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This is to certify that the
thesis entitled
The Preparation and Properties of
Some Scandium Compounds
presented bg
heed Farrar Riley
has been accepted towards fulfillment of the requirements for _P_hs_21_ degree in W
Major prof ssor [7)3te Agril 22 , 1251+
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DATE DUE DATE DUE DATE DUE
6/0I chlFIC/DateDue.p6&p.15 THE PREPARATION AND PROPERTIES
OF SOME SCANDIUM COMPOUNDS
BY
REED FARRAR RILEY
A THESIS
Submitted to the School of Graduate Studies of Michigan State College of Agriculture and Applied Science in partial fulfillment of the requirements for the degree of
DOCTOR OF PHILOSOPHY Department of Chemistry
1954
ACKNOWLEDGMENT
The author wishes to express his gratitude to Professor L. L. Quill, for his considerate and able guidance of this work, and to Marjorie Riley for her patience and aid.
******* ***** ##1‘ * vr Lx I \-I
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TABLE OF CONTENTS
I. Introduction ...... 1
II. Historical ...... 2.
III. Experimental ...... 5
A. Oxalates ...... 6
B. Fluorides ...... 25
C. Trichloroacetates, Acetates and Carbonates ...... 35 Oxides, Hydroxides and Sulfates...... 41
Polarographic Studies of Solutions Containing Scandium Salts ...... 48
IV. Summary ...... 59
V. Bibliography ...... 61 I. Introduction
Since the discovery of scandium in the latter portion of the nineteenth
century, work on this element has been both sporadic and limited, although,
according to V. M. Goldschmidt, scandium "occurs in the earth's crust to at least five parts per million, so that it is as common as arsenic, and nearly twice as common as boron". This neglect has probably two reasons: little commercial use at the present and scarcity of known
scandium ores. Too, since scandium has traditionally been included with the rare earths, chemists may have felt the same reserve in working with scandium.
However, as the lightest transition element and the bridge both between groups III and IV and between aluminum and yttrium within group
III, it should invoke modern chemical interest. The purpose of this thesis then shall be not only to present material gathered as isolated facts, but to correlate them within the framework of group 111. Since the scandium literature has been so well reviewed recently (53) , the organization of this thesis is a little different from the usual. In the section entitled "Historical" the development of scandium chemistry and the relation of scandium to
surrounding elements is reviewed in broad contour. The experimental section is divided into five subsections, each dealing with a separate phase of scandium chemistry. Every subsection is preceded by an analysis of pertinent work in that particular field, followed by the experimental work in this study. The last section sums up work herein completed.
II. Historical
The discovery of scandium in 1879 (45) by Nilson, who recognized it as ekaboron, provided more credence to Mendeleef's periodic arrange- ment of the elements. Much of the early work was done by Nilson and
Cleve. The compounds that scandium forms with a large number of anions, both simple and complex, were investigated around the turn of the century by Crookes. About 1910 much work, including solubilities and conduc- tivities, was done on the simple oxalates and sulfates by Meyer _et_a_l. and
Wirth. The formation of double sulfates and oxalates with group I and group I-like ions was also investigated. Later Sarkar and Urbain prepared more double scandium compounds with different anions, and commented too, as did the earlier authors,on the striking difference in basicity between it and the rare earths. The most modern and still extensive studies on scandium were made by the Sterba-Bbhms, who, among other things, studied simple and double oxalates and carbonates, the formates and acetates. More recent studies on the structure of scandium compounds, spot tests, etc. , have been made more for rounding out a series of measurements than for elucidating the chemistry of scandium.
The two characteristics, then, which loom large in any discussion of scandium are its weakly basic character and its tendency to form double salts. These are just the properties which have allowed it to be unjustly classified with the lanthanides and yttrium, because in the rare earth series, double salt formation is prevalent, and the yttrium group elements have distinctly more acidic properties. A number of researchers have pointed out that scandium differs from the rare earths (59, 67, 73, 74). Sarkar
(59), in a rather complete comparison of the rare earths and scandium, \J’ -3- has detailed these differences. If their fluorides are compared, one
sees, first of all, that scandium fluoride is insoluble in hydrochloric acid, whereas, in the same medium the rare earth element fluorides are relatively soluble . As a matter of fact, their solubility increases, in the cerium group at least, with decreasing basicity (increasing atomic weight) (41). More striking than the latter is the large solubility of scandium fluoride in excess of ammonium fluoride and the formation of well defined complex fluorides with ammonium and alkali metal ions (41,
65).
The acetylacetonates of the rare earths are not volatile as is scandium acetylacetonate. Further, the latter exists in organic solvents as a monomer, while the rare earth acetylacetonates may be dimeric in the same media.
With respect to the rigorous isomorphism found in the compounds of the rare earths, one finds no scandium salt isomorphous with the analogous rare earth compound. Covalent and ionic compounds of scandium are isomorphous with analogous compounds of Fe(III), C0(III), Cr(III), Al(III) and often Ga(III), In(III) and V(III). (59)
If we turn next to the carbonates, we find well developed normal salts for rare earth ions whose ionic radii areas small as yttrium (27). To the contrary, scandium probably does not form either a normal carbonate or basic carbonate (68).
Scandium sulfate is a very soluble compound and hence is different from the slightly soluble rare earth sulfates. The temperature coefficient of solubility of the scandium compound is positive, whereas, for the rare earth sulfates, the coefficients are negative. The double sulfate of scandium
I -4- being water insoluble (59) reminds one of the cerium group, but placement here is ruled out on the basis of ion size.
With the exception of ceric oxide, rare earth oxides dissolve comparatively easily in mineral acids, but scandium oxide is more refractory and solution takes place slowly even in boiling, concentrated acids. .
Lastly, one may mention some solubility anomalies. Scandium oxalate is slightly soluble in neutral and dilute acid solutions while the rare earth oxalates are known to be quantitatively insoluble under these conditions (64, 84). Recent studies have shown that the high solubility of scandium chloride in HCl-saturated: water, ethyl alcohol-water, and ether-water mixtures permits a separation of scandium from aluminum, yttrium and the lanthanides whose chlorides are insoluble in these media (I).
The similarity of the properties of scandium and aluminum is more striking than is generally pointed out. The fact that aluminum also forms an acetylacetonate, covalent enough to be distilled and monomeric in organic solvents, should be mentioned (62). Aluminum oxide, like scandium oxide, is quite refractory towards mineral acids. The isomorphism of their compounds has already been stated. In addition, Weiser and
Milligan (77) have pointed out the similarity between the structure of
Y-AlOCOH and hydrated scandium oxide. The similarity of their fluoride complexes is another consideration. However, it is puzzling that the strong amphoterism of aluminum hydroxide does not appear with scandium hydroxide.
Aluminum carbonate has never been isolated, analogous to the behavior of scandium (4). Too, aluminum sulfate is soluble although strongly hydrolyzed. As a last point, one may point out the formal analogy presented
-5- by the compositions of Sc c and A14C3 (20, 62). 4 3 In some ways one may think of scandium as a bridging element between the third and the fourth periodic groups. The analogies, naturally, are not as numerous as within group III, but some are: Zirconium and
scandium sulfates are soluble; zirconium, thorium and scandium form many basic and double salts; thorium and scandium fluorides are insoluble in HCl (62). Again, thorium forms an 8-quinolinolate with one molecule of 8-quinolinol of solvation (17) as does scandium oxinate (54). Wirth reports that the low solubility of scandium oxalate in concentrated sulfuric acid solutions (4. 32 N) is more like thorium oxalate than analogous rare earth compounds (84).
III. Experimental
The scandium used in this study was part of the stock at Michigan State
College. It was known to be free of all metals except a small (l-3%) amount of calcium, which was removed by precipitating the mixture with ammonium fluoride and then leaching out the scandium with excess of ammonium fluoride. Under these conditions, scandium fluoride forms a soluble complex fluoride while the calcium salt remains unchanged. The double fluoride solution was filtered, evaporated to dryness in platinum dishes, and the fluorides converted to sulfates by evaporation with concentrated sulfuric acid. The ammonium sulfate was then volatilized and the residue dissolved in boiling water, precipitated as oxalate and converted to oxide by ignition at 850°.
Because of the scarcity of scandium all residues were carefully saved and regularly recovered. The qualitative test using conchineal (53) was a great help in determining whether solutions contained a worth while amount.
-6-
Recoveries were usually made in two steps, and, if the contaminants were plentiful, the process was repeated. The first step was to precipitate hydrous scandium oxide at a pH of about 7 using a pyridine- nitric acid buffer (51), then to ignite the precipitate to the oxide. Step 2 involved redissolving the oxide, precipitating basic scandium tartrate with ammonium tartrate and ammonia (l8), and igniting the scandium tartrate to oxide.
Scandium analyses in the oxalate study were made by igniting the oxalate and weighing as scandium oxide (18). However, ignition was carried out at 850° instead of the 900° temperature recommended, because it was found that the oxide, when produced in this manner, dissolved more easily in mineral acids. Ignition at 800° netted unreproducible results.
Oxalate analyses were made by the permanganate method (80). Scandium trichloroacetate was analyzed for scandium by precipitation of the 8- quinolinolate (55) and. for chlorine by the method of Umhoefer (72). Ammonia analyses were run by the Kjeldahl method.
Powder diffraction photographs referred to later in this thesis were made using Cu K0. radiation. The camera diameter was 114. 59 mm. No film shrinkage correction was made for powder diagrams used for identification purposes.
A. Oxalates of Scandium
In the literature pertaining to the normal oxalate of scandium, one finds disagreement on the exact value of the hydrate formed by precipitation.
Nilson (46), basing his conclusions on a complete analysis, designated the hydrate as a hexahydrate. Crookes (8, 9), on the basis of a partial analysis, referred to mono, di, tri, and penta-hydrated oxalates.
