through the dolomitic limestones, containing varying amounts of
magnesium, to the impure varieties, which contain also ferrous and
manganous carbonates and siliceous compounds in variable proportions.
Many other minerals may be inclosed in limestones in small quantities,
and an exact qualitative analysis will often show the presence of
sulphides or sulphates, phosphates, and titanates, and the alkali or
even the heavy metals. No attempt is made in the following procedures
to provide a complete quantitative scheme which would take into
account all of these constituents. Such a scheme for a complete
analysis of a limestone may be found in Bulletin No. 700 of the United
States Geological Survey. It is assumed that, for these practice
determinations, a limestone is selected which contains only the more
common constituents first enumerated above.
DETERMINATION OF MOISTURE
The determination of the amount of moisture in minerals or ores is
often of great importance. Ores which have been exposed to the weather
during shipment may have absorbed enough moisture to appreciably
affect the results of analysis. Since it is essential that the seller
and buyer should make their analyses upon comparable material, it is
customary for each analyst to determine the moisture in the sample
examined, and then to calculate the percentages of the various
constituents with reference to a sample dried in the air, or at a
temperature a little above 100°C., which, unless the ore has undergone
chemical change because of the wetting, should be the same before and
after shipment.
PROCEDURE.--Spread 25 grams of the powdered sample on a weighed
watch-glass; weigh to the nearest 10 milligrams only and heat at
105°C.; weigh at intervals of an hour, after cooling in a desiccator,
until the loss of weight after an hour's heating does not exceed
10 milligrams. It should be noted that a variation in weight of 10
milligrams in a total weight of 25 grams is no greater relatively than
a variation of 0.1 milligram when the sample taken weighs 0.25 gram
DETERMINATION OF THE INSOLUBLE MATTER AND SILICA
PROCEDURE.--Weigh out two portions of the original powdered sample
(not the dried sample), of about 5 grams each, into 250 cc.
casseroles, and cover each with a watch-glass (Note 1). Pour over the
powder 25 cc. of water, and then add 50 cc. of dilute hydrochloric
acid (sp. gr. 1.12) in small portions, warming gently, until nothing
further appears to dissolve (Note 2). Evaporate to dryness on the
water bath. Pour over the residue a mixture of 5 cc. of water and
5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and again
evaporate to dryness, and finally heat for at least an hour at
a temperature of 110°C. Pour over this residue 50 cc. of dilute
hydrochloric acid (one volume acid (sp. gr. 1.12) to five volumes
water), and boil for about five minutes; then filter and wash twice
with the dilute hydrochloric acid, and then with hot water until
free from chlorides. Transfer the filter and contents to a porcelain
crucible, dry carefully over a low flame, and ignite to constant
weight. The residue represents the insoluble matter and the silica
from any soluble silicates (Note 3).
Calculate the combined percentage of these in the limestone.
[Note 1: The relatively large weight (5 grams) taken for analysis
insures greater accuracy in the determination of the ingredients which
are present in small proportions, and is also more likely to be a
representative sample of the material analyzed.]
[Note 2: It is plain that the amount of the insoluble residue and also
its character will often depend upon the strength of acid used for
solution of the limestone. It cannot, therefore, be regarded as
representing any well-defined constituent, and its determination is
essentially empirical.]
[Note 3: It is probable that some of the silicates present are wholly
or partly decomposed by the acid, and the soluble silicic acid must
be converted by evaporation to dryness, and heating, into white,
insoluble silica. This change is not complete after one evaporation.
The heating at a temperature somewhat higher than that of the water
bath for a short time tends to leave the silica in the form of a
powder, which promotes subsequent filtration. The siliceous residue
is washed first with dilute acid to prevent hydrolytic changes, which
would result in the formation of appreciable quantities of insoluble
basic iron or aluminium salts on the filter when washing with hot
water.
If it is desired to determine the percentage of silica separately, the
ignited residue should be mixed in a platinum crucible with about six
times its weight of anhydrous sodium carbonate, and the procedure
given on page 151 should be followed. The filtrate from the silica is
then added to the main filtrate from the insoluble residue.]
DETERMINATION OF FERRIC OXIDE AND ALUMINIUM OXIDE (WITH MANGANESE)
PROCEDURE.--To the filtrate from the insoluble residue add ammonium
hydroxide until the solution just smells distinctly of ammonia, but do
not add an excess. Then add 5 cc. of saturated bromine water (Note 1),
and boil for five minutes. If the smell of ammonia has disappeared,
again add ammonium hydroxide in slight excess, and 3 cc. of bromine
water, and heat again for a few minutes. Finally add 10 cc. of
ammonium chloride solution and keep the solution warm until it barely
smells of ammonia; then filter promptly (Note 2). Wash the filter
twice with hot water, then (after replacing the receiving beaker) pour
through it 25 cc. of hot, dilute hydrochloric acid (one volume dilute
HCl [sp. gr. 1.12] to five volumes water). A brown residue insoluble
in the acid may be allowed to remain on the filter. Wash the filter
five times with hot water, add to the filtrate ammonium hydroxide
and bromine water as described above, and repeat the precipitation.
Collect the precipitate on the filter already used, wash it free from
chlorides with hot water, and ignite and weigh as described for ferric
hydroxide on page 110. The residue after ignition consists of ferric
oxide, alumina, and mangano-manganic oxide (Mn_{3}O_{4}), if manganese
is present. These are commonly determined together (Note 3).
Calculate the percentage of the combined oxides in the limestone.
[Note 1: The addition of bromine water to the ammoniacal solutions
serves to oxidize any ferrous hydroxide to ferric hydroxide and to
precipitate manganese as MnO(OH)_{2}. The solution must contain not
more than a bare excess of hydroxyl ions (ammonium hydroxide) when it
is filtered, on account of the tendency of the aluminium hydroxide to
redissolve.
The solution should not be strongly ammoniacal when the bromine is
added, as strong ammonia reacts with the bromine, with the evolution
of nitrogen.]
[Note 2: The precipitate produced by ammonium hydroxide and bromine
should be filtered off promptly, since the alkaline solution absorbs
carbon dioxide from the air, with consequent partial precipitation
of the calcium as carbonate. This is possible even under the most
favorable conditions, and for this reason the iron precipitate is
redissolved and again precipitated to free it from calcium. When the
precipitate is small, this reprecipitation may be omitted.]
[Note 3: In the absence of significant amounts of manganese the iron
and aluminium may be separately determined by fusion of the mixed
ignited precipitate, after weighing, with about ten times its weight
of acid potassium sulphate, solution of the cold fused mass in water,
and volumetric determination of the iron, as described on page 66.
The aluminium is then determined by difference, after subtracting the
weight of ferric oxide corresponding to the amount of iron found.
If a separate determination of the iron, aluminium, and manganese
is desired, the mixed precipitate may be dissolved in acid before
ignition, and the separation effected by special methods (see, for
example, Fay, !Quantitative Analyses!, First Edition, pp. 15-19 and
23-27).]
DETERMINATION OF CALCIUM
PROCEDURE.--To the combined filtrates from the double precipitation of
the hydroxides just described, add 5 cc. of dilute ammonium hydroxide
(sp. gr. 0.96), and transfer the liquid to a 500 cc. graduated flask,
washing out the beaker carefully. Cool to laboratory temperature, and
fill the flask with distilled water until the lowest point of the
meniscus is exactly level with the mark on the neck of the flask.
