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This question already has an answer here:

How come atoms can have such a wide range of appearance, as well as atomic & chemical properties?

At one point you have something invisible & gaseous at room temperature, add a few protons/neutrons/electrons, and you have something solid & shiny. Seems like a huge difference for a small change.

Even within the same element you can have a wide spectrum of properties - carbon for example has a few quite different allotropes (e.g. diamond vs graphite vs Q-carbon).

I'm not so much interested on how these differences arise (e.g. valence electrons), but rather why they are so varied. (Which is why I don't think this is a duplicate of either Do current models of particle physics explain the chemical properties of elements/compounds? or What enables protons to give new properties to an atom every time one is added?)

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marked as duplicate by Qmechanic Dec 12 '17 at 23:29

This question has been asked before and already has an answer. If those answers do not fully address your question, please ask a new question.

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    $\begingroup$ What kind of answer do you expect to "why" so it is not covered by the answers to "how" in this case? $\endgroup$ – nasu May 29 '17 at 16:11
  • $\begingroup$ I don't really know myself what kind of answer I'm expecting. I'm guessing less of a "technical" one, more along the lines of r/explainlikeimfive ... when I started thinking about this, I was intrigued how varied elements are, considering it's "just" a slightly different mix of particles in each. $\endgroup$ – Orangenhain May 29 '17 at 20:39
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The most obvious issue is that the electronic configuration determines periodicity in chemical behavior. Further, each new proton also shrinks the mean orbital radius of the outermost electrons, while electrons entering a higher energy level expands said radius. Charge-to-size ratios (for atoms, never mind the more numerous kinds of molecule) thereby end up all over the place, with implications for density, electronegativity and melting and boiling point. All this results in characteristic behaviours of s- and p-block elements, becoming truer in later periods for left-hand groups and less true in later periods for right-hand ones, and molecules being nothing like their constituent atoms (as usually they have full outer electron shells). For example, later alkali metals (halogens) more (less) readily lose (gain) electrons because of smaller $Z/r^2$ values.

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Let's take the contrast between Helium and Lithium as an example. He has two protons and therefore two electrons, and they completely fill up the 1s orbital. So He cannot make any covalent bond in its atomic form. Add one proton to get Li, and therefore one electron as well. This electron will go on the next orbital, 2s, which has therefore room for another electron. So Li can make one covalent bond. The electron on the 2s orbitals can also "escape" their respective nuclei and make a "sea" of electrons shared by a macroscopic amount of atoms: that's why Li is a metal.

Now add some more protons to get to Neon. Its electrons distribute as follow between the atomic orbitals (number of electrons between parentheses): 1s(2) 2s(2) 2p(6). So all the three 2p orbitals are full, and Ne cannot make covalent bond. As He, and that's why they are on the top of each others in the Periodic Table.

Now add one more proton to get Na. The extra electron will go on the next available orbital, 3s, which has therefore room for another electron. As Li, it can make one covalent bond, or share that electron in the metallic manner. So they have very similar chemical properties, and that's why they are on the top of each others in the Periodic Table.

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