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Suppose I pour some sodium chlorine into water. So, what happens is that

$$ \ce{NaCl (s) -> Na+ (aq) + Cl- (aq)} $$

meaning that the ionic bond between Na and Cl breaks up (correct this far?)

Now, does this mean that the water actually contains separate charged Na and Cl particles? So... since chlorine boils at −34.04 °C according to Wikipedia, why is there then no chlorine gas evaporating? Because it is chlorine ions there, not chlorine atoms?

If I feed electrons some how into the solution, will chlorine gas start forming?

Also, could I use this so that I pour NaCl into water and get Na and Cl ions, and then (since they are separate) add something more to create Na[something] or Cl[something]? Something that would not form if I just mixed NaCl + [something else]. I hope I am not completely lost here..

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  • $\begingroup$ I think you got it mostly right, but as there are anions of chlorine not cations one should take electrons away (via electrolysis for example) to obtain chlorine gas $\endgroup$ – Mithoron Feb 8 '15 at 22:35
  • $\begingroup$ The effects of diffusion and electrostatic attraction of ions with water overcome ionic bonding between sodium and chlorine atoms -- so yes, the charged sodium and chlorine ions separate. This is a computer simulation which includes a model for van der Waals and electrostatic forces in real time (a tiny $\ce{NaCl}$ cube dissolving in 4 ps): youtube.com/watch?v=IPQzY34eCt4 $\endgroup$ – khaverim Apr 6 '17 at 1:45
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Most of what you wrote is correct. It is difficult to feed many electrons into the salt water because the Coulomb forces are extremely strong, see 1 and 2, but you can feed electrons in if you drain them out elsewhere so there is no net gain or loss of charge. that is called an electric current.

You can use that to electrolyze salt water, and get chlorine as you surmise; NASA has a nice lab. BTW, you would expect to get Cl gas at the positive electrode, and do get it, but you do not get Na metal at the negative, because hydrogen bubbles out before Na. You can electrolyze molten salt, though, to get a bit of Na (or at least you see it form and then catch fire with a yellow flame).

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Suppose I pour some sodium chlorine into water. So, what happens is that

NaCl(s) → Na+(aq) + Cl-(aq)

meaning that the ionic bond between Na and Cl breaks up (correct this far?)

That would be completely correct in an infinitely dilute solution of NaCl.

However, in real solutions of NaCl there are ion pairs in addition to the separately solvated ions.

There is an equilibrium between the aqueous ion pairs and the aqueous seperately solvated ions:

$$ \ce{NaCl(aq) <=> Na+(aq) + Cl-(aq)}$$

For a detailed consideration see Physical Electrochemistry of Strong Electrolytes Based on Partial Dissociation and Hydration: Quantitative Interpretation of the Thermodynamic Properties of NaCl(aq) from "Zero to Saturation" Journal of Electrochemical Society, Vol. 143, No. 6, June 1996, pp. 1789 - 1793. (no paywall link).

In summary, in a relatively concentrated solution of NaCl, there will be about 20-30% ion pairs with the remainder being seperately solvated ions.

If I feed electrons some how into the solution, will chlorine gas start forming?

You need to remove an electron from each of two Cl- to get chlorine gas, but yes you can form chlorine gas through electrolysis of NaCl solution.

Also, could I use this so that I pour NaCl into water and get Na and Cl ions, and then (since they are separate) add something more to create Na[something] or Cl[something]? Something that would not form if I just mixed NaCl + .

You could added silver nitrate solution to an NaCl solution and get AgCl precipitate.

For a video of the reaction see this University of Illinois site: http://www.chem.uiuc.edu/clcwebsite/AgCI.html

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