Atomic structure and electron configuration
Shells, subshells, and orbitals
- [Instructor] We've learned in other videos that the atom is in fact made up of even smaller constituent particles which is pretty amazing because atoms are already unimaginably small. And those particles are the protons which have positive charge, you have your neutrons which have neutral charge or no charge and then you have your electrons which have negative charge. Now, the big question physicist and chemists were facing over a hundred years ago is how are these things configure and they realized that the positive charge is concentrated at the center of the atom. In fact, most of the mass which is made up of the protons and the neutrons is concentrated at the center and so the early model for how an atom worked was maybe you have your protons and neutrons in the center so let's say, we're talking about a helium atom. A helium atom has two protons in the nucleus and a typical helium atom would have two neutrons as well so the nucleus might look something like that. And early physicists and chemists said, "All right, well, if the protons have a positive charge, "electrons have a negative charge, "so they'll be attracted to each other." Opposite signs, opposite charges attract. The same charge repels each other. So maybe the electron which has a negative charge orbits around the nucleus the way that a planet would orbit around its star. So, maybe it orbits something like this. So maybe one electron has an orbit that looks something like that, and then another electron, if we were talking about a neutral helium atom will have two electrons and two protons, well maybe, the other one orbits something like this. I'm trying to just draw an elliptic or a circular looking orbit. And so this is the idea that the electrons are in orbits. Now, it turns out that this is not exactly the case. Electrons are not in these well-defined circular or elliptical orbits. In fact, at any given point in time, it's not necessarily exactly right there, it could be there but there's some probability it's here, there's some probability that it's there, there's some probability that it's there's, some probability that it's over there. And so to describe where electrons are likely to be found, physicist and chemists introduced to the idea of an orbital and the best way to think about orbitals is to think about a hydrogen atom and actually the map for orbitals, it's hydrogen as the simplest atom and so the map for orbitals has been best completed for the hydrogen atom. So then a hydrogen atom especially the typical isotope of hydrogen found on Earth, the nucleus actually has no neutrons. You just have a single proton at the center. And if you have a neutral hydrogen atom, that one electron, instead of being in orbit around that one proton like that, we can really just think about the probabilities of where it might be. It could be here, it could be here at any given moment, it could be there at any given moment, it could be off the screen at some moment but it's more likely to be in certain regions of space around the nucleus and others. And we can visualize where it's most likely to be by saying, "All right, it looks like 90% of the time, "it's in a sphere that looks something like that." But once again, it could be here, it could be there, it could be there, could be there, could be there, could be out here, it could be anywhere. We're just saying where it happens to be 90% of the time. That's the visualization. Now, an interesting question is what if you were to give that electron a little bit more energy? Well, what does energy mean? Well, if you think about planets or rocket or satellite orbiting around, if you were to give it a little bit more energy, if you ere to give it a little bit push, it could have a larger orbit, it would look something like that. But quantum mechanics isn't about things happening gradually. Sometimes people think quantum means small or something like that. No, it really means that you're talking about discreet packets. So in quantum physics, quantum chemistry, if you add a certain amount of energy to an electron, instead of having a 90% chance of being found in this first shell, this first energy level, it could then be found, it would then jump into the next energy level or the next shell. And so now, it might be more of, 90% of the time, it's going to be found in this shell right over here. And then if you were to give it the right boost of energy, once again, just a little bit won't do, you have to give it enough so then it jumps into the next energy level, then it might form this weird patterns that looks kinda like dumbbells where 90% of the time, it's kind of you can view it as it's on the orbital that looks kind of like that dumbbell shape. I just did in kind of the horizontal direction. You can have it in a vertical direction. You could also have it on the in-out direction of this page. And if you were wondering where did these shapes come from and if you keep adding more and more energy, you get these more and more exotic shapes for orbitals, think about standing waves. That's my best hint I can give you that the quantum level, actually at all levels, but especially at the quantum level, you see things like electrons have both particle and wave-like properties. Imagine something like a standing wave where if I were to just take a rope and if I were to just shake it, I might get standing waves that look like that. If I were to take a some type of a membrane in two dimensions and if I were to push on one side right here if I were to drum on that, you might get, so this part dips down, and then that part dips up. And so when you get to three dimensions, you end up getting this dumbbell shape when you add more energy and then you get more and more and more exotic shapes, just to imagine what some of the first orbitals look like rendered by a computer, you see it right over here. So if you have your lowest energy electron, you are in what is called an S-orbital right over here and this one we would call 1s 'cause it is at the first shell, the one closest to the nucleus. If you give even more energy, then that electron might jump into the second energy level or the second shell and the orbital in that second shell which would be the default if it's the lowest energy in the second shell would be the 2s orbital. Once again, you have this spherical orbital, it's just a little, it's more likely to be found further out than the one, it was just in the one shell. Once again, if you add even more energy, you'll fall, you'll still be in the second shell but you will be into one of these orbitals that have higher energies so you could view this as the 2p orbital that is in the x-dimension. This could be the 2p orbital that is in the y-dimension as some people call that 2px. Some people would call that 2py. This you could view as the in and out of the page so you could view that as the z-dimension. So that is 2pz and the orbitals keep going. There is a d-orbital once you get to the third shell. Once you get to the fourth shell, there is an f-orbital. All we've talked about right now is an hydrogen. If you keep giving more energy to that one electron, what happens to it? What is the shape of the probabilities of where it might be into two-dimensional space? As you can imagine, if you have two electrons, it's not exactly the same but this is pretty good approximation. You can actually put two electrons in this 1s orbital but after that, you can imagine the electrons are repelling each other. So another electron doesn't wanna go there so the third electron that you add is going to end up in the 2s orbital. It's gonna be at that higher energy level and then that can fit two. So you can fit up to four electrons between the 1s and the 2s. And then the fifth one is going to have to go into one of these p-orbitals. Now, one last point, just to make sure you understand the terminology of orbitals and shells. So first of all, you have this idea of shells and sometimes the word shell will be used interchangeably with energy level, energy levels. And so in this diagram or this a visualization right over here, I've depicted the one shell and then I've also depicted the two shells. So this is a shell right over here. This is another shell. Now, you'll also hear the term, subshell, subshell, or sometimes people will say sublevels and that's where they're talking about s or p or d and eventually f so if I circle this, I'm talking about that first shell. Now, the first shell only contains one subshell and that's the 1s subshell and the 1s subshell only has one orbital. Once again, the 1s orbital. So for the first shell, the shell, the subshell, the orbital is all referring to the same thing, but as we get to the second shell, it's a little bit different. If we're talking about the subshells, in the second shell, there's s and p so this is a subshell, and then this is another subshell right over here. There's actually three orbitals in the p-subshell. So I'll leave you there. In the next video, we'll actually look at various atoms and think about their electron configurations. Where do their electrons sit? In which of these shells, subshells, and orbitals?
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