PHYSICOCHEMICAL
STUDIES ON GALLIUM(III)
SALT SOLUTIONS
999
SOME PHYSICOCHEMICAL STUDIES O S GALLIUM (111) SALT SOLUTIOXS' THERALD MOELLER A X D GLENDALL L. KING*
Noyes Chm&xd Laboratory, University of Illinois, Urbana, Illinois Received October 80, 1949 IKTRODUCTION
Among the characteristics of the gallium-indium-thallium family essential to its systematic study are trends and variations in basicity. The combined effects of small size and comparatively high charge among the tripositive ions of these elements should render all of them quite highly acidic, but significant differences should be discernible as these factors vary. Variations in basicity hare been studied as part of a general investigation of comparative properties in this family. I t is the purpose of this communication to summarize results obtained with gallium(II1) compounds by the techniques of electrometric titration with alkali and measurement of pH as a function of salt concentration and to compare them with existent data for indium(II1) compounds. Previously reported measurements of these types have been limited in number and lacking in agreement. Fricke and Meyring (3) obtained data, based upon a hydrogen electrode study of titration of gallium(II1) chloride solutions with sodium and potassium hydroxides, which indicate precipitation of the hydrous oxide (9) in the vicinity of pH 2.8-3.1. Von Bergkampf (15) reported precipitation from hydrochloric acid solution at pH 3.4 and from sodium hydroxide solution at pH 9.7 as a result of single titrations of 0.1 N gallium(II1) chloride solutions, a quinhydrone electrode being used. More recently Ivanov-Emin and Rabovik (5) found from a hydrogen electrode study that precipitation from 0.01 M gallium(II1) chloride solutions began at pH 3.48-3.50 with 2.6 equivalents of sodium hydroxide added and was complete at pH 6.7-6.8 with 2.8-2.9 equivalents. With 0.005 M gallium(II1) sulfate solutions, however, precipitation began at pH 3.4C-3.45 with 0.2 equivalent and was complete at pH 4.74.8with 2.5-2.6 equivalents. Dissolution of the hydrous precipitate at pH 9.6-9.7 after addition of 4.0-4.2 equivalents of alkali was considered evidence for the formation of the species [Ga(OH)J- in solution. Lacroix (8) obtained precipitation from 0.007 M gallium(II1) chloride solution at ca. pH 3.2 and dissolution of the hydrous oxide over a broad range above pH 6. Hydrolysis studies on gallium(II1) salt solutions appear limited to those of Fricke and his coworkers (2, 3). A preliminary measurement (2) of hydrogenion concentration showing 0.1875 N gallium(II1) chloride solution to be 32.9 per cent hydrolyzed was followed (3) by similar measurements as functions of Presented in part at the 114th Kational Meeting of the American Chemical Society, which was held in Portland, Oregon, September, 1948. Reproduction in whole or in part permitted for any purpose of the United States Government. 2 Present address: Los Alamos Scientific Laboratory, LOBAlamos, Kew Mexico.
