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blogfast25
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Quote: Originally posted by Nerro  | Are you guys sure about the reaction between NaOH and Mg? The equation used to "prove" the thermodynamics seems rather incomplete.
The correct equation would be:
2NaOH + 2 Mg --> 2Na + 2MgO + H2
and essentially be a combination of the following 3 reactions:
2NaOH --> Na2O + H2O
Mg + H2O --> MgO + H2
Na2O + Mg --> 2Na + MgO
Also there seems to be NO mention of the fact that a reaction that is thermodynamically unfavourable cán proceed at elevated temperatures. After all,
delta(G) = delta(H) - T(delta)S
delta(H) = 349.08 kJ/mole and delta(S) = 76.5 J/K
If delta(G) is calculated for room temperature it is 326.7 kJ/mole...
It does not become negative untill a temperature of 4365 K is reached. Evidently a different mechanism is driving this reaction.
[Edited on 18-12-2010 by Nerro] |
NaOH: HoF 298 K = - 426 kJ/mol
MgO: HoF 298 K = - 601 kJ/mol
Hess Law: reaction enthalpy independent of reaction path:
For NaOH + Mg ===> Na + MgO + ½ H2:
Heat of reaction at 298 K: ΔH = + 426 – 601 = - 175 kJ/mol
(break it down as NaOH === > Na + ½ O2 + ½ H2 (ΔH = 426 kJ/mol)
And Mg + ½ O2 ===> MgO (ΔH = - 601 kJ/mol) and use Hess)
Because there is gas evolved, ΔG is probably even more negative, because of Delta S. But - 175 kJ/mol is good enough...
This reaction proceeds with dangerous levels of heat produced.
[Edited on 18-12-2010 by blogfast25]
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blogfast25
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For those still interested in using lithium in thermite-like reductions, you really need to think fluorides. For example, the very
first bench scale productions of plutonium were achieved by PuF4 + 4 Li == > Pu + 4 Li.
Slightly more realistically  , Li should be capable of reducing the trifluorides
of the Rare Earths. Here’s an example: the HoF of NdF3 (at 298 K) is – 1,661 kJ/mol:
http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...
The HoF of LiF is – 617 kJ/mol (NIST value).
NdF3 + 3 Li == > Nd + 3 LiF: enthalpy of reaction = 1661 – 3 x 617 = - 190 kJ/mol, very exothermic.
Without having made any detailed thermochemistry calcs, I predict that that’s enough heat to heat the reaction products to above their respective
MPs and thus would be obtained in the molten state, allowing the heavier metal to coalesce out, neatly protected from oxygen by the LiF ‘blanket’.
Nd, for experimenters that want to see an Li reduction in action, also has the advantages of:
1. availability: see neodymium magnets (they sell them by the tonne on eBay nowadays),
2. NdF3 is insoluble in water and thus relatively easy to synthesise.
This does mean 'messing' with the inevitably dangerous fluorides. Read up before you start!
Magnesium, BTW, should aslo do the trick...
[Edited on 20-12-2010 by blogfast25]
[Edited on 20-12-2010 by blogfast25]
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Arthur Dent
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Hi guys,
This morning, I noticed that a small quantity of metallic lithium I had put aside in a jar filled with mineral oil looked kinda funny, so I decided to
get rid of it because it started to get very dark and foamy.
Since Lithium is still quite valuable, I decided to turn it into LiOH by dropping little chunks in distilled water. The metallic lithium was dabbed
with a paper towel to remove some of the mineral oil before being dropped in the water in small 1 cm chunks, but I guess it still had a fair amount of
oil on it.
After dissolving the entire stock, I was left with a fairly concentrated solution of Lithium Hydroxide, but unfortunately it seems that the oil was
somehow emulsified along in the solution. 2 hours after the procedure, the solution is quite turbid and there is a very thin layer of even
milkier-looking oil at the surface. Shouldn't aqueous LiOH be water clear?
My next step would be to filter off the solution. So my idea was to let the solution settle and maybe it would separate a bit more somehow, then use a
fine Whatman paper filter to filter off the solution, but i'm afraid that the oil will clog the filter paper. What would be my best way to acquire a
clean solution of LiOH with this hideous brew?
Robert
==========
UPDATE: I went ahead and chucked the whole shabang in a funnel with a plain coffee filter paper. Looked like it worked. All the oliy white gunk stayed
in the filter and the solution, still turbid, looks free of contaminants.
I have stored the solution in a HDPE bottle, it should keep well? Are there risks that the LiOH absorbs CO2 and slowly precipitates to a carbonate?
Would the chloride be more stable for storage?
Robert
[Edited on 31-12-2010 by Arthur Dent]
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plante1999
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| Quote: Originally posted by Arthur Dent | Hi guys,
This morning, I noticed that a small quantity of metallic lithium I had put aside in a jar filled with mineral oil looked kinda funny, so I decided to
get rid of it because it started to get very dark and foamy.
Since Lithium is still quite valuable, I decided to turn it into LiOH by dropping little chunks in distilled water. The metallic lithium was dabbed
with a paper towel to remove some of the mineral oil before being dropped in the water in small 1 cm chunks, but I guess it still had a fair amount of
oil on it.
After dissolving the entire stock, I was left with a fairly concentrated solution of Lithium Hydroxide, but unfortunately it seems that the oil was
somehow emulsified along in the solution. 2 hours after the procedure, the solution is quite turbid and there is a very thin layer of even
milkier-looking oil at the surface. Shouldn't aqueous LiOH be water clear?
My next step would be to filter off the solution. So my idea was to let the solution settle and maybe it would separate a bit more somehow, then use a
fine Whatman paper filter to filter off the solution, but i'm afraid that the oil will clog the filter paper. What would be my best way to acquire a
clean solution of LiOH with this hideous brew?
Robert
==========
UPDATE: I went ahead and chucked the whole shabang in a funnel with a plain coffee filter paper. Looked like it worked. All the oliy white gunk stayed
in the filter and the solution, still turbid, looks free of contaminants.
I have stored the solution in a HDPE bottle, it should keep well? Are there risks that the LiOH absorbs CO2 and slowly precipitates to a carbonate?
Would the chloride be more stable for storage?
Robert
[Edited on 31-12-2010 by Arthur Dent] |
I ave already made LiOH 4month ago from very fresh lithium from battery. the solution ave some cloudiness but my solution is colorless. I keep it in
dry HDPE bottle.
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garage chemist
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LiOH solutions always get cloudy due to uptake of atmpspheric CO2, forming sparingly soluble Li2CO3.
This problem is much more severe with e.g. Ba(OH)2 solution, which is generally always cloudy except immediately after filtration.
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