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You may notice that the end-notes don't necessarily start at "1". This article is an excerpt from a back issue of our newsletter.
Uranophane Experiment -
Ammonia Precipitation
by C. Thorsten
Analysis of Uranophane using Ammonia Precipitation
Safety reminder: Uranium salts are toxic. exercise due caution when handling radioactive materials. Do not ingest or inhale the particles of uranium compounds.Introduction:
Uranium's primary hazard is not its radiation (which is mild), but rather its toxicity. It is similar to the other "heavy metals" (lead, nickel, chromium) in this respect. Grinding up a uranium mineral for fusion presents an inhalation hazard. You must wear a cartridge respirator that can filter out harmful dusts. For safety, assume that soluble heavy-metal compounds will also be absorbed through your skin; exercise appropriate precautions.
One of our readers emailed a summary of an experiment he'd done using a sample
of uranophane from Canada. The general procedure was from the first
edition of Orsino Smith's Identification
and Qualitative Analysis of Minerals (1946). The sample was
crushed and fused with sodium carbonate, and then the fusion was dissolved
in distilled water. To the resulting solution was added some solid
ammonium chloride and then aqueous ammonia (i.e., ammonium hydroxide, NH4OH).
The solution was boiled and allowed to cool. The filtrate from
this ammonia step was then treated with ammonium oxalate and filtered;
the filtrate from the oxalate step was in turn treated with ammonium phosphate.
In this article we'll explore this procedure in some detail, especially the
first step (ammonia precipitation). To follow along with our reader's
experiment, we performed the same procedure using a small bit of uranophane.
As we'll discuss later, a thorough separation of uranium from other elements
can be a very involved process. In the current experiment we'll
look at the chemistry of a highly simplified procedure.Procedure:
1.) Crushing of the sample and mixing with sodium carbonate: This was accomplished by breaking off a pinky-nail sized piece of uranophane from a rock sample and pulverizing it in a steel mortar and pestle. Approximately three times the volume of powdered sodium carbonate (relative to the sample size) was added to this.
2.) Fusion of the sample: The uranophane-sodium carbonate mixture was placed in a shallow hole that had been drilled in a carbon block. The mixture was fused with a propane torch until it melted into a single bead.
3.) Dissolution of the sample: The bead was allowed to cool and placed in 5 mL of distilled water in a micro beaker. The block was checked for remnants using a Geiger counter; considerable residue was found in this way. It was scraped off carefully and added to the beaker.
The bead and other residue were allowed to stand in the water overnight.
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Figure 7. The sodium carbonate fusion
was dissolved in water. Notice the very slight coloration of the water
solution; note also, however, the considerable amount of insoluble
matter at the bottom. Some of the black came from carbon particles
scraped from the block. When this precipitate was left behind on the filter paper, it was checked with a Geiger counter and found to contain most of the detectable radioactivity of the sample. |
The next day, the remainder of the water-logged bead was broken apart with
the end of a glass rod.
The solution was stirred and allowed to settle again for a photograph (figure
7). It was then stirred up and filtered via gravity.
This filtrate (which we'll call the "water fraction") was collected in a micro
beaker and set aside. It had only a very slight yellowish tint, barely
noticeable against a white background.3a.) Insoluble residue remained in the filter paper. This was checked with a Geiger counter and found to contain most of the sample's radioactivity. This residue was treated with 6 M HCl; the resulting filtrate (which we'll call the "HCl fraction") was collected in another micro beaker (figure 8).
A drop of this solution crystallized on a microscope slide yielded distinct
crystals of NaCl, but only shapeless, vaguely-crystalline yellow masses of
uranium salts at this point.
Figure 8. The filtrate from step 3 did not contain much uranium, judging from its radioactivity and its color.
Most of the uranium was in the "HCl fraction", the solution obtained by dissolving the Step 3 precipitate with hydrochloric acid.
