STUDIES I N COPRECIPITATION. I V
THE COPRECIPITATION OF ALKALIIONS WITH CALCIUMOXALATE AND THE ADSORBENT PROPERTIES OF THE LATTER^ I. M. KOLTHOFF
AND
E. B. SANDELL
School of Chemistry, University of Minnesota, Minneapolis, Minnesota Received June 6 , 193.9
In a previous general discussion (1) of the phenomena of coprecipitation, the presence of foreign ions in the interior of a precipitate was quite generally ascribed to an adsorption of these ions during the growth of the crystalline precipitate. Therefore it may be expected that the order of the amount of coprecipitation in a series of cations or anions, respectively, will be the same as the order of adsorption at the external surface of the precipitate after it has been formed. In order to prove this conclusion, the adsorption of various electrolytes by aged calcium oxalate crystals has been determined. This study is preliminary in character. Many difficulties, especially of analytical nature, have been encountered, which have not yet been completely overcome; therefore amore extensive investigation of the adsorbent properties of calcium oxalate is reserved for the future. The results obtained in the present study are conclusive insofar as the relation between adsorbability and coprecipitation of ions is concerned. According t o the Paneth-Fajans adsorption rule, the adsorbability of a salt having an ion in common with the crystalline adsorbent should increase with decreasing solubility of the salt. J. s. Beekley and H. s. Taylor (2) in a study of the adsorption of silver salts by silver iodide obtained results in fair harmony with this rule; the agreement would have been still better if they had considered the deformation of the adsorbed anions. It follows from the Paneth-Fajans rule that the adsorbability of any individual ion will be larger the smaller the solubility of the compound that it can form with the ion of opposite charge in the adsorbing lattice. This conclusion, however, is not confirmed by the work of Sven Od6n (3) and Mlle. L. de Brouckere (4) ; a more exhaustive discussion must be postponed until more experimental data are available. The material covered in this and following papers has been taken from a thesis submitted by E. B. Sandell (Du Pont Fellow in Chemistry a t the University of Minnesota, 1931-1932) to the Graduate School of the University of Minnesota in partial fulfillment of the requirements for the degree of Doctor of Philosophy. 443
444
I. M. KOLTHOFF AND E. B. SANDELL ADSORPTION OF ELECTROLYTES BY CALCIUM OXALATE
Preparation of aged calcium oxalate suspensions Preparation I . 0.515 mole of calcium chloride hexahydrate dissolved in 1 liter of hot water was added slowly over a period of twenty minutes to 0.50 mole of ammonium oxalate dissolved in 3 liters of solution a t 90°C. The precipitate was washed by decantation until washings were entirely chloride-free. The precipitate contained traces of chloride that could not be washed out; the amount did not exceed 0.01 per cent. The preparation was allowed to stand under water for a week before use in the adsorption experiments. The average crystal diameter was estimated t o be 0.5 micron or less. Preparation I I . Precipitation was made in the same way as described under Preparation I, with the sole difference that at the end of the precipitation ammonium oxalate was a few per cent in excess. The preparation was washed by decantation thirteen times and kept under water for two weeks before being used. The average crystal diameter was estimated to be 0.5 micron or slightly more. Preparation I I I . 1.00 mole of calcium chloride in 1 liter of hot water was added in a slow stream to 1.02 moles of ammonium oxalate in 3 liters of boiling water containing 2 cc. of hydrochloric acid. The time of addition was fifteen minutes. The precipitate was washed by decantation twelve times. It was kept under water for two months before using. The crystal size was slightly larger than in the preceding preparations,
,
Procedure The carefully washed precipitate was mixed with water t o give a suspension containing approximately 3 g. of calcium oxalate in 50 cc. From this suspension, well shaken, 50 cc. were removed with a pipet and transferred to a dry glass-stoppered bottle. Then 50 cc. of the solution containing the substance to be adsorbed were mixed with the suspension, and the whole was agitated in a mechanical shaker. The time of shaking is recorded in table 1. The adsorption took place a t room temperature, 25f2"C. Because of the uncertainty involved in determining the small amount of substance adsorbed, the use of a thermostat was considered superfluous. After shaking, the mixture was filtered through paper, and 50 cc. of the filtrate was removed with a pipet and titrated with the suitable standard solution. I n the case of the oxalates, 0.02 N potassium permanganate was used. The end-point correction was obtained by adding potassium iodide t o the cold titrated solution and titrating the liberated iodine with 0.005 N sodium thiosulfate in the usual manner. The original solution containing the substance to be adsorbed was standardized in exactly the same way by taking 25 cc. In the case of the adsorption of
445
STUDIES I N COPRECIPITATION. IV
iodates, 0.02 or 0.002 N sodium thiosulfate served for titrating the iodine liberated by adding iodide to the acidified solution, starch being used as indicator. Bromate was titrated in the same way. For the titration of hydrochloric acid and oxalic acid, dilute carbonate-free sodium hydroxide was used. All salts used in the adsorption determinations were of the highest purity. The oxalate solutions were prepared immediately before use to preclude the formation of carbonate by photochemical decomposition. The results of these experiments will be found in table 1.
