T H E BROMINE-BROMIDE-TRIBROMIDE EQUILIBRIUM.

B Y R O B E R T OWEN GRIFFITH, ANDREW MCKEOWX, AND ALBERT GORDON WINN.

Received I 5th December, I 93 I .

In connection with the study of the thermal reaction between bromine and oxalates (see the following paper), a rather more detailed knowledge of the equilibrium constant (K3) of the equilibrium Br, + Br- + Bra- than that hitherto available was desired. Previous work on this equilibrium may briefly be summarised. I t was first studied by Jakowkin,' using the well-known partition method with carbon tetra- chloride as the second phase. Later, Worley investigated the equili- brium by allowing two aqueous solutions, the one containing bromine, the other bromine and potassium bromide, to come to equilibrium via the vapour phase, and then analysing the two solutions. Again, Jones and Hartmann 3 combining Jakowkin's method with measurements of electrical conductivity, carried out an elaborate study of the tribromide equilibrium at 0'. In their calculations of the equilibrium constant they made allowance for (a) the hydrolysis of bromine and (b) the formation of pentabromide (Br5-). In each of these investigations most of the data refer to solutions of high bromine content, in which, owing to allowance having to be made for pentabromide formation and also to possible deviations from the laws of dilute solution, the interpretation of results becomes needlessly complex and to some extent uncertain.

Lewis and Storch showed that the partial pressures of bromine above solutions of bromine in carbon tetrachloride are proportional to con- centration up to about o.zgM. (This corresponds to a concentration of free bromine in an aqueous layer in equilibrium with the carbon tetra- chloride layer of about o-oIM.) Above this concentration deviations from Henry's Law are to be expected. They also determined the value of K3 in presence of HBr a t 25' by the Jakowkin method. In measuring the partition coefficient of bromine between water and carbon tetra- chloride they made the aqueous phase N/IW with respect to H,SO, in order to cut down hydrolysis of the bromine. Lewis and Randall5 have recalculated some of Jakowkin's results obtained with KBr and weak bromine solutions, and Linhart has attempted to correct Worley's data (relating to solutions strong in bromine) by making allowance for pentabromide formation. Finally, Sherrill and h a r d 7 have carried out a series of measurements of K3 in the presence of HBr. They found that the values of this equilibrium constant (defined as K, = [Bra] [B!-]/[Br3.-]) decreases with increasing concentration of bromide, a variation which

1 Jakowkin, 2. physihd. Clem., 18,583, 1895 ; m, 19, 1896. 2 Worley, J.C.S., 87, 1107, 1905. 9 Jones and Hartmann, Trans. Amev. Electrochem. SOC., 30, 295, 1916.

Lewis and Storch, J . Anaev. Chem. Soc., 39, 2544, 1917. Lewis and Randall, J . Amer. Chem. SOC., 38, 2348, 1916.

* Linhart, J . Amer. C l e m . Soc., 40, 158, 1918. 7 Shemll and Izard, J . Amer. Chem. SOC., 50, 1665, 1928.

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I 02 THE BROMINE-BROMIDE-TRIBROMIDE EQUILIBRIUM

they were the first to note. investigations are summarised.

In Table I. the results of these previous

TABLE I.-VALUES OF THE EQUILIBRIUM CONSTANT K,.

Author.

Jones and Hartmann . Lewis and Storch . . Jakowkin (calculated by Lewis and

Worley (calculated by Linhak) : Worley Shemll and izard' 1

Randall) .

Bromide.

KBr HBr

KBr KBr KBr HBr

Temperature.

O0 2 9

25"

32.5" 2 5 O

26-5'

K3.

0.05 I 0.06 I

0.062 0.063 0.065

0.057-0.060

We have carried out fresh determinations of the equilibrium a t the two temperatures 16.5" and 21-5', using HBr, NaRr, KRr and LiBr, varying the concentration of electrolyte between 0.03 and 0.75 N and keeping the concentration of free bromine below O*OIM.

Experimental.

The method of Jakowkin was adopted, with carbon tetrachloride as the second solvent. Two preparations of bromine were used, the one from Kahlbaurn, the other sample purified by the method of A. Scott.s Also two preparations of CC1, were employed. The first of these was an A.R. grade sample which had been specially purified for photochemical work by refluxing with pure chlorine in an all-glass apparatus, washing with Na,CO, solution, drying over Na,CO, and then over Na, and finally fractionating. The second preparation, of the same origin, was subjected to similar treatment, with the difference that the refluxing was carried out with bromine instead of with chlorine and that the final drying with Na was omitted. The bromides were Kahlbaum or Merck preparations and were found to be free from impurities which react with bromine.

In all cases at least forty-eight hours was allowed for attainment of equilibrium between the aqueous phase and the carbon tetrachloride phase. The bromine content of each layer was determined by pipetting off samples, running into excess of sodium arsenite solution, and hack titrating with iodine.

