[iScience]

Don't know if this is an O-chem or P-chem question but here it is



Say i have aqueous NaHCO3; the bicarbonate and its counter ion (Na+) will dissociate, but as the water evaporates, if solid sodium bicarbonate is formed via the sodium ions falling into place near the R-O^-, shouldn't there be some bicarbonates with H2O (or H3O+) in place of the Na+? I mean, the sodium ions fall/lock into place via ionic interactions. the polar interaction with water wouldn't be as strong but it's still an interaction that's present. so then anytime i have a chemical of such configuration (anion and cation held together by electrostatic attraction but not quite an ionic bond) obtained via crystallization, shouldn't i always have a percentage of this crystal hydrated?

[Arkcon]

Bit hard to follow your question, but I didn't want it to get lost. So maybe you can try asking it a different way, and see if we can help you.  I've moved it here because, just to let you know, carbonates, bicarbonates-- all salts of carbonic acid, and carbon dioxide itself are inorganic compounds.  I know they contain carbon, and biochemically, are soon going to be converted into organic molecules.  But that's the definition we've all agreed to work with.

[iScience]

realized there was an easier way to state what i was trying to say:

if we were to evaporate the water from an aqueous sodium acetate solution(ie recrystallization), wouldn't the sodium acetate crystal be hydrated with not only H2O, but also H3O+ and OH- ions by spontaneous dissociation of water?

[Arkcon]

No.  Water of hydration is what we call water molecules that end up inside the crystal structures.  You seem to know that H⁺ doesn't exist, and have substituted the other one H₃O⁺.  Now let me disappoint you again -- that one and OH^- don't exist either.  Even NaCl doesn't really ionize into Na⁺ and Cl^-.  There are just models we use to describe the behavior of these substances in solution, they don't really exist.  Every so often, someone tries to use some sort of trick to "catch" an ion in solution inside of a solid.  And that can't be done -- because ions don't exist in solution either.

Ions in solution are associated with water molecules, several of them.  And those molecules have different behavior towards other water molecules, that water molecules usually have with each other.  And these water-ion clusters are responsible for the properties of ions in solution.  I'm sad to say, this is just what I heard, and I don't understand this concepts very well, this is very advanced chemistry.  But sooner or later, every chemistry student has to be told that the models we're using don't really describe reality. 

How could we ever prove an Na⁺ ion is in a liquid? How would we ever see the OH^- is in a crystal.  We can't see them.

[Arkcon]

You asked a similar question here:  http://www.chemicalforums.com/index.php?topic=72010.0

And another time, you got a good visual of what I was saying here, but I lost that particular posted image.

[iScience]

isn't it a bit extreme to state salts don't dissociate into their complementary parts? i'm having a bit of a difficult time picturing everything i've been taught about this being wrong.

would it perhaps be more accurate to say that NaCl does dissociate into Na+ and Cl- and then there is a hydration sphere formed around the ions?

[Arkcon]

Yes.  those are the words I was looking for    The hydration sphere is a more apt model for how ions behave in solution.  Do you remember that from your other thread?  The one I've lost track of?  There was a good picture there, I think it was from Inigo:

[snorkack]

In liquid, the molecules and ions are located in a random manner, are miscible in a wide range of concentrations and move rapidly about.

In crystals, they are stuck, and are stuck in a long range order.

If you cool a saturated saltwater solution - yes, Na⁺ and Cl^- will forsake the solution, break off from all hydration water they have in solution and form dry, cubic NaCl crystals, as long as the temperature is above about 0 degrees.

When the temperature is below that temperature, NaCl starts to form a crystal hydrate - NaCl·2H₂O. This is a substance with definite composition based on small integers, definite crystal shape and definite long-range order inside the crystal. It is just a different order than the one inside NaCl or in ice crystals.

If you grow NaCl·2H₂O crystals then you can put them to X-ray diffraction or neutron diffraction machines (provided you do not warm them over 0 degrees, where they would incongruently melt and be thereby destroyed) and find out where specifically the H₂O molecules exist in every unit cell, and where their protons are relative to the neighbouring O, Na and Cl nuclei.

Now, yes, neutral lioquid water can autodissociate, about 1 molecule out of 5 millions. Ice cannot because OH^- ion would not fit in the crystal long range order and ice does not have an unit cell of 5 million molecules. An OH^- ion would be a crystal defect, so it is rejected by crystal growth, concentrated in the remaining water, then neutralized by some H₃O⁺ ion also concentrated there.

If you freeze, say, HCl, then you can form crystal hydrates HCl·6H₂O, HCl·3H₂O, HCl·2H₂O and HCl·H₂O. HCl is a strong acid, and in aqueous solutions nearly completely dissociated into Cl^- and H₃O^-. Does this mean that the crystal hydrate should because of its long-range order be 100 % dissociated, because the crystal structure would be Cl^-·H₃O⁺·5H₂O - that any undissociated HCl in solution will dissociate on freezing because an HCl molecule would be a crystal defect? And can you examine a HCl·6H₂O crystal by x-ray and neutron diffraction and NMR of the 19 protons and the 6 O-17 nuclei in the unit cell to verify the bond lengths and find out which of the 6 oxygens in in H₃O⁺ ion rather than a H₂O molecule?

[Borek]

are miscible in a wide range of concentrations

I don't see how to apply "miscible" to solution of ions. "Miscible" refers to the way liquids mix.