| Pages:
1
2
3
4 |
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Neil:
With respect to a possibly easier way to make AlN, see the extract below from the US Patent 5226952 which apparently employs a replacement reaction
with calcium nitride and Al (and also a CaAl2 impurity) that forms AlN and pure Ca, the latter being the apparent goal of the process. Note, the
formation range of 700 to 1200 C.
"United States Patent 5226952
A process for refining calcium containing aluminum as an impurity in the form CaAl2. The calcium is nitrided to form Ca3N2, and that Ca3N2 is made to
react with CaAl2 initially present and with the particulate aluminum in such a manner as to form calcium, which is isolatable in the form of
high-purity calcium containing less than 0.1% by weight of aluminum, and aluminum nitride.
7. The process of any one of claims 1-3, in which the calcium is nitrided with the aid of nitrogen at a temperature of between 200° and 350° C.
9. The process of one of claims 1-3, in which the calcium nitride formed is made to react with the particulate aluminum introduced and with CaAl2
initially present in the calcium, at a temperature of between 700° and 1200° C., and the calcium formed is separated by distillation.
11. The process of one of claims 1-3, in which a calcium to be purified, containing n mols of calcium and m moles of CaAl2, is nitrided in such a
manner as to form n/3 mols of Ca3 N2, which is made to react with m mols of CaAl2 and (n-3m) 2/3mols of aluminum, so that aside from losses,
n+m mols of purified calcium and 2n/3 mols of AlN are obtained. "
You may be able to purchase Ca3N2, and then react it with particulate Al between 700 to 1200 C, if able and so inclined. Note, there is a parallel
employment of Calcium carbide to form carbides of other metals.
[Edited on 26-6-2011 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
blogfast25
Please read my post again.
I am not dictating any aluminate formation (perhaps you missed the word "not"), but I am suggesting an aqueous alkaline catalyst to the reaction of
Al2O3 and water (with parallels to AlN hydolysis). The fact that NH4OH is predominately NH3 and most importantly H2O makes it a good "aqueous alkalie
solution".
I also noted that this was not an original concept on my part.
I also found a similar comment on a chart on page 9 of a report entitled "Aluminum Compounds Review of Toxicological Literature Abridged Final
Report", prepared by Integrated Laboratory Systems, namely Al2O3 is "slowly soluble in aqueous alkaline solutions" given your concerns on the quality
of some MSDS statements.
Link:
http://www.scribd.com/doc/2895150/Aluminum
[Edited on 26-6-2011 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
My concern isn't with the 'quality of some MSDS statements' (although many are pretty crap), but with the fact that MSDS sheets aren't scientific
literature, rarely specify any conditions etc. They can be OK to glean the odd specific property of a compound or product from but not much besides
that. It's not really their function to provide precise chemical information.
How do solutions of unprotonated NH3 make "good aqueous alkaline solution"?
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
How do solutions of unprotonated NH3 make "good aqueous alkaline solution"?
I would guess pretty much the same way that it acts on AlN to initiate the hydrolysis action.
As far as I know, more details on the precise catalytic mechanism may, or may not, actually be known.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
LINK to Full Paper on "The Precipitation of Aluminium Hydrous Oxide and Its Solubility in Ammonia" by Prideaux and Henness
https://docs.google.com/viewer?a=v&q=cache:ydMv90ZG1p8J:...
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
If it’s unprotonated, it’s not an alkali: for something to qualify as alkali OH- ions must be present.
The paper is very interesting but glaringly makes my case for me: the amounts of dissolved/peptised/supersaturated hydrated alumina are very small and
tail off with increased ammonia concentration, despite the fact that the OH- increases slowly with increasing NH3 concentration.
