C. Remigius Fresenius.

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flask containing 100 c. c. dilute sulphuric acid (1 to 5), add about
1 grm. sodium 'bicarbonate, to produce carbonic acid and expel
the air, and then close the flask with an india-rubber stopper pro-
vided with an evolution tube, as shown in Fig. 84 ; c contains 20
or 30 c. c. water. Heat the flask at first gently, finally to gentle
boiling till the iron is dissolved. The clip b is open, and the
hydrogen escapes through the water in c. Meanwhile boil about
300 c. c. distilled water, to drive out all the air it contains, and
then allow it to cool. As soon as the iron is entirely dissolved,
remove the lamp and close the evolution tube with the clip. When
the iron solution has cooled a little loosen the clip and allow the
water in c to recede, pour the boiled water into c and allow this
also to recede till the solution nearly reaches the mark. Take out
the evolution tube and close the flask with an unperf orated stopper,
allow to cool to the temperature of the room, fill with water to the
mark, shake, and allow to stand, so that the particles of carbon
usually present may deposit. Now take out with a pipette 50 c. c.
of the clear and nearly colorless fluid (containing -J- of the iron

* ZeitscJir. /. analyt. Chem., I, 329.




weighed off), transfer to a 400 c. c. beaker, and dilute till the
beaker is half full. Place the beaker on a sheet of white paper,
or better, on a sheet of glass, with white paper underneath.

Fig. 84.

Fill a GAY-LUSSAC'S or GEISSLER'S burette of 30 c. c. capacity,
divided into 1 c. c. (see Figs. 23 and 24), up to zero with solu-
tion of potassium permanganate, of which take care to have ready
a sufficient quantity perfectly clear and uniformly mixed.

Now add the permanganate to the ferrous solution, stirring the
latter all the while with a glass rod. At first the red drops dis-
appear very rapidly, then more slowly. The fluid, which at first
was nearly colorless, gradually acquires a yellowish tint. From
the instant the red drops begin to disappear more slowly, add the
permanganate with more caution and in single drops, until the last
drop imparts to the fluid a faint but unmistakable reddish color,
which remains on stirring. A little practice will enable you
readily to hit the right point. As soon as the fluid in the burette
has sufficiently collected read off again and mark the number of
c. c. used. The reading off must be performed with the greatest
exactness (see 22) ; the whole error should not be more than
0*1 c. c.


The amount of permanganate solution used should be about 20
c. c. Repeat the experiment with another 50 c. c. of the iron solu-
tion. The difference between the permanganate used in the two
cases should not be more than 1 c. c. ; if it is, make one more
experiment, and when the results are sufficiently near take the mean.
Now calculate what quantity of iron is represented by 100 c. c. of
the permanganate. To this end first divide the iron weighed off
by 5, and then multiply by 0'996, since soft iron wire contains on
the average 0*4 per cent, carbon, &c. ; this gives the quantity of
pure iron contained in 50 c. c. of the solution. Suppose we took
1*05 grin, iron wire and used a mean of 21 -3 c. c. permanganate,
_!_<HL<L = 0-210, 0-210 X 0-996 = 0-20916, and then by rule
of three

21-3: 0-20916 :: 100: a; x = 0-98197;

therefore 100 c. c. permanganate = 0-98197 pure iron.

If there is a deficiency of free acid in the solution of iron, the
fluid acquires a brown color, turns turbid, and deposits a brown
precipitate (manganese dioxide and ferric hydroxide). The same
may happen also if the solution of potassium permanganate is
added too quickly, or if the proper stirring of the iron solution is
omitted or interrupted. Experiments attended with abnormal
manifestations of the kind had always better be rejected. That
the fluid reddened by the last drop of solution of potassium
permanganate added, loses its color again after a time, need create
no surprise or uneasiness; this decolorization is, in fact, quite
inevitable, as a dilute solution of free permanganic acid cannot
keep long undecomposed.

Ib. Titration ~by Ammonium Ferrous Sulphate.

Weigh off, with the greatest accuracy, about 1-4 grrn. of the
pure salt prepared according to the directions given in 65, 4,
dissolve in about 200 c.c. distilled water, previously mixed with
about 20 c.c. dilute sulphuric acid, and proceed as in aa.

By dividing the amount of salt weighed off by 7*0014 (or where
great accuracy is not required by 7) we obtain the quantity of iron

If the salt is not pure, for instance should it contain basic
radicals isomorphous with ferrous iron (manganese, magnesium,


etc.) ; or if it contains ferric iron, or is moist, the result will of
course be too high.