I .Vllll -7-
Meyer and Winter (40) assigned the precipitated scandium oxalate a
tetrahydrate value, while Wirth (83), precipitating the oxalate using
oxalic acid or ammonium oxalate in slightly acid solution, found a
pentahydrate. In the study made by Sterba-B’dhm and Skramovsky (66) ,
who precipitated the salt from slightly acid solution at 60° and allowed it
to dry in air for a day, a six-hydrate was found. In a recent study, which
involved precipitating the oxalate at 60° from a dilute scandium chloride
solution and air drying it two days protected from ammonia vapors, Klein
and Bernays (29) observed that the precipitate was non-homogeneous.
However, after the precipitate was dried over anhydrous calcium chloride,
a complete analysis showed. the hexahydrate was formed after a drying
period of two days.
Although some careful optical studies have been made with the hydrated
.rare earth oxalates (21, 85), the only measurements with hydrated scandium
oxalate are incomplete and conflicting. Sterba-BBhrn (64), in 1914, allowed
a saturated oxalate solution to crystallize very slowly at ordinary
temperature from which he got hexagonal quartz -like crystals (C51. 2484);
combination P (lOIl) withd’P (lOIO). The angle between (lOIl) and (10l0)
was reported as 34° 45‘. Other properties noted were that the PP(10T0)
faces are almost always striated and P(10Il) is usually very smooth and
shiny. In 1929 Sterba-Bma and Skramovsky (66) reported that the six-
hydrate crystallizes as plate -like microscopic rhombs or prisms which
are then able to attain macroscopic dimensions from a very slow recrys-
tallization. Few other properties for the six-hydrate are recorded. The
above authors found that it has a density of 2.14 and that it is not altered
in ordinary air, but, above anhydrous calcium chloride, it effloresces CI
",
U \J -8- after a certain period of time. Also,it is stated that drying at 100° forces the oxalate to lose four molecules of water forming the twins hydrate-
Wirth (84), on the other hand, stated that the monohydrate is formed at
100°. He began with the pentahydrate. Lastly, Dupuis and Duval reported that a brief horizontal stretch in the pyrolysis curve of freshly precipitated scandium oxalate between 67° and 89° corresponds to scandium oxalate decahydrate (14). Their results also indicate that at 277° no more than five molecules of combined water remain, but state it is impossible to discern the end of dehydration.
No normal oxalate of aluminum is formed, but the hydrated rare earth oxalates are both numerous and well known (62). In the cerium group the decahydrate is the most common, but in the yttrium group more diverse results have been obtained. Marsh (38) states that yttrium forms the six-, ten- and seventeen-hydrate, ytterbium forms the six-and seven-hydrate, dysprosium forms the ten—hydrate, and erbium forms the six-, nine-, ten-, twelve-, and fourteen-hydrate. The lower hydrates are formed at higher precipitation temperatures and the higher hydrates at lower temperatures. Wylie (85) found atwo- ‘31:] six-hydrate for lanthanum and yttrium oxalate and interstitial hydrates for lanthanum, yttrium and cerium over a composition range of nine to twelve moles of water per mole of oxalate. These oxalates react rapidly with cold 15% sodium hydroxide presumably to give the corresponding hydroxide.
Wylie, in the same paper, made dehydration and optical studies of the oxalates. The dehydration of the six-hydrates of both yttrium and lanthanum takes place rapidly at 180° forming the two-hydrates. It has been reported that anhydrous lanthanum oxalate is formed at 220° (21).
_9-
In the optical studies Wylie found the interstitial yttrium, lanthanum and cerium oxalates to be monoclinic and to demonstrate very similar optical properties. Gilpin has given a complete set of optical data for lanthanum oxalate decahydrate (21). Of more interest in this study are the optical pr0perties of yttr ium oxalate six-hydrate (85). The crystals are monoclinic, often pinicoidal parallel to the c axis and flattened parallel to the (110) plane. Optically negative, X along b, and Zj’c was equal to 40'. 2V was small. Refractive indices were as follows: a 1.47;
8 1. 61; ’1. 62.
The only investigation of the reaction between hydrated scandium oxalate and moist ammonia gas is by Sterba-B'dhm and Skramovsky (66) .
The reaction is quite slow, requiring several weeks for completion, and a complete analysis of the residue indicates the new compound fits the formula
(NH4)Sc2(CZO4)3- 8H20. However, the authors say this formula must be taken with some reservation.
Dehydration of hydrated oxalates
For the dehydration studies reported in this thesis the oxalates were prepared by a homogeneous precipitation method involving dimethyl oxalate
(22, 81). Solutions of the sulfate, nitrate and chloride, each containing the equivalent of two grams of scandium oxide, were made 0. 025111 in the appr0priate acid, and 5.67 g. of dimethyl oxalate (10% excess) was added.
The volumes were adjusted so the final volume was about 170 ml. The temperature of the scandium solutions was maintained at 70°:5°, as dimethyl oxalate solution was added via a dropping funnel at l to 3 drops per second.
The heating was continued for one hour after precipitation incidence; the solutions were then cooled and filtered. Ten m1. of saturated oxalic acid
-10- was finally added to make sure no scandium available for precipitation was left, but this procedure proved unnecessary except in the case of the sulfate solution. In the latter, a small amount of oxalate precipitated over an eighteen hour interval. The reason for this unequivalent behavior is unknown. After precipitation and addition of the oxalic acid, the pH values of all the supernatant solutions were about 0. 50 pH. The crystals were large and very granular in gross aspect as others have observed in similar cases.
The percentage recovery for homogeneous oxalate precipitations from nitrate and chloride solutions was evaluated and found to be 95-97% of that expected for a quantitative reaction. The percentage recovery was calculated by adding the weights of scandium oxide produced in analyses of the precipitated oxalate to the oxide produced by ignition of the unused scandium oxalate.) The ratio of this combined weight to the starting weight of oxide (2 g.) multiplied by 100 is the percentage recovery. The error in this latter value is estimated to be: 1%.
Originally, preliminary dehydration over calcium chloride was carried out for a two day period in order to obtain the’six-hydrate, for it is convenient to have a reference point. However, although the formula of the hydrate corresponded approximately to six after this time, the water content was decreasing rapidly. This continuing water loss suggested that drying studies over longer periods of time and with desiccants of different power should be carried out. In this and in all drying experiments involving desiccants, the oxalate was simply placed in a desiccator over a suitable desiccant and then periodically removed and capped preliminary to either weighing or analysis. After a compound had come to constant weight, an
-11- extra portion of desiccant was added and the drying period extended one week.
The results in Table 1 show that freshly precipitated scandium oxalate, when it is dried by 98% H2504 or Mg(ClO4)2 for several months, forms the dihydrate. Figure 1 shows that, of the two desiccants, only anhydrous magnesium per-chlorate caused dehydration to the pure dihydrate, although the oxalate which was dried over sulfuric acid had come to constant weight after 3. 5 months. Table 1 shows it analysis is slightly different from the pure dihydrate. Figure 1 shows the least efficient desiccant used, anhydrous calcium chloride, had not dried the oxalate to constant weight after three months. Because this desiccant is less efficient than 98% H2504, one assumes it is also unable to form the dihydrate. In Table l the abscissa,
"moles oxalate", refers to moles anhydrous oxalate.
Equilibrium vapor pressures taken from the International Critical Tables
(89) for the three desiccants are as follows: CaCl2 0. 34 mm. , 98% H2804 5.6x10-4mm., Mg(ClO 4’2 4.1xio’ 5 mm. In Table l the analyses corresponding to shorter drying periods over calcium chloride show that a hexahydrate is formed after approximately two days. However, further perusal of Table 1 shows that this hexahydrate is then rapidly dehydrated. Figure 1 gives the dehydration curves of freshly precipitated oxalate over three desiccants. The continuous change in slope and single plateau seen in each of these curves points to the dihydrate as the only stable phase under these conditions.
-12-
Table 1
Results of the Dehydration of Hydrated Scandium Oxalate
Analyses* Drying %: % Mole ratio : Mole ratio Pptg. Time Desmcant C204 SCZO3 SCZO3/C204 HZO/SCZO3 Soln.
2 days CaCl2 57.16 30. 06 l. 007/3 6. 3 Nitrate 6. 5 days " 61. 34 32. 49 1. 014/3 4. 3 Nitrate 3. 5 months H2504 66.71 34.93 1.003/3 2.4 Nitrate 2. 5 months Mg(C104)Z 67.68 35.26 0.998/3 2.0 Nitrate 2 days CaCl2 57.67 30.29 1.006/3 6.0 Chloride 4-5 days " 61. 30 31.73 0.990/3 4.6 Chloride 2 days .. 56. 86** ' Sulfate 3 days " 57. 72** Sulfate 10 days " 61. 62*“I 32. 45** 4. 3 Sulfate
C204 % SCZO3 '70
Theoretical for SCZ(C2O4)3°6HZO 57.14 29.85
Theoretical for Sc2(CZO4)3- ZHZO 67. 70 35. 36
*All analyses except otherwise noted are the average of two determinations.
**Only one analysis made.
I I .l .l 3k A \I .P l ::<
1 IOO
(987.) 42 75
36590423
H2504
MgCCIO) ColClZ
DESICCANTS
l
I so _ (DAYS)
HYDRATED
or
.FIGURE TIME 25
DEHYDRATION
31V‘IVXO 310w / OZH snow -14-
A preliminary thermal dehydration study on a partially dehydrated six-hydrate was carried out. Originally, it gave the following analysis:
C20: , 61. 28%; 50203, 32.49%. The results showed that 56. 5 hours at
80-81° gave a loss of weight corresponding to the formation of a two -hydrate.
To confirm this, hexahydrate samples were placed in a drying oven which was controllable to :1“. The experimental procedure was to regulate the temperature to a determined value and to allow the sample to come to constant weight. Then the temperature was increased to the next level, and the sample again brought to constant weight. This procedure was repeated as many times as was necessary. Since the composition of the original sample was known accurately, analysis was made only at the end of the run. The result of two such dehydration runs is given graphically in Figure 2.
