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Electron configurations of the 3d transition metals

The Aufbau principle predicts that the 4s orbital is always filled before the 3d orbitals, but this is actually not true for most elements! From Sc on, the 3d orbitals are actually lower in energy than the 4s orbital, which means that electrons enter the 3d orbitals first. In this video, we’ll discuss this in more depth and walk through all of the electron configurations for the 3d transition metals.

Created by Jay.

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  • piceratops tree style avatar for user Michael
    At , Jay says that the 4s orbital is the "highest in energy." However Sal in a previous video said that while the 4s orbital would be the "furthest away," in fact the 3d orbital is the "highest energy." He said this at in the video "electron configuration for d block element" while doing the electron configuration for Fe. Which is correct?
    (96 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      4s is higher in energy than 3d until you get to Ca. Then the relative energies of 4s and 3d switch.
      As you go from left to right across the Periodic Table, you are adding extra protons to the nucleus and extra electrons around the nucleus.
      The various attractions and repulsions in the atoms change as you do this. It is these attractions and repulsions that govern the energies of the various orbitals.
      For rather complicated reasons, once you get to scandium, the energy of the 3d orbitals becomes slightly less than that of the 4s.
      The "furthest away" argument is suspect. Nobody really knows how big a 4s or a 3d orbital is, except in an H atom where there is only one electron.
      (53 votes)
  • piceratops ultimate style avatar for user Utkarsh Sharma
    Why do Chromium and Copper behave so weirdly ?
    Moreover will their EC's be Cr:{Ar}4s1 3d5 only if we write it in actual ?
    Cu:{Ar}4s1 3d10 only if we write it in actual ?
    Moreover if this does not happen in real what is the point of showing and understanding that weird Orbital Notation ?
    (15 votes)
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    • orange juice squid orange style avatar for user Luke Yakielashek
      Cr and Cu are the two exceptions of electron configuration of atoms up to Kr. this is because a 1/2 or completely full D block has extra stability, therefore in the case of Chromium one electron will shift from the 4s block to fill the 3d block to exactly one half
      Cr: [Ar] 4s2 3d4 actually exists as: Cr: [Ar] 4s1 3d5 and;
      Cu: [Ar] 4s2 3d9 actually exists as: Cu: [Ar] 4s1 3d10

      This idea holds true for Mo and Ag in period 5 and so on
      (40 votes)
  • blobby green style avatar for user Sharan
    How can we write the electronic configuration for an element if the periodic table is not given to us ?
    (6 votes)
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    • piceratops ultimate style avatar for user Just Keith
      You must know the atomic number of the element. Without that, you cannot determine the electron configuration.

      Given the atomic number, there is a set of rules that allow you to determine the electron configuration. However, in the simplified version of these rules that is understandable to a general chemistry student, there are about 19 elements that are exceptions. These you just have to learn.

      What you do is you start assigning electrons to the subshells using the following pattern (you completely fill up one subshell before moving onto the next higher) and you keep going until the total number of electrons assigned is equal to the atomic number:
      1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
      There are no known elements that, in their ground state, have electrons in a subshell beyond 7p. This pattern will give you the correct configuration for all but about 19 elements.

      The actual filling order is more complicated. For example, the 4s and 3d have nearly the same energy level and so, in a way that varies according to which element it is, either of the two subshells could be the first to fill and there is no certainty that one will be completely filled before the other actually starts filling. There are similar irregularities in the filling of 5s and 4d as well as some other pairs of subshells.

      So, rather than going through the far more difficult to understand rules that do correctly predict the electron configurations for all elements, it is easier just to learn the elements that this simplified method doesn't work correctly for. Here they are:

      In the following elements, there is 1 electron, not the predicted 2, in the outermost s subshell, the "missing" electron is instead located in the outermost d subshell:
      Cu, Cr, Nb, Mo, Ru, Rh, Ag, Pt, Au

      In Pd (palladium) the outermost s subshell, predicted to have 2 electrons, is actually empty. Both of these electrons are instead in the outermost d subshell. Thus, the actual configuration of the Pd is [Kr] 4d¹⁰ (the 5s is empty).

