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Full text of "The oxygen pressures over mixtures of manganese dioxide and other metallic oxides with reference to catalysis .."

TH:^ OXYGEM PR33SUHSS OYER MIXTURES 0? 

MMGAKESE DIOXIDE AUD OTHiiH I/IETALLIG OXIDES 

WITH REJ^HEKCE TO CATALYSIS. 



DISSERTATION. 



Submitted to the Board of University Studies 
of The Johns Hopkins University in Conformity with the 
Requirements for the Degree of Doctor of Philosophy. 



^V 



■illiam S. English. 



Hay, 1921. 



i; 



0,'-/' 



Table of Contents. 

Acknowledgment iii 

Introduction 1 

Materials 2 

Apparatus 4 

Procedure 12 

Experimental data 13 

Discussion of results 16 

Summary and conclusion 26 

Biography 27 



ii 



Acknowledgment . 

This investigation was undertaken at the suggestion of 
Dr. J. C. W. Frazer, and the author desires to take this op- 
portunity to express sincere thanks for his constant interest 
and inspiration. 

The author also wishes to express his gratitude to Pro- 
fessors Patrick, Reid, Lovelace and Swartz and Doctors Thorn- 
ton and Milligan for valuable laboratory and class room in- 
structions received. 



iii. 



1. 



It has been frequently observed in catalysis that a mixture of 
several materials gives a catalyst which is better than that of any 
of the separate components for the particular reaction. Because of 
the remarkably low temperatures at which different catalytic oxida- 
tions could be carried out with mixtures of manganese dioxide and 
various other metallic oxides; i.e. a mixture of manganese dioxide 
and copper oxide will oxidize carbon monoxide to carbon dioxide at 
a temperature as low as -150C. ; it was thought that a study of the 
oxygen pressures over these materials might throw some light upon 
the phenomenon of their catalytic activity. According to Lewis^ 
manganese dioxide, platinum and metallic silver, all catalise the de- 
composition of silver oxide, in accordance with this view it was 
suggested in the case of the silver oxide and manganese dioxide 
mixture that the silver oxide acted as an intermediate oxygen carrier 
as follows : 

EAggO = 4Ag -I-' Og 

This was all the more plausible since it was assumed that the decom- 
position pressures over manganese dioxide were negligable as compared 
to those over silver oxide. According to Lewis the decomposition 
temperature of silver oxide, taking the atmospheric oxygen pressure 
as equal to sixteen centimeters of mercury, is 155°C; while according 
to Askenasy^ that of manganese dioxide is about 500°C. The results 
obtained, however, do not confirm the above suggestions. 

■LTeparation of Materials Used ; The manganese dioxide used 

tt - 

1, Lewis. Jour. Am. Chem. iioc. 28 , 13, (1906). / 

2. Askenasy and ELonomski. Zeit. iilect. Chem. 16, 107. 



was prepared according to the method of Fremy^. 100 grams of powder- 
ed potassium permanganate are slowly stirred into a cooled mixture of 
500 grams of sulphuric acid and 150 grams of water. The exact course 
of this decomposition is not understood, it has frequently been com- 
pared to the decomposition of the oxygen halogen compounds under simi- 
lar conditions, but Reynolds^ and others have pointed out that manga- 
nese belongs in the VIII group with iron, nickel, cobalt etc. and not 
in the VII group with the halogens, it has similar physical properties 
to members of the eighth group, specific gravity, melting point, etc. 
Chemically its salts are similarly constituted 
Me SO4 7H2O, M.e2S04.ke SO4.6H2O. 

It forms similar metallo and metalli-cynides and oxides and its oxygen 
salts K2Mn04 and EMn04 should bo compared with K2RUO4 and JLRUU4; and 
K2OSQ4 and KOSO4. 

According to Reynolds the eighth group is as follows: 

Mn Fe NiCo 

Eu Eh ir-d 

Os Ir i-t 

a 
Dry JiMn04 is fairly stable compound at ordinary temperatures, 

however on heating it gives off oxygen and red vapors of the higher 

manganese oxides. Reducing agents such as SO2. the lower alcohols, 

glycerine etc. reduce it to manganses dioxide. KMn04 is stable in 

1. Ji'remy. C. E. 82, 1213, il876j 

2. iieynolds. Chem. News. 96, 260 (1907} 



an alkaline solution. In a neutral or acid solution it decomposes 
spontaneously with the evolution of oxygen gas and with the precipi- 
tation of manganese dioxide. The rate of decomposition is greater 
the more concentrated the acid and with concentrated acids the de- 
composition may become explosive. It is for this reason that the 
150 cc of water are included in the above formula for the prepa- 
ration of Fremy oxide. During this work, hov/ever, even with this 
dilution of acid a decomposition of about two pounds of ifldnu^, after 
standing all afternoon in sunshine, exploded during the night with 
extreme violence, uf^times during these decompositions the volatile 
manganese oxides precipitated a fluffy manganese dioxide around the 
decomposition flask. The reaction is completed at the end of several 
days as is shown by the absence of the permanganate color when a 
pipetted sample from the reaction mixture-is poured into water. The 
manganese dioxide is then washed free from sulphates, to entirely 
free the manganese dioxide from sulphates is a long tedious process. 
A 400 gram sample of manganese dioxide was washed by deoantation with 
10 liters of water, night and morning for a month before a solution 

