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This is to certifg that the

thesis entitled

A hechanism for the Decomposition of Potassium Ferrate in

Date

0-169

Aqueous sodium hyuroxide

presented In]

Orville N. Hinsvark

has been accepted towards fulfillment

of the requirements for

JL_$L_ degree in W

’9 ,,

Major [rdfessor

May 27 , 1952

 

A MECHANISM FOR THE DECOMPOSITION
or
POTASSIUM FERRATEtVI)
IN
AQUEOUS SODIUM HIDROXIDE

By
Orville N. Hinevark

A THESIS

submitted to the School of Graduate Studies of Michigan
State College of Agriculture and Applied Science
in partial fulfillment of the requirements
for the degree of

MASTER OF SCIENCE

Department of Chemistry
1952

CKICMlI'I'HY DEFT.

ur—T" P" V A Q}
l V . C)”, ' /

 

AQNOWLEDGMENT

 

The author wishes to express his sincere
for the guidance offered to him by Dr. Andrew
Dr. Elmer Leininger throughout this work. He
to thank Dr. K. G. Stone, Jr. for the helpful
which lead to the conclusion and to thank his
devoted a great deal of time to the typing.

092m "a"?

appreciation
Timnick and
also wishes
suggestions

wife who

STRACT

A study was made on the decomposition of potassium fer-
rateWI) in highly alkaline media. Polarographic techniques
utilizing a rotating platinum microelectrode proved unsatis-
factory in determining the oxidation states through which
ferrate(VI) passes in its reduction. Alkaline solutions of
potassium ferrate(VI) were found to Obey Beer's Law at
500A747’; and through spectrophotometric measurements, the
kinetics of its decomposition could be determined. Utiliza-
tion of these kinetic data was made in prOposing a possible
mechanism for ferrate(VI) decomposition. In order to explain
the apparent unstable eQuilibrium, free hydroxyl radicals
were suggested as one of the products which caused a rever-
sal of the initial reaction by oxidation of the ferrite(III).
The subsequent oxidation of the free radicals or their re-
combination product, H202, by ferrate(VI) was suggested as
the primary explanation for the eventual completion of fer-

rate‘VI) decomposition in aqueous alkaline media.

Iggy; _O_F_‘ CONTENTS

PAGE
List of Tables . . . . . . . . . . . . . . . . . . . . 1
List of Figures. . . . . . . . . . . . . . . . . . . ii
Introduction . . . . . . . . . . . . . . . . . . . . . 1
Experimental Work. . . . . . . . . . . . . . . . . . . 3
Apparatus.................... 3
Reagents. . . . . . . . . . . . . . . . . . . . . 3
Solutions . . . . . . . . . . . . . . . . . . . . 3
Experimental Procedures . . . . . . . . . . . . . 4
Polarographic Studies. . . . . . . . . . . . 4
Colorimetric Studies . . . . . . . . . . . . 4
Kinetic Studies Using a Beckman DU Spectro-
photometer . . . . . . . . . . . . . . . . 4
Experimental Results and Interpretations . . . . . . . 6
Polarographio Data. . . . . . . . . . . . . . . . 6
Colorimetric Determination of Potassium Ferrate . 9
Kinetic Studies of Decomposition of Potassium
Ferrate . . . . . . . . . . . . . . . . . . . . 12
Substantiation of Proposed Initial Equilibrium
Reaction. . . . . . . . . . . . . . . . . ._. . 14
Free Hydroxyl Radical Reaction. . . . . . . . . . 22
Proposed Main Degradation Reaction. . . . . . . . 25
Discussion . . . . . . . . . . . . . . . . . . . . . . 28
Conclusions. . . . . . . . . . . . . . . . . . . . . . 34
Appendix . . . . . . . . . . . . . . . . . . . . . . . 35
Bibliography . . . . . . . . . . . . . . . . . . . . . 49

ngg or TAQLES
TABLE
1. Relationship of k1 to Normality of NaOH Solvent.

10.

11.

12.

Rate of Decomposition of K2Fe04 in 4N NaOH as
Found by Polarographic Amperometric Technique.

Rate of Decomposition of K2Fe04 in 4N NaOH as
Found by Polarographic Amperometric Technique.

Absorbtion Data of Potassium Ferrate in 8N NaOH
With a Beckman Spectrophotometer . . . . . . .

Absorbtion Data of Essentially Decomposed
Potassium Ferrate Solution - To Show Effect
of Decomposition Products. . . . . . . . . . .

Determination of Conformity of Potassium Ferrets
Solution in 8N NaOH to Beer's Law. . . . . . .

Kinetic Runs of the Decomposition of Potassium
Perrate in 8N NaOH Made on Beckman Spectro-
photometorat5000....oo.....ooo

Kinetic Runs of Potassium Ferrate Decomposition
Made on Beckman Spectrophotometer With Varying
Concentrations of NaOH . . . . . . . . . . . .

Kinetic Runs of the Decomposition of Potassium
Ferrate in 8N NaOH Saturated With Fe203-xfizo .

Kinetic Runs of the Decomposition of Potassium
FarrateianaOHatoce e so es s 0 see

Kinetic Runs of the Decomposition of Potassium
Ferrate in 8N NaOH With Glass Wool Added to

ReactionVassol...ooo.........o

Kinetic Runs of the Dgcomposition of Potassium
Ferrate in SN NaOH n the Absence of Light . .

PAGE

14

55

36

57

39

41

42

45

45

46

47

FIGURE

1.

II.

III.

VI.

VII.

VIII.

XI.

XII.

pIST pg; FIGURES

Graph Showing Rate of Deuomposétion of Potas-
sium Ferrate in 4N NaOH at O C as Determined
on the POlarOgrapho e e e e e e e or) e e e 0

Absorption Spectrum of Potassium Perrate in
8N NSOH o e e e e o e e e e

O O O O O O O O 0

Graph Showing Beer's Law Adherence of Potassium
Ferrate 1n 8N naOH. e e e e e e e e e o o e 0

Graph Showing Typical Rate of Decomposition
Curves of Potassium Ferrate in 8N NaOH . . .

Graph Showing Rates of Decomposition of
Potassium Ferrate in Varying Concentrations
Of NaOH 0 e e e

O O O O O O O O O O O O O O 0

Graph Showing Effect of NaOH Concentrations
0n FirSt Rate ConStant. e e o e e e e e e e e

Graph Showing Rate of Decomposition of Potas-
sium Ferrate in 8N NaOH Saturated With
Hydrous Ferric Oxide p. . . . ... . . . . . .

