The notion of "shells" is a simplification.

You are familiar, I imagine, with the idea of the s-, p-, d-, and f-orbitals, since you mention them in your post. These correspond to ℓ = 0, 1, 2, and 3 respectively in the state of an electron. And they come in different levels, given by the value of n, also part of the description of an electron. So for example, the 2s subshell corresponds to n = 2 and ℓ = 0.

As a broad rule, these subshells are filled in the following order:

• Subshells with lower n + ℓ are filled first.
• For subshells with the same n + ℓ, start with the lowest n.

That results in the order 1s, 2s, 2p, 3s, 3p for the first three rows (with n+ℓ values of 1, 2, 3, 3, 4, 4 respectively). But once you get to the next row, the first that contains d electrons, it goes 4s, 3d, 4p. It doesn't "skip" d, it's just that the d it's filling in the fourth row of the periodic table is the 3d subshell, not the 4d one. The noble gas in that row (krypton, as it happens) does in fact have its highest occupied d shell filled. It has the electron configuration [Ar] 4s^(2) 3d^(10) 4p^(6) - this 3d subshell is full.

The reason the 3d shell shows up between 4s and 4p here, even though the n is nominally an energy level, is that that these numbers describe the energies of an orbital in the absence of other electrons. But other electrons in the atom jostle energy levels quite a bit. It turns out that when lower orbitals are occupied, d- orbitals end up so high energy that they effectively get "bumped up a tier" of the table.

It doesn't have any 4d electrons yet, because 4d electrons would be much higher energy than the 4s, 3d, and 4p electrons it actually has.

The better way to think about this is in terms of the gaps in energy levels. Noble gases have a large gap between the energy level of the highest-energy electron they have and the next available electron slot. That makes sticking an electron to them hard, because that electron has to occupy a high-energy state. And it makes stripping an electron off of them hard too, because all the electrons they have occupy low energy states. Their configurations look like this (where the blue lines represent energy of occupied orbitals, and red represents unoccupied orbitals).

As you go down the periodic table, the notion of "shells" starts to become less useful, because the gaps between the shells shrinks enough that the gaps within the shells can start to cause them to spill over one another. So chemistry near the bottom of the periodic table becomes more complicated, and in fact it's generally believed that if element 118 - which is one of the noble gases by its position on the periodic table - would actually be a solid if it were stable enough to stick around and have any chemistry at all.