The Purification of NaOH
The Purification of NaOH
An Interesting diversion with some very ancient ‘crystal’ Drano
Recently I found a can of very ancient Drano drain clearer lurking in the depths of the garage, where it must have been for over 15 years. It was in
sad shape. The seam had split; festoons of feathery sodium carbonate decorated the side, and the can was nearly rusted through.
Now anyone in their right mind would have chucked the thing in the trashcan. But, being somewhat environmentally conscious, I hesitated. Then, seized
by a fit of science madness, I decided to examine the contents. After extracting them – no mean task – I was presented with a loathsome
putrescent greyish mud-like substance, liberally sprinkled and contaminated with flakes of rust. The outlines of the original prills floated in this
mess.
Anyone in their right mind would have diluted this liberally with water and flushed it down the toilet. But the madness was upon me. Could I extract
the remaining NaOH from this mess with a reasonable, say technical, degree of purity? I had reason to believe that I had used it as NaOH substitute
with #2 son during the course of his scanty chemical education (the school never seemed to provide any).
But on to the chemistry, lest this seem mere Whimsy.
The latest MSDS claims that Drano contains ~30% NaNO3. I cannot think of a single good reason for this. An earlier one contained Al turnings, which
sounds idiotic and an invitation to k3wls to make Coke bottle ‘bombs’. However, this ancient sample was originally colored green for some reason,
IIRC. Also I believe it was ~ 90% NaOH – once it behaved as such. I had assumed the green color was an organic dye, a sort of commercial attempt to
make the product look more effective or some such rot.
So I did the following:
(1) Dissolved what would dissolve in the minimum amount of water. The liquid got hot as it ought if substantial NaOH were present. Rust covered the
surface.
(2) Filtered through glass wool, a slow process. Added a little water if prills were still present. Most of the rust was trapped but some got through.
Refiltered through the same filter twice more. Allowed liquid to settle out rust and decanted.
(3) I was surprised to get a deep blue liquid. I still assumed this was a dye, and took a small amount and attempted to oxidize this by electrolysis,
small Fe cathode and much larger SS anode.
(4)There was a liberal release of oxygen at the anode but little hydrogen at the cathode which was puzzling. After a time the iron became covered with
a brown deposit. This proved to be copper, as tested by a flame test moistened with HCl (you need Cl ions to get the blue Cu flame – this is much
more sensitive than the green Cu flame color).
(5) What the hell was copper doing in this product? A sniff at the electrolyte told me – ammonia! So the blue coloration was a tetraamminecopper(II)
ion salt or the hydroxide of [Cu(NH3)4(H2O)2]++: Schweitzer’s reagent, which dissolves cellulose. That made a little sense; drain clogs of
vegetable fiber are quite common.
(6) The density of the blue liquid was about 1.35, equivalent to ~32% NaOH. At this level the solubility of Na2CO3 in NaOH/H2O solution is about 6.4
g/kg: See
http://www.solvaychemicals.com/Chemicals%20Literature%20Docu...
(I attempted to upload this to Scimadness but the ftp:// would not accept my password for some reason).
(7) Next: to destroy the tetraamminecopper(II) ion. Electrolysis is tediously slow. Heating to around 160C will do it. Hence, evaporate the 32%
solution, BP ~ 119C until the temperature reaches around 160C, when the NaOH concentration is ~ 60%.
To do this use a SS vessel, Fe or Ni. Glass, alumina, ceramic are all attacked by conc. hot NaOH. And be very careful. Hot NaOH is very
corrosive to skin! Eye/face protection, overalls, gloves, etc. Be aware that NaOH solutions sometime bump nastily when heated
(8) Heat further to around 200C (this heating can be done in either an oven or over a electric ring or propane flame. Don’t use a MW oven – metal
container!). On cooling a solid product, with light blue colorations, was obtained.
To get rid of the pesky copper, dissolve this impure NaOH in water to a concentration of about 45%w/w, let cool, filter rapidly through glass wool
(you’ll probably pick up some silicate). The Cu(OH)2 is precipitated. According to Solvay per above you should also have less than 1g/kg of Na2CO3
at this point.
(9) Repeat step 8, going all the way to red heat if you want. Pure NaOH melts at 323C. I don’t think one can ever get all the water out.
Product, slightly yellow hard solid (Fe?). Break up and bottle quickly!
...........
Notes: Sodium chloride and sulphate contamination should behave similarly. Sodium nitrate will not, AFAIK – it is far more soluble. This method will
not work with KOH contaminated by K2CO3 because of the extreme solubility of the carbonate
.........
Concurrently with this experiment I also concentrated a 5% solution made by electrolysis some time ago. Freezing to can take it to 8—16% but
supersaturation always occurs – add a small ice crystal and the ice will rapidly form. The eutectic forms at 19% and -28C. Straight evaporation is
best beyond 10%, then follow the above where applicable.
Regards, Der Alte
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