An odd thing happens when looking at Schrödinger’s math for atoms with more than one electron. It turns out that a set of orbitals with identical energy (like the three p orbitals on level 2, or like the 5 d orbitals on level 6, etc) are more stable when they are filled, half-filled or empty. That means that atoms with configurations that end in p3, or p6 or d5 or d10 are more stable than atoms whose configurations end in other numbers of d electrons. In fact, the stability of the p6 arrangement is why the noble gases are inert (non-reactive).
Because of this, a strange thing happens with elements whose electron configurations end with d4 or d9. Let’s focus specifically on two elements: chromium, [Ar]4s23d4, and copper, [Ar]4s23d9. Both of these configurations are ONE electron shy of stable.
If we also remember that the 3d orbital is just higher in energy than the 4s orbital, and that an s orbital is always stable because it is always either full, half-filled or empty.
So, nature can move an electron out of the 4s without losing stability and place it into the 3d’s thus creating stability.
That means that the electron configuration of chromium is [Ar]4s13d5 and the configuration of copper is [Ar]4s13d10.
You should assume that this occurs for all d4 and d9 elements.
The truth (as always) is more complicated, and there are a few other elements that do some shifting of electrons, but those examples are rare and not relevant at this point in your life and chemistry career.
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