606
INDUSTRIAL A N D ENGINEERING CHEMISTRY
Vol. 16, No. 6
T h e Determination of Manganese’ P a r t 11-Manganous
Oxalate as a Primary Standard
By R. W. Coltman THEUNIONCARBIDE & CARBON RESEARCH LABORATORIES, INC.,LONOISLANDCITY,N. Y.
A
LTHOUGH in recent standard, such as sodium Manganous oxalate is proposed as a primary standard in conYears an increasing oxalate, is very evident.” nection with the determination of manganese by Volhard’s method. realization of the the chlorate method, and similar processes-its preparation, properThe use of sodium oxalate merits of the bismuthate is naturally inadmissible in ties, and uses being described in detail. Among other things, it is method has ~esultedin its the case of a method in shown that the pure salt is best prepared from recrystallized potasadoption in Preference to which empirical standardsium permanganate and oxalic acid and that its use introltres only a the Volhard, Ford-Williams, ization is necessary, a mansimple treatment of a weighed sample with nitric or sulfuric acid. and Von Knorre methods, ganese salt of theoretical these procedures are still composition being required. employed in many laboratories and probably will con- Although anhydrous chloride or bromide of manganese can be tinue to be for many years to come. The greatest objec- prepared in a pure state,4 such complex apparatus and diffition to these procedures, all of which are based upon the pre- cult manipulations are required that the use of either is cipitation of manganese dioxide, lies in the fact that the precip- practically excluded. itate always contains less than the theoretical amount of oxyOf the many compounds of manganese, manganous oxalate, gen, necessitating the use of various empirical factors. The MnCz04.2Hz0, possesses in a high degree the requirements fact that no manganese compound having the requisite of a standard salt of manganese, and is proposed as a primary properties for use as a primary standard has heretofore been and working standard.6 I n spite of its water of hydration, available has led to much uncertainty as to the absolute this compound is neither hygroscopic nor efflorescent under accuracy of the results obtained. Such errors, relatively ordinary conditions, and as it possesses many interesting unimportant in the case of materials containing small amounts properties, most of which are not so well known as those of of manganese, are of considerable moment where products other manganese compounds, its preparation, properties, and having high percentages of this element are concerned. use as a standard will be described in some detail. These considerations make obvious the utility of a comNATURE OF MANGANOUS OXALATE pound suitable for use as a primary manganese standard. Manganous oxalate is formed by treating a solution of B PROPERTIES REQUIRED IN A PRIMARY STANDARD McBride2 has summarized as follows the conditions a sub- manganous salt with an alkali oxalate or with oxalic acid. It forms a dihydrate, MnCz04.2Hz0, and a trihydrate, stance must meet in serving as a primary and a working MnCz04.3Hz0. According to Graham,6 Liebig,’ and Souchay standard: and Lenssen,* manganous oxalate contains 2.5 molecules of ( a ) There must be reasonable ease of preparation and accurate water. Gorgeus showed, however, that the compounds reproducibility. (b) The purity must be determinable with sufficient accuracy, previously described had been mixtures, and stated that a and the Durified material must be stable under ordinarv condi- trihydrate precipitates from cold solutions and a dihydrate tions of {he laboratory. from hot solutions. Gorgeu’s work has been more recently (c) The use of the material in regular work must demand confirmed by Hauser and Wirth.lo neither complex apparatus nor difficult manipulation. This error in the formula for manganous oxalate has per( d ) Such precision must be obtainable that when it is used with ordinary care one determination (or a t the most a very sisted up to the present, the formula being given as MnC204.few) suffices for the fixing of the value of a standard solution. 