An Introductory Course of Quantitative Chemical Analysis
by
Henry P. Talbot
Part 3 out of 5
the weight of silver chloride thus obtained calculate the percentage
of chlorine in the sample of sodium chloride.
[Note 1: The washed asbestos for this type of filter is prepared by
digesting in concentrated hydrochloric acid, long-fibered asbestos
which has been cut in pieces of about 0.5 cm. in length. After
digestion, the asbestos is filtered off on a filter plate and washed
with hot, distilled water until free from chlorides. A small portion
of the asbestos is shaken with water, forming a thin suspension, which
is bottled and kept for use.]
[Note 2: The nitric acid is added before precipitation to lessen the
tendency of the silver chloride to carry down with it other substances
which might be precipitated from a neutral solution. A large excess of
the acid would exert a slight solvent action upon the chloride.]
[Note 3: The solution should not be boiled after the addition of the
nitric acid before the presence of an excess of silver nitrate is
assured, since a slight interaction between the nitric acid and the
sodium chloride is possible, by which a loss of chlorine, either as
such or as hydrochloric acid, might ensue. The presence of an excess
of the precipitant can usually be recognized at the time of its
addition, by the increased readiness with which the precipitate
coagulates and settles.]
[Note 4: The precipitate should not be exposed to strong sunlight,
since under those conditions a reduction of the silver chloride ensues
which is accompanied by a loss of chlorine. The superficial alteration
which the chloride undergoes in diffused daylight is not sufficient
to materially affect the accuracy of the determination. It should be
noted, however, that a slight error does result from the effect of
light upon the silver chloride precipitate and in cases in which the
greatest obtainable accuracy is required, the procedure described
under "Method B" should be followed, in which this slight reduction of
the silver chloride is corrected by subsequent treatment with nitric
and hydrochloric acids.]
[Note 5: The asbestos used in the Gooch filter should be of the finest
quality and capable of division into minute fibrous particles. A
coarse felt is not satisfactory.]
[Note 6: The precipitate must be washed with warm water until it is
absolutely free from silver and sodium nitrates. It may be assumed
that the sodium salt is completely removed when the wash-water shows
no evidence of silver. It must be borne in mind that silver chloride
is somewhat soluble in hydrochloric acid, and only a single drop
should be added. The washing should be continued until no cloudiness
whatever can be detected in 3 cc. of the washings.
Silver chloride is but slightly soluble in water. The solubility
varies with its physical condition within small limits, and is
about 0.0018 gram per liter at 18 deg.C. for the curdy variety usually
precipitated. The chloride is also somewhat soluble in solutions of
many chlorides, in solutions of silver nitrate, and in concentrated
nitric acid.
As a matter of economy, the filtrate, which contains whatever silver
nitrate was added in excess, may be set aside. The silver can be
precipitated as chloride and later converted into silver nitrate.]
[Note 7: The use of the Gooch filter commends itself strongly when a
considerable number of halogen determinations are to be made, since
successive portions of the silver halides may be filtered on the same
filter, without the removal of the preceding portions, until the
crucible is about two thirds filled. If the felt is properly prepared,
filtration and washing are rapidly accomplished on this filter, and
this, combined with the possibility of collecting several precipitates
on the same filter, is a strong argument in favor of its use with any
but gelatinous precipitates.]
!Method B. With the Use of a Paper Filter!
PROCEDURE.--Weigh out two portions of sodium chloride of about
0.25-0.3 gram each and proceed with the precipitation of the silver
chloride as described under Method A above. When the chloride is ready
for filtration prepare two 9 cm. washed paper filters (see Appendix).
Pour the liquid above the precipitates through the filters, wash twice
by decantation and transfer the precipitates to the filters, finally
washing them until free from silver solution as described. The funnel
should then be covered with a moistened filter paper by stretching it
over the top and edges, to which it will adhere on drying. It should
be properly labeled with the student's name and desk number, and then
placed in a drying closet, at a temperature of about 100-110 deg.C., until
completely dry.
The perfectly dry filter is then opened over a circular piece of
clean, smooth, glazed paper about six inches in diameter, placed upon
a larger piece about twelve inches in diameter. The precipitate is
removed from the filter as completely as possible by rubbing the sides
gently together, or by scraping them cautiously with a feather which
has been cut close to the quill and is slightly stiff (Note 1). In
either case, care must be taken not to rub off any considerable
quantity of the paper, nor to lose silver chloride in the form of
dust. Cover the precipitate on the glazed paper with a watch-glass to
prevent loss of fine particles and to protect it from dust from the
air. Fold the filter paper carefully, roll it into a small cone, and
wind loosely around !the top! a piece of small platinum wire (Note 2).
