3 in charge. Now, what would that look like? Well, if we're going to take into account three-dimensional drawings, we'd have our cobalt in the center. It would be connected to our 6 ammonias with 1 above and below the plane. Then we'd have 2 ammonias that are pointing into the paper, and then we'd have 2 ammonias pointing out of the paper. And then we put it in brackets with the overall charge on the outside. Geometries, we're geometries. And in because there's 2 possible geometries, we can use the electron configuration of the transition metal to determine which one is favored. So, with a coordination of 4, we're gonna say here that a transition metal with a d ten electron configuration complexes. So let's say something with a d ten electron configuration. A good example is the element of zinc. Zinc, its electron configuration is Argon 4s2 3d10. So here, its geometry would be tetrahedral. So, if we did an example we could have Zn and let's say we have OH4 and two. Alright, So if we were to draw this out, we have our Zn in the center. We're drawing tetrahedrals. So for tetrahedral, we'd have 1 OH on the plane of the paper, another OH on the plane of the paper, 1 OH pointing down into the paper, and then one pointing out from the plane of the paper. Brackets and the overall charge on the outside. Now, this is true whether we're dealing with neutral zinc or zinc2 ion. Because remember, if we had zinc2 ions, we'd lose electrons from the outer shell first, which would mean we lose those 2 4s electrons. So whether it's in its neutral form or its plus two form, zinc is still classified as having a d 10 electron configuration. So in both cases, it would give us a tetrahedral geometry. Alright. So now, for d 8 geometry, we could use nickel. Right? So if we have nickel, we'd say nickel is argon 4s2 3d8. And if we're dealing with nickel 2+ ion, it'd be argon, 4s2 those 2 electrons are lost, and it will be 3d8. So whether you're dealing with the neutral form or its plus 2 form, it's still a d 8 electron configuration. It would give us a square planar or planar geometry. So example here, we could use nickel. We could use CN4, 2 minus. So if we're going to draw a square planner we'd have nickel in the center. 2 of our CNs would point into the plane of the paper. And 2 of them would point out of the plane of the paper. Brackets, overall charge on the outside. So again, it's common for us to see complex ions where we have a coordination number of 2, 4 or 6. A coordination of 2 or 6 is pretty easy because they only have one geometry connected to them. But when it comes to a coordination of 4, then it's important to be able to determine the electron configuration of the transition metal. If we have a d ten electron configuration for the transition metal, then we're gonna form a tetrahedral, complex. And if we have a d 8, then we're going to form a square planar or planar, complex. So keep these structures in mind, keep in mind the electron configuration when it comes to coordination of 4 and you'll be able to draw any of these complex ions.

Coordination Complexes Exercise 3

So, here in this example it says determine the geometry for the following complex ion. Here, we have chromium as our central element. It's connected to what appears to be four ammonias and two chlorines. That's a total of six ligands. Pretty simple because, if our coordination number is six, then there's only one geometry that's possible, octahedral. Alright. So, if we're going for symmetry here, we have our two chlorines drawn on opposite ends so they're drawn anti to each other. And then we'd have our four ammonias. Now remember here, we'd have two ammonias pointing into the plane of the paper. And then we have the other two pointing out of the plane of the paper. And because it has a charge, we put brackets with the charge on the outside. So, this would be our structure for this given coordination complex. Now that you've seen this example, move down below and take a look at the next one.

So in this one, they're asking us to determine the geometry where we have palladium connected to four waters. Here, because there are four waters involved, that means our coordination number is four. So remember, what do you do when your coordination number is four to determine the correct geometry? Attempt this one on your own and come back and see does your answer match up with mine. Good luck, guys.

Determine the geometry for the following complex molecule:Pd(H_­2O)₄.

Determine the geometry for the following complex molecule. Here we have palladium connected to 4 waters. Now remember, we start off by determining the electron configuration of the transition metal. Palladium, if you were to look on the periodic table, it would appear as though its configuration is Krypton 5s24d8. But remember that palladium is a bit weird. The 5s electrons actually combine with the 4d electrons to give us krypton 4d10. Now because it's a d10 element, that means it's going to have a tetrahedral geometry. If there were a d8 element, then it would have a square planar geometry. Now because it's tetrahedral, we draw it as such. We'd have palladium here, we'd have on the plane of the paper are 2 water molecules, and then we'd have one that is out of the plane of the paper, and then we'd have one that's into the plane of the paper. So these would represent our tetrahedral complex molecule where we have palladium with 4.

Here’s what students ask on this topic:

What is a coordination complex in chemistry?

A coordination complex, also known as a coordination compound, is a structure consisting of a central transition metal atom or ion bonded to surrounding molecules or ions called ligands. These ligands can be neutral or negatively charged. The overall charge of the complex is balanced by counter ions. For example, in the complex [Ni(NH₃)₄]Cl₂, nickel is the central metal, ammonia molecules are the ligands, and chloride ions are the counter ions. The coordination number, which is the number of ligands attached to the central metal, determines the geometry of the complex.

What are the common geometries of coordination complexes?

The geometry of a coordination complex depends on its coordination number and the electron configuration of the central metal. Common geometries include:

• Linear: Coordination number 2 (e.g., [CuBr₂]^-)
• Tetrahedral: Coordination number 4 with d¹⁰ configuration (e.g., [Zn(OH)₄]^2-)
• Square Planar: Coordination number 4 with d⁸ configuration (e.g., [Ni(CN)₄]^2-)
• Octahedral: Coordination number 6 (e.g., [Co(NH₃)₆]³⁺)

These geometries are crucial for predicting the behavior and reactivity of coordination complexes in chemical reactions.

How do you determine the coordination number of a complex ion?

The coordination number of a complex ion is determined by counting the number of ligand atoms directly bonded to the central metal ion. For example, in the complex [Ni(NH₃)₄]Cl₂, the coordination number is 4 because there are four ammonia (NH₃) molecules bonded to the nickel ion. Common coordination numbers are 2, 4, and 6, which correspond to linear, tetrahedral or square planar, and octahedral geometries, respectively.

What is the role of counter ions in coordination complexes?

Counter ions in coordination complexes serve to balance the overall charge of the complex. They are not directly bonded to the central metal ion but are essential for maintaining electrical neutrality. For example, in the complex [Ni(NH₃)₄]Cl₂, the two chloride ions (Cl^-) are the counter ions that balance the +2 charge of the [Ni(NH₃)₄]²⁺ complex ion. Without these counter ions, the complex would not be electrically neutral.

How does the electron configuration of a transition metal affect the geometry of its coordination complex?

The electron configuration of a transition metal significantly influences the geometry of its coordination complex. For instance, a metal with a d¹⁰ electron configuration, like zinc in [Zn(OH)₄]^2-, typically forms a tetrahedral geometry. Conversely, a metal with a d⁸ configuration, like nickel in [Ni(CN)₄]^2-, usually forms a square planar geometry. Understanding the electron configuration helps predict the spatial arrangement of ligands around the central metal ion.