chloric- 

of it in hydrotdcC acid failed to give a precipitation with JiaCl2. This 
extreme washing is not necessary in the preparation of the actual 
catalysts. The material is stored wet. m preparing the mixtures 
studied a weight of the wet material was taken which would give the 
desired weight of manganese dioxide and to this was added a solution 
of a salt of the desired oxide, preferably the nitrate, and with 
vigorous stirring the whole was precipitated with either a UaOH or 
NagCOg solution. This material was washed by deoantation and filtered 



on asbestos since after drying in a des^ieator it is impossible to 
detach all of the filtering medium, if filter paper had been used it 
would have caused some reduction of the oxides on heating with the 
liberation of carbon dioxide and water. 

I'here are in general two methods, the static and the dynamic, for 
measuring the equilibrium concentrations of a substance between two 
phases, bince the static method was thought the more simple, it was 
used first, A dried mixture of 10 grams of MnOg and 6 grams of AggO 
was placed in a bulb attached to a closed manometer. The bulb was 
heated to 200°C. and the apparatus sv;ept out with dry oxygen for several 
hours, it was then slowly cooled in a stream of oxygen, quickly evacuated 
at room temperature to a pressure equal to 0.1 mm. of mercury and the 
following readings taken: 

Temperature i centigrade; 100 160 180 200 220 240 260 
JrTessures ascending 9.4 14.5 20.7 28.5 39.0 64.1 88.6 
Cm. of mercury descending 78.6 

At the higher temperatures in this run drops of water appeared in 
the tube in spite of a -^^2^5 ^^1^ which was between the reaction bulb 
and the manometer, in the presence of other gases water vapor is not 
readily taken up by ^o^^ especially if the gas space to be dried is 
separated from the ±"2^5 ^^ ^ length of tubing. During these experi- 
ments a P2O5 ^^It* was left open to the air for several months through 
about 6 inches of 4 mm glass tubing without becoming appreciably hydrated. 

In order to prevent the adsorption of gases from the flowing oxy- 
gen stream, which would register on the manometer at higher temperatures, 
the compact static apparatus sketched in i'ig. I was designed. 




Fiq. I 



The oxide, whose decomposition was to be measured was heated in bulb A 
in the electric furnace described by Jaorse and Frazer^, where the tem- 
perature is controlled by a careful regulation of the current, which 
flows through the furnace. The pressure was measured on the closed 
manometer B. E is a ±'£'^5 ^^1^» ^"^ operation stopcocks C and D are 
opened and the apparatus well evacuated by a Gaede pump attached to Z, 
as was shown by a McGleod gauge in the system. Stopcock D was then 
closed and oxygen admitted to the apparatus from the storage flask 
G through the two way stopcock F. This procedure of evacuating and 
admitting oxygen from G was repeated several times. The oxygen was 
prepared by heating KMn04 contained in bulb H. Any COg which might 
result from an oxidation of organic matter in the KMnu^ was removed 

by the soda lime bulb J, which also absorbed the volatile manganese 

g 

oxides formed by heating the KMnO^. This method according to Travers 



1. iiorse and i'razer. Am. Chem. Jour. 3£, 

2. Travers' book "Theory of Gases". 



93 a904i 



gives an oxygen of a high degree of purity. The evacuation for the 
following readings was carried out at 1500C. afterwhich the apparatus 
was allowed to cool in oxygen and quickly evacuated at room temperature 
to a pressure equal to 0.1 mm of mercury. The material studied was a 
mixture of 10 grams of manganese dioxide and 6 grams of silver oxide, 
which was prepared as described above. 

Temperatures iTessures Cm. of Hg 

(centigrade; ascending descending 

120 1.4 

140 6.8 IS.O 

160 9.4 15. S 

180 12.0 18.4 

200 15.2 £2.2 

220 19.5 28.1 

240 26.2 33. 6 

260 35.2 42.7 

280 50.2 

At 280°C. a thin film of water appeared on the inside of the tube 
just outside of the furnace. From this it was assumed that the appara- 
tus must be evacuated at tne highest temperature at which the decompo- 
sition pressure was afterwards to be measured in order to be sure that 
none of the pressure was due to water vapor. These results are con- 
tradictory to previous resuxts obtained in this laboratory where it was 
claimed that Fremy oxide could be completely dehydrated in a vacuum at 
160°C. Wright and Menlce-*- found that manganese dioxide retains 0.1 ^o of 