Hold the filter by the wire over a small porcelain crucible (which has
been cleaned, ignited, cooled in a desiccator, and weighed), ignite
it, and allow the ash to fall into the crucible. Place the crucible
upon a clean clay triangle, on its side, and ignite, with a low
flame well at its base, until all the carbon of the filter has been
consumed. Allow the crucible to cool, add two drops of concentrated
nitric acid and one drop of concentrated hydrochloric acid, and heat
!very cautiously!, to avoid spattering, until the acids have been
expelled; then transfer the main portion of the precipitate from the
glazed paper to the cooled crucible, placing the latter on the larger
piece of glazed paper and brushing the precipitate from the
smaller piece into it, sweeping off all particles belonging to the
determination.
Moisten the precipitate with two drops of concentrated nitric acid and
one drop of concentrated hydrochloric acid, and again heat with great
caution until the acids are expelled and the precipitate is white,
when the temperature is slowly raised until the silver chloride just
begins to fuse at the edges (Note 3). The crucible is then cooled in a
desiccator and weighed, after which the heating (without the addition
of acids) is repeated, and it is again weighed. This must be continued
until the weight is constant within 0.0003 gram in two consecutive
weighings. Deduct the weight of the crucible, and calculate the
percentage of chlorine in the sample of sodium chloride taken for
analysis.
[Note 1: The separation of the silver chloride from the filter is
essential, since the burning carbon of the paper would reduce a
considerable quantity of the precipitate to metallic silver, and its
complete reconversion to the chloride within the crucible, by means of
acids, would be accompanied by some difficulty. The small amount of
silver reduced from the chloride adhering to the filter paper after
separating the bulk of the precipitate, and igniting the paper
as prescribed, can be dissolved in nitric acid, and completely
reconverted to chloride by hydrochloric acid. The subsequent addition
of the two acids to the main portion of the precipitate restores the
chlorine to any chloride which may have been partially reduced by the
sunlight. The excess of the acids is volatilized by heating.]
[Note 2: The platinum wire is wrapped around the top of the filter
during its incineration to avoid contact with any reduced silver from
the reduction of the precipitate. If the wire were placed nearer the
apex, such contact could hardly be avoided.]
[Note 3: Silver chloride should not be heated to complete fusion,
since a slight loss by volatilization is possible at high
temperatures. The temperature of fusion is not always sufficient
to destroy filter shreds; hence these should not be allowed to
contaminate the precipitate.]
DETERMINATION OF IRON AND OF SULPHUR IN FERROUS AMMONIUM SULPHATE,
FESO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O
DETERMINATION OF IRON
PROCEDURE.--Weigh out into beakers (200-250 cc.) two portions of the
sample (Note 1) of about 1 gram each and dissolve these in 50 cc. of
water, to which 1 cc. of dilute hydrochloric acid (sp. gr. 1.12) has
been added (Note 2). Heat the solution to boiling, and while at the
boiling point add concentrated nitric acid (sp. gr. 1.42), !drop by
drop! (noting the volume used), until the brown coloration, which
appears after the addition of a part of the nitric acid, gives place
to a yellow or red (Note 3). Avoid a large excess of nitric acid, but
be sure that the action is complete. Pour this solution cautiously
into about 200 cc. of water, containing a slight excess of ammonia.
Calculate for this purpose the amount of aqueous ammonia required to
neutralize the hydrochloric and nitric acids added (see Appendix for
data), and also to precipitate the iron as ferric hydroxide from the
weight of the ferrous ammonium sulphate taken for analysis, assuming
it to be pure (Note 4). The volume thus calculated will be in excess
of that actually required for precipitation, since the acids are in
part consumed in the oxidation process, or are volatilized. Heat the
solution to boiling, and allow the precipitated ferric hydroxide to
settle. Decant the clear liquid through a washed filter (9 cm.),
keeping as much of the precipitate in the beaker as possible. Wash
twice by decantation with 100 cc. of hot water. Reserve the filtrate.
Dissolve the iron from the filter with hot, dilute hydrochloric acid
(sp. gr. 1.12), adding it in small portions, using as little as
possible and noting the volume used. Collect the solution in the
beaker in which precipitation took place. Add 1 cc. of nitric acid
(sp. gr. 1.42), boil for a few moments, and again pour into a
calculated excess of ammonia.
Wash the precipitate twice by decantation, and finally transfer it to
the original filter. Wash continuously with hot water until finally
3 cc. of the washings, acidified with nitric acid (Note 5), show
no evidences of the presence of chlorides when tested with silver
nitrate. The filtrate and washings are combined with those from the
first precipitation and treated for the determination of sulphur, as
prescribed on page 112.
[Note 1: If a selection of pure material for analysis is to be made,
crystals which are cloudy are to be avoided on account of loss of
water of crystallization; and also those which are red, indicating
the presence of ferric iron. If, on the other hand, the value of an
average sample of material is desired, it is preferable to grind the
whole together, mix thoroughly, and take a sample from the mixture for
analysis.]
