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Title: The reaction between manganese dioxide and potassium permanganate
Author: Arthur John Hopkins
Release date: July 29, 2024 [eBook #74154]
Language: English
Original publication: Baltimore, MD: Arthur John Hopkins, 1893
Credits: The Online Distributed Proofreading Team at https://www.pgdp.net (This file was produced from images generously made available by The Internet Archive)
*** START OF THE PROJECT GUTENBERG EBOOK THE REACTION BETWEEN MANGANESE DIOXIDE AND POTASSIUM PERMANGANATE ***
The
Reaction between Manganese Dioxide
and
Potassium Permanganate.
Dissertation
presented for the degree of
Doctor of Philosophy
to the Board of University Studies
of the Johns Hopkins University
by
Arthur John Hopkins.
1893.
Acknowledgment----
The work recorded in this paper is the result of a suggestion given by
Professor H. N. Morse. It has received his careful attention throughout
it’s course. For his instruction and exact criticism, I wish to offer
my acknowledgment and thanks. I wish also to express to Professor Ira
Remson my appreciation of his interest and instruction and to offer to
Dr. J. S. Ames my thanks for his instruction in Physics----
Table of Contents.
Page.
I Introduction 1
II Description of Apparatus 3
III Action of Manganese Dioxide on
Potassium Permanganate in acid solution 9
IV Action of a Black Oxide of Manganese on
Potassium Permanganate 24
V The Amount of Nitric Acid Neutralized 31
VI The Stability of Manganese Dioxide 34
VII The Effect of Varying Quantities of
Manganese Oxide and Potassium Permanganate 39
VIII The Reduction of Neutral Potassium Permanganate
by Manganese Oxide at Ordinary Temperature 45
IX Action of Potassium Permanganate upon a
Manganese Oxide Obtained from Manganese
Dioxide by Spontaneous Decomposition 50
Conclusion
_Introduction._
The usual laboratory solution of potassium permanganate must be
frequently restandardized. A slight loss in strength may be detected
after standing even a few days and this change becomes more rapid
as the decomposition proceeds. When one looks for the cause of this
increase in the rate of the decomposition of the permanganate, the
attention is naturally directed to the brown manganese oxide which
separates from the solution. A desire to ascertain whether the presence
of this oxide influences the rate of the decomposition suggested the
experiments here described.
The fact that potassium permanganate may react with certain so-called
peroxides with evolution of oxygen, has long been known and it appeared
possible that a similar reaction may take place between manganese
dioxide and potassium permanganate.
Morse and Allen investigated the reaction between lead dioxide and
potassium permanganate and there is embodied in Allen’s dissertation[1]
a statement of the earlier work in this line with references to the
literature on the subject.
[1] Johns Hopkins University 1892
They have shown that in the presence of a quantity of dilute nitric
acid which is equivalent to the potassium in the potassium permanganate
used, lead dioxide reduces the permanganate to manganese dioxide
without itself suffering reduction, while in the presence of an excess
of stronger nitric acid, e.g. normal to eight normal, the lead dioxide
is also reduced.
The following equations represent the reactions referred to:
I, when the nitric acid is equivalent to the potassium in the potassium
permanganate,
2 KMnO₄ + 3 PbO₂ = K₂O + 3 PbO₂ + 2 MnO₂ + 1½ O₂.
II, when the nitric acid is in excess,
2 KMnO₄ + 3 PbO₂ = K₂O + 3 PbO + 2 MnO₂ + 3 O₂.
The reducing action of manganese dioxide upon potassium permanganate
was suspected from the observed increase in the rapidity of the
decomposition of potassium permanganate solutions. But only a
suggestion of such a reaction could be found in the literature.
Thénard,[2] in 1856 states that manganese dioxide may act upon
potassium permanganate either as a reducing agent, in which act the
manganese dioxide is changed to a manganate, or its influence may be
catalytic, causing the evolution of oxygen. Again Mulder[3] in 1858,
ascribed the decomposition of potassium permanganate solutions to the
presence of some potassium manganate.
[2] Comptes Rendues 42, 382
[3] Jahresbericht 1858 p. 581
It could not be ascertained that any investigation of this matter had
ever been undertaken. It was therefore decided to study the question
and it appeared that the most satisfactory evidence could be obtained
by measuring the quantity of oxygen which is evolved when potassium
permanganate and manganese dioxide are brought together.
Description of Apparatus.
The apparatus[4] employed in this work consists of four parts:
“1) a flask, A, in which the reactions were conducted. A melting point
bulb of about 40 cc. capacity, the diameter of its neck being 20 mm.,
was used for this purpose.
[4] This apparatus was used by Allen and described by him. I wish to
acknowledge my indebtedness to him for the illustration and description
given here which is taken by permission word for word from his
dissertation.