1, V/right and Menke. Chem. Jour. 57 ,22 (1880) 



water even after strong ignition, .which small amount could be driven 
off on reheating the manganese dioxide with fused ii2Cr207. Small 
traces of substances in any material are always difficult to remove 
whether it he through difference in solubility in a precipitation or 
in an eztraction by a solvent; or whether it be through a difference 
in the boiling point in a distillation etc. , Lewis rightly objects 
to soma previous measurements made on the dissociation pressures of 
silver oxide claiming that part of the pressure measured was due to 
CO2 and that the last traces of 002 could not be removed at 300°C. 
i'rom an off handed study of the dissociation pressures of silver car- 
bonate this would hardly seem possible. The following measurements on 
the dissociation pressures of silver carbonate are due to uolson : 

Temperatures 13£ 167 182 210 218 
pressures Cm. of Hg ,6 9.9 17.5 54.7 75.2 

At 300° C. silver carbonate has a dissociation pressure equal to several 
atmospheres. 

The following results are from the second run made on the same 
sample of manganese dioxide and silver oxide used in the first run. 
The evacuation was carried out at 400°C. and the other procedures of 
cooling in oxygen etc. were the same as in the first run. 



1. Colson, C. H. 132 467 U901J 



Temperatures 


Pressures 


( centigrade i 


ascending 


170 


4,3 


210 


5.1 


240 


7.6 


315 


16.0 


370 


27.3 


425 


40.5 


500 





Cm. of Hg 
descending 

4.9 

17.4 

15.3 
27.5 



In the above run at 500°C. a small film of water appeared in the tube 
just outside of the furnace. 

The furnace was allowed to cool down to room temperature, momen- 
tarily evacuated to a pressure equal to 0.1 mm of mercury and the fol- 
lowing readings made: 

Cm. of Hg 
descending 

1.1 

2.3 

3.4 

4.9 

12.2 

22.2 

31.5 



Temperatures 


irTessures 


I centigrade; 


ascending 


100 


.1 


170 


.6 


210 


1.5 


240 


3.3 


315 


11.3 


370 


20.2 


400 


31.9 


450 


46.1 



9. 



An attempt was next made to measure the dissociation pressures 
of a mixture of 5 grams of CaO and 6 grams of Ag20, This mixture 
was prepared by precipitating a mixture of their nitrates with a 
solution of carbonate free sodium hydroxide. The procedures, of 
evacuating at the highest temperature, of cooling in oxygen, of 
evacuating at room temperature etc. as outlined in the previous runs 
were followed but on reheating, the material gave no oxygen pressure. 
The color of the material had changed from the dark gray of Ag£0 to 
a light gray which showed that the silver oxide had lost its oxygen 
during the evacuation and could not regain it while cooling in oxygen. 

iiecause it is impossible to entirely remove COg, HgO etc. by the 
above method without removing oxygen, and the latter reaction might be 
irreversible under the conditions of the experiment, the following 
dynamic apparatus was devised. The essential feature of this apparatus 
consists of passing an inert gas back and forth over the oxide whose 
decomposition pressures are to be measured and determining the oxygen 
pressures by an analysis of the gas mixture. A sketch of this appara- 
tus is shown in figure 2. Beam A is rocked by the lever and driving 
wheel B, which was driven by a small direct current motor, through a 
series of reducing gears, at about one revolution per minute. In the 
figure ^assume the v/heel to be moving counter clockwise, then bulb C 
will be raised and C lowered. These bulbs are partly filled with 
mercury and are connected by rubber pressure tubing to bulbs D and D' . 
Eow raising bulb C raises the mercury level in D. This forces gas out 
of D, through E which contains soda lime, through F which is filled 



^^^^- — w 




10, 



with glass beads wet with HgiiO^ , through the heating coils and bulb 
G which contains the material whose decomposition pressure^ asie to be 
measured, and back through a similar series of purifiers in a reversed 
order as above named to the gas space above the mercury level in D'. 
which is increasing because raising bulb C lowers bulb C by the same 
amount, which in turn causes the mercury level in iJ' to be lowered, 
un the reverse cycle gas will be forced from D' back to D. 

To prevent the added pressure due to the decomposition of the oxides 
from forcing mercury out of C and C ^and also to prevent any vibrations 
of the manometer H due to the tt&aymfrtry of the gas spaces in C and D with 
those of C and D'^or from bulbs C and C being at different distances 
from the center of beam A, 6 and C are connected by the rubber pressure 
tubing I, which is filled with water. 