The analyses of the dihydrate products are given in Table 2.
Table 2
Thermal Dehydration of Scandium Oxalate Hexahydrate
Drying Analyses Treatment Per cent Per cent Mole Ratio Mole Ratio
C204 ScZO3 HZO/SC203 SCZO3/CZO4
Temps. to 204° 67.48 35. 64 2. 0 1.011/3
Temps to 155° 67.64 35.19 2.1 0.996/3
Theoretical for Scz(CzO4)3 .ZHZO; C204, 67.70%; 56203, 35.36%.
.-15-
From Table 2 one sees that at temperatures to 204° the hexahydrate is thermally dehydrated to the dihydrate. Only at the higher temperatures is there some slight decomposition of the oxalate ion. This is evidenced by [SczO3]/[CZOZ] being larger than 1:3 in the sample heated to 204°.
Figure 2 points up the large temperature range of stability of the dihydrate.
The density of the twoz-hydrate prepared by drying at 155° was measured and found to be 2. 26 g. /ml. Drying rapidly by thermal means or drying slowly over a desiccant gives dihydrates having the same structure, as revealed by the similarity of the powder diffraction lines in Table 3.
The intensities of the lines in column 2 of Table 3 were generally less than those of the other column. The seventeen lines appearing in column 2 may all be matched to corresponding lines in column 11, Since exposure times of the two films were identical, one may say that the more intense pattern (column 1); is due to an ordering of the dihydrate at the higher temperature. There is no indication of any oxide lines in either column as one would expect from the analyses (Tables 1 and 2). The diagrams indicate a substance of lower symmetry than cubic or hexagonal. FIGURE 2 THERMAL DEHYDRATION 0F HYDRAT ED SC (C 0) Z Z 4- 3
200— 0 0 0
(D 150- 0 0 0 - 0 I00- 0 i 0 '1
m - ' U 50L- I I i
4 3 2 MOLES Hp / MOLE OXALATE -17..
Table 3
S * Distances ' for Sc2(CZO4)3 . ZHZO
SCZ(CZO4)3- ZHZO SC2(C204)3. ZHZO
(204° dried) ‘ (H2504 dried)
1.22 1.28 1.28 1.55 1.55 1.60 1.60 1.93 1.92 2.02 2.01 2.24 2.55 2.55 2. 63 2.87 2.87 2.97 2.97 3.22 3.21
3. 33 3.42 3. 51 3.60 3.60 3. 71 3.79 3.80 3.88 3.99 3.99 4.07 4.07 4.13 4.12 4.21 4.21 4.31 4.31 4.42 4.43
.------_---—------——--—-—------—------fl----- *5 is one -half the distance between corresponding lines in centimeters.
-t'
iii .a -18_
Optical studies on freshly precipitated scandium oxalate.
To further determine the phase which is precipitated when oxalic acid is added to scandium solutions, optical measurements of the precipitates were made. Solutions of oxalic acid and scandium chloride or nitrate, all in 1% strength, were added dropwise to acidified water, held at 50°, over a period of an hour or two. This method, which had been successful in the preparation of rare earth oxalates (21, 85), gave large well—formed crystals of hydrated scandium oxalate. Regardless of the modifications made in temperature, rate of addition, pH, and concentration of reagents, flat, platelike crystals were always obtained. Attempts to recrystallize the oxalate in contact with the saturated solution (neutral or 0. 33 Min acid) at the boiling point, lead to more poorly formed crystals. Larger and thicker rhombs were crystallized by cooling a saturated solution, but these crystals were so badly twinned identification of the faces was impossible.
All attempts to roll the flat crystals into other views under the microscope were unsuccessful. Thus, one has to be satisfied with a limited number of optical measurements.
A crystal of hydrated scandium oxalate together with its optical properties is shown in Figure 3. -19-
Figure 3 l
Crystal: scandium oxalate hexahydrate
Crystal System: monoclinic, optically negative
Refractive Indices (5893A;25°) : 6.1.495; 61.60;; {1.63
Extinction : 43° from KB—
2 V: small
If one allows the hydrated oxalate to stand for a time, the crystals
become opaque but do not lose their outline. This was not investigated
further. To analyze the freshly precipitated oxalate two different techniques
were tried. The first was to dry the crystals for a short time on a Buchner
funnel. The other was to dry the crystals for a short time on the funnel
and then to wash them with alcohol and ether. The oxalate is not soluble
in ethyl alcohol, so one would not expect solvation of the oxalate by alcohol.
Simply sucking the precipitate dry on the Buchner funnel did not lead to the hexahydrate, so all precipitates were dried with alcohol and ether. One
-20- precipitation was made at 100° to check .whether the same hydrate was precipitated over a wide temperature range. Thelexperimental procedure was to allow 3% solutions of oxalic acid and scandium nitrate to drop slowly into a solution of the appropriate pH and temperature. The data in Table 4
show that the six-hydrate is the phase precipitated in temperature range
50—95°. One sees, further, that increasing the acidity of the precipitating
medium from pH 1. 5 to pH 0. 8 decreases the water content of the oxalate.
In this case, the precipitated oxalate lies between a penta- and a hexahydrate.
The last preparation in Table 4 shows the precipitate, obtained from a
solution to which no acid was added, contained 18 m61es of water per mole of scandium oxalate. This precipitate was dried only half as long as the other, so the amount of water held chemically is unknown. However, if the
amount of water held is actually dependent upon the acid concentration of the
medium, this would explain the conflicting results obtained in other studies. -21-
Table 4
Compositions of Freshly Precipitated Scandium Oxalates
Analyses percent percent Mole Ratio Pptg. Procedure Drying Procedure C204 ScZO3 [58203]”(32041
Solns. at 45° and Suction for 30 min. 56. 62 29. 55 0.999/3 pH of 1. 5 then alc. and ether
Solns.at 50° and 1. Suction 20 min. 29.65 pH of 1. 5 then alc. and ether 2. Suction 30 min. 56. 20 29. 23 0.996/3 3. Suction 30 min. 29. 59 then alc. and ether
Solns. at 50° and 1. Suction 30 min. 31.02 pH of O. 8 2. Suction 15 min. 30. 93 then alc. and ether
5% H2C204 and Suction 15 min. 20. 4 Sc(NO3)3' added dropwise to 50 ml. H20 at 50
Theor. for Sc2(CzO 4)3'6HZO 57. 14 29.85
H H SC2(C 2O4)3-5HZO 59.49 - 31.02
-22.. Solid gas reaction of hydrated scandium oxalate and ammonia
Material corresponding to approximately scandium oxalate four- hydrate was exposed in a vacuum system to an atmosphere of dry ammonia
gas, and the weight increase due to ammonia sorption measured. The
results in Table 5 show that the ammonia is sorbed slowly.
Table 5 Sorption of Dry NH3 by Hydrated Scandium Oxalate
Absorption time Moles NH gained Desorption time Moles NH lost per mole 3oxalate ' per mole oxalate (hours) (minutes)
1 0. 0737 25 0. 118 12 0. 166 40 0. 132 25 0. 240 60 0. 139 29.5 0.284 203 0.254 35. 5 0. 341 40 0. 341 58. 5 0. 369
If this ammonia were simply held by surface adsorption, one would expect
it to be desorbed rapidly. However, the desorption studies in Table 5
show after 203 minutes a weight corresponding to 0.115 moles NH3 was
retained and the weight loss during the first 25 minutes was very rapid.
These results may be best explained as follows. During the first 25 minutes,
surface adsorbed ammonia and some hydrated water were removed. The
weight loss after this time correSponds to hydrated water being pumped out.
The desorption was effected by pumping out the system with a Cenco "HyvacH
vacuum pump.
Thus, the results indicate ammonia, dried by passage over barium
oxide, was absorbed by the hydrated oxalate, forming some type of compound.
Discussion of the nature of this compound is reserved until the results with -23-
Table 6
5* Distances
SC2(CZO4)3.6HZO (NH4)2C204.H20 SC2(CZO4)3'6HZO + NH3 exposure exposure exposure 1 hr. 2 days 6 days 0.98 0.58 1.40 0.64 1.44 0.79 1.51 0.99 1.62 1.08 1.90 1.14 2.10 1.39 1.38 1.38 1.39 2.36 1.48 2.42 1.59 2.64 1.72 1.72 1.72 1.72 2.76 1.86 2.83 1.94 1.94 2.96 2.05 2.03 3.01 2.13 3.09 2.34 2.34 2.34 2.33 3.18 2.49 2.49 2.48 3.31 2.53 3.48 2.59 2.59 2.58 3.66 2.73 2.73 2.73 2.72 3.79 2.92 2.91 2.92 2.90 3.89 3.13 3.12 3.13 3.11 3.96 3.29 3.30 4.04 3.36 3.36 3.36 3.35 4.09 3.45 3.45 3.45 3.44 4.17 3.54 4.27 3.65 3.64 4.36 3.69 3.68 3.70 4.47 3.72 4.52 3.78 4.60 3.82 3.80 4.66 4.04 4.74 4.13 4.82 4.21 4.21 4.21 4.20 4.91 4.38 5.03 4.44 5.16 4.49 4.49 4.49 4.48 5.26 ' 5.46 4.86 5.25 5.56 - 4.96 6.00
*S is one -half the distance between corresponding lines in cm. ."\
-24- moist ammonia gas are given.
When hydrated scandium oxalate on a microscope slide is exposed to 40% ammonia water in a desiccator, the crystal structure of the oxalate
completely disappears in less than 10 minutes indicating a very rapid
reaction. When exposed for longer times, clusters of acicular crystals
are noticed upon microscopic (examination of the material. X-ray
examinations were used to determine the changes in the reaction products.