      All the other exceptions are in the f block and just have to be learned on a case-by-case basis. However, it is unlikely you would be asked to know these in a general chemistry course. Here are the f-block exceptions:
      La, Ce, Gd, Ac, Th, Pa, U, Np, Cm, Lr
      Note: in one of his videos, Sal does the configuration of La, but he gets it wrong. He did not know that La is an exception to the rule.

      There is a simple chart that makes learning the filling order easy. Here is a link to it:

      Note: Nickel has two electron configurations that are commonly encountered over which there is a dispute amongst chemists as to which should be considered the ground state. I have assumed that the configuration that follows the rules is the ground state, since that is what is usually taught in general chemistry. However, this is not certain.
      (41 votes)
  • aqualine ultimate style avatar for user Srilakshmi Ajith
    At , Jay said that there are many other factors to consider, what are the other factors that need to be considered?
    (6 votes)
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  • leaf green style avatar for user Assaf Shomer
    You keep saying that 4s orbital electrons have higher energy than 3d orbital electrons (for scandium).
    In addition, I thought that electrons want to fill the positions with the least energy before they fill positions with higher energy.
    Combining those two things I do not understand why would the higher-energy 4s orbital fill with electrons before the lower-energy 3d orbital.
    Is it because the repulsion of the electrons filling the 3rd shell makes it "harder" for the electrons to occupy that state rather than the 4s state, despite the last is of higher energy state?
    Anyway, this remained unclear to me throughout the video.
    Thanks in advance
    (4 votes)
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    • piceratops ultimate style avatar for user Just Keith
      The 4s and 3d subshells have nearly the same energy level. How much of a difference, and which subshell is lower in energy, varies by element.

      In scandium and the other d block members of period 4, the 3d is lower in energy than the 4s. However, because the 3d orbitals are smaller and more closely packed than the 4s, the electron-electron repulsion that occurs in the 3d subshell is enough to make it possible for 1 or 2 electrons to have a lower energy by occupying the 4s. So, we have 3d and 4s co-filling rather than filling independently.

      So, with scandium, the first electron after completing the previous noble gas configuration goes into 3d. The difference in energy levels between 4s and 3d is small enough for the next two electrons to enter the 4s, because the electron-electron repulsion they feel if they enter 3d is enough to knock them into the 4s.

      The same process happens throughout the d-block. Sometimes the difference in energy levels between the s and d allows one electron to enter the s, sometimes it the difference allows 2 electrons to enter the s. And in palladium alone, the difference between d and s is so great that none of the electron are in the s.

      This leaves us with why you have s filling in Groups 1 and 2 instead of the d. That is simple. For the first few members to of Groups 1 and 2 there is no d to be filled. But with 4s-3d, 5s-4d, etc. there is the d to take into consideration. The explanation is simple: in Groups 1 and 2, but not in the d block, the s is lower in energy than the d. In the d block it is the other way around.

      Because the difference in energy levels between 4s-3d, 5s-4d, etc. varies by element, we have to look at each element on a case-by-case basis to see how these will fill.

      Similar irregularities appear in the f block, only now there are s, d and f subshells all with similar energy levels. So, those also have to be learned on a case-by-case basis.
      (9 votes)
  • marcimus pink style avatar for user eman
    Around , it is said that it is incorrect to imagine one of Cr's electrons to 'hop over' from the 4s subshell to 3d. Why or how is that incorrect? Isn't that exactly what's happening! What's the "real truth" that is too complicated for GenChem?
    (4 votes)
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    • piceratops ultimate style avatar for user Just Keith
      The real truth is that the 4s and 3d subshells have nearly the same energy level. The difference between these two levels varies by element. In the d block it is actually the 3d that starts filling first if you strip away the electrons from a single element and then start adding them back -- but this is not apparent if you look at the final configuration as you go across the d block.
      Because the 3d and 4s (the same thing applies to 4d and 5s) are so close together in energy level, the repulsion between electrons in the 3d can be enough to cause one or two electrons to enter the 4s without the 3d being filled completely. Thus, in most of the d block, 2 electrons are kept from entering the d subshell because of electron-electron repulsion and thus enter the next higher s subshell. But in some of the elements the difference between the d and the s is too great for 2 electrons to be "pushed" from the d into the s, so only one electron winds up in the s. And, in palladium alone, the difference between the s and d is too great and none is "pushed" into the s.