The HCl solution remained transparent yellowish but acquired a gelatinous
consistency on standing. This was presumed to be silicic acid.4.) [Ammonium Chloride and] Aqueous Ammonia: As discovered in step 3, the bead contained water-insoluble matter that had to be carried into solution with HCl. Because this introduced chloride ions into solution and thus formed NH4Cl, step 4 omitted the ammonium chloride.
About 3 mL of the HCl fraction were treated with 10% aqueous ammonia (presumed
to contain some dissolved CO2). At first the solution went
from pale yellow to colorless. More ammonia was added until the solution
had a slightly ammoniacal odor. At this point the solution turned yellow
again. On standing a muddy yellow, cloudy to slightly gelatinous precipitate
began to form (figure 9).
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Figure 9. The solution from Step
4. Notice the muddy-yellow, flocculent precipitate. At
this point, the solution was presumed to contain NH4Cl as well as ammonia,
since the aqueous NH3 had been added to a solution containing HCl. The photo shows the solution before boiling. |
5.) Boiling: The micro beaker containing the ammoniated solution was placed on a wire gauze on a ringstand. Five, PTFE boiling chips were added (to prevent "bumping"). The liquid was heated from underneath with an alcohol burner. The precipitate darkened upon heating. Boiling was continued for 5 minutes. The heat was removed and the beaker allowed to cool. A bright yellow, somewhat flocculent precipitate settled to the bottom (figure 10).
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Figure 10. Now we're getting somewhere:
the solution after Step 5. The Teflon chips are still in
there. Note the color change to bright canary yellow.
The yellow is probably UO2•2H2O.
In our sample the hydrous U oxide is mixed with silica,
silicic acid, insoluble Na silicates, etc. |
6.) Filtration: The solution was poured through filter paper in a funnel. The bright yellow precipitates on the funnel were washed with distilled water. The washings were collected in the filtrate. This solution was saved for Step 7. The yellow precipitates themselves were checked with the Geiger counter; predictably, they were radioactive. These precipitates were saved for Step 10.
Figure 11. The precipitate from Step 6:
bright yellow
and detectably radioactive. The precipitate is
made bulkier by the presence of silica and / or
silicic acid, which are separated in Step 10.
and detectably radioactive. The precipitate is
made bulkier by the presence of silica and / or
silicic acid, which are separated in Step 10.
Boiling down the Step 6 filtrate to 15 mL and trying again produced a definite, white precipitate on treatment with ammonium oxalate. Allowing the test tube to stand for a week produced fairly large, colorless crystals. The crystals appeared to be finished growing after two weeks.
Figure 12. Oxalate precipitate from Step 7. The test
tube was allowed to stand for two weeks to allow
complete precipitation of insoluble oxalates. These
crystals took quite a long time to grow to the stage
shown here.
Click for full-size image.
9.) Ammonium Phosphate: The liquid from step 8 was evaporated down to about 1/4 its volume by gentle heating, allowed to cool, and treated with (NH4)2HPO4 (which can be made by treating phosphoric acid with two molar equivalents of NH4OH).
The solution was allowed to stand for a few hours.
10.) Treatment of the Precipitate from Step 6: The water-washed precipitate on the filter was dissolved by running some 6 M HCl through the filter paper. The yellow filtrate was collected in a clean beaker. Left behind was a clear, gelatinous precipitate on the filter.
Figure 14. Silica / silicic acid remaining on the filter after
dissolving the uranium precipitates with 6M HCl.
If this had been Al(OH)3 or other metal hydroxide,
it would have dissolved in the strong acid.
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Figure 15. The solution from Step
10, made by dissolving the bright yellow uranium precipitates in 6M HCl and
washing with water, followed by more HCl. This solution at this stage is strongly acidic. It will contain some NH4Cl unless the precipitate at Step 6 was washed very thoroughly and perhaps boiled in distilled water to remove adsorbed ammonium compounds. |
Some of the Step 10 solution was placed on a microscope slide and allowed to evaporate at room temperature. Square, bright-yellow crystals formed; these were thought to be ammonium uranyl chloride. There were no appreciable amounts of other crystal species present in several of these slides.