[Note 2: When aqueous solutions of ferrous compounds are heated in the
air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in
the absence of free acid. The H^{+} and OH^{-} ions from water are
involved in the oxidation process and the result is, in effect, the
formation of some ferric hydroxide which tends to separate. Moreover,
at the boiling temperature, the ferric sulphate produced by the
oxidation hydrolyzes in part with the formation of a basic ferric
sulphate, which also tends to separate from solution. The addition of
the hydrochloric acid prevents the formation of ferric hydroxide, and
so far reduces the ionization of the water that the hydrolysis of the
ferric sulphate is also prevented, and no precipitation occurs on
heating.]
[Note 3: The nitric acid, after attaining a moderate strength,
oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an
intermediate nitroso-compound similar in character to that formed in
the "ring-test" for nitrates. The nitric oxide is driven out by heat,
and the solution then shows by its color the presence of ferric
compounds. A drop of the oxidized solution should be tested on
a watch-glass with potassium ferricyanide, to insure a complete
oxidation. This oxidation of the iron is necessary, since Fe^{++} ions
are not completely precipitated by ammonia.
The ionic changes which are involved in this oxidation are perhaps
most simply expressed by the equation
3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO,
the H^{+} ions coming from the acid in the solution, in this case
either the nitric or the hydrochloric acid. The full equation on which
this is based may be written thus:
6FeSO_{4} + 2HNO_{3} + 6HCl --> 2Fe_{2}(SO_{4})_{3} + 2FeCl_{3} + 2NO
+ 4H_{2}O,
assuming that only enough nitric acid is added to complete the
oxidation.]
[Note 4: The ferric hydroxide precipitate tends to carry down some
sulphuric acid in the form of basic ferric sulphate. This tendency is
lessened if the solution of the iron is added to an excess of OH^{-}
ions from the ammonium hydroxide, since under these conditions
immediate and complete precipitation of the ferric hydroxide ensues.
A gradual neutralization with ammonia would result in the local
formation of a neutral solution within the liquid, and subsequent
deposition of a basic sulphate as a consequence of a local deficiency
of OH^{-} ions from the NH_{4}OH and a partial hydrolysis of the
ferric salt. Even with this precaution the entire absence of sulphates
from the first iron precipitate is not assured. It is, therefore,
redissolved and again thrown down by ammonia. The organic matter of
the filter paper may occasion a partial reduction of the iron during
solution, with consequent possibility of incomplete subsequent
precipitation with ammonia. The nitric acid is added to reoxidize this
iron.
To avoid errors arising from the solvent action of ammoniacal
liquids upon glass, the iron precipitate should be filtered without
unnecessary delay.]
[Note 5: The washings from the ferric hydroxide are acidified with
nitric acid, before testing with silver nitrate, to destroy the
ammonia which is a solvent of silver chloride.
The use of suction to promote filtration and washing is permissible,
though not prescribed. The precipitate should not be allowed to dry
during the washing.]
!Ignition of the Iron Precipitate!
Heat a platinum or porcelain crucible, cool it in a desiccator and
weigh, repeating until a constant weight is obtained.
Fold the top of the filter paper over the moist precipitate of ferric
hydroxide and transfer it cautiously to the crucible. Wipe the inside
of the funnel with a small fragment of washed filter paper, if
necessary, and place the paper in the crucible.
Incline the crucible on its side, on a triangle supported on a
ring-stand, and stand the cover on edge at the mouth of the crucible.
Place a burner below the front edge of the crucible, using a low flame
and protecting it from drafts of air by means of a chimney. The heat
from the burner is thus reflected into the crucible and dries
the precipitate without danger of loss as the result of a sudden
generation of steam within the mass of ferric hydroxide. As the drying
progresses the burner may be gradually moved toward the base of the
crucible and the flame increased until the paper of the filter begins
to char and finally to smoke, as the volatile matter is expelled. This
is known as "smoking off" a filter, and the temperature should not be
raised sufficiently high during this process to cause the paper to
ignite, as the air currents produced by the flame of the blazing paper
may carry away particles of the precipitate.
When the paper is fully charred, move the burner to the base of the
crucible and raise the temperature to the full heat of the burner for
fifteen minutes, with the crucible still inclined on its side, but
without the cover (Note 1). Finally set the crucible upright in the
triangle, cover it, and heat at the full temperature of a blast lamp
or other high temperature burner. Cool and weigh in the usual manner
(Note 2). Repeat the strong heating until the weight is constant
within 0.0003 gram.
From the weight of ferric oxide (Fe_{2}O_{3}) calculate the percentage
of iron (Fe) in the sample (Note 3).
[Note 1: These directions for the ignition of the precipitate must be
closely followed. A ready access of atmospheric oxygen is of special
importance to insure the reoxidation to ferric oxide of any iron which
may be reduced to magnetic oxide (Fe_{3}O_{4}) during the combustion
of the filter. The final heating over the blast lamp is essential
for the complete expulsion of the last traces of water from the
hydroxide.]
[Note 2: Ignited ferric oxide is somewhat hygroscopic. On this account
the weighings must be promptly completed after removal from the
desiccator. In all weighings after the first it is well to place the
weights upon the balance-pan before removing the crucible from the
desiccator. It is then only necessary to move the rider to obtain the
weight.]