[Illustration]
This was closed with a two-hole rubber stopper, through one hole of
which passed:
2) A small glass tube, BB, running up from the flask about 40 cm., and
then bent twice at right angles. One limb of this tube was surrounded
by a small Liebig’s condenser, while the shorter turned down to meet--
3) a Schiff’s azotometer filled with mercury and connected at its top
with the tube above mentioned. In practice it was found desirable to
have a little water on the top of the mercury column.
4) Through the second hole of the rubber stopper closing the flask was
passed a short piece of glass tubing bent obliquely just above the
stopper and tightly connected with a rubber tube about 50 cm. long,
DD. This tube was clamped near its lower end with a Mohr’s pinch-cock,
while the upper was connected with the stem of a small funnel, F.
When an experiment was to be conducted, the azotometer C, was
connected to the tube, B, and tied tightly, a piece of good rubber
tubing being used for the connection. The reservoir of the azotometer,
M, was then raised until the mercury had driven out all the air in C,
and the stop-cock closed. Next the funnel F was filled with water, the
pinch-cock closing D opened and the water allowed to flow down and
drive out all the air in the tube. The latter was then clamped and
thus kept full of water. The apparatus being now ready, the flask A,
containing the substances in the desired quantity, was made fast to B
and D by means of its stopper. The stop-cock of C opened. The flask
A was heated in a water-bath for any desired time, the oxygen being
collected in C as fast as it was given off.
At the end of the experiment, all the gases which remained in A and B
were driven over into C by lowering the reservoir M and opening the
pinch-cock at D. When this was accomplished the stop-cock of C was
closed and the gases brought under atmospheric pressure by bringing the
mercury in C and M to the same level.
It will be seen that in this gas volume was included the air which A
and B contained at the beginning of the experiment.
The oxygen evolved from the contents of the flask was determined thus:
After reading off the total volume of air and oxygen and reducing
to normal conditions, the total volume of oxygen was obtained by
absorption with phosphorus or pyrogallol, and from the residual
nitrogen could be calculated from the volume of air to be deducted
from the total volume of gases. The remainder is the volume of oxygen
sought.”
Action of manganese dioxide on potassium permanganate in acid solution.
Three sets of apparatus like the one described were used.
The flask of apparatus No. I contained manganese dioxide and dilute
nitric acid.
The flask of apparatus No. II contained potassium permanganate and
dilute nitric acid.
The flask of apparatus No. III contained potassium permanganate and
dilute nitric acid and manganese dioxide.
The manganese dioxide used in these experiments was precipitated
from a solution of potassium permanganate by a dilute solution of
manganous sulfate in the manner described on page 34 and then washed
and dried at 100°C.
An illustration of the proportions in which the different substances
were brought together in each of the three flasks is to be found in the
following statement of the quantities used in the first experiment.
The variations from these proportions occurring in subsequent
experiments are noted in the table giving the summary of the results.
Usually about 150 m.g. of manganese dioxide were used and the solutions
in the three flasks were brought to the same volume by addition of
distilled water.
Experiment No. I
Flask No. I Manganese dioxide 150 m.g.
N/10 nitric acid 14.28 c.c.
[= Mn. in KMnO₄ in No. II or No. III]
Water, 27.12 c.c.
[Total solution = 41.40 c.c.]
Flask No. II
Potassium permanganate, 20 c.c. = 112.48 m.g. KMnO₄.
N/10 nitric acid 7.14 c.c. [= K in KMnO₄]
Water, 14.26 c.c.
[Total solution = 41.40 c.c.]
Flask No. III
Potassium permanganate, 20 c.c. = 112.48 m.g. KMnO₄.
N/10 nitric acid 21.40 c.c. [= K + Mn in KMnO₄]
Manganese dioxide 150 m.g.
The nitric acid in No. I is calculated to be the same as that remaining
free after the potassium of the permanganate in No. III has been
neutralized. It will be clear that the quantities of free nitric acid
are the same in flask No. I and in flask No. III provided the potassium
of the permanganate is appropriated by the nitric acid. Moreover each
flask contains the same quantity of manganese dioxide and the volume of
the liquid is the same in both.
On the other hand flask No. III contains permanganic acid. This
arrangement was selected in order to determine how far the evolution
of oxygen was influenced by the action of nitric acid on the manganese
dioxide or by the spontaneous decomposition of permanganic acid. It
was thought that if more oxygen should be obtained from flask No. III
containing a mixture of permanganate, manganese dioxide and nitric
acid, than from mixtures No. I and No. II, the larger volume must be
due to the reduction of the permanganate by the manganese oxide.
The three flasks after having been filled as described were attached to
the apparatus and then all submerged to the same depth in the boiling
water of a single glass water-bath. The stop-cocks of the azotometers
were then opened and the mercury allowed to fall to within about fifty
millimetres of its level in the reservoirs. This difference in level
was maintained as nearly as practicable by lowering the reservoirs
commenserably with the increase in gas volume. No action could be
observed in flasks No. I and No. II during the whole course of the
experiment. The contents of flask No. III however gave off bubbles of
oxygen immediately, the mercury fell in the azotometer and the dark
purple of the permanganate solution in flask No. III became rapidly
lighter. Within about five minutes it was reduced to a delicate pink
but this tint persisted for about ten minutes when it also disappeared.
