recrystallization from hot water. It must then be dried and cautiously
heated to fusion to expel the last traces of moisture, but not
sufficiently high to expel any oxygen. The pure salt thus prepared,
may be weighed out directly, dissolved, and the solution diluted in a
graduated flask to a definite volume. In this case no standardization
is made, as the normal value can be calculated directly. It is,
however, more generally customary to standardize a solution of
the commercial salt by comparison with some substance of definite
composition, as described below.
PROCEDURE.--Pulverize about 5 grams of potassium bichromate of good
quality. Dissolve the bichromate in distilled water, transfer the
solution to a liter bottle, and dilute to approximately 1000 cc. Shake
thoroughly until the solution is uniform.
To prepare the solution of the reducing agent, pulverize about 28
grams of ferrous sulphate (FeSO_{4}.7H_{2}O) or about 40 grams of
ferrous ammonium sulphate (FeSO_{4}.(NH_{4})_{2}SO_{4}.6H_{2}O) and
dissolve in distilled water containing 5 cc. of concentrated sulphuric
acid. Transfer the solution to a liter bottle, add 5 cc. concentrated
sulphuric acid, make up to about 1000 cc. and shake vigorously to
insure uniformity.
INDICATOR SOLUTION
No indicator is known which, like methyl orange, can be used within
the solution, to show when the oxidation process is complete. Instead,
an outside indicator solution is employed to which drops of the
titrated solution are transferred for testing. The reagent used is
potassium ferricyanide, which produces a blue precipitate (or color)
with ferrous compounds as long as there are unoxidized ferrous ions in
the titrated solution. Drops of the indicator solution are placed upon
a glazed porcelain tile, or upon white cardboard which has been coated
with paraffin to render it waterproof, and drops of the titrated
solution are transferred to the indicator on the end of a stirring
rod. When the oxidation is nearly completed only very small amounts
of the ferrous compounds remain unoxidized and the reaction with the
indicator is no longer instantaneous. It is necessary to allow a brief
time to elapse before determining that no blue color is formed. Thirty
seconds is a sufficient interval, and should be adopted throughout the
analytical procedure. If left too long, the combined effect of light
and dust from the air will cause a reduction of the ferric compounds
already formed and a resultant blue will appear which misleads the
observer with respect to the true end-point.
The indicator solution must be highly diluted, otherwise its own color
interferes with accurate observation. Prepare a fresh solution, as
needed each day, by dissolving a crystal of potassium ferricyanide
about the size of a pin's head in 25 cc. of distilled water. The salt
should be carefully tested with ferric chloride for the presence of
ferrocyanides, which give a blue color with ferric salts.
In case of need, the ferricyanide can be purified by adding to its
solution a little bromine water and recrystallizing the compound.
COMPARISON OF OXIDIZING AND REDUCING SOLUTIONS
PROCEDURE.--Fill one burette with each of the solutions, observing
the general procedure with respect to cleaning and rinsing already
prescribed. The bichromate solution is preferably to be placed in a
glass-stoppered burette.
Run out from a burette into a beaker of about 300 cc. capacity nearly
40 cc. of the ferrous solution, add 15 cc. of dilute hydrochloric acid
(sp. gr. 1.12) and 150 cc. of water and run in the bichromate
solution from another burette. Since both solutions are approximately
tenth-normal, 35 cc. of the bichromate solution may be added without
testing. Test at that point by removing a very small drop of the
iron solution on the end of a stirring rod, mixing it with a drop of
indicator on the tile (Note 1). If a blue precipitate appears at once,
0.5 cc. of the bichromate solution may be added before testing again.
The stirring rod which has touched the indicator should be dipped in
distilled water before returning it to the iron solution. As soon as
the blue appears to be less intense, add the bichromate solution in
small portions, finally a single drop at a time, until the point is
reached at which no blue color appears after the lapse of thirty
seconds from the time of mixing solution and indicator. At the close
of the titration a large drop of the iron solution should be taken for
the test. To determine the end-point beyond any question, as soon as
the thirty seconds have elapsed remove another drop of the solution
of the same size as that last taken and mix it with the indicator,
placing it beside the last previous test. If this last previous test
shows a blue tint in comparison with the fresh mixture, the end-point
has not been reached; if no difference can be noted the reaction is
complete. Should the end-point be overstepped, a little more of the
ferrous solution may be added and the end-point definitely fixed.
From the volumes of the solutions used, after applying corrections for
burette readings, and, if need be, for the temperature of solutions,
calculate the value of the ferrous solution in terms of the oxidizing
solution.
[Note 1: The accuracy of the work may be much impaired by the removal
of unnecessarily large quantities of solution for the tests. At the
beginning of the titration, while much ferrous iron is still present,
the end of the stirring rod need only be moist with the solution; but
at the close of the titration drops of considerable size may properly
be taken for the final tests. The stirring rod should be washed to
prevent transfer of indicator to the main solution. This cautious
removal of solution does not seriously affect the accuracy of the
determination, as it will be noted that the volume of the titrated
solution is about 200 cc. and the portions removed are very
small. Moreover, if the procedure is followed as prescribed, the
concentration of unoxidized iron decreases very rapidly as the
titration is carried out so that when the final tests are made, though
large drops may be taken, the amount of ferrous iron is not sufficient
to produce any appreciable error in results.
If the end-point is determined as prescribed, it can be as accurately
fixed as that of other methods; and if a ferrous solution is at
hand, the titration need consume hardly more time than that of the
permanganate process to be described later on.]
STANDARDIZATION OF POTASSIUM BICHROMATE SOLUTIONS
!Selection of a Standard!
A substance which will serve satisfactorily as a standard for
oxidizing solutions must possess certain specific properties: It must
be of accurately known composition and definite in its behavior as a
reducing agent, and it must be permanent against oxidation in the air,
at least for considerable periods. Such standards may take the form of
pure crystalline salts, such as ferrous ammonium sulphate, or may be
in the form of iron wire or an iron ore of known iron content. It is
not necessary that the standard should be of 100 per cent purity,
provided the content of the active reducing agent is known and no
interfering substances are present.
The two substances most commonly used as standards for a bichromate
solution are ferrous ammonium sulphate and iron wire. A standard wire
is to be purchased in the market which answers the purpose well, and
its iron content may be determined for each lot purchased by a number
of gravimetric determinations. It may best be preserved in jars
containing calcium chloride, but this must not be allowed to come
into contact with the wire. It should, however, even then be examined
carefully for rust before use.
If pure ferrous ammonium sulphate is used as the standard, clear
crystals only should be selected. It is perhaps even better to
determine by gravimetric methods once for all the iron content of a
large commercial sample which has been ground and well mixed. This
salt is permanent over long periods if kept in stoppered containers.
STANDARDIZATION
PROCEDURE.--Weigh out two portions of iron wire of about 0.24-0.26
gram each, examining the wire carefully for rust. It should be handled
and wiped with filter paper (not touched by the fingers), should
be weighed on a watch-glass, and be bent in such a way as not to
interfere with the movement of the balance.
Place 30 cc. of hydrochloric acid (sp. gr. 1.12) in each of two 300
cc. Erlenmeyer flasks, cover them with watch-glasses, and bring the
acid just to boiling. Remove them from the flame and drop in the
portions of wire, taking great care to avoid loss of liquid during
solution. Boil for two or three minutes, keeping the flasks covered
(Note 1), then wash the sides of the flasks and the watch-glass with
a little water and add stannous chloride solution to the hot liquid
!from a dropper! until the solution is colorless, but avoid more than
a drop or two in excess (Note 2). Dilute with 150 cc. of water and
cool !completely!. When cold, add rapidly about 30 cc. of mercuric
chloride solution. Allow the solutions to stand about three minutes
and then titrate without further delay (Note 3), add about 35 cc. of
the standard solution at once and finish the titration as prescribed
above, making use of the ferrous solution if the end-point should be
passed.
From the corrected volumes of the bichromate solution required to
oxidize the iron actually know to be present in the wire, calculate
the relation of the standard solution to the normal.
Repeat the standardization until the results are concordant within at
least two parts in one thousand.
[Note 1: The hydrochloric acid is added to the ferrous solution
to insure the presence of at least sufficient free acid for the
titration, as required by the equation on page 48.
The solution of the wire in hot acid and the short boiling insure the
removal of compounds of hydrogen and carbon which are formed from the
small amount of carbon in the iron. These might be acted upon by the
bichromate if not expelled.]
