The Electron Configuration: Exceptions (Simplified): Study with Video Lessons, Practice Problems & Examples

Electron orbital stability is enhanced when d subshells are half-filled or fully filled, following Hund's rule. Notable exceptions in electron configurations occur in elements like chromium (Cr) and copper (Cu), where electrons are promoted from the 4s to the 3d orbitals to achieve stability. For chromium, the configuration shifts from Ar4s²d⁴ to Ar4s¹d⁵. Similarly, copper transitions from Ar4s²d⁹ to Ar4s¹d¹⁰.

The Electron Configuration: Exceptions (Simplified) Concept 1

So when it comes to electron orbital stability, you just need to remember that d subshell orbitals are most stable when they are half-filled or totally filled with electrons because of symmetry. So what we mean by half-filled, remember, we're going to follow Hund's rule, which says that n orbitals that have the same energy that are degenerate are first half-filled. So we're dealing with up, up, up, up, up. Here we have a set of d orbitals electrons that are all facing up. So this set of orbitals is half-filled. Now, totally filled, we go up, up, up, up, up, then come back around, down, down, down, down, down. So here we have an example of a half-filled set of d orbitals, and here a set of totally filled d orbitals.

The Electron Configuration: Exceptions (Simplified) Concept 2

Now when looking at exceptions to the electron configurations, we're going to say starting from chromium, which is Cr, as the atomic number Z increases, exceptions to electron configurations can be observed. A memory tool we can have is that chromium has an atomic number of 24. So, think about that, 24. We're going to say that the exceptions happen with these 2 elements and with these 4 elements, so "24". We're going to start out with chromium; we know that's where it starts, and we're going to skip the next 4 columns. Right? So we start out with chromium and you skip the next 4. So skip Manganese, skip Iron, skip Cobalt, skip Nickel, and then you land on copper where the next group of exceptions can exist. So just remember, these are the 6 major types of elements where we're going to see exceptions to the electron configuration. Keep this in mind when we're looking at their electron configurations. Now that we know that these are the 6 to deal with, let's see how these exceptions arise. Click on the next video and let's see what happened.

The Electron Configuration: Exceptions (Simplified) Concept 3

So remember, exceptions start with chromium. Let's look at chromium. If we were to determine its electron configuration initially, we would see that it would look like argon 4s²3d⁴. Now here, what do we have here? We have 3d with 4 electrons within it. But remember, earlier we said that s and d subshells or sublevels have this urge to try to be half filled or totally filled. Now we're going to say, an s orbital electron can be promoted to create half-filled orbitals with d 4 electrons. So what we're saying here is if you're doing the electron configuration of chromium, you're going to end with a d 4. That's a key to tell you that, oh, d 4. We have only 4 electrons within these d orbitals. But if I could somehow get one more electron in there, those d orbitals will be half-filled. So what's going to happen here is we're going to take a d 5, and our 4s² just gave up an electron so it becomes 4s¹. So it now looks like this. This would be the correct electron configuration of chromium. So, again remember, chromium has this type of exception, and the driving force is trying to get a half-filled set of d orbitals. Okay. So here, we're not going to land, we're not going to stay as d 4, when it's neutral it's going to become d 5. Now that we've seen this with the first column, let's see what happens with the second column. So, click on the next video and let's see what happens with them.

The Electron Configuration: Exceptions (Simplified) Concept 4

So if we look at the 2nd column, let's look at copper. Now copper, if we were to look at the periodic table, we'd initially think it's argon s23d9. But if we look at the 3d9 orbitals, what we should notice is we just need one more electron here, and it will be completely filled. Remember, there's this need driven by your p and d sublevels or subshells to be either half-filled or totally filled. So we're going to say here, an electron from an s orbital can be promoted to create completely filled orbitals with d9 elements. Here, copper is a d9 element. It ends with d9. If you just get one more electron, it can become d10, and that's what's going to happen. We take one from the s orbital, from the 4s orbital, and we donate it over to the d. Doing this now, we only have one electron here within 4s, and this now becomes totally filled in and therefore more stable, so now this is s14d10. Just remember, we have these 6 elements we discussed previously in the videos, they have this drive where we're going to take 1 electron from an s orbital and promote it to 1 of the d orbitals, so that we can create either a half-filled d orbital or totally filled, completely filled d orbitals. Remember this driving force that causes the exceptions within these neutral elements.

