Non-metal chemistry. So as we move across the periodic table, of course, we're increasing the number of protons and we're increasing the number of electrons. What we are not doing is we are not increasing the number of orbitals. So as you get to carbon, you have as many electrons IE4 as you have orbitals and you form these electron precise compounds. By the time you get to nitrogen in group 15, you've now only got four orbitals but you've got five electrons. And that has significant consequences. With forming compounds that are no longer electron precise, they have an excess of electrons. And there are two main consequences of this chemistry. We move from nitrogen down to phosphorus and arsenic, then you'll see with nitrogen because nitrogen is small, because those p orbitals can overlap very efficiently. Then in the case of nitrogen, what we see is a propensity towards multiple bonds. Multiple bonds is one way of dealing with your extra electrons. The other way, of course, is that you have lots of lone pairs. Compounds with lone pairs are going to function as Lewis bases. So these electron excess compounds with more electrons than orbitals have lone pairs and where you've got lone pairs, you've got compounds that will function as Lewis bases. And that follows simply from just counting the number of electrons that are present in these systems. Okay, so the consequences of electron sets are either lone pairs and or multiple bonds. You've got always going to have lone pairs, whether or not you get multiple bonds, is going to depend on the size of your element. Right, so, well, do we form multiple bonds in the case of nitrogen? Well, this kind of chemistry is clearly going to be very familiar to you from carbon. So in the case of nitrogen, what if you've got, you've got one more electron than carbon. And in many ways, the N2 molecule is very similar to an alkyne molecule. In an alkyne molecule, you have a triple bond and you have CH bonds or CR bonds on the ends. In an N2 molecule, you're going to have a triple bond and you've got lone pairs on the end of the nitrogen molecule. How do you form these lone pairs? Well, if we consider each nitrogen's atom to be sp hybridized, then you have a linear sp hybrid pointing along the interatomic axis and that's going to form a sigma bond. We then have got two p orbitals remaining. And remember, those two p orbitals are going to be orthogonal to the interatomic axis and they're going to be orthogonal, meaning at right angles to one another. So nitrogen chemistry is quite interesting in the sense that what you get in nitrogen chemistry is every conceivable oxidation state. If we look at nitrogen, we can basically take nitrogen with a less electronegative element. Hydrogen is less electronegative than nitrogen. So formally, if we form a nitrogen, hydrogen bond, when we break it, the pair of electrons goes to the nitrogen so the nitrogen has a negative oxidation state. If we have three of those hydrogens, then the oxidation state of nitrogen in the ammonia molecule is minus three. Now, if we go to the opposite extent here, we take a nitrate anion, NO3 minus. Obviously, the charge on that species is one minus. Oxygen, what's the normal oxidation state of oxygen? So we have the oxidation state of nitrogen and three times minus two is equal to minus one. So by a bit of trivial algebra, we can see that the oxidation state of nitrogen in the nitrate anion must be plus five. And we'll see we get essentially everything in between. The obvious place to start when you surveying an element is with the elements in its elemental form. And there is only one alotrop of nitrogen. The N2 triply bonded molecule. It's an inert, relatively inert gas. Why is it relatively inert? Well, it's non-polar. If you're going to predict the reactivity of something, you have to look at its polarity. So nitrogen N2 is non-polar. Obviously, it's non-polar because the two nitrogen atoms have the same electronegativity. So N2 is a non-polar molecule. And it's triply bonded. Any reactivity of N2 requires us to break a triple bond, a non-polar triple bond, with a bond dissociation energy of plus 944 kilojoules per mole. That nitrogen, nitrogen bond, is the second strongest bond in chemistry.