-cc. T'di'<tti,ni Inj Oxalic Acid.

If solution of potassium permanganate is added to a warm
solution of oxalic acid, mixed with sulphuric acid, the liberated
permanganic acid oxidizes the oxalic acid to carbon dioxide and
water [511,0,0. + 2KMnO 4 + 3II,SO 4 - K a SO 4 + 2MnSO 4 +
10CO, + 8H,O]. For the oxidation of 1 mol. oxalic acid (II a C a () 4 )
and 2 at. iron (in the ferrous state) equal quantities of permanganic
acid are accordingly required; therefore, 126*048 parts (1 mol.)
of crystallized oxalic acid correspond, in reference to the oxidiz-
ing action of permanganic acid, to 111-8 parts (2 at.) of iron.

A solution of oxalic acid is altered by the action of light ; it is,
therefore, well only to dissolve as much as will be required for
immediate use. Dissolve 1 to 1'2 grm. pure acid prepared by
65, 1, to 250 c.c. ; 50 c.c. of this solution are introduced into
a beaker, diluted with about 100 c.c. water, from 6 to 8 c.c. cone,
sulphuric acid added, and the fluid heated to about 60. The beaker
is then placed on a sheet of white paper, and permanganate added
from the burette, with stirring. The red drops do not disappear
at first very rapidly, but when once the reaction has fairly set in,
they continue for some time to vanish instantaneously. As soon
as the red drops begin to disappear more slowly, the solution of
potassium permanganate must be added with great caution; if
proper care is taken in this respect, it is easy to complete the
reaction w r ith a single drop of permanganate ; this completion of
the reaction is indicated with beautiful distinctness in the colorless
fluid. To find the iron corresponding to the permanganate used,
multiply the amount of crystallized oxalic acid in the 50 c.c. by 8
and divide by 9.

If the oxalic acid was not perfectly dry, or not quite pure, the
result of the experiment will, of course, lead to fixing the strength
of the solution of potassium permanganate too high. Instead of
pure oxalic acid, SAINT-GILLES has proposed to use crystallized
oxalate of ammonium (NH 4 ) a C,O 4 + H a O). This can easily be pre-
pared in the pure state, keeps well, and can be weighed with
accuracy. 142-16 parts of the crystallized salt correspond to
111 '8 parts iron.

112.] FERROUS IRON. 317

Of tlie foregoing three methods of standardizing solution of
potassium permanganate, the first is the one originally proposed by
MARGUERITE. Ammonium ferrous sulphate was first proposed by
FR. MOHR, and oxalic acid by HEMPEL, as agents suitable for the
pin-pose. With absolutely pure and thoroughly dry reagents, and
proper attention, all three methods give correct results.

For myself, I prefer the first method, as the most direct and
positive, the only doubtful point about it being the question
whether the assumption that the iron w T ire contains 99*6 per cent,
of chemically pure iron is quite correct ; this, however, is of very
trifling importance, as the error could not exceed 0*1 or 0'2 per
cent.* The other two methods are, as may readily be seen, some-
what more convenient, but they are not so trustworthy unless you
can insure the purity and dryness of the preparations.

For the analysis of very dilute solutions of iron, e.g., chalybeate
water, in which the amount of iron may be very approximately
determined with great expedition, by direct oxidization with per-
manganate, a very dilute standard solution must be prepared.
Such a solution may be made by diluting the previous solution
with 9 parts of water or by dissolving O5 grm. crystals of potassium
permanganate in 1 litre of water. It is to be directly standardized
with correspondingly small quantities of iron, ferrous salt, or oxalic

In experiments of this kind, the fact that a certain quantity of
permanganate is required to impart a distinct color to pure acidi-
fied water (which is of no consequence in operations where the
concentrated solution is used) must be taken into consideration ; for
where the solution used is so highly dilute, it takes indeed a measur-
able quantity of it to impart the desired reddish tint to the amount
of water employed. In such cases, the volume of the solution of
iron used for standardizing the permanganate and the volume of
the weak ferruginous solution subjected to analysis should be the
same, and either the two solutions should contain about the same
quantity of iron, or by means of a special experiment, it is ascer-
tained how many 0*1 c. c. of the permanganate are required to
impart the desired pale-red color to the same volume of acidified
water. In the latter case, these 0*1 c. c. will be deducted from
the amount of permanganate used in the regular experiments.

* If you often make iron determinations, you may of course procure a
quantity of wire and determine the amount of the foreign matter in it.