Carefully remove any drops of water which are on the inside of the
neck of the flask above the graduation by means of a strip of filter
paper, make the solution uniform by pouring it out into a dry beaker
and back into the flask several times. Measure off one fifth of this
solution as follows (Note 1): Pour into a 100 cc. graduated flask
about 10 cc. of the solution, shake the liquid thoroughly over the
inner surface of the small flask, and pour it out. Repeat the same
operation. Fill the 100 cc. flask until the lowest point of the
meniscus is exactly level with the mark on its neck, remove any drops
of solution from the upper part of the neck with filter paper, and
pour the solution into a beaker (400-500 cc.). Wash out the flask with
small quantities of water until it is clean, adding these to the 100
cc. of solution. When the duplicate portion of 100 cc. is measured out
from the solution, remember that the flask must be rinsed out twice
with that solution, as prescribed above, before the measurement is
made. (A 100 cc. pipette may be used to measure out the aliquot
portions, if preferred.)
Dilute each of the measured portions to 250 cc. with distilled water,
heat the whole to boiling, and add ammonium oxalate solution slowly
in moderate excess, stirring well. Boil for two minutes; allow the
precipitated calcium oxalate to settle for a half-hour, and decant
through a filter. Test the filtrate for complete precipitation by
adding a few cubic centimeters of the precipitant, allowing it to
stand for fifteen minutes. If no precipitate forms, make the solution
slightly acid with hydrochloric acid (Note 2); see that it is properly
labeled and reserve it to be combined with the filtrate from the
second calcium oxalate precipitation (Notes 3 and 4).
Redissolve the calcium oxalate in the beaker with warm hydrochloric
acid, pouring the acid through the filter. Wash the filter five times
with water, and finally pour through it aqueous ammonia. Dilute the
solution to 250 cc., bring to boiling, and add 1 cc. ammonium oxalate
solution (Note 5) and ammonia in slight excess; boil for two minutes,
and set aside for a half-hour. Filter off the calcium oxalate upon the
filter first used, and wash free from chlorides. The filtrate should
be made barely acid with hydrochloric acid and combined with the
filtrate from the first precipitation. Begin at once the evaporation
of the solutions for the determination of magnesium as described
below.
The precipitate of calcium oxalate may be converted into calcium oxide
by ignition without previous drying. After burning the filter, it may
be ignited for three quarters of an hour in a platinum crucible at
the highest heat of the Bunsen or Tirrill burner, and finally for ten
minutes at the blast lamp (Note 6). Repeat the heating over the blast
lamp until the weight is constant. As the calcium oxide absorbs
moisture from the air, it must (after cooling) be weighed as rapidly
as possible.
The precipitate may, if preferred, be placed in a weighted porcelain
crucible. After burning off the filter and heating for ten minutes the
calcium precipitate may be converted into calcium sulphate by placing
2 cc. of dilute sulphuric acid in the crucible (cold), heating the
covered crucible very cautiously over a low flame to drive off the
excess of acid, and finally at redness to constant weight (Note 7).
From the weight of the oxide or sulphate, calculate the percentage of
the calcium (Ca) in the limestone, remembering that only one fifth of
the total solution is used for this determination.
[Note 1: If the calcium were precipitated from the entire solution,
the quantity of the precipitate would be greater than could be
properly treated. The solution is, therefore, diluted to a definite
volume (500 cc.), and exactly one fifth (100 cc.) is measured off in a
graduated flask or by means of a pipette.]
[Note 2: The filtrate from the calcium oxalate should be made slightly
acid immediately after filtration, in order to avoid the solvent
action of the alkaline liquid upon the glass.]
[Note 3: The accurate quantitative separation of calcium and magnesium
as oxalates requires considerable care. The calcium precipitate
usually carries down with it some magnesium, and this can best
be removed by redissolving the precipitate after filtration, and
reprecipitation in the presence of only the small amount of magnesium
which was included in the first precipitate. When, however, the
proportion of magnesium is not very large, the second precipitation of
the calcium can usually be avoided by precipitating it from a rather
dilute solution (800 cc. or so) and in the presence of a considerable
excess of the precipitant, that is, rather more than enough to convert
both the magnesium and calcium into oxalates.]
[Note 4: The ionic changes involved in the precipitation of calcium
as oxalate are exceedingly simple, and the principles discussed in
connection with the barium sulphate precipitation on page 113 also
apply here. The reaction is
C_{2}O_{4}^{--} + Ca^{++} --> [CaC_{2}O_{4}].
Calcium oxalate is nearly insoluble in water, and only very slightly
soluble in acetic acid, but is readily dissolved by the strong mineral
acids. This behavior with acids is explained by the fact that oxalic
acid is a stronger acid than acetic acid; when, therefore, the oxalate
is brought into contact with the latter there is almost no tendency to
diminish the concentration of C_{2}O_{4}^{--} ions by the formation of
an acid less dissociated than the acetic acid itself, and practically
no solvent action ensues. When a strong mineral acid is present,
however, the ionization of the oxalic acid is much reduced by the high
concentration of the H^{+} ions from the strong acid, the formation
of the undissociated acid lessens the concentration of the
C_{2}O_{4}^{--} ions in solution, more of the oxalate passes into
solution to re-establish equilibrium, and this process repeats itself
until all is dissolved.
The oxalate is immediately reprecipitated from such a solution on the
addition of OH^{-} ions, which, by uniting with the H^{+} ions of the
acids (both the mineral acid and the oxalic acid) to form water, leave
the Ca^{++} and C_{2}O_{4}^{--} ions in the solution to recombine to
form [CaC_{2}O_{4}], which is precipitated in the absence of the
H^{+} ions. It is well at this point to add a small excess of
C_{2}O_{4}^{--} ions in the form of ammonium oxalate to decrease the
solubility of the precipitate.
The oxalate precipitate consists mainly of CaC_{2}O_{4}.H_{2}O when
thrown down.]
[Note 5: The small quantity of ammonium oxalate solution is added
before the second precipitation of the calcium oxalate to insure
the presence of a slight excess of the reagent, which promotes the
separation of the calcium compound.]
[Note 6: On ignition the calcium oxalate loses carbon dioxide and
carbon monoxide, leaving calcium oxide:
CaC_{2}O_{4}.H_{2}O --> CaO + CO_{2} + CO + H_{2}O.
For small weights of the oxalate (0.6 gram or less) this reaction may
be brought about in a platinum crucible at the highest temperature of
a Tirrill burner, but it is well to ignite larger quantities than this
over the blast lamp until the weight is constant.]
[Note 7: The heat required to burn the filter, and that subsequently
applied as described, will convert most of the calcium oxalate to
calcium carbonate, which is changed to sulphate by the sulphuric acid.
The reactions involved are
CaC_{2}O_{4} --> CaCO_{3} + CO,
CaCO_{3} + H_{2}SO_{4} --> CaSO_{4} + H_{2}O + CO_{2}.
If a porcelain crucible is employed for ignition, this conversion to
sulphate is to be preferred, as a complete conversion to oxide is
difficult to accomplish.]
[Note 8: The determination of the calcium may be completed
volumetrically by washing the calcium oxalate precipitate from
the filter into dilute sulphuric acid, warming, and titrating
the liberated oxalic acid with a standard solution of potassium
permanganate as described on page 72. When a considerable number of
analyses are to be made, this procedure will save much of the time
otherwise required for ignition and weighing.]