The data are tabulated in Table 6 in which the S distances for ammonium oxalate monohydrate are given in the first column, for scandium oxalate hexahydrate in the second column, and the reaction products in the
remaining columns. It is noted that lines due to ammonium oxalate monohydrate were found in each of the diagrams for ammonia-treated
scandium oxalate.
The rapidity of the reaction of moist ammonia with hydrated scandium oxalate is shown by the absence of any diffraction lines due to the hexa— hydrate in column 3 of Table 6. In this same column one sees many
ammonium oxalate lines and a few new lines. Columns 4 and 5 show
ammonium oxalate lines, but the new lines have become far more numerous.
The reason for the delay in appearance of these latter lines is unknown.
These data lead one to postulate the following reaction for the action of moist ammonia on scandium oxalate hexahydrate.
Sc2(CZO4)3-6HZO + 2xNH3 + 3HZO—-) SC2(OH)2x(c204)3-x + x(NH4)2CZO4
In an attempt to find whether this reaction is completed to the right a
sample of the hexahydrate was allowed to stand over moist ammonia until
it reached constant weight (6 days). From time to time the oxalate had to
be recrushed because of the hard crust of reaction products which formed
-25- on its surface. Analyses showed the resultant compound to contain
12. 83% NH3 and 20. 76% ScZO3 corresponding to 5 moles of NH absorbed 3 per mole of ScZO3. However, the increase in weight of the exposed
sample indicated that roughly 11. 5 moles NH3/mole of oxalate were absorbed. Thus, 6. 5 moles of water must have been taken up in the reaction. No way to separate the reaction products could be devised,
so the formula of the scandium product could not be determined. Powder
'diffraction patterns taken after the reaction had proceeded for 6 days
showed no lines due to hydrated scandium oxide. 50 one tentatively assigns the lines not belonging to ammonium oxalate to basic oxalate intermediates and double salts. The latter are known to be very stable even in the presence of boiling, concentrated NaOH, so probably a certain percentage of the normal oxalate leaves the field of reaction without reacting with ammonia.
B. Fluorides
Scandium fluoride was first described by W. Crookes (7) who formulated it as the dimer Sc2F6. Meyer_e_t__al_. (41) later showed that scandium fluoride dissolved in excess alkali metal or ammonium fluoride.
Double fluorides of the type M3ScF6 are deposited on standing by the latter solutions. Sterba—B'dhm (65) pointed out that with ammonium fluoride an entire series of double fluorides may be produced and that the composition of the precipitated phase depends on the fluoride ion concentration in solution.
Complex fluorides of iron and aluminum are well known whereas complex rare earth fluorides are unknown (62). Aluminum and scandium trifluorides crystallize with. the same structure (78). The rare earth
-26- fluorides are rhombohedral, and a few dimorphic rare earth fluorides
are both rhombohedral and hexagonal (88). Scandium trifluoride has a
rhombohedral, pseudo-cubic structure with a = 4. 022, and u equal to
89° 34. 5', and has been described as an idealized W03 structure (49).
Recently much interest has been shown in the oxyfluorides of the
rare earths. Hund (24) points out that the simple fluorite structure should
exist for all those three-valent oxyfluorides whose cation to anion ratio is greater than 0. 73 because of the similarity in the size of the fluoride
and oxide ion. This has been amply shown (87, 88, 30, 16, 24, 25).
Zachariasen (87) has said, "since lanthanum and yttrium oxyfluoride are isostructural, all the 4f elements will be structurally analogous. And the fact that actinium and yttrium oxyfluoride are isostructural, actinium being the largest of the 5f elements, makes it possible to predict with assurance that oxyfluorides of all 5f elements will be isostructural. "
Zachariasen has found lanthanum and yttrium oxyfluorides dimorphic, tetragonal and rhombohedral. Actinium oxyfluoride is cubic. The preparation of these oxyfluorides involves four methods: a) melting the metal oxide with the corresponding fluoride in the correct stoichiometric
ratio, b) controlling the hydrolysis reaction of the trifluoride, c) precipi- tating the oxyfluoride using sodium fluoride and hydrofluoric acid, and d) changing the four-valent oxide to the three-valent oxyfluoride using fluorine.
Lanthanum and yttrium oxyfluoride have been produced by the first and
second methods (24, 87, 30). Actinium oxyfluoride has been produced by the second method (87). Cerium oxyfluoride has been produced by the first and the fourth method but not by the third (16) . Holmium oxyfluoride has been produced by the third method. It is of interest that dysprosium
-27- fluoride goes to the oxide without an oxyfluoride intermediate in ten hours at 500°. Naeser (44) has shown that neodymium oxyfluoride is produced by the hydrolysis of neodymium fluoride in air. The oxyfluoride of aluminum has never been isolated. Aluminum fluoride must be extremely inert, since water at 400° has only a very slow effect (62).
The end product for the hydrolysis reaction of the fluorides, reported so far, is the oxide. The rate and temperatures at which the oxide and oxyfluoride are produced appear to be extremely variable. For instance,
Zachariasen (87) reports actinium oxyfluoride is produced at 1200°, lanthanum oxyfluoride at 920°, and lanthanum oxide at 1130'. On the other hand, he reports yttrium oxyfluoride is formed at 500° and yttrium oxide at 900°. Naeser (44) reports neodymium oxyfluoride is made at
750° and 900° and the oxide at 1025°. It is reported that cerium oxy- fluoride hydrolyses to the dioxide at 300°, but the oxidation process complicates the situation (16). The hydrolysis of dysprosium fluoride has been mentioned previously.
The ionic radii of aluminum and scandium are 0. 57 and 0. 81 respectively
(42). Compared with the ionic radius of the oxide ion, 1. 40A, one finds cation to anion radius ratios of 0.41 and 0. 59, respectively, which are appreciably less than the lower boundary of 0. 73 for the fluorite structure.
Accordingly, the ability of scandium and aluminum to form oxyfluorides and the structures of these oxyfluorides becomes interesting.
A few mixed fluorides of the rare earths have been isolated. Zachariasen
(86) reports hexagonal sodium tetrafluorolanthanate and cubic and hexagonal potassium tetrafluorolanthanate. Nowacki (48) reports sodium tetrafluorolanthanate, and Hund (25) has made cubic sodium tetrafluoroyttriate.
Cubic ammonium decafluoroerbiate, -holmiate, and -thuliate are also reported.
(88). \_-1 -28-
In the search to produce some double sodium and potassium scandium fluorides by mixing them in the correct stoichiometric proportions and fusing, it was found that heating of these mixtures at 1050° for periods of one to three hours produced mostly scandium oxide. Because of this observation it became of interest to see whether or not the hydrolysis of scandium fluoride produced the oxyfluoride as an intermediate.
The first experiments involved mixing scandium oxide and scandium fluoride in a 1:1 mole ratio and heating by means of an oxy-hydrogen torch in an open platinum dish. The times of heating varied from five to twenty minutes. The diffraction patterns of these ignited materials showed that most of the product was scandium oxide, but some feeble scandium fluoride reflections were noted. Since hydrolysis to the oxide was proceeding too rapidly, we decided to try to control the reaction by heating scandium fluoride alone at lower temperatures.
The scandium fluoride was prepared by dropping dilute solutions of a scandium salt and of hydrofluoric acid into water held at 70°. All polythene equipment was used. After addition was complete, the fluoride was allowed to digest at 90° for varying lengths of time (up to one week). However, this did not appear to noticeably increase the particle size. After digestion, the fluoride was filtered and washed. This last process was difficult, because the fluoride was very gelatinous and easily peptized by washing.
This difficulty was circumvented by fewer washings and by accepting lower yields. Diffraction patterns were made for each batch of fluoride, after drying at 110° overnight, to insure uniform purity.
The first hydrolysis run was made at 550° and the loss in weight of the fluoride sample was followed by using a second sample alongside the first in the furnace. -29-
Table 7
Rate of Hydrolysis of Scandium Fluoride
Temperature Heating Time’l‘ % Conversion to SCZO3 (hrs.)
460 2 0.0 570 2 4.6 570 2.5 6.0 570 11.5 31.9 570 15.5 43.2 570 19.8 54.8 570 23.8 63.2 570 31.8 87.3 570 36.0 97.7 600 3 99 600 5 100 600 6 100
*Heating time is total time at that temperature.
Table 8 tabulates the powder diffraction lines of hydrolysed scandium fluoride together with those of pure scandium fluoride and oxide. Column
1 gives the lines due to pure scandium fluoride. Tabulated in columns 2, 3 and 4 are the lines for partially and fully hydrolysed scandium fluoride.
In the last column, the diffraction lines of pure scandium oxide are given.
Inspection of the table from left to right shows that all but three scandium fluoride reflections show up slightly displaced in the scandium oxide pattern.
Also, these displacements are about equal and occur gradually as scandium fluoride is hydrolysed. These facts lead one to conclude that a portion of the scandium atoms in SC203 have nearly the same relative positions as in
ScF3. \" . t x
o
.
. O
- ' I u
- I
' V o
. o \
. u
I V
\‘Jfis . a \ ,
.
| o A ' -‘ , I r
. ' - A
o
. . . ,
T A l
r . , . .
o .,
-30-
Table 8
Powder Diffraction Lines of Partially and Fully Hydrolysed ScF3
ScF3 ScZO +Heating time (hrs.) 23. 8 31. 8 Constant wt.