[Note 2: It is plain that all the iron must be reduced to the ferrous
condition before the titration begins, as some oxidation may have
occurred from the oxygen of the air during solution. It is also
evident that any excess of the agent used to reduce the iron must be
removed; otherwise it will react with the bichromate added later.
The reagents available for the reduction of iron are stannous
chloride, sulphurous acid, sulphureted hydrogen, and zinc; of these
stannous chloride acts most readily, the completion of the reaction
is most easily noted, and the excess of the reagent is most readily
removed. The latter object is accomplished by oxidation to stannic
chloride by means of mercuric chloride added in excess, as the
mercuric salts have no effect upon ferrous iron or the bichromate. The
reactions involved are:
2FeCl_{3} + SnCl_{2} --> 2FeCl_{2} + SnCl_{4}
SnCl_{2} + 2HgCl_{2} --> SnCl_{4} + 2HgCl
The mercurous chloride is precipitated.
It is essential that the solution should be cold and that the stannous
chloride should not be present in great excess, otherwise a secondary
reaction takes place, resulting in the reduction of the mercurous
chloride to metallic mercury:
SnCl_{2} + 2HgCl --> SnCl_{4} + 2Hg.
The occurrence of this secondary reaction is indicated by the
darkening of the precipitate; and, since potassium bichromate oxidizes
this mercury slowly, solutions in which it has been precipitated are
worthless as iron determinations.]
[Note 3: The solution should be allowed to stand about three minutes
after the addition of mercuric chloride to permit the complete
deposition of mercurous chloride. It should then be titrated without
delay to avoid possible reoxidation of the iron by the oxygen of the
air.]
DETERMINATION OF IRON IN LIMONITE
PROCEDURE.--Grind the mineral (Note 1) to a fine powder. Weigh out
accurately two portions of about 0.5 gram (Note 2) into porcelain
crucibles; heat these crucibles to dull redness for ten minutes,
allow them to cool, and place them, with their contents, in beakers
containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12). Heat
at a temperature just below boiling until the undissolved residue is
white or until solvent action has ceased. If the residue is white,
or known to be free from iron, it may be neglected and need not be
removed by filtration. If a dark residue remains, collect it on a
filter, wash free from hydrochloric acid, and ignite the filter in a
platinum crucible (Note 3). Mix the ash with five times its weight of
sodium carbonate and heat to fusion; cool, and disintegrate the fused
mass with boiling water in the crucible. Unite this solution and
precipitate (if any) with the acid solution, taking care to avoid loss
by effervescence. Wash out the crucible, heat the acid solution
to boiling, add stannous chloride solution until it is colorless,
avoiding a large excess (Note 4); cool, and when !cold!, add 40 cc. of
mercuric chloride solution, dilute to 200 cc., and proceed with the
titration as already described.
From the standardization data already obtained, and the known weight
of the sample, calculate the percentage of iron (Fe) in the limonite.
[Note 1: Limonite is selected as a representative of iron ores in
general. It is a native, hydrated oxide of iron. It frequently occurs
in or near peat beds and contains more or less organic matter which,
if brought into solution, would be acted upon by the potassium
bichromate. This organic matter is destroyed by roasting. Since a high
temperature tends to lessen the solubility of ferric oxide, the heat
should not be raised above low redness.]
[Note 2: It is sometimes advantageous to dissolve a large portion--say
5 grams--and to take one tenth of it for titration. The sample will
then represent more closely the average value of the ore.]
[Note 3: A platinum crucible may be used for the roasting of the
limonite and must be used for the fusion of the residue. When used, it
must not be allowed to remain in the acid solution of ferric chloride
for any length of time, since the platinum is attacked and dissolved,
and the platinic chloride is later reduced by the stannous chloride,
and in the reduced condition reacts with the bichromate, thus
introducing an error. It should also be noted that copper and antimony
interfere with the determination of iron by the bichromate process.]
[Note 4: The quantity of stannous chloride required for the reduction
of the iron in the limonite will be much larger than that added to the
solution of iron wire, in which the iron was mainly already in the
ferrous condition. It should, however, be added from a dropper to
avoid an unnecessary excess.]
DETERMINATION OF CHROMIUM IN CHROME IRON ORE
PROCEDURE.--Grind the chrome iron ore (Note 1) in an agate mortar
until no grit is perceptible under the pestle. Weigh out two portions
of 0.5 gram each into iron crucibles which have been scoured inside
until bright (Note 2). Weigh out on a watch-glass (Note 3), using the
rough balances, 5 grams of dry sodium peroxide for each portion, and
pour about three quarters of the peroxide upon the ore. Mix ore and
flux by thorough stirring with a dry glass rod. Then cover the mixture
with the remainder of the peroxide. Place the crucible on a triangle
and raise the temperature !slowly! to the melting point of the flux,
using a low flame, and holding the lamp in the hand (Note 4). Maintain
the fusion for five minutes, and stir constantly with a stout iron
wire, but do not raise the temperature above moderate redness (Notes 5
and 6).
Allow the crucible to cool until it can be comfortably handled (Note
7) and then place it in a 300 cc. beaker, and cover it with distilled
water (Note 8). The beaker must be carefully covered to avoid loss
during the disintegration of the fused mass. When the evolution of
gas ceases, rinse off and remove the crucible; then heat the solution
!while still alkaline! to boiling for fifteen minutes. Allow the
liquid to cool for a few minutes; then acidify with dilute sulphuric
acid (1:5), adding 10 cc. in excess of the amount necessary to
dissolve the ferric hydroxide (Note 9). Dilute to 200 cc., cool, add
from a burette an excess of a standard ferrous solution, and titrate
for the excess with a standard solution of potassium bichromate, using
the outside indicator (Note 10).
From the corrected volumes of the two standard solutions, and their
relations to normal solutions, calculate the percentage of chromium in
the ore.
[Note 1: Chrome iron ore is essentially a ferrous chromite, or
combination of FeO and Cr_{2}O_{3}. It must be reduced to a state of
fine subdivision to ensure a prompt reaction with the flux.]
[Note 2: The scouring of the iron crucible is rendered much easier if
it is first heated to bright redness and plunged into cold water. In
this process oily matter is burned off and adhering scale is caused to
chip off when the hot crucible contracts rapidly in the cold water.]
[Note 3: Sodium peroxide must be kept off of balance pans and should
not be weighed out on paper, as is the usual practice in the rough
weighing of chemicals. If paper to which the peroxide is adhering is
exposed to moist air it is likely to take fire as a result of
the absorption of moisture, and consequent evolution of heat and
liberation of oxygen.]
[Note 4: The lamp should never be allowed to remain under the
crucible, as this will raise the temperature to a point at which the
crucible itself is rapidly attacked by the flux and burned through.]
[Note 5: The sodium peroxide acts as both a flux and an oxidizing
agent. The chromic oxide is dissolved by the flux and oxidized to
chromic anhydride (CrO_{3}) which combines with the alkali to form
sodium chromate. The iron is oxidized to ferric oxide.]
[Note 6: The sodium peroxide cannot be used in porcelain, platinum, or
silver crucibles. It attacks iron and nickel as well; but crucibles
made from these metals may be used if care is exercised to keep the
temperature as low as possible. Preference is here given to iron
crucibles, because the resulting ferric hydroxide is more readily
brought into solution than the nickelic oxide from a nickel crucible.
The peroxide must be dry, and must be protected from any admixture of
dust, paper, or of organic matter of any kind, otherwise explosions
may ensue.]
[Note 7: When an iron crucible is employed it is desirable to allow
the fusion to become nearly cold before it is placed in water,
otherwise scales of magnetic iron oxide may separate from the
crucible, which by slowly dissolving in acid form ferrous sulphate,
which reduces the chromate.]
[Note 8: Upon treatment with water the chromate passes into solution,
the ferric hydroxide remains undissolved, and the excess of peroxide
is decomposed with the evolution of oxygen. The subsequent boiling
insures the complete decomposition of the peroxide. Unless this is
complete, hydrogen peroxide is formed when the solution is acidified,
and this reacts with the bichromate, reducing it and introducing a
serious error.]
[Note 9: The addition of the sulphuric acid converts the sodium
chromate to bichromate, which behaves exactly like potassium
bichromate in acid solution.]