[Note 3: The gravimetric determination of aluminium or chromium is
comparable with that of iron just described, with the additional
precaution that the solution must be boiled until it contains but a
very slight excess of ammonia, since the hydroxides of aluminium and
chromium are more soluble than ferric hydroxide.
The most important properties of these hydroxides, from a quantitative
standpoint, other than those mentioned, are the following: All are
precipitable by the hydroxides of sodium and potassium, but always
inclose some of the precipitant, and should be reprecipitated with
ammonium hydroxide before ignition to oxides. Chromium and aluminium
hydroxides dissolve in an excess of the caustic alkalies and form
anions, probably of the formula AlO_2^{-} and CrO_{2}^{-}. Chromium
hydroxide is reprecipitated from this solution on boiling. When first
precipitated the hydroxides are all readily soluble in acids, but
aluminium hydroxide dissolves with considerable difficulty after
standing or boiling for some time. The precipitation of the hydroxides
is promoted by the presence of ammonium chloride, but is partially
or entirely prevented by the presence of tartaric or citric acids,
glycerine, sugars, and some other forms of soluble organic matter.
The hydroxides yield on ignition an oxide suitable for weighing
(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).]
DETERMINATION OF SULPHUR
PROCEDURE.--Add to the combined filtrates from the ferric hydroxide
about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and
then add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess
and evaporate to dryness on the water bath. Add 10 cc. of concentrated
hydrochloric acid (sp. gr. 1.20) to the residue, and again evaporate
to dryness on the bath. Dissolve the residue in water, filter if not
clear, transfer to a 700 cc. beaker, dilute to about 400 cc., and
cautiously add hydrochloric acid until the solution shows a distinctly
acid reaction (Note 1). Heat the solution to boiling, and add !very
slowly! and with constant stirring, 20 cc. in excess of the calculated
amount of a hot barium chloride solution, containing about 20 grams
BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling for
about two minutes, allow the precipitate to settle, and decant the
liquid at the end of half an hour (Note 4). Replace the beaker
containing the original filtrate by a clean beaker, wash the
precipitated sulphate by decantation with hot water, and subsequently
upon the filter until it is freed from chlorides, testing the washings
as described in the determination of iron. The filter is then
transferred to a platinum or porcelain crucible and ignited, as
described above, until the weight is constant (Note 5). From the
weight of barium sulphate (BaSO_{4}) obtained, calculate the
percentage of sulphur (S) in the sample.
[Note 1: Barium sulphate is slightly soluble in hydrochloric acid,
even dilute, probably as a result of the reduction in the degree of
dissociation of sulphuric acid in the presence of the H^{+} ions of
the hydrochloric acid, and possibly because of the formation of a
complex anion made up of barium and chlorine; hence only the smallest
excess should be added over the amount required to acidify the
solution.]
[Note 2: The ionic changes involved in the precipitation of barium
sulphate are very simple:
Ba^{++} + SO_{4}^{--} --> [BaSO_{4}]
This case affords one of the best illustrations of the effect of an
excess of a precipitant in decreasing the solubility of a precipitate.
If the conditions are considered which exist at the moment when just
enough of the Ba^{++} ions have been added to correspond to the
SO_{4}^{--} ions in the solution, it will be seen that nearly all of
the barium sulphate has been precipitated, and that the small amount
which then remains in the solution which is in contact with the
precipitate must represent a saturated solution for the existing
temperature, and that this solution is comparable with a solution of
sugar to which more sugar has been added than will dissolve. It
should be borne in mind that the quantity of barium sulphate in
this !saturated solution is a constant quantity! for the existing
conditions. The dissolved barium sulphate, like any electrolyte, is
dissociated, and the equilibrium conditions may be expressed thus:
(!Conc'n Ba^{++} x Conc'n SO_{4}^{--})/(Conc'n BaSO_{4}) = Const.!,
and since !Conc'n BaSO_{4}! for the saturated solution has a constant
value (which is very small), it may be eliminated, when the expression
becomes !Conc'n Ba^{++} x Conc'n SO_{4}^{--} = Const.!, which is
the "solubility product" of BaSO_{4}. If, now, an excess of the
precipitant, a soluble barium salt, is added in the form of a
relatively concentrated solution (the slight change of volume of a few
cubic centimeters may be disregarded for the present discussion)
the concentration of the Ba^{++} ions is much increased, and as a
consequence the !Conc'n SO_{4}! must decrease in proportion if the
value of the expression is to remain constant, which is a requisite
condition if the law of mass action upon which our argument depends
holds true. In other words, SO_{4}^{--} ions must combine with some
of the added Ba^{++} ions to form [BaSO_{4}]; but it will be recalled
that the solution is already saturated with BaSO_{4}, and this freshly
formed quantity must, therefore, separate and add itself to the
precipitate. This is exactly what is desired in order to insure
more complete precipitation and greater accuracy, and leads to the
conclusion that the larger the excess of the precipitant added the
more successful the analysis; but a practical limit is placed upon
the quantity of the precipitant which may be properly added by other
conditions, as stated in the following note.]
