CR-Scientific
Minerals & Experimental
Science newsletter
Terms of
UseIssue #9
In this issue:
I. Purification of Hydrochloric Acid?
II. Qualitative Analysis - Manganoan Scapolite from Sparta, New Jersey
III. Qualitative Analysis - A Jade-Like Rock from Northwestern Washington
IV. Site News
While the information in this newsletter is thought to be accurate to the best of the authors' current knowledge, it is not guaranteed to be free of errors or to be suitable for any particular use. The procedures and experiments outlined within can be dangerous or even fatal if carried out improperly. If you choose to attempt any of them, you proceed entirely at your own risk.
I. Purification of Hydrochloric Acid?
Updated
Hydrochloric
acid is
useful in an enormous variety of procedures in the laboratory; it is
possibly
the most widely called-for chemical in qualitative mineral analysis. The avocational
scientist can readily obtain "muriatic acid", an industrial grade of
concentrated hydrochloric acid which is often used for cleaning
stonework and
tile. Muriatic acid fumes strongly when the container is opened (don't
do this
indoors!) and usually has a yellowish tint from metal ion impurities.
This is
mostly Fe+3 (who knows what else might be in there).
Transition
metal impurities such as this destroy the acid's usefulness in
qualitative
mineral analysis, except of course in the effervescence test for
carbonates,
and in the simple solubility tests.Aside from
the obvious danger of trying to distill something as corrosive as
muriatic acid,
there are other obstacles that would present themselves. The first is
the
gaseous nature of HCl itself. Unlike sulfuric acid, which is a liquid
at room
temperature and can exist as such without being dissolved in water,
"hydrochloric acid" is a water solution of hydrogen chloride (HCl)
gas. To turn HCl gas into a water solution of hydrochloric acid, the
gas is bubbled into a container of
distilled water, OR it's sent into a chamber through which a fine mist
of
distilled water is sprayed to dissolve the gas. This latter method
would
be difficult for an amateur chemist to accomplish but makes more sense
in that
a larger surface area of water is acting to dissolve the gas; whereas
in the
former method a single bubble at a time is welling up through liquid
and might
not even have the chance to dissolve before it reaches the top and
disperses
into the laboratory atmosphere. In either case,there's still another
problem:
what concentration of HCl does one have when this procedure is
finished? The
only way to be sure would be to do a careful titration against a
standardized
alkali solution. (An electronic pH meter would give a pretty good
approximation
as well, but only if maintained and calibrated properly.)The second
obstacle to
HCl distillation is the formation of an azeotrope or
constant-boiling
mixture. It's not an obstacle in the sense that it prevents
distillation altogether; it simply limits what the concentration
of HCl will be in the distillate. When one starts heating
concentrated
hydrochloric acid, the first
thing to come off is HCl gas 1.
When
enough HCl gas has left the muriatic acid solution that's being heated,
the
solution will reach a stage at which both HCl gas and water
vapor are
coming over into the receiving tube. At this point it's of course not necessary
to have water in the receiving flask, or else the HCl will become more
dilute
than it has to be. Unless something is done to "break" the azeotrope,
whatever comes over into the receiving tube
from this
moment on will be a specific but sub-maximal concentration of
hydrochloric acid. This
is
normally about 20 weight percent HCl in water, coming over at just
under 110
degrees C. Considering that concentrated hydrochloric acid is about 37%
HCl in
water, the 20% is probably strong enough for a lot applications. This
20% acid
is somewhere near 6 M concentration, one used often in the
laboratory.The third
obstacle is
really an offshoot of the first, but why not save the most spectacular
for
last? Here's the problem. Any time one tries to collect a gas by
dissolving it
in water, there's the danger that a partial vacuum will form in the
setup.