Immediately after the disappearance of the permanganate color the
solution was filled with a brown oxide which remained for about
thirty five minutes when it subsided leaving the supernatant liquid
clear. Notwithstanding the fact that the color of the permanganate
had entirely disappeared and the suspended oxide had subsided within
an hour after beginning the experiment, the flasks in the earlier
experiments were allowed to remain for two hours longer. By admitting
distilled water through the tube =D= [see figure] the gas was then
forced over into the azotometer. Finally, the oxygen was determined in
the manner previously described.
The following illustration from the first experiment may serve to show
the method of calculating the results.
I II III
Total air and oxygen = 42.60 c.c. 52.40 c.c. 60.60 c.c.
Nitrogen after absorption
of oxygen = 34.00 c.c. 40.90 c.c. 37.80 c.c.
Oxygen = 8.60 c.c. 11.50 c.c. 22.80 c.c.
Air = 79.8% nitrogen
20.2% oxygen
Oxygen of air 20.2
-------------- × ---- = 8.60 c.c. 10.353 c.c. 9.57 c.c.
Nitrogen found 79.8
Oxygen obtained = 0.0 1.147 c.c. 13.23 c.c.
Temperature = 21°C.
Barometer = 737 mm.
Oxygen obtained at Normal = 0 1.033 c.c. 11.913 c.c.
The 20 c.c. KMnO₄ = 112.48 m.g.
1 atom of oxygen = 7.967 c.c.
Atoms oxygen found I II III
from the molecule KMnO₄ = 0 0.129 1.496
As at the end of the preceding calculation the results are expressed
in the table following in terms of oxygen atoms obtained from each
molecule of potassium permanganate [KMnO₄] reduced. Thus an equivalent
of one atom of oxygen from one molecule of potassium permanganate
would equal 15.96/157.67 of the weight of potassium permanganate used
in the experiment. This weight divided by the weight of one cubic
centimetre of oxygen gives the number of cubic centimetres of oxygen at
normal, equivalent to one atom of oxygen from one molecule of potassium
permanganate. The results obtained are tabulated as follows:
Three-hour determinations.
-------------------+-------------------------+----------------------
Flask I | Flask II | Flask III
-------------------+-------------------------+----------------------
MnO₂ N/10 HNO₃| KMnO₄ N/10 HNO₃ | KMnO₄ N/10 HNO₃
[150 m.g.] No. of | 20 c.c. = No. of | 20 c.c. = No. of
molecules | 112.48 m.g. molecules | 112.48 m.g. molecules
-------------------+-------------------------+----------------------
1 0.00 2 0.13 1 1.495 3
2 0.00 ” ? 1 1.493 ”
3 ” 0.13 3 1.532* ”
4 0.00 ” 0.21 ” 1.466 ”
5 0.03 ” 0.10 ” 1.467 ”
6 0.02 ” 0.25 ” ”
7 0.06 ” 0.08 ” 1.526* ”
8 0.04 ” 0.15 ” 1.527 ”
9 0.04 ” 0.11 ” 1.504 ”
10 ” 0.20 ” 1.504 ”
11 0.03 ” 0.14 ” 1.516 ”
12 0.00 ” 0.18 ” 1.524 ”
13 0.03 ” 0.25 ” 1.559* ”
14 0.05 ” 0.26 ” 1.574 ”
15 0.02 1 0.32 2 1.552 2
16 0.03 ” 0.31 ” 1.583? ”
17 0.04 ” 0.31 ” 1.512 ”
18 0.01 ” 0.32 ” 1.466 ”
19 0.00 ” 0.18 ” 1.412 ”
20 0.01 ” 0.36 ” 1.411 ”
21 0.06 ” 0.22 ” 1.402 ”
22 0.00 ” 0.30 ” 1.495* ”
23 2 0.08 3 1.469 3
[5]Fifty minute determinations
24 0.04 1 0.06 2 1.564 2
25 0.04 ” 0.11 ” 1.491 ”
26 0.05 ” ? ” 1.512 ”
27 0.03 ” 0.05 ” 1.532 ”
[6]Fifteen minute determinations
28 0.05 1 0.08 2 1.340 2
29 0.04 ” 0.05 ” 1.325 ”
30 0.06 ” 0.05 ” 1.335 ”
31 0.00 ” 0.00 ” 1.347 ”
32 0.03 ” 0.03 ” 1.412 ”
33 0.02 2 0.02 3 1.363 3
34 0.02 ” 0.06 ” 1.377 ”
35 0.00 ” 0.02 ” 1.363 ”
[5] This marked the time when the oxides first settled, leaving a
colorless supernatant liquid.