Graph Showing Rate of Decomposition of Potas-
sium Ferrate in 4N NaOH at 0°C. . . . . . . .

Graph Showing Rate of Decomposition of Potas-
sium Ferrate in 8N NaOH, Glass Wool Added
to Reaction Mixture . . . .

Graph Showing Decrease in Potassium Ferrate
Concentration in 8N NaOH in Absence of Light.

Graph Showing Rate of Decomposition of Potas-
sium Ferrate in 8N NaOH .

O O 0 O O O O O O 0

Graph Showing Rate of Decomposition of Potas-
sium Ferrate in 8N Na0H as Compared With
calculated Curve. e e e o e e e e e e

11

PAGE

11

13

15

16

18

20

24

26

30

31

A MECHANISM FOR THE DECOMPOSITION
or
POTASSIUM FERRATEWI)
IN
AQUEOUS SODIUM HIDBOXIDS

W

Ircn in the higher oxidation states of plus four and six
has been known for a long time. Bohnson1 showed that the cata-
lytic decomposition of hydrogen peroxide was due to the forma-
tion of ferric acid which subsequently reacted with the hydrogen
peroxide to form oxygen and water. Bray and Gorin2 have sub-
mitted some evidence to show the presence of low concentrations
of ferryl ion, PeO‘z, in ferric-ferrous systems and in the
catalytic decomposition of hydrogen peroxide with ferrous iron.
Under certain limited conditions it is possible to form the
perferrite ion, re03=, but this appears to be unstable with
respect to decomposition into the plus three and six oxidation
states.5

Solutions involving the ferrate ion are difficult to
study because of their rapid decomposition unless the aqueous
medium is highly alkaline. The slightly soluble barium fer-
rate(VI), SaPe04, appears to be the most stable of its salts.

It can be suspended in dilute acetic acid with only slow de-

composition. In mineral aside, however, all ferrates decom-
pose rapidly with the formation Of iron in the plus three
oxidation state. If the acid is nonoxidizable under these
conditions, there is a rapid evolution of oxygen. In an
aqueous solution the following reaction takes place:6

4K2Fe04 + 10H20 ——’SKOH + momma 1- 302
In strongly alkaline media the stability of the ferrates is
increased to such an extent that a study of the kinetics of
its decomposition is possible.

The relative instability of the ferrates, coupled with
the difficulties involved in obtaining them in the pure state,
probably explains why the work pertaining to the study of this
oxidation state has been limited. The latter detriment to
this study has been overcome somewhat bypa recently published
paper7 which describes a method for preparing potassium ferrate
in a relatively pure state.

This study was made to ascertain the oxidation states
through which potassium ferrate passes in its decomposition
in an alkaline medium.

The report covering the work done in this study concerns
itself with a possible mechanism for the decomposition of
potassium ferrate and the experimental evidenca used to sub-

stantiate the prOposed mechanism.

EXPERIMENTAL WORK

Assassins.

A Sargent Hbdel XXI polarograph was used in an attempt
to ascertain the steps through which potassium ferrate passes
in its reduction. Since mercury is attacked by the potassium
ferrate solution, a dropping mercury electrode could not be
used; and a rotating platinum microelectrode was employed.
The electrode was made, with slight modifications, according
to specifications'given by Lingane4.

In the kinetic studies a Beckman D U spectrophotometer
with corex cells was used. Unless otherwise stated, the tem-
perature was maintained at 30° 2,0.20 by a constant temperature

bath.. The time was measured with a stopwatch, the readings
made within‘1_10 seconds.

Essssata.

The potassium ferrate was prepared according to Schreyer's7
procedure; and its purity was ascertained using Ichreyer's8
chronic oxidation method. The purity of the different prepara-
tions analyzed by this method varied between 62.5 and 95.5%
potassium ferrate, the impurities apparently being potassium

hydroxide and ferric oxide.

Salutisns.

The concentration of the potassium ferrate solutions was

approximately one millimolar with respect to potassium ferrate

(VI). The sodium hydroxide solutions used as the solvent
varied in concentration and were prepared by weighing C.P.

grade sodium hydroxide and dissolving it in distilled water.

Experimental Procedures,

__ '22lazggzaphig Studies: These studies were made in the

usual way. The solutions of potassium ferrate were introduced
into the polarographic cell; and the diffusion current, caused
by the reduction of the ferrate ion on a rotating platinum

microelectrode, was recorded on a polarogram.

Colorimetgig Studies: The aqueous sodium hydroxide solu-

tions were prepared and brought to the desired temperature,
50°C, in a constant temperature bath. Weighed quantities of
the analyzed potassium ferrate were added to the sodium hydrox-
ide solution flasks, and the optical density of each resulting
solution was measured at 500Am43five minutes, ;,10 seconds,
after dissolution of the potasSium ferrate. Using these

data, adherence to Beer's Law was observed.

Klgetig Studies Using A Beckman 2,_,§pectrophotometer:
The kinetic studies were performed in the following way:

 

(l) The aqueous sodium hydroxide solutions were prepared
and placed in 100 ml volumetric flasks.

(2) The sodium hydroxide solution was then placed in the
constant temperature bath and permitted to attain the desired

temperature. In this work the temperature was usually 30°

3.0.20.

(5) When the solvent had attained the desired temperature,
a weighed quantity of potassium ferrate was introduced into
the flask.

(4) The decomposition of the potassium ferrate was fol-
lowed colorimetrically by withdrawing samples from this flask
and introducing them directly into the absorption cell. When
this was done, the Optical density of the solution was meas—
ured at soommf. The Optical density and the time of the read-
ing were recorded.

(5) Since Beer's Law is obeyed and the optical density
is proportional to the concentration, the Optical density is
used directly in the calculations and graphs. The curves ob-
tained by plotting time versus optical density are found
throughout the report. The data used in the plotting of these

curves are found in the appendix.

W RESULTS m IyTERPRETgTIONS

The experimental results described in greater detail on
the following pages show the data used in arriving at the con-
olusions found in the report.

The first portion of the work involved the use of the
polarograph utilizing a rotating platinum microelectrode;
while the second and major portion concerns itself with kinet-

ic studies made with a Beckman D U spectrophotometer.