2.5 HzO in several reference books. (e) The accuracy obtained under ordinary conditions of its TRIHYDRATE, MnCz04.3Hz0 use in standardization must be at least as great as that required in the use of the solution to be standardized. The trihydrate is formed when a cold solution of a mangaCOMPOUNDS PREVIOUSLY PROPOSED nous salt is added to a soluble oxalate, or when manganous Few compounds of manganese can be prepared of theo- oxalate crystallizes from a cold solution. A pink precipitate retical composition without undue labor. Most of the is obtained, which under the microscope is seen to consist of salts containing water of crystallization are out of the ques- long prismatic needles. (Fig. 1) Under these conditions the tion, owing to their varying degrees of hydration and result- crystals do not always form twins, and the six-armed, staring instability. Although potassium permanganate can be shaped clusters, obtained in microchemical tests, are not purified to a high degree, it always contains a little water, present in any great quantity. The trihydrate crystals while anhydrous manganous sulfate, which has frequently polarize strongly, extinguish parallel to their length, and been proposed as a standard manganese salt, can be prepared exhibit brilliant polarization colors. At ordinary temperain a pure state only under very carefully regulated conditions. tures they change very slowly to the dihydrate, while a t Blum3 has studied the determination of manganese as elevated temperatures the transformation is exceedingly sulfate, and finds that manganous sulfate demands long4 Baxter and Hines, J . A m . Chem. SOC.,28, 1560 (1906). continued heating a t a definite temperature (480’ C.) in order 6 Rtist, Z . anal. Chem., 41, 606 (1902), proposed this salt for an oxidito free it from water, If this temperature is much exceeded metric standard, but evidently did not consider it as a standard for mangadecomposition takes place. He remarked that “the desira- nese. 6 Ann., 29, 2 (1839). bility of substituting for the manganous sulfate some other 71bid., 96, 116 (1855). 16, 58 (1924). Received October 18, 1923. Part I , THISJOURNAL, J . A m . Chem. SOC.,34, 393 (1912). 8 Ibid., 34, 1379 (1912). 1
a Ibid., 102, 47 (1862). Comfit. rend., 47, 629, 920 (1858). 10
J . fivakt. Chem., 79 (2), 358 (1909).
IiVDUSTRIAL A N D ENGIHEERING CHEMISTRY
June, 1924
rapid, 3 minutes' heating a t 80" C. being sufficientto convert the trihydrate to the dihydrate. Owing to this instability and the fact that the dihydrate is monotropic, there is no dnnger of it being contaminated by trihydratc if the solutions are kept hot during the preparation of the standard salt.
607
sium prsrmanganate was reduced with an excess of recrystallized oxalic acid in aqueous solution, and the manganous oxalate, which was formed according to the equation 2KMn04I- 8 H G 0 4 = ZMnCaO,
+ K&O, + 8H10ilocol
was filtered, washed thoroughly by decantation, and then brought onto a Biichner funnel. The precipitate was sucked as dry as possible, and then spread out on filter paper and allowed to dry in the air. One lot made in this manner was designated as Lot 1 4 . (Table I) Tbe dihydrate was also prepared according to the following procedure: Recrystallized potassium permanganate was reduced in the presence of sulfuric acid with slightly less than the theoret,ical quantity of recrystallized oxalic acid. The hydrated dioxide formed was filtered, the solution of manganous and potassium sulfate treated with ammonium carbonate, and the precipitate washed by decantation until practically free from sulfate. The manganous carbonate thus obtained was added little by little to a hot solution of o d i c acid in excess. The precipitate of manganous oxalate was washed free from acid by decantation with cold water, and then as much water as possible was removed by suction after bringing the precipitate onto a Buchner funnel. The precipitate was airdried as usual. This product was designated as Lot 2-A. A similar procedure was used in the precipitation of Lot 1-W, except that a solution of sulfur dioxide was used to reduce the permanganate. Precipitation as carhonate and conversion to oxalate by treatment with oxalic acid followed.
TESTSFOR PUXITY Fro. I-MnCzO.
