340
T H E J O U R N A L OF I N D U S T R I A L A N D E N G I N E E R I N G C H E M I S T R Y
The best blocks were given b y IO,OOO Ibs. and this was adopted as the standard pressure for preparing pellets for boosters. CONCLUSIONS
From the work which has been outlined i t appears t h a t the following conclusions were justified: ( I ) Hexil as a booster material is superior t o T N T , but somewhat inferior t o tetryl and TNA. ( 2 ) It is extremely stable and is safer t o handle t h a n either tetryl or TNA, and makes a satisfactory booster. (3) It can be manufactured by a simple process from sources which would not conflict with T N T manufacture, and because of the excellent yields obtained, and the cheapness of intermediates for it, its material cost should be less than for either tetryl or TNA. (4) On account of the simplicity of the process, the installation of a plant for its manufacture would be less expensive t h a n for an extension of t h e manufacturing facilities for either tetryl or TNA. ( 5 ) On account of the simplicity of operating methods, labor costs would be less than for either of the other materials. ACIDITY AND ACIDIMETRY OF SOILS.1 I-STUDIES OF THE HOPKINS AND PETTIT METHOD FOR DETERMINING SOIL ACIDITY AGRICULTURAL AND MECHANICAL COLLEGE, STI&LWATER, OKLAHOMA Received October 14, 1919
The Hopkins and Pettit method of determining soil acidity2 proposed in 1902 is essentially as follows: I O O g. of soil are shaken in a bottle of 400 cc. capacity with 2 5 0 cc. of 5 per cent commercial common salt solution for 3 hrs. 1 2 5 cc. of the clear liquid are taken off, boiled to expel carbon dioxide, and titrated using phenolphthalein as an indicator. The results are multiplied by 3 as a factor t o determine the total amount of base required. Later3 a normal solution of potassium nitrate was substituted for the 5 per cent commercial common salt and the factor 2 . 5 recommended. The modified method is still the provisional method of the A. 0. A. C. for determining the acidity of soils. Veitch* criticizes the Hopkins method upon the grounds t h a t it indicates only the apparent need for lime or the most urgent need, and claims further t h a t the acidity shown by this method is largely due t o aluminates. He also notes t h a t there is a great discrepancy between the Hopkins method and t h a t proposed by himself6 upon soils high in organic matter. , 1 From a thesis submitted t o the faculty of the University of Illinois in partial fulfillment of the requirements for the degree of Doctor of Philosophy. Acknowledgment is made of many helpful suggestions and criticisms from Prof C. G. Hopkins and Prof. A. H. Noyes. 2 Nineteenth Annual Proceedings, 0. A. C., U. S. Dept. of Agr., Bureau of Chemistry, Bulletin 73 (1902), 114. 8 U. S. Dept. of Agr., Bureau of Chemistry, Bulletin 107 (1908), 2 0 ; Hopkins, “Soil Fertility and Permanent Agriculture.” 1910, 566. 4 J. A m . Chem. Soc., 26 (1914), 637. 6 I b i d . , 24 (1902), 1120.
.
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No. 4
Harris’ claims t h a t the acidity shown by this method is due t o selective ion absorption by the soil colloids, basing his views upon the fact t h a t the acidity shown by the extract is dependent upon the character of t h e salts used. Freer2 also holds this view. Truoga strenuously combats the theory of colloidal absorption and brings evidence t o support the view of Hopkins t h a t the reaction is one of double decomposition between the acids or acid salts in the soil and the neutral salt solution. Parker4 concluded from analysis of extracts prepared by treating soils with potassium chloride and potassium acetate t h a t the base was absorbed t o a little greater extent than i t was liberated by the soil and t h a t the excess of the anion should b e accounted for by the presence of the corresponding acid. Brogue5 states t h a t i t has been repeatedly proven t h a t the base liberated by the soil is usually not merely equivalent t o the base absorbed from the solution. Sullivan,6 Morse and Curry,’l Abbott, Conn and Smalley,8 R ~ p r e c h t and , ~ others have noted the presence of aluminum and iron in salt extracts from acid soils. Ricelo concludes from hydrogen-ion concentration studies upon 31 soils using the indicator method of Sorensenll t h a t when so-called acid soils are shaken with salt solutions part of the cation of the salt is absorbed and an equivalent quantity of the base from the soil is given up t o the solution. It was t o test the above points t h a t the following investigations were made.