10a.) A sample of this Step 10 solution was placed in a crucible, evaporated to dryness with an alcohol lamp, and strongly heated for at least 5 minutes to drive off any ammonium salts. The residue, when cool, was redissolved in 7 M HCl. This new solution was also yellow, but this time it was impossible grow the beautiful, kaleidoscopic yellow squares.
|
Figure 16. A crystal of ammonium uranyl chloride, viewed at 400x.
These crystals displayed interesting and often remarkably beautiful patterns
within their square outlines. They were reminiscent of quilt patches
or kaleidoscope patterns. Click for full-size image. |
Discussion:
We've explored a highly simplified separation procedure for uranium which,
in the specific case of uranophane as the starting material, worked well enough
for qualitative purposes.
During the carbonate fusion, the uranophane sample was converted to a mixture
of the corresponding metal hydroxides and oxides; the fusion also contained
sodium hydroxide (from the decomposition of the carbonate), sodium
silicates, and some undecomposed sodium carbonate.
Treatment of the bead with water solubilized the Na, some of the Ca, and
the greater part of any Al and Pb that might have been present. The
bulk of the soluble silicates of Na were also dissolved at this step;
these gradually separated as insoluble silica on standing for a few days.
Some of the silicates, however, did remain undissolved in the water at step
3; these found their way into HCl solution during Step 3a.
The HCl-soluble portion of the bead (Step 3a) contained the U, the rest of
the Ca, and any Mg, as well as any other elements whose oxides or hydroxides
weren't soluble in alkaline solution. The carbon-block residue in Step
3 was presumed to contain uranium oxides, which would have formed during
intense roasting. This would explain their insolubility in the alkaline
solution resulting from the "water fraction" (Step 3).
Treatment of the HCl fraction with excess ammonia (Step 4) introduced NH4Cl
into the solution, as well as aqueous NH3. The NH4Cl
kept any Mg2+ in solution while the U2+
precipitated
at Step 5. It's important to note here that the uranium wouldn't
have
come down if there had been appreciable ammonium carbonate or vanadium
in
solution. (We can safely say, therefore, that the specimen
was indeed uranophane, not carnotite; the latter contains much V).
Boiling the ammoniated HCl fraction (Step 5) caused formation of hydrous uranium
oxide, as previously mentioned. Step 5 also brought about decomposition
of any alkali silicates that had dissolved in the HCl fraction. The
mixture of hydrous uranium oxide and silica gel / silicic acid stayed together
until Step 10, when the siliceous compounds were left behind on the filter
paper and the uranium compounds were carried into solution by HCl.
At Step 10 it appeared that uranium was the predominant metal cation in solution.
However, the author suspected that the yellow crystals were actually a double
salt of ammonium and uranyl ion. Some residual NH4Cl must
have remained adsorbed to the UO2•2H2O
and silicic acid precipitates from Step 5, continuing to adhere through the
washings of Step 6. Evaporating some of Step 10 solution to dryness,
heating intensely enough to drive off NH4 compounds, and re-dissolving
the residue in HCl gave a solution from which no yellow crystals could be
grown. At the temperatures used, Na compounds would not have volatilized.
Figure 17. The absence of other crystal types suggested a
lack of appreciable impurities by Step 10, although the
square, yellow crystals were probably a double salt with
ammonium chloride. Heating the solution to dryness
and volatilizing the NH4 salts, then redissolving in HCl,
gave a solution of what was expected to be uranyl chloride.
The putative uranyl chloride solution did not yield crystals on any attempt.
It seems the solution from Step 10 had just the right molar
proportions of ammonium chloride and uranyl ion to grow
these crystals. This was an accidental but pleasing
outcome of the experiment.