      There is no strict pattern to which elements have only one electron in the outermost s in the d block, so you just have to learn them.
      The following d block elements have one electron in the outermost s:
      Cu, Cr, Nb, Mo, Ru, Rh, Ag, Pt, Au
      And, as mentioned, palladium is unique, with no electrons being "pushed" into what would be its outermost s. Thus, Pd's configuration is:
      1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ (the 5s is empty) or in the noble gas configuration notation: [Kr] 4d¹⁰

      However, in Group 1 and Group 2, it really is the 4s and not the 3d (likewise for 5s and 4d) that has lower energy and fills first. But in the d block, they switch places with the d having lower energy.
      (6 votes)
  • starky tree style avatar for user Shreet Dave
    Why should Scandium have Energy level of 3d lower than 4s? As much as I know this thing called chemistry, every exception made in chemistry has a reason, an explanation. So why such an exception?
    (0 votes)
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    • piceratops ultimate style avatar for user Just Keith
      3d and 4s have nearly the same energy level. The difference between the two varies from element to element. So, if the two levels are close enough on a particular element, one or two electrons can get bumped up to 4s due to electron-electron repulsion being greater than the difference between the two energy levels. A similar situation happens in period 5 with 5s and 4d. While most of the elements of the d block have the relevant s and d of close enough energy for 2 electrons to get bumped up the the s, there are some in which the difference is not small enough and only one gets bumped up. In Pd none of the d electrons get bumped up to the s -- thus it is the only element which contains no electrons in the shell to whose period it belongs.

      Note: In Groups 1 and 2, 4s is lower in energy than 3d, but starting with Sc, 3d is slightly lower than 4s.
      (10 votes)
  • piceratops sapling style avatar for user Nick Winters
    Why do scientists not use original names for some elements? Wouldn't it be better to keep the abbreviations consistent with names?
    (3 votes)
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    • leaf red style avatar for user Richard
      Essentially there are different groups responsible for the name of the element and its chemical symbol. The chemical symbol is standardized by the IUPAC, while each country/region/language will have its own name for the elements. The reason is that these regions had their own history with the elements, particularly the earliest ones discovered, well before the science of chemistry or IUPAC were a thing.

      Iron is a good example. Humans have known about iron long before we knew about elements and atoms. Different groups of humans have their own history and way of referring to iron which differ from each other. The symbol Fe, refers to the Latin name for iron, ferrum. But that doesn’t mean everyone is going to refer to it by its Latin name, and they don’t. In French it is fer, Spanish is hierro, Italian and Portuguese is ferro. These are all the romance languages based on Latin and even though they sound similar to ferrum, they all have their own variation. Humans will generally use their own regional variation of a name for the elements as opposed what the symbols represent.

      Hope that helps.
      (5 votes)
  • leaf grey style avatar for user Stephen Marts
    Suddenly chemistry seems much less a predictable field of science than a "I hope this works and I don't know why but someone says it works" kind of thing. Like a "memorize this because it'll help you pass a test but it isn't true" sort of thing that keeps the world going spinning...
    (3 votes)
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    • leaf red style avatar for user Richard
      There is a logic to the madness. The issue is that the full explanations for certain chemistry phenomena, such as the energy differences between 4s and 3d subshells, requires a decent amount of math which mires most people. So to simplify things chemistry taught without a full mathematical explanation so as not to confuse students (or at least minimize confusion). Chemistry it taught to novice students as an edited version.
      (4 votes)
  • blobby green style avatar for user Hari Shankar Karthik
    how does the 4s orbital suddenly rise to a higher energy level than the 3d orbital while writing the electronic configuration for the d-block elements?
    (3 votes)
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Video transcript