Since it's highly soluble in water, HCl gas is especially prone to
going into
solution very suddenly... creating a pronounced and unexpected drop in
pressure. This causes the cool water in the receiving flask to be
sucked
rapidly back up the tubing and into the hot distilling flask, often shattering
it from the thermal shock! Having this happen
with acid is a
frightful possibility.The
astute reader might ask whether it's really necessary to
dissolve in water the HCl gas that first comes off the impure
acid. In fact, it's possible to obtain a
strong HCl distillate by simply distilling the crude acid over into
vessel containing no added water at all; the final concentration
will be at least 20%. If the condensed droplets of the 20% HCl-H2O
azeotrope are in contact
with the HCl gas for long enough, they will dissolve much of it
and result in a higher concentration. If the dissolution is
thorough enough (not always the case), it will mean the final
concentration of HCl will be
about equal to the initial concentration. The work of Kuehner et al. (1972) lends support to
this. If much of the HCl gas remains undissolved, however, it
will be
hanging around when you go to take apart the apparatus... As stated
before, distillation of HCl is not a very safe procedure. If one
attempts to capture the gaseous HCl by dissolving it in water, there is
the added hazard of a partial vacuum in the setup. To lessen the
danger there
must be one or more "traps" in the line, in which
the gas-carrying tube from the distilling flask goes part-way into an
empty
flask via a two-hole stopper, and another tube goes out the other hole
of the
stopper and into the next flask. If a sudden drop in pressure causes
the liquid
to be sucked out of the receiving vessel, the liquid will end up mostly
in a harmlessly cool "trap" flask. That's the idea, anyway. In
practice it's often necessary to have two or three "trap" flasks in
series before finally having the receiving flask where the distilled
water is
kept. It is surprising how the sudden drop in pressure can cause liquid
to jump
the air gaps in the "trap" flasks and somehow end up back in the hot
distilling flask
2, once again
presenting the danger of shattering it. These
materials are
alright if they stay inside the glassware 3,
but there's always the small but real
chance that any given piece of glassware will shatter during
heating--
if you heated it too fast, handled it roughly, or failed to follow
basic safety
precautions. If a breakage were to happen while heating acids or strong
alkalis, the consequences would be horrible to persons standing nearby.
Room-temperature hydrochloric acid may take a couple of seconds to
destroy the skin,
but the hot acid will corrode skin instantly.For anyone
willing to
accept the substantial risks associated with distilling the mineral
acids, it
is at least preferable (1) to use all-glass apparatus so that H2S
and other compounds from rubber stoppers will not contaminate the
product, and
(2) to do the procedure outside 4,
in an
area away from things and people that could be harmed in case of an
accident.
It is also wise to do acid distillations on as small a scale as is
practical. The mineral acids are entirely unforgiving; mistakes can
kill,
blind, or disfigure the careless. A face shield, goggles, adequate
ventilation,
and a nearby emergency shower are not optional- they are required. Do
not allow
any person or object to block one's path to the emergency shower. II. Qualitative Analysis - Manganoan Scapolite from Sparta, New Jersey
In a field
trip several
years ago to the now-defunct collecting locality at Lime Crest Quarry
in
Sparta, New Jersey, the author found a pink mineral that bore
superficial
resemblance to rhodonite. It did not, however, occur in what the author
understood should be the correct environment for rhodonite (the
"correct
environment" being a few miles down the road in the Franklin / Sterling
Hill orebodies). This intriguing specimen looked like a good candidate
for
qualitative mineral analysis.
Before any
tests
commenced, the author was fairly sure from the mineral's appearance
that it was
a silicate. He also had a strong suspicion it might be scapolite
(marialite
and/or meionite), which is fairly abundant at Lime Crest and can occur
in a
variety of colors. The appearance of the mineral's surface along a
fresh break
was also very similar to that of known scapolite specimens, not to
rhodonite;
only the color was similar to rhodonite. Of the tests that were done, the following produced useful or positive results:
1. Fusibility: A chip of the mineral fused fairly easily in the propane torch flame, bubbling into a white glass. This is typical behavior for scapolite, but this in itself isn't enough to make a positive identification. However, it does rule out rhodonite, bustamite, the feldspars, and a few other minerals the author has often heard collectors confuse with scapolite. Rhodonite and the rest of these minerals are more resistant to a propane torch flame. Unlike scapolite, none of them exhibits the bubbling and resulting formation of a glassy, white bead that the author observed here.
The pink
color also
disappeared entirely upon melting. This is certainly interesting, but
it's not
enough to rule out a transition metal ion as the color source yet (as
opposed
to some unusual artifact of optics or crystal structure); some colorful
metal
ions exist in oxidation states which aren't all that stable. 2. Solubility and Aqueous Chemical Tests:
A.
Concentrated nitric acid did not attack small chips of the
mineral,
even when boiled
for several minutes. The tube was let stand overnight just to make
sure. The
next day, the chips were still there-- unaffected.B.