[Note 3: Barium sulphate, in a larger measure than most compounds,
tends to carry down other substances which are present in the solution
from which it separates, even when these other substances are
relatively soluble, and including the barium chloride used as the
precipitant. This is also notably true in the case of nitrates and
chlorates of the alkalies, and of ferric compounds; and, since in this
analysis ammonium nitrate has resulted from the neutralization of the
excess of the nitric acid added to oxidize the iron, it is essential
that this should be destroyed by repeated evaporation with a
relatively large quantity of hydrochloric acid. During evaporation a
mutual decomposition of the two acids takes place, and the nitric acid
is finally decomposed and expelled by the excess of hydrochloric acid.
Iron is usually found in the precipitate of barium sulphate when
thrown down from hot solutions in the presence of ferric salts. This,
according to Kuster and Thiel (!Zeit. anorg. Chem.!, 22, 424), is due
to the formation of a complex ion (Fe(SO_{4})_{2}) which precipitates
with the Ba^{++} ion, while Richards (!Zeit. anorg. Chem.!, 23, 383)
ascribes it to hydrolytic action, which causes the formation of a
basic ferric complex which is occluded in the barium precipitate.
Whatever the character of the compound may be, it has been shown that
it loses sulphuric anhydride upon ignition, causing low results, even
though the precipitate contains iron.
The contamination of the barium sulphate by iron is much less in the
presence of ferrous than ferric salts. If, therefore, the sulphur
alone were to be determined in the ferrous ammonium sulphate, the
precipitation by barium might be made directly from an aqueous
solution of the salt, which had been made slightly acid with
hydrochloric acid.]
[Note 4: The precipitation of the barium sulphate is probably complete
at the end of a half-hour, and the solution may safely be filtered at
the expiration of that time if it is desired to hasten the analysis.
As already noted, many precipitates of the general character of this
sulphate tend to grow more coarsely granular if digested for some time
with the liquid from which they have separated. It is therefore well
to allow the precipitate to stand in a warm place for several hours,
if practicable, to promote ease of filtration. The filtrate and
washings should always be carefully examined for minute quantities of
the sulphate which may pass through the pores of the filter. This is
best accomplished by imparting to the filtrate a gentle rotary motion,
when the sulphate, if present, will collect at the center of the
bottom of the beaker.]
[Note 5: A reduction of barium sulphate to the sulphide may very
readily be caused by the reducing action of the burning carbon of the
filter, and much care should be taken to prevent any considerable
reduction from this cause. Subsequent ignition, with ready access
of air, reconverts the sulphide to sulphate unless a considerable
reduction has occurred. In the latter case it is expedient to add one
or two drops of sulphuric acid and to heat cautiously until the excess
of acid is expelled.]
[Note 6: Barium sulphate requires about 400,000 parts of water for
its solution. It is not decomposed at a red heat but suffers loss,
probably of sulphur trioxide, at a temperature above 900 deg.C.]
DETERMINATION OF SULPHUR IN BARIUM SULPHATE
PROCEDURE.--Weigh out, into platinum crucibles, two portions of about
0.5 gram of the sulphate. Mix each in the crucible with five to six
times its weight of anhydrous sodium carbonate. This can best be done
by placing the crucible on a piece of glazed paper and stirring the
mixture with a clean, dry stirring-rod, which may finally be wiped off
with a small fragment of filter paper, the latter being placed in the
crucible. Cover the crucible and heat until a quiet, liquid fusion
ensues. Remove the burner, and tip the crucible until the fused mass
flows nearly to its mouth. Hold it in that position until the mass has
solidified. When cold, the material may usually be detached in a lump
by tapping the crucible or gently pressing it near its upper edge. If
it still adheres, a cubic centimeter or so of water may be placed in
the cold crucible and cautiously brought to boiling, when the cake
will become loosened and may be removed and placed in about 250 cc.
of hot, distilled water to dissolve. Clean the crucible completely,
rubbing the sides with a rubber-covered stirring-rod, if need be.
When the fused mass has completely disintegrated and nothing further
will dissolve, decant the solution from the residue of barium
carbonate (Note 1). Pour over the residue 20 cc. of a solution of
sodium carbonate and 10 cc. of water and heat to gentle boiling for
about three minutes (Note 2). Filter off the carbonate and wash it
with hot water, testing the slightly acidified washings for sulphate
and preserving any precipitates which appear in these tests. Acidify
the filtrate with hydrochloric acid until just acid, bring to boiling,
and slowly add hot barium chloride solution, as in the preceding
determination. Add also any tests from the washings in which
precipitates have appeared. Filter, wash, ignite, and weigh.
From the weight of barium sulphate, calculate the percentage of
sulphur (S) in the sample.
