How would you account for the following:
$(i)$ of the $d^{4}$ species, $Cr ^{2+}$ is strongly reducing while manganese $(III)$ is strongly oxidising.
$(ii)$ Cobalt $(II)$ is stable in aqueous solution but in the presence of complexing reagents it is easily oxidised.
$(iii)$ The $d^{1}$ configuration is very unstable in ions.
$(i)$ $C r^{2+}$ is strongly reducing in nature. It has a $d^{4}$ configuration. While acting as a reducing agent, it gets oxidized to $Cr ^{3+}$ (electronic configuration, $d^{3}$ ). This $d^{3}$ configuration can be written as $t_{2 g }^{3}$ configuration, which is a more stable configuration. In the case of $Mn ^{3+}\left(d^{4}\right),$ it acts as an oxidizing agent and gets reduced to $Mn ^{2+}\left(d^{5}\right) .$ This has an exactly half-filled $d$ -orbital and is highly stable.
$(ii)$ $Co ( II )$ is stable in aqueous solutions. However, in the presence of strong field complexing reagents, it is oxidized to $Co (III)$. Although the $3^{\text {rd }}$ ionization energy for $Co$ is high, but the higher amount of crystal field stabilization energy $(CFSE)$ released in the presence of strong field ligands overcomes this ionization energy.
$(iii)$ The ions in $d^{1}$ configuration tend to lose one more electron to get into stable $d^{0}$ configuration. Also, the hydration or lattice energy is more than sufficient to remove the only electron present in the $d$ -orbital of these ions. Therefore, they act as reducing agents.
Which of the following is diamagnetic transitional metal ion
Which of the following statement is not true
In which of the following ionic radii of chromium would be smallest
The magnetic moment of a metal ion of first transition series is $2.83 $ $BM$. Therefore it will have unpaired electrons
Paramagnetism is exhibited by molecules