3H:O ( X 110)
DIHYDRATE, MnC20,.2HZ0 The dihydrate is formed when hot solutions of an oxalate and a manganous salt are employed, or, as mentioned above, by the spontaneous transformation of the trihydrate. It may also be prepared by adding manganous carbonate to a boiling solution of oxalic acid, or by the interaction of potassium permanganate and oxalic acid. The crystals usually appear as small, rhombohedral plates, optically active. (Fig. 2) However, the product obtained from hot solutions of oxalic acid and potassium permanganate shows many crystals in the form of long prisms, and the small plates are more perfect in form. In contrast to the pink color of the trihydrate, the dihydrate is pure white. The salt is exceedingly stable in air at room temperature, but dehydration and slight oxidation, as shown by the brown tinge that the crystals assume, commence a t temperatures slightly higher than 100" C. The pure anhydrous salt, obtained by heating the oxalate in nitrogen a t 120' C., possesses the same beautiful pink color characteristic of the anhydrous sulfate or chloride. Manganous oxalate is the most soluble of the omlates, the alkali metal oxalates excepted. According to Hauser and Wirth,'O 1 liter of water a t 25" C. dissolves 0.3120 gram of MnCeO&€*O. The solubility is increased by the presence of oxalic acid or ammonium oxalate. In sulfuric and nitric acid solution the solubility is increased still further, and hot, concentrated sulfuric or nitric acid decomposes the salt completely. On ignition, carbon monoxide and dioxide are evolvod, a residue of manganomanganic oxide, Mna04,being left if the heating is performed in air, or a residue of manganous oxide, MnO, if the ignition takes place in an atmosphere of nitrogen. PREPARATION OF TRE PUREDIHYDRATE Of the various possible methods of preparation, the following w&s found to he the simplest: Recrystallized potas-
METALSNOT PRECIPITATED BY AUMONIUM SIJLFIDE-TC~ grams of the sample were decomposed with nitric acid in a platinum dish, evaporated to fumes with sulfuric acid, and the resulting solution precipitated with ammonia and ammonium sulfide, a large excess of the former being used in order to obtain the green modification of manganese sulfide, which is much less bulky than the pink variety.
Fro. Z--MnC*O..2€LO
( X 110)
The precipitate was filtered and allowed to drain, the filtrate evaporated in platinum, and the residue ignited. A small amount of manganese escaping the first precipitation was thrown out from the sulfuric acid solution of the rcsidue with ammonium sulfide, and the filtrate from it again evapo-
INDUSTRIAL A N D ENGVNEERING CHEMISTRY
608
rated and ignited, A barely perceptible nonvolatile residue was obtained. HEAVY METALS-Fifteen grams of the sample were decomposed with nitric acid, the solution made ammoniacal, and hydrogen sulfide led in. The solution was then acidified with acetic acid. A small amount of pink manganese sulfide, which remained undissolved when the solution was acidified with acetic acid, was filtered off and dissolved in acid, the precipitation with hydrogen sulfide in ammoniacal solution being repeated. On the addition of acetic acid, the precipitate dissolved completely (absence of zinc), although owing to traces of the heavy metals the solution was not entirely colorless. TESTSFOR COMPOSITION The purity of the salt having been established, the composition was determined by heating it in nitrogen a t 120" C. and obtaining the value for the water of crystallization by the loss in weight. The amount of manganese was determined by converting weighed amounts of the oxalate to sulfate, the latter being dehydrated a t about 500' C. in an electric muffle furnace. The percentage of manganese was calculated from the weight of sulfate obtained by using the theoretical factor 0.3638.
Weight Sample Grams 3.2005 3.1807 3.2403 Weight Sample Grams 2,3099 2.2996 3.3102
TABLEI-ANALYSIS OF MnCz04.2HzO Lot I - A Water of Cvystulliration Loss Grams 0.6436 0.6403 0.6521
Weight MnSOI Grams 1.9493 1.9393 2.7923
Per cent Per cent Theory 20.11 ... 20.13 20.12 26: i 3 Total Manganese Weight Manganese Per cent Per cent Grams Manganese Theory 0.7091 30.70 0.7055 30.68 1.0135 30.69 36 69
...
