PERHALIDE EQUILIBRIUM I N NON-AQUEOUS SOLUTIONS BY ERNEST A. DANCASTER
The increased solubility of the halogens in water brought about by the presence of halide acids or salts has long been the subject of much speculation and experimental work, the most fruitful method of investigation being by means of the application of the Distribution Law of Nernst. This method was first used for the purpose by Roloff' but Jakowkin* was the first to make a systematic study of the equilibrium involved, and the view usually accepted a t the present time, that the halogen dissolved in a dilute aqueous solution of a metallic halide is present chiefly in the form of tri-halide is based mainly upon his work, and that of his successors. This investigator assumed that the thermal dissociation of the tri-iodide takes place according to the equilibrium MI.12SMI
+
12.
and further assumed that both the iodide atoms of a bivalent halide are equally active in combining with the dissolved iodine. The corresponding equilibrium constant is given by the expression K =
{ n a - (b - x ) ) ~ , b-x
where n is the valence of the metal under consideration, a is the original concentration of the halide, b is the concentration of the free halogen in the aqueous layer as determined by titration, and z is the concentration of the free uncombined halogen, all concentrations being given in gramme molecules per litre. Jakowkin's investigations have been confirmed and extended by Bray and MackayJ3Fedotieff4 and others. Herz and Kurzelj studied the case of barium iodide and iodine, keeping the concentration of the iodide constant, while they varied that of the iodine. They, however, assumed that one only of the two iodine atoms of the barium iodide combined with iodine, according to the equilibrium BaIIsBaI2 11,
+
the corresponding equilibrium constant being given by the expression
employing the same equation, Herz and Bulla6 investigated the iodides of calcium, strontium and barium. Although the values of K calculated from this equation, exhibit a fair degree of constancy it has been shown by Van IZ. physik. Chem.,
13, 341 (1894). 2 Z . physik. Chem., 13, 539 (1894); 18, j85 (189j); 20, 19 (1896). J. Am. Chem. Soc., 32, 1207 (1910). Z. anorg. Chem., 69, 122 (1910). 5Z. Electrochemie, 16, 869 (1910). 6 Z . anorg. Chem., 71, 2 5 (1911).
PERHhLIDE EQUILIBRIUM IN SON-AQCEOUS SOLCTIONS
'iI3
Xame and Brown' that a much greater constancy is obtained when Jakowkin's equation is used. These investigators, therefore, conclude that the latter is the correct one, a conclusion that has been confirmed by Pearce and Eversole.2 Van Xame and Brown have also shown that the constant] calculated by Jakonkin, holds not only for the halides of the bivalent metals, but also for the trivalent metal lanthanum. These investigators made a distinct advance by showing that the bromides and iodides of mercury and cadmium behave abnormally] while all the other halides which they examined] or which had been previously examined, gave a normal value for the dissociation constant. They also calculated the percentage of complex molecules and ions, on the assumption that these do not combine with the halogen. Although Jakowkin has studied the dissociation of a few perhalogen compounds of the type XC1Br2, XC112 and XBr12, very little work had been done on the perhalogen compounds containing two different halogens, until Priyadaranjan Rby and Pulin Vihari Sarkar3 investigated the formation and dissociation of the perhalogen acids HClBr2, HCl12 and HBrIz in aqueous solution, and Dancaste+ studied the chloro-perbromide equilibria] and found that the chloroperbromides of mercury and cadmium exhibited abnormalities to those shown by the perbromides and periodides of these metals, and showed that Jakowkin's constant held for the trivalent metal aluminium. Xone of the halides examined, except those of mercury and cadmium, showed any abnormality. Pierce and Eversolej investigated the distribution of iodine between carbon tetrachloride and aqueous solutions of barium iodide a t z j"C. The concentrations were determined on the basis of molecules per 1000 grammes of solvent, with the purpose of eliminating the effects due to variat,ions in the amount of solvent. These investigators found that the distribution ratio of iodine between water and carbon tetrachloride is independent of the concentration of the iodine for the range of concentrations used. This result is at, variance with that found by Jakowkin, and as all the work done by means of the distribution method of determining the dissociation constants of the perhalides in aqueous solution has been based upon Jakowkin's figures, Pearce and Eversole's results are of importance] and should be either confirmed or refuted. From the results obtained Pearce and Eversole conclude that the tri-iodide is the only periodide present in dilute solutions unsaturated with iodine, whilst a mixture of the tri-iodide and penta-iodide is present in solutions saturated with iodine. Dawson and his co-workers6 investigated the existence of periodides in nitrobenzene, and other nitro-organic aromatic compounds, and inferred from the results obtained that periodides as high as the enneaiodide, KI,, are probably present in the solutions. Jones' investigated the existence of hydrogen perbromide in acetic acid and aqueousacetic acid solutions at I 5°C. Most of the investigations have been carried ' A m . J. Sci, liv], 44, xojr (1917). (1924). J. Chem. SOC.,121, 1449 ( 1 9 2 2 ) . ' J. Chem. SOC.,125, 2038 (1924). Loc. cit. J. Chem. Sw., 79, 238 (1901); 81, 524 ( 1902); 85, 79 (1904); 93, 1308 (1908). ' ,I. Chern. Sor., 99, 403 (191I ) .