DETERMINATION OF MAGNESIUM
PROCEDURE.--Evaporate the acidified filtrates from the calcium
precipitates until the salts begin to crystallize, but do !not!
evaporate to dryness (Note 1). Dilute the solution cautiously until
the salts are brought into solution, adding a little acid if the
solution has evaporated to very small volume. The solution should be
carefully examined at this point and must be filtered if a precipitate
has appeared. Heat the clear solution to boiling; remove the burner
and add 25 cc. of a solution of disodium phosphate. Then add slowly
dilute ammonia (1 volume strong ammonia (sp. gr. 0.90) and 9 volumes
water) as long as a precipitate continues to form. Finally, add a
volume of concentrated ammonia (sp. gr. 0.90) equal to one third of
the volume of the solution, and allow the whole to stand for about
twelve hours.
Decant the solution through a filter, wash it with dilute ammonia
water, proceeding as prescribed for the determination of phosphoric
anhydride on page 122, including; the reprecipitation (Note 2),
except that 3 cc. of disodium phosphate solution are added before the
reprecipitation of the magnesium ammonium phosphate instead of
the magnesia mixture there prescribed. From the weight of the
pyrophosphate, calculate the percentage of magnesium oxide (MgO) in
the sample of limestone. Remember that the pyrophosphate finally
obtained is from one fifth of the original sample.
[Note 1: The precipitation of the magnesium should be made in as small
volume as possible, and the ratio of ammonia to the total volume of
solution should be carefully provided for, on account of the relative
solubility of the magnesium ammonium phosphate. This matter has
been fully discussed in connection with the phosphoric anhydride
determination.]
[Note 2: The first magnesium ammonium phosphate precipitate is rarely
wholly crystalline, as it should be, and is not always of the proper
composition when precipitated in the presence of such large amounts of
ammonium salts. The difficulty can best be remedied by filtering the
precipitate and (without washing it) redissolving in a small quantity
of hydrochloric acid, from which it may be again thrown down by
ammonia after adding a little disodium phosphate solution. If the
flocculent character was occasioned by the presence of magnesium
hydroxide, the second precipitation, in a smaller volume containing
fewer salts, will often result more favorably.
The removal of iron or alumina from a contaminated precipitate is
a matter involving a long procedure, and a redetermination of the
magnesium from a new sample, with additional precautions, is usually
to be preferred.]
DETERMINATION OF CARBON DIOXIDE
!Absorption Apparatus!
[Illustration: Fig. 3]
The apparatus required for the determination of the carbon dioxide
should be arranged as shown in the cut (Fig. 3). The flask (A) is
an ordinary wash bottle, which should be nearly filled with dilute
hydrochloric acid (100 cc. acid (sp. gr. 1.12) and 200 cc. of water).
The flask is connected by rubber tubing (a) with the glass tube (b)
leading nearly to the bottom of the evolution flask (B) and having its
lower end bent upward and drawn out to small bore, so that the carbon
dioxide evolved from the limestone cannot bubble back into (b). The
evolution flask should preferably be a wide-mouthed Soxhlet extraction
flask of about 150 cc. capacity because of the ease with which tubes
and stoppers may be fitted into the neck of a flask of this type. The
flask should be fitted with a two-hole rubber stopper. The condenser
(C) may consist of a tube with two or three large bulbs blown in
it, for use as an air-cooled condenser, or it may be a small
water-jacketed condenser. The latter is to be preferred if a number of
determinations are to be made in succession.
A glass delivery tube (c) leads from the condenser to the small U-tube
(D) containing some glass beads or small pieces of glass rod and 3 cc.
of a saturated solution of silver sulphate, with 3 cc. of concentrated
sulphuric acid (sp. gr. 1.84). The short rubber tubing (d) connects
the first U-tube to a second U-tube (E) which is filled with small
dust-free lumps of dry calcium chloride, with a small, loose plug of
cotton at the top of each arm. Both tubes should be closed by cork
stoppers, the tops of which are cut off level with, or preferably
forced a little below, the top of the U-tube, and then neatly sealed
with sealing wax.
The carbon dioxide may be absorbed in a tube containing soda lime
(F) or in a Geissler bulb (F') containing a concentrated solution
of potassium hydroxide (Note 2). The tube (F) is a glass-stoppered
side-arm U-tube in which the side toward the evolution flask and one
half of the other side are filled with small, dust-free lumps of soda
lime of good quality (Note 3). Since soda lime contains considerable
moisture, the other half of the right side of the tube is filled with
small lumps of dry, dust-free calcium chloride to retain the moisture
from the soda lime. Loose plugs of cotton are placed at the top of
each arm and between the soda lime and the calcium chloride.
The Geissler bulb (F'), if used, should be filled with potassium
hydroxide solution (1 part of solid potassium hydroxide dissolved in
two parts of water) until each small bulb is about two thirds full
(Note 4). A small tube containing calcium chloride is connected with
the Geissler bulb proper by a ground joint and should be wired to the
bulb for safety. This is designed to retain any moisture from the
hydroxide solution. A piece of clean, fine copper wire is so attached
to the bulb that it can be hung from the hook above a balance pan, or
other support.
The small bottle (G) with concentrated sulphuric acid (sp. gr. 1.84)
is so arranged that the tube (f) barely dips below the surface. This
will prevent the absorption of water vapor by (F) or (F') and serves
as an aid in regulating the flow of air through the apparatus. (H) is
an aspirator bottle of about four liters capacity, filled with water;
(k) is a safety tube and a means of refilling (H); (h) is a screw
clamp, and (K) a U-tube filled with soda lime.
[Note 1: The air current, which is subsequently drawn through the
apparatus, to sweep all of the carbon dioxide into the absorption
apparatus, is likely to carry with it some hydrochloric acid from
the evolution flask. This acid is retained by the silver sulphate
solution. The addition of concentrated sulphuric acid to this solution
reduces its vapor pressure so far that very little water is carried on
by the air current, and this slight amount is absorbed by the calcium
chloride in (E). As the calcium chloride frequently contains a small
amount of a basic material which would absorb carbon dioxide, it is
necessary to pass carbon dioxide through (E) for a short time and then
drive all the gas out with a dry air current for thirty minutes before
use.]
[Note 2: Soda-lime absorption tubes are to be preferred if a
satisfactory quality of soda lime is available and the number of
determinations to be made successively is small. The potash bulbs will
usually permit of a larger number of successive determinations without
refilling, but they require greater care in handling and in the
analytical procedure.]
[Note 3: Soda lime is a mixture of sodium and calcium hydroxides. Both
combine with carbon dioxide to form carbonates, with the evolution
of water. Considerable heat is generated by the reaction, and the
temperature of the tube during absorption serves as a rough index of
the progress of the reaction through the mass of soda lime.
It is essential that soda lime of good quality for analytical purposes
should be used. The tube should not contain dust, as this is likely to
be swept away.]
[Note 4: The solution of the hydroxide for use in the Geissler bulb
must be highly concentrated to insure complete absorption of the
carbon dioxide and also to reduce the vapor pressure of the solution,
thus lessening the danger of loss of water with the air which passes
through the bulbs. The small quantity of moisture which is then
carried out of the bulbs is held by the calcium chloride in the
prolong tube. The best form of absorption bulb is that to which the
prolong tube is attached by a ground glass joint.