- [Voiceover] We've already looked at the electron configurations for potassium and for calcium but let's do it again really quickly because it's going to affect how we think about the d orbitals and so we find potassium which is in the fourth period on the periodic table. If we do noble gas notation to save some time, we work backwards and the first noble gas we hit is argon, so we write argon in brackets. We know argon has 18 electrons and potassium has 19 electrons. Potassium has one more electron than argon and so we put that extra electron into a 4s orbital because for potassium the 4s orbital is lower energy than the 3d orbitals here. We have increasing energy and that electron goes into a 4s orbital so the complete electron configuration using noble gas notation for potassium is argon in brackets 4s 1. For calcium I should say. All right, we have one more electron then potassium and so that electron's going to go into the 4s orbital as well and so we pair our spins and we write the electron configuration for calcium as argon in brackets 4s 2. For the calcium two plus ion, so if you're thinking about forming an ion here, we're talking about the electron configurations for a neutral atom meaning equal numbers of protons and electrons. With the atomic number of 20, 20 protons and 20 electrons. If we lose two electrons, we have a net deposited two charge. We form the calcium to ion. The two electrons that we would lose to form the calcium two plus ion are these. These two electrons right here in the 4s orbital. The electron configuration for calcium two plus would be the same as the electron configuration for the noble gas argon here. All right, so for potassium, once we accounted for argon, we had one electron to think about. For calcium, once we counted for argon we had two electrons to think about. If we go to the next element on the periodic table, that's scandium. That's one more electron and calcium. We have three electrons to worry about once we put argon in here like that. This is where things get weird. Now we have to think about the d orbitals and once again things are very complicated once you hit scandium because the energies change. When you hit scandium even though these are very small energy differences, now the energy of the 4s orbital is actually higher than the energy of the 3d orbitals. We're talking about once again increasing energy and so that's pretty weird. All right, so if you think about these three electrons, where are we gonna put them? Well your first guess, if you understand these energy differences might be, okay, well I'm gonna follow Hund's rule. I'm gonna put those electrons in the lowest energy level possible here and I'm going to not pair my spins and so I'm going to write my electron configuration like that for scandium. You might think it would be argon 3d 3 but that's not what we observed for the electron configuration for scandium. Actually two of these electrons actually move up to the higher energy orbital so two of those electrons move up to the 4s orbital here like that. The electron configuration turns out to be 4s 2, 3d 1. It's actually 4s 2, 3d 1 or if you prefer 3d 1, 4s 2 once again with argon in front of it. Either one of these is acceptable. This is weird so like why did those electrons, why did those two electrons go to an orbital of higher energy? There's no simple explanation for this. All right, so even though it might be higher in energy for those two electrons, it must not be higher energy overall for the entire scandium atom. There are many other factors to consider so things like increasing nuclear charge. The scandium has an extra proton compared to calcium and then there are once again many more factors and far too much to get into in this video. Unfortunately there is no easy explanation for this but this is the observed electron configuration for scandium. How do we know this is true? How do we know that the 4s orbital is actually higher energy than the 3d orbitals? We know this from ionization experiments. For example if you form the scandium plus one ion, the electron configuration for the scandium plus one ion, so we're losing an electron from a neutral scandium atom. This turns out to be argon 4s 1, 3d 1 or once again you could write argon, 3d 1, 4s 1. Where did we lose that electron to form our ion? We lost that electron from the 4s orbital. We had 4s 2 here and here we have 4s 1. We lost this electron and that only makes sense if the 4s orbital is the highest in energy because when you lose an electron for ionization, you lose the electron that's highest in energy. That's the one that's easiest to remove to form the ion. The 4s orbital is actually higher in energy than the 3d orbitals. You don't see this a lot in text books and I think the main reason for that is because of the fact that if you're trying to think about just writing electron configurations. Your goal is to write, let's say you're taking a test and your goal is to write the electron configuration for scandium. The easiest way to do that ... Let me go ahead and use red here. The easiest way to do that if you want to write the electron configuration for scandium, you look at the periodic table and if you're doing noble gas notation, the noble gas that precedes it is of course argon right here. That takes care of the argon portion and then looking at the periodic table you would say this could be 4s 1, 4s 2, 3d 1. That gives you the correct electron configuration, argon 4s 2, 3d 1. But it's implying that the d orbitals, the 3d orbitals fill after the 4s orbital and is therefore a higher energy and that's not true actually. It does help you to just assume that's the case if you're writing an electron configuration