At the beginning of a run the relative mercury levels are fixed by 
running water into I from bulb J through the stopcock K. Bulb G was 
heated in an electric furnace made by wrapping no. 21 nichrome wire 
around a spirally corrugated aludum tube and the whole was insulated with 
11 inches of asbestos covering. The temperature was regulated by the 
expansion and contraction of a tightly drawn nichrome wire through the 
center of the furnace which acting through an electric relay made and 
broke the current flowing through the furnace. The first run on the 
mixture of 10 grams of manganese dioxide and 6 grams of silver oxide 
was made with this furnace. 

Jj'or the remainder of the runs a heavy copper pipe was put inside 
the aludum core. to improve the heat distribution of the furnace and 
the expansion and contraction of this pipe was used to make and break 
the heating current of the furnace, it is possible with this type of 



11. 



electric furnace to bold temperatures up to 600°C. constant within E 
degrees for days.'At ..the high temperatures considerable oxidation of 
the copper pipe takes place. The temperatures were measured with a 
platinum -platinum and 10%^ rhodium thermocoaiple which was checked 
against the boiling points of water and sulphur. The thermocomple 
was placed in the indenture in bulb G^which was easily made by pushing 
an iron rod into the hot glass. 

At the beginning of each ^run the apparatus was swept out by 
nitrogen gas which was let in- at stopcock L and out through the measuring 
aj)paratus iat the two way stopcock M. The compressed nitrogen used was 
freed from oxygen by passing it through an Ammonaical Ammonium chloride 
solution containing metallic copper in an apparatus similar to the one 
described by Badger.-^ 

The mercury apparatus is similar to the one designed by Prazer and 
Hoffman . In brief, the sample of gas is drawn into the gas burette P 
through two two-way stopcocks, IT and 0, by lowering the mercury leveling 
bulb R. The gas burette P is divided into two parallel limbs; one limb 
is a 2 cc. pipette, the other limb is made up of 5 bulbs of approximate 
capacity of 2 cc. each, as follows:- Ho. 1, 2.36 cc; Wo. 2, 2.77 cc; 
No. 3, 2.57 cc; No. 4, 2.71 cc; i>!o. 5, 2.81 cc ; and the volume from 
the top scratches over the bulbs and pipette to the stopcock N, which 
must be added to each reading is equal to 1.12 cc 

1. Badger. Jour. Ind. sixig, Chem. 11 , 1052 (1919). 
2. i'razer and Hoffman. Bull. No. 12; Bureau of iiines, 
Sampling and Analysis of ii'urnace Gases. 



12. 

In making a measurement after filling the gas burette, stopcocJc 
N is closed, the mercury level set at one of the scratches between the 
bulbs and stopcock 3 is then closed, stopcock is then turned connect- 
the gas burette with the compensator T. After the mercury levels in the 
compensator are balanced by raising and lowering the leveling bulb R 
stopcock S' is closed and the readings taken. The gas is then forced 
into the phosphorous pipette V. viihen all the oxygen has been absorbed 
by the phosphorous the gas is brought back into the gas burette and its 
volume remeasured as described above. Assume that the first reading 
was 10 CO., and the second 7 cc, then the partial pressure of oxygen 
in the gas is equal to 

-^ ■ '^■i" '^ 1 (barometer pressure-*- reading on manometer H) . 

The following measurements were made on a mixture of 10 grams of 
manganese dioxide and 6 grams of silver oxide, which mixture had been 
previously dried at 100°C in oxygen. The complete data which is 
necessary to calculate the decomposition pressures measured is given 
in the first table for the ascending readings only. A similar amount 
of data was used in calculating the descending pressures and all the 
decomposition pressures in the succeeding tables but it is not included 
in these tables because it would make them too bulky and add nothing 
to their value. 



13. 



Temp. 


1. J 

Pressure 
cm. Hg 


^ 

Volume of 


gas 


Oxygen 


Pressure 


Cent. 


Bar. 


Uan. 


start 


end 


absorbed 


ascending 


decending 


100 


75.5 


12.6 


15.03 


15.01 


0.02 


0.1 






150 


76.0 


17.4 


15.67 


14.81 


0.86 


5.1 




13.6 


200 


75.9 


6.1 


15.90 


12.79 


3.11 


16.0 




18.5 


256 


75.9 


10.4 


13.12 


10.10 


3.02 


20.0 




22.6 


315 


75.9 


20.4 


10.39 


7.30 


3.09 


28.6 




29.3 


370 


75.9 


30.2 


7.63 


4.60 


3.03 


41.8 




44.7 


400 





























Continuation of above table. Between each run the apparatus was 
swept out with nitrogen. 



Temp. 


Pres 


sures 


Centimeters of Mercury 




Cent. 


Run 
up 


2 

down 


Hun 
up 


3 
down 


Run 4 

up 


down 


150 


1.8 

3.8 
7.5 

13.0 

20.7 
H'emp. 

37 


£.6 

5.2 

7.5 

12.7 

23.0 
430°) 

.5 


0.8 
1.8 
4.6 
8.6 


e.8 

2.4 

5.0 

8.9 

17.1 

52.6 


0.2 
1.6 
3.6 
7.5 
17.3 

31.2 




200 




256 




310 




370 




450 





14. 