W441”.

If there is sufficient difference between the free ener-
gies of the oxidation states through which the ferrate passes
in its reduction, these states should be recorded on a polaro-
gram as a series of steps.

The electrode as described by Lingane could not be used
since apparently the insulating wax on the shaft was not stable
under the conditions employed. Leaks developed in the wax and
the cell formed caused an enormous increase in the recorded
current. An attempt was made to cover the shaft with glass
tubing through which a platinum wire lead protruded. This de-
sign caused considerable stirring which apparently prevented
the obtaining of reproducible diffusion currents. The polaro-
grams obtained with varying potassium ferrate and sodium hy-
droxude concentrations had the same general shape, but no breaks

were observed which would indicate a stable oxidation state

between the known plus three and plus six levels.

While reproducible results could not be obtained from
one run to the next, the diffusion current appeared to be pro- w
portional to the concentration of the ferrate in any given
solution; This observation is substantiated by the data de-
rived from kinetic runs on the decomposition of potassium
ferrate in four normal sodium hydroxide. The temperature was
maintained at 0°C in order to increase the stability. The
applied emf was set at v.15V versus 0.1N NaOH, Ego electrode,
and readings were taken at definite time intervals. This was
continued until all the potassium ferrate had decomposed. In
order to insure complete decomposition to obtain a base line,
hot water was added to the bath until the color, characteristic
of the ferrate ion, had disappeared. The polarographic cell
and contents were then cooled to 0°C and the diffusion current
measured. This diffusion current served as a reference line
from which step height measurements were made. The step height
at definite times was recorded; and when Log step height versus
time was plotted, a straight line resulted. The result of this
experiment is shown in Figure I. There is a great deal of dis-
persion of the points in the individual curves, but the slopes
appear to agree closely with rate constants of 6.76 X 10"":5min."1
and 6.83 X 10-3 min.'1. Midway through the rate run represented
by curve 2, there was a sudden displacement of the diffusion
current. This can probably be attributed to an inherent fault
in the electrode, possibly a difference in the stirring. 0n

the basis of the above observations, it is probable that if

 

 

 

 

 

 

'———————-

Time min '

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2.__ __
. I I I I I I I l I a
20 40 60 ST, . leO T20 I415 I60 ISO 200 2217 not

FIG.I GRAPH SHOWING RATE OF DECOMPOSITION

OF POTASSIUM FERRATE IN 4N NAOH
AT O'QAS DETERMINED ON THE
POLAROGRAPH.

stirring could be eliminated, the diffusion current would be
proportional to the concentration; and reproducible results

could be obtained.

WWQWW.

Since a solution of potassium ferrate is highly colored,
it seemed quite logical to attempt a colorimetric analysis
for the determination of ferrate concentration. This would
afford a convenient and rapid method for following its decom-
position in various media.

The absorbtion spectrum was determined to show the wave
length at which a solution of potassium ferrate had maximum
absorbency, the results can be seen in Figure II. The scan
was made by taking readings every lOnWon solutions of potas-
sium ferrate dissolved in 8N NaOH. The absorbtion spectrum
obtained in this way showed that the solution absorbed strongly
in the vicinity of 600074. By making a scan with a solution
of xzreo, which had essentially decomposed, the decomposition
products, as indicated by curve 2, were shown not to interfere.

By setting the wave length at SOqu and taking readings
of the optical density, D, of the analyzed ferrate solutions,
it is shown in Figure III that Beer's Law is obeyed over the
range of concentrations used. Because of the instability of
the solutions, it was necessary to make the readings at defi-
nite elapsed time intervals. If the elapsed time were five
minutes :,10 see. as in Figure III, the extinction coefficient
was found to be 5.15 l g‘lcm‘l. This was found to hold over

10

mk<mmwu 2:.mm4k0d do

WCQICIEQEEI Cu ....OIV O>M>>

 

 

.Iodz 20 Z.
Dahommm zo_._.n_momm< ”:9“.

 

 

 

 

 

 

goo am I D.» 06
l a a l l _ a _ _ l _ a _ fl _ l _
PI
I . ION
Q
I IS
ZOEbJOm QmmOdéoowO J. .
Q. 0. . &
I -0. Ian
0 .... .. O
, . O a
I G G/ I
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J; :4: m -IVIIW *w Mo *Ifl _ 4.? j b b:hv _ _ -c.

11

 

 

O 6 r T /IF
I'" —-J
0.4}— ._..
I
j 0
Qoi— _.
I ./
I r—- / —I
9’
//
02—— _.
OI —— /I
/
//
O I). IO

 

 

 

I
”— 0850/70 9///'ter
FIG-IIIIGRAPH SHOWING BEER'S LAW ADHERENOE
OF POTASSIUM FERRATE IN ON NAOH..

 

12

the concentration range from O to 0.125g KZFeO4/1. The
potassium ferrate in this expression is expressed as pure

potassium ferrate.

ggggti2.8tugigslgg Decompopition.gg Potassium Ferrgt .
As was stated previously, the purity of the prepared

material analyzed by Schreyer's method varied between 62.6
and 95.5% potassium ferrate(VI).

Kinetic runs of the decomposition were made using the
potassium ferrate samples of varying purity. The only dif-
ference in the experimental results was in the time required
to attain the equilibrium condition. On the basis of these
observations, it seems probable that the results should be
applicable to the pure material and the impurity is of little
significance in the reaction.

The shape of some typical decomposition curves plotting
log D versus time is shown in Pig. IV. As can be seen, a
straight line is observed over the first portion of the curve
indicating a first order reaction. The slope then gradually
decreases as competing reactions cause the reaction to become
complex. The curve then finally graduates into another stralght
line function.

The shape of these curves suggests the existence of a
first order reaction at the beginning, the rate of which is
rapid, relative to the other reactions. Since the SIOpe
gradually decreases and graduates into another straight line

whose slope is much less, it seems probable that the reaction

15

 

 

 

 

OS I I I

 

I21“—

 

 

“J IO JG]. :0 . “IL A BC
[me If? m/n.

FIG. IV GRAPH SHOWING TYPICAL RATE OF
DECOMPOSITION CURVES OF POTASSIUM
FERRATE IN ON NAOH.

