 In this video, we are going to study some of the properties of the transition elements and we will do that by focusing mostly on the first series which is the 3D series. So let's begin with the physical properties. We know that transition elements have metallic properties. They have height and style strength. They conduct heat and electricity. They have metallic luster and are also malleable and ductile which are the characteristic features of metals. With a couple of exceptions like zinc, cadmium, mercury and manganese, all of them have metallic lattice structures. Transition metals are also quite hard and have high melting and boiling points. And this high melting point is due to the participation of a large number of electrons. You see, we usually expect the outer valence electrons of an atom to be involved in bonding. And that's correct. But in the case of transition elements, what is the outer electronic configuration? It is n minus 1d, 1 to 10 and n is 1 to 2. Now the energy difference between the ns electrons and the n minus 1d electrons is quite low. And that is why electrons from the n minus 1d sub shell also participate in metallic bonding along with the ns electrons. Since these elements form metallic bonding, greater the number of unped electrons, stronger is their interatomic interaction and stronger is the metallic bonding. As a result, we need to provide more energy to break their lattice structure. And this is why transition metals have high melting points in general. The melting point reaches a maximum at the middle of each series. Now the maximum at the middle of these series indicate that one unped electron per d orbital is especially favorable for strong interatomic interaction. Thus, greater the number of valence electrons, better or stronger is the metallic bonding. Now the only exception in the 3d series is that of manganese. Despite having a d5 configuration, manganese has an unexpectedly low melting point. Now the other exceptions in the d block are zinc, cadmium and mercury. All of these d block elements have much lower melting point than the others. And this atypical behavior is because they have completely filled d sub shell. They do not have any unped d electrons that would increase or maximize their interatomic interaction. Now as we go from 3d series to 4d series to 5d series, the melting point increases in the subsequent series. That is melting point of 4d is greater than that of 3d and 5d is greater than that of 4d. This is because down the series we have more number of electrons which means stronger interatomic interaction. And overall if you see the melting point trend would be exactly same as that we observed in the 3d series. That is it increases to a maximum till d5 and then decreases as we go towards the end of the series. Now another consequence of more number of d electrons is the greater enthalpy of atomization. This is because as we mentioned in the case of melting point more number of d electrons or more number of unped electrons or valence electrons results in stronger metallic bonding. As a result the enthalpy of atomization increases. Just as we observed in the case of melting point the second and third series transition metals have greater enthalpy of atomization than the first series. That is 5d series has the highest values of atomization whereas the 3d series has the lowest values of atomization enthalpies. Let's now talk about the atomic size. In general atomic radius decreases as we move across the period with increase in nuclear charge. And in the case of d-block elements the radiate decrease from left to right. However as we reach towards the end of the series the size actually increases. Now in the d-block elements there are two interplaying factors. One is the shielding effect of the d orbitals which as we know is not very high because d orbitals do have poor shielding effect. And second thing is the nuclear attraction or the nuclear charge. As the d electrons have poor shielding effect the outer electrons experience greater nuclear attraction. And as a result the ionic radius decreases. Now within the series you will see slight variation. For example the ionic radius decreases steadily from scandium to manganese. But cobalt, ion and nickel have very similar ionic radii and the radius increases towards the end of the series. However this variation within the series is quite small and can be broadly attributed to the slight increase in the shielding effect with the number of d electrons. With more number of d electrons the shielding effect is expected to be slightly better. And as a result copper and zinc have a slight increase in their ionic radius. Now on comparing the 3d series with 4d and 5d series we once again notice an increase in radius from 3d to 4d series. That is 4d series elements have greater atomic radii as compared to the 3d series elements. However no change in radius is observed between 4d and 5d series. So what do you think is happening here? So as we go towards the heavier elements and if we look at the electronic configurations of the 5d series we can see that the 4f orbitals fill before the 5d orbitals. Now this filling of the 4f orbitals before the 5d orbitals results in a decrease in the atomic radius. And this is called lanthanide contraction. As a result of the lanthanide contraction the 2nd and the 3rd series that is the 4d and the 5d series have similar radii and also exhibit similar physical and chemical properties. The lanthanide contraction is caused by the imperfect shielding of the 4f electrons. So as we go down the group we can see that the atomic radius more or less remains the same but the atomic mass increases right? We know that the mass increases as we go down the group. So as we go down the group the atomic size decreases that is the size would be lesser than what is expected but the atomic mass increases. So what do we have here? Well mass divided by size or volume is nothing but density. So the decrease in atomic size or the metallic radius coupled with an increasing atomic mass in an increase in the density of these elements.