The topic of this video is electronic structure of atoms, otherwise known as electron configurations. The learning objectives are also on the screen, so go ahead and pause the video now to write those down in your notes. In a previous video, we discussed quantum numbers, and specifically the principal quantum number referred to the size and relative energy of the orbital, of the atomic orbital that we were looking at.
And really the principal quantum number defined what other orbitals were. available. You know with principal quantum number one there's only an s-type orbital available but as soon as you go up to the principal quantum number two now we could think about p orbitals and so on and so forth. When we want to describe the electronic structure of an atom using something like an electron configuration which we'll dive into in this video we need to really think about relative energy levels. Okay, and so this is heavily tied into principal quantum numbers because the higher the principal quantum number, the higher the relative energy of the orbital.
If we arrange s-type, p-type, d-type, and f-type atomic orbitals as a function of energy, we see that we'll start with at the lowest, this is the lowest energy down here, and this is the highest energy up here. If we were to fill electrons into these orbitals, we would start from the bottom up. We would start from a low energy point and work our way up to a high energy point. Now, if we were to put one single electron down here in the 1s, that is the most favorable place for it to go.
If you were to add another electron, you would have to also add it into the 1s because there's still room for one more electron down here. There's room for one more electron because the arrow up represents spin positive one-half, so that means we can also fit a second one in there with spin of negative one-half. You cannot fit another electron in here because there are only two types of spin, and we can't have, according to the Pauli exclusion principle, electrons with the same four quantum numbers. So these two electrons... don't have the same four quantum numbers, but they're close.
They're just different. They just differ in their spin. There's no more room for other electrons in this orbital.
But you can see that this is a really nice view because now we know, okay, if we want to add even more electrons, we would jump up to the 2s. And then after that, we would jump over to the 2p before filling the 3s. So the observation of starting to fill atomic orbitals from the lowest energy orbitals first is called the Aufbau principle.
The Aufbau principle states is such it can be stated such that each added electron occupies the subshell of lowest energy available subject to the limitations imposed by the allowed quantum numbers according to the Pauli exclusion principle. So all this is saying is that you fill from the lowest energy on up. That's it. This extra language about the Pauli exclusion principle is what I already outlined here. You can't fit another even though the 1s remains the lowest energy atomic orbital in this figure, we would violate the Pauli exclusion principle if we tried to fit a third electron into that.
one is atomic orbital. We can't because to do so means we'd have to have another electron that has the exact same four quantum numbers as either the spin up electron with the up arrow or the spin down arrow, spin down electron represented by the down single barbed arrow. Another really nice way to lay out the electron configurations of different atoms is using a periodic table. And so here, this is a, you hopefully recognize the overall shape and structure of this table, but notice that the typical stuff that you might see in a periodic table are removed. And instead, this is structured in such a way that we are filling starting from the 1s first, and then helium is now considered a 1s2 electron configuration.
Lithium would be now a 2s1. Beryllium would be a 2s2. So the periodic table is really just a roadmap for electron configurations.
That's all it really is. So let's pick an element that I worked with quite a bit as an undergraduate researcher, boron. And let's write out an actual electron configuration of boron and we can use this periodic table to help us do so.
So when you write out an electron configuration there's a few different things that we need. We need to start for the complete electron configuration, we need to start always with the 1s. So for boron we're already past the 1s orbitals so we know that those must be filled already.
So the 1s orbitals must be 1s2. because that's the maximum amount of electrons that an s orbital can accommodate. Not only have we completely filled the 1s here to get to boron, but we've also completely filled the 2s. So right next to that 1s2, we're going to write 2s2. So now we're saying that the first and second shells are, or orbitals, atomic orbitals, are completely filled with two electrons each.
When we get to p orbitals now, there are actually, since there are three p orbitals, each of which can accommodate two electrons. There are six total p electrons. For boron, since it's the very first atom here in the p orbital series, we only need to write 2p1.
So this right here is indeed the electron configuration for boron. Now, um... You can also depict this using what's called an orbital diagram. And so an orbital diagram just sort of uses boxes. So this box will represent 1s.
The next box will represent 2s. And then we can even have a box here for the 2p where we have 3p orbitals, all of which are at the same energy. So if we were to draw out the orbital diagram for boron again, so this is...
... for boron, we would put two electrons into the 1s, two electrons into the 2s, and only one electron into the 2p. So this would look like this. We have two electrons in the 1s, two electrons in the 2s, and only one electron in the 2p.
Let's go ahead and do one more example here. Let's look at the electron configuration of sodium. So sodium is right over here.