[Note 1: This alkaline fusion is much employed to disintegrate
substances ordinarily insoluble in acids into two components, one
of which is water soluble and the other acid soluble. The reaction
involved is:
BaSO_{4} + Na_{2}CO_{3}, --> BaCO_{3}, + Na_{2}SO_{4}.
As the sodium sulphate is soluble in water, and the barium carbonate
insoluble, a separation between them is possible and the sulphur can
be determined in the water-soluble portion.
It should be noted that this method can be applied to the purification
of a precipitate of barium sulphate if contaminated by most of the
substances mentioned in Note 3 on page 114. The impurities pass into
the water solution together with the sodium sulphate, but, being
present in such minute amounts, do not again precipitate with the
barium sulphate.]
[Note 2: The barium carbonate is boiled with sodium carbonate solution
before filtration because the reaction above is reversible; and it is
only by keeping the sodium carbonate present in excess until nearly
all of the sodium sulphate solution has been removed by filtration
that the reversion of some of the barium carbonate to barium sulphate
is prevented. This is an application of the principle of mass action,
in which the concentration of the reagent (the carbonate ion) is
kept as high as practicable and that of the sulphate ion as low as
possible, in order to force the reaction in the desired direction (see
Appendix).]
DETERMINATION OF PHOSPHORIC ANHYDRIDE IN APATITE
The mineral apatite is composed of calcium phosphate, associated with
calcium chloride, or fluoride. Specimens are easily obtainable which
are nearly pure and leave on treatment with acid only a slight
siliceous residue.
For the purpose of gravimetric determination, phosphoric acid is
usually precipitated from ammoniacal solutions in the form of
magnesium ammonium phosphate which, on ignition, is converted into
magnesium pyrophosphate. Since the calcium phosphate of the apatite
is also insoluble in ammoniacal solutions, this procedure cannot be
applied directly. The separation of the phosphoric acid from the
calcium must first be accomplished by precipitation in the form of
ammonium phosphomolybdate in nitric acid solution, using ammonium
molybdate as the precipitant. The "yellow precipitate," as it is often
called, is not always of a definite composition, and therefore not
suitable for direct weighing, but may be dissolved in ammonia, and the
phosphoric acid thrown out as magnesium ammonium phosphate from the
solution.
Of the substances likely to occur in apatite, silicic acid alone
interferes with the precipitation of the phosphoric acid in nitric
acid solution.
PRECIPITATION OF AMMONIUM PHOSPHOMOLYBDATE
PROCEDURE.--Grind the mineral in an agate mortar until no grit is
perceptible. Transfer the substance to a weighing-tube, and weigh out
two portions, not exceeding 0.20 gram each (Note 1) into two beakers
of about 200 cc. capacity. Pour over them 20 cc. of dilute nitric acid
(sp. gr. 1.2) and warm gently until solvent action has apparently
ceased. Evaporate the solution cautiously to dryness, heat the residue
for about an hour at 100-110 deg.C., and treat it again with nitric acid
as described above; separate the residue of silica by filtration on
a small filter (7 cm.) and wash with warm water, using as little as
possible (Note 2). Receive the filtrate in a beaker (200-500 cc.).
Test the washings with ammonia for calcium phosphate, but add all such
tests in which a precipitate appears to the original nitrate (Note 3).
The filtrate and washings must be kept as small as possible and should
not exceed 100 cc. in volume. Add aqueous ammonia (sp. gr. 0.96) until
the precipitate of calcium phosphate first produced just fails to
redissolve, and then add a few drops of nitric acid until this is
again brought into solution (Note 4). Warm the solution until it
cannot be comfortably held in the hand (about 60 deg.C.) and, after
removal of the burner, add 75 cc. of ammonium molybdate solution which
has been !gently! warmed, but which must be perfectly clear. Allow
the mixture to stand at a temperature of about 50 or 60 deg.C. for twelve
hours (Notes 5 and 6). Filter off the yellow precipitate on a 9 cm.
filter, and wash by decantation with a solution of ammonium nitrate
made acid with nitric acid.[1] Allow the precipitate to remain in the
beaker as far as possible. Test the washings for calcium with ammonia
and ammonium oxalate (Note 3).
[Footnote 1: This solution is prepared as follows: Mix 100 cc. of
ammonia solution (sp. gr. 0.96) with 325 cc. of nitric acid (sp. gr.
1.2) and dilute with 100 cc. of water.]
Add 10 cc. of molybdate solution to the nitrate, and leave it for
a few hours. It should then be carefully examined for a !yellow!
precipitate; a white precipitate may be neglected.
