Feb., 1945
THESYSTEM ALUMINUM SULFATE-ETHANOL-WATER
nate bonds. Charcoal seems to be the most suitable catalyst for practical purposes. In the presence of charcoal, aqueous ammonia displaces both chlorides from cis- [Co enrC12]+, giving preponderantly the trans product, [Co ena(NHa)l]+++. In the absence of the catalyst, only one ammonia molecule enters the coiirdination sphere, and the product [Co enz(NHa)C1]++ has the cis configuration, irrespective of whether the dichloro salt was cis or trans. Triethylenediamine chromic chloride is readily prepared by the reaction of hydrated chromic chloride and hydrated ethylenediamine in the presence of charcoal. In the presence of charcoal, sodium nitrite
179
reacts with hexammine cobaltic ion and ammonium hydroxide reacts with cobaltinitrite to give trinitrotriammine cobalt, which thus seems to be the most stable member of the cobalt nitro ammine series. Organic amines react with the dichlorodiethylenediamine cobaltic ion to give ions of the type [Co enz(RNH2)C1]++if no catalyst is present, but in the presence of charcoal, complete rearrangement to the extremely stable [Co en,] +++ ion takes place. The presence of. charcoal during reactions of optically active cobalt ammines does not alter the sign of rotation of the products. URBANA, ILLINOIS
RECEIVED MAY12,1944
[CONTRIBUTION FROM THE BUREAU OF MINES,EASTERN EXPERIMENT STATION]
The System Aluminum Sulfate-Ethanol-Water at 30' and 80' BY EDWINA. GEE* investigation. Aluminum sulfate solubility was determined by direct evaporation and ignition of the liquid sample or by precipitation with ammonia. Water was the 30 and 80' isotherms of the system aluminum calculated by difference. The density of the equilibrium sulfat-thanol-water and to determine the vari- solutions was measured in 50-milliliter pycnometers. Points near tLe plait point were determined by titrating ous hydrates present a t equiIibrium. two-phase liquid systems with water t o homogeneous soluA method for the preparation of iron-free alumi- tions. ntim sulfate using ethyl alcohol as a solvent was Identification of Solid Phases.-Analysis of the solid brought to the attention of the Bureau of Mines by phase was made by two methods. I n the first, the crystals were separated from the mother liquor by filtration after Dr. P. S. Roller of the Bureau staff. In the pilot- which they were successively washed with absolute alcohol plant development that followed, data on the and anhydrous ether. The water of hydration was then system aluminum sulfate-ethanol-water were determined by ignition. This method had a tendency to necessary to interpret certain phenomena that give low results in determination of the higher hydrates. was probably due t o the dehydration eKect of the had been observed. Systems of the general type This anhydrous alcohol and ether which were necessary to ob(inorganic salt-ethanol-water) have been previ- tain a rapidly drying product. The second method is esously investigated. sentially an application of the work of Foote and Sholes in which the hydrates of inorganic salts in equilibrium Experimental Methods with various alcohol mixtures are determined.' If a Solvent.-Ethanol-water mixtures were prepared from slightly soluble salt of known hydrate is placed in an alcodistilled water and purified ethyl alcohol. Commercial hol-water mixture the specific gravity of the alcoholic 96% alcohol was refluxed for several houis over lime and solution will increase or decrease according to the release or distdled; the distillate was redistilled to ensure complete absorption of water by the salt. If no change is noted the separation from the lime. Density determinations were hydrate is in equilibrium with that alcoholic solution. made to establish solution composition, pycnometers of Thus, by observing the change in specific gravity of the 50-ml. capacity being used. solvent the solid phase present at equilibrium with various Alumhum Sulfate.-Baker analyzed aluminum sulfate alcoholic mixtures can be identified. This method neceswas employed throughout. sitated the preparation of known hydrates which were preSolubility Determinations.-Saturation was effected in pared by precipitation of aluminum sulfate from alcohol250-milliliter bottles which were rotated four to six days water mixtures. The various hydrates were dried over in two water-baths operated at 30 0.01' and 80 * 0.2", correspondingly lower hydrates. The final hydrate comrespectively. When equjlibrium was approached from position was determined by ignition. undersaturation aluminum sulfate was added directly to The technique required slight modification for applicathe solvent. Approach from supersaturation was simu- tion to this investigation. As several hydrates are present lated by adding concentrated aqueous aluminum sulfate a t the same temperature, there must be a transition point solution t o the alcohol. After equilibrium was reached the at which two hydrates will be in equilibrium with the same mixtures were allowed to settle and were filtered through alcoholic solution. After the approximate location of this sintered glass of .fine size. transition point had been ascertained, its position could Alcohol content was determined by distillation to dry-. be established by adding a mixture of the two hydrates n&, the composition of the distillate being determined by involved and observing the change in composition. density. This method yielded alcohol percentages accurate within 1to 2% and this limited the over-all accuracy of the Results and Discussion
Introduction
This investigation was undertaken to establish
f
(1) Published by permhion of the Director, Bureau of Minu. U. 5. Department of the Interior. Not copyrighted. (2) h o d a t e physical chemist, Bureau of Mines, Eastern Experiment Station, College Park, Md.