* J. Phys. Chern., 28, 245
1714
ERNEST A. DANCASTER
out at 25'C., and there are very few figures available for the dissociation constant a t other temperatures. Also, with the exception of the work of Dawson and his co-workers and of Jones, the investigations have been limited to aqueous solutions. The objects of the present investigation were ( I ) to develop a method for the determination of the thermal dissociation of perhalides in non-aqueous solutions, ( 2 ) to ascertain whether the halides of mercury and cadmium in such solutions showed corresponding abnormalities to those shown in aqueous solution, ( 3 ) to extend the number of halides investigated so as to include those which cannot be employed in aqueous solution, (4) to ascertain whether or not any of these halides exhibit any abnormalities such as those shown by mercury and cadmium, and ( 5 ) t o carry out investigations in aqueous and non-aqueous solutions a t various temperatures in order to ascertain whether or not the dissociation constant varied with the temperature, and to compare the values of K obtained in aqueous and non-aqueous solutions. The work, of course, necessitated the determination of the distribution of bromine between the organic solvent and air at the temperature employed, and of the distribution ratio of bromine between water and carbon tetrachloride a t temperatures other than 25OC. Jakowkin's values were taken for the distribution of bromine between water and carbon tetrachloride at 2 5 T . Experimental.-The materials employed were the purest obtainable, and were usually of A. R. quality. The bromides of tin and aluminium were prepared from bromine and the respective metals. All materials were submitted to further purification. A. R. acetic acid gave very erratic results, owing to reaction with the bromine, but after purification by means of the method of Orton and Bradfield' the purified acid was not appreciably affected by bromine, and the results obtained were consistent. The method employed for the determination of the dissociation constant in an organic liquid was a modification of that employed by Jakowkin for the determination of the constant in aqueous solution, in which the distribution of a halogen between an organic solvent, or a solution of a halide in the organic solvent, and the atmosphere above the liquid is determined, instead of the distribution of the halogen between water, or an aqueous solution of a halide, and an organic liquid. I t was necessary that the organic liquid chosen for this purpose should be a fairly good solvent for the halides used, and a very good solvent for bromine, and that this liquid should not be attacked by bromine at the temperature of the experiments. Glacial acetic acid was found to satisfy these conditions better than any other organic liquid tried, and this acid was, therefore, used throughout the investigation, although the solubility of some of the halides was found to be very small. For this reason sodium chloride and aluminium chloride could be investigated in very dilute solution only, and cadmium chloride could not be used because it was found to be practically insoluble in acetic acid. I t was found possible to obtain solutions of from 0 . 0 2 to 0.1 gramme molecule per litre with all the other halides employed, and the chlorides of tin and antimony, and the broI
J. Chem. SOC.,125, 960 (1924).
PERHALIDE EQUILIBRIUM I N NON-AQUEOUS SOLUTIONS
1715
mide of tin, were found to be very soluble. Titanium chloride could not be employed because it was found to decompose acetic acid with explosive violence as soon as the two liquids were brought into contact. The halides investigated were the chlorides of lithium, sodium, potassium, mercury, iron, aluminium, tin and antimony, and the bromides of sodium, potassium, mercury, cadmium, aluminium and tin. At the time this method was developed the author was unaware that it had been previously employed, but has since discovered that the same method, differing only in details, was used by Jones' for the determination of solubility coefficients. The method employed for the determination of the dissociation constant in aqueous solution was the usual one, in which the aqueous solution and the solution of halogen in an organic solvent are shaken up in a stoppered bottle, and left in the thermostat to settle out. The amount of bromine in each layer is then determined by titration with sodium thiosulphate. The determination was carried out as follows. A solution of bromine in acetic acid, or in a solution of a halide in acetic acid, was placed in a wash bottle fitted with a ground-glass stopper, through which passed the inlet and outlet tubes, and the flask was left in the thermostat until the contents had attained the required temperature. I O C.C.of the solution were then withdrawn by means of a pipette, run into an excess of potassium iodide solution, largely diluted with water, and the liberated iodine titrated with N/zo sodium thiosulphate. An aqueous solution of potassium iodide was placed in a second wash bottle, and the two wash bottles were joined up to the rest of the apparatus, consisting of a tower containing calcium chloride and soda lime, an air chamber, an aspirator provided with a thermometer to register the temperature of the interior, and a manometer. The air chamber and the first wash bottle were kept within the thermostat. About z litres of air were slowly drawn through the apparatus. This air was dried and freed from carbon dioxide while passing through the tower, and brought to the temperature of the thermostat in the air chamber. The purified and heated air then bubbled through the solution in the first wash bottle, and the bromine laden air passed through the potassium iodide solution, liberating an equivalent amount of iodine which remained in solution in the second wash bottle. At the end of the experiment a second I O C.C.of the solution was withdrawn from the first wash bottle and titrated. The first and second titrations gave the initial and final bromine concentrations, and the mean of these two results was taken as the average bromine content of the solution during the experiment. The liberated iodine was also titrated in order to ascertain the amount of bromine removed, and the concentration of bromine in the atmosphere in contact with the solution was calculated from the volume of air passed through the apparatus, after applying corrections for temperature and pressure, and also for the vapour of water in the aspirator, the vapour pressure of acetic acid in the first wash bottle, and the volume of bromine in the atmosphere in contact with the solution. The distribution coefficient of bromine between acetic acid and the atmosphere in contact with the solution was determined at I~OC, zoo, 25', LOC.cit.
1716
ERNEST A. DANCASTER
30' and 40'. This atmosphere, of course, consists of a mixture of air, acetic acid vapour and bromine vapour in equilibrium with the liquid, but it will, in this paper be described simply as air. The results are shown in Table I, where B gives the concentration of the bromine in acetic acid, and G gives the concentration of the bromine in air. In both cases the concentrations are given in gramme molecules per litre.