[Note 10: If a standard solution of a ferrous salt is not at hand, a
weight of iron wire somewhat in excess of the amount which would be
required if the chromite were pure FeO.Cr_{2}O_{3} may be weighed out
and dissolved in sulphuric acid; after reduction of all the iron by
stannous chloride and the addition of mercuric chloride, this solution
may be poured into the chromate solution and the excess of iron
determined by titration with standard bichromate solution.]
PERMANGANATE PROCESS FOR THE DETERMINATION OF IRON
Potassium permanganate oxidizes ferrous salts in cold, acid solution
promptly and completely to the ferric condition, while in hot acid
solution it also enters into a definite reaction with oxalic acid, by
which the latter is oxidized to carbon dioxide and water.
The reactions involved are these:
10FeSO_{4} + 2KMnO_{4} + 8H_{2}S_{4} --> 5Fe_{2}(SO_{4})_{3} +
K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O
5C_{2}H_{2}O_{4}(2H_{2}O) + 2KMnO_{4} +3H_{2}SO_{4} --> K_{2}SO_{4} +
2MnSO_{4} + 10CO_{2} + 1 H_{2}O.
These are the fundamental reactions upon which the extensive use of
potassium permanganate depends; but besides iron and oxalic acid the
permanganate enters into reaction with antimony, tin, copper, mercury,
and manganese (the latter only in neutral solution), by which these
metals are changed from a lower to a higher state of oxidation; and it
also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid,
ferrocyanides, and most soluble organic bodies. It should be noted,
however, that very few of these organic compounds react quantitatively
with the permanganate, as is the case with oxalic acid and the
oxalates.
Potassium permanganate is acted upon by hydrochloric acid; the action
is rapid in hot or concentrated solution (particularly in the presence
of iron salts, which appear to act as catalyzers, increasing the
velocity of the reaction), but slow in cold, dilute solutions.
However, the greater solubility of iron compounds in hydrochloric acid
makes it desirable to use this acid as a solvent, and experiments made
with this end in view have shown that in cold, dilute hydrochloric
acid solution, to which considerable quantities of manganous sulphate
and an excess of phosphoric acid have been added, it is possible to
obtain satisfactory results.
It is also possible to replace the hydrochloric acid by evaporating
the solutions with an excess of sulphuric acid until the latter fumes.
This procedure is somewhat more time-consuming, but the end-point of
the permanganate titration is more permanent. Both procedures are
described below.
Potassium permanganate has an intense coloring power, and since the
solution resulting from the oxidation of the iron and the reduction of
the permanganate is colorless, the latter becomes its own indicator.
The slightest excess is indicated with great accuracy by the pink
color of the solution.
PREPARATION OF A STANDARD SOLUTION
!Approximate Strength 0.1 N!
A study of the reactions given above which represent the oxidation of
ferrous compounds by potassium permanganate, shows that there are 2
molecules of KMnO_{4} and 10 molecules of FeSO_{4} on the
left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of
Fe_{2}(SO_{4})_{5} on the right-hand side. Considering only these
compounds, and writing the formulas in such a way as to show the
oxides of the elements in each, the equation becomes:
K_{2}O.Mn_{2}O_{7} + 10(FeO.SO_{3}) --> K_{2}O.SO_{3} + 2(MnO.SO_{3})
+ 5(Fe_{2}O_{3}.3SO_{3}).
From this it appears that two molecules of KMnO_{4} (or 316.0 grams)
have given up five atoms (or 80 grams) of oxygen to oxidize the
ferrous compound. Since 8 grams of oxygen is the basis of normal
oxidizing solutions and 80 grams of oxygen are supplied by 316.0 grams
of KMnO_{4}, the normal solution of the permanganate should contain,
per liter, 316.0/10 grams, or 31.60 grams (Note 1).
The preparation of an approximately tenth-normal solution of the
reagent may be carried out as follows:
PROCEDURE.--Dissolve about 3.25 grams of potassium permanganate
crystals in approximately 1000 cc. of distilled water in a large
beaker, or casserole. Heat slowly and when the crystals have
dissolved, boil the solution for 10-15 minutes. Cover the solution
with a watch-glass; allow it to stand until cool, or preferably over
night. Filter the solution through a layer of asbestos. Transfer the
filtrate to a liter bottle and mix thoroughly (Note 2).
[Note 1: The reactions given on page 61 are those which take place in
the presence of an excess of acid. In neutral solutions the reduction
of the permanganate is less complete, and, under these conditions,
two gram-molecular weights of KMnO_{4} will furnish only 48 grams
of oxygen. A normal solution for use under these conditions should,
therefore, contain 316.0/6 grams, or 52.66 grams.]
[Note 2: Potassium permanganate solutions are not usually stable for
long periods, and change more rapidly when first prepared than after
standing some days. This change is probably caused by interaction
with the organic matter contained in all distilled water, except that
redistilled from an alkaline permanganate solution. The solutions
should be protected from light and heat as far as possible, since both
induce decomposition with a deposition of manganese dioxide, and it
has been shown that decomposition proceeds with considerable rapidity,
with the evolution of oxygen, after the dioxide has begun to form. As
commercial samples of the permanganate are likely to be contaminated
by the dioxide, it is advisable to boil and filter solutions through
asbestos before standardization, as prescribed above. Such solutions
are relatively stable.]
COMPARISON OF PERMANGANATE AND FERROUS SOLUTIONS
PROCEDURE.--Fill a glass-stoppered burette with the permanganate
solution, observing the usual precautions, and fill a second burette
with the ferrous sulphate solution prepared for use with the potassium
bichromate. The permanganate solution cannot be used in burettes with
rubber tips, as a reduction takes place upon contact with the rubber.
The solution has so deep a color that the lower line of the meniscus
cannot be detected; readings must therefore be made from the upper
edge. Run out into a beaker about 40 cc. of the ferrous solution,
dilute to about 100 cc., add 10 cc. of dilute sulphuric acid, and run
in the permanganate solution to a slight permanent pink. Repeat, until
the ratio of the two solutions is satisfactorily established.
STANDARDIZATION OF A POTASSIUM PERMANGANATE SOLUTION
!Selection of a Standard!
Commercial potassium permanganate is rarely sufficiently pure to admit
of its direct weighing as a standard. On this account, and because
of the uncertainties as to the permanence of its solutions, it is
advisable to standardize them against substances of known value. Those
in most common use are iron wire, ferrous ammonium sulphate, sodium
oxalate, oxalic acid, and some other derivatives of oxalic acid.
With the exception of sodium oxalate, these all contain water of
crystallization which may be lost on standing. They should, therefore,
be freshly prepared, and with great care. At present, sodium oxalate
is considered to be one of the most satisfactory standards.
!Method A!
!Iron Standards!
The standardization processes employed when iron or its compounds are
selected as standards differ from those applicable in connection with
oxalate standards. The procedure which immediately follows is that in
use with iron standards.
As in the case of the bichromate process, it is necessary to reduce
the iron completely to the ferrous condition before titration. The
reducing agents available are zinc, sulphurous acid, or sulphureted
hydrogen. Stannous chloride may also be used when the titration is
made in the presence of hydrochloric acid. Since the excess of both
the gaseous reducing agents can only be expelled by boiling, with
consequent uncertainty regarding both the removal of the excess and
the reoxidation of the iron, zinc or stannous chlorides are the most
satisfactory agents. For prompt and complete reduction it is essential
that the iron solution should be brought into ultimate contact with
the zinc. This is brought about by the use of a modified Jones
reductor, as shown in Figure 1. This reductor is a standard apparatus
and is used in other quantitative processes.
[Illustration: Fig. 1]
The tube A has an inside diameter of 18 mm. and is 300 mm. long; the
small tube has an inside diameter of 6 mm. and extends 100 mm. below
the stopcock. At the base of the tube A are placed some pieces of
broken glass or porcelain, covered by a plug of glass wool about 8 mm.
thick, and upon this is placed a thin layer of asbestos, such as is
used for Gooch filters, 1 mm. thick. The tube is then filled with the
amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc
is placed a plug of glass wool. If the top of the tube is not already
shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted
into the tube with a rubber stopper and the reductor is connected
with a suction bottle, F. The bottle D is a safety bottle to
prevent contamination of the solution by water from the pump. After
preparation for use, or when left standing, the tube A should be
filled with water, to prevent clogging of the zinc.
[Note 1: The use of fine zinc in the reductor is not necessary and
tends to clog the tube. Particles which will pass a 10-mesh sieve, but
are retained by one of 20 meshes to the inch, are most satisfactory.