1000
THERALD MOELLER AXD GLEKDALL L. KIXG
gallium concentration. Values of pH from 3.362 for 0.00314 M to 1.624 for 0.0951 M solutions were in agreement with a first hydrolysis constant, corresponding to the formation of Ga(OH)++ions, of 1.4 X lo+. Electrometric titration data for a second hydrolysis constant corresponding (3) led to a value of 3.5 X to the formation of Ga(OH)z+ ions. Correspondingly, the first and second stages in the hydrolysis of sodium gallate solutions were characterized, respectively, by the constants 3 X lo-* and 1.2 X Lack of agreement and comprehensiveness in these data warrants more nearly complete studies. The data reported here were obtained from electrometric titration studies on gallium(II1) sulfate, bromide, chloride, and nitrate solutions and from hydrolysis studies on bromide, chloride, and nitrate solutions. EXPERIMENTAL
A . Materials The gallium compounds used in this investigation were prepared from a sample of the metala containing only spectroscopictraces of other elements. Pure gallium(111) sulfate was obtained by dissolving a sample of the metal in a mixture of concentrated nitric and sulfuric acids and evaporating off the excess acids to leave a crystalline product. Anhydrous gallium(II1) bromide was prepared by passing nitrogen laden with bromine over gallium metal at 300°C. (7) and repeatedly subliming the product in dry nitrogen. Anhydrous gallium(II1) chloride was prepared by passing dry hydrogen chloride over gallium metal at 200°C. (6) and repeatedly subliming the product in dry nitrogen. Both bromide and chloride samples were sealed in nitrogen in glass ampoules which could be broken just prior to use. Gallium(II1) nitrate was obtained by dissolving the metal in dilute nitric acid solution, concentrating the resulting solution on the steam bath, cooling, and drying the crystallized product for several days in a desiccator over calcium chloride. Solutions obtained by dissolving these gallium compounds in carbon dioxide-free distilled water were standardized by gravimetric determination of gallium as either the oxide or the %hydroxyquinoline derivative (13). Stock solutions were diluted to desired concentrations with carbon dioxide-free water. Tenth-normal solutions of sodium hydroxide were prepared by dilution of a concentrated, carbonate-free solution with carbon dioxide-free water and were Standardized against the primary standard potassium acid phthalate. These solutions were protected from atmospheric carbon dioxide with soda lime.
B. Procedure For all measurements, solutions were maintained within Zt0.2"C. of the stated temperature by thermostating. Electrometric titration data were obtained by the standard procedure ( l l ) , using a Beckman Model H pH meter the glass electrode of which was checked frequently against a 0.05 M potassium acid phthalate a
Obtained from the Eagle-Picher Company, Joplin, Missouri.
\
PHYSICOCHEMICAL STUDIES ON GALLIUM(II1) SALT SOLUTIONS
1001
buffer. Hydrolysis measurements were made by determining pH values as functions of concentration (IO, 12). RESULTS AND DISCUSSION
A. Electromelric titration studies Data were obtained for electrometric titrations with sodium hydroxide of gallium(II1) sulfate solutions varying from 0.099 to 0.229 M in gallium ion at temperatures of lo", 25", and 40°C. Corresponding data were obtained at 25'C. for 0.0204.180 M gallium(II1) bromide solutions, 0.0104.150 M gallium(II1) chloride solutions, and 0.010-0.123 M gallium(II1) nitrate solutions. Some backtitration studies with standard acids were also carried out.
I I2
-1 0
t
I
I
1
I
I
3.0
4.0
5.0
'I
PPTN.
1.0
2.0
MOLE
6.0
R A T I O OH;/GA*++
FIG.1. Titration of gallium (111) sulfate solutions with sodium hydroxide
Representative data for 0.063 M gallium(II1) sulfate solutions (0.125 M in gallium ion) at the three temperatures are given in figure 1, plotted as pH us. the ratio of moles of hydroxyl ion added to moles of gallium ion initially present. Points at which precipitation began and dissolution of the precipitate was complete are indicated by small arrows. Data a t other concentrations gave exactly similar plots, the displacements of the curves with increasing temperature as shown in figure 1 being characteristic of all plots for solutions of a given concentration. Significant data characterizing these titrations are summarized in table 1. I t is apparent that, regardless of temperature, roughly 1 mole of hydroxyl ion to each mole of gallium ion was necessary for precipitation incidence. The average
1002
THERALD MOELLER AKD GLESD.ILL L. K I S G
values for these ratios are 1.23 at 10°C., 1.11 a t 25"C., and 0.93 at 40°C. Comparison of these ratios with the corresponding ones of 0.94, 0.84, and 0.73 for indium(II1) sulfate solutions (11) suggests a greater tendency for aggregation in the gallium solutions on addition of hydroxyl ion. I t was impossible to obtain precipitation a t lower ratios as reported by Ivanov-Emin and Rabovik ( 5 ) . Precipitation pH values are essentially independent of concentration a t a given concentration but decrease significantly with temperature increase. AverTABLE 1 Effect of gallzum-ion concentration on titrations o j gallzum(lll) sulfate solutions DISSOLUTION OF PRECIPITATE
PRECIPITATION INCIDENCE IhTRAL Ga++* CONCENTPATION
10°C. molcr/lr1rr
0 0 0 0
125 150 195 229
Average. . . . . . . . . . . . , . . . . . . . . . . .