[6] This marked the time when the solution was first decolorized.
Inspection of the table will show that whether the oxygen were
determined immediately after the subsidence of the oxides (i.e.
after fifty minutes) or after three hours the results obtained are
practically the same. The tendency of the potassium permanganate to
lose strength made the frequent preparation of new samples advisable.
Wherever a new sample has been used in the course of the experiments,
the fact has been noted in the table by an asterisk. It will be noticed
that where a new solution is used the figures for flask no. III are
immediately larger and nearer to one and one half atoms of oxygen
from each molecule of potassium permanganate.
A phosphorus gas-pipette was used for the absorption of the oxygen in
all the experiments the results of which are embodied in the foregoing
table. In subsequent experiments an alkaline solution of pyrogallol was
employed. It is now known that the variation in the composition of the
manganese oxide in use had some influence upon the results.
It has also been found that the manganese dioxide prepared by the
reduction of permanganate by manganese sulphate is much less stable
than was supposed at the time this work was begun. The dioxide
prepared in this way begins immediately to lose oxygen spontaneously
but recovers the same in the presence of an excess of potassium
permanganate. In the light of these facts it is easy to understand
why lower results were obtained when the oxygen was determined
immediately after the disappearance of the color of the permanganate
and before the suspended oxide had subsided. It appears that the
manganese oxide employed in these experiments was not, as was supposed
at the time, the dioxide but one containing a smaller proportion of
oxygen. If such is the case the first action of the permanganate upon
it would be to replace the oxygen which had been lost. The reduction of
the remaining permanganate would then probably be in accordance with
the equation,
2 KMnO₄ + 3 MnO₂ = 2 K₂O + 5 MnO₂ + 1½ O₂
At the time when the permanganate color disappears, all of the
manganese is in the dioxide condition and the further evolution of
oxygen, which is shown by the preceding experiments to take place
during the subsidence of the suspended oxides, is due to a partial
reduction of this manganese. Therefore the relation of the reduction
of the manganese oxide below the MnO₂ condition before the treatment
with permanganate to the reduction which follows the disappearance
of the permanganate color will determine whether the oxygen evolved
shall be more or less than one and one half atoms to each molecule of
permanganate.
Neither variations in the quantities of nitric acid used (from two to
three molecules in No. III) nor the very slight variations in the amount
of manganese dioxide used, seem to affect appreciably the amount of
oxygen obtained.
It appears that the action of manganese dioxide on potassium
permanganate is the same as that of lead superoxide[7] in the presence
of very dilute nitric acid. Both reduce it to manganese dioxide with
the evolution of one and one half atoms of oxygen to each molecule of
the permanganate.
[7] Allen’s thesis.
The evolution of oxygen from flask No. I containing manganese dioxide
and nitric acid is very slight. From flask No. II containing potassium
permanganate and one equivalent nitric acid, it is also slight but
usually greater than from flask No. I. The differences are much greater
in the case of those determinations in which the heating of the flask
was continued for three hours and in this fact is to be found further
evidence of the reducing action of manganese dioxide on potassium
permanganate.
The possibility of a reaction analogous to that between potassium
permanganate and lead dioxide in the presence of strong nitric acid
seems to be excluded by the fact that the higher oxides of manganese
may be prepared in the presence of concentrated nitric acid.
Action of a Black Oxide of Manganese on Potassium Permanganate.----
As an analysis of the manganese dioxide showed the probable presence
of some manganese in the manganous condition, a sample was treated on
the water-bath for some days with dilute nitric acid while distilled
water was added from time to time to keep the dilution nearly constant.
This was done for the purpose of determining whether the manganous
manganese might not be extracted in the form of manganese nitrate. The
light brown oxide which was used in the previous experiments had been
prepared by adding a dilute solution of manganous sulphate to a hot
solution containing potassium permanganate in excess and dilute nitric
acid. No analysis of the oxide thus prepared was made at the time of
its preparation, but it has been shown by the work of others that the
oxide must have had at that time a composition as regards the relation
of manganese to oxygen, very nearly equivalent to that of MnO₂. During
the course of the experiments previously recorded, a determination of
the ratio of the manganese to the available oxygen in the compound was
made and found to be as
1.03 : 1.00
The substance gave 54.69 percent manganese and 15.44 percent available
oxygen.
A second analysis made some weeks later gave a ratio of manganese to
available oxygen of 1.06: 1.00.--
The substance gave 54.76 percent manganese
15.02 ” available oxygen.
The manganese was determined by the method of Gibbs[8] and the
available oxygen by oxalic acid and potassium permanganate.