TABLE 1 Adsorption by calcium oxalate monohydrate
TIME
EQUILIB-
I
YIICROMOLES ADSORBED P E R QRAM OF CALCIUM
OF BOLUTION
Ammonium oxalate. . . . . . . Potassium oxalate. . . . . . . . . Sodium oxalate. .......... Sodium oxalate. . . . . . . . . . . Lithium oxalate.. . . . . . . . . . Magnesium oxalate. ....... Oxalic acid ................ Oxalic acid . . . . . . . . . . . . . . . . Ammonium iodate. . . . . . . . . Potassium iodate.. . . . . . . . . Potassium iodate.. . . . . . . . . Sodium iodate.. . . . . . . . . . . . Lithium iodate.. . . . . . . . . . . Calcium iodate.. . . . . . . . . . . Calcium iodate.. . . . . . . . . . . Iodic acid. . . . . . . . . . . . . . . . . Iodic acid.. . . . . . . . . . . . . . . . Iodic acid. . . . . . . . . . . . . . . . . Potassium bromate.. . . . . . . Sodium oxalate sodium iodate . . . . . . . . . . . . . . . . . .
+
ALATE MONOHYDRATE
LGITATION
CONCEN-
hours
molarity
1 1 1 3 1 1 1 3 1 1 1 1 1 1 3 1 2 2 1
0.0100 5.5 0.0100 0.0100 9.3 0.0100 9.8 0.0100 0.0014 6.5 0.0100 6.0(Ox-) 0.0100 6.2(Ox-) 0.00100 0.00100 0.58 0.0100 5.2 0.00100 0.0100 0.00100 1.7 0.00100 1.7 0.00100 0.00100 0,00082 0.0100
1
0.00100 0.00100
Preparation I1
3.7 4.8 7.5 7.2 3.5 7.9(Ox -) 8.0 (H+) 0.73 0.95 4.3 0.86 4.6 1.7
Preparation 111
1.3 (3 hours) 2.7 (3hours) 3.9 3.9
3.6 (Ox -)
1.32 (IO;) 1.57 (IOr) 18.0 (H+) 2.0 0.14 (IO;)
The millimoles of substance adsorbed per gram of calcium oxalate monohydrate equals NF W
[va- v,(Va 50.00 + 50.00)]
where N = the analytical normality of the solution used for titration, F = conversion factor (NF = molarity of solution of adsorbed ion), V z = cc. of standard solution required to titrate 50.00 cc. of solution containing substance to be adsorbed,
446
I. M. KOLTHOFF AND E. B. SANDELL
VI
cc. of standard solution required to titrate 50.00 cc. of filtrate after equilibrium, V , = cc. of water in 50.00 cc. of calcium oxalate suspension (determined by weighing 50.00 cc. of suspension and deducting from this value the weight of calcium oxalate monohydrate obtained by evaporating the suspension to dryness and heating the residue at 105"C.), and W = grams of calcium oxalate monohydrate in 50.00 cc. of suspension. =
DISCUSSION O F RESULTS
The adsorption of alkali oxalates by aged calcium oxalate is very small in amount. Even with the fine precipitates used in these determinations it is not easy to obtain reproducible results. The most strongly adsorbing preparation of calcium oxalate in equilibrium with 0.02 N sodium oxalate, the alkali oxalate most strongly adsorbed, contains less than 0.1 per cent of its weight of the foreign oxalate on the surface. It is evident that from the analytical point of view external adsorption cannot play any important r81e in the contamination of calcium oxalate, especially when it is remembered that even this small amount of adsorbed substance is partly removed on washing the precipitate. In analytical precipitations the crystals are much larger than those used in these experiments, so that in considering the causes of contamination, it is possible to leave out of sight entirely that due to external surface adsorption. The order of adsorption of the oxalates at room temperature at the same equilibrium concentration appears to be as follows: lithium < ammonium < potassium < oxalic acid, sodium < magnesium. Evidently the solubility of the oxalates does not alone determine the adsorbability, for then the order would be: potassium < oxalic acid < lithium < ammonium < sodium < magnesium. (Solubilities of oxalates : potassium 1.65, lithium 0.61, ammonium 0.39, sodium 0.28, magnesium 0.0020, oxalic acid 0.81 moles per thousand grams of water at 25°C.) Ammonium oxalate especially seems to be out of place, In agreement with the order of adsorbability, the amount of coprecipitation increases in the order ammonium < potassium < sodium < magnesium, as will be shown later. The mechanism of the adsorption of oxalic acid probably is somewhat different from that of the salts, since there is a possibility that some of it may be held at the surface in the form of bioxalate ions. In addition, it should be mentioned that the amount of iodic acid adsorbed, as found from the acidimetric determination, is much larger than that found from .the iodate analysis. The explanation is that the preparation of calcium oxalate used contained coprecipitated calcium hydroxide, which was neutralized by the hydrogen ions. I t was experimentally proved that in the