Material 20 grams Premy Oxide. Dried over CaGl2. 

PressuresAoentimetera of mercury. 

Temp. '. "'^ 

Run 1 Run S Run 3 Run 4 Run 5 
Cent. 

up down ^P down up down up down up down 

150 21.4 e.o 

200 24.5 0.2 5.9 

250 34.3 2.0 14.9 2.0 3.9 2.2 0.5 0.8 

300 34.5 7.7 16.6 6.5 3.5 4.7 2.2 3.7 

350 18.6 26.3 11.3 14.5 8.2 7.6 5.1 5.9 

400 33.1 35.0 20.2 23.1 14.1 17.0 12.8 12.8 

450 45.4 32.4 25.4 2^.7 



Material 10 grams CaO 8 grams AggO Dried over CaCig 



Tenfp. 


Pressures 


centimeters 


of mercury. 




Cent. 


Run 1 Run 2 
up down up 


down 


Run 3 
up down 


Run 4 
up down up 


down 


150 
200 

250 


23.0 
63.0 


52.7 
57.7 
61.7 


0.0 

0.6 

4.8 

51.2 

57.7 


47.5 
50.7 
53.5 

(T 


0.0 
0.4 




300 


1.0 




350 


1.0 








emp. 5201 2.8 





15. 

Material 10 gram Mn 2 grams CrO, Dried over CaClg. 
Temp. Pressures centimeters of mercury. 

Cent, 



^un 1 Run 2 Run 3 

up down up down up down 



150 4.5 

200 e.4 «> 

250 3.5 6.2 1.6 ^ 

300 7.1 8.3 4.0 

350 13.8 14.5 7.6 -^ 2.S 2.2 

400 21.7 20.6 11.8 +^ 3.6 3.6 

450 27.3 17.0 5.7 9.4 

500 31.3 12.1 



The following dissociation pressures of CuO were taken from the 
tables of Smyth and Roberts . 

Temp. Cent 983.5 1017.7 1043.0 1083.8 1165.7 1135.8 1204.4 1232.5 
Press. (gg ) 7.0 13.5 21.5 45.0 93.9 202 1609 4470 

Askenasy's values for the dissociation pressures of manganese 
dioxide are as follows: 

Temp. Cent, 382 406 448 478 536 

(Cm\ 
Hg) 2.3 7.7 12.5 20.0 104 



1. Smyth and Roberts. J. Imer. Chem. Soc. 42 2582, 1920 



16. 

The following dissociation pressures for AggO were calculated 
from Lewis's values of £0.5 atmospheres as the dissociation pressure 
over AggO at 302 C, by means of the simple Van't Hoff equation 

where ^rr. equilibrium constant; Q- heat of formation of 2 molecules 

of Ag20 or 12,800 calories; R r gas constant aild T the absolute 

temperature. 

Temp. Cent. 80 100 120 140 160 180 200 220 240 260 

Press. {^"^j 1.4 3.7 9.1 19.6 39.8 75.7 144 241 403 631 

The above data yt plotted on the following curves. The dissociation 
pressure curve for copper oxide is not shown on the dissociation pressure 
curves for the mixture of 10 grams of manganese dioxide and 2 grams of 
copper oxide because it would necessitate drawing these cjxrves to too 
small a scale. 

If it arS assumed that the oxygen pressures measured were due to 
a decomposition of a solid solution of the respective oxides, the curves 
are easily explained. In the case of solid solutions the shape of the 
curve is dependent on the initial composition of the solid solution and 
on the ratio of the amount of material to the gas space into which it is 
dissociating. It is for the latter reason that the curves from the 
static measurements are steeper, since the static apparatus had an inter- 
nal capacity of 50 cc, while the dynamic apparatus had a capacity of 
150 cc. 

In a mixture of two oxides A and B the partial pressure of oxygen 
over oxide A, assuming it is not affected by oxide B, 

and similarly the oxygen pressures over oxide B, assuming it is not af- 
fected by oxide A, where P amd P are partial pressures 



Oxygen pressure o^er /OjrFrerrif'Oxjc/e/.grOuO 




Tomp 



erature Cent 



%oo 



ZS^ 300 3Sd ^i^O" ^SO S'OO 

I Or^y^en Pressure. Qi^er lO^rCcfOj S/f^2,o 




73^ Joo tSd Joo -^so 



^00 4-So 




735^ ZOO ^^ 300 ' 3So 4^ 4S0 




^ 



17. 

and P and P are the pressures over the pure oxides respectively, 
a and b are the number of molecules of their respective oxides. If 
P , and P a and b should have such values that P w&a larger than 
P , on removing oxygen from the system, oxide A would dissociate 
solely until its partial pressure wae- equal to that of Oxide B, after 
which the two oxides wotild dissociate as a unit. 