Ammonia precipitation is a useful technique, but by itself it acts on perhaps
too wide a spectrum of ions unless one takes care to adjust pH and other
conditions. Assuming no other precipitations have been done first,
aqueous ammonia can in theory precipitate ions of Fe, U, Cr, Hg, Bi, Ti,
Zr, Th, Al, Be, Sn, Pb, Co, Cu, Zn, and Sb, mostly in the form of hydroxides
and/or hydrated oxides. The boiling, which is necessary to get all
the uranium precipitated, is also what brings down so many other ions.
For example, Cu2+ in ammonia will at first stay in solution as
the tetramminecopper (II) ion, but heating decomposes this and precipitates
CuO. An analogous situation can happen with Ag, Ni, Co, Cd, Zn, Cr,
and even Fe. Ammonia precipitation may also bring down some, but not
all, of the Mn, Mg, and even a few others.
The sheer number of ions that can be brought down by ammonia would seem too
great to render the operation very useful for general work, at least when
it is the first major treatment conducted on the sample. The chemists
who wrote qualitative analysis textbooks evidently thought so, too.
Their answer was to attempt to narrow the range of elements still in solution
by the time the experimenter made it to the ammonia precipitation. At least one,
fairly standard routine has emerged from this. Sorum's text (1953)
uses it, but many other manuals this author has seen appear to contain the
same basic scheme.
The first step in a standard qualitative analysis routine is precipitation
of the "silver group" using cold, dilute HCl. Silver (I), mercury
(I), and lead (II) should come down if present. In our case, where
we had a mineral suspected to be uranophane, we assumed there wasn't much
of any of these. Our assumption might have been wrong, especially in
the case of lead, but we skipped the "silver group" precipitation anyway.
(It is worth noting here that no precipitate of PbCl2 was noticed
in the HCl fraction from our experiment, although the amount of Pb2+
may have been small enough that the chloride remained dissolved; it
would then have been held in solution in the presence of NH4Cl
during Steps 4 and 5. Assuming, on the other hand, that Pb never made
it into solution from the filter paper in Step 3a, we might have wished to
soak the paper in aqueous NH4Cl and then test this solution for
lead, perhaps using cyanidin chloride).
Next in a standard qualitative analysis scheme
is to treat the acidified filtrate or decantate with hydrogen sulfide (danger-
highly toxic) to get even
more elements out of the way before going on to the ammonia step. The
so-called "hydrogen sulfide group" or "copper-arsenic group" will come down
with acidified H2S: tin, copper, arsenic, antimony, bismuth,
cadmium, and the remaining lead and mercury. Again, in our sample there
probably weren't any of these elements present, so we were also able to skip
the "hydrogen sulfide group" step. As stated before, however, one of the
stable decay products of uranium-238 is lead-206; it's therefore hard
to imagine a natural uranium sample being entirely free of lead. 1
Finally, when the qualitative analytical separation actually arrives at the
ammonia step, it often combines the NH4Cl and NH4OH
with ammonium sulfide. This precipitates the sulfides of manganese,
zinc, nickel, and ferrous iron, the bulk of which should not have come down
with H2S treatment. Furthermore, ammonium hydroxide / sulfide
treatment will precipitate the hydroxides of aluminum, chromium (III), and
ferric iron.
In the abbreviated scheme used in the present
experiment, however, we simply used aqueous ammonia and skipped a number
of steps. Though it wouldn't have been specific enough if a spectrum
of metal ions had been present, we weren't expecting the silver group, the
copper-nickel group, or the ammonium sulfide group in the sample. Because
the only metal ions we expected to encounter were those of uranium and calcium,
we had some basis for our omissions. 2
Even if a specimen definitely is uranophane (in this experiment, it was) and
therefore doesn't have most of the ions that would occur in a qualitative
analysis scheme, the fact that a few, similar-looking uranium minerals do
contain Al, V, and / or Mg should have some bearing on the analysis.