3 19 121
.I
3.03
1
1.23
1 ~
11 11 12 11
90 85 08 85
11.92
1
1
,
5 5 5 5
i
5.41
1
40 42 40 41
25°C. 2.85 2.78
0.125 0.150 0.195 0.229
__
'
1.14 1.05 1.11
____
Average. . . . , . , . . . , , . , . , . . . . . . . . , . .
2.80
11.40 11.48 11.85 11.76
5.41 5.17 6.31 6.07
11.62
5.74
10 57 10 65 10 85
4 92 4 90 5 00
40°C. 0,125 0.150 0.195 Average
1
2 69 2 65 2 65
0 95
I 2 6 8
1 0 9 3
~
0 93 0 92
~
1
1069
~
1 4 9 4
age values of 3.03 a t 10°C., 2.80 at 25OC., and 2.68 a t 40°C. are again to be compared with 3.56, 3.43, and 3.15, respectively, for indium(II1) sulfate solutions a t the same temperature (10). A somewhat greater acidity for the gallium(111) ion is thus indicated. Data characterizing complete dissolution of the precipitates are subject to somewhat greater variations. In general, between 5 and 0 moles of hydroxyl ion were required per mole of gallium ion to cause dissolution. These mole ratios are probably without real significance, for dissolution, like precipitation, was undoubtedly spread over a considerable interval. The lack of significant breaks
PHYSICOCHEMICAL STUDIES ON GALLIUM(III)
1003
SALT SOLUTIONS
in the titration curves in this region (see figure 1) may be cited as evidence in support of this conclusion. I t seems reasonable to expect considerable hydrolysis in gallate solutions, and the total quantity of hydroxyl ion required to effect complete dissolution is doubtless the sum of the quantity necessary for gallate formation and the quantity necessary to prevent precipitation by hydrolysis. The slight breaks in the titration curves (figure 1) in the vicinity of mole ratio 4 may indicate gallate formation involving four hydroxyls to one gallium. Ivanov-Emin and Rabovik (5) suggest this to be true, and the data of Fricke and Meyring (3) are not inconsistent with this conclusion. TABLE 2 Solubility product and solubility data, calculated f r o m titrattons of gallium(l1l) sulfate solutions TEMPEEATUEE
'C. 10
1 I I ~
I ~ ~ T I AGaL
CONCLXTEATION
SOLUBILITY PPODUCT
X 10:'
1
Average
SOLUBILITY
(molcr/lilcr)
0.17 0.08 0.10
0.125 0.150 0.195 0.229
~___________
I
1
molrr/lilsr
0.04
I
0.125 0.150 0.195 0.229
0 10
x
IOU
4.9 4.1 4.3 3.4
I
4 2 7.0 7.0 6.8 7.4
0.68 0.70 0.58 0.85 -I-
Average
7.1
0.70
0.125 0.150 0.195 Average,, , . , , , , , , . . , , . , . , , , , , , , , .
I I
8.4 7.3 12.5
I
I
9.4
I
13.1 12.6 14.4
I
13.4
~
The curves in figure 1 indicate completion of precipitation when the mole ratio of OH- to Gareaches 3, corresponding to the formation of the hydrous hydroxide or oxide. Formulation of the precipitation product as the latter is more consistent with the data of Milligan and Weiser (9), although for convenience in evaluation of the solubility product constant and water solubility, consideration of the material as hydroxide is useful. Application of the conventional treatment (e.g., 11, 14) without use of activities at various OH- to Ga+++ ratios in the precipitation regions of the titrations yielded the data summarized in table 2, the ion product of water being taken as 0.295 X lo-", 1 X lo-", and 3.02 X lo-" at lo", 25", and 40°C., respectively. Calculated solubility product constants increase markedly in magnitude with temperature but are of the By a similar approach using a chloride solugeneral order of magnitude of
1004
THERALD MOELLER AND GLENDALL L. KING
tion, Lacroix (8) obtained a value of 10-36,5.Corresponding water solubilities of the order of 10-lo gram-mole per liter agree quite closely with an average of 6.3 X l0-lo reported by Oka (14).Corresponding calculations (11) for indium hydroxide gave for the solubility product constant and gram-mole per liter for the water solubility, variation with temperature being a little less striking.