14

establishes an equilibrium with the relatively slow decompo-

sition of one of the products.

finbsisniisiiss.af.Proposes.Iniiialuissiliarissflasssiiass
In order to test the hypothesis concerning the unstable equi-

librium reaction, kinetic runs were made varying the sodium
hydroxide content of the potassium ferrate solutions. Sodium
hydroxide concentrations from four to nine normal were used.
(See Fig. V) The rates of reaction, found by taking the slope
over the first straight line portion, were determined; and
these rate constants were plotted against the normality of
sodium hydroxide. Figure VI shows that the rate constant de4
creases in direct proportion to the normality of the HaOH.
This suggests the existence of a relationship involving the
concentration of hydrOxyl ions to the first power.

TABLE
RELATIONSHIP OF k1 TO NORMALIT! OF NaOH SOLVENT

Con. x1 x 102 min.
on 2.92
as 3.61
7N 6.07
ON 7.68
an 8.66
4N 10.07

3 report the reaction:

Grube and Gmelin
ZNazFe04 t 6e‘ ZCFINagle204 + ---
to be reversible at an iron anode in 40$ NaOH at 50°C, hydrous
ferric oxide being quite soluble in concentrated alkaline

solutions.

15

 

00

0.01

 

 

{I

e 4N NaOH
O OH NaOH
5 6R Nam-I
-o- 7N" NaOH
c—SN NaOH
40 SN NaOH

 

 

 

—7€'me in min.
FIGV‘GRAPH SHOWING RATES OF DECOMPOSITION

TT
L‘s

 

.___).

OF POTASSIUM FERRATE IN. VARYING
CONCENTRATIONS OF NAOH.

16

 

 

 

 

._ \\ __..
_ I __
4L— \ ——
F o \ . —
aI— __
1’s I I Is I; I: (I I {'3 1]? Id III
—— Norma ht

FIC.VI:CRAPH SHOWING EF I-‘FECT OF NAOH

CONCENTRATIONS ON FIRST RATE

CONSTANT.

 

17

The validity of this reaction was tested.by saturating
an an flaOH solution with Fezoa-xHZO, adding 12Fe04, and sub-
sequently making a kinetics run on the ferrate decomposition.
From the results or this run, shown by Figure VII, it can be
seen that the initial rapid rate or decompositiOn, Observed
in ferrate solutions not containing the added FezosoxHQO, is
not apparent. The result is essentially a straight line from
the start. The slope of this straight line is the same as
the second straight line portion or the normal decomposition
curve. ’

The results obtained in these measurements apparently are
not in full agreement with the reaction as written and proposed
by Grubs and Gnelin to take place at an iron anode.

It appears that the amphoteric species of Fe203°xH20
exists not in the form or Fe204=, but rather as Feog'. This
conclusion is based.upon the apparent first order reaction
which occurs initially. If the dimer were formed immediately,
the reaction should.be more complex.

In order to further test this reaction, it would seal
that if the ferrite ion concentration were reduced to such an
extent that its concentration never became great enough to
establish the unstable equilibrium condition, the reaction
should continue to be first order with respect to the ferrate
decomposition throughout the whole reaction.

The solubility of the ferrite would be expected to de-
crease with decreasing alkalinity and with decreasing tempera-

ture. Taking advantage of these suppositions, a rate run was

 

 

l I I T T I
9'— .—
.8I’— _
.7— -
Bi? 0 _I

0
Q5? R ‘4
a
‘1
.4I— \\o-\ _
S 0-»- k
0%“ l ““““ Q 2
O~ "“0"" ‘
... I WITH F2203): H20 '—
2 NORMAL CURVE
I I0 I I

 

 

10 2e]— 40 +750 6U
Ime 3If min.
FIG VII“ GRAPH SHOWING RATE OF DECOMPOSITION

OF POTASSIUM FERRATE IN 8N NAOH
SOAXTUEATED WITH HYDROUS FERRIC
ID .

 

 

19

made at 0°C with.the potassium ferrate dissolved in four nor-
mal sodium hydroxide. As can be seen in Figure VIII, a straight
line is obtained when log D is plotted.against time. There is
considerable amount of dispersion Observed in the points. This
dispersion can‘be attributed to the moisture condensation which
takes place on the windows of the cells. The measurements made
in this run do not take into account the changes in the rate
constants, both forward and reverse, with changes in tempera-
ture.

Since increasing hydroxyl ion concentration causes an in-
crease in the stability of the ferrate solution, and since pre-
sumably only sodium hydroxide, water; ferrate, and ferrite are
present, water must be the substance being oxidized by the
Fe04'. This might be expected if the ferrate ion, De04', exists
as such in solution, the coulombic forces of repulsion would
be prohibitive for collision and subsequent reaction to take
place between the ferrate and hydroxyl icns..

The oxidation of water might be expected to occur in steps,
eg.

FeO4' .- 2320 =2:- reoz" + 503. I- OH"

This reaction fits the data presented thus far and can prob-
ably be further substantiated by referring to the preparative
method for K2Pe04.

Hypochlorite ions in a water solution can conceivably
exist as GlOoHZO', or represented in a different way, HOClOH’.

This could react as follows:

'l'lll I‘ll I; ‘11!"

2O

0.0 ...4 19.2 zv. z. uhdmmwu 23.mm<._.on.
....O zo_._._m0dzoouo mo mkdm 02.39% Inamc u=_>.0_.._

.CILk 9: WE‘
3% DESI PE. MOON amt our DC? DE 1: be om Lee ON

 

 

 

I I

 

_Ja4_a_fifl___la

 

 

 

21

(1) o/snocmn-ts 3/201“ + cos-

(2) OH" t 30H. 4- FeOg' z; Fe04= .- 21120
Reaction (2) shows a possible reversibility Of the reaction
for the formation of ferrate by a reaction with hydroxyl radi-
cals and ferrite.

The free energy of formation of hydrogen peroxide by com-
bination of two hydroxyl radicals is quite large', estimated
to be Al" . -21 kca1.5 On this basis, it might be expected
that the combination of hydroxyl radicals might be rapid and
result in the following reactions:

303- -%>-3/23202

no; . 2320 tareoz" «- 3/211202 .- or
This would indicate the existence of an equilibrium involving
the production of ferrate through the oxidation of ferrite by
hydrogen peroxide. This possible set of reactions is not sub-
stantiated by experimental observations. If the indicated equi-
librium is established, the addition of hydrogen peroxide should
stabilize the reaction mixture. The addition Of hydrogen per-
oxide to the alkaline solution of potassium ferrate greatly in-
creased the rate of decomposition of ferrate. The rapid evolu-
tion of oxygen prevented the possibility of following the de-
composition colorimetrically. The rate, however, appeared to

be immeasurable by this technique.