The zinc can be amalgamated by stirring or shaking it in a mixture of
25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric
acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution
should then be poured off and the zinc thoroughly washed. It is then
ready for bottling and preservation under water. A small quantity of
glass wool is placed in the neck of the funnel to hold back foreign
material when the reductor is in use.]
STANDARDIZATION
PROCEDURE.--Weigh out into Erlenmeyer flasks two portions of iron wire
of about 0.25 gram each. Dissolve these in hot dilute sulphuric acid
(5 cc. of concentrated acid and 100 cc. of water), using a covered
flask to avoid loss by spattering. Boil the solution for two or
three minutes after the iron has dissolved to remove any volatile
hydrocarbons. Meanwhile prepare the reductor for use as follows:
Connect the vacuum bottle with the suction pump and pour into the
funnel at the top warm, dilute sulphuric acid, prepared by adding 5
cc. of concentrated sulphuric acid to 100 cc. of distilled water. See
that the stopcock (C) is open far enough to allow the acid to run
through slowly. Continue to pour in acid until 200 cc. have passed
through, then close the stopcock !while a small quantity of liquid
is still left in the funnel!. Discard the filtrate, and again
pass through 100 cc. of the warm, dilute acid. Test this with the
permanganate solution. A single drop should color it permanently; if
it does not, repeat the washing, until assured that the zinc is not
contaminated with appreciable quantities of reducing substances. Be
sure that no air enters the reductor (Note 1).
Pour the iron solution while hot (but not boiling) through the
reductor at a rate not exceeding 50 cc. per minute (Notes 2 and 3).
Wash out the beaker with dilute sulphuric acid, and follow the iron
solution without interruption with 175 cc. of the warm acid and
finally with 75 cc. of distilled water, leaving the funnel partially
filled. Remove the filter bottle and cool the solution quickly under
the water tap (Note 4), avoiding unnecessary exposure to the oxygen of
the air. Add 10 cc. of dilute sulphuric acid and titrate to a faint
pink with the permanganate solution, adding it directly to the
contents of the vacuum flask. Should the end-point be overstepped, the
ferrous sulphate solution may be added.
From the volume of the solution required to oxidize the iron in
the wire, calculate the relation to the normal of the permanganate
solution. The duplicate results should be concordant within two parts
in one thousand.
[Note 1: The funnel of the reductor must never be allowed to empty.
If it is left partially filled with water the reductor is ready for
subsequent use after a very little washing; but a preliminary test is
always necessary to safeguard against error.
If more than a small drop of permanganate solution is required to
color 100 cc. of the dilute acid after the reductor is well washed, an
allowance must be made for the iron in the zinc. !Great care! must be
used to prevent the access of air to the reductor after it has been
washed out ready for use. If air enters, hydrogen peroxide forms,
which reacts with the permanganate, and the results are worthless.]
[Note 2: The iron is reduced to the ferrous condition by contact with
the zinc. The active agent may be considered to be !nascent! hydrogen,
and it must be borne in mind that the visible bubbles are produced by
molecular hydrogen, which is without appreciable effect upon ferric
iron.
The rate at which the iron solution passes through the zinc should not
exceed that prescribed, but the rate may be increased somewhat when
the wash-water is added. It is well to allow the iron solution to run
nearly, but not entirely, out of the funnel before the wash-water
is added. If it is necessary to interrupt the process, the complete
emptying of the funnel can always be avoided by closing the stopcock.
It is also possible to reduce the iron by treatment with zinc in a
flask from which air is excluded. The zinc must be present in excess
of the quantity necessary to reduce the iron and is finally completely
dissolved. This method is, however, less convenient and more tedious
than the use of the reductor.]
[Note 3: The dilute sulphuric acid for washing must be warmed ready
for use before the reduction of the iron begins, and it is of the
first importance that the volume of acid and of wash-water should
be measured, and the volume used should always be the same in the
standardizations and all subsequent analyses.]
[Note 4: The end-point is more permanent in cold than hot solutions,
possibly because of a slight action of the permanganate upon the
manganous sulphate formed during titration. If the solution turns
brown, it is an evidence of insufficient acid, and more should be
immediately added. The results are likely to be less accurate in this
case, however, as a consequence of secondary reactions between the
ferrous iron and the manganese dioxide thrown down. It is wiser to
discard such results and repeat the process.]
[Note 5: The potassium permanganate may, of course, be diluted and
brought to an exactly 0.1 N solution from the data here obtained. The
percentage of iron in the iron wire must be taken into account in all
calculations.]
!Method B!
!Oxalate Standards!
PROCEDURE.--Weigh out two portions of pure sodium oxalate of 0.25-0.3
gram each into beakers of about 600 cc. capacity. Add about 400 cc. of
boiling water and 20 cc. of manganous sulphate solution (Note 1).
When the solution of the oxalate is complete, heat the liquid, if
necessary, until near its boiling point (70-90°C.) and run in the
standard permanganate solution drop by drop from a burette, stirring
constantly until an end-point is reached (Note 2). Make a blank test
with 20 cc. of manganous sulphate solution and a volume of distilled
water equal to that of the titrated solution to determine the volume
of the permanganate solution required to produce a very slight pink.
Deduct this volume from the amount of permanganate solution used in
the titration.
From the data obtained, calculate the relation of the permanganate
solution to the normal. The reaction involved is:
5Na_{2}C_{2}O_{4} + 2KMnO_{4} + 8H_{2}SO_{4} --> 5Na_{2}SO_{4} +
K_{2}SO_{4} + 2MnSO_{4} + 10CO_{2} + 8H_{2}O
[Note 1: The manganous sulphate titrating solution is made by
dissolving 20 grams of MnSO_{4} in 200 cubic centimeters of water and
adding 40 cc. of concentrated sulphuric acid (sp. gr. 1.84) and 40 cc.
or phosphoric acid (85%).]
[Note 2: The reaction between oxalates and permanganates takes place
quantitatively only in hot acid solutions. The temperatures must not
fall below 70°C.]
DETERMINATION OF IRON IN LIMONITE
!Method A!
The procedures, as here prescribed, are applicable to iron ores in
general, provided these ores contain no constituents which are reduced
by zinc or stannous chloride and reoxidized by permanganates. Many
iron ores contain titanium, and this element among others does
interfere with the determination of iron by the process described.
If, however, the solutions of such ores are treated with sulphureted
hydrogen or sulphurous acid, instead of zinc or stannous chloride to
reduce the iron, and the excess reducing agent removed by boiling, an
accurate determination of the iron can be made.
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions
of about 0.5 gram each into small porcelain crucibles. Roast the ore
at dull redness for ten minutes (Note 1), allow the crucibles to cool,
and place them and their contents in casseroles containing 30 cc. of
dilute hydrochloric acid (sp. gr. 1.12).
Proceed with the solution of the ore, and the treatment of the
residue, if necessary, exactly as described for the bichromate process
on page 56. When solution is complete, add 6 cc. of concentrated
sulphuric acid to each casserole, and evaporate on the steam bath
until the solution is nearly colorless (Note 2). Cover the casseroles
and heat over the flame of the burner, holding the casserole in
the hand and rotating it slowly to hasten evaporation and prevent
spattering, until the heavy white fumes of sulphuric anhydride are
freely evolved (Note 3). Cool the casseroles, add 100 cc. of water
(measured), and boil gently until the ferric sulphate is dissolved;
pour the warm solution through the reductor which has been previously
washed; proceed as described under standardization, taking pains
to use the same volume and strength of acid and the same volume of
wash-water as there prescribed, and titrate with the permanganate
solution in the reductor flask, using the ferrous sulphate solution if
the end-point should be overstepped.
From the corrected volume of permanganate solution used, calculate the
percentage of iron (Fe) in the limonite.
[Note 1: The preliminary roasting is usually necessary because, even
though the sulphuric acid would subsequently char the carbonaceous
matter, certain nitrogenous bodies are not thereby rendered insoluble
in the acid, and would be oxidized by the permanganate.]
[Note 2: The temperature of the steam bath is not sufficient to
volatilize sulphuric acid. Solutions may, therefore, be left to
evaporate overnight without danger of evaporation to dryness.]