By Henry G. Knight OKLAHOMA
Vol.
EXPERINENTAL
Harris obtained different lime requirements for soils by repeated shaking with different salt solutions. These experiments were repeated in this laboratory using yellow-gray silt loam, and similar differences were obtained as was reported by Harris for different salts. As Hopkins claims t h a t the reaction between the neutral salt solution and the soil is one of equilibrium, the end reaction would be practically impossible t o realize by such a treatment. To overcome t h e objections which would arise from the above method provisions were made for forcing the salt solutions through the soil, so t h a t the soil particles would be continually bathed by fresh solutions. Twenty grams of yellow-gray silt loam12were placed upon a dry filter paper and the salt solution was allowed t o filter through. The filtrate was boiled and treated with 0.04 N potassium hydroxide a t room temperature using phenolphthalein as an indicator, with the results shown in Table I. Michigan Agr. College and Station, Bulletin 19 (1914). Penn. Dept. Agr., Bulletin 261 (1915), 106. 8 J . Phys. Chem., 20 (1916), 157. 4 THIS JOURNAL, 6 (1914), 831. 6 J. Phys. Chem., 19 (1915), 665. 6 U.S. Geol. Survey, Bulletin 312 (1907). 7 New Hampshire Agr. Station, Report 1906-08, 271. 8 Indiana Agr. Expt. Station, Bulletin 170 (1913). 8 Mass. Agr. Expt. Station, Bulletin 161 (1915). 10 J. Phys. Chem., 20 (1916), 214. b 11 Biochem. J . , 21 (1909), 131; Walpole, Biochem., 5 (1911), 2 0 7 ; 8 (1914), 628. If Sample No. 1. Subsoil from Southern Illinois. Lime requirement : Hopkins Method 4.2T., Veitch Method 5.6T, per acre of soil of 2,000,000lbs. 1
2
Apr., 1 9 2 0
T H E J O U R N A L OF INDUSTRIAL A N D E N G I N E E R I N G C H E M I S T R Y
TABLEI-ACIDITY OF DIFFERENT FRACTIONS OF VARIOUS SALT SOLUTIONS FILTERED THROUGH AN ACID SOIL Equivalent Salt --Cc. 0.04 N K O H Required foto Used 100 cc. 250 cc. 250 cc. 250 cc. Total T.CaCOa N KNOa .... 36.40 2.8 1.2 0.5 40.5 4.05 N KC1 . . . . . 35.95 4.3 2.1 0.5 39.95 3.99 N NaNOz.. . 26 50 9.2 1.8 1.4 38.9 3.89 N NaCl ...... 31.20 6.4 1.4 ... 39.0 3.90 N CaClz .... . 31.40 5.4 1.3 0.5 38.6 3.86
.. . .
The greatest difference shown is 0.19 T., calculated as calcium carbonate, which may easily be accounted for by errors in reading t h e end-point. I t will be noted t h a t the acidity of the sodium nitrate extract was quite marked even after 600 cc. had filtered through, while the first I O O cc. showed the lowest acidity. T h a t none of the extractions were carried t o completion is evident, but all, with the possible exception of t h a t with sodium nitrate, were carried t o a point beyond which it was impossible t o measure the acidity with any degree of accuracy by the ordinary indicator methods. The calcium salt extract would be expected t o show a slower reaction after the first surface reaction because of the greater insolubility of calcium compounds which would be formed upon the surface of t h e soil grains. INDICATOR EFFEcTs-In the first two series of extracts, a precipitate which had a rather marked effect upon the indicator was always formed in considerable quantity. The pink color produced by the addition of a slight excess of base disappeared after a time even when the titrating flask was tightly stoppered. By the further addition of base t h e color could be brought back. The end-point is also markedly influenced by t h e amount of indicator present. I t is quite apparent t h a t the indicator is absorbed t o a marked extent by t h e precipitate, and instead of t h e simple equilibrium HIn H+ Incolorless colored there must be taken into consideration the equilibrium with the absorbed indicator HIn H+ In-. HIn absorbed colorless colored Two equal quantities of a potassium salt extract of a n acid soil gave readings, as shown in Table 11; which clearly indicate the variation in results which may be obtained by using different amounts of indicator.
