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AJKOER
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Easy Sodium alum?
I recently placed (NH4)2SO4 (freshly prepared aqueous solution from the action of weak aqueous ammonia on MgSO4) in a
vessel with Aluminum foil strips, a piece of Copper metal, and NaCl. This was an experiment to create an Aluminum salt via an electrochemical
reaction. I am, in effect, creating a battery cell, as no external current is applied.
Within 6 hours, a milky solution with a evident white powder formed and on some of the Al strips, an interesting reddish substance attached (see
pictures). I believe this to be pure Copper displaced by the Al as per Wikipedia (http://en.wikipedia.org/wiki/Copper ) to quote:
"Pure copper is orange-red and acquires a reddish tarnish when exposed to air"
I am not exactly sure on the electrochemistry, but I expected the formation of OH- ions to form Al(OH)3, or an aluminate, or possibly a
double salt.
If there was insufficient ammonia used in creating the starting solution of Ammonium sulfate (but the evident smell of ammonia in the final reaction
mix may suggest otherwise), the reaction of MgSO4 with OH- could possibly form the insoluble Mg(OH)2. However, at the end of 9
hours, the reaction mix has cleared with a good amount of a white precipitate (resembling wet baking soda in texture). My experience with
Mg(OH)2 suggests it does not so readily fall out of solution and the texture in solution appears somewhat different. I suspect this salt to
be either Na2SO4.Al2(SO4)3 (Sodium alum), or a form of Al(OH)3.
Some of many possible reactions in the cell that may have resulted (?) in the formation of Sodium alum (to be confirmed) in a non-neutral solution:
8 H2O --> 4 H3O+ + 4 OH-
Al + 4 OH- --> Al(OH)4-
2 Al(OH)3 + 3 (NH)4)2SO4 + Na2SO4 -->
Na2SO4.Al2(SO4)3 (s) + 6 NH3 + 6 H2O
[Edited on 21-11-2013 by AJKOER]
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bfesser
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This is pretty unscientific, even more so than I had expected from you.
<img src="http://clipartist.info/Art/April/facepalm_facepalm-999px.png" width="14" alt="facepalm" />
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AJKOER
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Actually, this is an experiment in progress. A reaction, by the way, that has not been reported or performed previously, to the best of my research.
It was inspired by the reputed ability of Ammonium sulfate to attack certain Aluminum alloys including those containing copper. This suggested to me a
possible electrochemical basis for this observed, and important, corrosion reaction. And, guess what, there is apparently!
I also suspect now, after conducting more research, that the product may be more likely one form of Basic aluminum sulfate (see, for example,
discussion at http://www.google.com/patents/US4526772 ) and not, after all, Sodium alum. For those seeking chemical precision, may I suggest a salt other than
Basic aluminum sulfate as, to quote from the patent, it apparently exists "in various forms such as fibers, spheres, prisms and radially oriented
particles" and has an indefinite chemical formula to boot.
To be fair, I did pose the creation of Sodium album as only a possibility. My second sentence above, to quote: "This was an experiment to create an
Aluminum salt via an electrochemical reaction" and it certainly does.
The important point is that the electrochemical galvanic corrosion reaction ( http://tnb.ca/en/catalogues/online/cabletray/pdf/c10/09-corr... ) on Aluminum (see discussion on page 18 at http://books.google.com/books?id=KXwgAZJBWb0C&pg=PA18&am... ) involving Ammonium sulfate and Copper does apparently occur, and does so in
hours, not days, in an appropriate electrolyte! This thread suggests capitalizing on this fact to form Aluminum salts.
--------------------------------------------------
Now, I can understand why one may view some of my lack of details (like quantities employed for example) as unscientific. So, some more explanation on
my part.
> The NaCl is added to the solution to foster the cell reaction by providing a good electrolyte (it is not consumed in the reaction itself).
> The concentration of the Ammonium sulfate was not specified. This is due to one of differences from normal chemical reactions. Specifically,
electrochemical reactions proceed on differences on electro potentials and not concentrations (see interesting comments by the author of one of my
sources in my discussion of the so-called 'bleach battery' in a prior thread at http://www.sciencemadness.org/talk/viewthread.php?tid=24318 , where the author notes that in a HOCl based cell, even very small quantities of the
acid are preferentially searched out). To quote (source link: http://www.exo.net/~pauld/saltwater/ ):
"Even though there is relatively little HOCl in bleach, the latter reaction is more favored because of its large potential of 3.93 volts."