STUDIES I N COPRECIPITATION. IV
447
case of hydrochloric acid and of iodic acid, more calcium than oxalate went into solution when calcium oxalate was shaken with the acids, thus indicating that calcium hydroxide existed on the surface of the precipitate and that this reacted with the hydrogen ions. Therefore, the results of acidimetric analysis cannot be used here for obtaining the amount of acid adsorbed, It is possible that the inequality in the adsorption of the two ions is partly to be attributed t o an exchange adsorption between hydrogen and calcium ions. Calcium iodate is more strongly adsorbed than the alkali iodates, as would be expected from the precipitation rule. From 0,001 molar solutions the adsorption of the alkali iodates is of the same order of magnitude; more work should be done over a wider range of concentrations in order to learn whether sodium iodate is more strongly adsorbed than potassium iodate, and the latter more strongly than ammonium iodate. From the figures in table 1, it will be seen that the adsorption of the alkali iodates increases very strongly with the concentration up to an equilibrium concentration of a t least 0.01 molar. Since calcium iodate is much less soluble than calcium bromate (0.0078 and 3 moles per thousand grams respectively at 25OC.) a stronger adsorption of the former ion was anticipated. I n agreement herewith a much stronger adsorption of potassium iodate than of bromate was found (Preparation 11). Again, it was found in the study of the coprecipitation of anions that iodate is carried down to a much greater extent than bromate, if the precipitations are made under identical conditions. Very interesting is the result of the experiment in which calcium oxalate was shaken with a mixture of sodium oxalate and sodium iodate containing 0.001 moles of each per liter. Although the amount of iodate adsorbed is six times less than in the absence of oxalate, it is evident that the latter is not able to replace all adsorbed iodate a t the surface of the calcium oxalate. There is competition between the oxalate and iodate ions and since calcium oxalate is less soluble than calcium iodate, the former ion will be more strongly adsorbed than the latter, although oxalate is not able to prevent entirely the adsorption of the anion which is foreign t o the lattice. The adsorbability is not determined by the fact whether the ion is identical with, or of the same size as, that of the same sign in the lattice. The result found explains why a coprecipitation of iodate or some other anion is found if, during the precipitat'ion of calcium oxalate, an excess of oxalate ions is present in the solution, and why cations can be coprecipitated even if there is an excess of calcium present during the precipitation. The presence of the lattice ion does not entirely prevent the adsorption of foreign ions of the same.sign.
448
I. M. KOLTHOFF AND E. B. SANDELL COPRECIPITATION OF THE ALKALIES WITH CALCIUM OXALATE
The contamination of calcium oxalate by the alkalies was studied under the following heads: 1. Coprecipitation of alkali-sodium, potassium or ammonium-when the corresponding oxalate was used as a precipitant for calcium chloride. The results are given in table 3. 2. Special investigation of sodium coprecipitation. Sodium was present here as the chloride, not the oxalate, and ammonium oxalate was used as the precipitant for calcium. The results are given in table 4. 3. Investigation of alkali coprecipitation under analytical conditions. Ammonium oxalate was used as precipitant; the alkalies were present as chlorides. The results are given in table 5 .