The transition points of the manganese oxides in air are as follows: 

535**C 9ZZ^C 11600C 1,? 

MnOg — > Mn203 > 1^304 -^^ MnO 

(reversible in oxygen; 

Prom this table it might be assumed that manganese dioxide could 
be heated in air up to 5350C without the loss of oxygen. Accordingly 
Askenasy, using a static apparatus similar to the one described in this 
paper repeatedly evacuated and heated his manganese dioxide in oxygen 
up, to 400°C. in order to remove other gases, water, etc. before at- 
tempting to measure the oxygen pressure over his manganese dioxide. He 
was then measuring the decomposition pressures over an oxide of manganese 
which contained less oxygen than that required for manganese dioxide, 
ikanganese dioxide prepared by the method of i'remy cannot even be dried 
at room temperatures without the loss of oxygen. An analysis of the 
wet i'remy oxide, containing 75 % of water, using the oxalic acid method 
for determining the oxygen content in excess of MnO and the sodimn 
bismuthate method for determining the total manganese should be repre- 
sented by the emperical formula MnOg^^g . After drying the material 

1. Honda and Sone. Amer. Chem. Abs. 8__ 2648. 

2. keyers and Hotgers. Zeit. Anorg. Chem. 57 , 104. 



18. 



for several days in a dessicator over calcium chloride, it contair»d 
25 5^ of water and an analysis showed it to have a composition Mn0i^94. 
Results very similar to these were obtained in this laboratory by 
J^amaguchi . 

He prepared manganese dioxide by the electrolysis of a iain04 ..q 
solution. The permanganic acid formed at the anode decomposes spon- 
taneously into a manganese oxide which he found to have a composition 
Mn02 00 • ■'^ sample of this material after standing open to the air 
of the laboratory for one month was found to contain 20 % water and 
have a composition MnOi^QQ7 or as he expressed it 80 MQ02.1bi0.102H20. 

Another sample of the same material was dried in a despicator 
over 1^205 for 115 days. An analysis of this material showed it to 
contain 9 % of water and have a composition MnO-^^gg^ or again as he 
expressed it 29MnOgMn0.14H20. These high initial oxygen pressures 
over manganese dioxide fit in well with the views of jslliott and 
Storer,^ who give manganese dioxide the triple formula Mn20g.Mn0g, 
because they claim it to be oversaturated with oxygen from a study of 
its reduction products in hydrogen. vVright and Menke fix the lower 
temperature at which manganese dioxide starts dissociating in air 
at 21(&°C. it is probable that the extremely fine state of subdivision 
of manganese dioxide as prepared by the spontaneous decomposition of 
permanganic acid or by the action of sulphuric acid on potassium 



1. yamaguchi. Dissertation " ^n investigation of the uydrated Oxides 
of iiianganese derived from electrolytically prepared Permanganic acid. 

2. Slliott and Storer. Seit. ilryst. Mn. 48_ 447, 



19, 



permanganate, is accountable for this discrepancy in the lower tem- 
perature at which manganese dioxide started decomposing, as it is 
well known that the state of subdivision of a material is an import- 
ant factor in many of its physical properties; i.e. its solubility 
in different solvents, its vapor pressure, its decomposition pressure, 
etc. Manganese dioxide as prepared by the method of Fremy is so 
finely divided that when washed free from sulphates its jbrott^/#r7 
Movement will keep it from settling, in water, for days. 

In measuring the decomposition equilibrium of any material, 
hydrate, oxide etc. into a lower hydrate, oxide etc. there will be 
in general two cases which are illustrated by scheme II and III 
for the case of AggO 



® ® © © ^ ® & 

I Q Q Q Q O Q & 

(^ O @ ® Q Q '9 



® @ © 

n ® a ® a ® e t 
^ i£i e> 



III 



Q £> 9 

e a c i- 

^ 'f' i) 



B 9 (Bm 



20, 



Scheme I is a graphical representation of 7 molecules of silver 
oxide; where .^represents a silver atom and — represents an oxygen 
atom; since the atoms are supposed to be in continual motion. 

Scheme il is Hfe o w -i-se a s»aj?-8hat representation of the same seven 
molecules of Ag20 minus 4 atoms of oxygen. This scheme represents 
a solid solution where the dissociated molecules are so intimately 
mixed with the undissociated molecules that the undissociated mole- 
cules are forced to share their oxygen with the molecules that have 
already lost their oxygen. 

