 Compared to the compounds of SNP block elements that are almost always white, many of the ionic and covalent compounds of the transition metals are coloured in nature. And in this video, we will try to touch upon the reasons why compounds of transition elements appear coloured. Now, we all know the general phenomenon behind the appearance of colour, right? We know that when light passes through an object, some wavelengths of the light get absorbed. And if this absorption occurs in the visible region, then the colour that is complementary to the colour that got absorbed gets transmitted. For example, plants appear green because they mostly absorbed red light and some blue light while the green light gets reflected. Now, this absorption of light in the visible or the UV region occurs due to the changes in the electronic energy. That is, the energy that is associated with the promotion of electron from one energy level to the other. Now, usually the energy jump is so large that the absorption lies almost always in the UV region. But in some cases, it is possible to obtain small jumps in the electronic energy, which usually appear as absorption in the visible region. Now, things take an interesting turn when we talk about ions with incomplete D or F orbitals. You see, an isolated gaseous ion of a transition metal that is free from any external influence has 5 degenerate D orbitals. That is, all the orbitals here are of same energy. But in reality, these ions are not really isolated. They are usually surrounded by solvent molecules. For example, in the case of copper sulphate, hydrated copper sulphate where copper ions are coordinated or associated with 5 water molecules and this ion and this complex is actually blue in colour. So here, as you can see, the ions can be coordinated with solvent molecules. It can also be coordinated by other ligands in complexes like FECN64- or even other ions in crystal lattices like in the case of MNO4- So as you can see, the ions are not actually isolated but surrounded by different types of ions depending on whether it is in a crystal lattice or in a solution or in a complex. And these surrounding molecules will have some kind of effect on the transition metal ion, right? Of course, the surrounding groups affect the degeneracy of the D orbitals and as a result, the D orbitals split into two groups. Two groups of D orbitals of two different energy levels. Because of this, we can easily promote an electron from one D level to the other D level of higher energy. As the energy difference between the two D levels is not very high, the light is absorbed in the visible region. Now the colour that is transmitted depends on how big the energy difference is between the two D levels. And this further depends on the nature of the ligands, whether it is a strong ligand or a weak ligand, number of such ligands and the type of the complex that is formed. Now we cannot discuss more about these factors in this particular video. But to give you an overview, if we have a strong ligand, then the D energy levels would split more and therefore the delta E would be larger. Whereas if we have a weak ligand, the delta E would be smaller. And depending on this energy difference, we get the corresponding colour. For example, Ni H2O6 is green in colour, whereas Ni NO2 6 would be brown red in colour. So even though we have the same metal ion which is nickel, depending on the type or the strength of the ligands it is attached to, the nickel complexes can get different colours. Now what makes a ligand strong or weak is something that we cannot cover in this particular session, but will be covered in great detail in your next unit on coordination chemistry. Okay, for now let's understand that ligands can be classified into strong and weak. And depending on the strength of the ligands, the number of ligands and the shape or the type of the complex that is formed, colour shown by a particular ion can vary. However, we do have some white coloured compounds of transition elements too. For example, zinc sulphate and titanium oxide are white in colour. This is because in these compounds it is not possible to excite the electrons within the D level. In zinc sulphate, zinc is in plus 2 oxidation state, which means it has completely filled 3D orbitals. It does not have any empty D orbitals. On the other hand, titanium 4 plus gives you noble gas configuration. Here the number of electrons in D orbitals is zero. So in the case of zinc, there are no empty D orbitals available to promote the electrons. And in the case of titanium, there are no D electrons at all. So because of this, it is not possible to have a DD transition in these compounds. And therefore the complexes of these ions do not exhibit any colour. Now as I mentioned at the beginning of the video, S and P block elements are also colourless. They do not have any partially filled D orbitals. So obviously we cannot have any DD transitions here. On top of that, the energy required to promote an S or a P electron is much higher and as a result, the absorption occurs in the UV region. This is why the compounds of S and P block elements appear mostly white in colour.