Methods of determining the alkalies in the calcium oxalate precipitate Sodium. This cation was determined by the simple method of Barber and Kolthoff ( 5 ) . The presence of calcium as chloride does not interfere in the procedure given by these authors, so that their directions have been followed with minor changes to better accommodate the method to the present problem of determining sodium in calcium oxalate. The calcium oxalate precipitate was first dried and then ignited at incipient red heat t o convert it to calcium carbonate. The latter was dissolved in dilute acetic acid; the excess acid was removed by heating at 100OC. until solid calcium acetate remained. To the residue was added 10 cc. of zinc uranyl acetate reagent prepared according to the directions of Barber and Kolthoff. The resulting mixture was shaken until all the calcium acetate had been dissolved and was then allowed to stand overnight to insure the complete separation of the small amounts of sodium present. The precipitate was collected in a sintered glass filtering crucible and washed successively with the reagent and a saturated alcoholic solution of sodium zinc uranyl acetate; after washing several times with ether, it was dried by passing a stream of air through the crucible for five minutes and then weighed. Blanks with calcium oxalate to which had been added known small amounts of sodium gave satisfactory results. Only one small source of difficulty was encountered; sometimes after the ignition of the calcium oxalate, the residue contained small amounts of carbon. However, this substance could be removed by repeatedly evaporating the residue with small amounts of concentrated nitric acid. Potassium. It isnot easy to determine the potassium content of contaminated calcium oxalate precipitates. Sodium cobaltinitrite, the most sensitive precipitant for potassium, was chosen as the reagent most likely tjo give results satisfactory for the present purpose. The method finally
449
STUDIES IN COPRECIPITATION. I V
adopted consisted in weighing the air-dried precipitate of sodium potassium cobaltinitrite formed in alcoholic solution. The determination was made in the following manner. The washed and dried calcium oxalate precipitate was ignited a t dull redness to convert it entirely to the carbonate. The residue was dissolved in a slight excess of hydrochloric acid and the resulting solution was evaporated to dryness. The calcium chloride residue containing the traces of potassium sought was dissolved in 5 cc. of 95 per cent alcohol a t room temperature. When complete solution had been effected, 2 cc. of sodium cobaltinitrite prepared according to the directions of de Koninck (6) was added. In order to insure complete precipitation the mixture was allowed to stand overnight. The precipitate of sodium potassium cobaltinitrite was collected in a glass filtering crucible and washed successively with small portions of 70 per cent alcohol, 95 per cent alcohol and ether, and finally dried a t room temperature by drawing air through the crucible for a short time. In order to test the reliability of the method as well as to determine TABLE 2 POTASSIUN OXALATE TAKEN
OF THE
KaC?O4 m
mg.
0.20 0.40
0.70 1.2 2.0
0.18 0.18 0.24 0.23 0.20
POTASSIUM OXALATE
ERROR
FOUND
(Factor = 0.201
w.
w.
0.22 0.44 0.58 1.04 1.98
+0.02
4-0.04 -0.12 -0 18 -0.02 I
the amount of potassium in the precipitate, blanks containing small known amounts of potassium in the presence of the same amounts of pure potassium- and ammonium-free calcium chloride as were present in the actual determination were run in the way just described. The results will be found in table 2. It is to be noted that potassium in the blanks was present as the chloride. In the table, however, the amount of potassium added, and found, is expressed in terms of the oxalate, the form in which potassium occurs in the contaminated calcium oxalate. The second column of the table gives the value of the experimentally determined factor: KzCz04 Weight of precipitate
The values obtained are not constant, which is to be attributed in part, a t least, to the experimental errors necessarily associated with the handling and weighing of such small quantities of precipitate as are being dealt with here. The value of the factor adopted for use in the calculations
450
I. M. KOLTHOFF AND E. B. SANDELL
is 0.20. The third column of table 2 gives the amount of potassium oxalate found using the factor 0.20; the agreement between the amount taken and that found is quite satisfactory for our purpose. Ammonium. The ammonium content of calcium oxalate was determined colorimetrically with Nessler’s reagent? The washed precipitate of calcium oxalate was dissolved by warming with 2 cc. of concentrated hydrochloric acid diluted with a small volume of water. When all the calcium oxalate had dissolved, the liquid was cooled and diluted with approximately 75 cc. of water. To this solution was added 5 cc. of 6 N sodium carbonate in small portions with agitation. Finally, the volume was made up to exactly 100 cc. and after mixing, the liquid was filtered through a glass crucible. Ten cc. of the filtrate was taken and treated in a small vial of 20 cc. capacity with 0.5 cc. of sodium tartrate (50 grams in 100 cc. of water) and 0.5 cc. of Nessler’s reagent. The color comparison was made with a series of standards in identical vials made up fresh and treated with sodium carbonate and Nessler’s reagent simultaneously with the unknown. The standards usually differed by 0.002 mg. of ammonia. The color comparison was made a few minutes after mixing with Nessler’s reagent. On standing, the unknown solution slowly became turbid, but the color comparison could be made quite leisurely without interference from this cause. The presence of tartrate is indispensable; in its absence the unknown becomes turbid so rapidly that the colorimetric determination is rendered impossible. Blanks run with ammonium-free calcium oxalate in the manner described gave satisfactory results, e.g., 0.06, 0.10 and 0.16 mg. of ammonium, respectively, were taken, and 0.06-0.07, 0.10-0.11, and 0.16 mg. were found.
