Sciencemadness Discussion Board - Aqueus salts, alchohol: I understood it until I didn't

Aqueus salts, alchohol: I understood it until I didn't

jgourlay - 8-1-2014 at 04:37

Every few months one of my kids will get the jones to have me help them grow crystals: alum, copper sulphate, rock candy (which I have NEVER gotten to work), etc.

After, to recover the salt I dump some cheap vodka into the solution, it instantly precipates and everybody goes "oooh!!!" I have always explained this as the alchohol being soluble in water, but the salt not being soluble in ethanol, and thus the ethanol 'crowds out' the salt.

But I now I realize there is something I don't understand. Either that explanation is just wrong, or I'm missing the part about WHY the alchohol is able to 'crowd out' the salt even though the salt was there first.

Could you please explain this?

Zyklon-A - 8-1-2014 at 06:30

What salt are you talking about? The only way I've seen it, is where you have a hot solution of water and some salt (KNO3 works great) and then you pour very cold ethanol (or any other alchohol,) and because the salt is not as soluble in cold water, and insoluble in ethanol, and since the ethanol cools down the solution, the salt precipitates.

gravityzero - 8-1-2014 at 06:54

makes sense to me. LOL

DraconicAcid - 8-1-2014 at 08:23

The salt is very soluble in water because the dielectric constant of water is very high (it reduced the attraction between the ions). It's much less soluble in a mixture of alcohol and water because the mixture has a lower dielectric constant.

blogfast25 - 8-1-2014 at 09:17

Like chemical reactions, dissolution (whether it takes place or not) is ruled by whether the change in Gibbs Free Energy is negative or not:

ΔG = ΔH - TΔS

For dissolution, ΔS is usually positive because the dissolved state is less ordered (more probable) than the state where a solid salt just sits in a solvent (w/o dissolution), so entropy (S) increases on dissolution. So this usually counts towards making ΔG negative. And explains why salts that cool the solution on dissolving often still dissolve: ΔH < TΔS (in those cases).

When you introduce an ‘anti-solvent’ (like alcohol in this case) the precipitation is also caused by ΔG < 0.

[Edited on 8-1-2014 by blogfast25]

papaya - 8-1-2014 at 09:59

He asks why it's that the salt is precipitated rather than alcohol salted out!

DraconicAcid - 8-1-2014 at 10:15

Because the ethanol is much more soluble in water than the salt is- ethanol and water are miscible, but salt has a finite solubility in water. So the ethanol doesn't mind sharing the water with other solutes, but the salt does mind.

As an analogy, picture a subway car with some well-behaved, middle-class passengers. The bus stops, and a group of smelly, homeless drunks get on. There's plenty of seats, so it's not like there isn't room for the drunks, and the drunks aren't pushing people out of the car...and yet people precipitate out anyway.

jgourlay - 8-1-2014 at 11:40

Quote: Originally posted by Zyklonb

What salt are you talking about? The only way I've seen it, is where you have a hot solution of water and some salt (KNO3 works great) and then you pour very cold ethanol (or any other alchohol,) and because the salt is not as soluble in cold water, and insoluble in ethanol, and since the ethanol cools down the solution, the salt precipitates.

It works with both alum and copper sulphate. I know I've done this with rochelle salt, but can't remember if it worked.

jgourlay - 8-1-2014 at 11:55

Thanks for this answer. I understand this part.

Quote: Originally posted by blogfast25

For dissolution, ΔS is usually positive because the dissolved state is less ordered (more probable) than the state where a solid salt just sits in a solvent (w/o dissolution), so entropy (S) increases on dissolution. [Edited on 8-1-2014 by blogfast25]

What I don't understand is why the final state of a precipitated salt undissolved in the water/alchohol mixture is a more disordered, more 'probable', than the 'everything mixed' state? Why is S(all mixed) < S(precipitated)?

I appreciate the analogy about the smelly drunks downthread (and thank you for that), but it's less helpful because it doesn't explain how the 'decision' is made.