In estimating iron in mineral waters it is of course taken for
granted that the water contains no other substances, such as hy-
drogen sulphide, organic matter, nitrites, etc., that will reducr
the permanganate.

Fig. 85.

ft. Performance of the Analytical Process.
This lias been fully indicated in or. The compound to be ex-
amined is dissolved, preferably with application of a current of
carbon dioxide* (see Fig. 85), in dilute sulphuric acid, allowed
to cool in the current of carbon dioxide, and suitably diluted (if
practicable, the solution of a substance containing about 0*2 grm.
iron should be diluted to about 200 c. c.) ; if free acid is not pres-
ent in sufficient quantity, dilute sulphuric acid is added till about
20 c. c. are present altogether, and then standard permanganate
from the burette, to incipient reddening of the fluid. The vol-
ume of standard solution used is then read off. The strength of
the solution of permanganate being known, the quantity of iron
present in the examined fluid is found by a very simple calculation.
Suppose 100 c. c. of solution of potassium permanganate to corre-
spond to 0'98 grm. iron, and 25 c. c. of the solution to have been
used to effect the oxidation of the ferrous compound examined,

100 : 25:: 0-98 : a?; SB = 0-245.

* If commercial hydrochloric acid is u.M-d for the preparation of CO 2 by
action on marble, it must be free from sulphurous acid, an impurity which it
often contains.

112.] FERROUS IRON. 319

The quantity of ferrous iron originally present amounted
accordingly to 0*24:5 grm.

For the method of determining the total amount of iron
present in a solution containing both ferrous and ferric salts, I
refer to 113; for that of determining the amount in each con-
dition separately, to Section V.

y. Process to le used when titrating hydrocMoric-acid

solutions of Iron with Permanganate.

In titrating hydrochloric-acid solutions of iron with perman-
ganate, it is essential that the standardizing of the reagent and the
actual analysis be performed under similar conditions as regards
dilution, amount of acid, and temperature. Besides the proper
reaction lOFeCl, + 2KMnO 4 + 16HC1 = 5Fe 2 Cl 6 + 2KC1 +
2MnCl 2 + 8H a O, the collateral reaction 2KMnO 4 + 16HC1 =
2KC1 + 2MnCl 2 -f 8II 2 O -f 1001 also takes place, in consequence
of which a little chlorine is liberated. This chlorine does not
combine with the ferrous chloride to form ferric chloride in the
case of considerable dilution, but there occurs a condition of
equilibrium in the fluid containing ferrous chloride, chlorine, and
hydrochloric acid, which is destroyed by addition of a further
quantity of either body (LOWENTIIAL and LENSSEN"*). But since
it is difficult to observe the above conditions of obtaining correct
results, the determination in presence of hydrochloric acid is
always less trustworthy than it is in sulphuric acid solutions.

The following method I have, however, found f to give the
best results :

Standardize the permanganate by means of iron dissolved in
dilute sulphuric acid, make the iron solution to be tested up to \
litre, add 50 c.c. to a large quantity of water acidified with sul-
phuric acid (about 1 litre), titrate with permanganate, then again add
50 c.c. of the iron solution, and titrate again, &c. &c. The num-
bers obtained at the third and fourth time are taken. These are
constant, while the number obtained the first time, and sometimes
also the second time, differs. The result multiplied by 5 gives
exactly the quantity of permanganate proportional to the amount
of ferrous iron present.

J. PENNY'S Method (recommended subsequently by SCHABUS).

If potassium dichromate is added to a solution of a ferrous salt

* Zeitschr.f. analyt. Chem., I, 329. \ 2b., i, 361.


in presence of a strong free acid, the ferrous salt is converted into
ferric salt, whilst a potassium- and a chromic salt of the free acid
is formed (6FeSO 4 + K 3 Cr 3 O 7 + 7II a SO 4 = 3Fe a (SO 4 ), + K a SO 4
+ Cr a (S0 4 ) 3 + TII a O).

Now, with 29 '442 grin, potassium dichromate dissolved to 2
litres of fluid, 33*54: grin, iron may be changed from a ferrous to a
ferric salt (294 '42 being the inol. weight of K a Cr 2 O 7 , and 335*4
being 6 times the at. weight of iron) ; 50 c. c. of the above solution
correspond accordingly to 0-8385 grm. iron.

Care must be taken to use perfectly pure potassium dichro-
mate; the salt is heated in a porcelain crucible until it is just
fused ; it is then allowed to cool under the desiccator, and the re-
quired quantity weighed off when cold. Besides the above solu-
tion, another should also be prepared ten times more dilute and
cortaining hence 1-4721 grm. per litre.