Explanation of tables 3, 4, and 6: coprecipitation of alkalies The solutions of calcium chloride and alkali oxalate used in these precipitations were all 0.25 N. In the ordinary precipitations, the salt in excess was contained in the precipitating vessel and the other solution was added to it with hand shaking until the excess of the first solution was 20 per cent. The volume of water originally added was such that after the precipitation had been made the total volume amounted to 100 cc. in each case. Filtration was made immediately after precipitation; in those cases in which the precipitate was formed in the cold, the filtration and washing required several hours on account of the very finely divided condition of the precipitate. Precipitates obtained in the cold were washed ten times with small portions of cold water, and those formed in hot solution a like number of times with hot water, Precipitation according to Hahn has already been described (7). For the crystal size of the various precipitates see table 2 2 Nessler’s reagent: 3.5 grams of mercuric iodide, 25 grams of potassium iodide, and 15 grams of potassium hydroxide in 100 cc. of solution.
451
STUDIES I N COPRECIPITATION. IV
of a previous paper (7). The amount of alkali oxalate coprecipitated is expressed in two ways in table 3: (1) as milligrams of anhydrous alkali oxalate per gram of calcium oxalate monohydrate, and (2) as millimoles of alkali oxalate per gram of calcium oxalate monohydrate. The alkalies are actually coprecipitated as the oxalates, as is proved by the fact that calcium oxalate precipitated in the presence of much alkali chloride contains merest traces of chloride. In table 4 the results of a systematic study of the coprecipitation of sodium present in the solution as sodium TABLE 3 Copreeipitation of alkali oxalates with calcium oxalate PRECIPITATE NUM-
CONDITIONS OF PRECIPITATION
BER
CIPI-
1
2 3 4 5 6 7 8 9 10
IXCESS ION URING PREATION
20Oxto24Ca*, 55Hz0; R.T.; 1 minute 20 Ca to 24 Ox, 55 HzO;R.T.; 1 minute As in 1, but a t 100°C. As in 2, but at 100°C. Hahn's procedure. 21 Ca and 20 Ox; R.T.; 10 minutes Hahn's procedure. 20 Ca and 21 Ox; R.T. ;10 minutes As in 5, but at 100°C. As in 6, but at 100°C. As in 2, but digested a t 90100°C. for 18 hours As in 4, but in the presence of 1 g. of NaC1
1
I
I
AMOUNT OF ALKALI OXALATE COPRECIPITATED PER QRAM OR' CALCIUM OXALATE MONOEYDRATE
Sodium oxalate mg.
1 P","xta&m
millimoles
I
mg.
Ca
24
0.18
19
Ox
18
0.135 14
millimoles
0.115
mg.
5.5
millimoles
0.044
0.085 8.0 0.064
Ca Ox Ca
5 . 0 0.037 0.8 0.0051 0.8 0.0064 8 . 5 0.064 1.2 0.0072 1.05 0.0083 5.5 0.041 1.0 0.006 0.55 0.0044
ox
9 . 5 0.071
Ca ox
1.4 0.0081 1.15 0.0091
Ox
5 . 0 0.037 0 . 5 0.0031 0.1 0.0008 7.5 0.056 0.550.003: 0.1 0.0008 See t de 4 0.4 0.0024 0.2 0.0016
ox
0.45 0.002;
* 20 Ox = 20 CC.
of 0.25 N alkali oxalate; 20 Ca = 20 cc. of 0.25 N calcium chloride; R.T. = room temperature.
chloride, using ammonium oxalate as precipitating reagent, are reported. In table 5 some results on the coprecipitation of alkalies under conditions similar to those of the ordinary analytical practice are given. DISCUSSION O F RESULTS
1. The coprecipitation of sodium is much greater than that of potassium and ammonium no matter what the conditions of precipitation are; the difference is most marked in hot solution, especially if the precipitation is made according to Hahn. The coprecipitation of sodium is then still
TABLE 4 Coprecipitation of s o d i u m present as chloride
P;
w
2
E
0"
P cc.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23
24 25 26 27 28 29 30 31
32 33
-
cc.
20 24 20 20 20 24 24 As 7; stood 50
I
20 24 20 24 20 21 20 21
moles
cc.
legrees C.
24 Ca 0.05 55 ox 20 0.05 55 24 0.10 55 Ca Ca 24 0.05 55 Ca 24 0.01 55 0.10 ox 55 20 20 0.01 ox 55 five days before filtration 20 I Ox 1 30 0.01
R.T. R.T.
24 20 24 20 21 20 21 20
R.T. R.T. 100
I
Ca
ox Ca
ox ox Ca ox Ca
55 55 55 55 55 55 55 55
0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01
ma.