[8] Silliman’s Am. J. Science (2) 44.216
It was this oxide which was treated with nitric acid in the manner
described. In the course of this treatment the light brown color
changed to black and the oxide became heavier. The dilute nitric acid
in which the oxide had been digested was found to contain manganous
manganese, but analysis of the black product showed that the oxide had
suffered a still further reduction. The ratio of manganese to available
oxygen was found to have risen from 1.06: 1.00 to 1.157: 1.000. The
analytical data are given below:--
The substance gave { 60.49 percent manganese }
{ 60.46 ” ” }
and
{ 15.23 percent oxygen }
{ 15.19 ” ” }
This ratio would correspond very nearly to a compound having the
composition MnO·6MnO₂ but it is by no means certain that the reducing
action of nitric acid was finished. The action of this oxide on
potassium permanganate was determined under the same conditions as in
the case of the brown oxide as described on page 14. The following
illustration of the proportions used is taken from the first experiment.
Experiment No. I
Flask No. I Manganese oxide, black, = 150 m.g.
N/10 nitric acid = 14.93 c.c.
[= Mn in KMnO₄ in No. II or No. III]
Water = 27.46 c.c.
[Total solution = 42.39 c.c.]
Flask No. II Potassium permanganate = 20 c.c. = 117.71 m.g.
N/10 nitric acid = 22.39 c.c.
[= K and Mn in KMnO₄]
[Total solution = 42.39 c.c.]
Flask No. III Potassium permanganate = 20 c.c. = 117.71 m.g.
N/10 nitric acid = 22.39 c.c.
[= K and Mn in KMnO₄]
Manganese oxide, black = 150 m.g.
[Total solution = 42.39 c.c.]
The first two results are from experiments in which the action was
stopped at the first loss of the color of potassium permanganate.
The last two are the results from experiments in which the action
was stopped after the oxide had settled leaving a clear colorless
supernatant liquid. The amount of oxygen given off in any one case is
expressed in the table as the number of atoms of oxygen derived from
each molecule of potassium permanganate, as follows:
A = N/10 nitric acid. No. of molecules.
------------------+-------------------------+--------------------
Flask I | Flask II | Flask III
------------------+-------------------------+--------------------
Black | Potassium | Potassium
oxide A |permanganate A | permanganate A
150 m.g. | 20 c.c. = | 20 c.c.
| 117.71 m.g. | = 117.71 m.g.
----------------------+-------------------------+----------------
1 2 hours.
0.00 2 0.18 3 1.337 3
2 2¼ hours
0.02 ” 0.24 ” 1.335 ”
3 2½ hours
0.02 ” 0.44 ” 1.402 ”
4 4½ hours
0.00 ” 0.55 ” 1.362 ”
The volume of the oxygen evolved is considerably smaller than that
obtained when the brown oxide was used. It is nevertheless larger
than could have been evolved if all of the manganese had been left
in the dioxide condition at the close of the experiment. This fact
indicates that there was a slight reduction of the manganese dioxide
after the disappearance of the color of the permanganate. That such a
reduction of manganese dioxide does take place will be made apparent
by a comparison of the results of the first two experiments in which
the oxygen was measured soon after the color of permanganate had
disappeared, with the results of the last two experiments in which the
oxygen was determined after a longer interval. It will be noticed in
this as in the series of experiments with the brown oxide that the rate
of the evolution of oxygen liberated in flask No. II increases with the
time of heating.
The Amount of Nitric Acid neutralized.
A few experiments were made to determine how much nitric acid is
neutralized under the conditions which prevailed in the foregoing
experiments.
I. The flask contained brown oxide 150 m.g. N/10 nitric acid 6.98 c.c.
and was heated for three hours in the water-bath.
Acid neutralized = 4.51 c.c. = 64.60 percent.
II. The flask contained brown oxide 150 m.g. N/10 nitric acid 6.98 c.c.
and was heated for 50 minutes in the water-bath.
Acid neutralized = 3.09 c.c. = 44.20 percent.
III. The flask contained brown oxide 150 m.g. N/10 nitric acid 7.46
c.c. and was heated for fifty minutes in the water-bath.
Acid neutralized = 3.16 c.c. = 42.30 percent.
IV. The flask contained brown oxide 150 m.g. N/10 nitric acid 7.46 c.c.
and was heated for 15 minutes in the water-bath.
Acid neutralized 2.76 c.c. = 37.0 percent
V. The flask contained brown oxide 150 m.g. potassium permanganate 20
c.c. = 112.48 m.g. N/10 nitric acid 6.98 c.c. (= 2 molecules for each
molecule of KMnO₄) and was heated for three hours in the water-bath.
Acid neutralized = 0.35 c.c. = 5.0 percent.
VI. Repetition of No. V.
Acid neutralized = 0.44 c.c. = 6.3 percent.
VII. The flask contained brown oxide 150 m.g. potassium permanganate
the same as in numbers V and VI and was heated for 15 minutes in the
water-bath.
Acid neutralized = 5.2 percent.
VIII. Repetition of No. VII
Acid neutralized = 4.0 percent
IX. The flask contained black oxide 150 m.g. N/10 nitric acid 14.93
c.c. and was heated for two and one fourth hours in the water-bath.