% Conversion to SCZO3 63. 2 87. 4 100
*S distances (cm.) 1.74 1.73 2.00 2.22 2.22 2.20 2.20 2.19 2.43 2.44 2.83 2.82 2.83 3.16 3.14 3.13 3.13 3.13 3.47 3.46 3.64 3.64 3.63 3.634 3.89 3.88 3.87 3.86 3.86 4.06 4.08 4.08 4.08 4.17 4.31 4.29 4.29 4.30 14.52 4.52 4.65 4.70 4.69 4.69 5.09 5.09 5.06 5.06 5.25 5.25 5.24 5.24 5.41 5.42 5.42 5.61 5.61 5.60 5. 58 5. 59 5.76 5.76 5.74 5.75 5.92 5.91 5.91 6.08 6.04 6.07 6.08 6.25 6.24 6.24 6.24 6.40 6.40 6.39 6.39 6.58 6.57 6.55 6.55 6.56 6.69 6.71 6.85 6.86 7.04 7.03 7.01 7.00 7.01 7.13 7.14 7.48 7.48 7.60 7.60 7.57 7.58 7.74 7.74 7.78 7.73 7.90 7.88 7.86 7.87 8.02 8.01 8.09 8.33
+Hydrolysi 8 study was made at 570°. *5 distances are the distances between corresponding powder diagram lines. ca -31-
One sees from Table 8 that no separate and distinct ScOF phase appears, but that six new lines do appear which can not be attributed to either the oxide or fluoride.
In an effort to isolate, or at least to form a larger amount of the phase belonging to the extraneous diffraction lines, hydrolysis runs were made at several different temperatures. Also, diffraction patterns were taken at more frequent intervals. At 500°, the rate of hydrolysis is substantially slower, but only two of the new lines appear and these have completely disappeared after 40% conversion to SCZO3' At 400° the hydrolysis rate was proceeding so slowly that this run was discontinued.
At 750° the reaction proceeded so rapidly that 80% conversion to SCZO3 was found after heating only 97 minutes. As at 570°, only six lines of the new phase can be detected. The manner in which these lines appear is illustrated in Table 9
Table 9.
Appearance of New Diffraction Lines during Hydrolysis of ScF3
*Heating time (min.) 20 40 70 90
% Conversion to SczO3 28.1 41.5 61.6 75.5
S distances (cm.) 1.74 1.73 l. 99 2.01 2.44 2.44 2.83 2.83 2.82 2.83 - 3.47 3.46 4.17
*Hydrolysis study was made at 750°.
—32-
Again, one sees that the new lines appear in the range 60-80%
conversion. Theoretically, 66% conversion of a scandium fluoride
sample to scandium oxide corresponds to scandium oxyfluoride, so
assignment of these lines to a scandium oxyfluoride phase appearing in
small amount, is tentatively made.
When the hydrolysis reaction rate was plotted versus time for the
reactions proceeding at 570° and 750°, it was found that the rate is
independent of the concentration of scandium fluoride present. This is
exactly what one would expect for a reaction controlled by the rate of
diffusion of H20 vapor to the solid surface. At 500°, on the other hand,
the rate is linear with time only up to 75% conversiOn, whereupon it
slows, indicating that some other factor is coming into play. The rates
of the hydrolysis reaction are tabulated in Table 10.
Table 10
Temperature Dependence of the Rate of Hydrolysis of ScF3
Temperature Rate (°C.) (% Conversion/hr.)
400 0.17 500 0. 27 570 0. 50 750 40. 3
Since little aid was obtained by using different temperatures in order to find a point where the temperature dependence of the hydrolysis rate to ScOF would be favorable, it was decided that conversion to ScOF might be favored by slowing the hydrolysis rate down. Scandium fluoride was placed in a tube furnace at 780° and oxygen, dried by a pair of magnesium perchlorate towers, passed over it. Originally, the sample was allowed
-33- to stand one hour. When a diffraction pattern was obtained, it showed no scandium oxide lines but all of the lines hitherto attributed to the oxyfluoride. Also, a new line appeared, at d = 2. 22 which had been previously hidden by a line due to scandium oxide. These lines were all weak so another sample was heated three hours but no increase in the number or intensity of the extraneous lines appeared. So it is assumed that the hydrolysis was caused by the small amount of moisture which entered, when the sample was inserted into the furnace. All seven new diffraction lines are catalogued below in Table 11.
Table 11
New Diffraction Lines appearing in the Pattern of Hydrolysed, Scandium Fluoride d (A)
.10 .44 .65, .15 .59 22
NP’Nwqu-Ui .12
The conversion of scandium fluoride to scandium oxide with little oxyfluoride intermediate can be explained as follows. Wells (78) points out that scandium and rare earth oxides crystallize with a structure very similar to that of calcium fluoride. Since the oxyfluorides of the rare earths have the fluorite structure, too, an oxyfluoride intermediate would be crystallographically similar to the oxide. Hence, the oxyfluoride any form.
Scandium oxyfluoride is unable to form the fluorite structure, and steric considerations indicate it would have the rutile structure. It is because -34-
the rutile -type scandium oxyfluoride is crystallographically dissimilar
from the fluorite-type scandium oxide that the conversion to oxyfluoride
is hampered. I
The cubic lattice constants for scandium oxide and scandium
fluoride were calculated from powder diagrams and are as follows:
SCZO3, a.= 9. 88; ScF3, a.= 4. 022. These values are to be considered
approximate, as film shrinkage was not take into account. However,
comparison of these two cell distances with literature values is good.
Nowacki (49) gives 4. 022 A for the unit cell distance of ScF3 and
Goldschmidt (23) gives 9. 79 A for that of Sc O 23' In the last stages of this study another attempt was made to prepare
' mixed fluorides of scandium and univalent ions. Scandium and sodium
fluorides were heated in 1:1 mole proportions at 960° in a tube furnace under
dry oxygen for five minutes. The powder diffraction pattern showed
scandium fluoride lines due to unreacted fluoride and also many new lines
possibly due to a double fluoride. The exact composition of this new phase
was not determined, but, analogous to ammonium scandium double fluorides,
it should be NaScF Na ScF5 or Na3ScF6. 4’ 2 An attempt was made to prepare the complex ammonium scandium
fluorides by the "wet" method. That is, by dissolving freshly precipitated
scandium fluoride in ammonium fluoride in 1:3 mole proportion.
Recrystallization of this solution gave a mixture as determined under
the microscope. Needles, showing extinction parallel and perpendicular
to their long direction, rectangular rods of slight birefringence and some
nonbirefringent octahedra were found. The rods and octahedra appear to
have a refractive index of about 1.4. The powder diagram showed a great -35- profusion of lines. The identical procedure when used with sodium and potassium fluorides simply precipitated out most of the scandium immediately after addition of metal fluorides.
C. Trichloroacetates, Acetates and Carbonates of Scandium.
Cleve (6) mentioned the acetate of scandium without giving its composition. Crookes (8) described it with the analyses for the per cent of scandium oxide. Later he mentioned the mono- and dihydrate of scandium monobasic acetate. Sterba-B'dhm and Melichar(67) gave three methods of preparing the anhydrous acetate, and described monobasic scandium acetate. The acetates of the rare earths and of aluminum are well known; however, the latter cannot be produced in the presence of water because of extensive hydrolysis.
A trichloroacetate of scandium has not been previously found, but
Crookes (8), in an extensive study of the products formed between organic anions and scandium, found that all but the oxalate, lactate, and _r_n_-nitroben- zoate of scandium exist as basic salts. Of special interest is the fact that Crookes found Sc(OH)(CH2ClCOZ)2' ZHZO (‘10) and (CHZBrCHBrCOZI3Sc -
(CHzBrCHBrCOZ)ZScOH.
Both normal and basic, metallic salts of trichloroacetic acid are known and have been extensively reviewed (58). The normal trichloro- acetates are ionic salts which are highly soluble in water, alcohol and acetone, but practically insoluble in carbon tetrachloride and benzene.
The basic trichloroacetates are insoluble in water, but soluble in alcohol and are considered covalent. Solutions of metal trichloroacetates are known to give carbonates when they are heated. The kinetics of this type of reaction have been studied (75, 76). This method has been used to obtain
-36- pure, crystalline rare earth carbonates (58), because it has been found that precipitation of rare earth carbonates with ammonium or alkali metal carbonates gives poorly crystalline materials. In addition, these carbonates always contain some of the precipitating metal ion by occlusion or double salt formation. No trichloroacetate of aluminum has been prepared.
The first indication relative to the preparation of scandium carbonate was by Cleve (5). The author only stated that the voluminous precipitate gotten with a scandium ion solution and sodium carbonate redissolves on addition of excess of the latter. Crookes (9) reported scandium carbonate dodecahydrate by precipitation with sodium carbonate, but he could not get concordant results when analyses were made. The Sterba-B'dhm's
(68) could not repeat Crooke's work, but reported Sc(OH)CO3'H 0 instead. 2 However, they say that the precipitate is so soiled with sodium ion that the formula is not to be considered definite. The latter workers, having rejected direct precipitation methods of producing the carbonate, tried to add carbon dioxide tofreshly precipitated SCZO3-HZO and the dried oxide under pressure. Analyses showed no scandium carbonate was produced.
Pokras and Kilpatrick (53) attempted to produce scandium carbonate by dissolution of the oxide in excess trichloroacetic acid. However they found that several weeks of heating on the steam bath gave no precipitate, and concluded scandium carbonate could not be produced in this way.
Hydrolysis studies on the rare earth trichloroacetates, which involved metals with ionic radii as small as yttrium , have produced normal carbonates (27) .
\/ -37-
Preparation of scandium trichloroacetate.
The refractory nature of scandium oxide makes the use of boiling concentrated mineral acids necessary. With trichloroacetic acid,* a very strong organic acid, the same conditions are necessary except that longer dissolution periods are required. For instance, very little of the oxide dissolves into 25% TCA at 60°, while a large excess of 50% acid at 90° to 100° dissolves it almost completely.
Unfortunately, to prepare the metal trichloroacetate, this large excess of acid must be separated. Ifdone in the usual way by heating the solution, some of the primary hydrolysis product, chloroform, is hydrolysed and chloride ion is formed. To avoid using a large excess of the acid, the oxide used was formed by ignition at 520° instead of 900°.