In determining the distribution coefficient of bromine between CCI, and water in the absence of added bromides, the method of Lewis and Storch for limiting the hydrolysis of bromine in the aqueous phase was used, ie., the aqueous phase was made h7/1000 with respect to H,SO,. Table 11. gives the results of distribution measurements using this pro- cedure. The effect of hydrolysis of bromine in the aqueous layer is shown in the results of Table III., relating to experiments with much weaker bromine solutions in absence of added acid. In these and the following tables, concentrations are expressed in moles per litre, and D represents

the ratio concentration of Rr, in CCl, concentration of Br, in H,O'

Scott, J.C.S., 103, 847, 1913.

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R. 0. GRIFFITH, A. McKEOWN, AND A. G. WINN 103

m J C C * , - . iBr&,o * . D . .

TABLE 11.

t = 21.5~ t = 16.5" 7 7 /-----J+.--.- 7 -________-

0,1880 0.1881 0.1775 0.370 0.1843 0-1838

0.00682 0.00682 0.00646 0.01328 0-00700 0~00700

27.6 27.6 27.5 27.8 26-3 26-3

CBr21cc,, *

[Br21Hz0 . D .

t = 16.5'

0'1272 0.0748 0.03755 0.02360 0.01653

0.004950 0.002939 0.001490 0-000985 0~000698

25.8 25'5 25.2 24.0 23'7

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104 THE BROMINE-BKOMIDE-TRIBROMIDE EQUILIBRIUM

1Wio 7h-

0.109

that the ratio (log,, y) /p , where p is the ionic strength of the solution, is approximately constant for a given non-electrolyte and a given salt.

Clearly, what is required in the present instance is the effect of bromides on the activity of bromine, but, on account of the formation

Ratios.

-

0.2 0.4 FIG. I.

0' 6 P 4

of tribromide ion this cannot be determined directly. Yet, from Randall and Failey's data an approximately correct estimate of this effect may be obtained. From these authors' tables we extract the figures given in Table IV.

TABLE IV.-ACTIVITY COEFFICIENTS OF NON-ELECTROLYTES IN SALT SOLUTIONS.

(AFTER RANDALL AND FAILEY.)

Non-electrolyte. salt.

Na,SO, Na,SO, NaBr KBr

Na,SO, NaBr KBr

We have supplemented these data by determining the effect of Na2S04 on the activity of Bra, by finding the partition coefficient a t 21.5' between CCl, and an aqueous solution of Na,SO, of concentration M / 6 \p = 0. j). In this case we find D = 31.3 ; hence Y B ~ ~ in M / 6 Na,S04 = 31*3/27*5 =

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R. 0. GRIFFITH, A. McKEOWN, AND A. G. WINN 105

1.021 1.008

1.134 and (log yBrJ /p = 0.109. given in Table IV., we infer that :

Combining this figure with the ratios

For Br, in NaRr, (log y)/p = @.I09/1*24 = 0.088. For Br2 in KBr, (log y)/p = 0*1og/1.92 = 0.051.*

These estimates are probably reasonably accurate; i t is not so easy, however, to evaluate the effects of HBr and especially of LiBr on the

0.4527 0.1767

TABLE V.-VALUES OF THE EQUILIBRIUM CONSTANT K,.

1.070 1.035 1-014 1.014 I -007

0.4900 0'2409 0.09346 0*08050 0.04508

0.4903 0-1912 0.04633

0.009678 0~008810 0.003665

CBr - = I*.

0.7500 0.5000 0.5000 0.2500

0.1017 0.100g 0'1000 O'IOOO 0.06727 0.03683

0' 1000 0'1000

0.4627

0.04627 0.1851

0.7500 0.5000

0-0500 0'2000

CBr2.

0.05 150 0.0391 I 0.0 I 074 0-01059 0.04618 0.03123 0.01497 0.01266 0-oogg I g 0.005356 0.008375 0.003467

0-01 127 0'01 109 0-009961

0-05020 0.01047 0.01092 0.00760

KBr Solutions at 21-5'.

0.003782 0.004283 0'001 129 0~002060 0'02022 0.01289 0.00575 I

0.003 7 78 0-002004 0.003989 0.002 I 35

0'004792

1.105 1.070 1.070 1.035 1.014 1.014 1'014 re014 1.014 1.014 1.009 1.005

0.7023 0-4651 0.4904 0-2415 0'07404 0.08166 0.09248 0.09303 0.09386 0.09665 0.06288 0.03550

0.04772

0.00961 I 0.008530 0.02596 0.01834 0.0092 I 9 0.00 7 8 68 0.006141 0.003352 0.004386 0-001332

0.03483

HBr Solutions at 21.5'.

0-0 I 004 0~008401 0'004233

15~. 0.04630 0.009328 0.008369 0.003409

0.003899 0.001 142

0~004191 0.00255 I

0.0593 0.060 I

0.0573 0.0584

1.160 1.107 1.043 1-01 I

LiBr Solutions at 21.5'.

0-5000

0.0500 0'2000

0.01078

0.0075 9 7 0'01120

0.009574 0.008550 0.003409

0~001206 0.002650 0.004188

0.061 8 0.0593 0.0572

KRr Solutions at 16-5".