Acid neutralized = 4.87 c.c. = 32.46 percent.
X. Repetition of No. IX.
Acid neutralized = 38.10 percent.
XI. The flask contained black oxide 150 m.g. potassium permanganate 20
c.c. = 177.71 m.g., N/10 nitric acid 22.39 c.c. (= 2 molecules for each
molecule of KMnO₄) and was heated four and one half hours in the water
bath.
Acid neutralized = 100.00 percent.
No definite conclusions can be drawn from these neutralization
experiments since it is well known that manganese oxides prepared
as were those used in these experiments always contain some, though
probably not a constant quantity of alkali. It is stated by Bemmelen[9]
that manganese dioxide decomposed the salts of the alkalies with
liberation of acid.
[9] J. für Prak. Ch. (2) 23, 324-379.
The Stability of Manganese Dioxide.
Having found in the course of the work already described that
manganese dioxide prepared in the wet way is much less stable than was
supposed, it was decided to make some experiments upon the spontaneous
decomposition which it undergoes. To this end a fresh sample was
prepared in the following manner.
10 grams of potassium permanganate were dissolved in 500 c.c. of
distilled water and the solution allowed to settle for one day. The
liquid was them filtered through glass wool, heated to 65°C. and
treated with 320 c.c. N/10 nitric acid, having also a temperature
of 65°C. 20.5 grams of manganous sulphate dissolved in 2.5 litres of
water heated to 65°C, were now added, with stirring, to the acidified
solution of potassium permanganate. The precipitate was allowed to
settle and the supernatant liquid which still retained the permanganate
color, decanted. The residue was then treated with water, stirred and
allowed to settle. The water was decanted and the oxide filtered. The
filter used consisted of two platinum cones between which was placed a
small paper filter which did not quite reach to the edge of the outer
cone. This filter was placed in the bottom of a glass funnel and the
oxide poured upon it. The oxide was repeatedly washed with distilled
water and then transferred to a porous plate. It was afterwards
heated for several hours to a temperature of 65°C. At this temperature
it was found impracticable to bring the oxide to a constant weight,
but this fact was no serious obstacle in the way of the subsequent
work since it was only desired to ascertain what changes take place in
the ratio of the manganese to the available oxygen. The manganese was
determined by the method of Gibbs and the available oxygen by oxalic
acid and potassium permanganate. An analysis of the oxide made shortly
after its preparation in the manner described, gave
{51.98 percent}
Manganese { } = 51.86 pr. ct.
{51.74 ” }
{14.836 percent}
Oxygen { } = 14.837 ”
{14.839 ” }
showing a ratio of manganese to available oxygen of 1.018 : 1.000.
The above analysis was made May 1, 1892. The oxide was then placed in
a glass-stoppered weighing bottle and allowed to stand until November.
An analysis made on the tenth day of the latter month gave a ratio of
manganese to available oxygen of
1.198 : 1.000
corresponding very closely to a composition MnO·5MnO₂, in which the
calculated ratio is
1.200 : 1.000
Six month later (April 20, 1893) an analysis of the same material
showed a ratio of manganese to available oxygen of
1.300 : 1.000
This result was confirmed Mʳ Walker who made repeated analyses of the
substance at the same time.
Another sample of manganese oxide was prepared in the same manner
as that used in the previous experiments and placed under water. An
analysis made after
Eight days gave a ratio of Mn : available O = 1.013 : 1.000
Twenty ” ” ” ” ” ” ” ” = 1.037 : 1.000
Forty three ” ” ” ” ” ” ” ” = 1.040 : 1.000
The results here recorded clearly indicate that the brown oxide of
manganese which is prepared by the reduction of potassium permanganate
by manganese sulphate undergoes a spontaneous decomposition involving
a loss of oxygen. This loss of oxygen or the effect of varying
conditions upon this decomposition of manganese dioxide is now being
investigated by Mʳ Walker of this laboratory.
The Effect of Varying Quantities of Manganese Oxide on Potassium
Permanganate.
The following experiments were made for the purpose of determining the
effect of varying quantities of manganese oxide on the time required
for the reduction of potassium permanganate.
A. In the presence of nitric acid. The ratio of manganese to available
oxygen in the material used for these experiments was determined and
found to be
1.018 : 1.000
The experiments were conducted after the same manner as those
previously described.
Different portions of the manganese oxide were weighed out and treated
in the flasks used in the previous experiments with the quantities
of permanganate which were found by calculation to have the desired
molecular ratios to the oxide.
A quantity of N/10 nitric acid equivalent to the potassium of the
potassium permanganate was added. The flasks were then immersed in the
bath of boiling water and the time required for the disappearance of
the permanganate color noted.
The rations of the oxide to the permanganate and the times required for
the reduction of the latter are embodied in the following table.