DraconicAcid - 8-1-2014 at 12:01

Quote: Originally posted by jgourlay

Part of the reason that there is more disorder in the precipitated version is that the precipitation is exothermic. As the stuff precipitates, it gives off heat, and and creates disorder in the solution. Also, ions in solution tend to be surrounded by water molecules which stick with those anions, reducing their entropy.

alexleyenda - 8-1-2014 at 13:12

I think you can also see this in term of molecular interactions. Water and salt are both very polar so they have Keesom interactions. Water and EtOH both have OH groups that make Very strong hydrogen bridge interactions, so the salt gets kicked out so EtOH and water can make stronger interactions. Furthermore, EtOH is less polar than water because of the C-H bounds so once again it tends to repel the very polar salt.

By the way, I was experimenting yesterday and randomly found a quite awesome reaction of this kind. I had a saturated solution of Na2SO4 (sulfuric acid neutralised with NaOH) and added methanol. Both liquids are colorless at first and when you mix them you get a thick milky very white mix that acts as a liquid. The salt froms a kind of porous paste in the solution and moves with it. It takes a couple minutes before it precipitates. That made my day, you could try it!

I tried it after with NaCl, NaHCO3 and KCl and none of them gave an effect more than one tenth of the effect I got with Na2SO4 so yeah give it a try it is awesome :p

[Edited on 8-1-2014 by alexleyenda]

DraconicAcid - 8-1-2014 at 13:18

Quote: Originally posted by alexleyenda

By the way, I was experimenting yesterday and randomly found a quite awesome reaction of this kind. I had a saturated solution of Na2SO4 (sulfuric acid neutralised with NaOH) and added methanol. Both liquids are colorless at first and when you mix them you get a thick milky very white mix that acts as a liquid. The salt froms a kind of porous paste in the solution and moves with it. It takes a couple minutes before it precipitates. That made my day, you could try it!

You can also do saturated calcium acetate (12 g in 40 mL water) with ethanol (300 mL). It gels nicely.

alexleyenda - 8-1-2014 at 13:23

I guess i'll try this, I keep that in note for when I'll have free time . It could be interesting, especially concidering the fact that i'll have to make the calcium acetate as I don't have any, and it probably takes glacial acetic acid for that and I'll have to make that too.

[Edited on 8-1-2014 by alexleyenda]

blogfast25 - 8-1-2014 at 14:28

Quote: Originally posted by DraconicAcid

Part of the reason that there is more disorder in the precipitated version is that the precipitation is exothermic.

No.

For one there's a higher degree of disorder in the state 'dissolved salt in solvent' than for the state 'solid salt (undissolved) and pure solvent').

Imagine a chess board with an equal amount of white pawns and black pawns: what is the most probable situation; all black on one side and all whites on the other side or a scrambled, random mix? The latter, of course. The latter represents the solution in this analogy.

Entropy is of course temperature dependent but here it's the dissolution and rearrangement into a far more disorderly (and more probable, i.e. higher entropy) state (the solution) that causes the entropy change.

In the case of a precipitation, entropy decreases but if ΔH is sufficiently negative (exotherm) it will 'win' over - TΔS

This page doesn't do too bad a job of explaining it better than I can:

https://www.boundless.com/chemistry/solutions/properties-of-...

In the case of precipitation with an anti-solvent, there's an added complication: the mixing of the solvent and anti-solvent also causes an entropy increase, because the mixture solvent/anti-solvent is more disorderly than the solvent and anti-solvent separately. This entropy increase is what drives the mixing of miscible liquids.

[Edited on 8-1-2014 by blogfast25]

violet sin - 8-1-2014 at 15:12

alexleyenda : I made a bunch of calcium acetate for my tomato plants last summer( calcium deficient). just used slacked lime for plant use and some vinegar. I didn't calculate the amounts needed as I wanted a lot. I just dumped a nice big pile of lime in the bottom and kept adding vinegar(after decanting finished sol) till it was there was no more white powder in the bottom.

the thing about this method was there was a LOT of extra water because of the low % acetic acid. having quite warm weather I strained the solution into mason jars. I then placed a small stick across the top and tied a bit of rag to it, which dipped into the solution. I then placed 4 or so jars side by side in a plastic tub, under a 5gal bucket( small blocks for stand off underneath for airflow). the bucket had a black garbage bag tied tightly to it. the idea was( which worked wonderfully BTW), black bucket got really hot in 110'F sun and limited venting kept the warm air in. the wick greatly accelerated the water loss and I had big white puffy masses of calcium acetate all down them. easy peasy and my tomato's loved it! the puffy masses dissolved completely for rapid calcium feeding. wish I would have thought about crashing it out of sol and saving some time. I just wanted to be able to shelf some for later with out a spill hazard. I'll have to try that next time