> Some related references (given in my prior thread) on the electrochemistry of batteries: see http://sci-toys.com/scitoys/scitoys/echem/batteries/batterie... and also http://www.dtic.mil/dtic/tr/fulltext/u2/d019917.pdf
[Edited on 21-11-2013 by AJKOER]
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blogfast25
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Quote: Originally posted by bfesser  | This is pretty unscientific, even more so than I had expected from you.
<img src="http://clipartist.info/Art/April/facepalm_facepalm-999px.png" width="14" alt="facepalm" /> |
Personally I think he forgot to take his medication again. Does he even know how soluble (and thus difficult to crystallise) sodium aluminium alum
actually is? There's a reason why the K and NH4 alums are vastly preferred.
What he's doing is creating a witches brew, followed by a guessing game: 'what is going on in my black box?'
Seriously, try preparing an alum the old fashioned way: you [AJ] might actually learn something...
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AJKOER
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Actually, Blogfast25 thanks for bringing up the solubility point.
In fact, the name of this thread should have more appropriately been "Ammonium alum?" and not "Sodium alum?". This was a confusion on my part as the
solubility property (or, more precisely lack thereof) is an attribute of Ammonium aluminum sulfate (see Wikipedia discussion at http://en.wikipedia.org/wiki/Ammonium_aluminium_sulfate and also their solubility table). The excitement of forming a potentially useful,
insoluble, Aluminum salt got to me, where are my meds when I truly need them!
In fact, Ammonium aluminum sulfate, (NH4)2SO4.Al2(SO4)3·24H2O is a real candidate as it is prepared "from aluminium hydroxide, sulfuric acid and
ammonium sulfate" to quote from my Wikipedia reference. [EDIT] No longer a candidate as heating an aqueous suspension of the white salt does not
improve its solubility. So, back to Basic Aluminum sulfate (here is some solubility work on the salt http://pubs.aic.ca/doi/pdf/10.4141/cjss70-030 ) or even Magnesium hydroxide as I am not actually seeing significant corrosion on the Al strips (I
will repeat this experiment with an excess of the Ammonium sulfate relative to the Al and, in one variation, add H2O2 to assess any impact on salt
formation).
Blofast25, I am sorry you feel that their are just too many possible combinations in this brew. Perhaps I could recommend some meds?
Apparently, per Wikipedia, Sodium sulfate forms a couple of double salts. To quote Wikipedia ( http://en.wikipedia.org/wiki/Na2so4 ):
"Sodium sulfate displays a moderate tendency to form double salts. The only alums formed with common trivalent metals are NaAl(SO4)2 (unstable above
39 °C) and NaCr(SO4)2, in contrast to potassium sulfate and ammonium sulfate which form many stable alums."
[Edited on 21-11-2013 by AJKOER]
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blogfast25
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Something mildly innovative would be to prepare it from Al scrap, H2SO4 and strong ammonia. 1 mol Al + 2.1 (maybe even 2.2) mol H2SO4, dissolve. Then
neutralise the excess H2SO4 with 1/2 mol NH3. Filter hot (if needed), then chill.
Because ammonium alum is quasi insoluble at 0 C, it doesn't even matter what initial hot concentration you start from: recovery of alum crystals will
always be high. Recrystallise from 2 parts alum + 1 part water.
[Edited on 21-11-2013 by blogfast25]
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AJKOER
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Thanks Blogfast for the suggestion.
Actually, what I am really interested in is the ability of using widely available household chemicals (I am excluding H2SO4 also due to safety
concerns on storage, and even NaHSO4) in a attempt to prepare easily and in large quantities Aluminum sulfate. For example, via galvanic corrosion
with some simple items like scrap Al, Cu and Epsom salt.
With thermal decomposition of Aluminum sulfate one then has access to SO3, H2SO4,...
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MrHomeScientist
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I've prepared potassium alum from soda cans in a similar way, by first reacting the can with excess KOH solution, filtering, and then treatment with
H2SO4. It would be interesting to try with ammonia. I imagine I would have to react with acid first, as ammonia isn't powerful
enough to react away the metal (as far as I know). I prefer using base first, as this separates out any non-amphoteric alloying metals.
While I am glad to see AJ actually doing experimentation lately, rather than the usual wild speculation, I don't understand why he refuses to post
quantities for his experiments. You can't hope for repeatability or accurate analysis of results if you don't measure and report quantities. It's not
like we're going to steal your method or something.