After the potassium hydroxide is approximately half consumed in the
first bulb of the absorption apparatus, potassium bicarbonate is
formed, and as it is much less soluble than the carbonate, it often
precipitates. Its formation is a warning that the absorbing power of
the hydroxide is much diminished.]
!The Analysis!
PROCEDURE.-- Weigh out into the flask (B) about 1 gram of limestone.
Cover it with 15 cc. of water. Weigh the absorption apparatus (F)
or (F') accurately after allowing it to stand for 30 minutes in the
balance case, and wiping it carefully with a lintless cloth, taking
care to handle it as little as possible after wiping (Note 1). Connect
the absorption apparatus with (e) and (f). If a soda-lime tube is
used, be sure that the arm containing the soda lime is next the tube
(E) and that the glass stopcocks are open.
To be sure that the whole apparatus is airtight, disconnect the rubber
tube from the flask (A), making sure that the tubes (a) and (b) do not
contain any hydrochloric acid, close the pinchcocks (a) and (k) and
open (h). No bubbles should pass through (D) or (G) after a few
seconds. When assured that the fittings are tight, close (h) and open
(a) cautiously to admit air to restore atmospheric pressure. This
precaution is essential, as a sudden inrush of air will project liquid
from (D) or (F'). Reconnect the rubber tube with the flask (A).
Open the pinchcocks (a) and (k) and blow over about 10 cc. of the
hydrochloric acid from (A) into (B). When the action of the acid
slackens, blow over (slowly) another 10 cc.
The rate of gas evolution should not exceed for more than a few
seconds that at which about two bubbles per second pass through (G)
(Note 2). Repeat the addition of acid in small portions until the
action upon the limestone seems to be at an end, taking care to close
(a) after each addition of acid (Note 3). Disconnect (A) and connect
the rubber tubing with the soda-lime tube (K) and open (a). Then close
(k) and open (h), regulating the flow of water from (H) in such a way
that about two bubbles per second pass through (G). Place a small
flame under (B) and !slowly! raise the contents to boiling and boil
for three minutes. Then remove the burner from under (B) and continue
to draw air through the apparatus for 20-30 minutes, or until (H)
is emptied (Note 4). Remove the absorption apparatus, closing the
stopcocks on (F) or stoppering the open ends of (F'), leave the
apparatus in the balance case for at least thirty minutes, wipe it
carefully and weigh, after opening the stopcocks (or removing plugs).
The increase in weight is due to absorption of CO_{2}, from which its
percentage in the sample may be calculated.
After cleaning (B) and refilling (H), the apparatus is ready for the
duplicate analysis.
[Note 1: The absorption tubes or bulbs have large surfaces on which
moisture may collect. By allowing them to remain in the balance case
for some time before weighing, the amount of moisture absorbed on the
surface is as nearly constant as practicable during two weighings, and
a uniform temperature is also assured. The stopcocks of the U-tube
should be opened, or the plugs used to close the openings of the
Geissler bulb should be removed before weighing in order that the air
contents shall always be at atmospheric pressure.]
[Note 2: If the gas passes too rapidly into the absorption apparatus,
some carbon dioxide may be carried through, not being completely
retained by the absorbents.]
[Note 3: The essential ionic changes involved in this procedure are
the following: It is assumed that the limestone, which is typified by
calcium carbonate, is very slightly soluble in water, and the ions
resulting are Ca^{++} and CO_{3}^{--}. In the presence of H^{+} ions
of the mineral acid, the CO_{3}^{--} ions form [H_{2}CO_{3}]. This
is not only a weak acid which, by its formation, diminishes the
concentration of the CO_{3}^{--} ions, thus causing more of the
carbonate to dissolve to re-establish equilibrium, but it is also an
unstable compound and breaks down into carbon dioxide and water.]
[Note 4: Carbon dioxide is dissolved by cold water, but the gas is
expelled by boiling, and, together with that which is distributed
through the apparatus, is swept out into the absorption bulb by the
current of air. This air is purified by drawing it through the tube
(K) containing soda lime, which removes any carbon dioxide which may
be in it.]
DETERMINATION OF LEAD, COPPER, IRON, AND ZINC IN BRASS
ELECTROLYTIC SEPARATIONS
!General Discussion!
When a direct current of electricity passes from one electrode to
another through solutions of electrolytes, the individual ions present
in these solutions tend to move toward the electrode of opposite
electrical charge to that which each ion bears, and to be discharged
by that electrode. Whether or not such discharge actually occurs in
the case of any particular ion depends upon the potential (voltage) of
the current which is passing through the solution, since for each ion
there is, under definite conditions, a minimum potential below which
the discharge of the ion cannot be effected. By taking advantage
of differences in discharge-potentials, it is possible to effect
separations of a number of the metallic ions by electrolysis, and at
the same time to deposit the metals in forms which admit of direct
weighing. In this way the slower procedures of precipitation and
filtration may frequently be avoided. The following paragraphs present
a brief statement of the fundamental principles and conditions
underlying electro-analysis.
The total energy of an electric current as it passes through a
solution is distributed among three factors, first, its potential,
which is measured in volts, and corresponds to what is called "head"
in a stream of water; second, current strength, which is measured
in amperes, and corresponds to the volume of water passing a
cross-section of a stream in a given time interval; and third, the
resistance of the conducting medium, which is measured in ohms. The
relation between these three factors is expressed by Ohm's law,
namely, that !I = E/R!, when I is current strength, E potential, and R
resistance. It is plain that, for a constant resistance, the
strength of the current and its potential are mutually and directly
interdependent.
As already stated, the applied electrical potential determines whether
or not deposition of a metal upon an electrode actually occurs. The
current strength determines the rate of deposition and the physical
characteristics of the deposit. The resistance of the solution is
generally so small as to fall out of practical consideration.
Approximate deposition-potentials have been determined for a number
of the metallic elements, and also for hydrogen and some of the
acid-forming radicals. The values given below are those required
for deposition from normal solutions at ordinary temperatures
with reference to a hydrogen electrode. They must be regarded as
approximate, since several disturbing factors and some secondary
reactions render difficult their exact application under the
conditions of analysis. They are:
Zn Cd Fe Ni Pb H Cu Sb Hg Ag SO_{4}
+0.77 +0.42 +0.34 +0.33 +0.13 0 -0.34 -0.67 -0.76 -0.79 +1.90
From these data it is evident that in order to deposit copper from a
normal solution of copper sulphate a minimum potential equal to the
algebraic sum of the deposition-potentials of copper ions and sulphate
ions must be applied, that is, +1.56 volts. The deposition of zinc
from a solution of zinc sulphate would require +2.67 volts, but, since
the deposition of hydrogen from sulphuric acid solution requires only
+1.90 volts, the quantitative deposition of zinc by electrolysis from
a sulphuric acid solution of a zinc salt is not practicable. On the
other hand, silver, if present in a solution of copper sulphate, would
deposit with the copper.