31.4 34.0 12.7 9.9 5.9 19.7 9.7
100 100 100 100 100
8.2
100
100
R.T. R.T. 100 100
Acid* Acid* Acid* Acid* Hahn's procedure Hahn Hahn Hahn
As 7, but in presence of 5 cc. of glacial acetic I id As 21; digested at 100°C for 18 hours R.T. 5 cc. of glacial acetic acid 50 O.O1 24 2o Ox As 23; di sted at 100°C for 18 hours 0.01 R.T. Digested 18 hours 20 24 Ca 55 at 100°C. Digested 18hours 55 0.01 100 20 24 Ca a t 100°C. As in 26 ox 0.01 R.T. Digested 15 hours 24 20 55 at 100°C. As in 28 Digested 18 hours 0.01 100 24 ox 55 20 a t 100°C. Digested 18 hours ox 0.01 100 20 50 55 at 1 0 0 O C . Acidt 0.01 90 24 Ca 20 50 Acidt 0.01 90 24 ox 50 20
I
I
I
8.8 5.5 9.95 8.1 9.6 5.5 5.5 11.5 18.0 9.0 9.2 4.45 11.5 8.7 18.2 4.75 0.7 1.1 1.5 0.77 0.82 4.6 0.9 3.0 3.7
-of concentrated HC1 was added to volume of 100 cc., neutralized a t R.T. with dilute ammonia (2 minutes). t 5 cc. of concentrated HC1 was added t o volume of 100 cc., heated with 3 to 4 grams of urea a t 90°C. for 20 hours. 452
* 2 cc.
STUDIES I N COPRECIPITATION. IV
PRECIPI-
TATE NUMBER
1
453
TABLE 5 Coprecipitation of alkalies under analytical conditions 20 cc. of 0.25 N calcium chloride used in all experiments CONDITIONS
A. Sodium mg.
30 CC. of 1 molar sodium chloride present. Precipitation made a t 100°C. by adding 1 g. of ammonium oxalate in 50 cc. of water t o 250 cc. of solution containing 5 cc. of concentrated hydrochloric acid; then neutralized with 4 N ammonia added dropwise.
13.0
30 cc. of 1 molar sodium chloride present. One gram of ammonium oxalate in 50 cc. of water added to 250 cc. of calcium solution at 100°C., containing 1.0 cc. of concentrated ammonium hydroxide. Time, 1 minute.
8.2
30 cc. of 1 molar sodium chloride present. One gram of ammonium oxalate in 25 cc. of cold water added all at once to calcium solution having total volume of 250 cc. Precipitation made at room temperature. Then digested a t 100°C. for two days.
2.0
B . Potassium 4
0.020 mole of potassium chloride present. Precipitation made at 100°C. b y adding 1 g. of ammonium oxalate to 200 cc. of solution containing 2 cc. of concentrated hydrochloric acid and neutralizing with dilute ammonia added dropwise.
1.4
5
0.020 mole of potassium chloride present. Precipitation made a t 100°C. by adding rapidly dropwise 1 g. of ammonium oxalate in 50 cc. of hot water to 200 cc. of solution containing 1 cc. of concentrated ammonia.
0.3
6
One gram of ammonium oxalate in 50 cc. of water added to boiling calcium solution having a volume of 200 cc. Two cc. of concentrated hydrochloric acid present. Then neutralized with dilute (2 N ) ammonia added dropwise.
2.4
7
Precipitation made by adding 1 g. of ammonium oxalate in 50 cc. of water t o 200 cc. of hot calcium solution containing 1 CC. of concentrated ammonium hydroxide.
2.4
8
Exactly as in 1 except that 1.0 g. of ammonium chloride was present.
3.5
454
I. M. KOLTHOFF AND E. B. SANDELL
strong (0.5-1 per cent), while that of the other alkalies has decreased to less than 0.1 per cent and sometimes to only 0.01 per cent. At room temperature-ordinary precipitation-more potassium is coprecipitated than ammonium, but at 100°C. the respective amounts are nearly the same, with an apparent tendency for potassium to exceed ammonium. Precipitates obtained according to the method of F. L. Hahn are contaminated equally if they have been formed in cold solution, but in hot solution more potassium is coprecipitated. I n general, potassium tends to be coprecipitated to a greater extent than ammonium. T h e order of coprecipitation of alkalies i s therefore the order of adsorbability. 2. It is interesting to note that calcium oxalate formed according to Hahn’s procedure in hot solution in the presence of sodium contains as much of this element as when the precipitation is made in the ordinary way in hot solution, so that this special method of precipitation possesses no advantages. Hahn’s method gives better results with potassium, and especially with ammonium, but since those elements contaminate the oxalate precipitate formed in the cold only to the extent of 0.05-0.1 per cent of the weight of the precipitate, its application is unnecessary. 3. The order and amount of coprecipitation of sodium added as chloride under various conditions (table 4) correspond in general to the results obtained by precipitating sodium oxalate with calcium (table 3). The high sodium coprecipitation obtained here in Hahn’s method is due to the special way in which this procedure has been carried out, i.e., calcium and oxalate were added simultaneously to a solution of sodium chloride. The ratio of sodium to calcium or oxalate present a t any moment was therefore large and the amount of sodium coprecipitated was consequently large as expected. Whether calcium oxalate is precipitated in neutral or ammoniacal solution makes no difference in the amount of sodium carried down. Neutralization of a hydrochloric acid solution of calcium oxalate with dilute ammonia at room temperature or at 100°C. gives a precipitate which still contains as much sodium as that obtained by direct precipitation in neutral or ammoniacal medium, The ion in excess-calcium or oxalatemakes hardly any difference if a strongly acid solution is slowly neutralized. Crystals of calcium oxalate that are very large can be obtained by adding urea to a strongly acid solution of calcium and alkali oxalate and then digesting the mixture (procedure of H. H. Willard) ; the slow neutralization of the acid gives a slow precipitation of the oxalate with the consequent formation of large crystals. These, however, still contain fairly large amounts-analytically speaking-of sodium, although less than when the precipitation is made in the standard way. Only one method under most conditions will give calcium oxalate sufficiently free from sodium to allow its use for precise ana!ysis, and t,hat is