Scheme III is likewise a graphical representation of the same 
7 molecules of AggO minus 4 atoms of oxygen. In this scheme the 
dissociated molecules are pictured as being crystallized together 
away from the undissociated molecules.. In this scheme the decomposi- 
tion pressures over silver oxide would be independent of the composi- 
tion of the solid phase. Most decompositions are probably combinations 
of scheme II and scheme III as is shown by the difficulties met with 
in attempting to remove the last trace of any component from a com- 
pound. Although size of particles may play an important role in the 
initial high oxygen pressures over Premy oxide, the uniform drop of 
the oxygen pressure curves, as oxygen is removed, down to the values 
obtained by Askenasy and finally below his value and probably to 
zero pressure, is due mainly to solid solution formation as shown in 
scheme II. It is impossible to show from Askenasy's measurements of 
the dissociation pressures over manganese dioxide whether he was 
studying a solid solution or not because his highest measurement , is 
so far separated from his first four measurements, which do not lie 
on any smooth curve and furthermore his highest measurement is prob- 



21. 



ably in error due to other gases, water etc, which were not removed 
in his evacuation at 400^0. Smyth and Koberts in trying to confirm 
the work of V/8hler, who had prepared solid solutions of CuO and CugO, 
had their dissociated molecules of CU2O to crystallize out of the 
mixture as illustrated in scheme ill. ITrom the fact that Smyth and 
Boberts prepared their CuU from metallic copper which contained .01% 
impurities, it has been assumed by the writer that impurities play an 
important role in solid solution formation. In the case of Fremy oxide 
the impurit^f which cause the solid solution formation is potash , or 
in the case of the mixtures prepared the impurities are AggO, CuO etc. 
A sample of i'remy oxide made as described in this paper was analysed 
and found to contain 2.94 io of jigO, calculated to a dry basis, and 
after washing the material several times with strong HNO3 it still 
contained 2.5e;o KgO. In order to show how strongly manganese dioxide 
will hold potash, iVright and Menke dissolved a manganese dioxide, made 
from HNO5 and IMn04, containing 2.80/6 KgO in 50 parts of cold concen- 
trated HCl. The manganese dioxide obtained on reprecipitation with 
water still contained 1.16 :^ KgO. 

Le Chatelier found AggO to give a pressure of 10 atmospheres 
at 300 , and at this same temperature and 15 atmospheres pressure he 
was able to oxidize metallic silver up to 50 % AggO, while Lewis at 
this same temperature found AggO to give a decomposition pressure of 
19.6 atmospheres and had metallic silver separate out. jj'or some 
reason, difference in impurities if you like, Lewis was studying a system 
represented by scheme III while Chatelier was studying a system 
represented by scheme II. 

1. Le Chatlier. Zeit. x-hysik. Chem, 1 516 ia87. 



22, 



Guntz heated Ag AggO and KMnO^. in separate compartments of the 
same apparatus to 358° , obtaining an oxygen pressure of 49 atmospheres. 
On opening the tube the metallic silver had gained in oxygen and the 
silver oxide had lost oxygen so that both had the composition Ag^O. 
I'his again is a case of a solid solution, it is probable that a great 
many of these discrepancies are due primarily to the solid phase being 
represented by scheme II or scheme III. 

In order to determine how KgO, AggO CuO etc. are bound in manganese 
dioxide the following simple experiments were carried out. 'j.'he older 
literature contains many references to the so called manganites of the 
general formula M O.xMnOgyHgO, where M is any metal, and x and y numbers 
from 1 up to several hundred. Below are given a few of these references 
to the silver manganites^ alone and in some cases the formula and method 
of preparation is given. 

V<m BemmlenV however, thinks these compounds as well as the similar 
compounds of silica and tin to be adsorption compounds, balinger prepared 

1. Guntz. C. E. 128 . 996 1899. 

2. Silver Manganites. 

AggMnOg, WBhler P. A. 41^ 544 1837 



ii. Kose j:-. A. 101 229 

r 



23. 

a silver manganite from manganese nitrate, silver nitrate and ammonium 
persulphate which when dried at 150*^0. had a composition AggO.lOMnOg. 
SHgO. it lost water and 5 atoms of oxygen at a red heat and was then 
AggO.SlIngOg. The mixture of 10 grams of manganese dioxide and 6 grams 
of silver oxide, which had beea heated to 500*^0. in the static measurements 
after being heated in air at SOOOC, for several days was analysed by the 
following procedure. It was first treated with dilute HgSO^ and 
standard H2C2O4 with the idea that any metallic silver if present would 
remain t=, undissolved and could be filtered off and determined. The 
material, however, went entirely into solution. It was later shown that 
metallic silver will dissolve in dilute H2SO4 in the presence of MnOg . 
The solution was then made up to a definite volume in a volumetric flask. 
The total manganese was determined in a pipetted sample and the excess 
HgCgO^ and silver were determined in the renaining portion. The results 
of this analysis can be expressed^ by the formula 1.8 AggO.MnO.SMnOg • 

In order to study the silver ion concentrations in water over the 
manganese dioxide and silver oxide mixture, a concentration cell was set 
up consisting of a normal calomel half cell against various mixtures of 
manganese dioxide and silver oxide in a .INKNO5 solution as a conducting 
liquid. The voltages measured and the calculated silver ion concentrations 
are given in the following table. The temperature was kept constant at 
25°C. 