Molecular Ratio }
Oxide to KMnO₄ } 1:1 2:1 3:1 4:1 5:1 6:1 7:1
Time of decolorization
in minutes
1 ” 120 68
2 ” 110 65
3 ” 150 65
4 ” 150 80
5 ” 120 75
6 ” 120 75
7 ” 175(?) 78 53
8 ” 160 79 54
9 ” 145 41 30
10 ” 120 50 36
11 ” ? 47 32
12 ” 110 37 30
13 ” 64 [10]{20(?)
{20(?)
14 ” 72 5(?) 5(?)
15 ” 85 5(?) 5(?)
[10] Decolorized before time was taken.
The results appearing in the table on the same horizontal lines are
those of experiments made simultaneously, that is, the flasks in these
experiments were immersed at the same time and in the same water-bath.
These, that is the results appearing on the same horizontal line,
and these only are comparable, since it was found that very slight
differences in the temperature of the bath were of great influence on
the length of time required for the reduction of the permanganate.
It is clear that the rate of the reduction of the permanganate is
greatly increased by increasing the quantity of the manganese oxide.
After reduction of the permanganate, the contents of the flasks were
in every case filtered and the filtrate tested by Crum’s reaction
for the presence of manganese. No manganese was found. The absence of
manganese in the filtrate shows that the reduction of the permanganate
was not due to the formation of manganese nitrate, though it will be
clear from the results of subsequent experiments that the presence of
nitric acid exerts an influence on the rate of the reduction.
B. In neutral solution.
The condition were the same as under A, except that the nitric acid was
omitted. Those flasks containing potassium permanganate and one, two
and seven molecular equivalents respectively of the manganese oxide
were immersed in the water-bath as before.
No. I {containing one molecule of oxide and
{one molecule KMnO₄ was reduced in 57 hours
No. II {containing two molecules of oxide and
{one molecule KMnO₄ was reduced in
35 hours
No. III {containing seven molecules of oxide
{and one molecule KMnO₄ was reduced in
2¾ hours
C. In alkaline solution.
The conditions were the same as under B, except that five molecules of
potassium hydroxide were added for each molecule of the permanganate.
No. I {containing one molecules of oxide for
{each molecule KMnO₄ was reduced in
91¾ hours.
No. II {containing two molecules of oxide
{for each molecule KMnO₄ was reduced in
55(?) hours.
No. III {containing seven molecules of oxide for
{each molecule KMnO₄ was reduced in
10 hours.
It appears that potassium permanganate is less easily reduced by
manganese oxide in neutral than in acid solutions and in alkaline that
in neutral solutions.
The Reduction of Neutral Potassium Permanganate by Manganese Oxide at
Ordinary Temperature.
In these experiments, the manganese oxide used was prepared in the same
manner as that in which the proportion of manganese to available oxygen
was found immediately after its preparation to be
1.018 : 1.000
It was dried for several days at 65°C., but for almost a month after
its preparation was not used. Hence, what the ratio of the manganese
to the oxygen at the time when the permanganate was treated with it,
is quite uncertain. It however appears probable from the subsequent
experiments on the instability of manganese dioxide at 65°C. that the
ratio of manganese to available oxygen must have been at about 1.05:
1.00 corresponding to a formula MnO·20MnO₂.
Two solutions of permanganate were made, filtered through glass wool
and standardized.
Solution A.
A cubic centimetre of this was found to be equivalent to 5.278 m.g.
iron in the ferrous condition.
Solution B.
A cubic centimetre of this was found to be equivalent to 14.379 m.g.
iron in the ferrous condition.
Three 250 c.c. portions of each of the solutions A and B were placed in
clean glass-stoppered bottles. These bottles were labelled “a”, “a′”
and “a″” “b”, “b′” and “b″”.
a contained 250 c.c. potassium permanganate each cubic centimetre of
which was equivalent to 5.278 m.g. Fe″. It was closed and put away in
the dark and allowed to stand from June fourteenth until the fourteenth
of the following October. Its strength was then determined and found to
have declined 0.72 percent.
“a′” contained the same quantity of the same permanganate solution as
“a” and it was closed on the same date and placed in mildly diffused
light until October fourteenth. An analysis on the latter date showed
that its strength had also declined 0.72 percent.
“a″” contained the same quantity of the same permanganate solution as
“a” and “a′” and to this was added 0.5 gram of the manganese oxide
previously referred to. The bottle was then closed and placed in
diffused light beside “a′”. An analysis, October fourteenth showed that
its strength had declined 34.62 percent.
“b” contained potassium permanganate each cubic centimetre of which was
equivalent to 14.379 m.g. Fe″. It was closed and put away in the dark
beside “a”. An analysis October fourteenth showed that its strength had
declined 1.71 percent.
“b′” contained the same quantity of the same permanganate as “b”. It
was closed and placed in mildly diffused light beside “a′”. An analysis
October fourteenth showed a decline in strength of 2.41 percent.