-Violet Sin-

BromicAcid - 8-1-2014 at 16:10

It's interesting too that the opposite can happen. Once upon a time I decided I wanted to electrolyze a solution of sodium chloride in water with some acetone. The hope being that I would generate hypochlorite which would react with the acetone as formed and chloroform would precipitate out. The electrolysis was done with nickel electrodes. After 30 minutes the magnetic stirring was stopped. Two layers had formed, a top acetone layer and a bottom water layer

And there are other threads on this site, interesting being the thread on the purification of alcohol by dissolving various salts into an alcohol solution in water and having the two form layers (the opposite of what you talk about here), the thread can be read at:

https://www.sciencemadness.org/whisper/viewthread.php?tid=24...

DraconicAcid - 8-1-2014 at 16:22

Quote: Originally posted by BromicAcid

It's interesting too that the opposite can happen.

Yes- I remember doing that once with isopropanol and sodium chloride.

alexleyenda - 8-1-2014 at 22:50

@ BromicAcid The link is interesting. I really don't understand why the K2CO3 kicks ethanol out of the water though, in theory the hydrogen bridges of EtOH and water should be stronger but as it was said in the thread you sent, it looks like K+ ions make special interactions with water, that could explain it. The question is what is it.

@ violet sin Thanks for the suggestion, thought I'll have to try something else for the water loss as the temperature atm where I live is around -20°C so ...you get the point :p

jgourlay - 9-1-2014 at 06:12

Quote: Originally posted by DraconicAcid

Fella's, thanks. I understand it now.

I really regret not having folks like you as professors in college.

jgourlay - 9-1-2014 at 06:13

Quote: Originally posted by alexleyenda

I think you can also see this in term of molecular interactions. Water and salt are both very polar so they have Keesom interactions. Water and EtOH both have OH groups that make Very strong hydrogen bridge interactions, so the salt gets kicked out so EtOH and water can make stronger interactions. Furthermore, EtOH is less polar than water because of the C-H bounds so once again it tends to repel the very polar salt.

By the way, I was experimenting yesterday and randomly found a quite awesome reaction of this kind. I had a saturated solution of Na2SO4 (sulfuric acid neutralised with NaOH) and added methanol. Both liquids are colorless at first and when you mix them you get a thick milky very white mix that acts as a liquid. The salt froms a kind of porous paste in the solution and moves with it. It takes a couple minutes before it precipitates. That made my day, you could try it!

I tried it after with NaCl, NaHCO3 and KCl and none of them gave an effect more than one tenth of the effect I got with Na2SO4 so yeah give it a try it is awesome :p

[Edited on 8-1-2014 by alexleyenda]

I'll try it!

jgourlay - 9-1-2014 at 06:23

Quote: Originally posted by violet sin

I had big white puffy masses of calcium acetate all down them. easy peasy and my tomato's loved it! the puffy masses dissolved completely for rapid calcium feeding. wish I would have thought about crashing it out of sol and saving some time. I just wanted to be able to shelf some for later with out a spill hazard. I'll have to try that next time

-Violet Sin-

Why not pour the sol. right into the soil? (I grow 'maters too)

jgourlay - 9-1-2014 at 06:30

Quote: Originally posted by blogfast25

1.No.

2. For one there's a higher degree of disorder in the state 'dissolved salt in solvent' than for the state 'solid salt (undissolved) and pure solvent').

3 .In the case of precipitation with an anti-solvent, there's an added complication: the mixing of the solvent and anti-solvent also causes an entropy increase, because the mixture solvent/anti-solvent is more disorderly than the solvent and anti-solvent separately. This entropy increase is what drives the mixing of miscible liquids.

[Edited on 8-1-2014 by blogfast25]

1. "No" as in precipitation is not exothermic...or 'No' as in the whole explanation is wrong?
2. I understand this, good.
3. I understand this, but if the "no" above refers to it being exothermic, how is solvent+anti-solvent+solute less disordered than solvent+anti-solvent+precipate?