341
using the same quantity of phenolphthalein as a n indicator. The only difference between the duplicates was t h a t of temperature. TABLE 111-EFFECT OF TEMPERATURE UPON TITRATION END-PO~NT Temp. ' C. Cc.KOH 22 16.7 85 19.2
EFFECT
OF
TEMPERATURE
UPON
THE
AMOUNT
OF
S H O W N - Ithe ~ Hopkins method is a measure of the colloidal adsorption for the base by a soil a s maintained by Harris1 there should be a temperature effect which .could be measured. Travers2 ha5 shown t h a t the adsorption of carbon dioxide by charcoal decreases markedly with rise in temperature, and we may expect a similar change t o be shown by soils i n contact with neutral salt solutions. ACIDITY
;:i
+
__
+
TABLE11-EFFECT OF DIFFERENT AMOUNTS OF INDICATOR UPON TITRATION END-POINT Phenolphthalein 0.04 N KOH Used Required ' Drops cc. 4 62.7 10 61.1
I n several instances a point was reached which showed no visible color change with four or five drops of indicator, while upon the addition of larger quantities a marked color change was observed. To overcome as far as possible the variation due to the indicator the same quantity was used in each case unless otherwise stated. EFFECT O F T E M P E R A T U R E U P O N TITRATION-The temperature a t which the titration is carried out was found t o produce an effect which is shown in Table 111. Two equal quantities of potassium nitrate extract were titrated with 0.04 N potassium hydroxide
/
I
\
\
Fro. 1
To test this theory an apparatus, shown in Fig. L, was arranged in a constant temperature electrical oven of the Freas type. The apparatus was arranged in duplicate. a is the receptacle for the neutral salt solution, b a stopcock for regulating the flow of t h e neutral salt through tube d which passes through t h e ventilating openings c provided by the manufacturers in the stock oven. The bulb e serves t o bring the salt solution t o t h e temperature of the oven before i t runs into the receptacle f which contains the soil under investigation. The filtrate passes out of t h e apparatus through tube g , through ventilating openings in the oven a t c', and is caught in measuring flasks; 1 LOC.
2
Cil.
PYOC. Roy. SOC.London, 74 (1904),126.
T H E JOURlVAL O F I N D U S T R I A L A N D ENGINEERING C H E M I S T R Y
342
i is a thermometer placed in the center of the oven for noting the temperature Twenty grams of soil were placed in the receptacle, normal potassium nitrate solution allowed t o percolate through, and titrated at room temperature with results as shown in Table IV. TABLE IV-EFFECT
OF TEMPE$RATURE UPON REACTION BETWEEN ACID SOILAND NRUTRAL SALTSOLUTION
* --
e
TEMPERATURE------
-90' 1-' PERCOLATE 0.04 N KOH CaCOs cc. cc. T. 100 37.85 3.785 0.315 400 3.15 TOTAL500 41.00 4.1
*
-2.5' lo-0.04 N K O H CaCOs cc. T. 36.4 3.64 3 .'O 0.30 39 4 3 94
The temperature effect is very slight, and certainly does not indicate colloidal adsorption, unless this may b e an exceptional case. EFFECT OF S T R E N G T H OF S A L T soLvTIoN-Potassium nitrate solut-ions of different strengths were filtered through 20 g. samples of yellow-gray silt loam and the filtrate titrated with 0.04 N KOH, using phenolphthalein as an indicator with results shown in Table V. TABLEV-EFFECT
OF STRENGTH OF S A L T SOLUTION AMOUNTOF ACIDITY
-Cc.
Fractions cc 1-100 2-100 3-100 4-100 5-100 6- 100
.
N KNOs 35.0 3.5 1.2 0 8 0.5
TOTAL. . . ., . .,
a s T. CaCOn.. .
. .
..