:
, The figures in Table I show that the salt is the pure dihy-
drate. Assuming freedom from contaminating metals, the water content is naturally most important, and any yariation from lot to lot would preclude the use of the salt as standard. The values for water of crystallization obtained on other lots are given in Table 11. TABLE11-WATER Weight Sample Grams 3.7099 3.7930 3,4439 2.8348 3.3976 2.8648 3.6933 2.7790 3.0553 3,1987
*
OF
CRYSTALLIZATION I N MnCzOa 2 H 0 Lot 2 - A
Loss Grams 0.7468 0.7630 Lot T 0,6935 0.5708 0.6845 0,5770 Lot w 0,7434 0.5588 0.6150 0.6435
Per cent 20.13 20.12 20.14 20.14 20.15 20.14 20.13 20.11 20.13 20.12
The stability of the salt a t ordinary temperature upon exposure to atmospheres of varying humidity is shown by the data in Table 111. In these experiments the salt was weighed into flat-bottomed weighing bottles and exposed in desiccators containing sulfuric acid of the concentrations necessary to give the humidities desired. The oxalate had been kept for several months in a glass-stoppered jar, which, however, was not airtight. It will be noted that the variation in weight, which is definite only in the case of the greatest humidity, is no greater than that which any powdered material would undergo. The maximum variation (+0.018 per cent) is less than one
Vol. 16, No. 6
part in 5000, an amount decidedly less than the experimental errors involved in the methods used for manganese. TABLE 111-HUMIDITY EXPERIMENTS Relative Humidity Per cent
Weight Sample Grams 11.7577 11.8591 10.1631 10.1522
10
35 65 90
-Change 1 dav -0.0602 - 0 0008 -0.0006 10.0007
in Weight, Grams3 davs +0.0605 0.0000 10.0002 +0.0013
7 davs +0.0606 -0.0008 -0.0007 +0.0018
OTHERMETHODS OF PRODUCING THE DIHYDRATE The addition of sodium oxalate to a hot solution of manganous chloride or sulfate gives a precipitate of the dihydrate, which can readily be washed free from sodium. When ammonium oxalate is used as a precipitant, the manganous oxalate always contains a trace of ammonia. Precipitation with sodium oxalate is a very convenient method of preparing the salt, but the difficulty of obtaining manganese chloride or sulfate in a pure state is very great, as even the "chemically pure" salts contain small amounts of calcium, zinc, cobalt, etc. The use of potassium permanganate as a starting msterial for the preparation of manganous oxalate has the advantage that such metals are not likely to be present, and one recrystallization of the permanganate is generally sufficient to obtain a salt of satisfactory purity. The impossibility of preparing a pure oxalate from manganous solutions containing other metals-with the exception of the alkalies-is evident from Table IV. TABLE IV-SOLUBILITYOF METALLIC OXALATES Temperature Grams COMPOUNDS C. Mols per Liter per Liter AUTHORITY NanCz06 15.5 32.2O Souchay and Lensennl' 252.4" Engel12 KzC204 0 0.0055 Kohlrausch and CaCzC4.HzO 18 0.043'210-3 Rose13 SrCz04.2.RHzO 18 0 . 2 6 X 10-8 0.046 Kohlrausch and
...
Rose14
BaCrOc2HzO 18 MgC104.2HzO 18 ZnC?Oa 25 MnCz04.2HzO 25
... x lo-* x 10-5
2.7 7.7 0.218 X 10-2
0.089 0.303 0.0118 0.312
PbCzOi
0 . 5 4 X 10-8
0,0016 0.033 0.0253
CdCnOn
CudaO4 a
18 18 25
1.56
'x10-4
Per 1000 gram s water.