 

' The values of the free energies involving free hydroxyl
radicals must be considered as Latimer reports, 'strictly
tentative'.

22

These observations are substantiated by Latimer. As
Latimer reports, while the free energy of formation Of hydro-
gen peroxide by combination of two hydroxyl radicals is rela-
tively high, there is no experimental evidence to show that
the reaction rate is rapid.

The increase in the rate Of decomposition of potassium
ferrate by hydrogen peroxide may not be expected since Bohnson
showed that ferric acid is formed as an intermediate in the
catalytic decompOsition of hydrogen peroxide by plus three iron.
lhile this happens'in neutral solutions of hydrogen peroxide,
the addition of hydrogen peroxide to a strongly alkaline solu-
tion of hydrous ferric oxide produced no ferrate, as would be
evidenced by the appearance of the intense color characteristic

of ferrate.

In: M Miss; Miss: 0n the basis of the above
mentioned facts, it seems that while the combination of 03- may

take place, the effect is relatively small. If the concentra-
tion of free hydroxylradicals is high enough, one might ex-
pect the following reaction to take place:

5030 -*"3/402 9 5/2320
as a step preventing the reaction from attaining and maintain-
ing a condition of a stable equilibrium. This would probably
occur through combination of two hydroxyl radicals and the
subsequent oxidation by a third free radical. This hypothesis
is based upon the fact that the free energy Of formation of
atomic oxygen by interaction of two hydroxyl radicals is Quite

positive.

25

20H- —, 1120 .- o A r = 8.5 kcal.5
Since the free energy change is positive, this particular
reaction probably would not take place appreciably unless
the concentration Of free radicals were high or unless ex-
ternal energy were applied. Since the reaction between
hydroxyl radicals apparently takes place to some extent,_
there must be combination first and subsequent oxidation.

In order to determine the extent recombination plays in
the reaction, glass wool was added to the reaction vessel.
Radical combination is reportedly a surface reaction; there-
fore, this reaction should be enhanced by the increased sur-
face presented by the glass wool. The effect, an extension
of the time required to attain equilibrium, is seen in Fig. IX.

This effect would be expedted since the reversal of the
proposed first reaction is dependent upon the reaction between
ferrite and free hydroxyl radicals. One would expect, however,
that if this were the main reaction preventing the attainment
of a stable equulibrium condition, the effect would be more
pronounced than that whach was Observed, possibly preventing
attainment of the unstable equilibrium condition.

The above Observations appear to substantiate the previ-
ous claim that the reaction:

30H. -—>3/2H20 4- 3/402
is not Of great significance in the decomposition scheme of

potassium ferrateIVI).

 

3%—

0.8”—

 

IWITH GLASS WOOL
2NORMAL CURVE

 

 

02% 11‘) 7#210 1,1021% :10 7
[me Ion mm.
FIG. IX' GRAPH SHOWING RATE OF DECOMPOSITION

OF POTASSIUM FERRATE IN ON NAOH.

GLASS WOOL .ADDED TO REACTION
MIXTURE.

 

 

 

25

2:922:31 .lisin Degradation Beatles: 0n the basin of the

previously mentioned observations, there must be at least one
more reaction which would‘be the major contributor to the in-
stability of the equilibrium condition. The main degradation
reaction proposed is as follows:

Fe04= c SOHF-—ar 8/202 e Peoz' & OH" 7 320

This supposition is supported by the fact that the de-
composition is extremely light sensitive.

All of the previous reactions proceeded under essentially
constant illumination from tungsten lamps; however, when the
light was prevented from reaching the solution by using a
blacked out bottle, there was a pronounced change in the ob-
served results. This can be seen by referring to Figure x.
The initial slope was essentially the same; but the unstable
equilibrium condition was attained much more rapidly; and the
subsequent decomposition took place at a much slower rate.

According to Grotthusse's First Law of Photochemistry,
“only the rays which are absorbed can produce a chemical
change.‘I Since tungsten lamps emit primarily visible and
infra-red radiation and since only rays up to the relative-
ly near infra-red region would be able to penetrate the re-
action vessel, it would seem probable that the activation
frequency must be in the visible region Of the spectrum. In
the reaction mixture, the only material which absorbs in the
visible region is the potassium ferrate. This suggests that

the reaction affected would be the reaction between the ferrate

 

 

 

 

 

 

O.3I-
I ABSENCE OF LIGHT
2 NORMAL CURVE
I
0? IO 3I 4lU SIP 0.

 

2107..
lme 3In mm

FIG'Xi GRAPH SHOWING DECREASE IN POTASSIUM
FERRATE CONCENTRATION IN ON NAOH

IN ABSENCE OF LIGHT

 

27

ion and the hydroxyl radicals, and this would be the major
contributor to the degradation poocess.

No attempt was made to ascertain the effect of different
frequencies or to determine the quantum yield. Only the dif-

ference observable under the two sets of conditions were

noted and studied.

28

S USSIO

The kinetic studies undergone in the experimental work
indicate the following reactions to be the logical mechanism
by which potassium ferrate decomposes in strongly alkaline
media:

1) no," a- 2112034 I'eo2 e cano s- OH"

2) rec; + son _x:_,re02" .- 3/202 I» OH‘ + H20

:5) com ...—ksyO/éoz + o/znzo

The rate expression for the proposed reaction would be
developed as follows:

Placing: [3004f]: x and since

[DHT]constant, [OH-J k_1=k'1
dx - --k1x v x _1 [hoe ‘jfianfla- 3:: [33333 - 1:3[15333
313;; :1: - k _1 ['Ii'eOZij'OH']3 - Ixlejona3 - 3:33:03?
a) steady state approximation:
4 OH° ....O _k1x - x11 [reoz‘ Jfon-ja- xszOHJ5 - 1:3[OHCI5
at 03- 3 2 k1:
meoz'] I» :2: + :3

g;,: -2k1x kzx 9 k3

at k;1[p.oz;j v ‘2: + k5

Placing a : initial concentration of x and integrating

 