40.8 4.08
UPON THE
TOTAL
0.04 N KOH Required0 5 N KNOi 0 1 h' KNOi 33.3 18.3 4.7 7.6 4.3 1.5 0.9 ' 3.7 0.6 3 .3 .. 1.6 41.0 38.8 4.10 3.88
It will be noted t h a t with the stronger salt solutions the greatest acidity is shown in the first I O O cc. b u t rapidly falls while the weaker solutions show a higher acidity in succeeding fractions. Evidently end extractions would lead t o t h e same end results regardless of the strength of the salt solution. Further study of this reaction was made by extracting I O g. samples of the same soil with various strengths of potassium nitrate solution and titrating The t h e acid in the first I O O cc. of the percolate. results are tabulated in Table VI. TABLE VI-ACIDITY SHOWN IN FIRST100 Cc. FILTRATE FROM ACIDSOIL, USINGNEUTRAL SALTSOLUTIONS OF VARIOUS STRENGTHS 0.04 N K O H Required Calculated as Normality of cc. T. CaCOa KNOa 3.40 3.06 3.04 2.65 1.60 0.82 0.39 0.23
15.2 15.3 15.2 -13.25 8.00 4.10 1.95 1.17
1.00 0.50 0.25
The acidity of the first portion of t h e extract increases with increase in concentration of the neutral salt solution. TABLE VII-AMOUNTS
OF
LIMEABSORBED BY ACIDS O I L VARIOUSSTRENGTH
FROM SOLUTIONS
OF
Lime Added Calculated as T. CaC03 12 14 16 20 30 40
Absorbed by Soil T. CaC03 11.71 13.56 14.59 16.75 20.93 24.96
Left in Solution
T. CaCOs 0.29 0.64 1.41 3.25 9.07 15.04
Final A-ormality of Solution 0.000285 0.00064 0.0014 0.0032 0.009 0.015
EXAMPLES O F SO-CALLED C O L L O I D A L A B S O R P T I O N As examples of what are usually considered colloidal phenomena t h e following experiments are submitted.
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No. 4
EXPERIMENT 1-20 g. of yellow-gray silt loam were shaken with various amounts of lime contained i n zoo cc. of solution for 1 2 hrs. and aliquot parts of t h e clear liquid were titrated with results as shown in Table VII. A greater portion of the lime is absorbed from t h e dilute solutions t h a n from the more concentrated, thus apparently following t h e colloidal absorption law. EXPERIMENT 11-20 g. of soil were placed in an extracting apparatus ( b , Fig. z ) , and a 0.04 N calcium hydroxide solution allowed t o percolate through. I n t h e diagram a is t h e receptacle for t h e base, e a stopcock for regulating the flow of t h e base into 6 , and c a graduated receiver connected with tube d which equalize the pressure. d serves t o The apparatus is a closed system and duplicate results were easily obtained. 2 3 7 . 2 cc. of filtrate1 passed through before a pink color could be detected, with phenolphthalein as an indicator, representing a lime adsorption a t this point of 2 3 . 7 T. as calcium carbonate. At t h e close of t h e experiment j37.6 cc. of filtrate had percolated through, the last j o cc. being 0.0285 N base, while the soil had absorbed a total of 3 5.28 T. of lime as calcium carbonate. 15.82 T. of lime were washed out by the first 7 0 0 cc. of distilled water, the end fraction passing through 19.46 tons of 0.00242 N alkali. lime were still left in the soil. The details are given in Table VIII, in which A and B are duplicate FIG. 2 determinations.