The Electron Configuration: Exceptions (Simplified) Example 1

Here in this example question it says, based on the exceptions provide the condensed electron configuration for the silver atom. So we're going to say here that silver is a g. It has an atomic number of 47. Since we're dealing with an atom, it's the neutral form of it, so it has 47 electrons. Now looking at the periodic table, what we would see initially is we would see krypton s2d9. Remember, silver is one of the elements within the second column we discussed, and remember it's a d9 element. If it can become d10 those orbitals will be completely filled. In order to do this, we're going to take 1 electron from the s orbital and promote it up to the d set of orbitals. Doing this gives us now the correct exception to silver, which is Krypton s1d10. So this would be the correct electron configuration for silver. So just keep in mind these 6 elements we discussed earlier, all of them do do this, where we take 1 electron from the s orbital and promote it up, so that we either have half-filled d orbitals or completely filled in d orbitals like silver here.

Illustrate the exception to the electron configuration of molybdenum.

Here it says, illustrate the exception to the electron configuration of molybdenum. Molybdenum is Mo. It has an atomic number of 42, so it has 42 electrons. If you were to look at the periodic table, we're going to write down what we see initially. What we see here is Krypton Kr 5s2, and then we see 4d4. Molybdenum is right below chromium, and remember we said that in that column both chromium and molybdenum are exceptions to our typical electron configuration. Here it's a d four element. Remember, d is most stable when it's half filled. To do that, we're going to take one electron from the s orbital and promote it up to a d orbital, and that way make a half-filled set of d orbitals. So the correct electron configuration for molybdenum is actually Krypton Kr 5s14d5. So now we're going to illustrate that here with this electron orbital diagram. So here, this would be Krypton, this is 4s because remember, I'm in 5s because remember, the s sublevel has 1 orbital, and then these 5 would be our 4d set of orbitals. So we'd have only 1 here in the 5s orbital, and then we have 5 here in 4d. Following Hund's rule, it'd be up, up, up, up, and up. So, this would be the correct illustration for the electron orbital diagram for a molybdenum atom.

Which of the following is the correct electron configuration of gold? 

• A.
• B.
• C.
• D.
• E.

A comparison of the electron configurations of palladium (Pd) and silver (Ag) indicates that: 

• A. Ag has 2 more d electrons and the same number of s electrons as Pd.
• B. Ag has 1 more d electron and the same number of s electrons as Pd.
• C. Ag has 2 more d electrons and 1 less s electron than Pd.
• D. Ag has 1 more d electron and 1 less s electron than Pd.
• E. Ag has 1 more d electron and 1 more s electron than Pd.

Here’s what students ask on this topic:

Why are d subshells more stable when they are half-filled or fully filled?

d subshells are more stable when they are half-filled or fully filled due to electron-electron repulsion and exchange energy. When d orbitals are half-filled, each electron occupies a separate orbital, minimizing repulsion. This configuration also maximizes exchange energy, which is the energy released when electrons with parallel spins exchange positions. Fully filled d subshells are stable because they achieve a symmetrical distribution of electrons, reducing repulsion and increasing stability. This stability is a key factor in the exceptions observed in electron configurations for elements like chromium and copper.

What is the electron configuration exception for chromium?

The electron configuration exception for chromium (Cr) occurs because it is more stable for the d subshell to be half-filled. Instead of the expected configuration of [Ar] 4s² 3d⁴, one electron from the 4s orbital is promoted to the 3d orbital. This results in the configuration [Ar] 4s¹ 3d⁵. This half-filled d subshell configuration is more stable due to reduced electron-electron repulsion and increased exchange energy.

What is the electron configuration exception for copper?

The electron configuration exception for copper (Cu) occurs because it is more stable for the d subshell to be fully filled. Instead of the expected configuration of [Ar] 4s² 3d⁹, one electron from the 4s orbital is promoted to the 3d orbital. This results in the configuration [Ar] 4s¹ 3d¹⁰. This fully filled d subshell configuration is more stable due to reduced electron-electron repulsion and increased exchange energy.

How does Hund's rule relate to electron configuration exceptions?

Hund's rule states that electrons will fill degenerate orbitals (orbitals with the same energy) singly before pairing up. This rule is crucial in understanding electron configuration exceptions because it explains why half-filled and fully filled d subshells are more stable. In elements like chromium and copper, electrons are promoted from the 4s orbital to the 3d orbital to achieve a half-filled or fully filled d subshell, respectively. This minimizes electron-electron repulsion and maximizes exchange energy, leading to a more stable configuration.

Why do elements like chromium and copper have electron configuration exceptions?

Elements like chromium and copper have electron configuration exceptions because achieving half-filled or fully filled d subshells provides greater stability. For chromium, the configuration changes from [Ar] 4s² 3d⁴ to [Ar] 4s¹ 3d⁵ to achieve a half-filled d subshell. For copper, the configuration changes from [Ar] 4s² 3d⁹ to [Ar] 4s¹ 3d¹⁰ to achieve a fully filled d subshell. These configurations reduce electron-electron repulsion and increase exchange energy, leading to a more stable electron arrangement.