An excess of TCA was still necessary. This solution was evaporated to dryness in a vacuum desiccator. The solid produced was analysed for scandium. (Found: 13.28%; Theor. 8.45%). Subsequent extractions and recrystallizations with ethanol, ethyl acetate or acetone did not sufficiently purify the product to give good scandium trichloroacetate.
Accordingly, this method of preparation was dropped. A better method was to precipitate hydrated scandium oxide from nitrate solutions with ammonia, centrifuge and heat to rid of the last bit of ammonium nitrate.
This gave a scandium material which was almost 100% soluble in the stoichiometric amount of 50% TCA.
A typical preparation of scandium trichloroacetate was as follows:
SczO3(0. 5338 g.) was dissolved in concentrated HNO3 and evaporated to dryness twice. . The residue was taken up in a little water and washed into a centrifuge bottle. Ammonia was added in slight excess and the mixture
*Hereafter referred to as “TCA.
\Ii/ «- | a x'} -38-
centrifuged. After decanting, fresh water was added and the precipitate was slurried. It was again centrifuged and decanted. The bottle was then placed in an oven at 230° for ten hours. Ammonium nitrate
decomposes at an appreciable rate at this temperature (57). After heating, the hydrated oxide was allowed to stand overnight in contact with the stoichiometric amount of TCA, and all but a slight amount of dark
substance dissolved. The solution was then poured into an evaporating
dish and allowed to stand in a vacuum desiccator until the water was
removed. The trichloroacetate crystallized in long transparent rods.
The material was then scraped off the evaporating dish, crushed into a fine, white powder and dried two days over calcium chloride. The yield was 98% based on the amount of scandium oxide used. Scandium analyses
were made by the 8-quinolinol method, and chlorine analyses were made by refluxing the trichloroacetate with sodium and isopropyl alcohol and then titrating the chloride ion with potassium thiocyanate using ferric ion
as an indicator. The analyses correspond to the formula Sc(CCl3COZ)3.
Found: Sc 8. 59%; Cl 59.18%. Theory: Sc 8.45%; Cl 59.97%.
It is interesting that in one of the trichloroacetate preparations the hydrated oxide was inadvertently allowed to heat at 270°. This material would not fully dissolve in TCA, even when an excess was added, and after
six or seven days' standing.
The density of Sc(CCl3~COZ)3 was measured by the pycnometric method using xylene. As with the oxalate density measurements, the xylene was
stored over sodium ribbon and distilled just before using. Only the middle fraction was used. The density of the trichloroacetate was found to be 1. 80 g. /ml. The solubility of scandium trichloroacetate in the twenty-five
\/ -39-
milliliters of xylene used in the measurement was 8 mg. or 0. 2% of the total weight of solid used. This contributes negligible error.
The trichloroacetate did not melt or char up to 300°, but did seem to be decomposing at the higher temperatures.
Qualitative solubility measurements were made by adding approximately
4 mg. of the trichloroacetate at a time to 5 ml. of the particular solvent and agitating on a mechanical shaker. The results of these tests are given in Table 12.
Table 12
Solubility of Sc(CC13COZ)3 in Various Solvents
Solvent Solubility
Benzene (0.004 g.I5 ml. Xylene (0. 004 g. [5 m1. Carbon tetrachloride (0. 004 g‘. [5 m1. 1, 4-Dioxane immediate turbidity Ether M0.01g./5 m1. Acetone soluble Ethanol quite soluble
Moderately concentrated (M0. 1 M) solutions of scandium trichloro- acetate do not hydrolyse in the cold on standing. Hydrolysis of scandium trichloroacetate.
The behavior of scandium trichloroacetate solutions upon heating show that they hydrolyse to form hydrated scandium oxide with no evidence of a carbonate. Several trials which involved dissolving scandium oxide into excess TCA gave the same white gelatinous precipitate as heating a solution made by dissolving the trichloroacetate in water. In order to prove that the hydrolysate was really hydrated scandium oxide, powder
-4() ..
diffraction identification was used. The hydrolysate had the same powder pattern as the precipitate formed by ammonium hydroxide and scandium nitrate. Weiser and Milligan (77) have shown that ScZO3'HZO precipitates under the latter conditions. Using X-ray techniques, they gave the diffraction pattern of this substance, but did not elucidate the structure.
The lines are recorded in Table 13. The small discrepancies in the two patterns are due to the diffuse powder patterns which are given by these partially amorphous pr ecipitate s .
Table 13
Powder Diffraction Lines for Sc(CCl3COZ)3 Hydrolysate
SCZO3'HZO - Sc(CC13COZ)3 hydrolysate
S (cm.) S (cm.)
1. 36 l. 35 2.60 2. 59 3. 44 3.43 3. 60 3.60 4.48 4.48 4.72 ' 4.68 4.99 5. 00 5.64 5. 64 5.78 5. 76 6.33 6.31 6. 50 6.49 7.47 7. 47
To get some idea as to how the hydrolysis proceeded, l g. of scandium trichloroacetate was dissolved in 30 to 35 ml. of water and heated on a hot plate at 78° to 82°. These date are recorded in Table 14.
-41_
Table 14
Observations on the Hydrolysis of Sc(CCl3COZ)3
Hydrolysis time Observations
l. 25 hrs. COz evolution;CHCl3 odor; no AgNO3 test 2. 25 hrs. CO evolution; pH 2. 94 3. 25 hrs. No 0 ; faint AgNO test 5.25 hrs. No CH 1 odor; frot ' g 7. 75 hrs. Frothing; little pptn.; pH 4. 37 10. 75 hrs. Definite precipitation of SCZO3'HZO
The silver nitrate test in Table 14 shows that some Cl ion is formed
after 3. 25 hours' hydrolysis just as with rare earth trichloroacetates.
The long (11 hour) period before SCZO3-HZO began to precipitate indicates
that scandium trichloroacetate is harder to hydrolyse than the analogous
rare earth compound.
Solubility of scandium acetate.
Anhydrous scandium acetate was prepared by the method of Sterba- Bb'hm. It hydrolyses so readily that when a 0.1 lvi_ solution was allowed
to stand overnight, considerable basic acetate or hydrated oxide
precipitated. A qualitative estimation of the temperature coefficient of
solubility could not be made because of this strong hydrolysis.
D. Oxides, Hydroxides and Sulfates of Scandium.
Considerable study has been directed towards determining just what
specie precipitates, when a hydroxide is added to scandium ion solutions .
It haslbeen reported as scandium hydroxide (7), 2 Sc(OH)3-HzD (67) and
scandium oxide monohydrate (77). The latter reference showed
conclusively that hydrous scandium‘ oxide monohydrate was the specie precipitated between 25° and 100° and that it closely resembles I—AIZO3°HZO.
-42-
Also shown was that the oxide is not dimorphic below 900° as previously reported.
The structure of scandium oxide has been investigated (23, 52), and. was shown to have an almost cubic unit cell with 3 equal to 9.79A, and to have 16 ScZO3 molecules in the repeating unit. Aluminum oxide is dimorphic forming the a and )3 modifications. Both structures are formed by a network of closely packed oxygen ions with interstitial aluminum ions (78). The rare earth oxides are also dimorphic, one form being a high temperature modification (78). A
Regardless of what compound is precipitated by adding bases to scandium solutions, the precipitate comes down on the acid side of pH equallto seven. Sterba-oB'dhm and Melichar (67) have reported that the pH of precipitation from 0. 01 M scandium acetate solution was 6. 10 after
2. 25 equivalents of 0. 089 1:1 potassium hydroxide were added. Ivanov-Emin and Ostroumov (26) reported that with 5x10'3 M scandium chloride, precipitation begins at a pH equal to 4. 90 after one equivalent of 0. 1 1:1 sodium hydroxide, and is complete at a pH of 5.45 with 1. 8 equivalents of base. With 2. 5x10"3 _M_ scandium sulfate, precipitation begins at a pH of 4. 90 with 0. 75 equivalents of sodium hydroxide and is complete at a pH equal to 5. 10 with l. 25 equivalents of base. In another study, titration of 0.0075 _1\_/f scandium sulfate with 0. 2 N sodium hydroxide gave a precipitate at a pH of “5. 3 after two equivalents of base had been added (13).
Recently, excellent studies have been published on the nature of the scandium specie existent in perchlorate solutions, when the hydroxyl ion was increased (28). The results are amenable with the two step reaction \f’
-43-
+++ ++ + Sc(HzO)6 + HZO—-) Sc(HZO)5(OH) + H30 2 56(H20) 5(OH)++—) [Sc(HZOI 50H]?
In the same study, Pokras and Kilpatrick reported precipitation occurred at a pH of 4. 54 with 2.0x10'z _M_ scandium perchlorate after
1. 24 equivalents of sodium hydroxide had been added.
The hydrated sulfates, basic sulfates and double sulfates of scandium
are completely reviewed elsewhere (53). Anhydrous scandium sulfate may be produced by evaporating a dilute sulfuric acid solution of scandium to about 70% acid at 100°, and washing with alcohol and ether to remove any adsorbed sulfuric acid (71).
Of interest in this study are solutions of the normal sulfate, because the complex trisulfatoscandiate ion is reported to exist in them. Meyer
(41) reports ion migration experiments in sulfate solutions in which more than half of the scandium migration is anodic. All the rare earth sulfates tried behaved in this manner but to a lesser extent. Other authors have observed that cerium forms a complex sulfato -ion, Ce(SO4)f, in solution, but it has been reported only for very concentrated sulfuric acid media
(50, 61, 63, 32, 12). Early conductance studies of scandium sulfate
solutions showed a normal rise of conductance values with dilution which lead the authors to state, that the solutions were little influenced by hydrolysis (41). However, the pH value of 3.67 for a 0.00753 13L scandium
sulfate solution indicates considerable hydrolysis (1.3). Also, later Brauner and Svagr (2, 3) showed the conductance of scandium sulfate to be much the
same as that of the rare earth sulfates except at very high dilutions where it almost doubles the conductances of some rare earth sulfates. Conductance
\/ -44- readings taken immediately after dilution, differed from’those taken after a waiting period of several hours, giving further evidence for a hydrolysis reactio n which is not rapid.
