There are several exceptions to the octet rule found along the transition and rare earth metal sections. These exceptions are tied to the observed phenomenon that atoms are more stable if orbital systems are either whole or half filled. In other words, the electron configuration 3p2 is less stable than 3p3, and therefore there are times that atoms will shift electrons from one l to another l for reasons of stability. Two of the more common examples are copper and iron and different configurations help explain why these metals display multiple valences. Explain the changes observed in electron configuration in copper and iron that allow for these multiple valences.

Question

There are several exceptions to the octet rule found along the transition and rare earth metal sections. These exceptions are tied to the observed phenomenon that atoms are more stable if orbital systems are either whole or half filled. In other words, the electron configuration 3p2 is less stable than 3p3, and therefore there are times that atoms will shift electrons from one l to another l for reasons of stability. Two of the more common examples are copper and iron and different configurations help explain why these metals display multiple valences. Explain the changes observed in electron configuration in copper and iron that allow for these multiple valences.

Can someone help me? I don't know how to answer the question.

₂₈Ni= 1s2 2s2 2p6 3s2 3p6 3d8 4s2
Therefore, with Cu we would expect the configuration to be 3d9 4s2 but instead it is
₂₉Cu= 1s2 2s2 2p6 3s2 3p6 3d10 4s1 to make the d orbital completely filled and the 4s orbital half filled. The same kind of thing happens with Cr to make 3d orbital half filled and 4s half filled. Then for Cu(I) the one 4s electron is removed.

I never thought of Fe following this route; however,
₂₆Fe 1s2 2s2 2p6 3s2 3p6 3d6 4s2. Removing the two 4s electrons makes Fe(II) and removing one of the 3d6 to make 3d5 (and 3d is now half filled) makes Fe(III). I hope this helps.

Answer 1

In copper (Cu), the expected electron configuration based on the position of copper in the periodic table would be 3d9 4s2. However, the observed electron configuration is 3d10 4s1. This is because copper preferentially fills its 3d orbital to make it completely filled and its 4s orbital half-filled, which is a more stable configuration. This is an exception to the octet rule.

Similarly, in iron (Fe), the expected electron configuration would be 3d7 4s2. However, the observed electron configuration is 3d6 4s2. This is because iron preferentially removes its two electrons from the 4s orbital to form Fe(II), and then removes one electron from the 3d orbital to form Fe(III). This results in a half-filled 3d orbital, which is more stable.

So, in both copper and iron, the changes in electron configuration occur to achieve a more stable configuration by either completely filling or half-filling certain orbitals. These changes allow for the display of multiple valences in these metals.