[Note 3: The hydrochloric acid, both free and combined, is displaced
by the less volatile sulphuric acid at its boiling point. Ferric
sulphate separates at this point, since there is no water to hold
it in solution and care is required to prevent bumping. The ferric
sulphate usually has a silky appearance and is easily distinguished
from the flocculent silica which often remains undissolved.]
!Zimmermann-Reinhardt Procedure!
!Method (B)!
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions
of about 0.5 gram each into small porcelain crucibles. Proceed with
the solution of the ore, treat the residue, if necessary, and reduce
the iron by the addition of stannous chloride, followed by mercuric
chloride, as described for the bichromate process on page 56. Dilute
the solution to about 400 cc. with cold water, add 10 cc. of the
manganous sulphate titrating solution (Note 1, page 68) and titrate
with the standard potassium permanganate solution to a faint pink
(Note 1).
From the standardization data already obtained calculate the
percentage of iron (Fe) in the limonite.
[Note 1: It has already been noted that hydrochloric acid reacts
slowly in cold solutions with potassium permanganate. It is, however,
possible to obtain a satisfactory, although somewhat fugitive
end-point in the presence of manganous sulphate and phosphoric acid.
The explanation of the part played by these reagents is somewhat
obscure as yet. It is possible that an intermediate manganic compound
is formed which reacts rapidly with the ferrous compounds--thus in
effect catalyzing the oxidizing process.
While an excess of hydrochloric acid is necessary for the successful
reduction of the iron by stannous chloride, too large an amount
should be avoided in order to lessen the chance of reduction of the
permanganate by the acid during titration.]
DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE
INDIRECT OXIDATION
Pyrolusite, when pure, consists of manganese dioxide. Its value as an
oxidizing agent, and for the production of chlorine, depends upon the
percentage of MnO_{2} in the sample. This percentage is determined
by an indirect method, in which the manganese dioxide is reduced and
dissolved by an excess of ferrous sulphate or oxalic acid in the
presence of sulphuric acid, and the unused excess determined by
titration with standard permanganate solution.
PROCEDURE.--Grind the mineral in an agate mortar until no grit
whatever can be detected under the pestle (Note 1). Transfer it to a
stoppered weighing-tube, and weigh out two portions of about 0.5 gram
into beakers (400-500 cc.) Read Note 2, and then calculate in each
case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to
react with the weights of pyrolusite taken. The reaction involved is
MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} +
2CO_{2} + 4H_{2}O.
Weigh out about 0.2 gram in excess of this quantity of !pure! oxalic
acid into the corresponding beakers, weighing the acid accurately and
recording the weight in the notebook. Pour into each beaker 25 cc. of
water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the
beaker and its contents gently until the evolution of carbon dioxide
ceases (Note 3). If a residue remains which is sufficiently colored to
obscure the end-reaction of the permanganate, it must be removed by
filtration.
Finally, dilute the solution to 200-300 cc., heat the solution to a
temperature just below boiling, add 15 cc. of a manganese sulphate
solution and while hot, titrate for the excess of the oxalic acid with
standard permanganate solution (Notes 4 and 5).
From the corrected volume of the solution required, calculate the
amount of oxalic acid undecomposed by the pyrolusite; subtract this
from the total quantity of acid used, and calculate the weight of
manganese dioxide which would react with the balance of the acid, and
from this the percentage in the sample.
[Note 1: The success of the analysis is largely dependent upon the
fineness of the powdered mineral. If properly ground, solution should
be complete in fifteen minutes or less.]
[Note 2: A moderate excess of oxalic acid above that required to react
with the pyrolusite is necessary to promote solution; otherwise the
residual quantity of oxalic acid would be so small that the last
particles of the mineral would scarcely dissolve. It is also desirable
that a sufficient excess of the acid should be present to react with a
considerable volume of the permanganate solution during the titration,
thus increasing the accuracy of the process. On the other hand, the
excess of oxalic acid should not be so large as to react with more of
the permanganate solution than is contained in a 50 cc. burette. If
the pyrolusite under examination is known to be of high grade, say 80
per cent pure, or above the calculation of the oxalic acid needed may
be based upon an assumption that the mineral is all MnO_{2}. If the
quality of the mineral is unknown, it is better to weigh out three
portions instead of two and to add to one of these the amount of
oxalic prescribed, assuming complete purity of the mineral. Then run
in the permanganate solution from a pipette or burette to determine
roughly the amount required. If the volume exceeds the contents of a
burette, the amount of oxalic acid added to the other two portions is
reduced accordingly.]
[Note 3: Care should be taken that the sides of the beaker are not
overheated, as oxalic acid would be decomposed by heat alone if
crystallization should occur on the sides of the vessel. Strong
sulphuric acid also decomposes the oxalic acid. The dilute acid
should, therefore, be prepared before it is poured into the beaker.]
[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire
may be substituted for the oxalic acid. The reaction is then the
following:
2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O
The excess of ferrous iron may also be determined by titration with
potassium bichromate, if desired. Care is required to prevent the
oxidation of the iron by the air, if ferrous salts are employed.]
[Note 5: The oxidizing power of pyrolusite may be determined by other
volumetric processes, one of which is outlined in the following
reactions:
MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O
Cl_{2} + 2KI --> I_{2} + 2KCl
I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI.
The chlorine generated by the pyrolusite is passed into a solution of
potassium iodide. The liberated iodine is then determined by titration
with sodium thiosulphate, as described on page 78. This is a direct
process, although it involves three steps.]
IODIMETRY
The titration of iodine against sodium thiosulphate, with starch as an
indicator, may perhaps be regarded as the most accurate of volumetric
processes. The thiosulphate solution may be used in both acid and
neutral solutions to measure free iodine and the latter may, in turn,
serve as a measure of any substance capable of liberating iodine from
potassium iodide under suitable conditions for titration, as, for
example, in the process outlined in Note 5 on page 74.
The fundamental reaction upon which iodometric processes are based is
the following:
I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}.
This reaction between iodine and sodium thiosulphate, resulting in
the formation of the compound Na_{2}S_{4}O_{6}, called sodium
tetrathionate, is quantitatively exact, and differs in that
respect from the action of chlorine or bromine, which oxidize the
thiosulphate, but not quantitatively.
NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE
If the formulas of sodium thiosulphate and sodium tetrathionate are
written in a manner to show the atoms of oxygen associated
with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and
Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are
five atoms of oxygen associated with sulphur, instead of the four
in the two molecules of the thiosulphate taken together. Although,
therefore, the iodine contains no oxygen, the two atoms of iodine
have, in effect, brought about the addition of one oxygen atoms to the
sulphur atoms. That is the same thing as saying that 253.84 grams of
iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8
grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97
grams of iodine should be contained in one liter of normal iodine
solution. By a similar course of reasoning the conclusion is reached
that the normal solution of sodium thiosulphate should contain,
per liter, its molecular weight in grams. As the thiosulphate in
crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight
is 248.12 grams. Tenth-normal or hundredth-normal solutions are
generally used.
PREPARATION OF STANDARD SOLUTIONS
!Approximate Strength, 0.1 N!
PROCEDURE.--Weigh out on the rough balances 13 grams of commercial
iodine. Place it in a mortar with 18 grams of potassium iodide and
triturate with small portions of water until all is dissolved. Dilute
the solution to 1000 cc. and transfer to a liter bottle and mix
thoroughly (Note 1).[1]
[Footnote 1: It will be found more economical to have a considerable
quantity of the solution prepared by a laboratory attendant, and to
have all unused solutions returned to the common stock.]
Weigh out 25 grams of sodium thiosulphate, dissolve it in water which
has been previously boiled and cooled, and dilute to 1000 cc., also
with boiled water. Transfer the solution to a liter bottle and mix
thoroughly (Note 2).
[Note 1: Iodine solutions react with water to form hydriodic acid
under the influence of the sunlight, and even at low room temperatures
the iodine tends to volatilize from solution. They should, therefore,
be protected from light and heat. Iodine solutions are not stable for
long periods under the best of conditions. They cannot be used in
burettes with rubber tips, since they attack the rubber.]
[Note 2: Sodium thiosulphate (Na_{2}S_{2}O_{3}.5H_{2}O) is
rarely wholly pure as sold commercially, but may be purified by
recrystallization. The carbon dioxide absorbed from the air by
distilled water decomposes the salt, with the separation of sulphur.
Boiled water which has been cooled out of contact with the air should
be used in preparing solutions.]
INDICATOR SOLUTION
The starch solution for use as an indicator must be freshly prepared.
