Sciencemadness Discussion Board - Suitable drying agents

Suitable drying agents

CHRIS25 - 2-10-2012 at 02:27

I have read that Sodium hydroxide, Calcium Chloride, calcium Oxide and silica gel make excellent drying agents: I knew they were hygroscopic, but what I also read was that NAOH is particular suited when it comes to absorbing acidic gases? Is this so? I was surprised that Potassium Hydroxide was not on the list. It would be nice if someone could clarify the above. I have calcium oxide and potassium and sodium hydroxide. I have noticed, subjectively, that Potassium hydroxide is very much more hygroscopic than sodium hydroxide, but then I have 99% KOH, and the NAOH is commercial.

[Edited on 2-10-2012 by CHRIS25]

weiming1998 - 2-10-2012 at 03:08

Quote: Originally posted by CHRIS25

I have read that Sodium hydroxide, Calcium Chloride, calcium Oxide and silica gel make excellent drying agents: I knew they were hygroscopic, but what I also read was that NAOH is particular suited when it comes to absorbing acidic gases? Is this so? I was surprised that Potassium Hydroxide was not on the list. It would be nice if someone could clarify the above. I have calcium oxide and potassium and sodium hydroxide. I have noticed, subjectively, that Potassium hydroxide is very much more hygroscopic than sodium hydroxide, but then I have 99% KOH, and the NAOH is commercial.

[Edited on 2-10-2012 by CHRIS25]

Potassium hydroxide would be (slightly) better at absorbing acidic gases, because theoretically, KOH is slightly more basic than NaOH. But practically, this difference is so slight it can be dismissed. But by comparing the hydroxides of two other alkali metals, CsOH and LiOH, we can see that the further down on the periodic table you go in group 1, the more basic the hydroxide is. Potassium hydroxide is also more soluble in water than sodium hydroxide, so that contributes to its hygroscopy. Its negative heat of hydration (slightly more negative than NaOH) also contributes to its hygroscopy.

Practically, though, I would go with what's cheaper. Both KOH and NaOH are extremely efficient at absorbing acidic gases. Heck, a sample of NaOH or KOH left outside in an open bottle for a few days quickly turns to the carbonate due to the CO2 in the air! But I would definitely not use CaO, due to it being a much weaker base, and a solid-gas reaction occurs much slower than a solvent-initiated one (NaOH and KOH contains traces of water due to its hygroscopy. However, by all means use CaO as a drying agent for organic solvents, since NaOH and KOH (especially KOH) has the habit of dissolving into organic solvents.

[Edited on 2-10-2012 by weiming1998]

CHRIS25 - 2-10-2012 at 05:32

Thankyou for that informative explanation. So a base is something that accepts the hydrogen ions, and therefore any metal oxide is a base? Plus ammonia, though I will not be using these of course, but even though ammonia is a base, left for a few weeks the ammonia disappears and you are left with water, However if in the presence of an acid in a tightly concealed air tight container, does then the ammonia (in solution) accept the Hydrogen ions and become acidified? I know I am going off topic, but just trying to understand a few more things.

Siggebo - 2-10-2012 at 06:14

Not all metal oxides are basic; read up on amphoteric oxides for example.
A Brönstedt base is a species that accepts protons, yes. There are other acid-base definition, for example Lewis acid/base. Wikipedia will help here also!.

Ammonia is a base, and when acid is present will accept hydrogen ions ("protons") and become ammonium ions.

weiming1998 - 2-10-2012 at 06:43

Quote: Originally posted by CHRIS25

Thankyou for that informative explanation. So a base is something that accepts the hydrogen ions, and therefore any metal oxide is a base? Plus ammonia, though I will not be using these of course, but even though ammonia is a base, left for a few weeks the ammonia disappears and you are left with water, However if in the presence of an acid in a tightly concealed air tight container, does then the ammonia (in solution) accept the Hydrogen ions and become acidified? I know I am going off topic, but just trying to understand a few more things.

This is a rather complex issue. Let me explain:

There are many definition for acids and bases. In the case of metal oxides, some are purely (weak) bases (as you said), others are both acidic and basic. Take aluminum (III) oxide, for example. It can act as a base (i.e, proton acceptor, as in: Al2O3+6H+---->Al(3+)+3H2O.), but it can also act as an acid (proton donor) as demonstrated: Al2O3+3H2O<---->Al(OH)3. Al(OH)3+H2O<----->Al(OH)4- +H+. This is all related to rather complex coordination chemistry and the like.