TABLE VIII-AMOGNT OF L I M E ABSORBSDB Y ACIDS O I L CONSTANT STRENGTH PERCOLATING THROUGH A
FROM SOLUTION OF THE SOIL
Ca(0H)z c _ _ _ _ L _ A _ _ _ _ - ~ - _
Percolate c c. 236.5 27.0 28.5 45.0 51.4 48.6 50.0 50.5 TOTAL 537.5 238.0 43.2 58.0 50.5 50.0 50.0 50.0 TOTAL 539.7
Calculated as T.CaC03 Absorbed 23.65 2.50 1.95 1.97 1.61 0.93 1.04 1.45 35.10
Per cent Absorbed 100.0 92.6 68.4 43.7 31.3 19.1 20.8 28.7
Calculated as T . CaCOa in Percolate 0.00 0.20 0.90 2.53 3.53 3.93 3.96 3.60
B
23.80 4.17 3.04 1.24 1.08 1.15 0.99 35.47
100.0 96.6 52.2 24.5 20.2 23.0 19.8
0.00 0.15 2.76 3.81 3.92 3.85 4.01
At the completion of t h e experiment water was added and t h e first 7 0 0 cc. carried through t h e equivalent of 16 tons of lime as CaC03, or nearly half t h e amount absorbed. 1
Figures given are average of results
A and B below.
T H E J O U R N A L OF I N D U S T R I A L A N D ENGINEERING CHEMISTRY
Apr., 1920
E X P E R I M E N T 111-Twenty gram samples of yellow-gray silt loam were shaken for 3 hrs. with zoo cc. of 0.04 N calcium hydroxide and potassium hydroxide, respectively. Since the solution containing the potassium hydroxide would not settle, 2 5 cc. of normal potassium nitrate were added t o both the calcium and potassium hydroxide solutions, and filtered. Titrations of I O O cc. of filtrate with 0.04 N hydrochloric acid gave the results included in Table I X .
TABLE IX-AMOUNT
OF BASENOTABSORBED BY AN ACID SOIL FROM SOLUTIONS OB
EQUIVALENT STRENGTH
0.04 N HC1 Required t o Neutralize Filtrate BASE USED cc. Ca(0H)S.. .................... 7.7 KOH.. . . . . . . . . . 31.9
As the potassium hydroxide solution was neutralized, a copious precipitate of aluminum hydroxide was formed. No precipitate was noted upon neutralizing t h e calcium hydroxide solution. I t would appear from this experiment t h a t the soil has a greater absorption power for calcium than for potassium, which is not indicated by other experiments. A chemical difference in the action of the two bases seems the more simple explanation. Potassium aluminate is soluble while calcium aluminate is not. Both are unstable except in the presence of a base. Since the potassium aluminate passes into solution i t is titrated above while the calcium aluminate is precipitated around the soil particles. This apparently throws doubt up.on the magnitude of the colloidal adsorption effects which may be assumed from Expts. I and I1 above. The probable cause would seem t o be precipitation effects. Hydrolysis will account for the rapid washing out of the lime when water is added in Expt. 11. E X C H A N G E O F BASE
There is considerable doubt whether, when a n acid soil is shaken up with a neutral salt, there is a complete exchange of base. Rice1 comes t o the conclusion t h a t there is an equivalent exchange and t h a t the acidity is due t o t h e aluminum salts. He further claims t h a t the ordinary methods of analysis are too crude t o determine this accurately. Sharp and Hoagland2 by use of the hydrogen electrode show t h a t soil acidity is due t o a n excess of hydrogen ions, and t h a t the acidity is increased by the presence of certain neutral salts. Analysis was made of the potassium nitrate extract of yellow-gray silt loam with results shown in Table X. POTASSIUM NITRATEEXTRACT OF AN ACID SOIL 627 Cc. of 0.04 N Acid = Gram 0.04741 Si02 . Pa06 0.00576 0,38822 AlaOs ...................... Trace Fen03 ...................... 0.06025 C a O . . ..................... . . . . 0,04839 0.19901
TABLE X-ANALYSIS
OB
The acid combined with the alumina would. be equivalent t o 5 7 0 cc. of 0.04 N acid leaving an excess of 5 7 cc. t o be accounted for in other ways. 1
LOG.
cit.
* J . Agr. Res.,
[3]7 (1916). 124.
