 So, when it comes to structure, what we're going to need to decide is what are the atoms in the structure, but just looking at individual atoms is fairly pointless. So, we're going to need to consider how and why atoms interact, how strongly they interact, and why some atoms prefer to interact with other atoms. You've probably taken this in various forms, either in upper secondary school or at university, but I'm going to introduce this in a slightly different way, possibly. This is more chemical than physical way. At the end of the day, the reason why any atoms interact is due to electrons, and electrons is in principle complicated. You should solve the Schrodinger equation, but if you're a chemist then we need to be able to have an understanding. We tend to use something called orbitals, which is a schematic way to display the way electrons would interact with each other so that we can hand wave about how bonds work. But make no mistake, this is an approximation. And if you haven't heard of orbitals before, you definitely heard of it indirectly in upper secondary school. We talked about electrons in shells. You might remember you had this first S shell, and then three P electrons, and then D electrons. And this is largely the same concept it has to do with various energy levels. In each of these shells, you can have two electrons. And the reason we have two of them is that electrons are paired. You might also remember that you talk about some sort of spin-up property and some sort of spin-down property. Of course, that has nothing to do with up and down, but you can think of a sine equation. A sine is sometimes positive and sometimes negative. But if you now take the sine equation and push it 180 degrees pi, you're now going to have another sine wave that's also sometimes positive and negative, but it's offset. So one one is positive, the other one is negative. So this is roughly the same concept. The electron maintains its charge. It has to do with density and where the density probability is highest. But that also means for the S orbital, that's easy. That's rotationally spherically symmetric. But for these other orbitals, there is a very clear spatial favorability here. The P orbitals are, think of that as Cartesian axis, X, Y, Z, they are orthogonal. And for the D orbitals, it's even more complicated. So why do I tell this? It turns out that this concept of orbitals is going to tell us when atoms are likely to bind. What happens if two atoms get closer to each other? When will they repel each other versus when will they bind? And it's also going to explain lots of concepts such as hydrogen bonding. What would happen if you had two of these with spin up? So if I erase that spin down for a second and then I put it as spin up. Well, what then would happen is that you would have two electrons that have exactly the same quantum mechanical ground state as they approach each other. And that leads to something called the Pauli exclusion principle. And you could even prove that's not that hard mathematically, that that would lead to an exponentially strong repulsion. So it's going to be very, very bad in terms of energy as they get close. So I'll erase that one. So I don't have something bad written on screen. In practice, it gets slightly more complicated because sometimes if you're a chemist, you can have double bonds and things like that. I'll worry about that when we get there. We might not get there that much in this class. So can we use this to understand binding? Let's try. This is something very simple. It's two hydrogen atoms, I think. And each hydrogen atom has one proton and one electron. And it's not, well, at very high temperatures in the plasma or something, they would be happy that way. But hydrogen is a fairly reactive molecule. So under natural conditions, what's going to happen is that you have these two hydrogens. And as they approach each other, what will happen is that if both of these electrons start, instead of going, if you have one hydrogen atom there, and then one hydrogen atom there, if this electron and that electron, instead of being separated in the two hydrogen atoms being unpaired, what can happen is that we have our hydrogens. And then we bring them really close. That might be a little bit too close. But basically you would have both these electrons encompass both of the hydrogens. What will then happen is that around the left hydrogen atom here, you're effectively going to have both spin up and spin down. And around the right atom, you're also going to have both spin up and spin down, because they will pair up that way. So what's now going to happen is that instead of having two incomplete atoms, you're going to have two atoms that each feel that they're complete. And this is going to lead to a chemical bond. What complicates things now is that this gets difficult to explain to a teenager or something. And then we start inventing all these things. We have covalent bonds. We have ionic bonds. This is what I would call an ionic bond. Sorry, covalent bond, my bad. But in reality in life, this is a sliding scale. It's not going to be one fixed distance when it starts being one bond and it cuts over into being another type of bond. But I'm going to go through a couple of different ways these electrons interact at different distances and see if we can learn a little bit more about that. And I'm not going to go that much deep into the orbitals for now. We'll come back to the orbitals once we start entering the hydrogen bonds.