“b″” contained the same quantity of the same permanganate as “b” and
“b′”. To it was added 0.5 grams of the manganese oxide. The bottle
was closed and placed in diffused light beside “a′”, “a″”, and “b′”.
An analysis, October fourteenth showed a decline in strength of 78.86
percent.
The conclusions to be drawn from the foregoing results are (1) that
dilute solutions of potassium permanganate are more stable both in
diffused light and in darkness than the more concentrated ones, and (2)
that the presence of manganese oxides hasten to an enormous degree the
reduction of potassium permanganate even at summer temperatures. It
will also be observed that this reducing action of the manganese oxide
is greater in the stronger than in the more dilute solutions. It should
be remarked that the summer of 1892 in which these experiments were
made, was one of unusual heat.
Action of Potassium Permanganate upon a Manganese Oxide Derived from
Manganese Dioxide by Spontaneous Decomposition.
The oxide used in these experiments was one which had been prepared
in the manner heretofore described by adding manganese sulphate to an
excess of potassium permanganate. Soon after its preparation the ratio
of manganese to available oxygen in it was found to be
1.018 : 1.000
At the time when it was used in these experiments, the ratio of
manganese to available oxygen was found to be
1.198 : 1.000
Corresponding very nearly to the ratio of manganese to available oxygen
in MnO·5MnO₂.
Experiment I.
230 m.g. of the oxide was treated with the quantity of permanganate
calculated to be necessary for its oxidation to manganese dioxide.
The mixture was kept at a temperature of 22°C. until the color of the
permanganate disappeared. This required about ten hours.
An analysis of the oxide then gave a ratio of manganese to available
oxygen
1.027 : 1.000
Experiment II
A repetition of experiment I. An analysis of the oxide after
decolorization of the permanganate gave a ratio of manganese to
available oxygen of
1.022 : 1.000
Experiment III
The conditions are the same as in the preceding experiments except that
twice the amount of permanganate calculated for the conversion of the
oxide into manganese dioxide was added.
An analysis of the residue after decolorization gave a ratio of
manganese to available oxygen of
0.994 : 1.000
Experiment IV
A repetition of experiment III. An analysis of the residue gave a
ratio of manganese to available oxygen of
0.995 : 1.000
Experiment V
A repetition of experiments III and IV. The ratio of manganese to
available oxygen in the residue was
0.995 : 1.000.
The foregoing results indicate (experiments III, IV, and V) that the
oxygen lost spontaneously by manganese dioxide prepared in the wet way,
is fully recovered when the reduced oxide is treated with excess of
potassium permanganate.
The ratio of the manganese to the available oxygen was determined in
the following manner. The contents of the flask were filtered through
asbestos and the flask and the precipitate washed with cold water.
The contents of the filter were then returned to the flask and
treated with an excess of a standard solution of oxalic acid and dilute
sulphuric acid. After reduction the excess of the oxalic acid was
determined by a standard solution of potassium permanganate.
The solution was then filtered, well washed and the manganese in the
filtrate precipitated by the method of Gibbs. From the total amount of
manganese found, there was deducted the quantity which was introduced
in determining the excess of oxalic acid by potassium permanganate.
Summary of Conclusions--
I
Potassium permanganate in weakly acid, in neutral, and in alkaline
solutions, is reduced by manganese dioxide which has been prepared in
the wet way, with evolution of one and one half atoms of oxygen for
each molecule of the permanganate. In other words, the permanganate is
reduced to manganese dioxide.
II
This reduction is most rapid in acid and slowest in alkaline
solutions.
III
The rate of reduction is greatly increased by increasing the proportion
of manganese dioxide.
IV
In concentrated solutions of potassium permanganate, the reduction by
manganese dioxide is relatively more rapid than in dilute.
V
Manganese oxide, prepared in the wet way i.e. by the addition
of manganese sulphate to a hot solution containing an excess of
permanganate, has the composition MnO₂ as regards the ratio of
manganese to available oxygen.
VI
This oxide is unstable. It loses oxygen spontaneously even at ordinary
temperatures.
VII
The oxygen which is thus lost is recovered when the oxide is again
treated with an excess of potassium permanganate.
VIII
It is suggested that a partial explanation of the reduction of
potassium permanganate by manganese oxide may be found in the
instability of the latter; in other words that the reduction may
consist in alternate reduction and reoxidation processes. But when the
rapidity of the reduction is taken into account, it seems probable that
a reaction analogous to that between permanganic acid and hydrogen
superoxide must also take place.
Arthur John Hopkins was from September 20, 1864 in Bridgewater,
Massachusetts. Receiving his preparation for college in the public
schools of that town, he entered Amherst College in 1881 and graduated
with the degree of A. B. in 1885. The five years following were spent
in teaching in the States of Massachusetts and New York. He entered the
Johns Hopkins University in 1890, was assistant in the quantitative
laboratory of that university during the academic year of 1891-92, and
the following year was appointed Fellow in Chemistry.
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