TABLEI The Distribution of Bromine between Acetic Acid and Air B Ratio G At 40'
0.39274 0.26092 0.23 I34 0.15i91 0.12658 0.12077
0.10969 0,10331
0.09808 0.08414 0.06863 0.04696
At 30'
0.29630 0,23382 0.21410 0,17931 0.16350 0.14762 0 .'3659 0.12125
0.113c)O 0.09804 0.07906 0.05506 0.04926 0.04017
At
25'
0.89242 0.6417 5 0.42800 0.22809 0.18107 0 .I5354 0.14911 0.13846
0.001422
0.00093I 2 0.0008295 0.0005640 0.0004583 0.0004333 0.0003 982 0.0003930 0.000353 7 0,0003049 0.0002469 0.0001706 0 .0007056 0.0005577 0 .ooo5102
0.0004269 0.0003922 0.0003518 0.0003284 0.0002894 0,00027 I 3 0.0002334 o .0001885
276.2 280.2 278.9 280.0
276.2 278'7 275.5 277.0 277.3
275.8 278 .o 275.3 Average 277.4 419.9 419.2 419.6 420.0 416.9 419.6 415.9 418.9 419.8 420.0
419.4
0.00013 I I
420.0
0,0001 182
416.7 418.9 Average 419.I
o.00009589 0.001707
0 , O O I23 7
0.0008228 0.0004400 0.0003498 0.000295I 0.0002880 0.0002673
522.8 518.9 520.2
518.4 517.6 520.3 517.7
518.0
PERHALIDE EQUILIBRIUM IN NON-AQUEOUS SOLUTIOSS
1717
TABLE I (continued) The Distribution of Bromine between Acetic Acid and Air G Ratio B 0.12j38 0.0002414 519.4 0.12303 0.0002369 519.3 0 .I 1643 0.0002242 519.3 0.000206j 522.8 0.10795 At
zjo
0.10360 0.09470 0.0900j 0.08544 0.08161 0.0686j 0.04272
0.0002000
ji8.0
0.0001830 0 . O O O I74j O.OOOIj7j
517.3 516.0 519.4 518.2
0,0001324
518.5
0.0001645
519.2
o.00008228
Average 5 19. o At
20'
0 .j8243 0.28653 0.22421 0.17513 0 .I3388
0.0009129 0.0004466 0.0003485
O.II25j
o 1023j 0.09453 0.08639 0.07937 0.07158 o.oj893 0.04605
At
Ijo
638. o 641.6
0.0002754
643.3 638. o
0.0002088
641.1
0.0001 760 0.0001 599
639.2 640.I
0.0001484 0.0001347 o.0001240 0 .OOOII I 8 0.00009225
637. o 641.3 640.I 640,~ 638.8 o.oooo7~1o 638' 7 Average 639.8
0.59393
o .0007480
0.35998
0 .00045 j4
794.0 790.5 792 I
0.19186 0.13791 0.12326 0.10970 0.1046j 0.09215
0.0001560 0.0001384 o.00013I 8 o ,0001 166
790.1 792.6 794.0 790.3
0.08846
0.0001 I I7
79r .D
0.07766 o.oj806
o.0000978~
0.04712
0.0002422 0,000I 74 I
'
79".
1
793.5 793.8 O.OOOOj945 792.6 Average 792 ,3
0.00007314
I t is evident that this ratio is constant at any given temperature for all concentrations of bromine within the range of the experiments. The solu-
1718
ERNEST A. DANCASTER
bility of bromine in acetic acid therefore follows Henry's law. This ratio is, therefore, a measure of the partial vapour pressure of bromine a t different temperatures, and the temperature variation over the range in which the latent heat of vaporization is constant should be given by
Q LogioR = 4.6
I+C . T
It was found that between 15' and 4o'C. 1650
Logl,R = - - 1.828.
T
The values of R obtained by experiment agree fairly well with those calculated from this equation, as shown below. Temperature "A 288 293 298 3 03 313
R (observed) 792 640 519 419 277
R (calculated) 796 63 6 513 415 2 78
The agreement between the values of R obtained by experiment and those obtained by calculation is fairly satisfactory. The latent heat of bromine in acetic acid solution, obtained from the above values, is 48 calories, a value which is not far from that of liquid bromine (45.6 calories). In order to ascertain the effect due to the presence of an indifferent salt, i.e. one that does not combine with bromine to form a perhalogen compound, solutions of sodium acetate, disodium hydrogen phosphate and potassium acetate in acetic acid were employed in place of the pure acid. The sulphates and nitrates of sodium and potassium were also tried, hut proved to be too insoluble in acetic acid to be of any value. Tho results obtained are given in Table 11, where A is the concentration of the salt, B is the concentration of tho biumine in the salt solution, and G is the concentration of the bromine in air. All concentrations are given in gramme molecules per litre. It is evident from the figures given in Table I1 that the dissolved salt has no appreciable effect upon the distribution coefficient. The ratios given in Table I can, therefore, be used for the calculation of the dissociation constants of the perhalides examined. The distribution of bromine between solutions of lithium chloride in acetic acid and air at 2 j"C was now determined. The concentration of the lithium chloride was varied from o 09367 to o 04152 gramme molecules per litre and the concentration of the bromine in the solution from o 12354 to o 08685 gramme molecules per litre. Although the perhalide formed in
PERHALIDE EQUILIBRIUM IN NON-AQUEOUS SOLUTIONS
1719
TABLEI1 The Distribution of Bromine between a Solution of a Salt in Acetic Acid and Air a t 2 j o B
A
Sodium acetate 0.97760 0.60128 -
0.0003 19I
0,16462 0.15938
515.9 518.6 520.9 517.7 519.2
0.0003070
0.00033IO 0.0003246 0.0003234 0.0003I 60
0.17242
0.16803 0,16786 0.16349
0.48830 -
Ratio
G
517.2
Average 5 I 8 . 2 Disodium hydrogen phosphate 0.23550 0.10309
-
0.0001987
518.8
0,09991 0.09662
0.0001925 519.0 0.0001861 5'9.1 Average j 19.o
0,11147 0.10684
0.0002143
jzo.I
o.000206I 0.0003990 0.00039I9 0.0002906 0.0002836 Average 5 I 8 .6
518.4
Potassium acetate 0.80212
0.51255 -
0.20694 0.20258 0.15094 0.14700
0.12814 -
518.7
516.9 519.4 518.3
dilute aqueous solution is the tri-halide, it does not follow that this will also be the case in acetic acid solution. Therefore, four dissociation constants were calculated, according to the equilibrium LiCl
+ mBrt-LiCl.Brz,,
(m
= I, 2,
3, 1).
The corresponding equilibrium constant is given by
K =
b-x m
where the concentration of the free bromine (x) is found by multiplying the observed concentration in air by the distribution coefficient. The results are shown in Table 111.