I
I
I
I
%:I
I
I
12-
OISSOLN.
.-
P PTN.
4
0
1.0
2.0
3.0
MOLE R A T I O
4.0
5.0
6-0
OH~GA+++
FIG 2. Titration of gallium(II1) bromide solutions with sodium hydroxide
Plotted in figures 2, 3, and 4 are representative titration data for gallium(II1) bromide, chloride, and nitrate solutions, respectively, at 25°C. Xumerical values significant to these titrations are summarized in table 3. While the curves obtained for the bromide, chloride, and nitrate are similar to those obtained with sulfate solutions, precipitation incidence is markedly different. With bromide, chloride, and nitrate, nearly 3.0 moles of hydroxyl ion were required per mole of gallium ion before visible precipitation occurred, yet the plateaus in the titration curves indicate steady consumption of hydroxyl ion up to these points. At the points of precipitation incidence shown in figures 2 , 3 , and 4 and recorded in table 3, complete flocculation occurred, but it was never possible to obtain indications of cloudiness in solutions containing less hydroxyl ion. It appears that either excessive ion aggregation occurs under such conditions or that peptization by excess gallium(II1) ion takes place as hydroxyl ion is added, the bromide, chloride, and nitrate ions having negligible flocculating effects on the sols produced. It is significant that Ivanov-Emin and Rabovik (5) reported precipitation from gallium(II1) chloride solutions only after addi-
PHYSICOCHEMICAL STUDIES ON GALLIUM(III)
SALT SOLUTIONS
I2
10
8
PH 6
4
2
0
LO
3.0
2.0
MOLE
RATIO
40
5.0
6.0
OH;/GA'*+
FIG.3 . Titration of gollium(II1) chloride solutions with sodium hydroxide
0
I 1.0
2.0
3.0
MOLE R A T I O
4.0
5.0
8.0
OH/GA*++
FIG.4. Titration of gallium(II1) nitrate solutions with sodium hydroxide
1005
1006
THERALD MOELLER AND GLEKDALL L. KING
TABLE 3 Effects of gallium-ion concentration on titrations of bromide, chloride, and nitrate solutions at 56°C. PRECIPITATION INCWESCE
IhlTIAL GaCONCENTPATION
pH
I
OH-/Ga+'+
1
1
DISSOLUTION OF PPECIPITATE
pH
I
OH-/Ga+++
Gallium(II1) bromide moleslliler
0.020 0.030 0.060 0.100 0.125 0.150 0.180 Average. .
5.60 5.30 5.05 5.09 5.02 4.99 5.00
.,
.,,I
5.15
3.63 3.63 3.50 3.14 2.98 2.90 2.81
i
3.23
i
11.43 11.28 11.42 11.57 11.50 11.65 11.63
6.05 5.62 6.05 4.92 4.96 5.53 4.81
11.50
5.42
Gallium(II1) chloride 0.010 0.015 0.020 0.030 0.040 0.060 0.100 0.125 0.150
5.50 5.25 5.15 5.22 5.00 4.91 4.80 4.71 4.51
3.21 3.19 3.14 3.13 3.12 3.07 2.98 2.94 2.97
10.90 11,05 11.11 11.05 11.10 11.25 11.50 11.62 11.68
Average., . . . . . . .