343
Since, as has been shown by Blum,l alumina is completely precipitated before the hydrogen-ion concentration drops t o a value of IO-' and, cdnversely, alumina will not pass into solution until the hydrogen-ion concentration reaches a value higher t h a n IO-', i t is reasonable t o conclude t h a t there must be some absorption of base before the alumina will pass into solution. If this is true we must assume t h a t the dissolution of alumina is a secondary reaction, and the above analysis certainly points in this direction. I t was noted t h a t strongly acid extracts of this soil were highly colored with iron. T o test the question of why iron is not also taken out in larger quantities by a similar secondary reaction, I O O g. samples of yellow-gray silt loam were shaken with 2 5 0 cc. of 0.04 N acids and a partial analysis made of 1 2 j cc. of the filtered extracts. The results are given in Table X I . TABLEXI-PARTIAL ANALYSISOF WEAKLYACID EXTRACTS OF ACID SOIL A1203 ACID USED Gram Acetic. . . . . . . . . . . . . . . . . . 0.00383 Nitric. . . . . . . . . . . . . . . . . . . 0.069 13 Hydrochloric. . . . . . . . . . . . 0.07049
.
AN
Fez03 Trace Trace Trace
The solutions were still acid and upon neutralization white precipitates of aluminum hydroxide were formed in the hydrochloric and nitric acid extracts but none was noted in the acetic acid axtract until i t was neutralized and boiled. It would seem t h a t when dilute acids are allowed t o act upon an acid soil alumina is first brought into solution, and t h a t neutral salts, when brought into contact with a n acid soil, show the properties of a weak acid in this respect. DIALYSIS
Normal potassium nitrate solution was shaken with yellow-gray silt loam and allowed t o settle. The supernatant liquid was drawn off into a collodion bottle and subjected t o dialysis with the results given in Table XII. TABLEXII-DIALYSIS
OF A
NEUTRAL SALTIN CONTACTWITH AN ACID SOIL 0.04 N KOH t o Neutralize
1st water fraction.. . . . . . . . . . . . . . . . . . . . . . . . 2nd water fraction.. ....................... 3rd water fraction. . . . . . . . . . . . . . . . . . . . . . . . . Left in flask.. . . . . . . . . . . . . . . . . . . . . . . . . . . . .
TOTAL ...............................
cc. 35.1 12.06 24.9 24.9 96.96
74 per cent of the titratable acid had passed through the membrane. AS the liquid left in the dialyzing flask was being titrated a heavy precipitate formed, while t h a t which passed through remained clear upon neutralization. It was evident t h a t the acid passed through while the aluminum hydroxide did not. This, however, was t o be expected, as this is one of t h e recognized methods for the preparation of colloidal aluminum hydroxide.2 DISTILLATION O F SOIL EXTRACT
Attempts were made t o remove acid from potassium nitrate and chloride extracts of soil by prolonged distillation with steam but without success. Better success followed the distillation of the potassium 1 J . A m . Chem. Soc., 7 (1916), 1282. * G r a h a m , A n n . , 121 (1866), 41.
T H E J O U R N A L OF J N D U S T R I A L A N D ENGTNEERING C H E M I S T R Y
344
acetate extract of the yellow-gray silt loam as shown in Table XIII. TABLEXIII-DISTILLATIONOF A POTASSIUM ACETATEEXTRACT OF ACID SOIL, AND OF STOCKSOLUTION OF POTASSIUM ACETATE ( A c i d i t y in Terms of 0.04 N B a s e ) Potassium Acetate Potassium Acetate Extract--Stock Solution-Distillate Residue Distillate Residue 72.9 Cc. 21.9 Cc. 8.6 Cc. 1.95 c c .
AN
--
Phenolphthalein was used as an indicator. This experiment shows the presence of appreciable quantities of acetic acid in the soil extract. About threefourths of the acid shown by the extract was distilled over. Walter Cruml prepared colloidal aluminum hydroxide by separating the acetic acid by heating, but as there were only traces of aluminum salts carried by the potassium acetate extract2 it can hardly be conceived t h a t the phenomena may be accounted for by the presence of these salts, but rather t h a t there is a n excess of acid. COMPARISON OF CATION AND ANION ABSORPTION
To compare the cation and anion absorption of the yellow-gray silt loam from neutral salt solution a 0.0358 N solution of calcium chloride was allowed t o percolate through 2 0 g. of the soil in t h e apparatus shown in Fig. 2 . T h e extract was analyzed for calcium and chlorine. The calcium was determined in an aliquot portion of the extract which had been freed from iron and aluminum by first precipitating as the oxalate and titrating t h e precipitate with a standard potassium permanganate in the presence of dilute sulfuric acid. The chlorine was determined by the Volhard method3 using 0.04 N solutions. The results are tabulated in Table XIV.