The results of the investigation a t 30 and 80' are recorded in Table I and Fig. 1and in Table I1 (3) H. W. Foote and S.R. Sholea, Tars JOURNAL, 88,1809 (1911).
Vol. 67
EDWINA. GkB
180
TABLE I Total composition, % A. S EtOH
2.8 3.3 3.1 1.0 1.8 2.7 1.0 3.5 2.4 1.6 5.2 2.0 1.7 3.2 4.1 4.8 3.3 2.9 5.2 6.0 6.1
97.1 (G) 93.0 91.8 (R) 91.2
EQmLIBarvM.DATA FOR 30' ISOTHERM, A o m p o s i t i o n of solution% A, S. % EtOH Density
0.1 .1
.1
.o .0 .o .o
88.6 80 5
76.0 66.8 06.6 66.0 60.8 (Q) 62.0 60.4 56.5 54.5 52.3 50.8 49.1 45.3 41.8 38.2
.o
$0
.o '
0
.o .1 .2 $2
.3 .4 .i
1.0 1.4 2.8
99.8 96.0 95.1 92.6 90.8 83.5 77.5 70.1 69.1 67.8 66.4 64.0 62.4 59.4 58.8 56.8 53.5 51.8 49.9 46.0 40.9
0.781 0.790 .791 .800 ,804
.824 .a1 .861 .863 .867 .870 ,878
.881 .882 .892
.898 .908 .913 919 .924 .935
24.7 (EandF) 27.7
Solid phase
Aln(SO4)s Al:(SOdi Al:(SO4), and Als(SO4):.lOH;O Alr(SO4)r.lOHtO Aln(SO4)r*1OH%O Alz(SO4):.10H10 Alz(m4)r.lO&O A~:(SO~)~.~OHSO ~,(S04)r'lOH*O Ah(SOda*lOH& Alr (S04):.10HI0 and A12 (S04):. 16HzO A1&3Or)a~16HrO Alr(S04):.16H10 A1&304):*16HzO A b (S04):. 16H& Al~(S04):*16HsO Al3(SO4):*16H10 Als(SO4)a*16HtO A l r (SO4h. 16HzO Alr(SOl)t*l6HzO A1z(SO4),*16&0
Double layer'
c r
13.2 10.2
30'
% A . S.
% EtOH
3.8 6.9
38,l 33.0
Density '
0.947 (F)
0.957
% A.S.
11.4 15.3
Bottom %EtOH
26.1 20.0
Duuity
1.038 (E; 0.890
Solid phase
19.5 15.5 1.162 24.7 24.2 11.4 20.3 13.0 1.186 22.5 8.1 1.223 24.2 7.3 26.6 2.0 1.293 27.2 2.0 27.7 0.0 1.302 !a.7 0.0 (D) Aluminum sulfate hexadecahydrate is the solid phase throughout. 13.2
&(SOdr*16HIO At(S04h.16HIO &(mdr*1fWO AL(S03a.16HIO Ali(SOc)a*l6HiO
and Fig. 2, respectively. The table headings In Fig. 1 the solubility of AIp(S04)s.16Ht0 in Total Composition, and Composition of Solution, alcohol-water solvents is represented bp the curves refer to the composition of the initial over- or DE aqd FQ, which are separated by the 2 liquid undersaturated mixture and to the final solution in equilibrium with solid phase, respectively.
curve EPF, E and F being theisothermally invariant conjugate solutions in equilibrium with
Fig. 1.-Phase
Fig. 2.-Phase
diagram of the system A4/SO&-EtOHHtO a t 30'.
diagram of the system AL(S0')rEtOHHz0 at 80".
Tmz SYSTEM ALUMPJU SULFATE-ETFUNOL-WATER ~~
Feb., 1945
181
TABLEI1 Totalcomposition 7 A. S. &O&
EQurLmWrras DATA FOR 80' ISOTHERM, 80" --Composition of mlution%A. S. % EtOH Density
96.8 (G)
3.0 2.0 2.4 2.6
0.9 .7
90.9 84.3 81.0
3.1
17.6 (9)
.o
2.1
73.3
.1
0.746 .759 ,776 .783 .790 .812
98.9 92.0 88.3 83.0 80.0 74.9
.&
.1
Solid phase
Alz(S03: Alr(SO4)a
AI:(SOd: M:(SOda Aln(SO4)a and Al:(SOJa.lOHiO
Double layera
30.1
22.8 16.3 27.7 21.1 14.0 25.7
19.5 13.0 18.2
TOP
Bottom
%AS.