After all, that mineral thought to be "uranophane" could easily be tyuyamunite,
carnotite, or one of several others. Therefore it is wise for us to
consider a few additional points, even though we used a heavily simplified
procedure:- Al(OH)3 and certain other precipitates will come down at first from ammonia treatment, but they'll redissolve in part if too much ammonia is added (excess NaOH redissolves Al(OH)3 completely). This will cause noticeable traces of Al to carry into subsequent steps. According to Hillebrand and Lundell (1929), the pH should be adjusted to about 7.0 to bring down Al as completely as possible. It should stay below 7.5 to prevent re-dissolution. In other words, don't add so much ammonia that the solution becomes noticeably alkaline. If one wants to re-dissolve the hydroxides of Al, Zn, Cr, Sb, Sn, and Pb on purpose, NaOH is more effective than NH3; the latter doesn't form as much OH- in solution. (One can then attempt to separate sodium aluminate, sodium zincate, etc. from these solutions; see Sorum, 1953). Uranophane doesn't contain aluminum in its formula, but a number of related minerals do. It is probable that the matrix rock contains at least some Al3+.
- Uranium ions will not precipitate completely if ammonium carbonate
is present in the aqueous ammonia, because ammonium uranyl carbonate forms.
Enough ammonium carbonate will therefore keep all U completely in solution.
In our experiment there may have been minor amounts of ammonium carbonate
already in solution, since we started with a sodium carbonate fusion and
then had it in solution with ammonia (even though most of the carbonate was
driven off by the intense heat of fusion, some was probably reabsorbed from
the air on standing). One way to get around this would be to roast
the fusion very intensely in a crucible to drive off the CO2 (leaving
hydroxides and oxides) then let it cool in a CO2-free atmosphere
or a vacuum desiccator having soda-lime in the bottom. Taking
up this cooled fusion with freshly-prepared aqueous ammonia (made with
degassed, deionized water and pure NH3) would minimize the amount
of ammonium carbonate in the resulting solution. For qualitative work,
though, a brand-new bottle of NH4OH is sufficient.
- Vanadium ions will also tend to keep uranium ions in solution,
causing them to carry over into the next step. (This can happen if, instead
of "uranophane", the sample is actually carnotite)
- Certain ions will co-precipitate species that wouldn't normally
come out of solution. For example, the ammonia will cause calcium phosphate
to come down from solution if the mineral sample had any phosphate in it.
(This can happen if the mineral is autunite instead of uranophane)
- Magnesium ions would precipitate during the ammonia step too,
if it weren't for the ammonium chloride that Smith recommends adding to the
solution prior to adding the ammonia. A person is not born knowing
such things.
With the complexities of ammonia precipitation, it seems hard to be sure
if anything in our experiment either stays in solution entirely or precipitates
completely. Since the focus here is qualitative rather than quantitative
analysis, however, it is not as critical if a small amount of some ion carries
into subsequent steps.
If, by the way, we're sure a sample contains uranium-- for example, it's both
yellow and radioactive-- but we get no yellow precipitate at the ammonia
step, then we know that sample either contains vanadium with the uranium,
or else we must have ammonium carbonate in the solution. Perhaps both
are present. Chances are that we will have some carbonate left
over from the fusion unless we roast the sodium carbonate to the point where
it decomposes to NaOH (somewhere about 1000-1200°C). Even
then it will reabsorb significant amounts of atmospheric CO2 as
it cools. Unless the cooling happens in an atmosphere free of
carbon dioxide, the NaOH will contain much Na2CO3; when
this is dissolved in aqueous ammonia, there will consequently be the requisite
ammonium and carbonate ions to hold uranium in solution.
There is also the matter of CO2 absorbtion by the ammonia solution
itself; a flexible plastic bottle of aqueous NH3 that has
sat on a shelf for five or ten years is probably going to have absorbed some
CO2. Again, this can cause ammonia precipitation of uranium
ions to be incomplete.
