Question

Why is copper `3d^10 4s^1`??

Answer 1

For copper? Because the `3d` orbitals are significantly lower in energy for copper (thus making a doubly-occupied `4s` orbital unfavorable enough!), AND because it can fill all the quantum states available in the `n = 3` quantum level, as seen in its electron configuration of:

`[Ar] 3d^10 4s^1`

So, it's only natural for copper to "prefer" filling its `3d` orbitals rather than its `4s`.

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Consider the data here (Appendix B.9):

• `E_(4s,Cu) = -"8.42 eV"`
• `E_(3d,Cu) = -"13.47 eV"`

The `3d` orbitals of copper are about `"5.05 eV"` of energy lower than the `4s` orbitals of copper (which is about `"487 kJ/mol"`, almost `5` times the strength of a chemical bond!).

You can see the significant increase in favorability when you look at the big picture (graphed from the data here [Appendix B.9]), and see that the `3d` orbitals drop in energy faster than the `4s` orbitals do as the atomic number increases:

So, copper would have a more stable electron configuration by

• filling its `3d` orbitals to occupy all possible `bb(n = 3)` quantum states.
• taking advantage of the particularly low `bb(3d)` orbital energies compared to the `4s` orbital.

Answer 2

You mean to ask why is the electronic configuration of `"ATOMIC COPPER"` is........

Explanation:

You mean to ask why is the electronic configuration of `"ATOMIC COPPER"` is........`[Ar]3d^(10)4s^1` rather than `[Ar]3d^(9)4s^2`?

For copper, `Z=29`; and thus in the atom there are 29 electrons to distribute in the neutral atom. Eighteen of these electrons are distributed in a core the same as that of atomic argon, i.e. `1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)`. Because these are inner core, and do not really participate in the chemistry of the proceeding elements, we can represent this configuration by `[Ar]` rather than go thru the pfaff of writing it out.

So to your question, why is the electronic configuration of `"ATOMIC COPPER"` ........`[Ar]3d^(10)4s^1` rather than `[Ar]3d^(9)4s^2`? This is probably a combination of Hund's rule, which gives a special electronic stability to HALF-FILLED SHELLS, i.e. the `4s` orbital is half-filled; and also to the special stability of FILLED ELECTRONIC shells, i.e. the `3d` shell is FULL at `3d^10`.

The electronic stability to HALF-FILLED SHELLS, is also apparent in the atomic configuration of the `Cr` atom, which has a configuration of `[Ar]4s^(1)3d^5`. I would expect a 2nd year inorganic student to know these specific, special configurations,