Electron configurations for the second period

- Let's do electron configurations for the second period. So we find the second period on the periodic table and we go across and the first element we see is Lithium with three electrons. So three electrons to worry about for Lithium. Let's think about the first two electrons for Lithium. The first two electrons for Lithium are going to go into the first shell. So we talked about this in the last video. The first shell when n is equal to one, the only possible value for l is zero. So we're talking about an s orbital. And there's only one s orbital in the first shell here so I can draw in, let me go ahead and draw that orbital in. So here's the one s orbital in the first shell. So Lithium has three electrons. The first two electrons for Lithium are going to go into this one s orbital. So we pair up our spins like that. So writing the electron configuration for Lithium, let me go ahead and we'll start writing it right here. So we have one s two so far. Well Lithium has three electrons but the first shell is full, it's closed. So we have to move on to the second shell to add in Lithium's third electron. So in the second shell, n is equal to two. What are the allowed values for l? L could be equal to zero or l could be equal to one. So we talked about that again in the videos on quantum numbers. So when l is equal to zero, we're talking about an s orbital. So in the second shell, in the second energy level, we also have an s orbital and we also have one of them and we also have to think about l is equal to one, that's talking about a p orbital. The allowed values for ml would be negative one, zero and positive one. So three possible values means three p orbitals here. So we have three p orbitals in the second energy level as well. So let's draw those in on our orbital diagram over here. So we already drew in this s orbital in the first shell. Next let's draw in this s orbital in the second shell, the second energy level. It's of higher energy. So we draw it in here. This is the two s orbital. Then we also have p orbitals in the second energy level, we have three of them. So we draw in our p orbitals in the second energy level. They're of higher energy so here are the two p orbitals and there are three of them. So one of them is, it would be two px, one of them would be two py and one of them would be two pz. Doesn't really matter which one is which. We'll just draw them in there like that so far. All right, so Lithium. We've taken care of the two of it's three electrons. It's third electron has to go into this next highest orbital in terms of energy, so that would be the two s orbitals. We have energy going up this way. So as you get higher and higher, energy increases. So Lithium's next electron, as you build up the Lithium atom must go into this orbital here. The two s orbital. So therefore, Lithium's electron configuration is one s two, two s one and remember what these numbers mean. So this means that there is one electron and an s orbital in the second energy level. So we have one s two, two s one for Lithium's electron configuration. Let's do the next element. So that's Beryllium. Beryllium has four electrons to worry about. So for Beryllium, if you look at the diagram, Let's see if we can just make a different color here for Beryllium, so let's make Beryllium red here. So one more electron. So we can put Beryllium's fourth electron into this orbital and pair up our spins. So let's write the electron configuration for Beryllium. So it would be one s two and then we have two electrons and then two s orbitals, so we would write two s two here like that. Now, we've filled the two s orbitals. Remember, each orbital can hold a maximum of two electrons. We filled the two s orbitals so when we move on to the next element, which is Boron over here. So Boron has five electrons. So let's write the electron configuration for Boron. Well so far we have one s two, two s two but that only takes care of four electrons, we need five. So where does the fifth one go? The fifth one goes into the next available orbital here so we're going to put the electron in, the fifth electron for Boron goes into a two p orbital. So we would write two p one indicating that the fifth electron for Boron went into a p orbital in the second energy level. So one s two, two s two, two p one, is the full electron configuration for Boron. All right, so let's do Carbon. So next we have Carbon. Let's use blue for Carbon here. So Carbon has a total of six electrons. We have one more to think about. So we know it's going to go into a p orbital, a p orbital in the second energy level. The question is which one of these p orbitals do we put this next electron for Carbon? So we have to think about something called Hund's rule. I'm never going to pronounce the German properly so Hund's rule tells us that our goal is to minimize electron repulsion here. So let's think about... Let me just go ahead and draw the p orbitals down here. So we already have one electron right here. Well it doesn't make any sense to put an electron into the same orbital here because that puts the electrons really close together in space. So if you're thinking about a p orbital, remember a p orbital is shaped like a dumbbell so I'm just saying we have a p orbital on this axis let's say. So we all ready have, let me use, I'll just use blue here. So we already have one electron in there, it doesn't make any sense to add