When crushed
into dust and covered with concentrated hydrochloric acid, the
mineral
did dissolve after standing overnight. A white, insoluble residue of
silica
(powder, not gel) remained in the bottom of the tube; the HCl solution
was
siphoned off carefully with a Pasteur pipet and saved.C.
This HCl
solution was tested first with aqueous ammonia (NH4OH),
which
produced a flocculent, white precipitate that did not re-dissolve in
the
presence of excess ammonia. It also did not dissolve in ammonium
chloride
solution (as would zinc hydroxide).D.
The
precipitate was re-dissolved with HCl and was then tested for cations
which
would come down in the presence of oxalate; the possibilities
included
Th, Sc, the rare-earth elements (Ce, La, etc), and Ca; and to a lesser
extent,
Ba and Sr. A white precipitate did come down, but further tests (as per
Smith,
1953) did not indicate any of the "oxalate group" cations except
calcium in the sample.E.
When the
filtrate was evaporated to dryness with HNO3 and the residue
ignited
to produce oxides, there was vague evidence that Fe or Mn might be
present in
the sample, judging from the color of the ignited oxides (brown to
almost
black). However, attempting to re-dissolve the residue and perform
specific
tests for Mn, Fe, and other transition metals did not produce any
concrete
results. This oxide residue was pretty sparse to begin with; it is
possible
that the amounts involved were too small, or that the solution had
become too
dilute, or that there were too many losses due to spattering when the
crucible
was heated. 3. Bead tests 5 were at first ambiguous until the author decided to use large amounts of mineral sample relative to the amounts of bead reagent used.
There was
a careful
effort to pick from the crushed mineral only the most richly-colored
fragments,
leaving behind anything that looked as if it might contain the green,
matrix
mineral. This sorting was accomplished with some patience and a pair of
sharp
forceps. The actual crushing was done with a hammer against steel, and
the
fragments were dumped onto a carbon block for sorting and incorporation
into
bead tests.
The bead tests resulted in the following colors:
Borax - gave a sort of violet-pink (when hot) and pale red-violet (when cold) upon treatment in the oxidizing flame (O.F.); the bead was essentially colorless from the reducing flame (R.F.). The color of the bead from the O.F. was much more brilliant when a speck of potassium nitrate was added to the borax.
Sodium Carbonate - blue in the O.F., quite brilliant in the presence of KNO3. In the R.F., the color is very faint to non-existent unless KNO3 is present, and then it isn't really a "reducing" flame anyway; it's still an oxidizing flame. The author thought he saw the bead assume a pink color after heating in the R.F. during one trial, but this could not be repeated.
Sodium Fluoride - pale turquoise blue in the O.F., much less pronounced in the R.F. The NaF bead was checked for fluorescence in both SW and LW UV; there was none.
4. Other observations: Between each bead test, the platinum wire loop was cleaned by melting some sodium carbonate on it, removing the molten compound, and putting the wire in dilute HCl. When the last traces of bead chemicals were being cleaned off by burning away in the torch flame, the usual sodium-yellow flame color was observed-- but it was followed by a distinct red-orange flame typical of calcium. Plain HCl did not do this. The pure bead chemicals (sodium carbonate, etc.) did not do this either.
Conclusions and Discussion: The "unknown" mineral from this experiment is probably scapolite (marialite-meionite), with the pink color caused by manganese impurities.