% EtOH
Density
% A.S.
% EtOH
2.0 1.9 2.0 2.0 3.0 3.4 4.6
63.5 62.1 60.8 80.5
0.840 (F)
38.0 36.4 3288 32.9 32.2 29.4 28.5 27.2 25.0 23.2
6.9 7.2 9.1 9.0 9.3 11.5 11.5 12.8 14.0 15.5
19.1 (EandF) 29.2 38.9 17.7 26.9 36.4 16.4 26.0 33.7 23.2
6.6 6.6 8.4
.848
.864
.w .860 .871 .903 ,912 .918 .964
53.5 62.4
50.0 47.3
44.0 38.8
.
Density
1.346 (E) 1.320 1.286 1.280 1.230 1.210 1.200 1.152 1.I43 1.132
Solid pbax
24.0 16.3 1.131 17.6 17.6 22.0 1.042 12.3 12.3 31.0 0.990 40.6 39.4 3.6 43.6 0.0 (D) 42.2 0.0 Aluminum sulfate decahydrate is the solid phase throughout.
24.0
16.3 22.0 31.0 3.6
... ...
the solid. The curve QR represents solutions in equilibrium withethe deca-hydrate and RG solutions saturated with the anhydrous salt, Q and R being the isothermally invariant solutions saturated in each case with two solids, hexadeca- and deca-hydrate, and deca-hydrate and anhydrous salt, respectively. I n Fig. 2 the solubility of AMSOhlOHsO in alcohol water solvents is represented by the curves DE and FR, which are separated by the 2 liquid curve EPF corresponding to Fig. 1. The curve RG represents solutions in equilibrium with the anhydrous salt, R being the isothermally invariant solution saturated with deca-hydrate and anhydrous salt. The data on the identification of the solid phases are given in Table 111. The results are an average of a t least five determinations. TABLE111 IDENTIFICATION OF SOLID PEASE,Moms HnO, 30" Direct ignition
15.2 * 0.3 9.7 0.2 0.2 f O . 1 f
9.0 0.1
* 0.6 * 0.1
Indirect density
Average
15.80 16.4 * 0 . 3 10.30 10.9 * 0.4 0.10 0.0 a o . 1 Moles &O, 80" 10.1 * 0.4 9.65 0.2 * 0.1 0.15
Probable formula
16.0 10.0 0.0 10.0 0.0
Dirsculty was experienced in identification of the solid phases. As shdwn in'Table 111, con-
Alr (Sod)vlOH10 Ah (SO4)I' lOH2O Aln(SO,)a.lOH*O Als(SOJ,.lOH~O Al:(SO~)~.lOH~O
siderable deviation occurred between the two methods. The direct separation of solid hydrates from equilibrium solutions was complicated by the presence of aqueous alcohol. The use of anhydrous alcohol and ether as a wash obviated this difficulty to a large degree as it permitted rapid drying of the solid. However, this procedure gave low results with the higher hydrates for the aforementioned reasons. The application of the method of Foote and Sholes was fairly simple a t a room temperature but density determinations a t the higher temperatures were difficult. The method of wet residues was not applicable to this study because of insuflicient precision of the alcohol determination a t low concentrations. The literature gives numerous references for the hydrate in equilibrium with solution a t room temperature.' The majority of investigators give the octadeca-hydrate as the stable form. These data have been recently reviewed by N. 0. Smith and a series of careful measurements made which establish the stable form as the heptadecahydrate at 2 5 O . 4
Inasmuch a s t h e latter's work was a t a lower temperature, the hexadFcahydrate which coincides most closely with the experimental results was chosen as the most probable solid phase a t 30'. As the investigation was of a preliminary nature, the possibility that other hydrates may exist within narrow fixed limits of alcohol strength must be admitted. (4)
N. 4.Smith, Tzus JOURNAL, 64, 4 1 (1942).
JAMES J.
