Stability of the Ferrate(V1) Ion in Aqueous Solution J. M. SCHREYER
'
L. T. OCKERMAN', University of K e n t u c k y , Lexington, K y .
AND
calibrated in this manner (technical data supplied with electrode). The solutions were agitated continuously and a constant temperature of 26" & 0.5" C. was maintained during the course of the experiment. The concentrations investigated covered a range from 0.020 to 0.086 molal.
E C E N T publications (6,6) have reported the use of the ferrate ion as a new oxidizing agent in alkaline solution. In connection with the development of quantitative oxidations involving the ferrate ion, an investigation of the factors influencing its stability was necessary. When potassium ferrate is added to water, oxygen is evolved and hydrous ferric oxide is precipitated. The resulting solution is strongly alkaline. The decomposition reaction has been reported to be as follows ( 4 ) : 4FeOa--
Results are shown in Figure 1. Readings were made a t 1-fiinUte intervals, but many observations were omitted from the figure for brevity. These data are reproducible only if eonstant conditions are maintained. Changes of conditions were found not to change the general shape of the curves, but only the time required to reach the plateaus of constant pH and highest pH. The slopes of the pH-time curves were observed to be a function of the concentration. The curves for concentrations less than 0.030 molal contained plateaus of relatively constant pH before the attainment of the highest pH. The pH of the solutions containing concentrations greater than 0.030 molal increased rapidly and no plateaus were observed prior to the highest pH. I t is apparent that a 0.030 molal solution represents tt critical concentration, as evidenced by the nonieproducibility shown in curves 111and IT'. At the end of 60 minutes the solutions of 0.020 and 0.025 molal concentrations still contained approximately 89% of the original ferrate ion when analyzed by the chromite method. Analysis of the entire solution at the highest p H in the case of the 0.030 to 0.086 molal concentrations revealed that complete decomposition had occurred. The observations of pH on the dilute concentrations were not continued throughout the time required for complete decomposition, but it is probable that these curves would likewise show a second rapid rise in p H prior to the attainment of the highest pH, as is shown in Figure 2 on a solution of approximately the same concentration. These plateaus of relatively constant p H during the decomposition of dilute solutions of potassium ferrate are unexplained a t present. In agreement with Moser ( 2 ) ,these data revealed that the more dilute solutions of the ferrate ion are more stable.
+ 4H20 +2Fe20a + 8 0 H - + 30,
Rloser ( 2 ) reported that dilute solutions of the ferrate ion were more stable than concentrated solutions. The addition of potassium hydroxide, sodium hydroxide, potassium chloride, potassium bromide, potassium nitrate, potassium carbonate, potassium chlorate, or sodium chlorate was said to retard the decomposition of the ferrate ion. Salts of calcium, strontium, and magnesium, metals and their oxides, peroxides, and organic materials Lvere said to accelerate the decomposition. Rose ( 3 )stated that the addition of potassium chloride, potassium sulfate, sodium carbonate, sodium nitrate, potassium nitrate, or sodium tetraborate retarded the decomposition reaction, while sodium chloride appeared to accelerate the decomposition. The conclusions of these investigators were based upon the results of nonquantitative experiments. In addition, both the available solid samples of potassium ferrate and solutions of this compound were contaminated to such an extent that such specific effects of additive salts seemed questionable. The availability of both highly purified samples of the soluble salt, potassium ferrate (7), and suitable methods of analysis ( 5 , 6 ) made possible a quantitative investigation of the stability of the ferrate ion in aqueous solutions. INFLUENCE OF CONCEhTRATION ON STABILITY
In order to test Moser's report that dilute solutions of the ferrate ion were more stable than concentrated solutions, it was decided to study the rates of decompesition by determining the changes in hydroxyl ion activity as the decomposition proceeds in solutions of differ'- I e n t c o n c e n t r a t i o n of the ferrate ion. Weighed samples of potassium ferrate of known purity, a6 analyzed hy the arsenite (5) and chromite (6) methods, w e r e p l a c e d in a 50-ml. beaker. A weighed amount of conductivity water, at a temperatureof26" f 0.5" C., was added to the beaker containing the potassium ferrate. Zero time was arbitrarily assumed to be the time of contact between liquid and sample. Periodic observations of pH were made using the Beckman Model G meter with a Beckman 1190-E high pH electrode. The instrument was set a t 9.5 in a pH 10.00 + 0.05 buffer and a correction of 1-0.5 was applied to all esperimental r e a d i n g s . T h e m a n u f a c t u r e r states that this electrode is useful up to a pH of 13.5 when
' Deceased
April 11. 1950.
