Which of the following species is diamagnetic?

To determine which species are diamagnetic, it is essential to understand the concept of electron configuration and how it relates to magnetic properties. A species is considered diamagnetic when it has all of its electrons paired, resulting in no net magnetic moment.

Starting with Co²+, this ion (Cobalt with a +2 charge) has an electron configuration of [Ar] 3d⁷. When cobalt loses two electrons to form Co²+, it will typically lose its two 4s electrons first, leaving a configuration of 3d⁷. In this case, there are unpaired electrons present in the 3d subshell, making Co²+ paramagnetic rather than diamagnetic.

On the other hand, O²⁻ (the oxide ion) has an electron configuration of 1s² 2s² 2p⁶ after gaining two additional electrons. In this configuration, all the electrons are paired – specifically, the 2p orbitals have two electrons each, which results in no unpaired electrons. This makes O²⁻ a diamagnetic species.

Therefore, since O²⁻ is diamagnetic due to its fully paired electrons and Co²+ is not, the answer that states both Co²