E R S E S T A . DAPI’CASTER
I720
TABLE111 The Dissociation of Lithium Chloro-Perbromides at
B
G a = 0.09367M - LiCl
s
K,
Kz
0.11163 O . O O O I ~ O ~0.08854 0.2707 0.05579 o.ro8jo 0.0001662 0.08626 0.2661 0.05468 0.10585 0.0001613 0.08371 0 . 2 7 0 5 0.0551~
Average 0.2691 0,05519 a = 0.07862M
25’
KI
0
0
00775 00740
0 00721
0
00745
Kt
0.000936 0,000869 0.000822
0.000876
- LiCl
0.09072 0.0001416 0,07348 0.2616 0.04395 0 00503 0,08879 0.0001393 0 . 0 7 2 3 0 0.2681 0.04359 0 oojoz 0.08685 0.0001361 0.07064 0 . 2 7 2 0 0.04344 o 00478
Average 0.2672 0.04366
0
00494
0,000503
0.000496 0.000460 0.000486
a = 0,05968M - LiCl 0.11188 o.oo01850 0.09602 0,2653 0.0601j o 00910 0 . 0 0 1 2 0 0.10983 0.0001813 0.09409 0.2625 0.05828 o 00864 0,001I I 0.10j32 0.0001jj6 0.09217 0 . ~ 7 0 9 0.05839 o 00847 0.00106
Average 0.2662 o ,0589j a = 0.04152 M - LiCl 0.12354 0.0002146 0.11138 0.2689 0.12056 o.000~091 0.10852 0.26j7 0,11749 0.0002040 0.10588 0.2672
o
00874
0.00112
0.08020 o 01278 0.00195 0.06945 o 01196 0.00Ijj 0.06902 0 01154 0.00167
Average 0.2673 0.07289 o 01209 o.00180 The values of K2, Ks and K I increase rapidly with increase in the concentration of the halogen, and as the concentration of the lithium chloride decreases these values fall to a minimum a t a concentration of 0.07862 gramme molecules per litre, and then rise with further decrease in the concentration. On the other hand, the value of K1 remains fairly constant, and is evidently unaffected by alterations in the concentration of either the bromine or the halide, within the limits of the experimental conditions. It is, therefore, obvious that the principal perhalide formed is the trihalide, LiClBrZ, and i t was not considered necessary to calculate the values of K other than KI in the case of the remaining chloro-perbromides examined. In order to save space the remaining data for the chloro-perbromides are summarised in Table IV. The values quoted for K1 are the mean of usually three experiments, and the variation in KI is shown in the last column.
PERHALIDE EQCILIBRIL‘M IK NON-AQUEOUS SOLUTIONS
I721
TABLE IV The Dissociation of Chloro-Perbromides at 2 5’ Substance
X’aClBrZ 1,
KClBrz f,
A1Cl3Br6 f,
HgClzBrl 11
FeC13Brs
,, ft
SnCI,Br8
,, ,,
SbClsBrlo 1
Range
K,
A B molality molality 0.1oo50 0,1393 - 0.1456
0.262
0.261
-
0.00948 0.1193 -
0.260
0.258
-
0.261
-
0.1262
0.01225
0.1317
-
0.01008
O.IOOO
-
0.1071
0.03794 0.14995 0.01897 0.18785 -
0.1584
0.02676 0.1675 0.02674 0,1553 0.01531 0.0812 0.01482 0.0664 -
0.1782 1.422 1.422 0.1625 1.419 1 . 4 1 1 1.056 1 . 0 5 2 0.0852 0.0708 1.022 0.994
-
0.1866 3,0835 3.070 0.1612 2.391 2.386 0.11joj 1.749 1.692
0.02640
0.1302
0 . 0 ~ ~ 50 5. 1 4 1 1
0.05078 0.1824 0.02734 0.155;
-
-
0.01102
0.1118
0,05120
0.1121 0,1060 -
0.02550
0.1364 0.263 0.1496 0.262 0.1383 0.267 0.271
0.25;
0.265 0.264
3.4465 3.368 0.1959 2.994 2.972
0.11jj 0.1117
3.159 3.138 2.408 2.3945
-
-
-
0.263 0.261
0.2675 0.266
0.269 0.276 3’5475 3.008 1.423
1,424 1.059 1.0665
-
3.097 2,405 1.777
-
3.190
-
2.417
It is evident that the chlorides of sodium, potassium and aluminium behave “normally,” i.e. they give a value of the dissociation constant K which agrees with that given by lithium chloride, at any rate at the concentrations examined, which were the greatest that could be obtained owing to the low solubility of these chlorides in acetic acid. On the other hand, the chlorides of mercury, iron, tin and antimony all give values that are considerably higher than those given by the “normal” halides. The values of K given by ferric chloride are distinctly lower than those of the other abnormally behaving halides examined, though they are still considerably higher than the normal value. In each case the value of K increases with increase of concentration of the halide. The distribution of bromine between solutions of potassium bromide in acetic acid and air at 25’ was next determined. The concentration of the potassium bromide was varied from 0.04199to 0.00840 gramme molecules per litre, and that of the bromine from 0.22371to 0.08573gramme molecules per litre. -4s in the case of lithium chloride, four dissociation constants were calculated, corresponding to KRr
+ mBr2
The results are given in Table V.
KBrZ,+,,
(m =
I,
2,3,4).
