| Pages:
1
2
3
4 |
497
National Hazard
Posts: 778
Registered: 6-10-2007
Member Is Offline
Mood: HSbF6
|
|
how exactly would one do that?
[Edited on 14-10-2007 by 497]
|
|
|
UnintentionalChaos
International Hazard
Posts: 1454
Registered: 9-12-2006
Location: Mars
Member Is Offline
Mood: Nucleophilic
|
|
Problem there is that real tiny traces of sodium will discolor the flame if it is aluminum hydroxide....If any aluminate is present, washing it until
there would be no sodium carbonate traces left will convert all of it to aluminum hydroxide and makes a flame test useless.
[Edited on 10-14-07 by UnintentionalChaos]
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
|
|
|
Antwain
Hazard to Others
Posts: 252
Registered: 21-7-2007
Location: Australia
Member Is Offline
Mood: Supersaturated
|
|
You could titrate it against an acid. I can't find anything which I could use to tell you the pH of aluminium hydroxide, but from memory the hydroxide
is not stable (will form aluminates or aluminium salts) outside the range ~4.5-9.5, while the aluminate is "strongly alkaline"
In one sense it is a moot point, since if you take pure hydrated aluminium hydroxide [Al(OH)3(H2O)3]3- and start adding OH- ions they will start to
replace the water ligands sequentially, and conversely if you add H3O+ ions the hydroxide ligands will be replaced by water ligands. For this reason I
have always found aluminium compounds to be aesthetically displeasing. Obtaining a 'pure' hydrated aluminium compound is damn near impossible. The
alums are the only aluminium species I know of that crystalise even remotely decently.
And what the fuck is with the spell checker thinking that "aluminium" is "aluminum". I mean, sure, the americans can't pronounce it properly, but even
they know how to spell it last time I checked. No offense to any americans here 
|
|
|
497
National Hazard
Posts: 778
Registered: 6-10-2007
Member Is Offline
Mood: HSbF6
|
|
hah well... i find that whole naming thing funny....
"By 1812, Davy had settled on aluminum, which, as other sources note, matches its Latin root. He wrote in the journal Chemical Philosophy: "As yet
Aluminum has not been obtained in a perfectly free state." But the same year, an anonymous contributor to the Quarterly Review, a British
political-literary journal, objected to aluminum and proposed the name aluminium, "for so we shall take the liberty of writing the word, in preference
to aluminum, which has a less classical sound.""
while you're at it why dont you change platinum to platinium and molybdenum to molybdenium??? 
so... no definitive way to identify Al hydroxide? i cant say i could disagree with you about the aesthetically displeasing aspect of it. these damn
alumiNUM compounds really are beginning to annoy me..
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
It's likely Al(OH)3. The sodium aluminates are only stable under very alkaline conditions, CO2 will split them up. The freshly precipitate hydroxide
is highly absorbent, both compounds that interact with the -OH groups such as dyes and metal ions such as Na will be absorbed onto the precipitate;
and it has a high surface area. This means that simple tests, such as checking for sodium, are just showing that something was absorbed in a
mechanical and/or hydrogen bonding fashion. That also means that the white colour of Al(OH)3 may be modified by anything coloured in the reaction mix
Because Al compounds hydrolyse so readily you can assume that any white precipitate is going to be either a basic salt or Al(OH)3, if in moderately
alkaline conditions then Al(OH)3 is the choice.
|
|
|
497
National Hazard
Posts: 778
Registered: 6-10-2007
Member Is Offline
Mood: HSbF6
|
|
thanks! that clears things up. i just did a larger scale measured test. more on that later, but your discription of Al(OH)3 fits. grey clumpy ppt,
absorbent. it doesnt dry well at all... 300F in the oven for an hour and it was still a little damp...
|
|
|
chemkid
Hazard to Others
Posts: 269
Registered: 5-4-2007
Location: Suburban Hell
Member Is Offline
Mood: polarized
|
|
Aluminium and HCl gray insoluble precipitant
I reacted aluminium foil with muriatic acid until no more aluminium would react. The reaction formed a gray foam of bubbles. I left the beaker out
overnight with a watch glass over it to allow any remaining aluminium flakes to react. In the morning i found there was still a gray precipitant. It
measured a very low pH (cabbage indicator solution and pH paper). The precipitant did not dissolve with water or HCl. The precipitant seemed to float
on the solution but i am not sure of that. Filtration yielded a gray filtrate.
To the best of my knowledge aluminium chloride is yellow and soluble in water. I don't think it was aluminium because that would have reacted upon the
addition of more acid. Perhaps an impurity in the metal. Any insight on what the precipitant may be and if this is normal would be appreciated.
Furthermore , why does the reaction of aluminium and HCl turn gray?