Meyer (41) reports that oxalic acid precipitates scandium oxalate. from sulfate solution slowly and incompletely, while under the same conditions, a rapid and almost complete separation from nitrate and chloride solutions results. '
The only other reference to the scandium specie existent in solution is that mentioned previously by Pokras (28) for perchlorate solutions.
Kraus and Smith (31) reported that filter paper electromigrations appear to be a rapid means of determining the charge of complex ions and for obtaining estimates of the stability constants of complexes. Acid
solutions as strong as l N HCl with potentials as high as 30 volts have been used (37), but McDonald (39) reports, when the ionic strength of the carrier exceeds 0. 1, it becomes difficult to achievelinear movement of the migrant with time. Strain 313.1. (60) have reported on filter paper migration of scandium and rare earths in lactic acid. They have shown that scandium migrates to the cathode, as do the rare earths, in solutions of lactic acid up to 1. 5 M. Scandium leads the rare earths in 0.1 M lactic acid but gradually falls behind as the lactic acid concentration increases up to l. 0 M, where it trails.
The complex nature of scandium sulfate.
It is difficult to study the species existent in scandium sulfate using ordinary methods, because it does not absorb in the region 2200 to 10,000 A
(105), and no known compound, extractable in organic medium, appears to show reproducible absorption in the organic phase (53).
-45-
Anhydrous scandium sulfate was prepared by dissolving scandium oxide in sulfuric‘acid and concentrating the acid solution to about 70% sulfuric acid. Under these conditions, scandium sulfate is almost completely insoluble. The sulfate was then washed three times with alcohol and dried with ether to get rid of excess acid. This was repeated four times. The sulfate was heated at 400° for three—quarters of an hour, and showed only 0. 25% weight loss. Analysis of this sulfate by the 8-quinolinol method gave the following values for per cent scandium:
Found 24.04%, 24.10%; theory 23. 80%.
Filter paper ion migration studies of scandium sulfate solutions appeared to be a simple yet conclusive way of studying the charge on the scandium ion, if only a good color test for scandium could be devised.
Much time was spent in looking for a suitable reagent, but no 100% efficient test was found. Aurinetricarboxylic acid, conchineal, and morine were all eliminated, and sodium alizarine sulfonate was finally used. The spot test, using the latter dye, went as follows: After the electrodes were disconnected, the filter paper was dipped into a 1% dye solution, exposed to ammonia vapors to precipitate the scandium, and then exposed to hydrogen chloride gas. Short treatment with hydrogen chloride changes all the basic form of the indicator (violet) to the acidic form (yellow) except that which is adsorbed by the scandium. Unfortunately, this test does not give good results if the supporting electrolyte is very concentrated. The test breaks down if a fluoride is used as the supporting electrolyte .
-46 -
When migration runs were actually made, another limitation on the usefulness of this technique was discovered. To get as rapid a movement as possible, as large a voltage as possible was used. This meant that, at higher sulfuric acid concentrations 90. 51:1), the power dissipated along the (paper strip was enough to evaporate the water solvent, concentrate the acid, and char the paper. Various schemes were used such as making the run under a bell jar in an atmosphere saturated with water vapor, but all were more or less unsuccessful. The results, despite all the limi- tations mentioned, showed that the entire amount of scandium was transported to the cathode in a sulfuric acid supporting electrolyte up to 0. 25 M. in sulfate ion. However, in the lower sulfuric acid concen- trations especially, two color bands appeared indicating a part of the scandium moves with a different velocity.
Applied potentials of 300 v. (0.011! H2504) down to 30 v. (0. SN H2504) were generated with a regulated D.C. supply of the type which is ordinarily used for transference studies. The filter paper, 10w 1. 2 cm. , was supported laterally by three glass rods attached to a metal frame, and, at either end, 1"ears” hung down into beakers of the carrying electrolyte.
Copper clamps from the power supply were connected to the paper through platinum strips,
In a further effort to find whether appreciable scandium is tied up as
Sc(SO4):, scandium sulfate solutions, with and without added sulfate ion
(added as ammonium sulfate), were titrated using sodium hydroxide.
Figure 4 shows the result of the titration with added ammonium sulfate.
The second plateau is due to the titration of ammonia by the base.
FIGURE 4 TITRATION OF 0.0I_M_ 3960,)3 WITH I.OO8_M_ NaOH ONE MOLE (NH4IZSO4ADDED PER MOLE sczcsq),
I l0-
PPTN. INCIDENCE CI.|9 EQUIVS.)
I l J I LOO 2.00 3.00 4.00 - EQU IVS. Na 0H -48-
The experimental conditions were to titrate 0. 01 M scandium sulfate
with 1. 008 1:1 sodium hydroxide using a microburette readable to 0. 005 m1.
Careful titration showed that the addition of 0. 066 g. of ammonium sulfate
(one mole of ammonium sulfate for each mole of scandium sulfate) did not affect the pH at which precipitation incidence was noted. That is,
in both titrations, precipitation began at a mole ratio OH-/SC+H=1.19
and at a pH of 5. 27. The end point of the titration in which no added
sulfate ion was used came at OH-/SC+H=2.70 at a pH of 8.42. If scandium
ion were able to be complexed through an equilibrium reaction, appli-
cation of the mass action principle shows that added sulfate ion would
reduce the concentration of free scandium ion, and that more hydroxide
ion would be needed to exceed the solubility product of the precipitated
specie. Thus, one is fairly safe in stating that little, if any, complex
formation occurs between sulfate and scandium ions.
E. Polarographic Studies of Solutions Containing Scandium Salts.
The first application of polarography to solutions of scandium salts
was made by Noddack and Brukl (47). In the course of a study on the
reduction of rare earth sulfates using no supporting electrolyte or maximum
suppressor, scandium sulfate was electrolysed and a double wave found.
The first wave at -l. 63 v. and the second wave at -1.79 v. (tangent
potentials) were ascribed to the following processes.
SC+H+ {in—)SCH SoH + rot-2.56o
Leach and Terry's (35) studies of scandium chloride solutions also
stowed a double wave. However, the first or prewave was nicely defined \t -49- by increasing the hydrogen ion concentration, and so they ascribed the first wave to the reduction of hydrogen. The second wave is then scandium ion being reduced to the metal with E% equal to -l. 80 v. versus the standard calomel electrode. The same authors found a linear relation between concentration and diffusion current in solutions as dilute as 10-4_I\_/i_ in scandium ion, if sufficient hydrochloric acid was added to define the prewave. The study was run in several supporting electrolytes, lithium chloride being the best. However, low galvanometer sensitivities (l to Spa/mm.) were used. At these sensitivities, the current passing through the solution is high enough to cause serious complications in the electrode reactions.
West, Dean, and Breda (79) found the scandium wave to be completely obliterated in a l millimolar scandium chloride solution to which 0. 5 moles per liter NaF, as supporting electrolyte, have been added.
The early polarographic work on rare earth solutions was carried out by Noddack whose experiments were performed without supporting electrolytes or maximum suppressor, so the discussion will be confined to the excellent and recent studies of Glockler et a1. (15, 56, 69, 70) and
Laitinenitjl; (33, 34).
Of all rare earths studied polarographically, europium, ytterbium, and samarium show stepwise reduction, the others, single step reduction.
The E5“ values for the reduction of all rare earths to the metal are not separated to any great extent with the exception of europium and ytterbium.
Apparently, in all the studies, a small amount of acid must be added to develop the hydrogen prewave. When this is done, the polarographic wave of the metal ion is well separated and well defined. Without this \J -50- small amount of added acid, the diffusion current depends on both the scandium and hydrogen ion concentration. As one increases the acid concentration, a definite shift of E; to more negative potentials is noticed. 2 The following data are taken from a paper by Swenson and Glockler (69).
Table 15
Variation of E; with Acid Concentration 2 Pr2(SO4)3(6. 75 millimolar) in 0. 1 N LiCl and 0. 01% gelatin with H2504 added.
pH :ELVB.S.C.E. 2 5.88 -1.85 v. 3.08 -1.86 v. 2.40 -l.93 v.
Also, as one increases the concentration of rare earth ion, El shifts 2 to more negative potentials. The following data are from (56).
Table 16
Variation of E; with Gd Ion Concentration 2 Gd2(SO4)3 in 0.1 E1 KCl and 0.01% gelatin.
Millimoles Gd2(SO4)3/ liter E_1_ versus S.C.E. 2 0. 80 -1.74 v. 1.60 -1. 74 v. 4.00 -1.77 v.
Glockler states that in all cases the number of electrons involved in the reduction is three, but in (56) he states that when log(i/id - i ) is plotted versus Ed e , irreversibility is indicated by the fact that the slope is not as steep as one would expect from a three electron reduction.
-51-
The polarographic behavior of aluminum is well known (36) . One
finds well developed waves (Eé~ is equal to -1. 75 vs. S.C.E.) using
lithium chloride as a supporting electrolyte, if one does not use a
hydrogen ion concentration which exceeds the aluminum ion
concentration. Large quantities of acid completely mask the aluminum
wave.
Experimental.
The polarograms were automatically recorded on a Sargent model XXI
polarograph. The type of H cell used has been described by others (70).
A 0.4% agar 2 molar potassium chloride plug connected the saturated
calomel cell and the H cell electrically. The resistance of the H cell
and standard cell was assumed to be about 200 ohms, and a correction for the IR drop was applied when it was significant. Previous to each
run, water pumped nitrogen was bubbled directly from a cylinder for ten minutes through each solution. Since tests showed no interference
due to the oxygen in commercial nitrogen, its removal was not deemed necessary. No attempt was made to control the temperature with a bath.
E% and id values were evaluated from the polarograms by the method described in Williard, Merritt and Dean (82).