The
presence of a white,
flocculent precipitate during the ammonia step (2C, above) rules out
wollastonite, CaSiO3. Aqueous ammonia will throw down Al,
Be, and
Zn, but not Ca. If the precipitate is caused by beryllium, it will
dissolve in
ammonium carbonate solution; if caused by zinc, it will dissolve in
ammonium
chloride solution (Smith, 1953).The
presence of calcium
is suggested by the precipitation with oxalic acid and by the definite
red-orange coloration that appeared in the flame after the sodium had
"burned off". No other tests the author did suggested anything
besides Ca was thrown down by the oxalate.The
presence of
manganese was established with bead tests. If cobalt had been
responsible for
the mineral's color, the borax bead would have come out blue, not
violet, when
it cooled. Brush and Penfield (1926) indicate that the sodium carbonate
bead,
on the other hand, will be blue in the presence of Mn-- as it
was in
this case. There is no mention in Brush and Penfield of the daylight
color that
would appear in an NaF bead. Smith
(1953) lists five
elements that would make the NaF bead fluorescent in either SW or LW
UV:
bismuth, columbium (niobium), titanium, tungsten, and uranium; the lack
of
fluorescence in the NaF bead suggests the mineral didn't contain
significant
amounts of any of these.This
mineral specimen
gave us the opportunity to exercise a few of the different techniques
that are
available to the mineral hobbyist. In this case the author reached a
point of
reasonable certainty about the mineral's identity, but an experimenter
could
certainly do more tests if desired. Why not search for beryllium,
magnesium, or
barium in the sample-- or look for elements that, had they been present
in
abundance would have showed up in a bead test, but instead might have
been
present in such small traces that only specialized chemical tests would
reveal
them? Even with a modestly-equipped lab setup, there are numerous
possibilities
that one can try if safety is observed at all times.Please note: as with all other information on this website, the reader assumes full responsibility for any consequences that might arise if he / she chooses to attempt the procedures given herein. We cannot control what you do or how you do it; therefore, neither we nor the author(s) accept any responsibility for what may happen as a result.
III. Qualitative Analysis - A Jade-Like Rock from Northwestern Washington
Recently
the author
obtained a green, mottled rock which had come from a beach in the state
of
Washington. The collector who had originally found it stated that in
many years
of walking up and down that beach, he'd found no other rocks like it.
Having
very little familiarity with northwestern Washington and being
therefore unable
to recognize its mineral assemblages simply by looking at them, the
author
decided to resort to some fairly simple qualitative tests.
|
1. Visual and Mechanical Properties.
Simple
scratch tests
suggested the an overall hardness of 6 to 7; this measure was
difficult
to obtain precisely, even when studying the surface of the specimen
with a
magnifier to see if a hardness sample really did make a scratch, or
whether it
was simply a polished streak that looked like a scratch. The
stone was
able to scratch orthoclase (H=6) with difficulty; it took a few tries.
Orthoclase was not able to scratch the unidentified stone, however (so,
let's
say for now that H=6.5). The unidentified stone also had a high overall
toughness and resistance to crushing, as evidenced by its
behavior under
the hammer. Minerals
make up any
given rock, but a rock is not always made of a single mineral;
sometimes,
different minerals are in grains too small to tell apart. In the case
of this
unidentified beach rock, the zones of green and of white seemed to be
of a
different composition from the darker regions, which appeared to be
alteration
products of the green / white mineral(s). These dark regions seemed to
exist
along grain interstices and old micro-fractures where water could have
exerted
its effect through the ages. The brown
mineral was
examined at 10x and 30x. It had a greasy or waxy luster; the overall
appearance
was reminiscent of something from the serpentine or chlorite groups. A
sharp
needle was used under the microscope in an attempt to scratch some of
this
material, which proved to be very soft: estimated hardness perhaps 2,
no more
than 3. It smeared in very much the same manner as talc or
serpentinized
alteration minerals that one might find in certain marble or skarn
deposits.The green
mineral was
harder (as stated before, around 6.5). Inspection at 10x and 30x
suggested it
was not at all amorphous; rather, it seemed to be a typical
massive-form
silicate made of interlocking crystal grains. Though the material
didn't
exhibit cleavage, this is typical of massive-form minerals, even if
they
otherwise show excellent cleavage when present in well-developed
crystals.
There was, however, a vaguely splintery fracture habit which could be
seen
clearly with and even without magnification-- a fracture habit the
author had
never observed in chalcedony (chert, jasper, or agate). The picture
below
doesn't adequately capture the play of light that gives away the
individual
crystal grains, nor does it show off the splintery fracture, but these
were
visible in real life. 
2. Fusibility: a chip consisting mostly of the greenish material fused only along thin edges in the propane torch flame (somewhere between 1600 and 1900°C at the hottest portion). These edges melted fairly quickly, but continued heating did not melt the thicker portions at all. The flame assumed a bright sodium-yellow with a fine but persistent border of red to red-orange (Ca or possibly Sr). The mineral's green color disappeared, leaving an ivory to pale yellow color.
3. Closed Tube Test: a short piece of 5 mm borosilicate glass tubing, sealed off at one end by heating, served as the closed tube. The powdered mineral sample did give off some water upon prolonged heating. No other sublimates were observed.