The foregoing examples suffice to illustrate the application of the
principle of deposition potentials, but it must further be noted
that the values stated apply to normal solutions of the compounds in
question, that is, to solutions of considerable concentrations. As the
concentration of the ions diminishes, and hence fewer ions approach
the electrodes, somewhat higher voltages are required to attract and
discharge them. From this it follows that the concentrations should be
kept as high as possible to effect complete deposition in the least
practicable time, or else the potentials applied must be progressively
increased as deposition proceeds. In practice, the desired result is
obtained by starting with small volumes of solution, using as large an
electrode surface as possible, and by stirring the solution to bring
the ions in contact with the electrodes. This is, in general, a more
convenient procedure than that of increasing the potential of the
current during electrolysis, although that method is also used.
As already stated, those ions in a solution of electrolytes will first
be discharged which have the lowest deposition potentials, and so
long as these ions are present around the electrode in considerable
concentration they, almost alone, are discharged, but, as their
concentration diminishes, other ions whose deposition potentials are
higher but still within that of the current applied, will also begin
to separate. For example, from a nitric acid solution of copper
nitrate, the copper ions will first be discharged at the cathode, but
as they diminish in concentration hydrogen ions from the acid (or
water) will be also discharged. Since the hydrogen thus liberated is a
reducing agent, the nitric acid in the solution is slowly reduced to
ammonia, and it may happen that if the current is passed through for a
long time, such a solution will become alkaline. Oxygen is liberated
at the anode, but, since there is no oxidizable substance present
around that electrode, it escapes as oxygen gas. It should be noted
that, in general, the changes occurring at the cathode are reductions,
while those at the anode are oxidations.
For analytical purposes, solutions of nitrates or sulphates of the
metals are preferable to those of the chlorides, since liberated
chlorine attacks the electrodes. In some cases, as for example, that
of silver, solution of salts forming complex ions, like that of
the double cyanide of silver and potassium, yield better metallic
deposits.
Most metals are deposited as such upon the cathode; a few, notably
lead and manganese, separate in the form of dioxides upon the anode.
It is evidently important that the deposited material should be so
firmly adherent that it can be washed, dried, and weighed without
loss in handling. To secure these conditions it is essential that the
current density (that is, the amount of current per unit of area of
the electrodes) shall not be too high. In prescribing analytical
conditions it is customary to state the current strength in "normal
densities" expressed in amperes per 100 sq. cm. of electrode surface,
as, for example, "N.D_{100} = 2 amps."
If deposition occurs too rapidly, the deposit is likely to be spongy
or loosely adherent and falls off on subsequent treatment. This places
a practical limit to the current density to be employed, for a given
electrode surface. The cause of the unsatisfactory character of
the deposit is apparently sometimes to be found in the coincident
liberation of considerable hydrogen and sometimes in the failure of
the rapidly deposited material to form a continuous adherent surface.
The effect of rotating electrodes upon the character of the deposit is
referred to below.
The negative ions of an electrolyte are attracted to the anode and are
discharged on contact with it. Anions such as the chloride ion yield
chlorine atoms, from which gaseous chlorine molecules are formed
and escape. The radicals which compose such ions as NO_{3}^{-} or
SO_{4}^{--} are not capable of independent existence after discharge,
and break down into oxygen and N_{2}O_{5} and SO_{3} respectively. The
oxygen escapes and the anhydrides, reacting with water, re-form nitric
and sulphuric acids.
The law of Faraday expresses the relation between current strength and
the quantities of the decomposition products which, under constant
conditions, appear at the electrodes, namely, that a given quantity
of electricity, acting for a given time, causes the separation of
chemically equivalent quantities of the various elements or radicals.
For example, since 107.94 grams of silver is equivalent to 1.008 grams
of hydrogen, and that in turn to 8 grams of oxygen, or 31.78 grams of
copper, the quantity of electricity which will cause the deposit of
107.94 grams of silver in a given time will also separate the weights
just indicated of the other substances. Experiments show that a
current of one ampere passing for one second, i.e., a coulomb of
electricity, causes the deposition of 0.001118 gram of silver from a
normal solution of a silver salt. The number of coulombs required to
deposit 107.94 grams is 107.94/0.001118 or 96,550 and the same number
of coulombs will also cause the separation of 1.008 grams of hydrogen,
8 grams of oxygen or 31.78 grams of copper. While it might at first
appear that Faraday's law could thus be used as a basis for the
calculation of the time required for the deposition of a given
quantity of an electrolyte from solution, it must be remembered that
the law expresses what occurs when the concentration of the ions in
the solution is kept constant, as, for example, when the anode in
a silver salt solution is a plate of metallic silver. Under the
conditions of electro-analysis the concentration of the ions is
constantly diminishing as deposition proceeds and the time actually
required for complete deposition of a given weight of material by
a current of constant strength is, therefore, greater than that
calculated on the basis of the law as stated above.
The electrodes employed in electro-analysis are almost exclusively
of platinum, since that metal alone satisfactorily resists chemical
action of the electrolytes, and can be dried and weighed without
change in composition. The platinum electrodes may be used in the
form of dishes, foil or gauze. The last, on account of the ease of
circulation of the electrolyte, its relatively large surface in
proportion to its weight and the readiness with which it can be washed
and dried, is generally preferred.
Many devices have been described by the use of which the electrode
upon which deposition occurs can be mechanically rotated. This has an
effect parallel to that of greatly increasing the electrode surface
and also provides a most efficient means of stirring the solution.
With such an apparatus the amperage may be increased to 5 or even 10
amperes and depositions completed with great rapidity and accuracy. It
is desirable, whenever practicable, to provide a rotating or stirring
device, since, for example, the time consumed in the deposition of the
amount of copper usually found in analysis may be reduced from the
20 to 24 hours required with stationary electrodes, and unstirred
solutions, to about one half hour.
DETERMINATION OF COPPER AND LEAD
PROCEDURE.--Weigh out two portions of about 0.5 gram each (Note 1)
into tall, slender lipless beakers of about 100 cc. capacity. Dissolve
the metal in a solution of 5 cc. of dilute nitric acid (sp. gr. 1.20)
and 5 cc. of water, heating gently, and keeping the beaker covered.
When the sample has all dissolved (Note 2), wash down the sides of the
beaker and the bottom of the watch-glass with water and dilute the
solution to about 50 cc. Carefully heat to boiling and boil for a
minute or two to expel nitrous fumes.
Meanwhile, four platinum electrodes, two anodes and two cathodes,
should be cleaned by dipping in dilute nitric acid, washing with water
and finally with 95 per cent alcohol (Note 3). The alcohol may be
ignited and burned off. The electrodes are then cooled in a desiccator
and weighed. Connect the electrodes with the binding posts (or other
device for connection with the electric circuit) in such a way that
the copper will be deposited upon the electrode with the larger
surface, which is made the cathode. The beaker containing the solution
should then be raised into place from below the electrodes until the
latter reach nearly to the bottom of the beaker. The support for the
beaker must be so arranged that it can be easily raised or lowered.
If the electrolytic apparatus is provided with a mechanism for the
rotation of the electrode or stirring of the electrolyte, proceed as
follows: Arrange the resistance in the circuit to provide a direct
current of about one ampere. Pass this current through the solution
to be electrolyzed, and start the rotating mechanism. Keep the beaker
covered as completely as possible, using a split watch-glass (or other
device) to avoid loss by spattering. When the solution is colorless,
which is usually the case after about 35 minutes, rinse off the cover
glass, wash down the sides of the beaker, add about 0.30 gram of urea
and continue the electrolysis for another five minutes (Notes 4 and
5).