oa
les
I L L
CurveX 0.Q25 CurveIE0.030 CurvelZ0.030 C u r v e P 0.040 C urvoH0.086
12.0 12.0
MOLAL MOLAL MOLAL MOLAL
MOLAL
11,s 0
fi
11.0 1
w
1
I
.
I
I IO
0
I
I
~
I 20
I
I
I
30
1
1
1
40
I
I
I
x1
1
1
1
'so
TIME (MIN.1
Figure 1.
Effect of Concentration of Potassium Ferrate u p o n Decomposition i n Aqueous Solutions 1312
0.030
0.025 Molar H,l=eO.,. P r e p a r e d /rom 98.41 Per Cent Pure Sample. Temperature 2 6'C.
126124
*
12 2
0-0-
120
-
118
-
11.6
-
114
.
P"
---+--.-.-*
0-.
-\
0 .
0 .
- 0.024 L W
-0.OILi
p H Vs. Time Concentration in Moles/Liter Vs Time -0012
ss"
0~o~-o-o-o.o~ooo-o~-o
112
-z
o ~ o ~ o ~ o . o - o . ~ ~
-
p'o-bo-~ 0-0
0.006
composition occurred after 50 minutes. The curve of pH us. time showed two distinct areas of relat,ively constant pH. The initial rapid rise in p H and the rapid rise after 50 minutes correspond to a rapid decrease in ferrate in solution. EFFECT O F ADDED IMPURITIES UPON STABILITY
In order to study the effect of added iinpurities on the decomposition of potassium ferrate, solutions of 0.040 M potassium ferrate were used, as the use of more dilute solutions would involve long periods of time as evidenced by
The marked instability of the fwrate ion in the presence of hydrous ferric oxide cxplains the rapid decomposition of ferrate solutions after quantities of hydrous ferric oxide are produced by the decomposition reaction. STABILITY OF FERRATE(V1) IOY IN BUFFERED SOLUTIONS
Figure 3.
Stability of Potassium Ferrate in Aqueous Solution in Presence of Added Impurities
Preliminary investigation of the stability of the ferrate ion in the p H 7 boric acidborax buffer recommmded by Palitzsch (1) showed marked instability. The p H 7 phosphate mixture of Sgrenson (1) showed very little s t a b i l i z i n g effect on the ferrate ion. Clark and Lubs's phosphate buffers ( 1 ) of p H '7 and 8 showed marked retardation of the decomposition of the ferrate ion. Clark and Lubs's buffers of pH 7 and 8 were prepared. Solutions of 0.050 X potasbium ferrate were prepared
'
1314
0.05
0: 0.04
w
t J
?z‘II 0.03 0 2
9Ls 0.02
Y
ANALYTICAL CHEMISTRY
.zz - 10.60 - 10.50
4e
5
potassium ferrate deconipose- in phosphate buffer solutions
10.40
0.050MOLAR K.F.0. TEMPERATJRE-26 C 0 CWNG€SINpHlNpH 7.0BUFFER (1) 0 CnANOESINCONC.INpH 7.0BUFFERtX) CHANGES IN pH IN pH 8.OBWFER 0 CHbNGES IN CONC.IN pH 6.0 BUFFER (IY)
0.01
0.00 c
Figure t . pII and Concentration cs. Time
- 10.70
I
I
I
1
using these buffer solutions and simultaneous determinations of pH and concentration of potassium ferrate us. time were made using the procedure previously dePcribed (Figure 4). The solutions of potassium ferrate a8 shown by these data increased in pH, and showed two plateaus of constant pH. The initial decomposition as shown by the concentration-time curves was rapid, attaining a relatively conqtant concentration corresponding to the area of constant pH. The potassium ferrate solution prepared from the pH 8 buffer R a5 more stable than the solution prepared from the pH 7 buffer. The solution prepared from the p H 7 buffer solution contained 49yo of the original potassium ferrate after 8 hours; the solution prepared from the pH 8 buffer solution contained 71.4% of the original potassium ferrate after 10 hours. As these solution^ did not remain buffered a t p H 7 and 8, and no btabilizing effect was observed until a high pH was attained, it appears that the major factor influencing stability of the ferrate ion in aqueous solution is a high alkalinity. It is believed that the presence of the phosphate ion has some stabilizing effect. CONCLUSIONS
