 So we are going to explore the group 13 elements or the boron family in this video. We will look at the various trends in the properties especially atomic and physical properties of the group 13 elements. So where do we begin? We will begin by looking at the electronic configurations of the elements here. So as you already know the outer electronic configuration of group 13 elements would be NH2, NP1 that means there are a total of 3 electrons in the valence shell and as we go down the group we can see the presence of completely filled D and F orbitals in the heavier members of the group correct? Now what can you comment about the nature of these elements? Well that you can see from the diagram here. The green elements are non-metals while the blue ones are metals and the orange ones are metalloids. So that means except boron every other member of the group 13 is a metal. Now there are some controversies around boron. You see many people tend to look at boron as a metalloid because it does exhibit some properties which are similar to metals. For example if you take a sample of boron you will see that just like metals it has a very shiny appearance. Now just like metals boron also has a very strong crystalline structure because of which it has a high melting point. Nevertheless we will study boron as a non-metal. This is because the boron chemistry is predominantly non-metallic in nature. For example boron does not form cations in aqua solutions. And not just that because of its high charge to size ratio boron compounds are predominantly covalent in nature. Anyway we will get back to the boron chemistry later but needless to say that we will study boron as a non-metal in this course. Alright now before we go ahead and discuss the various properties of group 13 elements we have to remember that all the variation that you see in the trends of these properties is because of the intervening D and F orbiters. To be more precise it is because of the poor shielding effect of the D and F orbiters. So let's take a quick recap of what they do. So here I have the example of gallium as we know gallium is the first element of the group 13 which has completely filled 3D electrons. Now we said that D and F electrons have poor shielding effect correct? So what happens is because of this it cannot screen or shield the outer electrons effectively from the nucleus. So basically D and F electrons act like a thinly veiled curtain. You know if you are preparing for an exam and decide that you don't want to get distracted with what's happening outside. But if you cover the window with a thin curtain does that serve your purpose? No right you can still see everything that's happening and you're more likely to get distracted. Now this is not the exact analogy but what I'm trying to tell here is that the valence electrons cannot escape the pull of the nucleus. Because of this these valence electrons get pulled closer towards the nucleus especially the NS valence electrons and all of these properties that you see or the variation in these properties arises because of this particular effect. And we're going to look at all of these properties and their variation in this video. So let's begin with the atomic radius. We know that as we go down the group the atomic radius should increase right? Because of the addition of newer shells the outer electrons tend to go farther and farther away from the nuclear pull. But is that what's happening here? Here we can see that from aluminum to gallium the atomic radius actually decreases. Now why is that? Well for starters you can see that the major difference between aluminum and gallium is a completely filled 3D electrons. And what do we know about D electrons? We know that D electrons have very poor shielding effect. So that means it cannot screen the outer electrons from the nuclear pull and this results in a decrease in the atomic radius. Now something quite similar occurs with indium and thallium as well. As you can see from indium to thallium we should expect a much larger increase in the atomic radius. But here the increase is very very nominal almost similar to be honest. Now this is again due to the poor shielding effect of the F as well as D electrons. Because of this the electrons get pulled closer to the nucleus and this completely offsets the expected increase in the atomic radius. Now if you extend the effect of atomic radius or the variation of the atomic radius you will see something quite similar happening with density as well. And we know that density is nothing but atomic mass by volume. And what is the general trend in the density as we go down the particular group? It should increase because the atomic mass increases with more number of protons and neutrons and atomic volume or the atomic size also increases. So in general density increases down the group. But if you look at a table here you will see a much sharper increase in density from aluminium to gallium and similarly from indium to thallium. Now this is because while the atomic mass increases as we go down the group the atomic size or the volume decreases from aluminium to gallium. As we just saw due to the poor shielding effect of the D electrons. And that means the volume decreases. So what is the effect on density? As the volume decreases density increases. Now something quite similar happens between indium and thallium as well. The atomic size does not increase proportionally to the atomic mass. And here again we notice a sharp increase in the density from indium to thallium. Alright moving ahead let's look at another important property which is the ionization enthalpy. Ionization energy in general decreases down the group. This is because the valence electrons are far away from the nuclear attraction. That means it is very easy to knock these electrons off. But is that what we are seeing here? Here you can see that there is a distinct increase in the ionization enthalpy from aluminium to gallium. It further increases from indium to thallium. By now I am sure that you know why this happens. Yes again due to the poor shielding effect of the intervening D and F orbitals. So because of the poor shielding effect of the D and F electrons. The outer electrons experience higher nuclear attraction or higher effective nuclear charge. That means you need to provide a lot more energy to knock off these electrons. And this is why in contrast to what we expect ionization enthalpy increases from aluminium to gallium and indium to thallium. Now if I extend this trend on the ionization enthalpy a little further. We can kind of conclude that it becomes difficult to lose the electrons as we go down the group. And this is exactly what happens even with the reducing property. Reducing property depends on how easily an element can lose electrons. And in general when we talk about reducing property it increases down the group. That means it becomes easier to lose an electron as we go down the group. Right? But in the case of group 13 elements this reducing property actually decreases. If the property decreases that means it becomes more difficult to lose the electrons. Right? And this again happens due to the D and F electrons. You see because of their poor shielding effect they actually help the nucleus hold on to their balance electrons strongly. And this is why the heavier members become reluctant to lose their electrons. In other words they show decreased reducing property. Okay now before wrapping up the video let me discuss one more property which is the electronegativity. We know that electronegativity is a tendency of an atom to attract a shared pair of electrons towards itself. Right? And in general we know that electronegativity decreases down the group. However in group 13 you will see that from aluminium instead of decreasing electronegativity values we actually find them increasing although nominally. Now this again can be attributed to the poor shielding effect of the D and F electrons. You see we just saw that the nuclear charge increases with increasing the atomic number. Right? Because we have more number of protons and neutrons in the nucleus. But because the D and F electrons do a thinly veiled job of shielding the nucleus is still capable of attracting an external pair of electrons towards itself. Now this is what we again see here. Although you can still see that the increase is not very drastic but it is still important to notice that electronegativity in fact increases in group 13 elements. As you can see almost all of the variations in these properties occurs as we move from aluminium to gallium and from indium to thallium. Now all of this can be attributed to the poor shielding effect of the D and F orbitals. In the next video we will look at the chemical properties and the trends in the reactivity of the group 13 elements. Do you think that the poor screening effect of the D and F electrons will also affect the chemical reactivity? Let's find out.