4. Solubility: A fragment placed in concentrated HCl did not seem to dissolve at all, but the solution turned from colorless to strong yellow.
Aqueous
ammonia was
added to this yellow solution, and a considerable amount of rust-brown
precipitate quickly formed. The yellow color disappeared. When the
mineral chip
was taken out of the liquid and washed, it was evident that the brown,
serpentinaceous mineral had dissolved in the HCl but the green mineral
had been
mostly unaffected.5. Bead Tests: Perhaps to the relief of some readers, but to the disappointment of others who were hoping the CR-Scientific Newsletter's apparent monomania for element 25 might never end, the bead tests did not suggest the presence of manganese in the sample. Furthermore, there were noted differences this time between the "hot" and "cold" states of the beads.
Following are the observed results:
Borax - R.F.- Yellow (hot); Nearly Colorless with a faint aqua tinge (cold)
Borax - O.F. - Bright Yellow (hot); Colorless (cold)
Salt of Phosphorus - R.F. - Colorless (hot); Colorless (cold)
Salt of Phosphorus - O.F. - Bright Yellow (hot); Colorless (cold)
Sodium Carbonate - R.F. - Muddy Yellow-Brown (hot); Pale Green (cold)
Sodium Carbonate - O.F. - Muddy Yellow Brown (hot); Dirty Yellow (cold)
6. Chromate Flux - powdered sample was fused on the carbon block with chromate flux (1 part KHSO4, 1 part K2CrO4, 2 parts sulfur; as per Smith, 1953). The coating near the assay was pale white; same color, hot or cold.
7. Iodide Flux (1 part KHSO4, 1 part KI, 2 parts sulfur; as per Smith, 1953) The coating near the assay was dense white, fading to pale white at some distance from the center. This outer coating was volatile in the reducing flame, but the inner coating was very stable. There were some tiny but definite yellow spots also noted. The coating looked the same, hot or cold.
8. Aqueous Chemical Tests
A.
To prepare the
sample for dissolving, it was mixed with an equal volume of sodium
carbonate
and fused on the carbon block. The fused mass was allowed to cool for
about
20-30 minutes and then crushed into powder.B.
The resulting
powder was placed in a micro
beaker and covered with dilute HCl. Profuse bubbling occurred. A
flocculent
and highly insoluble residue now floated in the liquid, eventually
settling to
the bottom. The solution remained colorless, and there were still
undissolved
lumps at the bottom.C.
In order to
dissolve the remaining sample, concentrated (37%) HCl was carefully
added to
the micro beaker. The bubbling became severe, and the solution quickly
turned
bright yellow with the slightest influence of green. The color remained
this
way.D.
The flocculent
material was allowed to settle. The yellow solution was siphoned off
carefully
with a dropper for further tests. E.
A drop of
this solution was added to a spot plate containing potassium iodide.
This caused it to turn a deep and pure shade of yellow (as opposed to
having a
hint of greenish, which the HCl solution had before this). Adding
strong NaOH
to this destroyed the yellow color completely. F.
About 1/2 mL
of the same solution (from 8D) was placed in a small test tube. The
addition of
strong sodium hydroxide caused a cloudy, white precipitate to
form.
Excess NaOH failed to make this precipitate redissolve, as far as could
be
seen. A separate aliquot of solution was tested with aqueous ammonia
instead of NaOH; this also caused a white precipitate which turned
light brown
on standing. Excess ammonia did not dissolve the precipitate. G.
Another
aliquot of solution from 8D was placed in a clean test tube and brought
up,
using NaOH, to a pH just below neutral. The addition of sodium
ammonium
phosphate (salt of phosphorus) caused a cloudy white precipitate,
presumably MgNH4PO4 • 6H2O. H.
A drop of the
solution from 8D was added to a spot plate well containing potassium
chromate. The solution immediately turned bright orange and cloudy;
more
HCl made the cloudiness disappear, but the orange color persisted.
Bringing the
pH back up to around 3-4 (by adding NaOH) made the orange compound
precipitate
out again. I.