5.01
3.08
11.24
Gallium(II1) nitrate 0.010 0.020 0.030 0.040
0.050 0.060 0,070
0.082 0.103 0.123 Average
__
5 25 5 25 5 35 5 38 5 25 5 15 5 08 4 78 4 80 484 5 11
~
__
3.02 2.95 2.92 2.92 2.78 2.93 2.91 2.86 2.85 2.86
10.80 11.20 11.22 10.90 11.12 11.05 10.98 11.18 11.05 11.10
2.90
11.06
1 1 ~
1
'
4.80 4.94 5.06 4.72 5.06 4.89 5.26 5.19 5.13 5.01
1
449 452 4 48 4 48 452 4 31 4 34 4 45 4 32
1
4.44
~
tion of 2.6 equivalents of sodium hydroxide, although Fricke and Meyring (3) obtained precipitates shortly after the addition of single equivalents. An investigation of these phenomena will be reported in a subsequent communication. Because of the uncertainties characterizing precipitation from bromide, chlo-
1007
PHYSICOCHEMICAL STUDIES O N GALLIUM(II1) SALT SOLUTIOXS
ride, and nitrate solutions, evaluations of solubility product constants were not attempted. The major inflections in the curves in figures 2 , 3, and 4 again indicate completion of precipitation at an OH- to Ga+++ ratio of 3.0 and suggest that the hydrous oxide (or hydroxide) is the ultimate precipitate. Data on dissolution of the precipitates in excess hydroxyl ion are in excellent agreement with those given for the sulfate solutions. Alterations in initial gallium-ion concentration are without major effect at any place in the titrations. A few back-titrations with standard acid were carried out to cast light on the nature of the gallate solutions. Plotted in figure 5 are data for titration of 0.02
,
MOLE R A T I O O ;.
i
"/OH0.rO
Oi40
0.160
"t 210
1
I I
I
10 .
2.0
MOLE
I
I
3.0
4.0
RATIO
OHyGA++*
I
I
5.0
6.0
1
FIG.5. Titration and back-titration of a gallium (111) nitrate solution
M gallium(II1) nitrate solution with 0.1 N sodium hydroxide followed by backtitration of the strongly alkaline solution with 0.1 N nitric acid. As acid is added, hydrogen ion is consumed regularly by excess hydroxyl ion. However, the curve breaks significantly just before the hydrous precipitate reappears, and the position of this break suggests rather conclusively that four hydroxyls are associated with a single gallium in the gallate. The remainder of the curve is essentially a reproduction, although at slightly lower pH values, of the original titration curve. Studies at other concentrations gave the same results.
B. Hydrolysis studies Hydrolysis studies were limited to gallium(II1) bromide, chloride, and nitrate solutions at 25'C. Bromide, chloride, and nitrate were selected as being suffi-
1008
THERALD MOELLER AKD GLEND.4LL L. KING
ciently xeakly basic in character to contribute comparatively nothing to any observed alterations in the hydrogen-ion concentrations in the solutions. Under these conditions then, any such alterations could be ascribed to the acidic characteristics of the gallium(II1) ion. It may be assumed, as a first approximation, that the governing hydrolytic process in dilute solutions may be formulated as either Ga+++
+ H20
+ H+
$ GaOH-
(1)
or Gaff+
+ HzO
GaO+
+ 2H+
(2)
ionic hydration and complex-ion formation being neglected in all cases. The first of these reactions is characteristic of a number of other tripositive ions, among them indium (4,10, 12), and has been suggested for gallium (3, 8). The second is probably never a primary process, but a product such as GaO+ (or its hydrate, Ga(0H)z+) may result from the secondary hydrolysis of GaOH++. Applications of conventional treatments (1) to these equations yields the expression
for equation 1 and the expression
for equation 2, where Kh and K: are the respective hydrolysis constants, c the molar concentration of gallium ion, and z the degree of hydrolysis. The degree of hydrolysis, in turn, may be evaluated for reactions 1 and 2 as
x
=
c,+/c
and
x
=
cp/2c