but that's not what's happening in reality. We need to think about the other elements here. We just did scandium. Next let's move on to titanium. Thinking about titanium, so the next element in the periodic table if your question on the test was write the electron configuration for titanium, the easiest way to do it is just once again to think about argon. Put argon in brackets and then think to yourself, this would be 4s 1, this would be 4s 2, this would be 3d 1 and this would be 3d 2. You could write 4s 2 and then 3d 2 or once again you could switch 3d 2 and 4s 2. Once again this is implying the d orbitals fill after the 4s orbital which isn't true but it does get you the right answer. It's useful to think about it both ways. It's useful to think about the energy levels properly but the same time if your goal is to get the answer the fastest way possible, looking at the periodic table and running through the electron configuration might be the best way to do it on test. Let's look at some of these other elements here so we've just talked about scandium and titanium. All right, so let's go down here. Let's look at this little setup here. All right, so we just did scandium and titanium. All right, so scandium was argon 4s 2, 3d 1. We talked about two electrons in the 4s orbital, one electron in the 3d orbital. We just did titanium 4s 2, 3d 2 or once again you could switch any of these. When you're doing orbital notation, adding that second electron to a d orbital. Here's the electron that we added so we didn't pair up our spins. We're following Hund's rule here. Next element is vanadium so we do the same thing. One more electron, we add that electron to a d orbital but we add it to, we don't add it to one of the ones that we've already started the fill here, we add that electron to another d orbital, so once again following Hund's rule. Things get weird when you get to chromium. Let me use a different color here for chromium. If you're just thinking about what might happen for chromium, chromium one more electron to think about than vanadium. You might think, let's just add that one electron to a 3d orbital like that and then be done with it. If you think about it, you might guess 4s 2, 3d 4. Let's go ahead and write that. 4s 2, 3d 4, so question mark but that's not actually what we get. We get 4s 1, 3d 5. That electron, this electron here, let me go ahead and use red. So you could think about this electron. We expect it to be there, we expect it to be 4s 2, 3d 4. It's like that electron has moved over here to this empty orbital to give you this orbital notation. All right, so that's just an easy way of thinking about it and in reality that's not what's happening if you're building up the atom here because of the different energy levels. But just to make things easier when you're writing electron configurations, you can think about moving an electron from the 4s orbital over to the last empty d orbital here. Some people say that this half filled d subshell, let me go and circle it here. This half filled d subshell is extra stable and that might be true for the chromium atom but it's not always true so it's not really the best explanation. The real explanation is extremely complicated and actually just way too much to get into for a general chemistry course. Next element is manganese. All right, let me go ahead and stick with blue here. Manganese, one more electron than chromium here. Chromium we had six electrons here, and manganese we need to worry about seven electrons. This is kind of what we expect, just going across the periodic table. Let me go ahead and do this for manganese. Let me use green here. You might say okay, that's 4s 1, that's 4s 2 and then 3d 1, 3d 2, 3d 3, 3d 4, 3d 5. You might say to yourself 4s 2, 3d 5. This precedes how we would expect it to. All right, and the same thing with iron, so 4s 2, 3d 6. All right, so that takes care of iron and once again now you can start to pair up your spins. We've seen that in earlier electron configurations. Next cobalt, one more electron to worry about. All right, so 4s 2, 3d 7 makes sense and you can see here would be the electron that we added and we paired up our spins again. Nickel, same trends. We add one more electron, 3d 8. That makes sense, here's the electron that we added and once again we got a weird one. All right, so when we get to copper. So copper you might think ... Let me use red for copper so we know copper's red. We think about it, writing one more electron. If we took the electron configuration here for nickel, we added one more electron. You might guess that would be the orbital notation for copper but that's not what we see. We've taken this electron here and moved it over to here, like that. This gives us a filled d subshell here. Once again one explanational see for that is extremely stable for copper and that might be true for copper. All right, and that leaves us only one electron here in our 4s orbital. The electron configuration is 4s 1, 3d 10 but all these general chemistry explanations are just a little bit too simple for reality but if you're just starting out, they're pretty good way to think about it. Then finally zinc, zinc makes sense. We're adding one more, writing one more electrons. We just took care of copper. For zinc we have one more electron and so you could think about this being 4s 2 right here and then we have 3d 10, one, two, three four, five, six, seven, eight, nine, 10. 4s 2, 3d 10 or 3d 10, 4s 2 with argon in front of it gives you the complete electron configuration and you can see, you've now filled your 4s orbital and your 3d orbitals like that. Once again pretty complicated topic and hopefully this just gives you an idea about what's going on.