[Note 1: Magnesium ammonium phosphate, as noted below, is slightly
soluble under the conditions of operation. Consequently the
unavoidable errors of analysis are greater in this determination than
in those which have preceded it, and some divergence may be expected
in duplicate analyses. It is obvious that the larger the amount of
substance taken for analysis the less will be the relative loss or
gain due to unavoidable experimental errors; but, in this instance, a
check is placed upon the amount of material which may be taken both by
the bulk of the resulting precipitate of ammonium phosphomolybdate
and by the excessive amount of ammonium molybdate required to effect
complete separation of the phosphoric acid, since a liberal excess
above the theoretical quantity is demanded. Molybdic acid is one of
the more expensive reagents.]
[Note 2: Soluble silicic acid would, if present, partially separate
with the phosphomolybdate, although not in combination with
molybdenum. Its previous removal by dehydration is therefore
necessary.]
[Note 3: When washing the siliceous residue the filtrate may be tested
for calcium by adding ammonia, since that reagent neutralizes the
acid which holds the calcium phosphate in solution and causes
precipitation; but after the removal of the phosphoric acid in
combination with the molybdenum, the addition of an oxalate is
required to show the presence of calcium.]
[Note 4: An excess of nitric acid exerts a slight solvent
action, while ammonium nitrate lessens the solubility; hence the
neutralization of the former by ammonia.]
[Note 5: The precipitation of the phosphomolybdate takes place more
promptly in warm than in cold solutions, but the temperature should
not exceed 60 deg.C. during precipitation; a higher temperature tends to
separate molybdic acid from the solution. This acid is nearly white,
and its deposition in the filtrate on long standing should not be
mistaken for a second precipitation of the yellow precipitate. The
addition of 75 cc. of ammonium molybdate solution insures the presence
of a liberal excess of the reagent, but the filtrate should be tested
as in all quantitative procedures.
The precipitation is probably complete in many cases in less than
twelve hours; but it is better, when practicable, to allow the
solution to stand for this length of time. Vigorous shaking or
stirring promotes the separation of the precipitate.]
[Note 6: The composition of the "yellow precipitate" undoubtedly
varies slightly with varying conditions at the time of its formation.
Its composition may probably fairly be represented by the formula,
(NH_{4})_{3}PO_{4}.12MoO_{3}.H_{2}O, when precipitated under the
conditions prescribed in the procedure. Whatever other variations may
occur in its composition, the ratio of 12 MoO_{3}:1 P seems to
hold, and this fact is utilized in volumetric processes for the
determination of phosphorus, in which the molybdenum is reduced to
a lower oxide and reoxidized by a standard solution of potassium
permanganate. In principle, the procedure is comparable with that
described for the determination of iron by permanganate.]
PRECIPITATION OF MAGNESIUM AMMONIUM PHOSPHATE
PROCEDURE.--Dissolve the precipitate of phosphomolybdate upon the
filter by pouring through it dilute aqueous ammonia (one volume of
dilute ammonia (sp. gr. 0.96) and three volumes of water, which
should be carefully measured), and receive the solution in the beaker
containing the bulk of the precipitate. The total volume of nitrate
and washings should not much exceed 100 cc. Acidify the solution with
dilute hydrochloric acid, and heat it nearly to boiling. Calculate the
volume of magnesium ammonium chloride solution ("magnesia mixture")
required to precipitate the phosphoric acid, assuming 40 per cent
P_{2}O_{5} in the apatite. Measure out about 5 cc. in excess of this
amount, and pour it into the acid solution. Then add slowly dilute
ammonium hydroxide (1 volume of strong ammonia (sp. gr. 0.90) and 9
volumes of water), stirring constantly until a precipitate forms. Then
add a volume of filtered, concentrated ammonia (sp. gr. 0.90) equal to
one third of the volume of liquid in the beaker (Note 1). Allow the
whole to cool. The precipitated magnesium ammonium phosphate should
then be definitely crystalline in appearance (Note 2). (If it is
desired to hasten the precipitation, the solution may be cooled, first
in cold and then in ice-water, and stirred !constantly! for half an
hour, when precipitation will usually be complete.)
Decant the clear liquid through a filter, and transfer the precipitate
to the filter, using as wash-water a mixture of one volume of
concentrated ammonia and three volumes of water. It is not necessary
to clean the beaker completely or to wash the precipitate thoroughly
at this point, as it is necessary to purify it by reprecipitation.
[Note 1: Magnesium ammonium phosphate is not a wholly insoluble
substance, even under the most favorable analytical conditions. It
is least soluble in a liquid containing one fourth of its volume of
concentrated aqueous ammonia (sp. gr. 0.90) and this proportion should
be carefully maintained as prescribed in the procedure. On account of
this slight solubility the volume of solutions should be kept as small
as possible and the amount of wash-water limited to that absolutely
required.
A large excess of the magnesium solution tends both to throw out
magnesium hydroxide (shown by a persistently flocculent precipitate)
and to cause the phosphate to carry down molybdic acid. The tendency
of the magnesium precipitate to carry down molybdic acid is also
increased if the solution is too concentrated. The volume should not
be less than 90 cc., nor more than 125 cc., at the time of the first
precipitation with the magnesia mixture.]
