Explain these facts: (a) Copper and iron have several oxidation states, whereas zinc exists in only one. (b) Copper and iron form colored ions, whereas zinc does not

Oxidation states play a pivotal role in understanding the chemistry of transition metals. They indicate the degree of oxidation of an atom in a compound, which is essentially a hypothetical charge that an atom would have if all its bonds were ionic. For transition metals like copper and iron, the oxidation states can vary due to the presence of d and s electrons that can participate in bonding.

For example, iron can exhibit oxidation states of +2 and +3. In the case of Iron(II), the atom loses its two 4s electrons, resulting in the +2 oxidation state. When it steps up to Iron(III), an additional 3d electron is removed, leading to the +3 state. Copper follows a similar pattern. With Copper(I), the oxidation state is +1 as it has lost one 4s electron and with Copper(II), the oxidation state is +2 due to the loss of both 4s electrons.

Zinc stands apart from these metals as it possesses a completely filled 3d sub-shell, allowing for only one stable oxidation state, which is +2. When zinc achieves this state through the loss of two 4s electrons, the energy required to remove any more electrons from the completely filled 3d sub-shell is significantly high, effectively preventing further oxidation. This characteristic of zinc explains its constancy in oxidation state.

The fascinating colors of transition metal ions are not just visually pleasing but also scientifically profound. These colored ions occur due to their unique ability to absorb and emit certain wavelengths of light, a phenomenon closely tied to their electron configurations. Transition metals like copper and iron can form various colored compounds, while zinc compounds typically lack color.

This difference comes down to the arrangement of the d-electrons in these metals. Copper and iron have partially filled d sub-shells that allow for the absorption of light, promoting electrons from lower to higher energy d orbitals—what is known as 'd-d transitions.' When light passes through a solution containing these metal ions, certain wavelengths are absorbed to facilitate these electron jumps, and the color observed is the complement of the absorbed light.

Zinc, however, has a full d sub-shell, which does not allow for this type of electron excitation, and therefore, no color is observed. This is because the energy required to promote an electron in a filled d sub-shell is significantly higher and usually outside the visible spectrum, resulting in colorless zinc ions.

d-d electron transitions are at the heart of the color phenomena displayed by transition metal compounds. These transitions are electron jumps between different energy levels within the d sub-shell. When a transition metal ion forms a complex ion, the d orbital's degeneracy is broken, creating a gap between the higher and lower energy d orbitals.

In complexes of copper and iron, for instance, this energy gap corresponds to visible light's wavelength. When such light is shined upon the complex, electrons in lower energy d orbitals can absorb photons and leap to the higher energy orbitals. The specific wavelength—or color—of the absorbed light determines the energy of the d-d transition. The light not absorbed is reflected or transmitted, and it is this remaining light that gives the complex its characteristic color.

Zinc complexes, however, cannot undergo such transitions since all the d orbitals are uniformly filled, and higher energy is required to accomplish these transitions. Hence, zinc remains colorless in its compounds. The intriguing aspect of d-d transitions is that the color we see is the result of a spectrum of light minus the absorbed frequencies—therefore, the study of these colors can reveal much about the electronic structure of the metal ion.