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We know that if Cl and Na get too close, they produce ionic bonding.

Cl has 17 proton and 17 electrons and is considered stable.

Na has 11 protons and 11 electrons and is considered stable.

I understand that in Na, we have 1 valence electron and in Cl, we have 7 valence electrons. Now, if they get too close, 1 valence electron is transferred from Na to Cl. I don't understand why. Na was already stable. I get that part that Cl, since it has more protons, it would attract the electron from Na more than Na would do for itself due to less protons(11), but when ionic bond happens, Cl would end up having 17 proton and 18 electrons. Wouldn't Cl itself become unstable as there're more electrons than protons ? It seems like that before bonding, they both were stable, but after bonding, Cl ended up not stable.

I'm trying to understand it, but I get a point that explanation with quantum would be a waste of time as I have no knowledge in quantums. So would appreciate the explanation in terms of attraction forces and why after bonding, Cl doesn't become unstable. Note that octet rule and staff like this is something I don't really need to know as they're rules and not explanation in terms of attractions. I checked this why do atoms want to have a full outer shell but doesn't help.

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Here is an easier explanation not relying on "stability" or quantum mechanics.

You start out with some sodium ions and some chloride ions in aqueous solution. Each sodium ion is surrounded by water, and each chloride ion is surrounded by water. Then, you let the water evaporate, the sodium and chloride ions come together and form a crystal, where each sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six sodium ions.

We know that if Cl and Na get too close, they produce ionic bonding.

While ionic bonding is often explained as a transfer of electrons, it would be hard to find chlorine and sodium as elements. It is much more common that you already have sodium ions (as in sodium acetate or sodium hydroxide), and already have chloride ions (as in magnesium chloride) or make them from a molecular compound (as when hydrogen chloride gas comes in contact with water).

If you start with ions, you just have to figure out how they form an ionic solid under the right conditions (e.g. taking away water as a solvent).

Na has 11 protons and 11 electrons and is considered stable.

Elemental sodium is highly reactive. When it comes in contact with air or water, it will form ionic compounds or solvated ions. As you can read in the comments, "stable" becomes meaningful by saying stable with respect to some condition or in combination with which other substances.

I'm trying to understand it, but I get a point that explanation with quantum would be a waste of time as I have no knowledge in quantums.

If you want to have an explanation rather than a collection of rules and facts, you have to engage with quantum mechanics. All other explanations are incomplete. There is a perfect explanation of the strength of ionic bonds using pre-1900 physics, but to explain why the ions form in the first place, you need quantum mechanics.

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    $\begingroup$ Thanks so much. So when we have neutral sodium and on google, it says it's stable, it actually is stable in vacuum when only that sodium exists, but in respect to the world(putting it in the real world), it actually is not that stable anymore. Right ? I wondered why we can't explain it with attractions. when 2 neutral atoms come(sodium and chloride), maybe because chloride has more protons in it, its attraction for that valence electron is bigger than that valence electron is attracted by 11 protons in sodium, hence attraction dominates and Cl takes Na's electron. Is this wrong ? $\endgroup$
    – Chemistry
    May 4, 2023 at 15:58
  • $\begingroup$ "because chloride has more protons in it, its attraction for that valence electron is bigger than that valence electron is attracted by 11 protons in sodium, hence attraction dominates and Cl takes Na's electron. Is this wrong ?" That might work for NaCl, but does not work for NaF, right? $\endgroup$
    – Karsten
    May 4, 2023 at 16:10
  • $\begingroup$ you're correct. my logic is wrong. $\endgroup$
    – Chemistry
    May 4, 2023 at 16:25

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