Are you saying that the controlling phenomena here is not thermodynamics but the previously mentioned hydrogen bonding crowding out?

AJKOER - 9-1-2014 at 07:29

The reaction between CH3OH, Na2SO4 and H2O is interesting as not so widely known, but mentioned in one place in Wikipedia (but not in its Na2SO4 article), the reaction between Na2SO4 and H2O is better represented by:

Na2SO4 + 10 H2O --cooling--> Na2SO4(H2O)10

and not the usual hydrated salt (namely, Na2SO4.10H2O). Here is a quote from Wikipedia on the topic of water of crystallization (http://en.wikipedia.org/wiki/Water_of_crystallization ):

"Glauber's salt, Na2SO4(H2O)10, is a white crystalline solid with greater than 50% water by weight.

Consider the case of nickel(II) chloride hexahydrate. This species has the formula NiCl2(H2O)6. Crystallographic analysis reveals that the solid consists of [trans-NiCl2(H2O)4] subunits that are hydrogen bonded to each other as well as two additional molecules of H2O. Thus 1/3 of the water molecules in the crystal are not directly bonded to Ni2+, and these might be termed "water of crystallization"."

Now, in the context of the current discussion, namely dissolution and corresponding change in Gibbs Free Energy, I suspect, the nature of the bonding may be material when explaining the effects observed with Glauber's salt.

[EDIT] Some comments from Wikipedia (http://en.wikipedia.org/wiki/Dissolution_(chemistry) ) on the chemistry of dissolution to quote:

"The outcome of the process of dissolution (the amount dissolved at equilibrium, i.e., the solubility) is governed by the thermodynamic energies involved, such as the heat of solution and entropy of solution, but the dissolution itself (a kinetic process) is not. Overall the free energy must be negative for net dissolution to occur. In turn, those energies are controlled by the way in which different chemical bond types interact with those in the solvent."

And also:

"Compounds in a fluid state may also dissolve in another liquid depending on the compatibility of the chemical and physical bonds in the substance with those of the solvent. Hydrogen bonds play an important role in aqueous dissolution."

Also, per Wikipedia (http://en.wikipedia.org/wiki/Heat_of_solution ), the energy change is composed of three parts, to quote:

"1. Breaking solute-solute attractions (endothermic), see for instance lattice energy in salts.
2. Breaking solvent-solvent attractions (endothermic), for instance that of hydrogen bonding
3. Forming solvent-solute attractions (exothermic), in solvation.

where the final value of the enthalpy of dissolution (or so-called heat of solution) is the sum of the above steps.

[Edited on 9-1-2014 by AJKOER]

blogfast25 - 9-1-2014 at 08:50

Jgourlay:

‘No’, as in ‘the precipitation is exothermic but that doesn’t cause the entropy change’.

Re. your point 3, there’s only so far you can go with the ordered/disordered thingy. For the solute/solvent/anti-solvent system, you’d have to calculate the enthalpic and entropic changes and from there ΔG to determine what will happen.

“Are you saying that the controlling phenomena here is not thermodynamics but the previously mentioned hydrogen bonding crowding out?”

That’s the wrong way of looking at it. Take a simple example:

H2 + Cl2 === > 2 HCl

A lot is happening there: sigma hydrogen bonds are broken (which costs energy), sigma chlorine bonds are broken (which costs energy) and sigma H-Cl bonds are formed which releases energy. Entropy also changes from left to right. Overall ΔG = ΔH - TΔS is negative and the reaction proceeds (kinetic obstacles aside).

So the energies and entropies are determined at the atomic level and thermodynamics (very, very simply put) is the 'accountant' who decides what can go ahead and what not.

Read up on Gibbs Free Energy: fascinating stuff but not an easy concept. I'm at great risk of over simplifying it here, for our specific purpose.

[Edited on 9-1-2014 by blogfast25]

blogfast25 - 9-1-2014 at 08:57

Quote: Originally posted by AJKOER

but the dissolution itself (a kinetic process) is not.

In plain English: 'without stirring nothing happens' (unless you want to rely on diffusion only).