 

between the limits a x, this result is obtained:

k_1a e- 1} (k5 .- kza) xT-{kll ~13 - 1‘1ng e- kzagz -2k1t
k3 (1:3 I- kth i;— 13 1- k2:

 

 

M

I
..
I
' I ~ r. -‘
r _
I I
I -. 7 - .
l G
. _, e
9
t - ..-

29

The rate constants could not be determined individually; how-
ever, if the third reaction is neglected, the following dif-
ferential equation can be obtained:

A; -.- Zkikzxz
dt k; Fe 2 e 12;

If this is integrated between the limits a x, the fellow-
ing result is Obtained:

55.1. 1-23.34 lngg2k1t

kg x k x

As an be see The rate constants kll and k2 are present
as ratios in this expression. Since k1 is known from initial
slope of reaction curves, the ratio til/k2 can be determined
by taking the slope of the curve obtained by plotting F‘s04=
versus time and placing this slope equal to dx/dt. (lee
Figure XI) By using the differential equation and solving
for the ratios, k_1/k2, a value can be obtained.

If the suppositions made in the report are correct, placing
this value into the integrated expression, and plotting this ex-
pression, the resultant curve should correspond quite closely to
the observed curve.

2:1: 3.6x10'2

fig. : slope = 8.0 x 10“

x = .328
a = .620
[PeOé]= a - x = .292

til/k2 = 10.8

The result Obtained by performing this Operation can be seen

in Figure XII.

v-

30

 

O4@—- _a

 

 

035”“

 

 

 

 

03 I i I l l I I l
IO 20 ’30 4O ’50 60 7O

_77noe In ffllfl ”—

FIG. XI GRAPH SHOWING RATE OP DECOMPOSITION
OF POTASSIUM FERRATE IN 8N NAOH

31

 

 

 

 

I I I
09*— _i
oa__ _J
0.7L:- ._

Took: fl
I ”‘9
Qw- "b
a \. ...4
-\\ u ..IIC curve
0
I ‘0
IOAI— ° _
O L
‘0
obso c a Que o o
o x O _, \4
oce— _
0.2 ...-.. I I0 I I
0 I0_ ‘_"‘2II-Q_/;/77 40 50 6

 

FIG.X|I GRAPH “MSHOWING RATE OP DECOMPOSITION

OP POTA USIM FERRATE IN 8N NAOH
A5 ARED WITH CALCULATED

CURVCOE

 

32

Over the first segment of the graph, the calculated curve
follows the Observed curve closely. lfter this initial reac-
tion rate is passed, however, the calculated curve deviates
considerably, the slope being less than the observed curve.
When this deviation is passed, the two curves again agree
closely.

If the reactions proposed are accurate, this curve might
be expected. In the first portion of the curve where the
concentration of hydroxyl radicals is high, one would expect
reaction (3) to exhibit its effect most strongly. As the
concentration Of hydroxyl radicals decreases and the ferrite
concentration increases, one would expect this reaction to ex-
hibit a lesser effect and finally might become negligable.

It is probable that the mass kinetics so employed in
this work do not present the true picture of the over-all re-
action, ie., the following (or somewhat analogous) reactions
are probably present in the decomposition of Fe04':

1' no: 0» 2e- ~+reotz or FeOa'

2' PeO’z e e' -—d> Feoz‘

Since plus four iron seems to be unstable with respect to
decomposition into the plus three and plus six oxidation
states, it seems that the reverse reaction of 1' could be
immeasurably fast while the same would hold for the forward
reaction of 2'. Essentially, the observed reaction would in-

volve the plus three and plus six species.

53

This study uncovered nO facts substantiating this supposi-
tion; however, by utilizing these reactions lower moleculari-

ties are permitted, increasing the probability of a reaction
taking place.

were

34

QONCLUSION§

In this study Of potassium ferrate the following facts
observed:

1) Highly alkaline solutions Of potassium ferrate do
not appear to be applicable to polarographic studies.

2) Potassium ferrate solutions can be measured by col-
orimetrid methods; and at a wave length of SOQWAM, the
solutions obey Beer's Law over the concentration ranges
used in this work.

3) The kinetics of the reaction show the decomposition
of potassium ferrate in alkaline media to be complex.
The rate of this decomposition is definitely a function
of the concentration of hydroxyl ions.

4) The decomposition is greatly accelerated by light.
On the basis of the Observed data, the following mechanism

was proposed to explain the decomposition of potassium ferrate

in highly alkaline media at 30°C:

(1) FeO4' e 2320 f; PeOz" 1- con- .. on“
(2) F004. e SOH'I~—> F002- r 3/202 9 OH. 9 H20
(a) con. fia- 3/402 . 3/2320

The kinetic data Observed in the study seem to substan-

tiate these reactions. The data indicate that reaction (2)

is the main reaction taking place after the first equilibrium

reaction has been established.

’3

gPENDI;

TABLE 2
RATE or DECOMPOSITION or too IN 4N NaOH
AS FOUND BY POLAROGRAPHIC AMP neutrals TECHNIQUE
e 0°C (0.01213 KzreO‘IZSOO 4N Neon)*
Time In Min. Step Height Time In.Min. Stepcgeight

:lO Sec. Cm :10 Sec.

5 12.90 100 5.55
10 12.90 105 5.55
20 12.95 111 7.10
50 12.15 125 6.05
55 10.90 155 5.50
40 11.05 145 5.50
ea 10.85 155 5.42
50 9.700 155 5.15.
55 9.700 175 4.50
70 9.100 195 4.15
75 5.050 195 4.00.

200 4.55

‘ Data used for plotting Curve 1 in Fig. I.

TABLE 5

RATE OF DECOMPOSITION 0F Kggeafi IN 4N HaOH
0 T

A3 FOUND BY POLAROGRAPHIC AMP

e 0°C (0.00953 121‘504/2500 a NaOH)’
Step Height
Gm

Time In Min.
310 Sec.

15
2O
25
50
55
4O
50
60
70
95
100
110

* Data used in plotting Curve 2 in Fig. I.

7.20
7.10
7.00
6.65
5.85
5.65
5.60
5.45
4.95
4.55
5.15
5.50

Time In Min.
1:10 Sec.

