Part 3

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

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 ”

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 ”

This marked the time when the oxides first settled, leaving a colorless supernatant liquid.

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 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.

Allen’s thesis.