A soluble starch is obtainable which serves well, and a solution of
0.5 gram of this starch in 25 cc. of boiling water is sufficient. The
solution should be filtered while hot and is ready for use when cold.
If soluble starch is not at hand, potato starch may be used. Mix about
1 gram with 5 cc. of cold water to a smooth paste, pour 150 cc. of
!boiling! water over it, warm for a moment on the hot plate, and put
it aside to settle. Decant the supernatant liquid through a filter
and use the clear filtrate; 5 cc. of this solution are needed for a
titration.
The solution of potato starch is less stable than the soluble starch.
The solid particles of the starch, if not removed by filtration,
become so colored by the iodine that they are not readily decolorized
by the thiosulphate (Note 1).
[Note 1: The blue color which results when free iodine and starch
are brought together is probably not due to the formation of a true
chemical compound. It is regarded as a "solid solution" of iodine in
starch. Although it is unstable, and easily destroyed by heat, it
serves as an indicator for the presence of free iodine of remarkable
sensitiveness, and makes the iodometric processes the most
satisfactory of any in the field of volumetric analysis.]
COMPARISON OF IODINE AND THIOSULPHATE SOLUTIONS
PROCEDURE.--Place the solutions in burettes (the iodine in a
glass-stoppered burette), observing the usual precautions. Run out 40
cc. of the thiosulphate solution into a beaker, dilute with 150 cc. of
water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate
with the iodine to the appearance of the blue of the iodo-starch.
Repeat until the ratio of the two solutions is established,
remembering all necessary corrections for burettes and for temperature
changes.
STANDARDIZATION OF SOLUTIONS
Commercial iodine is usually not sufficiently pure to permit of its
use as a standard for thiosulphate solutions or the direct preparation
of a standard solution of iodine. It is likely to contain, beside
moisture, some iodine chloride, if chlorine was used to liberate the
iodine when it was prepared. It may be purified by sublimation after
mixing it with a little potassium iodide, which reacts with the iodine
chloride, forming potassium chloride and setting free the iodine. The
sublimed iodine is then dried by placing it in a closed container over
concentrated sulphuric acid. It may then be weighed in a stoppered
weighing-tube and dissolved in a solution of potassium iodide in a
stoppered flask to prevent loss of iodine by volatilization. About 18
grams of the iodide and twelve grams of iodine per liter are required
for an approximately tenth-normal solution.
An iodine solution made from commercial iodine may also be
standardized against arsenious oxide (As_{4}O_{6}). This substance
also usually requires purification by sublimation before use.
The substances usually employed for the standardization of a
thiosulphate solution are potassium bromate and metallic copper. The
former is obtainable in pure condition or may be easily purified by
re-crystallization. Copper wire of high grade is sufficiently pure
to serve as a standard. Both potassium bromate and cupric salts in
solution will liberate iodine from an iodide, which is then titrated
with the thiosulphate solution.
The reactions involved are the following:
(a) KBrO_{3} + 6KI + 3H_{2}SO_{4} --> KBr + 3I_{2} + 3K_{2}SO_{4} + 3H_{2}O,
(b) 3Cu + 8HNO_{3} --> 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O,
2Cu(NO_{3})_{2} + 4KI --> 2CuI + 4KNO_{3} + I_{2}.
Two methods for the direct standardization of the sodium thiosulphate
solution are here described, and one for the direct standardization of
the iodine solution.
!Method A!
PROCEDURE.--Weigh out into 500 cc. beakers two portions of about
0.150-0.175 gram of potassium bromate. Dissolve each of these in 50
cc. of water, and add 10 cc. of a potassium iodide solution containing
3 grams of the salt in that volume (Note 1). Add to the mixture 10 cc.
of dilute sulphuric acid (1 volume of sulphuric acid with 5 volumes of
water), allow the solution to stand for three minutes, and dilute to
150 cc. (Note 2). Run in thiosulphate solution from a burette until
the color of the liberated iodine is nearly destroyed, and then add 1
cc. or 2 cc. of starch solution, titrate to the disappearance of the
iodo-starch blue, and finally add iodine solution until the color
is just restored. Make a blank test for the amount of thiosulphate
solution required to react with the iodine liberated by the iodate
which is generally present in the potassium iodide solution, and
deduct this from the total volume used in the titration.
From the data obtained, calculate the relation of the thiosulphate
solution to a normal solution, and subsequently calculate the similar
value for the iodine solution.
[Note 1:--Potassium iodide usually contains small amounts of potassium
iodate as impurity which, when the iodide is brought into an acid
solution, liberates iodine, just as does the potassium bromate used as
a standard. It is necessary to determine the amount of thiosulphate
which reacts with the iodine thus liberated by making a "blank test"
with the iodide and acid alone. As the iodate is not always uniformly
distributed throughout the iodide, it is better to make up a
sufficient volume of a solution of the iodide for the purposes of the
work in hand, and to make the blank test by using the same volume of
the iodide solution as is added in the standardizing process. The
iodide solution should contain about 3 grams of the salt in 10 cc.]
[Note 2: The color of the iodo-starch is somewhat less satisfactory in
concentrated solutions of the alkali salts, notably the iodides. The
dilution prescribed obviates this difficulty.]
!Method B!
PROCEDURE.--Weigh out two portions of 0.25-0.27 gram of clean copper
wire into 250 cc. Erlenmeyer flasks (Note 1). Add to each 5 cc. of
concentrated nitric acid (sp. gr. 1.42) and 25 cc. of water, cover,
and warm until solution is complete. Add 5 cc. of bromine water and
boil until the excess of bromine is expelled. Cool, and add strong
ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates
the presence of an excess. Boil the solution until the deep blue is
replaced by a light bluish green, or a brown stain appears on the
sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp.
gr. 1.04), cool under the water tap, and add a solution of potassium
iodide (Note 3) containing about 3 grams of the salt, and titrate
with thiosulphate solution until the color of the liberated iodine
is nearly destroyed. Then add 1-2 cc. of freshly prepared starch
solution, and add thiosulphate solution, drop by drop, until the blue
color is discharged.
From the data obtained, including the "blank test" of the iodide,
calculate the relation of the thiosulphate solution to the normal.
[Note 1: While copper wire of commerce is not absolutely pure, the
requirements for its use as a conductor of electricity are such that
the impurities constitute only a few hundredths of one per cent and
are negligible for analytical purposes.]
[Note 2: Ammonia neutralizes the free nitric acid. It should be added
in slight excess only, since the excess must be removed by boiling,
which is tedious. If too much ammonia is present when acetic acid is
added, the resulting ammonium acetate is hydrolyzed, and the ammonium
hydroxide reacts with the iodine set free.]
[Note 3: A considerable excess of potassium iodide is necessary for
the prompt liberation of iodine. While a large excess will do no harm,
the cost of this reagent is so great that waste should be avoided.]
!Method C!
PROCEDURE.--Weigh out into 500 cc. beakers two portions of 0.175-0.200
gram each of pure arsenious oxide. Dissolve each of these in 10 cc. of
sodium hydroxide solution, with stirring. Dilute the solutions to 150
cc. and add dilute hydrochloric acid until the solutions contain a few
drops in excess, and finally add to each a concentrated solution of
5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the
beakers before adding the bicarbonate, to avoid loss. Add the starch
solution and titrate with the iodine to the appearance of the blue of
the iodo-starch, taking care not to pass the end-point by more than a
few drops (Note 1).
From the corrected volume of the iodine solution used to oxidize the
arsenious oxide, calculate its relation to the normal. From the
ratio between the solutions, calculate the similar value for the
thiosulphate solution.
[Note 1: Arsenious oxide dissolves more readily in caustic alkali than
in a bicarbonate solution, but the presence of caustic alkali during
the titration is not admissible. It is therefore destroyed by the
addition of acid, and the solution is then made neutral with the
solution of bicarbonate, part of which reacts with the acid, the
excess remaining in solution.
The reaction during titration is the following:
Na_{3}AsO_{3} + I_{2} + 2NaHCO_{3} --> Na_{3}AsO_{4} + 2NaI + 2CO_{2}
+ H_{2}O
As the reaction between sodium thiosulphate and iodine is not always
free from secondary reactions in the presence of even the weakly
alkaline bicarbonate, it is best to avoid the addition of any
considerable excess of iodine. Should the end-point be passed by a few
drops, the thiosulphate may be used to correct it.]