I'm circling it in red. So now we know, let me go ahead and erase what I have written before here. So we're looking at sodium. I'll put an X in sodium.
So we know that our 1S is completely filled, our 2S is completely filled, our 2P is completely filled before we get to sodium. So here I'm going to write out, I'm going to change back to white, 1s2 is filled, 2s2 is filled, 2p6 is filled, and now we're at 3s, and there's only one electron at sodium's position. One electron in the 3s orbital.
We could draw out the entire orbital diagram here, so this is the 1s2s. This is going to be our 2p orbitals. And then also we have our 3s. So this is 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11 electrons total. Now you can see that this would get tedious.
We're only up to sodium, which is not one of the heavier elements. You can imagine that if you had to do this, for example, xenon or even barium, the electron configuration is going to be really really tedious. It's going to be redundant with many other atoms and it takes a while to write it all out.
So thankfully there is a designation for how to handle this. Let's take a look at sodium. So for sodium in particular I'm going to rewrite the electron configuration over here.
1s2 2s2 2p6 and 3s1. There's a break here where the principal quantum number jumps from 1 to 2 and 2 to 3. But the outermost one, the one that has the largest size or the highest relative energy is the highest principal quantum number value of 3. All of the electrons before that break, that principal quantum number 3, these are the core electrons. Okay.
And the outermost electrons, outermost electrons have a special name. They are referred to as valence electrons. Okay, so what we have, what we can do is we can use this shorthand notation to sort of get, we don't have to write out all the core electrons every single time.
What we can do is we can actually take a look and see what element has the same electron configuration as these core electrons. So which element has the same configuration as 1s2, 2s2, 2p6? That is the element, erase these again.
The element there is going to be, so we have 1s2 is gone, 2s2 is good, 2p6 is neon. So what we can do is we can actually say, okay, forget writing out all those core electrons. We can just summarize all of those using a bracket neon, a closed bracket notation, followed by the valence electrons, 3s1. So this is now the abbreviated electron configuration.
On your screen now, I have the electron configuration table where the valence electrons are included with each element. And what makes the periodic table of elements periodic? and what makes it a very useful predictor of chemical properties is the fact that it is arranged in such a way that each element or all the elements are grouped according to similar valence electrons. Finally, we're going to end on electron configurations of ions. So if you are asked to write the electron configuration of ions, you will presumably be given the ion.
So that's over here. The ions are here that will be... practicing with. And the first step is to write out the electron configuration of the atom itself, of not the ion, but just the atom.
Okay, so this is electron config of atom. I have chosen to use the abbreviated notation here because I don't feel like writing out all of the core electrons. The next question is to ask yourself, To form the ion in question were electrons added or removed.
So if we have cations, which are the first three samples here, that means that the ion is inherently positively charged. In order to become inherently positively charged, it must have lost a negative charge. And so these all three cases, the cations, to get to the cation from the neutral atom, we must remove electrons.
How many electrons do we remove? Well, that depends on the magnitude of the charge. Sodium only has a one plus charge, so it lost one electron.
So the ion electron config would just be removing one electron. And there's only one electron to remove from the valence shell, the 3s1. So literally the only thing that you would need to write for the abbreviated electron configuration of the sodium ion is bracket neon.
The next ions are both iron. ions, iron 2 plus, iron 3 plus. The rule for transition metals is that you remove electrons from s orbitals first. So iron 2 plus, we need to remove two electrons.
And since it's a transition metal, we are going to remove them from the 4s2 first. So the electron configuration for iron 2, for the iron 2 plus ion is going to be argon core electrons 3v6 as our valence electrons. Now The reason that I wanted to look at iron 3 plus is because now we can essentially say, well, similar to iron 2 plus, we're going to remove the 4s2 electrons first. But we also need to remove one more because this is a 3 plus cation.
So we actually have the argon core 3d5 for iron 3 plus. Now in the case of chloride, which is an anion, in order to get from a neutral atom to a negatively charged ion, we need to add. one or more electrons. This only has a single negative charge, so we needed to add one electron to this.
If we're going to add the electron, we need to add it to the subshell that can accommodate it. The 3s subshell is already occupied completely, but the 3p can occupy one more, or can accommodate one more electron. So for chloride, we would write the neon core, and you can write...
3s, 2, 3p, 6. But now you notice now we have a complete shell here. where the 3P subshell is completely filled. So you could write it this way, neon core 3S2 3P6. Alternatively, that is also the valence of argon.
And so we could also write this as just the argon noble gas configuration.