an electron to that exact same p orbital. That puts them really close together in space and electrons repel. So that doesn't make any sense so we need to take that electron out of there. That doesn't make any sense. We need to add an electron to another p orbital. So we'll take this electron out of there like that. So remember, there are other p orbitals on these other axis here. So here's another p orbital and then here's another p orbital. So we need to add an electron to another one of these. Whichever one, it doesn't really matter. Let's just say we're adding one here. So we're adding an electron to a different p orbital, whichever one it is, px, py, or pz. And it turns out that keeping the spins parallel helps to minimize the electron repulsion for pretty complicated reasons and I think they might still be doing research on this and so we put the electron in a different orbital and we keep the spins parallel which helps to lower energy for the atom here. And so that's where we're going to put Carbon's sixth electron. We're going to put it into a different p orbital and we're going to keep the spins parallel like that. So we can go ahead and write Carbon's electron configuration. Just read off everything we have on our orbital notation here. So we have one s two, we have two s two and we have two p two here. So two electrons in the p orbitals in the second energy level for Carbon. Next we have Nitrogen so let's use green here for Nitrogen. So Nitrogen has seven electrons, so one more electron to think about. Let's put Nitrogen right here. So we have so far, one s two, two s two. Now let's think about Nitrogen. So we need to add one more electron to our diagram. Once again, we're going to follow Hund's rule here. We're not going to add the electron to one of the already occupied orbitals, we're going to add this electron for Nitrogen to an unoccupied orbital and we're going to keep the spins parallel to keep everything lower in energy. And so we have three electrons for Nitrogen, in the two p orbitals. So we write two p three. So we have one, two, three. So we have one s two, two s two, two p three, would be the full electronic configuration for Nitrogen. Let's move on to Oxygen. So let's pick, let's see here, what color should we pick for that? Let's use orange here for Oxygen. So we have eight, eight total electrons. So for Oxygen, let's see, let's put Oxygen right here. So so far we have, one s two, two s two. So how many more electrons do we need for Oxygen? Oxygen has a total of eight electrons, we all ready represented four, so we need to represent four more. Oxygen's eighth electron, now that all of our orbitals are occupied, we can start to pair our spins. So we put Oxygen's eighth electron in there. So we can start to pair up our spins. We have four electrons in the two p orbitals for Oxygen, so we write two p four. So one s two, two s two, two p four. Notice if you add these together, two and two and four then you get eight which is the total number of electrons that we had to represent for the electron configuration for Oxygen. Let's move on to Fluorine. So let's use a different green here for Fluorine. So let's say Fluorine right here. Nine total electrons. So once again, we're pairing up our spins. So we add Fluorine's ninth electron to there and we can go ahead and write it right here. So for Fluorine we would write one s two, two s two, and notice we have five electrons now. So two p five would be for Fluorine. And then finally, let's go ahead and do Neon. So Neon has ten electrons. So we have one more electron to account for. We have one more space right? The last electron for Neon would go into a two p orbital here. So for Neon we would write, one s two, two s two, two p six. Notice we have no more places to put electrons, in the first or the second energy levels. We are completely full. So the second shell is now full and if you wanted to add another electron, you would have to open up a new shell. You would have to go to the third energy level. And so you notice a pattern here emerging on the periodic table. So we said that Hydrogen's electronic configuration over here was one s one. Then we went over here to Helium was one s two. And then we moved on to the second energy level. So this was Lithium here ended in two s one. And Beryllium ended in two s two. And then we filled the s orbital and moved on to the p orbitals. And notice we have over here, Boron's last electron was two p one, Carbon's two p two, Nitrogen's two p three, Oxygen's two p four, Flourine's two p five and Neon is two p six. So notice we have these six boxes over here on the periodic table. Those represent our p orbitals. And then over here on the left we have these two boxes representing our s orbital. And so that's the idea. The s orbital, we have one of them, holds a maximum of two electrons. We have these two boxes on the periodic table. Over here on the right, we had these six boxes which is the maximum number of electrons we can put into the p orbitals because we have three p orbitals, each one can hold two. So noticing these patterns on the periodic table helps you when you are writing electron configurations. You can just sit down and look at the periodic table and write them out after you've had enough practice. So make sure to do all of these again and think about electron configurations. Where you're putting your electrons and think about how it relates to the structure of the periodic table.