respectively. Both equations 3 and 4 are of such form that if the hydrolytic process be represented by either equation 1 or equation 2, linear relationships should exist between pH and log c. Data showing the existence of such relations a t least up to c = 0.1 are plotted for bromide, chloride, and nitrate solutions in figure 6. -4greement among bromide, chloride, and nitrate data is excellent and suggests the three anions to be of equal influence or lack of influence in governing hydrolysis. Deviations a t higher concentrations are not unexpected since activities have been neglected, while those a t very low concentrations doubtless result from lack of absolute accuracy in pH measurement. Decision as to the mechanism governing the hydrolytic process is best based upon the constancy of the hydrolysis constant as calculated from experimental data. In all cases a greater degree of constancy is obtained by application of equation 3, and it may be assumed, therefore, that reaction 1 is governing. Data tabulated in tables 4, 5 , and G summarize experimental observations on bromide, chloride, and nitrate solutions and calculated values for the degree of
PHYSICOCHEMICAL STUDIES ON GALLIUM(III)
CONCENTRATION
GA***+MOLES
1009
SALT SOLUTIOKS
PER LITER
FIG.6. Hydrolysis of gallium(II1) bromide, chloride, and nitrate solutions
TABLE 4 Hudrolzisis o ialliuni(II1 womide solutions at W C . " " IXlTIAL
Cia+-+ CONCEXTPATION
1
''
DEGREE OF EMROLYSIS
pH
EYDROLYSIS CONSTAPI
x
(XI
IO'
-~ ---I
molcr/lilrr
0.9000 0.8000 0.7000 0.6000 0.5000 ,
1.18 1.28 1.39 1.51 1.66
;$:
f :i
0 2000 0 1500 0 1250
~
2 10 2 19 2 24
0,103 0.102 0.0741 0,0657 0.0581 0,0515 0.0438
~
1
0 0 0 0
0394 0382 0398 0429 0 0459
1
144. 123. 52.6 36.8 25.1 16.7 10.0 6.66 5.61 4 80 3 80 3 26 2 87 2 i6
0.1000 0.0800 0.0600 0.0400 0.0300 0.0200 0.0150 0.0100 0,0080 0.0060 0.0040 0.0020 0.0010
2.29 2.33 2.38 2.45 2.50 2.58 2.62 2.67 2.70 2.75 2.81 2.93 3.08
0.0512 0.0583 0.0695 0.0888 0.106 0.131 0.160 0.214 0.249 0.297 0.388 0.585 0.830
2.86 2.90 3.10 3.48 3.70 4.03 4.43 5.72 6.60 7.92 12.0 13.7 34.4
hydrolysis and hydrolysis constant based on application of the principles outlined above t o equation 1. Agreement among average values over the concentration range 0.0044.28 M of 4.G X 4.1 X and 3.7 X for the hydrolysis constants for bromide, chloride, and nitrate solutions, respectively, is excellent. Indium(II1) bromide and chloride solutions, by comparison, are uni-
1010
THERALD MOELLER Ah-D GLENDALL L. KING
formly less hydrolyzed at 25°C. and are characterized by mean hydrolysis conand 1.2 X respectively (10, 12), again illustrating the stants of 1.4 X greater acidity of the gallium(II1) ion. The data of Fricke and his coworkers (2, 3) indicate considerably greater hydrolysis for gallium(II1) chloride solutions than do the data presented here, but those values were obtained with solutions TABLE 5 Hydrolvsis of q a l l i u m ( I I I ) chloride solutions at W C . INITIAL
Gah+
DEGREE OF HYDROLYSIS (t)
PH
CONCEZTRATIOZ
1
~
1.81 1.86 1.94 2.03 2.13 2.20 2.28 2.34 2.38 2.42 2 46
0.0465 0,0470 0.0460 0,0467 0,0494 0.0505 0.0525 0.0572 0.0596 0.0633 0.0694
'
I
'
3.35 2.91 2.77 2.65 2.57 2.59
0.0080
0.0060 0.0040 0.0020
0.0010
2 49 253 2 59 264 2 71 2.74 2.79 1 2.84 2.97 3.11 1
INITIAL
Ga+*+
PH
DECREE OF EYDROLYSIS
1
pH
(x)
TR.4TION
-
x
lo'
0.0983
0.1285 0.153 0.193 0.227 0.270 0.363 0.535 0.776
1
DEGREE OF EYDRPI;YSIS
1
EYDPOLYSIS COKSTAYT
x
10'
molcs/lilcr
tnalcs/liler
0.2620 0.2450 0.2050 0.1640 0,1230 0,1030 0,0820 0,0700 0,0600 0,0500
EYDPOLYSIS CONSTANT
2.86 3.22 3.78 4.12 4.73 5.36 6.00 8.25 12.3 18.6
0.0810
-Hydrolysis of galZium(lI1) nitrate solutions at M°C. CONCENTRATION
,
l-
moles/lilcr
0.3334 0.3000 0.2500 0.2000 0.1500 0.1250 0.1000 0 0800 0 0700 0.0600 0.0500
1
2.09 2.10 2.15 2.21 2.27 2.31 2.34 2.37 2.39 2.42
-
0.0310 0,0324 0.0345 0.0376 0.0437 0.0476 0.0557 0.0610 0.0679 0.0760
2.61 2.65 2.53 2.41 2.44 2.45 2.70 2.78 2.96 3.13
0.0400 0.0300 0.0200 0.0100 0.0080 0.0060 0.0040 0.0020 0.0010
2.46 2.50 2.55 2.67 2.72 2.80 2.88 2.99 3.10
0,0868 0.105 0.141 0.214 0.238 0.265 0.330 0,510 0.794
3.30 3.72 4.62 5.83 6.00 5.74 6.61 10.6 30.6
containing free hydrochloric acid, which may not have been completely corrected for, and were not obtained at any stated temperature. SUMMARY
1. The acidic characteristics of gallium(II1) ion in aqueous solution have been studied by electrometric titrations with sodium hydroxide and by hydrogen-ion determinations as functions of concentration.