It is always advisable to test the correctness of the standard
solution of potassium dichromate by oxidizing with it a known
quantity of pure iron dissolved to a ferrous salt (see 112, 2, aa).

The ferrous solution is sufficiently diluted, mixed with a suf-
ficient quantity of dilute sulphuric acid, and the standard solution
of potassium dichromate slowly added from the burette, the liquid
being stirred all the while with a thin glass rod. The fluid, which
is at first nearly colorless, speedily acquires a pale green tint, which
changes gradually to a darker chrome-green. A very small drop
of the mixture is now from time to time taken out by means of
the stirring-rod, and brought into contact with a "drop of a solution
of potassium ferricyanide (free from ferrocyanide) on a porcelain
plate, which has been spotted with several of such drops. When
the blue color thereby produced begins to lose the intensity which
it exhibited on the first trials, and to assume a paler tint, the
addition of the solution of potassium dichromate must be more
carefully regulated than at first, and towards the end of the process
a fresh essay must be made, and with larger drops than at first,
after each new addition of two drops, and finally, even of a single
drop; drops must also ho left for some time in contact before the
observation is taken. When no further blue coloration ensues, the
oxidation is terminated. From the remarkable sensitiveness of the
reaction, the exact point may be easily hit to a drop. To heighten
the accuracy of the results, the dilute (ten times weaker) standard
fluid should. jn>t at the end of the process, be substituted for the

113.] FERRIC IRON. 321

concentrated solution of potassium dicliromate ; the iron solution
may besides be diluted to measure 250 c. c., 50 c. c. of this being
used for making the approximate determination, then another
50 c. c. being taken for the determination proper, thus minimizing
the loss incidental to the method.

Thus if exactly 0'84 grm. of a substance has been dissolved,
the number of half c. c. of the standard solution will show the per
cents., while the diluted solution will show the 0~1 per cents, of
iron present.

For the manner of proceeding in presence of ferric salts,
I refer to 113. If there is a deficiency of free acid in the
solution, brown chromic chromate may form, upon which the
solution of ferrous salt exercises no longer a deoxidizing action.

Of the two methods the first affords the advantage that the
end of the operation is at once known by the red color acquired
by the liquid, thus requiring no special testing; the advantage
possessed by the second method is that the standard solution of
potassium chromate is easily prepared and may be readily pre-
served unchanged. Since it has been found that the titration of
iron in hydrochloric-acid solution does not afford entirely satisfac-
tory results with permanganate, recourse has been had again of
late to titration with potassium chromate, which has been iieg-
lected for some time. Attention may also be called to the fact
that where the analyst is free to choose between a solution in
hydrochloric or sulphuric acid, when making a volumetric esti-
mation, preference should always be given to the sulphuric-acid
solution, because it is affected far less by atmospheric oxygei?
than the former (PATTINSON*).


a. Solution.

Many ferric compounds are soluble in water. Ferric oxide
and most ferric compounds which are insoluble in water dissolve
in hydrochloric acid, but many of them only slowly and with
difficulty; compounds of this nature are best dissolved in con-
centrated hydrochloric acid, in a flask, with the aid of heat;

* Zeitschr. f. analyt. Chem., ix, 512,


which, however, should not be allowed to reach the boiling-
point; the compound must, moreover, be finely powdered, and
even then it will often take many hours to effect complete solu-
tion. Sometimes, as in the case of strongly ignited ferric oxide,
the substance is dissolved in potassium disulphate at the fusion
point, or in a mixture of 8 parts sulphuric acid and 3 parts water.
It is frequently advisable, also, to reduce the ferric oxide to the
metallic form by prolonged ignition in hydrogen, and then to
dissolve the metal. Iron-containing silicates which are not decom-
posable by hydrochloric acid, are treated according to 140, b.

b. Determination.

The iron of ferric compounds is usually weighed as ferric
oxide, but sometimes as ferrous sulphide (81). It may, how-
ever, be estimated also indirectly, as well as by volumetric analysis,
both directly and after reduction to ferrous iron. The conver-
sion of compounds of iron into ferric oxide is effected either by
precipitation as ferric hydroxide, preceded in some cases by pre-
cipitation as ferrous sulphide, or as basic ferric acetate, succinate,
or formate, or by ignition. While the volumetric and the now
seldom used indirect methods are applicable in almost all cases,
we may convert into


a. By Precipitation as Ferric Hydroxide.