ERNEST A. DANCASTER
I722
TABLE V The Dissociation of Potassium Perbromides a t B
G
X
KI
Kz
25'
KI
K,
a = 0.04199 ILI - KBr. 0.11294 o.0001160 0.06020 -0.o119 0.00220 o.oo0307 o.0000~91 0 . 1 1 1 2 2 o.0001148 0.05958 - O . O I I I 0.00221 o.oo0305 o.0000284 0.09829 O . O O O O ~ ~ Z0.05149 -0.00529 0 . 0 0 2 1 1 0.000~31o.000018~ I 0.09709 0.0001008 0.05232 -0.00325 0.00239 0.000260 0.0000206 0.09592 0.00009720 0.05045 -0.00386 0.00215 0.000227 0.0000174 Average -0.00708 0.00221 0.000266 o.oooozz7 a = 0.04170M - KBr. 0.22371 0.0002854 0.14813 -0.00664 0.00227 0.00213 0.22000 0.0002790 0.14480 -0.00645 o.00~280.00201 0.21439 0.0002688 0,13951 -0.00618 0.00221 0.00182 Average - 0.00642 0.00225 0.00199
o.000575 o.oo0540 o.oo04g1 0.000535
a = 0.02359iL1 - KBr. 0.09101 o.0001183 0.06140 -0.0125 0.00224 0.000322 o.0000311 0.08835 0.0001147 0.05953 -0.o108 0.00226 o.oo0307 o.0000285 0.08573 0.0001111 0.05766 -0.0091~ o.00~260.000~91o.0000261 Average - 0.0I 08 0.00225 0.000307 0.0000286 a = o.02100 M - KBr. 0.15769 0.0002340 0.12145 -0.0511 0.15495 0.000~~840,11854 -0.o502 0.15223 0.0002245 0.11652 -0.0480 0.11320 0.0001568 0.08138 -0.0277 O . I I I I Z o.0001540 0,07993 -0.0262 0.10873 0.0001476 0.07661 -0.0267 Average - 0.0383 a = 0.00840M - KBr. 0.0002658 0.13795 -0.0613 0.14902 0,0002582 0.13400 -0.0586 Average - 0.0599 0.15307
0.00234 0.00216 0.00238
0.00132
0.000287
0 . 0 0 1 ~ o.000258 ~
0.000249 0.000528 o.0000718 0 . 0 0 2 ~ 10.000521 o.0000691 0.00181 0.000436 o.0000556 0.00217 0.000872 0.000165 0.00121
0.00212
0.00211 0,00210
0.00211
0.00175 0.000443 0.00159 0.000386 0.00167 0.00041g
I t is obvious that in the case of potassium bromide the values of KI, Kt and K 4 vary greatly. The values of K, are all negative, and decrease rapidly with increase of bromine concentration, whilst the values of K Band K 4are positive, and increase under these circumstances. The variations in the concentration of the halide also appear to have an effect upon these values, but, as shown below, if the concentration of the bromine is kept constant, an increase in the concentration of the halide causes a corresponding increase in the value of K1 and a decrease in the values of KB and K4.
PERHALIDE EQUILIBRIUM IN NON-AQUEOUS SOLUTIONS
Concentration of Br in solution
Concentration of Halide
0.00840
KI
I723
KS
K4
0,15307 0,15495
-0.0613 -0.0502
o.oo17j
0.000443
0.02100
0.00122
0.0002 58
0.02100
0.11112
-0.0262
0,04199
0.11122
-O.OIII
o.oooj21 o.ooo3oj
o.0000691 o.0000284
The value of K Zis not appreciably affected by variations in the concentration of either the halide or the bromine, and remains fairly constant. I t therefore appears that in t,he case of potassium bromide it is the pentabromide that is the principal product in solutions moderately dilute with respect to bromine, and not the trihalide, as in the case of the chloro-perbromides. The values of K,, Kz and K 3 were calculated in the case of all the other bromides examined. The bromides of sodium and aluminium were found to behave in the same manner as potassium bromide, the value of Kz remaining unaffected by variations in the concentrations of either the halide or the bromine, whilst’ the values of K1 and Ka vary considerably, and in the same manner as in the case of the potassium salt. These halides are, therefore, “normal” salts. The bromides of cadmium, mercury and tin give abnormal values for K1, Kz and K3. The data for the perbromides are summarised in Table VI. In order to save space the values of K, and K8 are not, shown, and the values quoted for K Pare the mean of usually three experiments. The variation in Kz is shown in the last column. TABLE VI The Dissociation of Perbromides a t 25’ A molality
KI
B molality
Range
0.02419
0.1905
- 0.1994
0.00226
0.002Ij -
0.00964
0. I478
-
0.I5395
o .00228
0.00224
0.00233 - 0.00234
0.0327 I
o ,00226
0.00226
- 0.00227
0.0022j
0.30773
0,1735 - 0,1759 0.1658 - 0 . 1 7 1 0 0.1548 - 0.1602
0.00223 0.00219 -
0,02977 0.01488
0.1306 - 0.1376 0.1707 - 0.1788
0.257 0 .I775
0.241
- 0.265
0.172
- 0.184
0.038jog 0.038705 0.03413 0.01935 0.017065
0.17075 0.07875 0.1676 0.0916 0 .I 528 -
1.696 1.640 1.5515 1.2555 1.187
-
0.200
0.433 0.295 0.193
0.152
0.145
-
0 .or478
0.11558 0.04973 0.024865 0.01243
0.1770
1.747
0.0805
I ,689
0.1762 0.0962 0,1567
561 ,280 1 .I94
- 0.1211 - 0.1288 0.1088 - 0.1148 0.1376 - 0 . 1 4 1 5 0.1151
0.1229
0.00223
I . I
0,437 0.310
0.00228
0.00228
1.7915 1.748
1.567 1.307 1.201
0.441
0.319 0.206 0.159
1724
ERNEST A. DANCASTER
The dissociation constants of lithium chloro-perbromides and potassium perbromides in acetic acid solutions at various temperatures were now determined, the temperatures chosen being 40°, 30°, 20' and 15'. In each case four dissociation constants were calculated, corresponding to the four constants already calculated at 2 5 O , the results obtained were similar to those found with these two compounds a t the latter temperature. The data for lithium chloro-perbromide are summarised in Table VII, and those for potassium perbromide in TableVIII. Thevalues of Klin the former case and of Klin the latter case are the mean of usually four experiments, The values of K2, Ka and K, for lithium chloro-perbromide and of KI, Ks and K, for potamium perbromide have been omitted.