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ScienceSquirrel
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I think the best way to go would be to mix a warm solution of ammonium sulphate with a warm solution of aluminium sulphate and cool.
Both are readily available so it should be just cooking chemistry.
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blogfast25
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| Quote: Originally posted by ScienceSquirrel | I think the best way to go would be to mix a warm solution of ammonium sulphate with a warm solution of aluminium sulphate and cool.
Both are readily available so it should be just cooking chemistry. |
Sure but it's a bit more adventurous starting from simpler materials. I used to make 1 kg batches of NH4 and K alum for sales and then it's definitely
cheaper to start from the simplest precursors possible. I ended up buying them in of course.
@Mr HS: yes, ammonia doesn't dissolve aluminium at all. So you need to start by dissolving the metal in excess acid, then neutralise the excess with
NH3. I've done this with KOH too, works very well.
[Edited on 22-11-2013 by blogfast25]
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AJKOER
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| Quote: Originally posted by blogfast25 | ......
@Mr HS: yes, ammonia doesn't dissolve aluminium at all. So you need to start by dissolving the metal in excess acid, then neutralise the excess with
NH3. I've done this with KOH too, works very well.
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Practically, this is a true statement. Technically, with time, placing freshly created strips of Aluminum foil in a solution of aqueous ammonia (or
even NaCl) results in an eventual penetration of the protective Al2O3 layer (blistering) allowing an interaction between the reactive Aluminum and
water:
2 Al + 6 H2O --> 2 Al(OH)3 (s) + 3 H2 (g)
This reaction can present a bursting issue for sealed jars or plastic containers. The form of the Aluminum oxide precipitate may vary, but with
ammonia, a clear gel is observable.
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blogfast25
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What could be interesting to do is use the aluminium reflective coating of CDs. I used to peel off the paint + Al to get transparent diffraction
gratings (for spectroscopes). You stick some gaffer tape to the labelled side, then pull it off suddenly, hoping the label + Al coating come off.
Except... usually some (sometimes most of it) of the Al coating still sticks to the CD. I used to remove that using NaOH solution. Rarely have I seen
Al react so quickly with an alkali, it's almost instantaneous removal (dissolution).
The foil is of course very thin and very clean. It could be worth comparing the action of equivalent solutions of NaOH and NH3 on that type of foil.
[Edited on 23-11-2013 by blogfast25]
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plante1999
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I'm actually impressed AJKOER did a real experiment at all bfesser. As been a while I waited for that. I would have taught bleach would be more used
though...
I never asked for this.
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AJKOER
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Talk about experimenting, relative to precursor chemicals, no need for NaOH to dissolve Aluminum. Consider the hydrolysis of Na2CO3:
Na2CO3 + H2O <---> NaOH + NaHCO3 [EDIT correction]
And the following source (link: https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ) that notes the reactions, to quote:
" 2 Na+ + 2 Al + 2 OH- + 6 H2O → 2 Na+ + 2 Al(OH)4- + 3 H2 (3)
and
2 Na+ + 2 Al(OH)4- → 2 Na+ + OH- + 2 Al(OH)3↓ (4)"
Another source (http://eng.sut.ac.th/metal/images/stories/pdf/02_Aluminium%2... ) specifies the decomposition of NaAlO2 at around 50 C per the reaction:
2 NaAlO2 + 3 H2O --> 2 NaOH + Al2O3. 3H2O
The experiment I performed was simply placing Aluminum foils strips in a strong solution of Na2CO3 (prepared from heating up dry household NaHCO3).
What happens is nothing, until the solution is heated to boiling. Then, remove the heat and the solution will continue to boil with the generation of
steam vapors, apparently hydrogen and the attack of the Al. Wait a few minutes until the reaction stops and re-heat. The reaction procedures as it did
originally, but, I suspect, it is apparently the newly created Sodium aluminate now attacking the Al. A white suspension of Al(OH)3 is also formed
(and diluting the solution with distilled water also forms the precipitate). Some more references on the stability issues of NaAlO2, please see http://www.usalco.com/wp-content/uploads/2013/05/Sodium_Alum... and http://www.us.edu.pl/uniwersytet/jednostki/wydzialy/chemia/a...
[Edited on 25-11-2013 by AJKOER]
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bfesser
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And <em>K</em> is?