120
150
140
150
160
175
180
190
200
210
220

RIC TECHNIQUE

4.90
5.20
5.00
4.65
5.00
4.40
4.05
5.25
4.10
5.65
5.10

56

Step Height
Gm

TABLE 4*

ABSORBTION DATA OF POTASSIUM FERRATE
IN 8N NaOH MADE WITH A BICKMAN SPECTROPHOTOMETER

Wave Slit 5 Tranl- Optical

Length Width,d, fiiasion Density

In.nw%/ In mm 1 D
520 0.602 15.8 .725
550 0.440 19.2 .614
540 0.545 25.8 .540
550 0.280 55.9 .447
560 0.218 46.2 .555
570 0.205 57.0 .244
330 0.178 65.5 .197
590 0.157 65.2 .186
400 0.140 65.8 ..190
410 0.125 60.5 .220
420 0.117 56.8 .246
450 0.104 52.9 .274
440 0.100 49.7 .504
450 0.094 47.0 .529
460 0.089 44.8 .550
470 0.084 42.7 .570
480 0.080 41.0 .587
490 0.075 40.0 .599
500 0.072 59.5 .402

510
520
550

550
560
580
600
620

“600
620

680
700
720
740
760
780
800

TABLE 4 (continued)

a
0.069
0.068
0.007
0.000
0.069
0.071
0.077
0.100
0.109
0.100
0.190
0.170
0.102
0.120
0.119
0.110
0.100
0.102
0.099
0.090

T
41.0
42.0
40.1
44.3
40.7
47.1
02.2
00.0
67.0
74.0
09.0
07.0
70.7
02.9
01.7
79.7
77.9
70.0
70.1
76.0

* Data need in plotting Gurvel in Fig. II.
'* Changed phototube

.007
.070
.000
.000
.040
.020
.290
.202
.174
.128
.220
.170
.121
.000
.007
.100
.100
.117
.110
.116

58

TABLE 5*

ABSORBTION DATA 0! ESSENTIALLI DECOMPOSED
POTASSIUI FERRATE SOLUTION-
TO SHOW EFFECT OF DECOMPOSITION PRODUCTS

Wave 0110 % Trane- 0191:1001

Length lidth,d, mission Density

Inrwafiy In In T D
500 1.975 75.5 .124
510 1.010 74.0 .105
520 0.645 82.9 .085
550 0.470 85.6 .068
540 0.570 87.5 .058
550 0.295 90.0 .046
560 0.247 92.7 .052
570 0.215 95.1 .021
580 0.186 96.5 .015
590 0.162 96.5 .015
400 0.145 96.5 .016
410 0.152 95.6 .019
420 0.120 95.0 .025
450 0.111 94.5 .025
440 0.104 94.5 .025
450 0.096 95.7 .027
460 0.090 95.4 .029
470 0.086 92.6 .055

480
490
500
510
520
540
560
580
600
625
.5625
650
675
700
725
750
775
800

TABLE 5 (continued)

6
0.081
0.077
0.074
0.071
0.069
0.069
0.072
0.080
0.101
0.150
0.176
0.145
0.150
0.118
0.110
0.104
0.099
0.095

T
92.5
92.1
92.2
92.5
92.9
94.0
94.2
95.0
95.5
97.1
97.7
98.7
99.0
99.0
98.7
98.4
98.2
98.7

* Data used in plotting Curve 2 in Fig. II.
*9 Changed Phototubc

D
.054
.055
.054
.055
.052
.027
.026
.022
.020
.012
.010
.006
.005
.005
.006
.007
.008
.006

40

41

TABLE 6*

DETERMINATION OF CONFORMITY 0F POTASSIUM FERRATE
SOLUTION IN 8N NaOH T0 BEER'S LAW

30803 g: lfxln d In §** D**
0714 n ”M7 m
0.075 500 .046 57.5 .241
0.109 500 .046 47.2 .555
0.154 500 .046 50.5 .498
0.186 500 .046 25.8 .596

* Data used in plotting Fig. III.

** Readings made after a five minute period had elapsed
from time of mixing.

42

TABLE 7*

KINETIC RUNS OF THE DECOMPOSITION 0F POTASSIUM FERRATE
IN 8N NaOH MADE 0N BECKMAN SPECTROPHOTOMETER AT 50°C.

Run #1 0. 016g 02. 05 12!.04/10000 an NaOH
s0.004 24- 000mg
t Min D T t D T
0.0 0.001 20.1 00.0 0.000 44.2
0.0 0.011 00.7 07.0 0.040 40.0
7.0 0.469 04.0 40.0 0.000 44.0
10.0 0.440 00.9 40.0 0.000 44.0
10.0 0.411 00.9 40.0 0.000 40.1
20.0 0.002 41.4 00.0 0.004 46.2
20.0 0.067 42.9 00.0 0.020 47.0
00.0 0.009 40.0 70.0 0.016 40.0
Run.#2 0.0154g .04 (90. 0%)[10000 00 NaOH
4:0.00 2~=0oom4
t 0 r t 0 T
4.0 0.400 00.1 00.0 0.291 01.1
0.0 0.410 00.0 40.0 0.200 02.1
10.0 0.079 41.0 40.0 0.277 00.0
10.0 0.040 40.0 00.0 0.271 00.0
22.0 0.021 47.0 00.0 0.204 04.0
20.0 0.000 00.1 00.0 0.260 00.0
00.0 0.004 49.0 00.0 0.209 00.1
70.0 0.201 06.1

* Data used in plotting Fig. IV.

TABLE 8*

KINETIC RUNS 0F POTASSIUM FERRATE DECOMPOSITION
MADE 0N BECKMAN SPECTROPHOTOMETER
WITH VARIING CONCENTRATIONS 0F NaOH

Run #1 0.0110g 62.5% 12F004/1000c 41! NaOH
t Min. 1 D t T D
2.0 00.0 0.246 15.0 04.9 0.071
0.0 65.8 0.197 17.0 07.0 0.001
7.5 70.7 0.151 -20.0 00.5 0.000
10.0 74.0 0.126 27.0 94.0 0.027
10.0 01.0 0.090 00.0 90.5 0.040
Run.#2 0.01003 02.0% KgFeO‘/10006 50 NaOH
t T D t 1 D
2.5 05.0 0.250 17.0 77.0 0.112
5.0 62.1 0.207 20.0 02.7 0.000
7.0 07.9 0.100 22.0 05.0 0.071
10.0 71.0 0.140 25.0 00.1 0.064
12.0 70.2 0.124 27.0 00.0 0.062
15.0 70.0 0.100 00.0 00.0 0.004
Run.#0 0.00903 02.0% léreo4/10000 6N Naflfi
t I D t T D
2.0 01.0 0.290 17.0 79.0 0.100
0.0 07.2 0.242 20.0 00.9 0.090
7.0 00.2 0.107 22.0 01.4 0.004
10.0 69.0 0.161 20.0 00.0 0.077
12.5 70.0 0.107 27.5 00.0 0.071
10.0 70.0 0.117