DETERMINATION OF COPPER IN ORES
Copper ores vary widely in composition from the nearly pure copper
minerals, such as malachite and copper sulphide, to very low grade
materials which contain such impurities as silica, lead, iron, silver,
sulphur, arsenic, and antimony. In nearly all varieties there will be
found a siliceous residue insoluble in acids. The method here given,
which is a modification of that described by A.H. Low (!J. Am. Chem.
Soc.! (1902), 24, 1082), provides for the extraction of the copper
from commonly occurring ores, and for the presence of their common
impurities. For practice analyses it is advisable to select an ore of
a fair degree of purity.
PROCEDURE.-- Weigh out two portions of about 0.5 gram each of the
ore (which should be ground until no grit is detected) into 250 cc.
Erlenmeyer flasks or small beakers. Add 10 cc. of concentrated nitric
acid (sp. gr. 1.42) and heat very gently until the ore is decomposed
and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of
concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then
add about 7 cc. of concentrated sulphuric acid (sp. gr. 1.84) and
evaporate over a free flame until the sulphuric acid fumes freely
(Note 2). It has then displaced nitric and hydrochloric acid from
their compounds.
Cool the flask or beaker, add 25 cc. of water, heat the solution
to boiling, and boil for two minutes. Filter to remove insoluble
sulphates, silica and any silver that may have been precipitated as
silver chloride, and receive the filtrate in a small beaker, washing
the precipitate and filter paper with warm water until the filtrate
and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x
12 cm.) into triangular form and place it on edge in the beaker. Cover
the beaker and boil the solution (being careful to avoid loss of
liquid by spattering) for ten minutes, but do not evaporate to small
volume.
Wash the cover glass and sides of the beaker. The copper should now be
in the form of a precipitate at the bottom of the beaker or adhering
loosely to the aluminium sheet. Remove the sheet, wash it carefully
with hydrogen sulphide water and place it in a small beaker. Decant
the solution through a filter, wash the precipitated copper twice by
decantation with hydrogen sulphide water, and finally transfer the
copper to the filter paper, where it is again washed thoroughly, being
careful at all times to keep the precipitated copper covered with the
wash water. Remove and discard the filtrate and place an Erlenmeyer
flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr.
1.20) over the aluminium foil in the beaker, thus dissolving any
adhering copper. Wash the foil with hot water and remove it. Warm this
nitric acid solution and pour it slowly through the filter paper,
thereby dissolving the copper on the paper, receiving the acid
solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc.
of saturated bromine water (Note 3) through it and finally wash the
paper carefully with hot water and transfer any particles of copper
which may be left on it to the Erlenmeyer flask. Boil to expel the
bromine. Add concentrated ammonia drop by drop until the appearance of
a deep blue coloration indicates an excess. Boil until the deep blue
is displaced by a light bluish green coloration, or until brown stains
form on the sides of the flask. Add 10 cc. of strong acetic acid (Note
4) and cool under the water tap. Add a solution containing about 3
grams of potassium iodide, as in the standardization, and titrate with
thiosulphate solution until the yellow of the liberated iodine is
nearly discharged. Add 1-2 cc. of freshly prepared starch solution and
titrate to the disappearance of the blue color.
From the data obtained, calculate the percentage of copper (Cu) in the
ore.
[Note 1: Nitric acid, because of its oxidizing power, is used as a
solvent for the sulphide ores. As a strong acid it will also dissolve
the copper from carbonate ores. The hydrochloric acid is added to
dissolve oxides of iron and to precipitate silver and lead. The
sulphuric acid displaces the other acids, leaving a solution
containing sulphates only. It also, by its dehydrating action, renders
silica from silicates insoluble.]
[Note 2: Unless proper precautions are taken to insure the correct
concentrations of acid the copper will not precipitate quantitatively
on the aluminium foil; hence care must be taken to follow directions
carefully at this point. Lead and silver have been almost completely
removed as sulphate and chloride respectively, or they too would
be precipitated on the aluminium. Bismuth, though precipitated on
aluminium, has no effect on the analysis. Arsenic and antimony
precipitate on aluminium and would interfere with the titration if
allowed to remain in the lower state of oxidation.]
[Note 3: Bromine is added to oxidize arsenious and antimonious
compounds from the original sample, and to oxidize nitrous acid formed
by the action of nitric acid on copper and copper sulphide.]
[Note 4: This reaction can be carried out in the presence of sulphuric
and hydrochloric acids as well as acetic acid, but in the presence
of these strong acids arsenic and antimonic acids may react with the
hydriodic acid produced with the liberation of free iodine, thereby
reversing the process and introducing an error.]
DETERMINATION OF ANTIMONY IN STIBNITE
Stibnite is native antimony sulphide. Nearly pure samples of this
mineral are easily obtainable and should be used for practice, since
many impurities, notably iron, seriously interfere with the accurate
determination of the antimony by iodometric methods. It is, moreover,
essential that the directions with respect to amounts of reagents
employed and concentration of solutions should be followed closely.
PROCEDURE.--Grind the mineral with great care, and weigh out two
portions of 0.35-0.40 gram into small, dry beakers (100 cc.).
Cover the beakers and pour over the stibnite 5 cc. of concentrated
hydrochloric acid (sp. gr. 1.20) and warm gently on the water bath
(Note 1). When the residue is white, add to each beaker 2 grams of
powdered tartaric acid (Note 2). Warm the solution on the water bath
for ten minutes longer, dilute the solution very cautiously by adding
water in portions of 5 cc., stopping if the solution turns red. It
is possible that no coloration will appear, in which case cautiously
continue the dilution to 125 cc. If a red precipitate or coloration
does appear, warm the solution until it is colorless, and again dilute
cautiously to a total volume of 125 cc. and boil for a minute (Note
3).
If a white precipitate of the oxychloride separates during dilution
(which should not occur if the directions are followed), it is best to
discard the determination and to start anew.
Carefully neutralize most of the acid with ammonium hydroxide solution
(sp. gr. 0.96), but leave it distinctly acid (Note 4). Dissolve 3
grams of sodium bicarbonate in 200 cc. of water in a 500 cc. beaker,
and pour the cold solution of the antimony chloride into this,
avoiding loss by effervescence. Make sure that the solution contains
an excess of the bicarbonate, and then add 1 cc. or 2 cc. of starch
solution and titrate with iodine solution to the appearance of the
blue, avoiding excess (Notes 5 and 6).
From the corrected volume of the iodine solution required to oxidize
the antimony, calculate the percentage of antimony (Sb) in the
stibnite.
[Note 1: Antimony chloride is volatile with steam from its
concentrated solutions; hence these solutions must not be boiled until
they have been diluted.]
[Note 2: Antimony salts, such as the chloride, are readily hydrolyzed,
and compounds such as SbOCl are formed which are often relatively
insoluble; but in the presence of tartaric acid compounds with complex
ions are formed, and these are soluble. An excess of hydrochloric acid
also prevents precipitation of the oxychloride because the H^{+} ions
from the acid lessen the dissociation of the water and thus prevent
any considerable hydrolysis.]
[Note 3: The action of hydrochloric acid upon the sulphide sets free
sulphureted hydrogen, a part of which is held in solution by the acid.
This is usually expelled by the heating upon the water bath; but if it
is not wholly driven out, a point is reached during dilution at which
the antimony sulphide, being no longer held in solution by the acid,
separates. If the dilution is immediately stopped and the solution
warmed, this sulphide is again brought into solution and at the same
time more of the sulphureted hydrogen is expelled. This procedure must
be continued until the sulphureted hydrogen is all removed, since it
reacts with iodine. If no precipitation of the sulphide occurs, it
is an indication that the sulphureted hydrogen was all expelled on
solution of the stibnite.]
[Note 4: Ammonium hydroxide is added to neutralize most of the acid,
thus lessening the amount of sodium bicarbonate to be added. The
ammonia should not neutralize all of the acid.]
[Note 5: The reaction which takes place during titration may be
expressed thus:
Na_{3}SbO_{3} + 2NaHCO_{3} + I_{2} --> Na_{3}SbO_{4} + 2NaI + H_{2}O +
2CO_{2}.]
[Note 6: If the end-point is not permanent, that is, if the blue of
the iodo-starch is discharged after standing a few moments, the cause
may be an insufficient quantity of sodium bicarbonate, leaving the
solution slightly acid, or a very slight precipitation of an antimony
compound which is slowly acted upon by the iodine when the latter is
momentarily present in excess. In either case it is better to discard
the analysis and to repeat the process, using greater care in the
amounts of reagents employed.]