Kunschert" Hauser and Wirth'o Kohlrauschls Kohlrausch'o Abegg and Schaeferz'J
Figures for the solubility of nickel, cobalt, and ferrous oxalates were not obtainable, but preparations made with manganous solutions to which salts of these metals had been added showed that they were precipitated in large part with the manganous oxalate.
PROCEDURE FOR PREPARATION AND USE OF
THE
DIHYDRATE
Five hundred grams of oxalic acid, H&zO4.2HzO, are dissolved in 1 liter of hot water with the addition of 5 cc. of concentrated hydrochloric acid. The solution is filtered through hard paper to remove dust, etc., and the filtrate is allowed to crystallize in the cold, The acid thus obtained is brought onto a Bdchner funnel, sucked as dry as possible, and washed twice with a little ice-cold distilled water. The crystals are spread out on a large filter and allowed to dry a t room temperature in a place free from dust. Two hundred and fifty grams of C. P. potassium permanganate are treated with 800 cc. of hot distilled water in a large Ann., 99, 33 (1886). Bull. SOC. chrm., 46, 318 (1886). 15114 2. physik. Chem., 12, 234 (1893). 16 Ber., 34, 3513 (1901). 16 Z . physik. Cheni., 60, 356 (1904). 17 Compl rend., 116, 939, 1028 (1892). 18 2. physzk. Chem., 64, 159 (1908). 10 Ibid., 44, 107 (1903). 20 Z . anoyg. Chem., 45, 293 (1903). 11
1%
June, 1924
INDUSTRIAL AND ENGIhTEERING CHEMISTRY
beaker. The liquid is heated to effect solution of the crystals as far as possible, and then filtered through purified asbestos. The filtrate is allowed to cool over night in a porcelain evaporating dish, and the crystals are separated from the mother liquor t)y filtering through asbestos. They are then dried a t room tomperature for 24 hours and a t 100’ C. for 2 hours. The oxalate is made from these purified compounds in the following manner: Two hundred and fifty grams of the oxalic acid are dissolved in about 500 cc. of hot distilled water in a large beaker, and a solution of 70 grams of the permanganate is added slowly and cautiously, the oxalic acid solution being kept near the boiling point. The permanganate color is quickly destroyed after each addition, and finally manganous oxalate begins to settle out. Towards the end the permanganate must be added very slowly. When all the permanganate has been added and the manganous oxalate has settled, the supernatant liquid is decanted without cooling the solution. The manganous oxalate is washed by decantation with cold water until free from the excess of oxalic acid, brought) onto a Biichner funnel, and sucked as dry as possible. The crystals, which should be pure white without the least
609
pink tint, are spread out on filter paper to dry in the air. They may be further dried a t room temperature in a current of dry air. The salt so obtained is bottled, and will keep indefinitely under ordinary conditions. If the oxalic acid solution is allowed to become too cold some trihydrate will form. I n this case the manganese oxalate, after having been filtered, is heated with water to about 90’ C. until no pink tinge is perceptible. The crystals are then again brought onto a filter and washed and dried as outlined. The use of the salt in practice is exceedingly simple. The desired quantity is weighed on a watch glass (1 gram of MnCz04.2Hz0 is equal to 0.3069 gram Mn) and transferred to a beaker or flask. Evaporation to fumes with sulfuric acid will form manganous sulfate for use with Volhard’s method, or treatment with concentrated nitric acid will give a solution of manganous nitrate for the chlorate or bismuthate method. ACKNOWLEDGMENT The author desires to express his thanks to J. A. Holladay for assistance in the preparation of the manuscript, and to T. W. B. Welsh for the photomicrographs.
