 Okay, so let's now talk about the carbon family. So in this video we will basically discuss quickly the trends of some of their important atomic properties like atomic radius, ionization enthalpy, electronegativity and so on. Now what do we know about these properties in general? Well, we know that atomic radius in general increases down the group, right? This is because with successive periods newer and newer shells get added and that means the valence electrons tend to move farther and farther away from the nucleus, correct? So because of this atomic radius, which is nothing but the distance from the nucleus to the valence shell also increases. Now based on this, what can you comment about the nuclear charge or nuclear attraction that this valence electron experiences? Well, we can easily say that if the valence electrons are in the two-wit shell, obviously it would experience greater nuclear charge, right? Because it is very close to the nucleus. But as you go away from the nucleus and if we have a valence electron like let's say in the sixth shell or the fifth shell, then this will experience much lesser nuclear attraction, right? Because obviously it's much farther away from the range of the nucleus. So from here we can see that you need to provide very less energy to knock off these electrons as compared to a 2s or a 2p electron. So from here we can see that ionization energy in general decreases down the group because as electrons move farther away it becomes easier to knock them off from the nuclear attraction. Now this can be extended to the electronegativity as well. Electronegativity is nothing but the tendency to attract a shared pair of electrons. And now here again as the atomic size increases it becomes difficult for the nucleus to attract any shared pair of electrons as well. So this means electronegativity also decreases down the group in general. So folks, what we're talking about is the general trends or these are our general expectations of how these properties should be. But in group 14 we'll observe that the trends are not exactly like this. There are some variations. So let's now see how these properties vary within the carbon family and why do they do so. Okay so in order to do that we first need to know who the members are of this particular family. So these are carbon, silicon, germanium, tin and lead. And from this image we can see that carbon is a non-metal, silicon and germanium are metalloids while tin and lead are completely metallic in nature. So let's quickly also look at the electronic configurations and what do we notice here? By now we know that almost everything that we need to know about a particular element can be found or can be seen in their electronic configurations. So what is the common thing that you observe in these elements? We'll see that they all have the same number of valence electrons right? Two electrons in the s orbital and two electrons in the p orbital. So that is they have four valence electrons in their outer shell. Now as we go down the group we also notice the presence of completely filled d and f orbitals. And we know that the d and f orbitals tweak the chemistry of the elements significantly. It is the poor shielding effect of the d and f orbitals that creates all the anomalies in their general behavior. For example, atomic radius increases substantially from carbon to silicon. But from silicon to lead you will see that this increases not very big. It's very nominal. I mean it's not as large as we would normally expect. This is because the poor shielding effect of the d and f orbitals causes the valence electrons to experience more nuclear attraction. And because of this they get pulled closer towards the nucleus. So what happens is the atomic radius does not increase as much as we would normally expect. Just like that if you look at the values of the ionization enthalpy we will see that the decrease in ionization enthalpy from carbon to silicon is quite sharp. But from silicon to germanium it's not a big decrease and again you will see that from tin to lead there's actually an increase in the ionization enthalpy. This once again is due to the poor shielding effect of the d and f orbitals. So because of which the outer valence electrons experience much stronger or effective nuclear charge and that means it becomes more difficult to knock off these electrons. And from tin to lead the ionization energy actually increases. This is because you see lead is a pretty big atom. It's the heaviest member of the family right. It has a huge nucleus and that means its nuclear charge is also very high. But the intervening 4f and 5d orbitals do a very poor job of shielding. So that means the nuclear charge can still attract the outer valence shell electron and can pull it closer towards a nucleus. And that means since this experiences higher effective nuclear charge you need to provide more energy and that is why your ionization enthalpy consequently increases. And what about electronegativity? As we would normally expect electronegativity decreases from carbon to silicon. But unlike our expectation it does not decrease. It remains constant from silicon to tin and in fact increases in the case of lead. So this once again can be attributed to the poor shielding effect of the d and f orbitals. Because the d and f orbitals do such a poor job of shielding the nuclear force or the nuclear charge can still attract an external shed pair of electron towards itself. And this is why although electronegativity decreases down the group it is just a nominal decrease and almost remains constant for the heavier members.