TABLE VI1 The Dissociation of Lithium Chloro-Perbromide a t Different Temperatures Temperature
A
molality 40'
30° 20° ZOO
I 5'
0.06581 0,04396 0.09367 0.07545 0~05540
B
molality
0.09285 0.0762 0,1713 0.1820 0.1239
- 0.1065 - 0.0832 - 0.1856 - 0.1885 - 0.1289
K,
0,266 0.268 0.268 0.266 0.268
TABLE VI11 The Dissociation of Potassium Perbromide at Different Temperatures Temperature
A molality
40° 30'
0,03299
20'
0.02090 0.02492
15'
0.02250
B
K1
molality
0.1067 - 0 . 1 1 7 7 5 0.1167 - 0.1259 0.1614 - 0.1696 0.1328 - 0.1345
0.00218 0.00216 0.00227
0.00223
It is evident from these results that the dissociation constant does not vary with the temperature within the range of the experiments, and that over the whole range of temperatures considered the chloro-perbromides exist principally as the tri-halides, whilst the perbromides exist mainly as pentahalides. In order to compare the results obtained in acetic acid with those obtained in water, the dissociation constants of lithium chloro-perbromide and potassium perbromide in aqueous solution were now determined. As already stated, this determination was carried out by the method of Jakowkin. Determinations were made a t 4ob, 30°, 25', 20' and IS', but as satisfactory results had already been obtained with potassium perbromide at 25' it was not thought necessary to carry out any further experiments a t this temperature with the latter compound. Before the values of K could be calculated it was
PERHALIDE EQUILIBRIUM Ix’ XON-AQUEOUS SOLUTIOSS
I725
necessary t o determine the distribution ratios of bromine between water and carbon tetrachloride at these temperatures, at any rate for a range of concentrations somewhat exceeding the limits of those employed in the determination of the dissociation constants. These distribution ratios are given in Table IX. The ratio a t 2 5’ was not determined, Jakowkin’s figures being used instead. The data for lithium chloro-perbromide and potassium perbromide are summarised in Tables X and XI, the values quoted for KI being the mean of usually four experiments.
TABLE IX The Distribution of Bromine between Water and Carbon Tetrachloride a t Different Temperatures At 40’
B
G
o 02088
o 6j767 o 60117 0.50610 0.39840 0,34221 0.30066 0.22656 0.186jo
01929 o 01602
0
0.01302 0 . 0 1I 2 0 0 .OIOOI
0.007601 0.006229 0.005820
29.81
0.17240
29.94 29.62
0.004245
0.122jS
28.92
0.00345 5
0.10124
29.30
0 . 0 0 2I52
0.06287 o .oj569 0.03514 0,03079
29.22
0.99282
33,56 31.84
0.001 920 0,001 208
o.001060
4 t 30’
Ratio
32.46 31 16 31 60 30.61 30.55 30.05
0.02959 0.0238 j
29.00
29.15 29.04
0.007730
0.75958 0 . 6 j194 0.60064 0 44849 0.38200 0.35117 0,27345 0.22070
0.0072 j I
0.20962
0.005912
0 .I 6860
28.52
0.004533 0.003528 0.002626 0.002065
0 . 1 2 7 73
28.18
0.09826 o.oj301
27.85 2 7 .oo
o.oj786
28.02
O.OZIj4
0.01972 0 .OI533 0.013I3 O.OI2Oj
0.009427
’
30.27
30.45 29.26 29.I O
29.15 29.00
28.55 28.91
1726
ERNEST A. DANCASTER
TABLE IX (Continued) The Distribution of Bromine between Water and Carbon Tetrachloride a t Different Temperatures B G htio At 20' 0.77206 0.02535 30.45 0.68113 0.52807 0.35849 0.21687 0.20653 0,17344 0.13826 0.13766
0.02320
0.01865 0.01301
0.008125 0.007845
0.006504 0.0052 58
0.00529I 0.004202 0.003 236
At 15'
29.36 28.32 27.55
0.08486
26.69 26.33 26.67 26.29 26.02 26.21 26.22
I . 12675 0,95247 0.81538 0.66542 o ,60013
31.47 30.42 28.80 28.49 27.33
0.55512
27.21
0.51684 0.47580 0.44815 0.40880 0.32756 0,28768 0.26089
26.89
0.1101g
0.03580 0.03I 3 I 0.03018 0.02336 0.02196 0.02040 0.01922 0.01762 0,01595 0.01521
0.01219 0.01096 0.0099I I
27 .oo 27.27
26.88 26.87 26.26 26.32
TABLE X The Dissociation of LiCIBrs in Aqueous Solution at Different Temperatures Temperature
40° 30° 25"
20° 1So 1
so
A
molality
0.05458
0.06043 0 .IO2 54 0.08849 0.09358 0,08978
€3
molality
0.0031440.0063910.005280 0.009785 0.01124 0.007046 -
o.oroj3 0.01120 0.01302
0.02783 0.02279 o.01081
KI
0.7455 0.747 0.747 0.745 0.742 0.745
TABLE XI The Dissociation of KBr3 in Aqueous Solution at Different Temperatures Temperature
A molality
40°
0.0~001
30' 20°
0.0471 2 0.02036 0.06704
1 5 O
B
molality
- 0,01909 - 0.02104 0.003980 - 0 . 0 1 2 1 3 0.007610 - 0.02935 0,005819 0.01030
K1
0.06325 0.0632 0.0628 0.0630
PERHALIDE EQUILIBRIUM IN NON-AQUEOUS SOLUTIONS
1727
Discussion of Results The method of determining the dissociation constants of perhalides in non-aqueous solutions used in this research has been found to give satisfactory results. I t is, therefore, a practical method, and can be used in other cases where it is desired to investigate complex molecules and ions in nonaqueous solutions by means of the distribution method. The method is, of course, limited to those substances which are sufficiently soluble in the liquid chosen, and also in air. The results obtained for the distribution of bromine between water and carbon tetrachloride a t different temperatures show that in each case the ratio increases as the concentration of the halogen is increased, and, therefore, agree with those obtained by Jakowkin a t 2 5 ’ . The curves obtained at the various temperatures are approximately parallel at the comparatively low concentrations investigated; thus indicating that at these concentrations the increase in the distribution ratio of bromine is approximately constant at all temperatures between 15’ and 40’. The distribution of bromine between acetic acid and air does not show this increase, but remains constant with increase in the concentration of the bromine. The values obtained for the dissociation constant of perhalides in acetic acid solution show that in the case of the chloro-perbromides the principal halide existing in the solution is the trihalide, as it is in the case of these compounds in aqueous solution. The results obtained with the perbromides, however, indicate that these compounds exist in acetic acid solution principally as pentabromides, whilst in aqueous solution, under similar conditions, they exist mainly as tribromides. Although the dissociation of periodides in acetic acid solution could not be determined by the method employed because of the low solubility of iodine in air, it appears probably that they, too, would be found to exist as pentd- or higher perhalides in this solvent. The results obtained by Dawson and his co-workers1 are not inconsistent with the view that in some organic solvents the perhalides tend to form compounds richer in halogen as we pass up the series from the chloro-perbromides to the periodides. Where the same halides have been investigated in both aqueous and acetic acid solutions it has been found that those which give normal values of the dissociation constant in the one case also give normal values in the other, and those which giie abnormal values in aqueous solution also give abnormal values in acetic acid. It is, therefore, probable that the other “normal” halides investigated in aqueous solution, but not in acetic acid, would also behave normally in this respect in the latter solvent; and that because aluminium perbromide has been found to give a normal value in acetic acid, it would also be found to give a normal value in aqueous solution. When the values of the dissociation constants obtained with the “normal” halides in aqueous and acetic acid solutions are compared, it is found that those ob-
‘ LOC.cit.
