Colour and paramagnetism in transition metal compounds generally arise due to the presence of unpaired d-electrons.

• In $\text{CuF}_2$, the oxidation state of Cu is $+2$. Its electronic configuration is $[\text{Ar}] 3d^9$, which has one unpaired electron. Thus, it is paramagnetic and coloured due to d-d transitions.
• In $\text{K}_2\text{Cr}_2\text{O}_7$, the oxidation state of Cr is $+6$. Its electronic configuration is $[\text{Ar}] 3d^0$, which has zero unpaired electrons. Thus, it is diamagnetic (though it is coloured due to ligand-to-metal charge transfer).
• In $\text{KMnO}_4$, the oxidation state of Mn is $+7$. Its electronic configuration is $[\text{Ar}] 3d^0$, which has zero unpaired electrons. Thus, it is diamagnetic (also coloured due to charge transfer).
• In $\text{K}_4[\text{Fe(CN)}_6]$, the oxidation state of Fe is $+2$ ($3d^6$). Since $\text{CN}^-$ is a strong field ligand, all six d-electrons pair up in the lower energy $t_{2g}$ level, leaving zero unpaired electrons. Thus, it is diamagnetic.