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woelen
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Sodium carbonate only is a weak base, so the concentration of hydroxide ion will be low. I can imagine though that boiling the solution does dissolve
the Al. By taking away hydroxide, the equilibrium is drawn to the right, but I also can imagine that the reaction becomes more and more difficult as
more Al has dissolved, due to formation of more and more insoluble Al(OH)3. With NaOH instead of Na2CO3 you will not have the complication of
formation of Al(OH)3, because excess NaOH will dissolve Al(OH)3.
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blogfast25
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It’s very misleading to write it that way: the sodium ions play no part. And there's a sodium too many on the right.
The hydrolysis (1st step) of carbonate anions is:
CO<sub>3</sub><sup>2-</sup> + H<sub>2</sub>O < === > HCO<sub>3</sub><sup>-</sup> +
OH<sup>-</sup>
For the second step, substitute the carbonate by bicarbonate.
The pKa (or pKb) of both steps you’ll find in Wiki.
Strong soda solutions contain enough OH<sup>-</sup> to attack aluminium and form aluminate (the [OH<sup>-</sup>] can be
estimated easily, even though it doesn't depend on carbonate concentration much). This is well known and hardly controversial.
Reaction 4 is another clumsy way of writing the hydrolysis equilibrium of aluminate. That paper, BTW, is a real hoot. No wonder you get confused: you
don’t understand the basics but want to get knowledge from obscure (and largely incorrect) sources like that one.
G-d help you. Because I doubt if science will…
[Edited on 25-11-2013 by blogfast25]
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AJKOER
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OK, a few points.
First, my computer crashed and I am sharing another machine (my response time may be longer than usual). I also have some better references, but
everything have to be rediscovered. However, my last link is an evident mistake augmented by my being in a hurry as well for which I apology.
Now, I have since added a reference, and it is interesting in that it suggests that starting with aqueous NaOH is not necessarily without danger of
forming Al2O3.3H2O if the reaction temperature is over 50 C. To quote from above:
"Another source (http://eng.sut.ac.th/metal/images/stories/pdf/02_Aluminium%2... ) specifies the decomposition of NaAlO2 at around 50 C per the reaction:
2 NaAlO2 + 3 H2O --> 2 NaOH + Al2O3. 3H2O"
-----------------------------------
On the question of K, I will comment that one normally sees the reaction written in the reverse order. Change of reaction conditions in aqueous Sodium
aluminate in the presence of Aluminum (removing any formed NaOH) appears to change reaction direction. Basically, aqueous Sodium aluminate attacks
aluminum as confirmed per this source (http://www.usalco.com/wp-content/uploads/2013/05/Sodium_Alum... ), page 9, to quote:
"Non-ferrous metals, such as aluminum, copper and brass, will deteriorate rapidly when exposed to Liquid Sodium
Aluminate. "
----------------------------------
Last comment, it is one cool reaction (please try it). After all, when does one generate Hydrogen gas and dissolve a metal relatively quickly with
safe household ingredients!
[Edited on 25-11-2013 by AJKOER]
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AJKOER
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| Quote: Originally posted by woelen | | Sodium carbonate only is a weak base, so the concentration of hydroxide ion will be low. I can imagine though that boiling the solution does dissolve
the Al. By taking away hydroxide, the equilibrium is drawn to the right, but I also can imagine that the reaction becomes more and more difficult as
more Al has dissolved, due to formation of more and more insoluble Al(OH)3. With NaOH instead of Na2CO3 you will not have the complication of
formation of Al(OH)3, because excess NaOH will dissolve Al(OH)3. |
Yes, you comments concur with my observations that removing heat, the reaction tappers down. However, I did at this point given the limited amount of
solution remaining, add some distilled water (which makes a precipitate, this event is discussed in the literature).
Then, upon further heating, the hydrolysis of the NaAlO2 in the presence of Al fuels a rapid start-up of an additional reaction creating primarily
Al2O3. 3 H2O. This reaction is conditional on there being an excess of Aluminum.
Also, your comment on employing an excess of NaOH is discussed in Patent 2734796 (link: http://www.google.com/patents/US2734796 ), to quote:
"it is known that these solutions of sodium aluminate may be made more stable and the tendency of the aluminum hydroxide to precipitate may be
suppressed or prevented by the presence of a large excess of caustic over that theoretically required to react with the aluminum trihydrate to produce
sodium aluminate."
which suggests to me the need for a large excess of NaOH, which may be a factor when considering the economics of this path.
