 In this video, we will look at the various trends in the properties of the group 15 elements. So how do we begin? Obviously, we will begin by looking at the electronic configurations of the various elements that are present in the group 15, right? Because by now we know that almost everything that we need to know about an element can be figured out from the electronic configurations. And what do we see here? We can see that all of these elements have the same number of valence electrons that is 2 electrons in the s orbital and 3 electrons in the p orbital. So this gives us a general outer electronic configuration of Ns2, Np3, correct? But wait a minute. We can see that here we have a half-filled electronic configuration. And we know that half-filled configurations are extra stable, right? So to quickly take a recap, half-filled electronic configurations are extra stable due to two factors. One is symmetry and the other one is exchange energy. So as you can see here, we have one electron in each of the px, py and pz orbitals. Now having equal distribution of electrons in these orbitals kind of decreases the overall energy of the system. And we also know that by the rule of nature, symmetry almost always gives extra stability. Now the second factor that contributes to this stability is the exchange energy. When the orbitals are half-filled, they all have the same spin in degenerate orbitals. What do we mean by degenerate orbitals? Orbitals that have the same energy. So here you can see that although the three p orbitals are oriented or they orient themselves along the three different axes, they all have the same energy. That is why px, py, pz orbitals are called degenerate orbitals. Now this also means that these electrons can exchange with each other, right? But remember, only two electrons of the same spin can exchange their positions and in the process they release energy. And this is called the exchange energy. So if we have more of such exchangers, greater will be the energy released and more will be the stability, right? And this is why half-filled and fully filled configurations are extra stable. An important consequence of this extra stability is the increased ionization enthalpy of the group 15 elements. Now as you can see from here, the group 15 elements have higher ionization enthalpy as compared to the group 14 elements. And this higher ionization enthalpy is due to the extra stability offered by the half-filled electronic configuration. So now that we have covered electronic configurations part, let's go ahead and look at the different properties or the trends in the properties of the group 15 elements, alright? So let's look at the first property which is the covalent radius. As we would normally expect, the atomic radius increases considerably from nitrogen to phosphorus. But only a small increase is observed from arsenic to bismuth. And as we all know by now, this effect can be attributed to the poor shielding effect of the intervening D and F orbitals in the heavier elements. Similarly, if you look at the electronegativity and the ionization enthalpy values, you will see that both of these values decrease as we go down the group. But again, the decrease is not very large in the heavier members. So that's mostly about the atomic properties like atomic radius, electronegativity, ionization enthalpy and so on. Now what can we comment about the physical nature of these elements? Well, we know that nitrogen is the only gas whereas all of the other elements, phosphorus, arsenic, antimony and bismuth, all of them are solids. And not just that, while nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids and bismuth is actually purely metallic in nature. So in a way, we can also say that as we go down the group, the metallic nature or the metallic character increases, correct? Now metallic character or metallic nature is synonymous to the ability of elements to lose their electrons. So more metallic an element is, easier will it be able to lose the outer electrons. Now this is because as we go down the group with increased atomic size, the valence electrons are farther and farther away from the nucleus, right? So that means they experience lesser nuclear attraction and therefore it becomes easier to knock off these electrons. Let's now look at one more property before wrapping up this video, which is the oxidation state. So we have 5 valence electrons in group 15 and the most common oxidation states that they show are plus 3, plus 5 and minus 3. In fact from group 15 onwards, you will start seeing many elements in negative oxidation states because you see group 15 is quite close to the end of the periodic table, right? I mean you are really very close to the noble gas configuration. So as you go towards the end of the period to attain a stable octet configuration, all you need is to accept or gain electrons, right? And when you gain or accept electrons, your oxidation state becomes negative. And this is why negative oxidation states become very common in group 15, 16 and 17. Anyway, coming back to our oxidation states, we know that the stability of the higher oxidation state or the group oxidation state which is plus 5 decreases as we go down the group. Whereas the stability of the lower oxidation state which is plus 3 increases as we go down the group. Now this preferential stability of plus 3 oxidation state over plus 5 in the heavier members as due to something that we already know which is the inert pair effect, right? Now we won't go into the details of the inert pair effect because we have seen it a number of times in our previous videos. So if you have forgotten by any chance, then please go back and watch this video, okay? So we discussed the case of plus 3 and plus 5 states. What about minus 3, our negative oxidation state? Well, it turns out that just like plus 5, the tendency to show minus 3 oxidation state also decreases as we go down the group. Minus 3 oxidation state means that the element is gaining 3 electrons, right? Now as we go down, what did we say about the metallic character? We saw that metallic character increases, right? That means the heavier elements are more likely to lose their valence electrons than gain additional electron because the nuclear attraction is not strong enough to hold or attract additional electrons. So in a way, we can say that the stability of the minus 3 oxidation state decreases because of the increase in the size of the atom as well as the increase in the metallic character or the tendency to lose electrons. In fact, the heaviest member of Bismuth almost never shows minus 3 oxidation state. Now if you look at nitrogen specifically, you will see that it is highly versatile. It exhibits a wide range of oxidation states as you can see in these compounds like plus 5 in nitric acid, plus 4 in NO2, plus 2, plus 1 and minus 3. Now this is not the only thing that makes nitrogen special. You see compared to the rest of the members, nitrogen actually enjoys a lot more attention. In the next video, we will look at some of the important anomalous properties or anomalous behavior of nitrogen.