Starch in Sorghum Sirup’” By J. J. Willaman and F. R. Davison DIVISION O F AGRICULTURAL BIOCHEMISTRY, UNIVSRSITY
ORGHUM sirup is typically rather viscous, and OCcasionally the viscosity is so great as to cause a semisolid consistency of the sirup. I n 1920 the writers received from several scattered sources samples of sirup that were jelly-like in consistency, like a soft and stringy fruit jelly. This was so marked, and the output of one Minnesota sirup factory during this season was so consistently of this character, that it was decided to investigate the cause of the condition. The writers have been negligent in publishing the results, however, and meanwhile Sherwood has published two papers on the same ~ u b j e c t . ~Since the writers used methods in their work somewhat different from those used by Sherwood, it was thought desirable to publish their results as well. The first suggestion was that starch was the cause of the jellying, and Sherwood found this to be the case. He has reviewed the previous work ’on the occurrence of starch in sorghum juice. He found by analysis from 0.14 to 0.85 per cent of starch in various samples of sorghum juice, with an average of 0.33 per cent. By treatment of the heated juice with malt diastase he could remove the starch entirely, and the resultant sirups were free from sliminess and gelatinous properties. Added starch could be circumvented in the same way. He concluded that starch was responsible for the jelly-like nature of some samples of sorghum sirup.
S
MATERIAL AND METHODS Several samples of jellied sorghum sirup, from various sources and made from unknown varieties, were obtained, as well as some samples made from Minnesota Early Amber grown in M i n n e ~ o t a . ~ 1 Presented under the title “The Gums of Sorghum Sirup” before the Division of Sugar Chemistry a t the 66th Meeting of the American Chemical Society, Milwaukee, Wis., September 10 to 14,1923. 2 Published with the approval of the Director, as Paper 419, Journal Series, Minnesota Agricultural Experiment Station. * THISJOURNAL, 15, 727, 780 (1923). 4 Thanks are due to the Waconia Sorghum Mills, Inc., Waconia, Minn., for many large samples of sirup.
OF
MINNESOTA, ST.
PAUL,
MINX.
Analysis for “crude gum” was made by diluting the sirup with five volumes of water, precipitating with two volumes of 95 per cent alcohol, filtering, drying, and weighing the precipitate. Starch was determined by analyzing the crude gum, with or without drying, by the diastase method. Dry matter in the sirup was determined by an Abbe refractometer. EXPERIMENTAL RESULTS Table I gives the analyses of twelve samples of sirup for total solids, crude gum, and starch. One sample of cornstalk sirup is included for comparison, since this sirup contains no starch. It will be noted that the jellying property is not dependent on the density of the sirup. I n general, the jellying is proportional to the amount of crude gum present, although the last column shows that the percentage of starch in the crude gum is of more importance than the total amount of gum. Furthermore, the data show that any sirup having 1.3 per cent or more starch is likely to be more or less jellylike. T A R II-ANALYSGS ,~ OF SORGHUM A N D CORNSTALK SIRUPS Solids
7 .-” 84.0
3.10 0.80 6.92
1.17 0.32
37.8 40.7
1.39 4.95
0.0
0.0
..
4.30
0.70
16.3
.. .. ..
4.30
0.80
18.6
3.54
1.33
37.8
3.56
1.90
53.4
4.23
0.79
18.7
4.01
1.18
28.7
.
8 9 10
11
Normal, hiinn., September 11, 1921 Moderately jelly-like, Minn,, Septemher 12 1921 Moderately jelly!-li ke , Minn ., October 10 1921 Nqcmal,_. I d n n . , October 11, l t l i s l
12
Normal,
Minn.,
12, 1921
October
Starch Starch in in Crude Sirup Gum
7 ,0 3.36
SamDle DESCRIPTION 7 .”n 1 Very jelly-like Minn. 1920 8 2 . 5 2 Moderately jeliy-like, Ihinn., 1m n 77.6 3 N;r-&k, Minn., 1920 80.7 4 Very jelly-like, Ohio, 1920 . 5 Cornstalk sirup 1921 and 1922 (average; 75.0 6 Very jelly-like, Wis., 1920 68.0 Normal, Minn., September 7
10 1921
Crude Giim
.. ..
41,
4.00
.. ..
.. ..