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the method of recrystallization of a finely divided precipitate formed at room temperature. By precipitating calcium oxalate in a volume of 50 to 75 cc. a t room temperature and digesting, the fine crystals first precipitated undergo an entire structural change due to transformation of the higher hydrates, and grow a t the expense of the finest, with the result that a t the end of a day or so the crystals are a few microns in diameter and then contain much less sodium than originally. It is true that such a precipitate filters considerably more slowly than one obtained under the usual conditions of analytical precipitation. However, by the use of glass filtering crucibles such a precipitate can be filtered and washed in approximately half an hour. It seems to us that this method of procedure is far superior to that of double precipitation of calcium oxalate, which would otherwise be necessary if a n y considerable amount of sodium is present. The application of the results to analytical work will be discussed in a forthcoming paper. 4. When calcium oxalate is precipitated with excess ammonium oxalate (table 4), no other alkali ion being present, under analytical conditions either by neutralizing a hot acid solution with dilute ammonia or by adding ammonium oxalate directly t o a faintly ammoniacal solution, then the amount of ammonium oxalate coprecipitated is of the order of a few tenths of a per cent. In the presence of much ammonium chloride this value becomes larger. Thus, precipitation from acid solution by neutralizing with ammonia in the presence of one gram of ammonium chloride gave a contamination of 0.35 per cent. This amount will not be serious in ordinary analytical work, especially since this error will be compensated by a slight solubility loss in washing with water if the determination is to be made volumetrically with permanganate. If the determination is made according to the gravimetric procedure, a coprecipitation of ammonium is, of course, immaterial. 5. In the precipitation of calcium oxalate from cold solutions, a larger coprecipitation of alkalies is found in the presence of an excess of calcium than of oxalate (table 3, experiments 1 and 2), a result in opposition to the rules of coprecipitation. The explanation of this more or less abnormal behavior is relatively simple. In a study of the formation of higher hydrates of calcium oxalate (previous paper (7)) it has been shown that a t room temperature fairly large amounts of the higher hydrates are formed. The latter are not stable when left in contact with the mother liquor and are transformed into the monohydrate, this process being completed within twenty-four hours. The higher hydrates are stabilized by an excess of calcium, whereas oxalate ions promote the transformation. This transformation is accompanied by an internal structural change and a recrystallization resulting in a purification of the primary precipitate; this purification takes place much more quickly in a presence of an excess of oxalate
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(experiment 2, table 3) than of calcium (experiment 1). The filtration and washing of the precipitates obtained a t room temperature is slow, more than two hours being required. The amount of coprecipitated alkali found is, therefore, more or less accidental (see, for example, experiments 1 and 2, table 4, where more sodium was coprecipitated if oxalate was in excess); it may be expected that the primary precipitate formed in the presence of an excess of oxalate will contain more sodium than that separated from a solution containing an excess of calcium. That the above interpretation is correct will be shown in a discussion of the coprecipitation of iodate with calcium oxalate (next paper); and further, in a paper dealing with the internal structural changes on the aging of freshly prepared calcium oxalate, it will be demonstrated that the fresh precipitate formed in the cold is porous and permits fairly rapid effusion of foreign constituents from, or diffusion into, its interior. If no higher hydrates are formed (precipitation according to Hahn; experiments 5 and 6, table 3; 16 and 17, table 4), the coprecipitation of the alkalies is in harmony with the general rules, Le., more alkali in the precipitate if the oxalate is in excess. 6. In the precipitation of calcium oxalate a t higher temperatures invariably more alkali is carried down in the presence of an excess of oxalate than of calcium, a result in harmony with the rules of coprecipitation. The coprecipitation of sodium increases with increasing excess of oxalate in the medium. Since no higher hydrates of calcium oxalate are formed at higher temperature, this normal behavior was anticipated. 7. The amount of coprecipitation of alkalies a t 100°C. is much less than a t room temperature. At the present time it is impossible to make any predictions regarding the influence of the temperature upon the amount of coprecipitation, since the conditions of crystal growth, the speed of adsorption, and the adsorbability of the ions a t the two temperatures are quite different, and no quantitative information is available for a systematic analysis of these factors a t different temperatures. It is interesting to note that in contrast to the alkalies, the coprecipitation of the anions is invariably higher a t 100°C. than a t room temperature (compare with next paper). 8. Coprecipitation of one alkali ion does not prevent the carrying down of another alkali in the same precipitate. Thus it is shown by experiment 10 (table 3) that sodium and potassium are simultaneously coprecipitated, whereas comparison of results of experiments in tables 3 and 5 indicates that ammonium and potassium are carried down a t the same time. This result is not surprising, because during the growth of the precipitate competition between the various foreign ions to be adsorbed a t the surface of the growing particles takes place and the coprecipitation of the ion with the highest adsorption potential will predominate. A coprecipitation of only one type of ion if other kinds of the same sign are present, as D.