O.OOgg52 0.009078 0.006540 O.OIg50 0*004920

0.001 I 18 0'002 I02 0.003850 0.01327 0.005880

0.01 107 0.01 I 18 0-0103g 0.03277 0~01080

0.0550 0.0557 0.0550 0.0548* 0.0539

0-5000 0.2500

0-0500

0'1000 0'1000

NaBr Solutions at I 5.5'.

o-oroSI O*OII39 0.007980

1.107 1.043 1-01 I

o~oo1132 0.002 5 80 0.0043 I5

0.5000

0.0500 0'2000

0.0574 0.0560 0,0548

*Justification of this procedure follows from the derivation of Debye and McAuley (Physzkal. Z., 26, 22, 1925) of the effect of electrolytes on the activity coefficients of non-electrolytes. See also S h e d and Izard ( J . Amer. Chews. SOC., 53, 1667, 1g31), who have applied a correction analogous to that attempted here to their data on the equilibrium Q, + C1- =+ Cl,-.

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106 THE BROMINE-BROMIDE-TIBROMIDE EQUILIBRIUM

activity of Br,. From Randall and Failey's data, i t is found that (log y) /p for non-electrolytes in HC1 is in general about one-fifth that in NaC1. Assuming the same ratio for HBr and NaRr, we obtain (log y ) / p = 0.018 for the effect of HBr on 7 ~ ~ ~ . For Br, in LiRr, the admittedly uncertain value (log y ) / p = 0.042 was chosen.

In Table V. are summarised our determinations of K, in presence of HRr, NaRr, KBr and LiBr a t 2 1 . 5 ~ and 16.5". The values of y~~~ used in each case are given. The values of D employed were 27.5 for the experiments a t 21.5" and 26.3 a t 16*5", except in the three asterisked cases, relating to experiments with a somewhat high content of free bromine. For these the value of D used was 27-8 a t 21.5" and 26.6 a t 16.5". The values of K , given in the above Table are shown plotted against the ionic strength (p) in Fig. I .

Discussion of Results.

The following points call for comment :- (a) From the results with 0.1 N KBr solutions a t 2 1 * 5 O , i t is seen that

a good constant K, is obtained with EBr2 varying nine-fold, no correction for pentabromide formation having been applied. It would appear that for solutions of low bromine content such as here employed pentabromide formation is not appreciable ; further, i t is likely that previous estimates of its extent (in stronger Br, solutions) are too high,

In solutions of NaBr, KBr, and (probably) LiBr the constant at first increases, passes through a maximum in the neighbourhood 0-4-0.6 N , and then decreases. The salt effect is greatest in the case of LiBr, least for KBr.

This is in agreement with the results of Sherill and Izard, though the magnitude of the fall we ob- tain is somewhat less than theirs.

(d) The values of K3 here presented are slightly less than those of previous workers. This is due to the correction made here for the activity coefficient of bromine.

(e ) The rise of K, with increasing p in the cases of the alkali halides is somewhat unexpected. Thus with NaBr at 2 1 * 5 O , K3 rises from 0.0568 a t p = o to 0.0601 at p = 0.5, a rise of about 508 per cent. Since, however, over the same range YBrp increases by about 10.7 per cent.,

(b) K3 varies with ionic strength.

(c) With HBr, K , falls with increasing p.

cent. greater at p = 0.5 than a t p = 0. Since by definition aBrp ' aBr-

is constant, this means that the activity coefficient of the tribromide ion is approximately 16 per cent. greater than that of the bromide ion when the cation is Na+ and the ionic strength 0.5. That this result is not occasioned by an incorrect assumption regarding the activity of bromine in solutions of NaBr niay be demonstrated as follows. The constant K3 was determined a t 21.5" for a solution (A) which was o.05M with respect to NaBr and 0.15M with respect to Na,SO,, ie., with a total ionic strength p = 0.5 ; the value obtained was K3 = 0 . ~ 5 7 5 . In the calculation the value Y B ~ ~ = 1-134 was that obtained (see above) in M / 6 Na$O, solution €or which also p = 0.5. Combining these values of K3 and yBq, it follows that in solution (A) the ratio Y B ~ ~ - / Y B ~ - is ap- proximately 1 5 per cent. greater than iii a solution of zero ionic strength, in good agreement with the result for O*SM NaBr solution. This con-

aBr3-

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R. 0. GRIFFITH, A. McKEOWN, AND A. G. WINN 107

elusion is not in agreement with that reached by Lewis and Randall," who postulated that the activity coefficient of tribromide ion is less than that of the bromide ion.

We wish to thank the Department of Scientific and Industrial Research for a grant to one of us (A. G. W.) while this work was in progress. We are also indebted to Imperial Chemical Industries, Ltd., for a grant defraying part of the cost of this investigation.

ildusprait Laboratory of Physical and Eleclvochemistry, University of Liverpool.

11 Lewis and Randall, Thermodynamics, p. 520.

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