TABLE 8 (continued)

Run #4 0.01003 62.5% 520004/10000 7N NaOH
t Min. '1' D t T D
2.5 ' .07.9 0.200 20.0 01.0 0.092
0.0 60.0 0.197 22.5 01.2 0.005
.7.0 64.7 0.190 20.0 02.9 0.002
10.0 70.2 0.150 27.5 04.6 0.070
12.5 09.7 0.100 00.0 00.5 0.000
17.0 77.1 0.112 00.0 00.2 0.065

Run.#0 0. 01003 62. 00 Kano4d /10000 05 0400

050° C.,/s:50m7. 630.055m

t r t T D
2.5 49.0 0.006 17.0 02.5 0.204
5.0 02.1 0.204 20.0 00.0 0.100
7.5 04.9 0.201 22.5 05.0 0.100
10.0 00.0 0.207 25.0 06.2 0.100
12.0 00.0 0.221 27.5 66.0 0.170
10.0 61.2 0.210 00.0 07.5 0.171

Run #6 0.0105g 62.5% K2FeO4/10000 9N NaOH
t 7 0 ’ t 7 0
2.5 00.0 0.270 17.0 04.0 0.194
0.0 00.2 0.200 20.0 00.0 0.100
7.0 05.0 0.250 22.5. 64.1 0.190
10.0 09.0 0.229 20.0 67.8 0.170
12.0 00.9 0.210 27.5 60.2 0.100
15.0 02.0 0.200 00.0 60.7 0.102

Data used in plotting Fig. V.

TABLE 9*

KINETIC RUNS OF THE DECOMPOSITION 0F POTASSIUM FERRATE
IN 8N NaOH SATURATED WITH F020301320

d=0.054 kw: 500014

Min. T D t T D
0.0 24.0 0.014 00.0 20.1 0.501
0.0 24.0 0.600 00.0 20.4 0.040
0.0 20.4 0.097 40.0 20.0 0.541
10.0 20.0 0.094 40.0 29.0 0.001
10.0 26.0 0.000 00.0 29.9 0.520 ,
20.0 25.0 0.070 60.0 00.0 , 0.010
20.0 27.0 0.000 70.0 01.1 0.500

Data used for plotting Fig. VII.

45

TABLE 10*

KINETIC RUNS OF THE DECOMPOSITION 0F
POTASSIUM FERRATE IN 4N NaOH AT 0°C.

0.01203 02.0% 027.04/10000 4N 0400

d = 0.052 mm z~= 000 mrt/

t n10. T 0 t T 0
15.0 42.0 0.070 90.0 40.7 0.010
20.0 42.0 0.070 100.0 01.5 0.209
20.0 42.9 0.060 105.0 40.0 0.017
00.0 44.2 0.054 120.0 00.2 0.000
00.0 40.0 0.010 100.0 01.9 0.270
05.0 40.0 0.010 100.0 02.5 0.200
70.0 49.1 0.000 100.0 00.2 0.200
90.0 40.2 0.000

2 Data used for plotting Pig. VIII.

47

TABLE 11"

KINETIC RUNS OF THE DECOMPOSITION 0F POTASSIUM FERRATE
IN 8N NaOH WITH GLASS WOOL ADDED TO REACTION VESSEL

0.0157g 92.5% K29004/10000 8N NaOH

* Data used for plotting Fig. IX.

6. a 0.055 mm )0: 500 ”II/

t 210. T 5 t T n
0.0 20.0 0.000 00.0 41.0 0.007
0.0 20.0 0.000 40.0 42.9 0.002
7.5 00.9 0.012 01.0 44.0 0.052

10.0 02.0 0.490 60.0 40.0 0.040
10.0 00.1 0.400 70.0 40.2 0.000
20.0 07.0 0.420 00.0 47.1 0.026
20.0 09.4 0.400) 95.0 40.0 0.019

TABLE 12*

48

KINETIC RUNS OF THE DECOMPOSITION 0F POTASSIUM FERRITE

t Min.
2.5
5.0
7.5

10.0
20.0
55.0

IN 8N NaOH IN THE ABSENCE OF LIGHT

d 3 0.019 mm

T
25.0
27.5
27.5
28.0
29.7
50.5

D
0.600
0.565
0.565
0.555
0.558
0.517

,?~= 009,2«.7/

t T 0
50.0 01.0 0.009
60.0 01.0 0.000
00:0 02.0 0.490

100.0 02.0 0.409
120.0 00.0 0.400
140.0 00.0 0.400

* Data used for plotting Fig. X.

49

BIBLIOGRAPHY

Bohnson, V. L. I'Cs.‘l‘:a1ytic': Decomposition 0! Hydrogen
Peroxide,“ J. Phys. Chem.,‘§§, 19 (1921). ,

Bray, wm. C. and Gorin, M. H. “Ferryl Ion A Compound
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Grubs, G. and Gmelin, H. "The Electrolytic Formation
of the Alkali Salts of Ferrous Oxide and Ferric Oxide,“
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Kolthoff, I. M. and Lingane, J. J. Polagographi.
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Latimer, Wendell M.. The Ozigatiog States 3; the Element

ggg Their Potential: in Agugoug Solutions. New York:
Prentice-Hall, Inc., 1958. ,

Parks, 0. D, and Mellor, J. W. Mellor's Modern.lngt-
gagig'Chgmigtzz. New York: Longman, Green and 00.,
1959. . _ . . .

Schreyer, J. M., Thompson, G. W. and Ockerman, L. T.
“Oxidation of Chromium III with Potassium Ferrate VI,"
Anal. Chem., 22, 1426 (1950).

Thompson, G. W., Qckerman, L. T. and Schreyer, J. M.
”Preparation and Purification of Potassium Ferrate VI,“
J. Am. Chem. 500.. 25. 1079-00 (1901).

 

 

 

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