It confirms that using NH3 solution to precipitate hydrated alumina from Al<sup>3+</sup> bearing solutions is good practice and easier
than using strong alkalis, where overshooting the required OH- concentration is often a problem.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is an explanation why NH4OH is the proper formula for ammonia in water:
"nitrogen likes to form 3 bonds to complete it valence shell, but it also has that lone pair of unbonded electrons hovering in that p orbital. In a
water environment, nitrogen's electron orbital will actually dissociate water into H+ and OH-. The H+ is of course attracted to nirtogen's negatively
charged electron cloud and it forms the ammonium ion, NH4+, which ionically bonds to the remaining OH- hydroxide ion from the water"
source: Kayak on Ask.com
|
|
|
barley81
Hazard to Others
Posts: 481
Registered: 9-5-2011
Member Is Offline
Mood: No Mood
|
|
Yes, but in ammonia solutions, only a small amount of NH4+ is formed. In household ammonia, the pH is around 11-12 (~0.01M OH-). The great majority of
ammonia exists as NH3.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by AJKOER  | Here is an explanation why NH4OH is the proper formula for ammonia in water:
"nitrogen likes to form 3 bonds to complete it valence shell, but it also has that lone pair of unbonded electrons hovering in that p orbital. In a
water environment, nitrogen's electron orbital will actually dissociate water into H+ and OH-. The H+ is of course attracted to nirtogen's negatively
charged electron cloud and it forms the ammonium ion, NH4+, which ionically bonds to the remaining OH- hydroxide ion from the water"
source: Kayak on Ask.com |
Yeah, never forget to ask ‘kayak’ from ask.com, he’ll know!
It just shows how little you know about chemistry and how you’re scraping the barrel.
‘NH4OH’ is a very old way of notation of NH3 solution, dating back to before Bronsted acid-base theory.
NH3 does have an extra electron pair which it uses to bond protons (from water) with (that makes it also an electron donor). The reaction is thus as
indicated ad nauseam above:
NH3(aq) + H2O(l) < === > NH4(+)(aq) + OH-(aq)
… which proceeds only feebly (the product of [NH4(+)] and [OH-], divided by [NH3] is only about 1.78 x 10<sup>-5</sup>
What little NH4(+) forms is completely dissociated from OH-. There never was any ‘NH4OH’ to begin with but once upon a time its believed existence
was a reasonable explanation for ammonia’s weak alkalinity.
Sigh… 
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
"Ammonia is not a base when using the more widely known Arrhenius definition (which states that a base releases OH- and an acid releases H+. Instead,
it is considered a base when using the Lewis definition, which states that a base is an electron pair donor, and an acid is an electron pair acceptor.
The structure of ammonia, NH3, has a pair on the nitrogen atom, and it is this lone pair that acts as a electron pair donor.
Also, when mixed with water, ammonia forms ammonium hydroxide (NH4OH), which is a base according to the Arrhenius definition."
Link:
http://wiki.answers.com/Q/Why_does_NH3_give_OH-_ions_on_diss...
Also here is a reference from "Chemistry Explained", a teaching website:
Ammonia (NH 3 ) is a weak base, and although it does not have OH − ions in its formula, it produces the ion on reaction with water.
NH 3 ( aq ) + H 2 O ( l ) ⇆ NH 4 + ( aq ) + OH − ( aq ) (8)
Link:
http://www.chemistryexplained.com/A-Ar/Acid-Base-Chemistry.h...
Apparently, this view is generally held, for example:
"Similarly, ammonia dissolves in water and partly dissociates in the presence of water into ammonium ions. As we all know that water dissociates into
H(+) and OH(-). The NH3 reacts with H(+)to form NH4(+). So we can write a chemical reaction:
NH3 + H(+) + OH(-) <-----> NH4(+) + OH(-)
The above reaction is an equilibrium reaction. So the amount reacts depends on the equilibrium constant which is dependent on temperature."
Link:
http://www.cheresources.com/invision/topic/9750-ammonia-wate...
While all interesting, the argument on NH4OH is not relevant since NH4OH is generally considered a weak base and I believe this is what was precisely
being reference on the MSDS and you, I suspect, do not.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Widely held views aren’t necessarily true. Truth isn’t arrived at by majority rule.