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No. 4
a-The acidity of the salt extract of a n acid soil is independent of the temperature within the range from z j o t o g o o C. 3-The precipitate formed in titrating the soil extract obtained by the Hopkins method absorbs the indicator t o a marked extent. The end result depends upon the temperature, time, and amount of indicator used. 4-The acidity of the first portions of the neutral salt extracts of an acid soil increases with increase in concentration of the neutral salts. 5-The difference in absorption of calcium and potassium from solutions of their bases b y an acid soil may be accounted for by precipitation effects. 6-There is a marked basic exchange when a neutral salt solution is added t o an acid soil, by which alumina is carried into solution. This, however, does not account for the total acidity of t h e solution. 7-When acid soil is extracted with potassium acetate solution, a portion of acetic acid may be distilled off from the extract, showing the presence of free acid. 8-Exchange of acid radicals when an acid soil is treated with a neutral salt solution was not noted. THE USE OF CUPFERRON IN QUANTITATIVE ANALYSIS1 By G. E. F. Lundell and H. B. Knowles BUREAUOF STANDARDS, WASHINGTON. D. Received September 8, 1919
c,
INTRODUCTION
The increasing use of cupferron (the ammonium salt of nitrosophenylhydroxylamine, CeHs.N.N0.0NH4) for the determination of zirconium in its ores and metallurgical products, as well as for minor purposes such TABLEXIV-ANALYSIS OF FRACTIONS OF EXTRACT OF AN ACIDSOILBY A NEUTRAL SALTSOLUTION TO DFTERMINE CATIONAND ANIONABSORPTION as the separation of iron and titanium from manganese Extract Fractions Calcium Chlorine Acid and aluminum in limestone analysis,2 and other deNormality Normality Normality c c. terminations for which i t has been recommended, 0.0348 0.0046 0.0180 1-50 0.0358 0.0037 0.0314 2-50 makes a review oE the possibilities and limitations of 0.0358 0.0028 0.0332 3-50 0.0358 0.0022 0.0335 4-50 this reagent highly desirable. This paper presents a The cation is absorbed t o a measurable extent, for review of the literature dealing with the use of cupwhich the change in acidity fails t o account. There ferron as a quantitative precipitant, and gives the remust, therefore, have been an exchange of base. sults of many tests which were performed a t this This confirms the fact t h a t there is a basic exchange Bureau in connection with a n attempt t o adapt the cupferron method t o the determination of zirconium regardless of the neutral salt used. (See analysis of in its ores and metallurgical products. potassium nitrate extract.) The anion is absorbed little or not a t all. The slight absorption shown in the first 5 0 cc of extract is probably due t o the wetting of the particles and t o a slight dilution of the extract by moisture in the air-dried soil. A small amount of precipitate was formed in each case upon neutralizing the extract. SUMMARY
I-When normal solutions of potassium nitrate, potassium chloride, sodium nitrate, sodium chloride, and calcium chloride were percolated through an acid soil all gave the same end titrations, using phenolphthalein as a n indicator. This corroborates Hopkins’ statements. 1 Ann.,
89 (1854), 168. Compare Conner, THESJOURNAL, 8 (1916). 35. 8 Ann., 190, 1 .
2
GENERAL PRINCIPLES
Cupferron precipitates are salts in which the ammonium radical of the reagent has been replaced by metals. Precipitations are performed in cold solutions containing free mineral or organic acids. Cold solutions must be employed t o prevent decomposition of the reagent into various organic substances, such as nitrobenzene, and the temperature of precipitation is usually specified as “cooled in ice water.’’ A 6 per cent water solution of the reagent is used and complete precipitation is indicated by the formation of a temporary flash of a fine, white precipitate which redissolves, as contrasted with the flocculent insoluble cupferron Published by permission of the Director of the Bureau of Standards. Private communication from J. A. Holladay, Electro-Metallurgical Co.,Niagara Falls, N. Y. 1 2