182
LINGANE
Vol. 67
The data indicate two hydrates, the hexadeca tion. The rdative precision of the 80' isotherm and deca as well as anhydrous salt exist at 30'. is thus limited to *3%. One hydrate, the deca, and anhydrous salt were S = w found to exist a t 80'. The 30 and 80' isotherms for the system The precision of measurement was limited by the alcohol determination. As no special pre- aluminum sulfate-ethanol-water have been detercautions were taken in these determinations the mined. Anhydrous salt, hexadecahydrate and relative precision of the 30' isotherm is *2oJ,. decahydrate were found to be present a t 30' and At 80" difficulty was experienced in the sampling anhydrous salt and decahydrate a t 80'. of solutions of high aluminum sulfate concentra- COLLEGE PARK,MD. F~ECEIVSD AUGUST26, 1944
[CONTRIBUTION FROM THE
MALLINC~~ODT CHBXICAL LABOWTORY OF HARVARD UNIVERSITY ]
Polarographic Characteristics of Vanadium in its Various Oxidation States BY JAMES J. LINGANE The lirst polarographic study of vanadium compounds was carried out by Zeltzer,' who investigated the reduction of vanadium trichloride from acid solution, and the reduction of vanadate (+5 vanadium) from acid, strongly alkaline, and ammoniacal media. Unfortunately the vanadium trichloride solutions used by Zeltzer were more or less oxidized and not standardized, so that the polarograms that he obtained cannot be interpreted. Neither is it possible to deduce half-wave potential data for the reduction of vanadate, or to definitdy assign oxidation states to the various waves of vanadate, from the incomplete polarograms given by Zeltzer. The reduction of vanadate at the dropping electrode from acidic, strongly alkaline, and ammoniacal media has since been studied by Stackelberg, et aLJ2 Thanheiser and Willems,' Vori~ k o v aand , ~ Page and RobinsonJKbut, since these authors were completely preoccupied with the devefopment of empirical methods for determining vanadium in various materials, very little fundamental infomation can be gleaned from their papers. Indeed, their conclusions in several respects are bewildering. For example, Stackelberg, et uZ.,~ concluded that vanadate ion is reduced to the +3 state from ammoniacal medium, whereas Zeltzer' had concluded that reduction proceeds to 'he +2 state. In the text of his paper Voriskova4 assumes reduction t o the +2 state, but in an accompanying table he presents an equation showing reduction to the +3 state. Furthermore, although the polarograms of vanadic acid in dilute hydrochloric acid presented by Zeltzer' clearly show a double wave, those published by Stackdberg, et al., display only a single wave, and Thanhdser and Willems' state that they observed only a single wave fiom slightly (1) S. zCl&er, CoU. C&%boslw. C h . Conrmur.. 4, 818 (1032). (2) M. P. StpfLelberg, P. glinger, W. Koch and E. h t h , Forrchurtsbsrkhle Tack. Miff. Krrrpp. Earen, %, 59 (1989). (3) G. Thanheism and J. Willems, Arch. Eismhilitcnm., la, 19 (1939). (4) M. Vorislrwa, Coll. Cseckorh. Chem. Commun., 11,688 (1939). (5) J. E. Page and F. A. Robinson,A r o l y r f . 68,260 (1943).
acid solutions (PH2 to 6) of +5 vanadium a t about -0.9 v. us. the S.C.E. These latter conclusions are surprising and difficult to reconcile with the fact that +5 vanadium in acid medium is a rather strong oxidant, easily reduced to the +4 state, and with the additional facts that the standard potentials of the V+K 4 V+' and V+44V+8couples in acid medium differ by nearly 0.7 v., and those of the V+' 3 V+' and V+a + V+2couples in acid medium are about 0.6 v. apart, so that conditions are very favorable for stepwise reduction. These contradictory conclusions, and the lack of any information a t all on the Polarographic behavior of vanadium in its lower oxidation states, prompted the present investigation.
Experimental Ammonium xneta-te, p d e d by recrystallization, served as the source of +S vanadium. An 0.08 M stock solution was prepared in 1 N sulfuric acid and standardized by the sulfurous acid-permanganate procedure.' A stock solution of vanadyl sulfate in 0.6 N sulfuric acid was prepared by reducing a 2M)cc. w o n of th.e stock ammonium metavandate solution mth sulfur dioxide gas. The excess sulfur dioxide was expelled by boiling the solution, while a stream of purified nitrogen was swept through it, until the duent.gas ceased to reduce permanganate ion. The solution was cooled and diluted to exactly 500 cc. The concentration of +4 vanadium in this solution, determined by p-te titration, checked the value expected from the concentration of the ammonium vanadate solution used in its preparation, and the solution remained unchanged over a period of six weeks in contact with air. A stock solution of vanadic s a t e , 0.01 M in respect to V+++,was prepared in 0.1 N sulfuric acid from a sample of pure V * ( ~ ~ ) * . l O &which O was kindly furnished by Professor Gtlnnell Jones. The preparation and analysis of this salt have been described by Jonesand Colvin.' Since vanadic ion is rather easily air-oxidized,care was taken to prepare and store the solution in an atmosphere of purified nitrogen. The concentration of +3 vanadium in this solution determined by permsnganimetric titration was 0.3% larger than the value computed from the weight of (6) W. F. Hillebrand and 0. E. F. Lundell. "Applied Inorganic Analysis," John Wiley and Sons, Ine., New York, N.Y.,1929. p. 359. (7) G. Jones and J. H. Colvin. TEISJOURIUL, 66, 1568, 1573 (1944).