Sorry polverone if the intent wasn't for me just to post over here, but this thread isn't terribly helpful considering it discusses the reaction of
aluminium with NaOH and sodium carbonate. Furthermore, even the information presented about aluminium hydroxide does nothing to explain why that would
form over aluminium chloride. Are you proposing that this mysterious precipitant may also be aluminium hydroxide?
Chemkid
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Commercial metals are not analytically pure or anything. The gray stuff is most likely silicon and other impurities, present in the 1-2% range.
Tim
|
|
|
Antwain
Hazard to Others
Posts: 252
Registered: 21-7-2007
Location: Australia
Member Is Offline
Mood: Supersaturated
|
|
As 12AX7 says, impurities. Or as I call it, "crap". You always get that from Aluminium foil of any brand. I think if you filter it properly you will
find that the particles are just really small, and that the distillate is actually clear. Best way to do it is actually to let it settle, maybe for
days, in a beaker and then pour off as much solution as you can without disturbing the 'mud' at the bottom.
If its aluminium chloride you are after then good luck, I mean it! I was never able to make this to my exacting standards. As it concentrates, it
hydrolises- which both gasses you with HCl and leads to a product which is not AlCl3. If you come up with a way to work that out I would be interested
to know.
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
| Quote: | | The most common foil alloys - the 1000, 3000 and 8000 series - contain between 0.5% and 1.5% iron, 0.1% and 0.7% silicon and 0.02% to 1.5% manganese.
Up to 0.2% copper may be added when additional strength is required. |
http://www.azom.com/details.asp?ArticleID=1434
1000 series alloys are nearly pure aluminium, around 99% and in some cases 99,5 or better; the digits following the 1 indicate the fractional purity
above 99%, 1030 is +99,3% Al. The 3000 series has manganese as its main alloying element, while the 8000 series uses lithium and is fairly rare.
Almost all aluminium (alumium for you Davy fans) alloys have a few tenths of a percent each of silicon and iron.
Contrary to a popular belief, cooking foil is generally not coated with any polymer layer. Even the fluorocarbon plastics break down some at the
temperatures the foil is exposed to during use, most other plastics go sooner; in any case not something you want coating your food. The differing
appearance of the two sides of foil is due to the manufacturing process, the differences in forces applied and rolling/sliding contact with the
equipment.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
I believe the "nonstick" foil is anodized, FWIW.
Tim
|
|
|
not_important
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
Reynolds makes a non-stick foil that has a silicone release agent on the matte side (USP 6696511 - had to look that up out of curiosity) that was put
on the market in 2002 or 2003. The coating can cause problems in recycling and using methane from landfills, so some areas are considering banning
such products. Higher temperatures do cause off-gassing, using such foil as heat containment around glassware can cause your lab to acquire a thin
coating of silicones.
Any additional processing adds to the cost, cheap generic food foil is likely to be just plain Al and thus the best for most uses in the home lab.
Foil intended for construction or industrial applications has a better chance of not being a 1000 series alloy, but also is more
likely to have information on which alloy it is made of.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
| Quote: | | To the best of my knowledge aluminium chloride is yellow and soluble in water. I don't think it was aluminium because that would have reacted upon the
addition of more acid. Perhaps an impurity in the metal. Any insight on what the precipitant may be and if this is normal would be appreciated.
Furthermore , why does the reaction of aluminium and HCl turn gray? |
AlCl3 is a white powder, the hydrated salt AlCl3.6H2O also is white, somewhat transparent, crystalline. A solution of AlCl3 in hydrochloric acid is
clear and colorless, not yellow.
As others have stated, the grey stuff is crap, impurity in the metal foil. I have reagent grade Al-needles and dissolving these in acid give perfectly
clear and colorless solutions. I also have pyrotechnics grade Al-powder, and that gives dark grey crap as well (carbon??).
Making anhydrous AlCl3 is not only difficult, as Antwain writes, it is near impossible for a moderately equipped homelab. Anhydrous AlCl3 is nasty
stuff. It heavily fumes in air, and is very corrosive. Not something you can make at home.
Hydrated AlCl3 on the other hand should not be that difficult. If you dissolve aluminium in a large excess amount of 10% HCl (use colorless acid, not
the yellow stuff sold as muriatic acid) and let the grey stuff settle (which may take a long time, at least several hours), then the clear liquid can
be allowed to evaporate in a petri dish. Don't heat it, just put it in a nice dry and warm place (e.g. above a heat radiator, with a paper tissue
loosely covering the dish in order to prevent dust from entering it). After a few days you should have crystals of solid AlCl3.6H2O.
|
|
|
chemkid
Hazard to Others
Posts: 269
Registered: 5-4-2007
Location: Suburban Hell
Member Is Offline
Mood: polarized
|
|
That sounds excellent woelen. I think i will try that with a desiccator. (i have found that a simple coffee can and calcium chloride is extremely
effective). Currently i am making aluminium hydroxide though. I have been able to filter the impurities so i will continue on that route. Perhaps i
will acquire some high purity aluminium as well.