About 1/2 mL
of the solution from 8D was placed in another small test tube. Adding ammonium
sulfide solution caused a cloudy to almost
gelatinous
precipitate that was dark gray to black. It is not certain whether this
was one
compound or two; around the edges the material was white, suggesting
maybe a
hydroxide mixed with a sulfide ((NH4)2S can
precipitate
both if the right ions are present). Boiling the suspended precipitate
down in
a crucible
caused
it to turn rust-brown in the final stages of heating. When it was cool,
this
residue was taken up in a few drops of dilute HCl. Addition of sodium
hexacyanoferrate (II) caused the color to turn deep green to
blue-green
(this was not the typical deep blue of Prussian Blue; it's possible the
color
was caused by the the similar iron hexacyanoferrate compound
known as
Prussian
Green / Berlin Green, which could form under the right conditions). J.
Another 1/2
mL of the solution from 8D was placed in a clean test tube and brought
to just
below pH 7 with NaOH (the point at which the hydroxide precipitate
redissolved
with difficulty). Saturated aqueous oxalic acid was then added
and the
solution boiled for several minutes. No precipitate formed at first,
but the
next day there was a white, crystalline precipitate at the bottom. This
was
washed several times in distilled water, placed in a crucible and
covered with
conc. HNO3, and evaporated to dryness (fume hood!!) to
destroy the
oxalate. The residue was then taken up in HCl and made alkaline with
aqueous
ammonia. There was no precipitate noted, suggesting a lack of Th, Sc,
and the
rare-earth elements but not ruling out Ca, Ba, or Sr (as per Chapter
VII,
Procedure 4 in Smith, 1953) K.
Another 1/2
mL of the solution from 8D was tested with metallic zinc. No
color
change was observed, even after boiling. This suggests the absence of
Ti. L.
Just to be
sure of no Mn, Co, or Ni, an aliquot of solution from 8D was made
alkaline with
ammonia. In the presence of NH4Cl (as would be the case when
neutralizing HCl with ammonia solution), these three will not come out
of
solution until boiled with H2O2 (Smith,
1953).
This treatment caused no observable changes in the test solution. Conclusions and Discussion
The
rock's properties
and the test results suggest a variety of jade; the elements
established with
reasonable certainty are Si, Fe, Na, Ca, and Mg. The following is a
list of the
steps which gave the best evidence: Silicon......step 8B
Iron......step 8I and step 5, especially the sodium carbonate bead. There might have been interfering ions which prevented the green color from coming out in the R.F. beads of borax and salt of phosphorus-- possibly by hindering the reduction of Fe+3 to Fe+2.
Sodium....... the yellow flame test
Calcium......step 8J and the red-orange flame test
Magnesium......steps 8F and 8G
Another
element
suggested in the tests was Pb (steps 7, 8E, 8H). The fact that NaOH
caused the
yellow precipitate in the KI test to disappear is actually not
surprising:
first of all, PbI2 is soluble in alkali (check any suitable
reference, such as the CRC Handbook of Chemistry and Physics);
second,
lead forms an amphoteric hydroxide, meaning that it will dissolve in
excess
NaOH to form a colorless, complex ion. Sorum (1953) indicates this is
an anion
with the formula Pb(OH)4-2. As for the formation
of an
orange precipitate rather than a yellow one in the chromate test (8H),
recall
that lead chromate can be either yellow or orange depending on what
conditions
formed it (example: crocoite). The colors observed in the borax and
salt of
phosphorus bead tests (step 5) don't at all contradict the possibility
of lead
in the sample, either.What
additional tests
might have been useful? There are quite a few: specific gravity,
optics, and
quite a few more chemical tests; the possibilities are limited only by
what
equipment the experimenter has available. Note that we didn't even test
the
flocculent residue from step 2 for W, Ta, or Nb, for example; nor did
we
attempt to precipitate the suspected Pb+2 ions using cold
HCl, as
one would do in a typical qualitative analysis scheme 6. The presence of Mo, V, and a few
other
metals
could have been sought as well. Once again, the
author performed enough
tests
to give at least a plausible idea of the subject's identity: a variety
of jade,
probably an amphibole and most likely nephrite 7,
a mineral whose green color is caused by
ferrous iron, and which may contain traces of lead and other elements
that were
present either as part of the primary mineral or as constituents of
accessory
minerals present in grains too small to pick out (but which found their
way
into the crushed sample and therefore the chemical tests). Hopefully
this has been
an informative journey for our readers... it was certainly enlightening
to this
writer, who originally thought the rock was green jasper!IV. Site News
It looks
like we should
soon have the Belomo
loupes back and in fairly steady supply.Micro-burettes
are now in stock, and they will be in the on-line catalog as
soon as
possible. These micro burettes are of high precision and can be used
for
titrations and other areas of quantitative analysis. They come with a
plastic
base that is best fastened to the desktop with screws or bolts. The
micro-burettes can also be held with an extension
clamp.Teflon®
boiling chips are now available,
eliminating
the worry of contamination that's introduced by using makeshift boiling
stones
fashioned from silicon carbide grinding wheels, broken pieces of
quartz, etc.