All salts of inorganic or volatile organic acids and soluble in
water, and likewise those which, insoluble in water, dissolve in
hydrochloric acid, with separation of their acid.

b. By Precipitation as Ferrous Sulphide.
All compounds of iron without exception.

c. By Ignition.

All ferric salts of volatile oxygen acids.


All compounds of iron without exception.

The method 1, <?, is the most expeditious and accurate, and is
therefore preferred in all cases where its application is admis-
sible. The method 1, #, is the most generally used. The
methods 1, 5, and 2, serve principally to effect the separation of
the iron from other bases ; they are resorted to also in certain in-
stances where a is inapplicable, especially in cases where sugar or

113.] FERRIC IRON. 323

other non-volatile organic substances are present; and also to de-
termine iron in ferric phosphates and borates. For the manner
of determining iron in ferric chromate and silicate, I refer to
130 and 140. The volumetric methods for estimating the
iron of ferric compounds are used in technical work almost to
the exclusion of all others, and are very frequently employed in
scientific analyses. The methods of precipitating iron in the form
of basic salts Mall be given in Section V.

1. Determinatio?i as Ferric Oxide.

a. By Precipitation as Ferric Hydroxide.

Mix the solution in a porcelain dish (a glass beaker does not
answer so well) with ammonia in excess, heat nearly to boiling,
decant repeatedly on to a filter, wash the precipitate carefully
with hot w r ater, dry thoroughly (which very greatly reduces the
bulk of the precipitate), and ignite in the manner directed in 53.

For the properties of the precipitate and residue, see 81.
The method is free from sources of error. The precipitate, under
all circumstances, even if there are no fixed bodies to be washed
out, must be most carefully and thoroughly washed, since, should
it retain any traces of ammonium chloride, a portion of the iron
would volatilize in the form of ferric chloride. It is also highly
advisable to dissolve the weighed residue, or a portion of it, in
strong hydrochloric acid, or to fuse it with potassium disulphate
and dissolve the melt in dilute hydrochloric acid to see whether
it is quite free from silicic acid (if any is present it will remain
undissolved). The solution is most readily effected in hydro-
chloric acid if the oxide is previously reduced to metallic iron by
ignition in hydrogen.

I). By Precipitation as Ferrous Sulphide.

The solution, in not too large a flask, is mixed with ammonia
till all the free acid is neutralized. (In the absence of organic,
non-volatile substances, this leads to the precipitation of a little
ferric hydroxide, which, however, is of no consequence.) Add
ammonium chloride, if not already present in sufficient quantity,
then colorless or yellowish ammonium sulphide in moderate ex-
cess, and lastly water, till the fluid reaches to the neck of the flask.
Cork it up and stand in a warm place till the precipitate has
subsided and the supernatant fluid has a clear, yellowish appear-


ance (without a tinge of green). Then wash the precipitate if at all
considerable, first by decantation, then on the filter, using water
containing ammonium sulphide and gradually decreasing quan-
tities of ammonium chloride. "When decanting pour the liquids
into a flask, and not into a filter, and when the washing is com-
plete, then filter the mixed fluids, bring the precipitate on to the
filter, and continue washing uninterruptedly, while the funnel is
kept covered with a glass plate. Neglect of any of these precau-
tions will occasion some loss of substance, the ferrous sulphide
gradually combining with the oxygen of the air and passing thus
into the filtrate as ferrous sulphate. As this sulphate is reprecip-
itated by the ammonium sulphide present, the filtrate assumes,
in such cases, a greenish color, and gradually deposits a black
precipitate, the separation of which is greatly promoted by addi-
tion of ammonium chloride.

When the operation of washing is completed, the moist pre-
cipitate (if it is not dried and determined according to 2) is put,
together with the filter, into a beaker, some water added, and then
hydrochloric acid, until the whole is redissolved. Heat is now
applied, until the solution smells no longer of hydrogen sulphide ;
the fluid is then filtered into a flask, the residual paper carefully
washed, incinerated, the ash treated with warm strong hydrochloric
acid. The solution thus obtained (if yellowish) is added to the
main filtrate, which is next heated with nitric acid (see 112, 1);
the solution (now ferric) is finally precipitated with tinnnoni:i, ;is
in a.

If a solution of potassium ferric, ammonium ferric, or sodium
ferric tartrate contains a considerable excess of alkali carbonate,
the precipitation of the iron as sulphide is prevented to a greater
or less extent (BLUMENAu). In such cases the fluid must therefore
be nearly neutralized with an acid, before the precipitation with

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