CHLORIMETRY
The processes included under the term !chlorimetry! comprise
those employed to determine chlorine, hypochlorites, bromine, and
hypobromites. The reagent employed is sodium arsenite in the presence
of sodium bicarbonate. The reaction in the case of the hypochlorites
is
NaClO + Na_{3}AsO_{3} --> Na_{3}AsO_{4} + NaCl.
The sodium arsenite may be prepared from pure arsenious oxide,
as described below, and is stable for considerable periods; but
commercial oxide requires resublimation to remove arsenic sulphide,
which may be present in small quantity. To prepare the solution,
dissolve about 5 grams of the powdered oxide, accurately weighed,
in 10 cc. of a concentrated sodium hydroxide solution, dilute the
solution to 300 cc., and make it faintly acid with dilute hydrochloric
acid. Add 30 grams of sodium bicarbonate dissolved in a little water,
and dilute the solution to exactly 1000 cc. in a measuring flask.
Transfer the solution to a dry liter bottle and mix thoroughly.
It is possible to dissolve the arsenious oxide directly in a solution
of sodium bicarbonate, with gentle warming, but solution in sodium
hydroxide takes place much more rapidly, and the excess of the
hydroxide is readily neutralized by hydrochloric acid, with subsequent
addition of the bicarbonate to maintain neutrality during the
titration.
The indicator required for this process is made by dipping strips of
filter paper in a starch solution prepared as described on page 76,
to which 1 gram of potassium iodide has been added. These strips are
allowed to drain and spread upon a watch-glass until dry. When touched
by a drop of the solution the paper turns blue until the hypochlorite
has all been reduced and an excess of the arsenite has been added.
DETERMINATION OF THE AVAILABLE CHLORINE IN BLEACHING POWDER
Bleaching powder consists mainly of a calcium compound which is a
derivative of both hydrochloric and hypochlorous acids. Its formula is
CaClOCl. Its use as a bleaching or disinfecting agent, or as a source
of chlorine, depends upon the amount of hypochlorous acid which it
yields when treated with a stronger acid. It is customary to express
the value of bleaching powder in terms of "available chlorine," by
which is meant the chlorine present as hypochlorite, but not the
chlorine present as chloride.
PROCEDURE.--Weigh out from a stoppered test tube into a porcelain
mortar about 3.5 grams of bleaching powder (Note 1). Triturate the
powder in the mortar with successive portions of water until it is
well ground and wash the contents into a 500 cc. measuring flask
(Note 2). Fill the flask to the mark with water and shake thoroughly.
Measure off 25 cc. of this semi-solution in a measuring flask, or
pipette, observing the precaution that the liquid removed shall
contain approximately its proportion of suspended matter.
Empty the flask or pipette into a beaker and wash it out. Run in the
arsenite solution from a burette until no further reaction takes place
on the starch-iodide paper when touched by a drop of the solution of
bleaching powder. Repeat the titration, using a second 25 cc. portion.
From the volume of solution required to react with the bleaching
powder, calculate the percentage of available chlorine in the latter,
assuming the titration reaction to be that between chlorine and
arsenious oxide:
As_{4}O_{6} + 4Cl_{2} + 4H_{2}O --> 2As_{2}O_{5} + 8HCl
Note that only one twentieth of the original weight of bleaching
powder enters into the reaction.
[Note 1: The powder must be triturated until it is fine, otherwise the
lumps will inclose calcium hypochlorite, which will fail to react with
the arsenious acid. The clear supernatant liquid gives percentages
which are below, and the sediment percentages which are above, the
average. The liquid measured off should, therefore, carry with it its
proper proportion of the sediment, so far as that can be brought about
by shaking the solution just before removal of the aliquot part for
titration.]
[Note 2: Bleaching powder is easily acted upon by the carbonic acid in
the air, which liberates the weak hypochlorous acid. This, of course,
results in a loss of available chlorine. The original material for
analysis should be kept in a closed container and protected form the
air as far as possible. It is difficult to obtain analytical samples
which are accurately representative of a large quantity of the
bleaching powder. The procedure, as outlined, will yield results which
are sufficiently exact for technical purposes.]
III. PRECIPITATION METHODS
DETERMINATION OF SILVER BY THE THIOCYANATE PROCESS
The addition of a solution of potassium or ammonium thiocyanate to one
of silver in nitric acid causes a deposition of silver thiocyanate as
a white, curdy precipitate. If ferric nitrate is also present, the
slightest excess of the thiocyanate over that required to combine with
the silver is indicated by the deep red which is characteristic of the
thiocyanate test for iron.
The reactions involved are:
AgNO_{3} + KSCN --> AgSCN + KNO_{3},
3KSCN + Fe(NO_{3})_{3} --> Fe(SCN)_{3} + 3KNO_{3}.
The ferric thiocyanate differs from the great majority of salts in
that it is but very little dissociated in aqueous solutions, and the
characteristic color appears to be occasioned by the formation of the
un-ionized ferric salt.
The normal solution of potassium thiocyanate should contain an amount
of the salt per liter of solution which would yield sufficient
(CNS)^{-} to combine with one gram of hydrogen to form HCNS, i.e.,
a gram-molecular weight of the salt or 97.17 grams. If the ammonium
thiocyanate is used, the amount is 76.08 grams. To prepare the
solution for this determination, which should be approximately 0.05
N, dissolve about 5 grams of potassium thiocyanate, or 4 grams of
ammonium thiocyanate, in a small amount of water; dilute this solution
to 1000 cc. in a liter bottle and mix as usual.
Prepare 20 cc. of a saturated solution of ferric alum and add 5 cc. of
dilute nitric acid (sp. gr. 1.20). About 5 cc. of this solution should
be used as an indicator.
STANDARDIZATION
PROCEDURE.--Crush a small quantity of silver nitrate crystals in a
mortar (Note 1). Transfer them to a watch-glass and dry them for an
hour at 110°C., protecting them from dust or other organic matter
(Note 2). Weigh out two portions of about 0.5 gram each and dissolve
them in 50 cc. of water. Add 10 cc. of dilute nitric acid which has
been recently boiled to expel the lower oxides of nitrogen, if any,
and then add 5 cc. of the indicator solution. Run in the thiocyanate
solution from a burette, with constant stirring, allowing the
precipitate to settle occasionally to obtain an exact recognition
of the end-point, until a faint red tinge can be detected in the
solution.
From the data obtained, calculate the relation of the thiocyanate
solution to the normal.
[Note 1: The thiocyanate cannot be accurately weighed; its solutions
must, therefore, be standardized against silver nitrate (or pure
silver), either in the form of a standard solution or in small,
weighed portions.]
[Note 2: The crystals of silver nitrate sometimes inclose water which
is expelled on drying. If the nitrate has come into contact with
organic bodies it suffers a reduction and blackens during the heating.
It is plain that a standard solution of silver nitrate (made by
weighing out the crystals) is convenient or necessary if many
titrations of this nature are to be made. In the absence of such a
solution the liability of passing the end-point is lessened by setting
aside a small fraction of the silver solution, to be added near the
close of the titration.]
DETERMINATION OF SILVER IN COIN
PROCEDURE.-- Weigh out two portions of the coin of about 0.5 gram
each. Dissolve them in 15 cc. of dilute nitric acid (sp. gr. 1.2) and
boil until all the nitrous compounds are expelled (Note 1). Cool the
solution, dilute to 50 cc., and add 5 cc. of the indicator solution,
and titrate with the thiocyanate to the appearance of the faint red
coloration (Note 2).
From the corrected volume of the thiocyanate solution required,
calculate the percentage of silver in the coin.
[Note 1: The reaction with silver may be carried out in nitric acid
solutions and in the presence of copper, if the latter does not exceed
70 per cent. Above that percentage it is necessary to add silver in
known quantity to the solution. The liquid must be cold at the time of
titration and entirely free from nitrous compounds, as these sometimes
cause a reddening of the indicator solution. All utensils, distilled
water, the nitric acid and the beakers must be free from chlorides,
as the presence of these will cause precipitation of silver chloride,
thereby introducing an error.]
[Note 2: The solution containing the silver precipitate, as well as
those from the standardization, should be placed in the receptacle for
"silver residues" as a matter of economy.]