[Note 2: The magnesium ammonium phosphate should be perfectly
crystalline, and will be so if the directions are followed. The slow
addition of the reagent is essential, and the stirring not less so.
Stirring promotes the separation of the precipitate and the formation
of larger crystals, and may therefore be substituted for digestion in
the cold. The stirring-rod must not be allowed to scratch the glass,
as the crystals adhere to such scratches and are removed with
difficulty.]
REPRECIPITATION AND IGNITION OF MAGNESIUM AMMONIUM PHOSPHATE
A single precipitation of the magnesium compound in the presence of
molybdenum compounds rarely yields a pure product. The molybdenum can
be removed by solution of the precipitate in acid and precipitation
of the molybdenum by sulphureted hydrogen, after which the magnesium
precipitate may be again thrown down. It is usually more satisfactory
to dissolve the magnesium precipitate and reprecipitate the phosphate
as magnesium ammonium phosphate as described below.
PROCEDURE.--Dissolve the precipitate from the filter in a little
dilute hydrochloric acid (sp. gr. 1.12), allowing the acid solution to
run into the beaker in which the original precipitation was made (Note
1). Wash the filter with water until the wash-water shows no test for
chlorides, but avoid an unnecessary amount of wash-water. Add to
the solution 2 cc. (not more) of magnesia mixture, and then dilute
ammonium hydroxide solution (sp. gr. 0.96), drop by drop, with
constant stirring, until the liquid smells distinctly of ammonia. Stir
for a few moments and then add a volume of strong ammonia (sp. gr.
0.90), equal to one third of the volume of the solution. Allow the
solution to stand for some hours, and then filter off the magnesium
ammonium phosphate, which should be distinctly crystalline in
character. Wash the precipitate with dilute ammonia water, as
prescribed above, until, finally, 3 cc. of the washings, after
acidifying with nitric acid, show no evidence of chlorides. Test both
filtrates for complete precipitation by adding a few cubic centimeters
of magnesia mixture and allowing them to stand for some time.
Transfer the moist precipitate to a weighed porcelain or platinum
crucible and ignite, using great care to raise the temperature slowly
while drying the filter in the crucible, and to insure the ready
access of oxygen during the combustion of the filter paper, thus
guarding against a possible reduction of the phosphate, which would
result in disastrous consequences both to the crucible, if of
platinum, and the analysis. Do not raise the temperature above
moderate redness until the precipitate is white. (Keep this precaution
well in mind.) Ignite finally at the highest temperature of the
Tirrill burner, and repeat the heating until the weight is constant.
If the ignited precipitate is persistently discolored by particles of
unburned carbon, moisten the mass with a drop or two of concentrated
nitric acid and heat cautiously, finally igniting strongly. The
acid will dissolve magnesium pyrophosphate from the surface of the
particles of carbon, which will then burn away. Nitric acid also aids
as an oxidizing agent in supplying oxygen for the combustion of the
carbon.
From the weight of magnesium pyrophosphate (Mg_{2}P_{2}O_{7})
obtained, calculate the phosphoric anhydride (P_{2}O_{5}) in the
sample of apatite.
[Note 1: The ionic change involved in the precipitation of the
magnesium compound is
PO_{4}^{---} + NH_{4}^{+} + Mg^{++} --> [MgNH_{4}PO_{4}].
The magnesium ammonium phosphate is readily dissolved by acids, even
those which are no stronger than acetic acid. This is accounted for
by the fact that two of the ions into which phosphoric acid may
dissociate, the HPO_{4}^{--} or H_{2}PO_{4}^{-} ions, exhibit the
characteristics of very weak acids, in that they show almost no
tendency to dissociate further into H^{+} and PO_{4}^{--} ions.
Consequently the ionic changes which occur when the magnesium ammonium
phosphate is brought into contact with an acid may be typified by the
reaction:
H^{+} + Mg^{++} + NH_{4}^{+} + PO_{4}^{---} --> Mg^{++} + NH_{4}^{+} +
HPO_{4}^{--};
that is, the PO_{4}^{--} ions and the H^{+} ions lose their identity
in the formation of the new ion, HPO_{4}^{--}, and this continues
until the magnesium ammonium phosphate is entirely dissolved.]
[Note 2: During ignition the magnesium ammonium phosphate loses
ammonia and water and is converted into magnesium pyrophosphate:
2MgNH_{4}PO_{4} --> Mg_{2}P_{2}O_{7} + 2NH_{3} + H_{2}O.
The precautions mentioned on pages 111 and 123 must be observed with
great care during the ignition of this precipitate. The danger here
lies in a possible reduction of the phosphate by the carbon of the
filter paper, or by the ammonia evolved, which may act as a reducing
agent. The phosphorus then attacks and injures a platinum crucible,
and the determination is valueless.]
ANALYSIS OF LIMESTONE
Limestones vary widely in composition from a nearly pure marble
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 deg.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 deg.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 deg.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 deg.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 deg.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 deg.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 deg.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 deg.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 deg.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
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