Problem 34 - Chemistry: An Atoms-Focused Approach

Question

Unlike the compounds of most transition metal ions, those of \(\mathrm{Ti}^{4+}\) are colorless. Why?

Step-by-Step Solution

Answer

Answer: The compounds of \(\mathrm{Ti}^{4+}\) ions are colorless because there are no electrons in their d-orbitals. As a result, no electronic transitions between the d-orbitals can occur, which is responsible for the color in transition metal compounds. Since no part of the visible light spectrum is absorbed due to these electronic transitions, the compounds appear colorless.

Step 1: Find the electronic configuration of \(\mathrm{Ti}^{4+}\)

To determine the electronic configuration of \(\mathrm{Ti}^{4+}\), we first need to know the atomic number of titanium (Ti) which is 22. The ground state electronic configuration is \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^2 4s^2\). Since we are looking at the \(\mathrm{Ti}^{4+}\) ion, it means that 4 electrons have been removed from the neutral atom. Titanium loses its two 4s electrons followed by the 3d electrons. So, the electronic configuration of \(\mathrm{Ti}^{4+}\) is: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^0\)

Step 2: Analyze d-orbital electron configuration

For \(\mathrm{Ti}^{4+}\), as seen in Step 1, the electron configuration in the 3d orbital is \(3d^0\). This means there are no electrons in the d-orbital. The color phenomena in transition metal ions occur due to the electronic transitions between the split d-orbitals, which requires the presence of at least one electron in the d-orbital.

Step 3: Explain the absence of color in \(\mathrm{Ti}^{4+}\) compounds

Since there are no electrons in the d-orbitals of \(\mathrm{Ti}^{4+}\), no electronic transition between the d-orbitals can occur. As a result, the compounds of \(\mathrm{Ti}^{4+}\) ions are colorless, since they do not absorb any part of the visible light spectrum due to these electronic transitions.