Which statement is not correct with respect to transition metals? (a) The colour of the hydrated ions is due to transitions of electrons from different 'd'orbitals of the same principal quantum number (b) Most transition metals should, on thermodynamic considerations, liberate hydrogen from acids (c) The complex formation of transition metal ions is often accompanied by a change of colour and sometimes by a change in the intensity of the colour (d) The compounds of the elements in low oxidation states are generally ionic

Transition metals are intriguing elements that occupy the center of the periodic table. One of their most characteristic features is their ability to display vibrant colors. This colorful display is attributed to electron transitions in 'd' orbitals.

These metals have unfilled 'd' orbitals that can accept and donate electrons. When light shines on a transition metal compound, some wavelengths are absorbed to promote electrons from a lower-energy 'd' orbital to a higher-energy one. This is fundamentally an electron transition within the same principal quantum number, as correctly stated in the given exercise. The energy difference between these orbitals corresponds to the energy of the absorbed light, and the remaining transmitted light gives the compound its color. This phenomenon is elegantly described by the crystal field theory, which explains the color changes as a result of the splitting of 'd' orbitals in the presence of a field created by surrounding ligands.

Understanding these transitions is vital not only for explaining the vibrant colors but also for the study of various chemical reactions and technologies such as catalysts and sensors.

The formation of transition metal complexes is another fascinating aspect of these elements. In a complex, the central metal atom is bonded to one or more ligands, which are molecules or ions offering at least one pair of electrons for bonding. This interaction is pivotal in determining the properties of the complex, including its color.

When transition metals form complexes, the 'd' orbitals split into different energy levels, leading to the absorption of specific wavelengths of light and often a change in the intensity of the colour, as observed in statement C of our exercise. These changes can be tunable depending on the ligands involved and the geometry of the complex. This tunability finds applications in industrial processes, where complexes are used as catalysts, as well as in everyday products like paints and pigments.

By altering the ligands or the metal center, chemists can design complexes with desired properties, showcasing the adaptability and utility of transition metal chemistry.

Lastly, the varying oxidation states of transition metals play a pivotal role in their chemistry. These metals can adopt multiple positive oxidation states, which influences the types of compounds they form. While high oxidation states can lead to ionic compounds due to the high charge on the metal drawing electrons towards it, low oxidation states can indeed lead toward covalent bonding, as electrons are more likely to be shared than transferred.

Thus, the notion that compounds in low oxidation states are generally ionic, as proposed in statement D of our exercise, is incorrect. Covalency increases in lower oxidation states due to lesser electrostatic attraction between the metal ion and the electron pairs belonging to the ligands or other atoms. Therefore, we can see that the chemistry of transition metals is rich and varied, driven by their ability to participate in both ionic and covalent bonding, which in turn affects their reactivity, stability, and other chemical behaviors.