Thank you.
|
|
|
Antwain
Hazard to Others
Posts: 252
Registered: 21-7-2007
Location: Australia
Member Is Offline
Mood: Supersaturated
|
|
A much cheaper way of doing that would be to buy alum from a garden shop and hit it with just about any base, say sodium carbonate or hydroxide
|
|
|
chemkid
Hazard to Others
Posts: 269
Registered: 5-4-2007
Location: Suburban Hell
Member Is Offline
Mood: polarized
|
|
Are your refering to aluminum sodium sulfate or one of the many other "Alums"?
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
The most common alum is KAl(SO4)2.12H2O, plain potassium alum.
Aluminium ion gives a hydroxide on addition of sodium carbonate. Aluminium carbonate cannot be precipitated from water. It is hydrolysed at once.
When you use sodium hydroxide, then you must adjust the amount carefully. If you add too much hydroxide, then the aluminium hydroxide redissolves
again, giving aluminates. Aluminiumhydroxide is amphoteric and can act as base, and as acid. On addition of excess sodium hydroxide it acts as acid.
When you try to evaporate a solution of an aluminate, then, however, it hydrolyses again, giving aluminium hydroxide.
|
|
|
kclo4
National Hazard
Posts: 916
Registered: 11-12-2004
Location:
Member Is Offline
Mood: No Mood
|
|
Hmm, sorry if i missed a conclusion on this already, but it is only releasing H2 and NOT CO2. Because it produces Sodium Bicarbonate 
Right?
|
|
|
chemkid
Hazard to Others
Posts: 269
Registered: 5-4-2007
Location: Suburban Hell
Member Is Offline
Mood: polarized
|
|
what is "it"? (if your talking about the sodium carbonate and water discussion a while back, i was totally confused and sodium carbonate does not
react with water)
Chemkid
|
|
|
kclo4
National Hazard
Posts: 916
Registered: 11-12-2004
Location:
Member Is Offline
Mood: No Mood
|
|
Yeah, earlier.
The reaction between sodium carbonate, water, and Al does not make CO2
it makes NaHCO3 instead.
|
|
|
497
National Hazard
Posts: 778
Registered: 6-10-2007
Member Is Offline
Mood: HSbF6
|
|
im pretty sure it ended up with the same amount of Na2CO3 as it started with. so no it doesnt make CO2 (or very much anyway). but it'd be easy to test
if you wanted to know for sure.
[Edited on 24-10-2007 by 497]
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Al, Na2CO3 and water will produce H2 and Al(OH)3, the net reaction being
2Al + 6H2O --> 2Al(OH)3 + 3H2
The carbonate only is used catalytically, it is required to make some OH(-) by means of hydrolysis, which in turn causes the Al to form H2 and
hydroxide.
This reaction, however, will be painfully slow. Maybe a lot of heating of Al in a concentrated solution of Na2CO3 will give a somewhat higher reaction
rate.
|
|
|
497
National Hazard
Posts: 778
Registered: 6-10-2007
Member Is Offline
Mood: HSbF6
|
|
oh yes its quite slow at room temp. but at 95C it goes quite nicely, ~6g Al reacted (not quite completely) in an hour or two. concentration of Na2CO3
was about 65g in 1 liter. i think the concentration could be higher and would probably go even faster.
[Edited on 25-10-2007 by 497]
|
|
|
ShadowWarrior4444
Hazard to Others
Posts: 226
Registered: 25-4-2008
Member Is Offline
Mood: Sunlight on a pure white wall.
|
|
I have recently been testing various gas generation mixtures, and came across something that I am not entirely sure of.
The mixture was simply Draino Kitchen Crystals and aluminum foil, with a bit of water. Very crude, however the reaction vessel and subsequent gas
containment and piping were rigorously air-tight.
Draino adds NaNO3 to their product to produce NH3 in normal operation, however I have found that the gas produced from the reaction with foil
contained very little NH3. Only on standing do the reaction byproducts emit a detectable amount of ammonia.
The main observation of interest is this, though:
16 hours after the completion of the reaction a white precipitate with quite a bit of volume formed from the previously clear liquid, (discounting the
grey iron/silicon impurities.) I suspect this white precipitate to be aluminum hydroxide, so then what does it form from? The hydrolysis of sodium
aluminate? The clear solution remained heavily basic throughout, and no evaporation took place.
|
|
|
UnintentionalChaos
International Hazard
Posts: 1454
Registered: 9-12-2006
Location: Mars
Member Is Offline
Mood: Nucleophilic
|
|
It forms from the absorbtion of carbon dioxide by the solution. Carbonic acid being a stronger acid than aluminum hydroxide displaces it from sodium
aluminate forming sodium carbonate and the aluminum hydroxide precipitates.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
|
|
|
| Pages:
1
2
3
4 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