ERNEST A. DANCASTER
1728
tained in the latter case are much smaller than those obtained in the former. Tab!e XI1 shows the mean values of Kl obtained with the chloro-perbromides a t 25’ in aqueous and acetic acid solutions.
TABLE XI1 Mean Values of K1 obtained with Chloro-Perbromides in Aqueous and Acetic Acid Solutions a t 25’ Salt
LiCIBrp KClBrz AlC13Bre
KI (aqueous)
K, (acetic acid) 0.75 o 267 o 73 (Jakowkin) o 2 6 3 0.72
o 269
This difference in the value of K1 indicates a smaller degree of dissociation of the perhalides in acetic acid than takes place in aqueous solution. In both aqueous and acetic acid solution the value of K given by the “normal” perhalides is found to be constant at temperatures ranging from 15’ to 40’. Jakowkinl found the value of K at 36.5’ for KBr3to be 0.069, whilst the value at 2 5 ’ for the same compound was only 0.063, and, therefore, concluded that the dissociation constant of the perhalides increases with increase of temperature. This is a result which might be expected, but it may be pointed out that the value of K at the higher temperature was not determined with anything approaching the accuracy of that at 2 s 0 , and the differences between individual values obtained a t the latter temperature are sometimes as great as that found between the values obtained a t the two temperatures. The dissociation constants of certain halides were found to have a much greater value than that given by “normal” compounds, whether the chloroperbromides or the perbromides of the metals in question were chosen for investigation. The compounds of cadmium, mercury, iron, tin and antimony all showed this abnormal behaviour, which corresponds closely with that shown by the chloro-perbromides, perbromides and periodides of cadmium and mercury in aqueous solution. As in the case of the “normal” perhalides, the chloro-perbromides of these abnormally behaving metals evidently exist in acetic acid solution mainly as the trihalides, whilst the corresponding perbromides exist as the pentahalides. In each case the value of K increases considerably with increase in the concentration of the halide. The mercury salts show the greatest variation from the normal values. The only compounds which have shown this abnormal behaviour in aqueous solution are the perhalides of cadmium and mercury, but it is evident from the results obtained in this investigation that the number of “abnormal” compounds must be extended to ferric, stannic and antimonic perhalides, and it is probable that there are still others. The three new “abnormal” salts could not be examined in aqueous solution because they are not stable in water. 1
Z. physik. Chem. 20, 19 (1896).
PERHALIDE EQUILIBRIUM I N NON-AQUEOUS SOLUTIONS
I729
TABLE XI11 The “Active Fraction” of Abnormally Behaving Chlorides in Acetic Acid Solution at 25’ Chloride
[ZClBr,]
lBr~l
[ZCI] IZCl]+[ZC1Br2] “Active calculated fraction” %
Mercuric chloride, HgC12 0,03794
0.01897
0.00318 0.00325 0.00312 0.00232
0.00225 0.0022 I
0.155’8
0.00543
0.15072
0.00571
0.14683 0.19358 0.18964 0.18564
0.00563 o ,003I 7 0.00314 0.00315
0.16966 0.16406 0,15944 0 .I 5466 0.15113 0.14771 0.08190 0.07801 0.06814 0.06602 0.06373
0.01337
0.17563 0.17169 0.15461 0.15191 0.14926 0,11439 0.11184 0.10909
0.01659 0.01647
0.10909 0.10665 0.10396 0.10629 0.10344 0.10090
0,02089 0.02058 0.02085 0.01339
0.00861 o ,00896 0.00875 0,00549 0.00539 0.00536
11.34 11.82 11.53 14.47 14.21 14.13
0.02193 0.02186 0.02154
27.32 27.23 2 7 .83 26.49 26.64 26.33 30.35 30.35 29.26 31.24 30.75
Ferric chloride, FeCl, 0.026760
0.026740
0.015309
0.014820
o ,00856 0.00836 0.00809 0.00783 0.00776 o ,00756 0.00329 0.003 17 0.00266 0.00277
0.00265
0.01350
0,01345 0.01342 0.01361 0.01356 0.01065 0.01077
0.01035 0.01112
0 .OIIOZ
0.02125
0.02137 0.02112
0,01394 0,01394 0.01301 0.01389 0.01367
Stannic chloride, SnCh 0.050782
0,027342
O.OIIOZO
0.01099 0.01067 0.00662 0.00656 0.00644 0.00266 0.00261 o ,00267
0 .OI135
0.01144 0.01143 0.00616 0.00618 0.00649
0.02758
0.02714 0.01797 0 .or800 0.01787 0.00882 0,00879 0.00916
‘3.58 ‘3.36 16.43 16.46 16.34 20.01
19.94 20.78
Antimony chloride, SbCls 0.051196
0.025502
0.00860 0.00828 0.00818 0.00537
0.00528 0.00~12
0.01353
0.02949 0.02886 0.02903 0.01876 0.01881
0.01345
0.01857
11.52 11.27
11.34 14.71 14.75 14.57
1730
ERNEST A. DANCASTER
TABLE XIV The “Active Fraction” of Abnormally Behaving Bromides in Acetic Acid Solution a t 25’ Bromide
fzBrjl
lBrnl*
IB-I
calculated
[ZBrl +pBrr]
Active fraction %
Cadmium bromide, CdBrz 0.02977
0.00360 0,00343 0.00354
0.021412 0.016194 0.015259 0.029119
0 .ooog8 0.00048
0.027755
0.00033
o.oz6j45
0.00033
0.00398 0.00391 0.00406 0.00438 0.00446 0.00424
0.01488
0.00407
0 .OOOIO
0 .oo142
I .83
0 .OOOIO
0.00135
0 . 0 0 0 27
0.030424 0.029404 0.028296 0.006379 o ,0061I 7
1.74 1.99 0.53 0.48
0.00131
0.030121
0.00010
0.00154 0.00041 0.00037 0.001 41
0.00123
o ,028716
0.00010
0.00133
0.00117
0.027322
0.00010
0.00127
0.00028
0.009149
o.00007
0.0003j
0.00025
0.008727
0.00031
o.00025
0.00829j 0.024129 0.022952
o.00006 o.00006 o.00006 o.00006
1.95 I .86 0.90 0.80
0.00031
0.80
0.00074 0.00069
2.17
0.00235 0.0023j o .oo146 0.00135 0.0013’/
0.01281
2.77
o ,0119j o ,00964 0.00876 0.00853 0.00639 0.0058j
2.59
0.00413 0.00391
0.00052 0 ,0003I
6.68 6.57 6.82 14.72
14.99 14.25
Mercuric bromide, HgBrz o ,038705
o ,00132 0.00125 0.00127
0.038705 0.034129
0.019350
0.017065
o.00030
0.00068 0.00063
0.0002
7
0.00011
0.00010