24. 





Materials 


parts by weight 




AggO 


1 


MnOgSAggO 


1 


" 2 " 


1 


" 1 " 


5 


" 1 n 


5 


" 1 " 


10 


" 1 " 



Voltage 
measured 
e.3075 
0,£874 
0.2965 
0.2960 
0.2828 
0.2705 
0.2520 



Cone. Ag. 
gram lone/liter 
2.29 X 10 "^ 
1.07 X 10 '^ 
1.41 Z 10 -4 
1.41 2 10 -^ 
0.93 X 10 -4 
0.57 X 10 -4 
0.25 X 10 -4 



The silver ion concentrations were calculated from the formula 
S 5 ilgg^i^ 4- RT log. C 
where E = voltage measured; Eq^^^^- voltage of a normal calomel half 
cell or 0.521 volts; Era constant and T - the absolute temperature, 
The product R and T for 25°C. ^ 6.059. If the voltage 0.3075 for 
the cell 



Ag 



,1NXN0, 



.1 EWO, 



Calomel 
half cell 



be substituted in the above formula their results 

0.3075 = 0.521 +.0591og.C 
a solution of this equation gives a value of C, the silver ion con- 
centration, equal to 2.29 X 10 '^ gram ions per liter. If complete 
dissociation of AggO is assumed for this dilution then the solubility 
of AggO ; # X 2.29X10 '^ or 1.14 mols per liter, since each silver 



25. 

oxide molecule will give two silver ions if completely dissociated. 
These results are in excellent accord with the accepted solubility 
of AggO , From Seidel's solubility tables the solubility of AggO in 
water at 25°C is equal to 0.91X10 "'^ mols per liter. From the above 
tables it is seen that each addition of manganese dioxide to silver 
oxide reduces the silver ion concentration of the silver 6xide;an ad- 
dition 6f 10 parts by weight of manganese dioxide to 1 mol of AggO 
or 30 mols of manganese dioxide to 1 mol of AggO reduces the silver ion 
concentration to 1/10 of its value. Again it is to be noticed that 
these results fit in very well with the proposed theory that the mix- 
ture of manganese dioxide and silver oxide is a solid solution. 

Because so many theories have been proposed for catalysis and 
the mechanism of catalysis, which have little theoretical or practical 
value, it seems out of place to propose any new theory of catalysis 
based on the scant data at present available on mixtures of manganese 
dioxide with various other oxides as to their catalytic activity, it 
does not seem out of place, however, to call attention to the advantages 
of a solid solution as an intermediate step in catalysis. 

In a solid solution oscillating back and forth between scheme I 
and scheme II as illustrated for silver oxide, all atoms or aggregates 
of atoms take part in the reaction since they are all vitally affected 
by any change in the solid phase which:. would not be the case if part 
of the atoms were crystallized aut as illustrated in scheme III. 
Impurities decrease the surface energy of the different aggregates and 
in this paper it has been assumed that they play an important role in 
solid solution formation. Materials of extreme purity are usually 
unreactive, which bears out this idea. V/right and Menke volatilized 



26. 



the potash out of a sample of MTOg by heating it in a current of 
hydrogen. If the potash was completely removed the yellow'MnO 
would not take up oxygen again. However if these impurities are 
easily fusible they will sj>op up the capillaries in the catalyst 
and fornthis reason prevent the material from being catalytic. It 
is probably for this reason that the copper oxide manganese dioxide 
mixture is a better catalyst than the silver oxide manganese dioxide 
mixture. 

The advantage of the higher dissociation pressure of silver oxide 
over that of copper oxides is more than counter balanced by the greater 
ease of fusibility of metallic silver as compared to the fusibility of 
CugO. 

SUMMARY. 

1. A new method is described for measuring the dissociation *- 
pressures of metallic oxides in which' it is unnecessary to free the 
oxide from different adsorbed gases, water etc. before making measure- 
ments. It is also applicable to other dissociations where none of the 
gaseous components are condensable at ordinary temperatures. 

2. Fremy Oxide (LIn02 prepared from lUnO^ and H2SO4) mixtures of 
it with AggO, CuO etc. indicate by a study of their dissociation 
pressures that they are solid solutions. 

3. The advantages of a solid solution as an intermediate compound 
in catalysis is discussed. 



27. 



Biography. 

William Smith iinglish was born October 2» 1896 in 
iPaulaboro, New Jersey. His early education was received 
in the public and high school of this town. In 1918 he 
received the B. S. degree from the University of Peansyl- 
vania. 



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