Henry P. Talbot.

An Introductory Course of Quantitative Chemical Analysis With Explanatory Notes online

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Fit the funnel into the stopper of a filter bottle, and connect the
filter bottle with the suction pump. Suspend some finely divided
asbestos, which has been washed with acid, in 20 to 30 cc. of water
(Note 1); allow this to settle, pour off the very fine particles, and
then pour some of the mixture cautiously into the crucible until an
even felt of asbestos, not over 1/32 inch in thickness, is formed. A
gentle suction must be applied while preparing this felt. Wash the
felt thoroughly by passing through it distilled water until all fine
or loose particles are removed, increasing the suction at the last
until no more water can be drawn out of it; place on top of the felt
the small, perforated porcelain disc and hold it in place by pouring a
very thin layer of asbestos over it, washing the whole carefully;
then place the crucible in a small beaker, and place both in a drying
closet at 100-110°C. for thirty to forty minutes. Cool the crucible
in a desiccator, and weigh. Heat again for twenty to thirty minutes,
cool, and again weigh, repeating this until the weight is constant
within 0.0003 gram. The filter is then ready for use.

Place the crucible in the funnel, and apply a gentle suction, !after
which! the solution to be filtered may be poured in without disturbing
the asbestos felt. When pouring liquid onto a Gooch filter hold the
stirring-rod at first well down in the crucible, so that the liquid
does not fall with any force upon the asbestos, and afterward keep the
crucible will filled with the solution.

Pour the liquid above the silver chloride slowly onto the filter,
leaving the precipitate in the beaker as far as possible. Wash the
precipitate twice by decantation with warm water; then transfer it
to the filter with the aid of a stirring-rod with a rubber tip and a
stream from the wash-bottle.

Examine the first portions of the filtrate which pass through the
filter with great care for asbestos fibers, which are most likely to
be lost at this point. Refilter the liquid if any fibers are visible.
Finally, wash the precipitate thoroughly with warm water until free
from soluble silver salts. To test the washings, disconnect the
suction at the flask and remove the funnel or filter tube from the
suction flask. Hold the end of the tube over the mouth of a small test
tube and add from a wash-bottle 2-3 cc. of water. Allow the water to
drip through into the test tube and add a drop of dilute hydrochloric
acid. No precipitate or cloud should form in the wash-water (Note 16).
Dry the filter and contents at 100-110°C. until the weight is constant
within 0.0003 gram, as described for the preparation of the filter.
Deduct the weight of the dry crucible from the final weight, and from
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°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°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


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Online LibraryHenry P. TalbotAn Introductory Course of Quantitative Chemical Analysis With Explanatory Notes → online text (page 9 of 17)