2.07
2.02
Stannic bromide, SnBrl 0.115j78 0,049730
0.024865
0.012432
0.01046 0.00960 0.00818 0.00741 0.007I 6 0.00529 0.00481 o ,00468 0.00527
0.00467
0.01003I 0.009189 0.012638 0.012358 o ,011788 0.010862 0.0104j3 0.009898 0.017142 0.016461
0.001IO 0 .oo104 0.00106 0.00069 0.00064
0.00574
0.00596 0.0053I
4.85 4.40 4.29 6.42 5.88 5.77 11.98 10.68
Van Name and Brown‘ attribute the abnormal behaviour of the salts of cadmium and mercury towards the trihalide equilibrium to the power to unite with the halogen being limited to the normal molecules and ions, and have shown that if we accept this view we can calculate that portion of the total halide concentration which is in the form of simple molecules, which they call the “active fraction.” If we make the reasonable assumption that ‘Am.J. &i., [iv], 44, 105 (1917).
PERHALIDE EQUILIBRIUM I N NON-AQUEOUS SOLUTIONS
‘73’
the abnormality shown by these halides in acetic acid solution may be attributed to the same reason as that shown by them in aqueous solution, we can calculate the “active fraction” in this case also. The chloro-perbromides exist in acetic acid solution as the trihalides; the calculation is, therefore, the same as in the case of the salts in aqueous solution, except that Brl is obtained by multiplying G by the distribution coefficient of bromine in the former case. The perbromides, however, exist as the pentahalides in acetic acid solution, and must be calculated from the equation.
where [ZBr] and [ZBrs] represent the total concentrations of the bromide and pentabromide radicles. Table XI11 shows the “active fraction” of the abnormally behaving chloro-perbromidw, and Table XIV shows those of the abnormally behaving perbromides. Owing to the low solubility of cadmium and mercuric halides in acetic acid the results obtained are not so reliable as some of those obtained in aqueous solution. I t is, however, obvious that in all the abnormal salts examined the magnitude of the “active fraction” decreases with the concentration of the halogen, and increases as the concentration of the halide is reduced. This result agrees with that found in aqueous solution, and is to be expected. I t is also found that where the chloride and bromide of the same metal have been examined, e.g. in the case of mercury and tin, the magnitude of the “active fraction” is much larger in the case of the chloride than it is in that of the bromide. This again agrees with the results found in aqueous solution. I n this case it was also found that the “active fraction” was greater in the case of the bromide than in that of the iodide. The magnitude of the “active fraction,” therefore, decreases as the atomic weight of the halogen in combination with the metal increases. The value of the “active fraction” in pure solutions of the halides could not be calculated because a sufficient range of concentrations of bromine, expecially of very low concentrations, had not been investigated to enable this value to be obtained. S-arg
A modification of the distribution method of investigating the dissociation of the perhalides, in which the atmosphere in contact with the solvent takes the place of one of the liquids employed in the usual method, has been developed and utilized for the examination of the dissociation of chloro-perbromides and perbromides in glacial acetic acid. It has been found possible by means of this method to extend the investigation to a number of halides which cannot be examined in aqueous solution, e.g. the chlorides of iron, tin and antimony, and the bromide of tin. It has been found that those halides which give a normal value of K in aqueous solution also give a normal value in acetic acid, but that the magnitude of K is much less in the latter case. I t has also been shown that whilst
I732
ERNEST A. DANCASTER
the chloro-perbromides exist principally as trihalides in both aqueous and acetic acid solution, the perbromides exist mainly as tribromides in aqueous solution, but as pentabromides in acetic acid. The list of “normal” perhalides investigated has been extended to aluminium bromide and lithium chloride, which had not been previously investigated. The halides of cadmium and mercury, which give abnormal values of K in aqueous solution, also give abnormal values in acetic acid, and in addition to these, the chlorides of iron, tin and antimony, and the bromide of tin were also found to give abnormal values. In every case the value of K is much greater than the normal value, and it increases with the concentration of the halide. The percentage of simple molecules and ions in acetic acid solutions of the abnormally behaving salts has been calculated by Van Name and Brown’s method, and it has been found that this percentage decreases with decrease in the concentration of the halogen, and increases with decrease in the concentration of the halide. The percentage is greater in the case of the chlorides of these metals than in that of the bromides. The values of K for “normal” halides have been determined in both aqueous and acetic acid solutions, a t temperatures ranging from 15’ to 40°, and it has been found in both cases that there is no appreciable change in the magnitude of these values over the whole range of temperatures employed. The author wishes to express his thanks to Dr. G. Senter and Dr. S. Sugden for their valuable advice and guidance throughout the work. Birkb@ Co e, uniuerat 03On&m, London,5.c. 4.