The answer to your second question is largely based on equilibrium. Ammonia dissolved in water forms NH4+ cations, OH- anions, and free ammonia molecules. They exist on an equilibrium as: NH3(aq)+H2O<---->NH4+(aq)+OH-(aq), with the free NH3 usually being the dominant species. When ammonia is placed outside, as ammonia is more volatile than water, it will escape the solution. But when an acid is added to it (say HCl), the free OH- ions will react with the H+ ions in HCl. As the OH- is consumed, more and more NH4+ and OH- is formed due to the removal of OH- shifting the equilibrium to the right.
Finally, after all OH- is consumed, only NH4+ is left, as the ammonium salt of an acid.

In anhydrous conditions, however, the ammonia molecules directly accepts the proton in the acid (hydrochloric acid as an example), forming NH4+ and Cl-. This is due to the proton being attracted to the lone pair at in the central nitrogen atom in an ammonia molecule. So yes, ammonia will accept the hydrogen ions and become acidified (or neutralised).

CHRIS25 - 2-10-2012 at 11:25

Right, a lot to digest here then. Nice to read clarity and then be able to work from here. thankyou Welming, and sigebo.

vmelkon - 2-10-2012 at 11:52

A list of drying agents + recommendations
http://www.erowid.org/archive/rhodium/chemistry/equipment/dr...

unionised - 2-10-2012 at 12:06

Calcium oxide is a much stronger base than sodium hydroxide.

It is used commercially for removing CO2 from gas streams and it works just fine even though CO2 isn't a strongly acidic oxide.

http://en.wikipedia.org/wiki/Carbon_dioxide_scrubber

KOH is not a significantly better drying agent than NaOH, but it's more expensive.
http://www.mercury-ltd.co.il/admin/userfiles/image/Informati...

CHRIS25 - 2-10-2012 at 14:01

Quote: Originally posted by vmelkon

A list of drying agents + recommendations
http://www.erowid.org/archive/rhodium/chemistry/equipment/dr...

But I do not see a category for Nitrates, (copper nitrate), it's not for drying but for removing enough water and all the acid to leave me with: Cu(NO3)2.xH20

And welming mentioned not to use CAO because it is a weaker base.

[Edited on 2-10-2012 by CHRIS25]

weiming1998 - 2-10-2012 at 16:55

Quote: Originally posted by CHRIS25

Quote: Originally posted by vmelkon

A list of drying agents + recommendations
http://www.erowid.org/archive/rhodium/chemistry/equipment/dr...

Sorry, my bad. CaO would indeed be a stronger base than either NaOH or KOH, because it contains the species O(2-). Water can act as a very weak acid, hydrolysing in an equilibrium. The first one is H2O<--->H(+)+OH-. But the OH- can further deprotonate, as: OH- <------>O(2-)+H+. The second reaction does not (I think) happen in any measurable extent in neutral water at all. Because water is a diprotic acid, the OH- would be its acid salt, and the O(2-) be its (fully neutralised) salt. Thus, CaO is more basic even though it contains the Ca(2+) ion, which is more acidic than the Na+ or K+ ion.

An example showing how CaO is the stronger base is comparing NaHSO4 and CaSO4. They're both salts of H2SO4, NaHSO4 being the acid salt, and CaSO4 being the fully neutralised salt. NaHSO4 is clearly more acidic than CaSO4, even though it contains the Na+ ion, because the HSO4- can donate a proton, but the SO4(-2) can't. It's the same for NaOH and CaO. The OH- can donate a proton (in extreme circumstances), but the O(2-) can't. Thus, the O(2-) is a much stronger base than the OH-, but the [Ca(H2O)6]2+ is only slightly more prone to hydrolysis than the [Na(H2O)6]+, so CaO is the stronger base.

That aside, NaOH and KOH solution are used in factories to remove CO2 and other acidic gases as well as CaO.

And anhydrous Cu(NO3)2, or in fact, most transitional metal nitrates, are extremely difficult to procure, because the hydrates can't be dehydrated by heating (it would decompose instead). So using them as a drying agent is expensive and a waste. Also, these also have a tendency to dissolve in (or react with) various organic solvents.

CHRIS25 - 3-10-2012 at 00:25

<<<<<<And anhydrous Cu(NO3)2, or in fact, most transitional metal nitrates, are extremely difficult to procure, because the hydrates can't be dehydrated by heating (it would decompose instead). So using them as a drying agent is expensive and a waste. Also, these also have a tendency to dissolve in (or react with) various organic solvents.>>>>>>

Ok, but leaving the copper nitrate in a dessicator bag with NAOH in a seperate beaker the same height as the beaker containing the copper nitrate would at least remove a lot of the acidic and water solution? Or maybe I will throw in the calcium oxide as well and see which one collects the water first.