The initial concentration of potassium ferrate in aqueous solution exerts a marked influence upon the decomposition of the ferrate ion in solution. The more dilute solutions of the ferrate ion are more stable.
1
- 1030
- 10.20 -1010 1
The addition of potassium chloride or potassium nitrate as an impurity in solutions containing the ferrate ion accelerated the initial decomposition of the ferrate ion, but had the effect of stabilizing a small quantity of ferrate ions. Sodium chloride and hydrous ferric oxide as impurities caused complete decomposition of the ferrate ion in solutions a t a rapid rate. Solutions of the ferrate ion in Clark and Lubs’s buffers of pH 7 and p H 8 attained a relatively constant concentration after 10 minutes. The solution prepared from the p H 8 buffer was more stable than the solution prepared from the pH 7 buffer. LITERATURE CITED
Clark, IT. )I., “Determination of Hydrogen Ions,” p. 106, Baltimore, Williams B: Xilkins Co., 1927. hfoser, L., J . prakt. Chenz., (2) 56, 425 (1897). Rose. H., Pogg. Ann., 59, 315-24 (1843). Schreyer, J. AI., thesis, ‘‘Higher Valence Compounds of Iron,” 01egon State College, Corvallis, Ore., 1948. Schreyer, J. M., Ockerman, L. T., and Thompson, G. IT.,ANAL. CHEM..22. 691 119.50). Schreyer, J. hl., Thompson, G. IT.,and Ockerman, L. T., ~ N A L . CHEW, 22, 1426 (1950). Thompson, G. IT.,Ockerman, L. T., and Schreyer, J. hl., J . Am. Chem. Soc., 73, 1379 (1951). RECEIVED October 26, 1950.
Presented in abstract before the Diviaion of Physical and Inorganic Chemijtry a t the 118th Meeting of the AMERICAS C H E M I C ASOCIETI-, L Chicago, Ill.
Separation of Cotton and Rayon or Cottonand Acetate for Analytical Purposes-Addendum
1 directions
N RESPONSE
to numerous requests for amplification of the for preparation of the zincate solution [Heim, cHEN., 22, 360 (1950) the following details are
Oskar, provided : In making the stock solution containing 2o parts by weight of sodium hydroxide, of zinc oxide, and of water, the oxide is slurried with 20 parts of water, and all the sodium hydroxide, preferably the technical flakes, is added stirring until dissolved. The heat evolved by these proportions, especially when several pounds are prepared, is sufficient to effect solution.
Pouring a few drops of the milky solution into a tube holding a tenfold volume of water determines conipleteness of solution. Scanning the bottom of the reaction vessel, preferably of stainless steel, through a flat-bottomed tube (Nessler, etc.) conveniently reveals any undissolved zinc oxide agglomerations, which will then go into solution upon either more stirring or heating; heating is rarely necwsary. The remaining 31 parts of water are added finally. OSKARHEnf 5943-48th Ave. Woodside, N. Y .