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Balarew (8) claims to be the case, is contrary to what we may expect from our knowledge of adsorption and coprecipitation phenomena and to experimental evidence. 9. The coprecipitation of sodium with calcium oxalate does not increase linearly with the concentration of the alkali ion in the solution, but seems to be an exponential function of the latter (compare experiments 3, 4, and 5, table 4). Thus with 0.01 mole of sodium chloride in the solution 5.9 mg. of sodium oxalate were coprecipitated; with 1.0 mole in the same volume, 12.7 mg. of sodium oxalate were coprecipitated. Such an exponential relation may be expected from our knowledge of the relation between amount adsorbed and equilibrium concentration in the solution. Therefore, a relatively large coprecipitation of an ion which is strongly adsorbed may be expected even if i t is present in fairly small concentrations. 10. Digestion of a freshly formed precipitate results in a purification of the crystals. The most effective method of obtaining pure calcium oxalate is by precipitation from a weakly acid solution (pH 4 to 6; compare paper V of this series) a t room temperature and digestion for eighteen hours or so thereafter. 11. The reproducibility of the coprecipitation experiments is fairly good, if the precipitations are repeated under identical conditions. SUMMARY
1. The order of the adsorption of alkali ions by, and coprecipitation with, calcium oxalate is the same, sodium being more strongly adsorbed and coprecipitated than potassium and the latter more strongly than ammonium. 2. Ammonium oxalate is less adsorbed and coprecipitated than the potassium salt, a result contrary to the adsorption rule of Paneth-Fajans. 3. The presence of an excess of the lattice ion in the solution decreases, but does not necessarily prevent, the adsorption and coprecipitation of ions of the same sign, foreign to the lattice. 4. The coprecipitation of alkali ions with calcium oxalate is larger if there is an excess of oxalate instead of calcium in the solution during the precipitation. An apparent exception to the coprecipitation rule is found if precipitation takes place from relatively concentrated solutions a t room temperature. This abnormal behavior is explained by the fairly rapid transformation of the higher hydrates accompanied by radical structural changes and recrystallization after the precipitation. 5. A study has been made of the coprecipitation of alkali ions under analytical conditions. Sodium, especially, can seriously contaminate the precipitates. 6. By far the purest calcium oxalate is obtained if the predipitation is made a t room temperature from relatively concentrated solutions a t a
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pH of 4 to 6 and the mixture is digested for about twenty hours on the steam bath. 7. The coprecipitation of the alkalies a t 100°C. is much smaller than at room temperature. 8. Coprecipitation of one alkali ion does not prevent the carrying down of another alkali in the same precipitate. 9. The amount of sodium and other alkalies coprecipitated is an exponential function of the alkali concentration in the solution. Relatively speaking, there is already a marked coprecipitation in dilute solutions. REFERENCES (1) (2) (3) (4) (5) (6)
(7) (8)
KOLTHOFF, I. M.: J. Phys. Chem. 36, 860 (1932). J. S., AND TAYLOR, H. S.: J. Phys. Chem. 29,942 (1925). BEEKLEY, O D ~ NSVEN: , Arkiv. Kemi. Mineral. Geol. 7, No. 26 (1920). PINKUS,A,, AND BROUCKBRE, L. DE.: J. chim. phys. 26, 605 (1928). DE BROUCK~RE: J. chim. phys. 26,250 (1929); 27, 543 (1930). BARBER, H. H., AND KOLTHOFF, I. M.: J. Am. Chem. SOC.60, 1625 (1928). DE KONINCK: 2.anal. Chem. 49, 53 (1910). SANDELL, E. B., AND KOLTHOFF, I. M.: J. Phys. Chem. 37, 153 (1933). Compare BALAREW, D.: Kolloidchem. Beihefte 30, 249 (1930).