The Teflon chips are re-usable and are unaffected by practically any
kind of
liquid. "Teflon" is a registered trademark of E.I. DuPont. Collectors
of
fluorescent minerals who want the brightest in short-wave hand-held UV
lamps
should check out the new UltraLight
5000-S, now in stock. This powerful lamp illuminates a wide area
and really
brings out the colors in those subtly-responding minerals that tend to
get
overlooked with smaller lamps.That concludes this issue of the CR Scientific newsletter.
Until next time, stay safe and have fun.
Notes:
1 If one started boiling a <20% solution of HCl in water, the first thing to come off would be water vapor, and then the constant-boiling solution of approx. 20% HCl. In the case of concentrated acid, it is HCl vapor that comes off first, and then the constant-boiling solution of approx. 20% HCl. Do not confuse weight percents with mole percents or vise versa, or you could have a serious accident. Never sniff or otherwise inhale vapors thinking they're "just water"... you might've miscalculated a percent somewhere and could be inhaling corrosive vapors. Most good chemistry habits arise from common sense. Back to article
2 This seems to defy logic but has happened to the author despite careful set-up of the apparatus, including having the inlet and outlet tubes at different heights. The notion to distill muriatic acid was thus abandoned. Back to article
3 Another reminder, in case you forgot... NEVER breathe acid vapors. They will cause irreparable lung damage and possible death. Back to article
4 In a laboratory fume hood is better. Outside, there's always the danger that the wind could shift and blow acid vapors in your face. A face shield saved the author from severe eye damage when this happened once. However, don't overlook the considerable danger of getting the acid vapors in your lungs-- some types of respirators are designed to filter out HCl gas, etc. One of these would be mandatory for anyone working with concentrated or hot HCl but lacking access to a lab fume hood. Back to article
5 The samples were not suspected at all to contain As, Sb, or other materials which could form brittle alloys with the platinum wire. The sodium carbonate bead with a manganese-containing sample, pointed out by Dana Martin Morong as possibly harmful to the platinum (reference at http://www.rockhounds.com/rocknet/archive/messages/19247.shtml), produced no damage to the platinum wire, though repeating the experiment too many times probably isn't a good idea from the viewpoint of making a platinum wire last as long as possible.
Back to article
6 Instead, any Pb present would've remained in solution, thanks to the strong HCl added to the pulverized fusion (which also released heat). To precipitate lead as the chloride, it requires cold, dilute HCl. Back to article
7 While nephrite (actinolite) doesn't contain sodium in its formula, we've already explored the seawater possibility. Alternatively, our mineral could be one of the many, many amphiboles or pyroxenes similar in composition to actinolite but off by only a sodium atom here and there... the number of different silicate minerals that can possibly contain Fe, Ca, Mg, and Na is impressive. Back to article
Works Cited:
Brush, George, and Penfield, Samuel. Determinative Mineralogy and Blowpipe Analysis, 16th ed. New York: John Wiley & Sons, 1926.
Kuehner, Edwin C., Robert Alvarez, Paul J. Paulsen, and Thomas J. Murphy. "Production and Analysis of Special High-Purity Acids Purified by Sub-Boiling Distillation". Analytical Chemistry 44: 2050-2056 (1972).
Smith, Orsino C. Identification and Qualitative Analysis of Minerals, 2nd ed. Princeton, New Jersey: D. Van Nostrand Co., 1953.
Sorum, C.H. Introduction to Semimicro Qualitative Analysis, 2nd ed. New York: Prentice-Hall, 1953.
While the information in this newsletter is thought to be accurate to the best of the authors' current knowledge, it is not guaranteed to be free of errors or to be suitable for any particular use. The procedures and experiments outlined within can be dangerous or even fatal if carried out improperly. If you choose to attempt any of them, you proceed entirely at your own risk.
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