If stationary electrodes are employed, the current strength should be
about 0.1 ampere, which may, after 12 to 15 hours, be increased to 0.2
ampere. The time required for complete deposition is usually from 20
to 24 hours. It is advisable to add 5 cc. of nitric acid (sp. gr. 1.2)
if the electrolysis extends over this length of time. No urea is added
in this case.
When the deposition of the copper appears to be complete, stop the
rotating mechanism and slowly lower the beaker with the left hand,
directing at the same time a stream of water from a wash bottle on
both electrodes. Remove the beaker, shut off the current, and, if
necessary, complete the washing of the electrodes (Note 6). Rinse the
electrodes cautiously with alcohol and heat them in a hot closet until
the alcohol has just evaporated, but no longer, since the copper is
likely to oxidize at the higher temperature. (The alcohol may be
removed by ignition if care is taken to keep the electrodes in motion
in the air so that the copper deposit is not too strongly heated at
any one point.)
Test the solution in the beaker for copper as follows, remembering
that it is to be used for subsequent determinations of iron and zinc:
Remove about 5 cc. and add a slight excess of ammonia. Compare the
mixture with some distilled water, holding both above a white surface.
The solution should not show any tinge of blue. If the presence of
copper is indicated, add the test portion to the main solution,
evaporate the whole to a volume of about 100 cc., and again
electrolyze with clean electrodes (Note 7).
After cooling the electrodes in a desiccator, weigh them and from the
weight of copper on the cathode and of lead dioxide (PbO_{2}) on the
anode, calculate the percentage of copper (Cu) and of lead (Pb) in the
brass.
[Note 1: It is obvious that the brass taken for analysis should be
untarnished, which can be easily assured, when wire is used, by
scouring with emery. If chips or borings are used, they should be well
mixed, and the sample for analysis taken from different parts of the
mixture.]
[Note 2: If a white residue remains upon treatment of the alloy with
nitric acid, it indicates the presence of tin. The material is not,
therefore, a true brass. This may be treated as follows: Evaporate the
solution to dryness, moisten the residue with 5 cc. of dilute nitric
acid (sp. gr. 1.2) and add 50 cc. of hot water. Filter off the
meta-stannic acid, wash, ignite in porcelain and weigh as SnO_{2}.
This oxide is never wholly free from copper and must be purified for
an exact determination. If it does not exceed 2 per cent of the alloy,
the quantity of copper which it contains may usually be neglected.]
[Note 3: The electrodes should be freed from all greasy matter before
using, and those portions upon which the metal will deposit should not
be touched with the fingers after cleaning.]
[Note 4: Of the ions in solution, the H^{+}, Cu^{++}, Zn^{++}, and
Fe^{+++} ions tend to move toward the cathode. The NO_{3}^{-} ions and
the lead, probably in the form of PbO_{2}^{--} ions, move toward the
anode. At the cathode the Cu^{++} ions are discharged and plate out as
metallic copper. This alone occurs while the solution is relatively
concentrated. Later on, H^{+} ions are also discharged. In the
presence of considerable quantities of H^{+} ions, as in this acid
solution, no Zn^{++} or Fe^{+++} ions are discharged because of their
greater deposition potentials. At the anode the lead is deposited as
PbO_{2} and oxygen is evolved.
For the reasons stated on page 141 care must be taken that the
solution does not become alkaline if the electrolysis is long
continued.]
[Note 5: Urea reacts with nitrous acid, which may be formed in the
solution as a result of the reducing action of the liberated hydrogen.
Its removal promotes the complete precipitation of the copper. The
reaction is
CO(NH_{2})_{2} + 2HNO_{2} --> CO_{2} + 2N_{2} + 3H_{2}O.]
[Note 6: The electrodes must be washed nearly or quite free from
the nitric acid solution before the circuit is broken to prevent
re-solution of the copper.
If several solutions are connected in the same circuit it is obvious
that some device must be used to close the circuit as soon as the
beaker is removed.]
[Note 7: The electrodes upon which the copper has been deposited
may be cleaned by immersion in warm nitric acid. To remove the lead
dioxide, add a few crystals of oxalic acid to the nitric acid.]
DETERMINATION OF IRON
Most brasses contain small percentages of iron (usually not over 0.1
per cent) which, unless removed, is precipitated as phosphate and
weighed with the zinc.
PROCEDURE.--To the solution from the precipitation of copper and
lead by electrolysis, add dilute ammonia (sp. gr. 0.96) until the
precipitate of zinc hydroxide which first forms re-dissolves, leaving
only a slight red precipitate of ferric hydroxide. Filter off the
iron precipitate, using a washed filter, and wash five times with hot
water. Test a portion of the last washing with a dilute solution of
ammonium sulphide to assure complete removal of the zinc.
The precipitate may then be ignited and weighed as ferric oxide, as
described on page 110.
Calculate the percentage of iron (Fe) in the brass.
DETERMINATION OF ZINC
PROCEDURE.--Acidify the filtrate from the iron determination with
dilute nitric acid. Concentrate it to 150 cc. Add to the cold solution
dilute ammonia (sp. gr. 0.96) cautiously until it barely smells of
ammonia; then add !one drop! of a dilute solution of litmus (Note 1),
and drop in, with the aid of a dropper, dilute nitric acid until the
blue of the litmus just changes to red. It is important that this
point should not be overstepped. Heat the solution nearly to boiling
and pour into it slowly a filtered solution of di-ammonium hydrogen
phosphate[1] containing a weight of the phosphate about equal
to twelve times that of the zinc to be precipitated. (For this
calculation the approximate percentage of zinc is that found by
subtracting the sum of the percentages of the copper, lead and iron
from 100 per cent.) Keep the solution just below boiling for fifteen
minutes, stirring frequently (Note 2). If at the end of this time the
amorphous precipitate has become crystalline, allow the solution to
cool for about four hours, although a longer time does no harm (Note
3), and filter upon an asbestos filter in a porcelain Gooch crucible.
The filter is prepared as described on page 103, and should be dried
to constant weight at 105°C.
[Footnote 1: The ammonium phosphate which is commonly obtainable
contains some mono-ammonium salt, and this is not satisfactory as a
precipitant. It is advisable, therefore, to weigh out the amount of
the salt required, dissolve it in a small volume of water, add a drop
of phenolphthalein solution, and finally add dilute ammonium hydroxide
solution cautiously until the solution just becomes pink, but do not
add an excess.]
Wash the precipitate until free from sulphates with a warm 1 per cent
solution of the di-ammonium phosphate, and then five times with 50 per
cent alcohol (Note 4). Dry the crucible and precipitate for an hour at
105°C., and finally to constant weight (Note 5). The filtrate should
be made alkaline with ammonia and tested for zinc with a few drops of
ammonium sulphide, allowing it to stand (Notes 6, 7 and 8).
From the weight of the zinc ammonium phosphate (ZnNH_{4}PO_{4})
calculate the percentage of the zinc (Zn) in the brass.
[Note 1: The zinc ammonium phosphate is soluble both in acids and in
ammonia. It is, therefore, necessary to precipitate the zinc in a
nearly neutral solution, which is more accurately obtained by adding
a drop of a litmus solution to the liquid than by the use of litmus
paper.]
