 This video is going to be an introduction to the group 16 elements also called the chalcogens. So what constitutes the chalcogens? Well these are oxygen, sulfur, selenium, tellurium and polonium and from their electronic configurations we can see that they all have NH2, NP4, outer electronic configuration. That means they have a total of six valence electrons. So clearly as we are nearing the end of the periodic table, we move closer towards the noble gas configuration right or the octet configuration. That means these elements just need two more electrons to satisfy their octet and this is why they also tend to be more reactive. All right so let's now quickly look at the general trends in the properties of the group 16 elements. As we have seen in all the previous cases the atomic radius or the covalent radius of these elements also increase as we go down the group. I mean oxygen has an exceptionally small size. Just like that their ionization enthalpy values and electronegativity values also decrease as we go down the group with increase in atomic size. So that's what you can see here as well. But in the case of ionization energy if you compare these elements with their immediate predecessors which is the group 15 elements you can see that the ionization enthalpies of group 16 elements are lower than that of the group 15 elements especially for the lighter elements. And this can be attributed to the exceptional stability of the half filled electronic configuration of the group 15 elements right. This is something that we've already discussed in the previous video. So clearly you can see that the ionization enthalpy values of the group 15 elements are higher than either of the elements on their neighboring sites that is both group 14 and group 16. Now talking about electronegativity values as we've already discussed these values also decrease as we go down the group. But the important thing to note here is that oxygen is one of the most electronegative elements. In fact it is the second most electronegative element after fluorine. It can easily accept two electrons to attain stability. Okay it's now time to discuss another very important atomic property which is the electron gain enthalpy. Electron gain enthalpy is nothing but the amount of energy that is released when an electron is added to an isolated atom. And in general the electron gain enthalpy values become more negative that is more amount of energy is released when the electron is added to a neutral atom as we go across the periodic table. And this is because the effective nuclear charge increases as we go from left to right that is it is easier to add an electron to a smaller atom since the added electron on an average would be closer to the nucleus right. Like if you're adding an electron to the second shell then it is closer to the nucleus than say if you're adding one to a larger shell like n is equal to 3. Here the added electron would be farther from the nuclear attraction correct. And by extension the electron gain enthalpy values would become less negative as we go down the group because here the atomic size increases and as I said before the added electron would be farther away from the nuclear attraction that means less energy would be released and the value becomes less negative. However in the case of group 16 elements there is an exception. Contrary to our expectation the electron gain enthalpy value is less negative in oxygen when you compare it with that of sulfur. So this means that when you're adding an electron to oxygen it is releasing lesser energy than when you add one electron to sulfur. Now this is also strange because electronegativity value of oxygen is much greater than that of sulfur so this is indeed an exception. But we do have a pretty valid reason for why this is observed and that has got to do with the atomic size of oxygen. You see because it is small size of oxygen atom when we add an extra electron to the neutral oxygen atom it experiences a lot of dripple shun from the other valence electrons. Because the atomic size is really small right all the electrons are very close to each other. But if you compare it with the larger sulfur atom you can see that the valence electrons here are in the third shell that is the added electron is now present in a larger orbital or a larger region of space. So that means an incoming electron is more comfortable or more welcome to go to the third shell or to sulfur than it is to oxygen. Now something very similar happens in the case of the halogens as well. That is the group 17 elements about which we'll study in the subsequent videos. You see there again the electron gain enthalpy value of fluorine is much less negative as compared to chlorine and this is again because of the small size of fluorine due to which the incoming electron experiences a higher electron-electron repulsion. Now lastly let's come to the oxidation states. Group 16 elements exhibit a varied number of oxidation states minus 2 being the most common one. Obvious right it is just two electrons away from the noble gas configuration. However as we've discussed in the case of the nitrogen family the stability of the negative oxidation states in this case minus 2 oxidation state decreases as we go down the group. You see if you look at the physical nature of these elements we know that oxygen and sulfur are non-metals selenium and tellurium are metalloids whereas polonium is a metal. In other words we can see that the metallic character increases as we go down the group that is the heavier larger atoms are much more likely to lose electrons and gain electrons and this is why polonium almost never shows minus 2 oxidation state. It shows only positive oxidation states that corroborate with what we just discussed that is that ability or tendency to lose electrons in the heavier elements or as we go down the group. Now as these elements have six valence electrons in the outer shell plus six is also that group oxidation state and as you can see from here except oxygen and polonium all the other elements show plus six oxidation state in some compounds or the other. Now let me ask you something why do you think oxygen does not show plus six oxidation state? Obviously because you're asking a highly electronegative element to give up electrons right and not just one or two literally six of its valence electrons that's not in a reasonable expectation so obviously oxygen will not show a plus six oxidation state but why does polonium not show plus six oxidation state? It should be able to give up its electrons easily right? Well theoretically yes and practically no. As we've discussed numerous times by now in the heavier elements of the P block the S electrons generally prefer to remain paired. Yes the famous inert perifet kind of extends to the group 16 as well. This is why the stability of plus six decreases down the group and plus four becomes more stable for polonium. So as you can see from here the more we learn about elements the more we see some discrepancies some anomalies and some exceptions but there's really nothing to be afraid about of to be frustrated about because if you understand the inherent nature of an element all of these exceptions would seem very logical. For instance why does oxygen not show plus six oxidation state? It's logical right because the second most electronegative element it cannot lose so many electrons so with that in mind let's look at some anomalous behavior of oxygen in the next video.