 So let's think about the hybridization of the atomic orbital diagram of nitrogen, just the isolated nitrogen, and then hybridizing to the molecular form of nitrogen where it actually combined to those three hydrogens. Is everybody OK with what we're doing? OK, so the isolated nitrogen, if you recall, so when we're doing this, you remember the valence electrons are the only electrons that participate in bonding. So those are the ones that we're going to emphasize. So the valence electrons in nitrogen are all the two electrons. So in this case, we have 2s and 2p. Remember, three orbitals in the 2p. We look up at the periodic table at nitrogen. How many valence electrons does it have? Five, right? Five. So how do we put those in? Remember all our rules, off-bout, arms, poly exclusion. One, two, three, four, five. What do we know about ammonia? So ammonia makes three, our nitrogen and ammonia makes three bonds. In this situation, you see three half-dilled orbitals. Everybody sees that. So potentially, nitrogen could make three bonds as it stands. But what we know is that the bonds of nitrogen are not 120 degrees apart, right? So what we need to do is figure out, so what do we know about nitrogen and ammonia? We know that the structure of ammonia looks something like that, right? OK, remember the relative bond angle. So it's very similar to methane, which is, remember the bond angle of methane? 109.5, right? Do you remember the bond angle of ammonia? Anybody? 107.3, right? So very close to that 109.5. But you remember the electronic arrangement around that nitrogen atom, right? What was the electronic arrangement? Tetrahedral, right? So you remember tetrahedral, too. So tetrahedral gives you kind of that around 109.5 bond angle. So since all of those electron groups are equivalent, because of those bond angles being equivalent, we know that all these orbitals have to be equivalent. So that's where we get the fact that it's got to be hybridized, OK, that and the bond angles, OK? So anyways, or hybridized. So we're going to mix these guys up, or hybridized, or whatever you want to call it, OK? And we've got 1s and 3p orbitals that we're mixing. So we're going to get, so we're putting in four orbitals. So how many are we getting now? You guys remember? Put in four, we get out four. They're all of equal energy. And what are they called? SP3, right? SP3. So when we do that, we fill it up using our puns rule and polyexclusion. 1, 2, 3, 4, 5, OK? So if you look here, this is our lone pair electrons. Is everybody OK with that being the lone pair? And notice bonding electron, bonding electron, bonding electron. Is everybody OK with that kind of analysis a little bit? OK, again, we know this due to what Vesper theory tells us about the bond angles and the equivalency of those electron groups. Is everybody OK with thinking like that? So if we were to look at one of these NH bonds, right, what would N, what orbital would N be using to make that bond? What is that orbital called? SP3, right? SP3. So it's got its one electron. What is the H using for its orbital? 1s. 1s, SP3. How many electrons does hydrogen have? 1. So it can only use the 1s orbital, right? So the 1s orbital. That's the bonds that are the orbitals that are used to make that bond. Is everybody OK with that? I see some confusion. You OK with that? 1s for the hydrogen. It's the only orbital it's got, right? Any other questions? Or besides my question? OK, good.