As such, it is possible that someone with a lot of scrap Al and limited amount of NaOH, may find this path to an Aluminum salt more appealing.
[Edited on 26-11-2013 by AJKOER]
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bfesser
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| Quote: Originally posted by AJKOER | | After all, when does one generate Hydrogen gas and dissolve a metal relatively quickly with safe household ingredients! |
Any time an idiot puts aluminium utensils, pots, and pans in an automatic dishwasher. Explaining this to people has been a great
source of frustration for me, over the years.
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AJKOER
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| Quote: Originally posted by bfesser | | Quote: Originally posted by AJKOER | | After all, when does one generate Hydrogen gas and dissolve a metal relatively quickly with safe household ingredients! |
Any time an idiot puts aluminium utensils, pots, and pans in an automatic dishwasher. Explaining this to people has been a great
source of frustration for me, over the years. |
I guess less likely, more but dramatic, the same person smoking at the time of opening the dishwasher, may come to this forum to ask why his
dishwasher apparently exploded (I am referring to the detonation hazard of any formed Hydrogen gas).
Perhaps not always a 'cool' experiment at that.
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blogfast25
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“Then, upon further heating, the hydrolysis of the NaAlO2 in the presence of Al fuels a rapid start-up of an additional reaction creating
primarily Al2O3. 3 H2O. This reaction is conditional on there being an excess of Aluminum.”
Utter tosh. Gobbledigook.
”"Another source (http://eng.sut.ac.th/metal/images/stories/pdf/02_Aluminium%2... ) specifies the decomposition of NaAlO2 at around 50 C per the reaction:
2 NaAlO2 + 3 H2O --> 2 NaOH + Al2O3. 3H2O"”
Again, this is no more than the consequence of the aluminate equilibrium:
Al(OH)<sub>3</sub> + OH<sup>-</sup> ↔ Al(OH)<sub>4</sub><sup>-</sup>
Dilute an aluminate solution sufficiently and the equilibrium shifts to the left. Dilution reduces [OH<sup>-</sup>] and dictates that.
Practical experience shows that aluminate is difficult to keep in solution below pH 11.
If alkaline enough aluminate solutions do not shed Al(OH)3 on heating alone. That’s based on my very real and repeated experience.
”Basically, aqueous Sodium aluminate attacks aluminum as confirmed per this source (http://www.usalco.com/wp-content/uploads/2013/05/Sodium_Alum... ), page 9, to quote:
"Non-ferrous metals, such as aluminum, copper and brass, will deteriorate rapidly when exposed to Liquid Sodium
Aluminate. "”
Yawn. Of course, by definition aluminate solutions are strongly alkaline!
As regards dish washing machines exploding due to hydrogen evolution from Al kitchenware, I’d really like to see one single credible report on this.
‘Theoretically’ this is possible, in reality it’s extremely unlikely. Modern washing up liquids aren’t near as alkaline as good old soaps or
washing soda. And for an explosion to occur you would still need to reach the lower of the explosivity limits for hydrogen + air, smoking or not.
I also conducted my own mini-experiment. About 3 g of sodium carbonate (anh.) was dissolved in about 10 g of water, at RT. To the solution in a test
tube was added a strip (about 5 cm by 1 cm) of thin gauge, ‘no frills’ Al kitchen foil (it weighed less than 100 mg). After half an hour bubbles
of hydrogen could be clearly seen. The rate of reaction was slow but very perceptible. After some two hours the formed precipitate of
Al(OH)<sub>3</sub> became also visible:
This is what it looks like nearly 24 hours later:
No heat was applied or observed to evolve during the entire experiment. The tube is cool to the touch. So much for: “What happens is
nothing, until the solution is heated to boiling.”…
[Edited on 26-11-2013 by blogfast25]
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bfesser
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Washing up liquids?
I was talking about washing soda. If you look at the ingredients for an automatic dishwasher pellet or powder, it's primarily
Na<sub>2</sub>CO<sub>3</sub>. While I would agree that there's zero chance of explosion—that's just nonsense
AJKOER made up—the reaction certainly occurs.
Long, boring, and slightly off-topic explanation:
I worked at a grocery store, and in addition to my regular boring stock duties, I'd often help out with maintenance. In the deli/bakery kitchen, they
had two commercial automatic dishwashers. The principles of operation are identical to that in a household automatic dishwasher; react and blast the
food residue away with hot aqueous sodium carbonate (saponifies lipids) spraying at high pressure out of nozzles in a spinning bar. The major
differences are the size and materials chosen for construction. A commercial dishwasher is made almost exclusively of food grade stainless steel,
which is practically inert to the cleaning solution.