Anhydrous scandium sulfate was prepared as described elsewhere
in this thesis. A 0.1_M stock solution was prepared by dissolving the
stoichiometric amount of sulfate in distilled water. Scandium perchlorate
was prepared by dissolution of the oxide into the stoichiometric amount of
concentrated perchloric acid. The negligible amount of oxide which did
not dissolve was filtered off, the solution evaporated almost to dryness
on a hot plate and made up to volume. The solutions for polarography
-52- were prepared by dilution of this 0.l_l\_d_ stock solution. A 0.1 M stock
solution of scandium chloride was prepared by dissolution of the oxide into excess concentrated hydrochloric acid and evaporation to crystals several times to remove excess acid.
Polarography in scandium chloride solutions.
To check the work of Leach. and Terry (35), 5 millimolar solutions of scandium chloride were prepared, and varying amounts of hydrochloric acid added to them. Polarography revealed E_1_ increases in the negative 2 direction with increased acid concentration. However, E; values found . 2 in this study are higher than previously reported.
Table 17
E; of ScCl3 Solutions of Varying Acidity 2 ScCl3 in 0.01% gelatin and 0.1 M- LiCl
ScCl3 concn. E; vs. S.C. E. HCl Concn. 2 5 millimolar ’ l. 68 (single wave) 1 millimolar 5 " 1. 83 6 ' " 5 " 1. 88 11 ”
Table 17 shows the acid concentration needs to be greater than 1 millimolar to define the hydrogen wave under the conditions of this experiment. Six millimolar acid concentrations define the hydrogen wave, so two waves are found. However, both waves end in inflection points and not in plateaus, so the waves are ill-defined and Eé‘ difficult to obtain exactly.
In separate experiments the effect of a small amount of chloride ion
-53..
upon E; was studied. These data are shown in Table 18. 2
Table 18
E; of ScCl3 Solutions of Varying Acidity 2 ScCl3 in 0.01% gelatin and 0.1 _M_ LiClO4
ScCl3 concn. E; vs. S. C. E. HClO4 concn. 2 5 millimolar 1.69 (single wave) 1 millimolar 5 " 1. 80 7 " 5 H 1. 83 9 H 5 " 1. 88 13 "
The results in Table 18 are quite similar to those in Table 17. Increase
in acid concentration again leads to a negative increase in E%' Further
study of the polarograms showed that the waves ended in inflection points
so were ill-defined.
Plots of logi/id -i versus potential for the polarograms in Tables
17 and 18 have. slopes corresponding to less than one electron drops. For
the reaction SCH to scandium metal the electron change is three. This
shows the electrode process is quite irreversible in these solutions.
Figure 5 gives a typical set of data.
Polarography in sulfate solutions.
Polarographic studies were originally undertaken to determine the
extent of complex formation in scandium sulfate solutions. However,
since the results in scandium chloride solutions showed the electrode
process was definitely irreversible, a quantitative study of complex
. formation was impossible.
Cursory polarographic study of sulfate solutions indicated the system
was too complex to reveal significant results. Consequently, this study
was dropped.
FIGURE 5 DEPENDANCE OF LOGO/gs) ON EMF ' SXIO‘BM ScCI,,IO"I_/_I to, iO“°/. GELATIN, LIX IO‘ZM HCL
0.4 ~-—-
I
.0 o
(i/id-i)
LOG I I m ,o I I I 180 |85 '90 E MF (VOLTS) _55-
Polarography in scandium perchlorate solutions.
Polarographic studies were next made on scandium perchlorate in
wholly perchlorate media. The perchlorate ion is known to be a poor
complexing agent, so the reducible specie should be Sc(HZO)ZH or some
form of [SC(HZO)6(—9H) x]( 3 -x)+ .
Table 19
E; of Sc(ClO4)3 Solutions of Varying Acidity 2 Sc(ClO4)3 in 0.01% gelatin and o.1_1\_,t_ Lic104 Sc(ClO 4 ) 3 concn. E; 2 vs. S.C.E. HClO 4 concn. 2 millimolar -l. 76 v. None 2 " -1.79 v. 2. 5 millimolar 2 " -1. 83 v. 5.0 " *5 ” —1.88 v. 5.0 " *5 ,H -1.90 v. 7.5 " * 5 " —1. 94 V. 10 "
Table 19 shows E% in creases negatively with increasing acidity as
was the case in chloride media. In the above polarograms neither the
hydrogen wave nor the scandium wave ended in a plateau, but simply
terminated in inflection points. When larger acid concentrations and
lower sensitivities were used, the hydrogen wave terminated in a short
plateau, but E’é‘ of the scandium ion was displaced beyond -l.90 v. versus
5. C. E. , where the discharge of hydrogen became visible. These facts,
together with the lowered sensitivity useable, dictated against recording
any data at higher acid concentrations.
Studies using tetramethylammonium brorride as a supporting electrolyte
are shown in Table 20. The statements in the above paragraph on the
*NaClO4 was used as a supporting electrolyte. \g’ -56-
nature of the polarographic waves hold true here.
Table 20
E; of Sc(ClO4)3 Solutions of Varying Acidity 2 3 Millimolar Sc(C104)3 in 0. 01% gelatin and 0. l _M_ (CH3)4 Br.
E; vs. S.C.E. HClO4 Concn. 2 -l. 73 None -1. 77 2. 5 millimolar -1. 83 5.0 "
Plots of log i/id-i versus potential for the data show the slope is considerably less than the theoretical. The fact that these slopes correspond to less than a one electron drop, shows the electrode reaction is quite irreversible. These results are identical to those found in chloride media.
In order to avoid using acidic scandium solution, solutions of ammonium scandifluoride were employed. Scandium fluoride was completely dissolved in ammonium fluoride (1 moles ScF3/3 moles of NH4F), and the resulting solutions diluted to contain five millimoles of scandium fluoride/liter. This was made 0.1 _M_ in sodium perchlorate, and 0.01% gelatin was added. A study of the polarographic behavior of this solution showed the scandium wave was indeed obliterated. This indicates that the concentration of the scandium specie, which is ordinarily reduced, is lessened to a polarographically insignificant amount. This, of course, is further evidence for the existence of quite strong fluoride complexes of scandium.
It is also of interest to find whether the scandium ion concentration
-57- is proportional to the diffusion current in solutions which are not acid enough to require very low galvanometer sensitivities or dischrage hydrogen simultaneously with the scandium wave. The data are shown in Table 21.
Table 21
Diffusion Current Versus Concentration
Sc(ClO4)3 in 0.01% gelatin, 0.1 _M_ LiClO4 and 2. 5 millimolar HClO4 Galvanometer . . . Sc(C10 4 ) 3 concn. E; a vs. S.C.E. i d pH Sen81t1v1ty
0.2 )1 a/mm 0. 5 millimolar -1.76 v. 6.0 p a 2.69 0.15pa/mm 1.0 " -l.,78 v. 13.2 pa 2.64 0.3 pa/mm 2.0 "' -l.82 v. 24.9)13. 2.57 0.8 pa/mm 6.0 " -1.93 v. 81.6ya 2.39
A plot of id versus concentration is given in Figure 6, and one sees that the dependence is linear except at the highest scandium perchlorate concentration. Table 21 shows the acid concentration corresponding to this is considerably different than for the other points. This explains why the id is larger than a linear relation predicts. FIGURE 6
CONCENTRATION DEPENDANCE OF DIFFUSION CURRENT 90— 80— O 70 60 50'
§ 40
\ \l VVI ‘30 I I I 5X/0'4/0‘3 5X/0'3 CONCENTRATION CM/L) -59-
IV. Summary.
(1) Freshly precipitated scandium oxalate dried over calcium
chloride for two days yields a material corresponding
approximately to a hexahydrate. On extended drying, a
material forms which is more nearly a dihydrate in composition.
Dehydration of the hexahydrate over magnesium perchlorate
gives the dihydrate as does‘heating at temperatures of 80°
to 200°.
(Z) The densities of the hexahydrated and dihydrated scandium oxalates are 2.15 and 2. 26, respectively.
(3) Scandium oxalate hexahydrate crystallizes in flat, diamond shaped, monoclinic crystals with refractive indices as
follows: 0.1. 495; [31. 602;?1. 63.
(4) The reaction of hydrated scandium oxalate with moist ammonia
gas is shown to be simply the reaction of the salt of a weak
base with hydroxyl ions.
(5) Evidence for the formation of scandium oxyfluoride as an
intermediate in the hydrolysis of scandium fluoride is given.
The rate of this hydrolysis at several temperatures is
tabulated.
(6) The lattice constants for scandium fluoride and scandium
oxide we chermined and found to be in good agreement with
literature values. They are: SCF3, a = 4.022; SCZO 3’ a = 9. 88.
_60-
(7) The new compound, scandium trichloroacetate Sc(CCl3COZ)3,
was prepared. Solutions of this salt hydrolyse to hydrated
scandium oxide and not to scandium carbonate as is the case
with the rare earths. Some physical properties of Scandium
trichloroacetate are. given.
(8) Evidence is given which indicates scandium sulfate solutions
contain no anionic scandium species as previously reported.
(9) The polarographic E% for scandium perchlorate and scandium chloride solutions is shown to vary with acidity and scandium
ion concentration as has been reported for the rare earths.
Graphs of log i/id-i versus potential for this media show
considerably less than the theoretical slope which shows the
electrode process is quite irreversible. The plot of scandium
concentration versus diffusion current is shown to be roughly
linear, at a hydrogen ion concentration of 3 millimolar, for
the perchlorate media. \I.’
-61 -
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CEHRQSTRYIJBRARY
Riley, Reed Farrar The preparation and properties of some scandium compounds.
T645.9 CHEMISTRY LIBRARY R573 Riley, Reed Fbrrcr ‘7) )_.1 I. .13 The preparation and properties of some scandium compounds.