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I teach school Physics. And am trying to get a reasonable explanation of the root cause of the energetic instability of unfilled outermost electron shells that leads to atoms bonding. For most of what I teach, I am able to tie things back to the four fundamental interactions/forces. But unfilled valence/outer shells being the reason why atoms bond beats me. I get Pauli exclusion. I think I get Schrodinger's equation. Yet, to me these don't provide a fundamental explanation of why a shell 'filled' by bonding is more energetically favorable than neutral atom with an unfilled outermost shell. Do I just accept this as fundamental property of our universe?

Santosh Zachariah
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    At the level of fundamental interactions there are no shells. Dozens of approximations has to be introduced starting from the fundamentals on the way to orbitals and shells. And understanding these approximations is essential to understand the stability of species with filled valence shells and instability of species with unfilled valence shells. – Wildcat Feb 24 '16 at 15:43
  • I'm not sure this question 'why' can actually be answered. It just 'is'. Wildcat's comment does not help (while true) and Spontification's answer speaks to myriad exceptions. – Lighthart Feb 24 '16 at 17:06
  • @Lighthart I object that my comment does not help. Once a person understand the orbital approximation it is rather trivial to understand why it is energetically more favourable to fill a particular shell before starting to fill the next one. – Wildcat Feb 24 '16 at 17:45
  • The comment it fails to actually address 'why', but rather discusses what is needed to answer the original question. This is why I find the comment unhelpful. – Lighthart Feb 25 '16 at 00:01
  • @Wildcat I get that it is energetically more favorable to fill a particular shell before starting to fill the next one. My question is "in terms of chemical bonding, why is a filled outermost shell more energetically favorable than an unfilled shell?" Or simply put, why do atoms bond? Lighthart are you saying I just tell my students "It is"? In which case, from the perspective of grade-school science I should teach them that to understand the physical sciences they need to understand there are four fundamental interactions, the atomic/standard model, and the "unhappiness of atoms." – Santosh Zachariah Feb 25 '16 at 17:45
  • @SantoshZachariah, until now I didn't thought your question is about bonding... Well, anyway. The simplest explanation of a chemical bonding is that, loosely speaking, if occupied orbitals of few atoms are collectively higher in energy than occupied orbitals of a molecule, then its energetically favorable to form this molecule. Fundamentally, one might say that bonding is driven by the principle of minimum potential energy. – Wildcat Feb 25 '16 at 18:12
  • Unfilled shells aren't "filled by bonding". When a bond is formed, atomic orbitals are fully reorganized into molecular orbitals. This a much more complicated process than a simple filling of unfilled shells. And besides, all this is just a model of reality. In reality, all you have is electron density being redistributed in energetically more favorable way. – Wildcat Feb 25 '16 at 18:20
  • But using the orbital model, you can explain why atoms with unfilled outer orbitals usually bond while those with filled do not as follows. For a couple of atoms with unfilled outer shells, you will populate only the bonding MO which is energetically favorable. But when you have two atoms with fully filled shells, not only bonding MO, but antibonding as well will be populated, which is energetically unfavorable. See for details this question. – Wildcat Feb 25 '16 at 18:42

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The idea of a "full" outer shell is a bit of a misnomer, for instance 3rd row has 3d orbitals but doesn't bother to occupy them until 4th row. Does Argon have a full outer shell, even though it has a $3d^0$ configuration?

Generally there's a large energy gap between the p-orbitals in one row and the s-orbitals in the next row, because of the increase in shielding and such factors, which acts as a kind of natural barrier, or a step, to adding any more electrons than what appears to be a 'full' valance shell.

But within any row, the amount of shielding and associated factors (penetration, nuclear charge, etc) remain largely similar, so if the valence shell isn't 'full' there is still sufficient attraction between the electron and a neutral atom that a negative ion may be stable. Adding more than this, the energy gap is too large to overcome under usual circumstances so it stops at a 'full' shell.

There are always exceptions. Noble gases lower down can form compounds even though they are "stable" "full" outer shells, and quirks like some transition metals having $s^1d^{10}$ configuration go to show that some of the ideas about electron filling and stable shells are rules of thumb, each individual case really needs to have its energy levels calculated to explain them.

Spontification
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  • Alright but why is it energetically favorable for a halogen to gain an electron. Your explanation doesn’t work because it isn’t going down a principal quantum number by gaining an electron. – EatingTechnobladesRemainsAt3am Apr 25 '23 at 08:10