Anyway, the commercial dishwashers I've seen always have two doors, on opposing sides. One for trays going in, one for clean stuff coming out. The
doors operate either on a linked lever system, or with a complex system of pulleys, stainless steel cable, and a counterweight. We never had any
trouble with the door-lifting levers, but the pulley system was another matter...
I'd get called over to the kitchen every so often to make repairs on the dishwasher with the complaint that the doors weren't opening together. The
problem was always the same; between one and all four of the cables had somehow come detached from the doors where they are looped and crimped to a
stainless tab on the bottom inside of the door. So, I'd spend a couple of hours getting everything re-aligned, and re-attached, in 100% humidity,
cramped quarters, and unbearable heat (had to crawl halfway inside the thing to repair it).
For a while, the actual maintenance technician for the store thought that the issue must have been the employees slamming the doors down too hard and
breaking the aluminum crimps on the cables. Then, on day, it hit me (it's hard to have rational thoughts when you're inside a dishwasher or covered
in grease and much afterward). The aluminum crimps were dissolving in a reaction with the dishwashing powder! I ran over to the kitchen, and asked
to see a jar of the stuff, and sure enough, Na<sub>2</sub>CO<sub>3</sub> was listed right on the label as the major
ingredient.
It took a while, but I finally managed to convince the real maintenance guy that he needed to order stainless clamps to replace the Al crimps, because
they were dissolving. He finally relented, when I found and showed him a tiny piece of what was left of a mostly dissolved one that I had pulled out
of the drain strainer. I never had to repair the thing again, after the aluminum parts were replaced with stainless steel. And that is why I get so
grumpy when people don't accept that Al dissolves in aqueous Na<sub>2</sub>CO<sub>3</sub>.
tl;dr: Chemistry is <em>sometimes</em> the answer.
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AJKOER
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Thanks Blogfast for performing part of the reaction. Nice pictures!
However, you didn't perform what I feel are some of the most interesting reactions.
Namely, after boiling away the solution with an excess of Aluminum, stop applying heat and note the continuation of the reaction till cessation. This
suggests the exhaustion of the NaOH (from the hydrolysis of Na2CO3) with the solution now primarily consisting of aqueous NaAl(OH)4 (and NaHCO3).
Reaction:
2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2 (g)
Now, add distilled water to the concentrated remaining solution. Observe the formation of a white precipitate. Reaction based on the disruption of a
solubility equilibrium:
2 NaAl(OH)4 ---> 2 NaOH + 2 Al(OH)3↓
And last but what I find most interesting, reheat and observe the vigor of the reaction and compare to the initial one. The last reaction is, in my
opinion, per above:
2 NaOH + 2 Al + 6 H2O → 2 NaAl(OH)4 + 3 H2 (g)
with perhaps more of:
2 NaAl(OH)4 --Heat--> 2 NaOH + 2 Al(OH)3↓
but, I may be off. I will repeat and post further comments.
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blogfast25
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bfesser:
I've no trouble accepting that story. Al and hot alkali are a bad combination. Fabricators of parts for machines processing alkaline liquids should
definitely know better than to use Al parts.
| Quote: Originally posted by AJKOER | Namely, after boiling away the solution with an excess of Aluminum, stop applying heat and note the continuation of the reaction till cessation. This
suggests the exhaustion of the NaOH (from the hydrolysis of Na2CO3) with the solution now primarily consisting of aqueous NaAl(OH)4 (and NaHCO3).
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It does 'suggest' that but suggestion isn't good enough. For instance in my experiment I'll never reach the point where all carbonate is converted to
bicarbonate, simply because there isn't enough Al for that to happen.
If you want that to happen you need to set up the experiment that way: moles of carbonate < moles of aluminium. Are you sure these were your
conditions?
Also, bicarbonate is alkaline too, so the reaction would in any case continue albeit probably at a lower rate because of the lower alkalinity of the
bicarbonate.
Then you write with ill-deserved confidence:
... when my cold experiment very predictably shows Al(OH)3 precipitation. I'm working on a simple model that might apply in that narrow region where
Al(OH)3 and Al(OH)4(-) might coexist.
Time allowing I'll run a 'hot' experiment tonight.
[Edited on 27-11-2013 by blogfast25]
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