After the ammonia precipitation step, Smith's procedure calls for treatment
of the filtrate with ammonium oxalate and then ammonium phosphate.
Let's examine why.
Not counting Na+ and K+, which would have entered the
sample in huge amounts by way of the fusion anyway, Smith's NH4OH
/ NH4Cl treatment should hold Ca++, Mg++,
and Ba++ in solution-- so long as there is no co-precipitation
and we're fairly sure there aren't unaccounted-for metal ions in solution.
Of these three remaining species, uranophane contains only Ca++
in appreciable amounts.
The calcium ions should come down as insoluble calcium oxalate when the ammonium
oxalate is added. Oxalate precipitations become complete only slowly;
unless the solution is boiled for a few minutes, it can take days to
reach completeness. Normally, most of the Ba++ and some
of the Mg++ would also come down; however, according to
procedure P-19 in the second edition of Smith (1953), making the filtrate
strongly alkaline with an excess of NH4OH before adding the ammonium
oxalate should keep most of the Mg++ in solution, provided it's
kept in the cold. The oxalates of Ba and Ca, along with any Sr that
might be present, should therefore be the only ones to come down appreciably.
If we used uranophane as our mineral sample, we might expect the oxalate
step to be the last.
If multiple Group II ions are present in the initial sample, the only one
that should remain in solution by the last stage is Mg++.
The final, phosphate treatment of the filtrate should thus bring this down
as the insoluble magnesium ammonium phosphate (struvite), MgNH4PO4
• 6H2O. A sample of this
white precipitate (which we don't expect to have if our sample is pure uranophane
and we do the procedure correctly) should produce a pinkish color when moistened
with cobalt nitrate solution and heated on the carbon block (Smith, 1953).
The mineral sample in this experiment behaved mostly as expected with regard
to the various tests. It is interesting that a failure to wash out
all the NH4Cl at Step 6 created a useful result. It will make a good
future project to try duplicating the reagent concentrations necessary to
allow crystallization of ammonium uranyl chloride.
For any future experiments involving uranium, we remind the reader once again
to exercise caution. Again, uranium's primary hazard is one of heavy
metal toxicity; wear a double layer of vinyl or latex gloves when handling
soluble uranium compounds, and of course avoid breathing in dusts that may
contain uranium.Notes - "Ammonia Precipitation"
1 As with the silver group, we may have to revisit this step if it appears our mineral sample contained more than we expected.
Back to the article
2 Impatient lab partners of the world, rejoice. A bit of a digression here: a number of students, even at the graduate level, make it a point to work as rapidly as they can, cutting corners as much as possible in an all-but-overt race to see who can "just get out of there" the fastest. Some of these unfortunately go on to design or prescribe the medicines people take every day. A few go on to handle surgical scalpels. The analytical mindset, especially when applied to any kind of work where someone's life or future depends on the outcome, should have no room for the kind of habitual impatience this writer has seen on so many occasions.
Editorializing aside, it's still not very thorough or scientific
to cut out whole chunks of the qualitative analysis scheme if we want to
allow that our mineral might not be uranophane at all. Time permitting,
we'd want to look for every common metal ion, at least the ones expected
to occur in yellow, radioactive minerals that could be mistaken for uranophane.
Back to the article
References - "Ammonia Precipitation"
Hillebrand, W., and Lundell, G. Applied Inorganic Analysis. New York: John Wiley and Sons, 1929.
Smith, Orsino C. Identification and Qualitative Analysis of Minerals, 1st ed. Princeton: Van Nostrand, 1946.
Smith, Orsino C. Identification and Qualitative Analysis of Minerals, 2nd ed. Princeton: Van Nostrand, 1953.
Sorum, C.H. Introduction to Semimicro Qualitative Analysis. New York: Prentice Hall, 1953.
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