[Note 2: The precipitate which first forms is amorphous, and may have
a variable composition. On standing it becomes crystalline and then
has the composition ZnNH_{4}PO_{4}. The precipitate then settles
rapidly and is apt to occasion "bumping" if the solution is heated to
boiling. Stirring promotes the crystallization.]
[Note 3: In a carefully neutralized solution containing a considerable
excess of the precipitant, and also ammonium salts, the separation
of the zinc is complete after standing four hours. The ionic changes
connected with the precipitation of the zinc as zinc ammonium
phosphate are similar to those described for magnesium ammonium
phosphate, except that the zinc precipitate is soluble in an excess of
ammonium hydroxide, probably as a result of the formation of complex
ions of the general character Zn(NH_{3})_{4}^{++}.]
[Note 4: The precipitate is washed first with a dilute solution of the
phosphate to prevent a slight decomposition of the precipitate (as a
result of hydrolysis) if hot water alone is used. The alcohol is added
to the final wash-water to promote the subsequent drying.]
[Note 5: The precipitate may be ignited and weighed as
Zn_{2}P_{2}O_{7}, by cautiously heating the porcelain Gooch crucible
within a nickel or iron crucible, used as a radiator. The heating
must be very slow at first, as the escaping ammonia may reduce the
precipitate if it is heated too quickly.]
[Note 6: If the ammonium sulphide produced a distinct precipitate,
this should be collected on a small filter, dissolved in a few cubic
centimeters of dilute nitric acid, and the zinc reprecipitated as
phosphate, filtered off, dried, and weighed, and the weight added to
that of the main precipitate.]
[Note 7: It has been found that some samples of asbestos are acted
upon by the phosphate solution and lose weight. An error from this
source may be avoided by determining the weight of the crucible
and filter after weighing the precipitate. For this purpose the
precipitate may be dissolved in dilute nitric acid, the asbestos
washed thoroughly, and the crucible reweighed.]
[Note 8. The details of this method of precipitation of zinc are fully
discussed in an article by Dakin, !Ztschr. Anal. Chem.!, 39 (1900),
273.]
DETERMINATION OF SILICA IN SILICATES
Of the natural silicates, or artificial silicates such as slags and
some of the cements, a comparatively few can be completely decomposed
by treatment with acids, but by far the larger number require fusion
with an alkaline flux to effect decomposition and solution
for analysis. The procedure given below applies to silicates
undecomposable by acids, of which the mineral feldspar is taken as a
typical example. Modifications of the procedure, which are applicable
to silicates which are completely or partially decomposable by acids,
are given in the Notes on page 155.
PREPARATION OF THE SAMPLE
Grind about 3 grams of the mineral in an agate mortar (Note 1) until
no grittiness is to be detected, or, better, until it will entirely
pass through a sieve made of fine silk bolting cloth. The sieve may be
made by placing a piece of the bolting cloth over the top of a small
beaker in which the ground mineral is placed, holding the cloth in
place by means of a rubber band below the lip of the beaker. By
inverting the beaker over clean paper and gently tapping it, the fine
particles pass through the sieve, leaving the coarser particles within
the beaker. These must be returned to the mortar and ground, and the
process of sifting and grinding repeated until the entire sample
passes through the sieve.
[Note 1: If the sample of feldspar for analysis is in the massive or
crystalline form, it should be crushed in an iron mortar until the
pieces are about half the size of a pea, and then transferred to a
steel mortar, in which they are reduced to a coarse powder. A wooden
mallet should always be used to strike the pestle of the steel mortar,
and the blows should not be sharp.
It is plain that final grinding in an agate mortar must be continued
until the whole of the portion of the mineral originally taken has
been ground so that it will pass the bolting cloth, otherwise the
sifted portion does not represent an average sample, the softer
ingredients, if foreign matter is present, being first reduced to
powder. For this reason it is best to start with not more than the
quantity of the feldspar needed for analysis. The mineral must be
thoroughly mixed after the grinding.]
FUSION AND SOLUTION
PROCEDURE.--Weigh into platinum crucibles two portions of the ground
feldspar of about 0.8 gram each. Weigh on rough balances two portions
of anhydrous sodium carbonate, each amounting to about six times the
weight of the feldspar taken for analysis (Note 1). Pour about three
fourths of the sodium carbonate into the crucible, place the latter on
a piece of clean, glazed paper, and thoroughly mix the substance and
the flux by carefully stirring for several minutes with a dry glass
rod, the end of which has been recently heated and rounded in a flame
and slowly cooled. The rod may be wiped off with a small fragment of
filter paper, which may be placed in the crucible. Place the remaining
fourth of the carbonate on the top of the mixture. Cover the crucible,
heat it to dull redness for five minutes, and then gradually increase
the heat to the full capacity of a Bunsen or Tirrill burner for
twenty minutes, or until a quiet, liquid fusion is obtained (Note 2).
Finally, heat the sides and cover strongly until any material which
may have collected upon them is also brought to fusion.
Allow the crucible to cool, and remove the fused mass as directed on
page 116. Disintegrate the mass by placing it in a previously prepared
mixture of 100 cc. of water and 50 cc. of dilute hydrochloric acid
(sp. gr. 1.12) in a covered casserole (Note 3). Clean the crucible and
lid by means of a little hydrochloric acid, adding this acid to the
main solution (Notes 4 and 5).
[Note 1: Quartz, and minerals containing very high percentages of
silica, may require eight or ten parts by weight of the flux to insure
a satisfactory decomposition.]
[Note 2: During the fusion the feldspar, which, when pure, is a
silicate of aluminium and either sodium or potassium, but usually
contains some iron, calcium, and magnesium, is decomposed by the
alkaline flux. The sodium of the latter combines with the silicic acid
of the silicate, with the evolution of carbon dioxide, while about two
thirds of the aluminium forms sodium aluminate and the remainder
is converted into basic carbonate, or the oxide. The calcium and
magnesium, if present, are changed to carbonates or oxides.
The heat is applied gently to prevent a too violent reaction when
fusion first takes place.]
[Note 3: The solution of a silicate by a strong acid is the result of
the combination of the H^{+} ions of the acid and the silicate ions
of the silicate to form a slightly ionized silicic acid. As a
consequence, the concentration of the silicate ions in the solution is
reduced nearly to zero, and more silicate dissolves to re-establish
the disturbed equilibrium. This process repeats itself until all of
the silicate is brought into solution.
Whether the resulting solution of the silicate contains ortho-silicic
acid (H_{4}SiO_{4}) or whether it is a colloidal solution of some
other less hydrated acid, such as meta-silicic acid (H_{2}SiO_{3}),
is a matter that is still debatable. It is certain, however, that the
gelatinous material which readily separates from such solutions is of
the nature of a hydrogel, that is, a colloid which is insoluble in
water. This substance when heated to 100°C., or higher, is completely
dehydrated, leaving only the anhydride, SiO_{2}. The changes may be
represented by the equation:
SiO_{3}^{--} + 2H^{+} --> [H_{2}SiO_{3}] --> H_{2}O + SiO_{2}.]
[Note 4: A portion of the fused mass is usually projected upward by
the escaping carbon dioxide during the fusion. The crucible must
therefore be kept covered as much as possible and the lid carefully
cleaned.]
