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I read that the electronic configuration of uranium is [Rn] 5f³ 6d¹ 7s² . Given that the subshells fill in the order 5f --> 6d, why is the 5f subshell only partially filled? Why do electrons fill the 5f subshell partially then proceed to fill the 6d subshell?

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I am sure you are familiar with the rules for assigning electron orbitals, I will briefly descibe them here:

Electrons fill orbitals in a way to minimize the energy of the atom. Therefore, the electrons in an atom fill the principal energy levels in order of increasing energy (the electrons are getting farther from the nucleus). The order of levels filled looks like this:

Afbau build up

Pauli Exclusion Principle

The Pauli exclusion principle states that no two electrons can have the same four quantum numbers. The first three (n, l, and ml) may be the same, but the fourth quantum number must be different. A single orbital can hold a maximum of two electrons, which must have opposing spins; otherwise they would have the same four quantum numbers, which is forbidden.

Hund's Rule

When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy (also referred to as degenerate) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. This explains the behaviour of Chromium: Z:24 [Ar] 3d54s1 (note here the the one electron in 4s orbital while the d orbitals are occupied with single electrons of one spin direction)

Exceptions

Although the Aufbau rule accurately predicts the electron configuration of most elements, there are notable exceptions among the transition metals and heavier elements. The reason these exceptions occur is that some elements are more stable with fewer electrons in some subshells and more electrons in others and a notable example is uranium, for it to acquire maximum stability is usually having this ground state: Uranium: Z:92 [Rn] 7s2 5f3 6d1

References

  1. Rules for Assigning Electron Orbitals
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  • $\begingroup$ Ah so uranium is an exception to this rule. What is it about this specific configuration that makes it so stable? $\endgroup$ – Jonathan Smith Mar 17 '17 at 15:30
  • $\begingroup$ Its not only Uranium, read again I have also mentioned Chromium. There are other elements for example Copper, Niobium, Palladium, Silver, Thorium etc which deviate from this trend. The reason like described is partly based on the combination of the rules. Remember that in a ground state of an element the electron configuration has its lowest energy. The lower the energy the more the stability. In some cases this kind of stability can only be obtained when there are fewer electrons in a particular orbital say Uranium configuration. $\endgroup$ – xavier_fakerat Mar 17 '17 at 15:35

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