Are Na and Cl ions from NaCl in water separate charged particles?

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..




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](http://physics.bu.edu/~duffy/py106/Charge.html) and [2](https://chem.libretexts.org/Core/Physical_and_Theoretical_Chemistry/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Intermolecular_Forces/Coulombic_Forces), 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](http://aquarius.umaine.edu/cgi/ed_act.htm?id=47). 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](http://en.wikipedia.org/wiki/Ion-association) 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"](http://www.jh-inst.cas.cz/%7Erheyrovs/jes1996.htm) 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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