PHYSICOCHEMICAL STCDIES
ox
QALLIUM(III) SALT SOLUTIONS
1011
2 . In the presence of sulfate ion, electrometric titration has shown the hydrous oxide to begin precipitating when the mole ratio of added hydroxyl ion t o gallium ion approximates 1 and to be complete when this ratio reaches 3. At 25°C. the average pH for precipitation is 2.80. 3. In the presence of bromide, chloride, and nitrate ion, apparent precipitation is delayed until the OH- to Gaff+ ratio reaches 3, where complete flocculation occurs. The pH values characterizing these changes are not too consistent. 4. In all cases, dissolution of the precipitates in excess hydroxyl ion occurs only after addition of 5-6 moles of hydroxyl ion, although the gallate ion has the composition [Ga(OH)J-. 5 . Solubility product constants for the hydrous hydroxide of the order of and corresponding water solubilities of the order of M have been evaluated from data on sulfate solutions. 6. Hydrolysis in dilute gallium(II1) bromide, chloride, and nitrate solutions has been shown to be governed by formation of GaOH++ ions and to correspond respectively, to hydrolysis constants of 4.6 X low4,4.1 X lo-*, and 3.7 X a t 25°C. 7. Gallium(II1) ion has been shown by these data to be somewhat more acidic than indium(II1) ion. The authors wish to express appreciation to the Office of Kava1 Research for support received during this investigation. REFERENCES (1) BRITTOK:Hydrogen Ions, 3rd edition, Vol. I, pp. 208-11. Chapman and Hall, Ltd., London (1942). (2) FRICKEA N D BLEEKE:Z . anorg. allgem. Chem. 145, 183 (1925). (3) FRICKE A N D MEYRING: Z.anorg. allgem. Chem. 176, 325 (1928). (4) HATTOX . ~ N D DE VRIES:J. Am. Chem. soc. 68, 2126 (1936). (5) IVASOV-EMIN AND RABOYIK: J. Gen. Chem. (C.S.S.R.) 14, 781 (1944). (6) JOHSSOK ASD HASKEW: Inorganic Syntheses, 1'01. I , pp. 26-7. McGraw-Hill Book Company, Inc., S e w York (1939). (7) KLEMM AND TILK:Z . anorg. allgem. Chem. 207, 161 (1932). (8) LACROIX: Dissertation, University of Paris, Paris, France (1948). AND WEISER:J. Am. Chem. Soc. 69, 1670 (1937). (9) MILLIGAK (10) MOELLER:J. Am. Chem. SOC.63, 1206 (1941). (11) MOELLER:J. Am. Chem. Soc. 88,2625 (1941). (12) MOELLER:J. Am. Chem. Soc. 64, 953 (1942). (13) MOELLER AND COHEN:Anal. Chem. 22, 686 (1950). (14) OKA:J. Chem. SOC.Japan 69,971 (1938). (15) YON BERQKAMPF: Z. anal. Chem. BO, 333 (1932).