Teaching sites simplify things accordingly. My daughter passed (I hope) her GCSE chemistry exam today: on that level they’re still writing reaction
equations in the simplest of forms. Actual chemists really only do so when writing overall reactions or for stoichiometry determination. And the
definition of ‘catalyst’ she was taught is comical to someone like me…
Water based solutions contain almost no free protons (H<sup>+</sup>:
these are heavily solvated to H<sub>3</sub>O<sup>+</sup> (hydronium or oxonium ions).
NH3 is both a Lewis AND a Bronsted base. Arrhenius acid-base theory is now largely obsolete.
The question ‘why does NH3 give OH- when dissociated in water’ is badly formulated. NH3 isn’t ‘dissociated’: it’s (weakly) protonated in
water. The every small proportion of water molecules thus deprotonated become OH- ions.
Do you think you could go and read quietly now?
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
blogfast25:
Another point is I agree with your opinion that any alleged ammonium aluminate is best described (to quote you) as "dissolved/peptised/supersaturated
hydrated alumina".
However, a subtle legalistic and possibly important point is that there is an alluded to process by which ammonium aluminate has been prepared, namely
"In this connection, it is interesting to consider the evidence presented by C. Renz (Ber., 36, III, 2751 (1903)). This author dismisses the
possibility of the existence of an ammonium aluminate, even though by an indirect method (viz., solution of Al(OH)3 in Ba(OH)2 and subsequent addition
of (NH4)2SO4) he was able to obtain a clear solution free from Ba ++ and SO4-, 50 cc. of which contained 0.1 g. Al2O3." Reference: "Journal of the
American Chemical Society", Volume 38, page 1287 (link previously supplied).
The potential legal significance of this is that to quote from US Government Patent Office ( per 2173.05(t) Chemical Formula - 2100 Patentability), "A
compound may also be claimed in terms of the process by which it is made without raising an issue of indefiniteness", where indefiniteness is a cause
to dismiss a patent claim. Interestingly, I have seen ammonium aluminate cited in a patent, possibly since one cannot claim its "indefiniteness".
Link: http://www.uspto.gov/web/offices/pac/mpep/documents/2100_217...
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Regarding AlN, Brauer’s ‘Preparative Inorganic Chemistry’ gives three methods for its preparation (p.827, available in SM’s library):
Aluminum
A1N
Nitride
= A1N
41.0
A nickel boat is filled with very pure aluminum powder which
has been degreased and dried either by extraction with ether or
by heating to 150°C in a stream of nitrogen. The boat is placed
in a quartz or porcelain tube and heated in an electric furnace
while purified nitrogen is passed over it. Even though the nitride
starts to form on the surface below 650°C, the reaction proper
begins only at 820°C, when the entire mass begins to glow. At
this point the flow of nitrogen should be increased to prevent the
Ns pressure from decreasing owing to the rapid reaction. When
the reaction is essentially complete, the mass is allowed to cool
in a stream of nitrogen. Since the product still contains some
unreacted metal, it is pulverized and reheated under nitrogen for
1-2 hours at 1100-1200°C. The product obtained is nearly white
and has a nitrogen content not far below theoretical.
II. Al + NH3 = A1N + 3/2H2
27.0 17.0 41.0
To obtain silicon-free A1N, aluminum powder pretreated as
above is placed in a trough of molybdenum sheet inside a nickel
reaction tube, and NH3 is led through while the tube is heated to
1300°C in an electric furnace.
III. AlCls • NH3 = A1N + 3 HC1
150.4 41.0 109.4
The reaction is performed in the apparatus shown in Fig. 246,
which consists essentially of a thick-wall Pyrex tube with an
enlargement in the middle and four necks at the top. A thin
glass tube (nitrogen inlet) passes through the middle neck and
reaches nearly to the bottom. The two side necks contain silver
wire leads to a tungsten heating coil w suspended in the reaction
tube at the level of the bulb. The fourth neck is an outlet for the
gas. The reaction tube is thoroughly dried and A1C13- NH3, prepared
|
|
|
Neil
National Hazard
Posts: 556
Registered: 19-3-2008
Member Is Offline
Mood: No Mood
|
|
Ah, that would make my plans a fools errand. Thank you for that 
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Welcome!
|
|
|
| Pages:
1
2
3
4 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