[Note 5: A gritty residue remaining after the disintegration of
the fused mass by acid indicates that the substance has been but
imperfectly decomposed. Such a residue should be filtered, washed,
dried, ignited, and again fused with the alkaline flux; or, if the
quantity of material at hand will permit, it is better to reject the
analysis, and to use increased care in grinding the mineral and in
mixing it with the flux.]
DEHYDRATION AND FILTRATION
PROCEDURE.--Evaporate the solution of the fusion to dryness, stirring
frequently until the residue is a dry powder. Moisten the residue with
5 cc. of strong hydrochloric acid (sp. gr. 1.20) and evaporate again
to dryness. Heat the residue for at least one hour at a temperature
of 110°C. (Note 1). Again moisten the residue with concentrated
hydrochloric acid, warm gently, making sure that the acid comes into
contact with the whole of the residue, dilute to about 200 cc. and
bring to boiling. Filter off the silica without much delay (Note 2),
and wash five times with warm dilute hydrochloric acid (one part
dilute acid (1.12 sp. gr.) to three parts of water). Allow the filter
to drain for a few moments, then place a clean beaker below the funnel
and wash with water until free from chlorides, discarding these
washings. Evaporate the original filtrate to dryness, dehydrate at
110°C. for one hour (Note 3), and proceed as before, using a second
filter to collect the silica after the second dehydration. Wash this
filter with warm, dilute hydrochloric acid (Note 4), and finally with
hot water until free from chlorides.
[Note 1: The silicic acid must be freed from its combination with
a base (sodium, in this instance) before it can be dehydrated.
The excess of hydrochloric acid accomplishes this liberation. By
disintegrating the fused mass with a considerable volume of dilute
acid the silicic acid is at first held in solution to a large extent.
Immediate treatment of the fused mass with strong acid is likely
to cause a semi-gelatinous silicic acid to separate at once and to
inclose alkali salts or alumina.
A flocculent residue will often remain after the decomposition of the
fused mass is effected. This is usually partially dehydrated silicic
acid and does not require further treatment at this point. The
progress of the dehydration is indicated by the behavior of the
solution, which as evaporation proceeds usually gelatinizes. On this
account it is necessary to allow the solution to evaporate on a steam
bath, or to stir it vigorously, to avoid loss by spattering.]
[Note 2: To obtain an approximately pure silica, the residue after
evaporation must be thoroughly extracted by warming with hydrochloric
acid, and the solution freely diluted to prevent, as far as possible,
the inclosure of the residual salts in the particles of silica. The
filtration should take place without delay, as the dehydrated silica
slowly dissolves in hydrochloric acid on standing.]
[Note 3: It has been shown by Hillebrand that silicic acid cannot be
completely dehydrated by a single evaporation and heating, nor by
several such treatments, unless an intermediate filtration of the
silica occurs. If, however, the silica is removed and the filtrates
are again evaporated and the residue heated, the amount of silica
remaining in solution is usually negligible, although several
evaporations and filtrations are required with some silicates to
insure absolute accuracy.
It is probable that temperatures above 100°C. are not absolutely
necessary to dehydrate the silica; but it is recommended, as tending
to leave the silica in a better condition for filtration than when
the lower temperature of the water bath is used. This, and many other
points in the analysis of silicates, are fully discussed by Dr.
Hillebrand in the admirable monograph on "The Analysis of Silicate and
Carbonate Rocks," Bulletin No. 700 of the United States Geological
Survey.
The double evaporation and filtration spoken of above are essential
because of the relatively large amount of alkali salts (sodium
chloride) present after evaporation. For the highest accuracy in the
determination of silica, or of iron and alumina, it is also necessary
to examine for silica the precipitate produced in the filtrate by
ammonium hydroxide by fusing it with acid potassium sulphate and
solution of the fused mass in water. The insoluble silica is filtered,
washed, and weighed, and the weight added to the weight of silica
previously obtained.]
[Note 4: Aluminium and iron are likely to be thrown down as basic
salts from hot, very dilute solutions of their chlorides, as a result
of hydrolysis. If the silica were washed only with hot water, the
solution of these chlorides remaining in the filter after the passage
of the original filtrate would gradually become so dilute as to throw
down basic salts within the pores of the filter, which would remain
with the silica. To avoid this, an acid wash-water is used until the
aluminium and iron are practically removed. The acid is then removed
by water.]
IGNITION AND TESTING OF SILICA
PROCEDURE.--Transfer the two washed filters belonging to each
determination to a platinum crucible, which need not be previously
weighed, and burn off the filter (Note 1). Ignite for thirty minutes
over the blast lamp with the cover on the crucible, and then for
periods of ten minutes, until the weight is constant.
When a constant weight has been obtained, pour into the crucible about
3 cc. of water, and then 3 cc. of hydrofluoric acid. !This must be
done in a hood with a good draft and great care must be taken not to
come into contact with the acid or to inhale its fumes (Note 2!).
If the precipitate has dissolved in this quantity of acid, add two
drops of concentrated sulphuric acid, and heat very slowly (always
under the hood) until all the liquid has evaporated, finally igniting
to redness. Cool in a desiccator, and weigh the crucible and residue.
Deduct this weight from the previous weight of crucible and impure
silica, and from the difference calculate the percentage of silica in
the sample (Note 3).
[Note 1: The silica undergoes no change during the ignition beyond the
removal of all traces of water; but Hillebrand (!loc. cit.!) has shown
that the silica holds moisture so tenaciously that prolonged ignition
over the blast lamp is necessary to remove it entirely. This finely
divided, ignited silica tends to absorb moisture, and should be
weighed quickly.]
[Note 2: Notwithstanding all precautions, the ignited precipitate of
silica is rarely wholly pure. It is tested by volatilisation of the
silica as silicon fluoride after solution in hydrofluoric acid, and,
if the analysis has been properly conducted, the residue, after
treatment with the acids and ignition, should not exceed 1 mg.
The acid produces ulceration if brought into contact with the skin,
and its fumes are excessively harmful if inhaled.]
[Note 3: The impurities are probably weighed with the original
precipitate in the form of oxides. The addition of the sulphuric
acid displaces the hydrofluoric acid, and it may be assumed that the
resulting sulphates (usually of iron or aluminium) are converted to
oxides by the final ignition.
It is obvious that unless the sulphuric and hydrofluoric acids used
are known to leave no residue on evaporation, a quantity equal to that
employed in the analysis must be evaporated and a correction applied
for any residue found.]
[Note 4: If the silicate to be analyzed is shown by a previous
qualitative examination to be completely decomposable, it may be
directly treated with hydrochloric acid, the solution evaporated to
dryness, and the silica dehydrated and further treated as described in
the case of the feldspar after fusion.
A silicate which gelatinizes on treatment with acids should be mixed
first with a little water, and the strong acid added in small portions
with stirring, otherwise the gelatinous silicic acid incloses
particles of the original silicate and prevents decomposition. The
water, by separating the particles and slightly lessening the rapidity
of action, prevents this difficulty. This procedure is one which
applies in general to the solution of fine mineral powders in acids.
If a small residue remains undecomposed by the treatment of the
silicate with acid, this may be filtered, washed, ignited and fused
with sodium carbonate and a solution of the fused mass added to the
original acid solution. This double procedure has an advantage, in
that it avoids adding so large a quantity of sodium salts as is
required for disintegration of the whole of the silicate by the fusion
method.]