[Edited on 3-10-2012 by CHRIS25]

weiming1998 - 3-10-2012 at 16:35

Quote: Originally posted by CHRIS25

<<<<<<And anhydrous Cu(NO3)2, or in fact, most transitional metal nitrates, are extremely difficult to procure, because the hydrates can't be dehydrated by heating (it would decompose instead). So using them as a drying agent is expensive and a waste. Also, these also have a tendency to dissolve in (or react with) various organic solvents.>>>>>>

Ok, but leaving the copper nitrate in a dessicator bag with NAOH in a seperate beaker the same height as the beaker containing the copper nitrate would at least remove a lot of the acidic and water solution? Or maybe I will throw in the calcium oxide as well and see which one collects the water first.

[Edited on 3-10-2012 by CHRIS25]

Dessicator bag would probably work. I would go with CaO, since it is more hygroscopic than NaOH (It can cause thermal burns in contact with water), but I haven't used a dessicator bags before since I always dry things outside in the sun, so put both in. One thing I'll be certain of is that it will remove both nitric acid residues and water, since both of them are strong bases.

CHRIS25 - 3-10-2012 at 23:43

Ah the sun....stories about the sun were handed down to me by my great great great grandfathers, about a time long ago when something yellow hung in the sky and people could not look at it...I always believed that that was just a myth....until now....I heard it on the chemistry forum, where something has been empirically observed - in another part of the world.

And the calcium oxide turns to hydroxide and then heat it up in the oven to convert back to CAO, so that I can re-use it?

[Edited on 4-10-2012 by CHRIS25]

weiming1998 - 4-10-2012 at 03:24

Quote: Originally posted by CHRIS25

Ah the sun....stories about the sun were handed down to me by my great great great grandfathers, about a time long ago when something yellow hung in the sky and people could not look at it...I always believed that that was just a myth....until now....I heard it on the chemistry forum, where something has been empirically observed - in another part of the world.

And the calcium oxide turns to hydroxide and then heat it up in the oven to convert back to CAO, so that I can re-use it?

[Edited on 4-10-2012 by CHRIS25]

Technically, yes. But I don't think an oven is hot enough. Wikipedia lists the dehydration temperature of Ca(OH)2 to be 500 degrees Celsius. Maybe even a bit over that to get it dehydrated fast enough.

CHRIS25 - 4-10-2012 at 04:07

Ok, my blow torch should do the trick then, 1050 degrees c.

weiming1998 - 4-10-2012 at 06:00

Quote: Originally posted by CHRIS25

Ok, my blow torch should do the trick then, 1050 degrees c.

You can't heat on it directly, since the constant CO2 and H2O from the flame will convert it to CaCO3, which requires a very long calcination at 1000 degrees to convert to CaO.

I'm not sure if you heat it through a heat-conductive surface, though. Fires can generally reach over 1000 degrees, but something like a stove can only go up to 400 degrees max.

CHRIS25 - 4-10-2012 at 08:34

I think the idea here would be not to convert it back to CAO entirely, but to drive off most of the water would at least make it somewhat re-useable again. I would place it in an army aluminium container I suppose, they are strong and I have heated them up to 500 degrees when melting other stuff.

Pyro - 4-10-2012 at 11:20

P2O5 is a great dessicator, but it is expensive. and I don't think you can recuperate it easily

CHRIS25 - 4-10-2012 at 11:48

How is anyone supposed to get Phosphorous oxide?

[Edited on 4-10-2012 by CHRIS25]

[Edited on 4-10-2012 by CHRIS25]

unionised - 4-10-2012 at 11:59

The usual method is to react phosphorus with oxygen.

CHRIS25 - 4-10-2012 at 12:56

Yes I know, but I meant that phosphorous is not something sold that readily, you need a permit.

weiming1998 - 4-10-2012 at 16:55

Quote: Originally posted by CHRIS25

Yes I know, but I meant that phosphorous is not something sold that readily, you need a permit.

P2O5 is usually brought directly from lab suppliers. It reacts with water to form phosphoric acid, which cannot be dehydrated back to P2O5 by conventional means. Thus, it cannot be regenerated. It is also fairly expensive. But it makes up for that with its extreme attraction towards water. It can dehydrate H2SO4 to SO3, and char organic materials on contact.

But I think using P2O5 for your purposes is a bit over the top. Simple NaOH, KOH or CaO will work. If you plan on using the dessicator bag, there would also be gas build-up if you use P2O5, as it will dehydrate nitric acid to N2O5, which decomposes to form NO2 and O2.