OX THE COMPOSITION OF ACID BORIC ACID-DIOL COMPOUNDS BY J. BOESEKEN AXD N. VERMAAG
Introduction For some time the Boric Acid Problem has been studied according to different physical methods in the Laboratory of Organic Chemistry a t Delft. The article by Wilder D. Bancroft and Herbert L. Davis,’ whose conclusions depart entirely, not only from the results obtained a t Delft during the last few months, but also from all that had already been attained before by one of us and his pupils and by several other investigators, induces us to publish some of our results now. Before proceeding to the discussion of the cause of the variation of acidity, we may be allowed to draw attention to a few passages of this paper which are, in our opinion, incorrect, or which, through their incompleteness, leave room for an erroneous conclusion. On the significance of the preparation of optically active BRz-anion (R = diol radical) by which the proof is furnished not only of the em’stence of these complexes, but also of their configuration, no more is said than the sentence: “One of the most recent developments in this field is the isolation of substances which appear to be compounds in which boron must be assumed to have a coordinated valence of four and to possess optical activity,” while on p. 2482 it is stated that “the isolation of these compounds seems to be attended with some difficulties,” whereas in reality a number of these compounds has been isolated without any difficulty, if in the polyols the hydroxyl groups are favourably situated, and the salts are not too readily soluble in water or in salt solutions. When we consider that these salts are obtained from aqueous solutions it is natural to assume that the anions BRZ are also present in aqueous solutions and to ascribe the increase of the conductivity of boric acid by the polyols to the formation of complex acids. Referring to what will follow presently, where we have also furnished the proof of the existence of these anions in aqueous solution we may point out that these complex acids will be very strong acids, because they are formed by absorption of the electron of the H-atom, through which in the completely formed compound the hydrogen will be present entirely as proton; in aqueous solutions these acids will, however, be hydrolysed for an important part; hence it is to be expected that the free acids HBR, could be isolated from their aqueous solutions only if they should be much less readily soluble than boric acid or the other component, and there is not much chance of this, exactly with these strong acids. We meet here with analogous circumstances as with the ammonium-bases, which must cede an electron to the hydroxyl group and which are in consequJ. Phys. Chem.,
34, 2479 (1930).
J. BOESEXEN A N D N. V E R M A A S
I478
ence very strong bases. These can, indeed be isolated, if no H-atom is bound to the N; if this is the case they are dissociated for the greater part in aqueous solution, making the impression of being weak bases. It is, acordingly, not astonishing that Bancroft and Davis have not succeeded in isolating complex acids, and notably not with glycerine, where the acid is almost completely hydrolysed (p. 2486). As here the formation of strong acids is in question, the reactions, just as H20 S ammonium base is established, will when the equilibrium amine bear the character of ion reactions, hence take place practically instantaneously. We may, therefore, by no means compare the establishment of the equilibrium: z diol boric acid S complex HzO with the formation of esters of alcohols with weak organic acids (p. 2468 and conclusion 6 pp.
+
+
+
2505-2j06).
I n consequence of the fact that these complexes and notably glycerine boric acid, remain almost entirely hydrolysed, the heat effect will be slight (2486). Glycerine happens to be a very unfavourable object to prove the formation of compounds by another way than by the increase of the conductivity. Meulenhoff* has succeeded in preparing free dipyrocatecholboric acid by heating its dry aniline salt in vacuo; in this it is noteworthy that the acid itself is very little volatile, as may be expected from such a molecule with polary bonds; it carbonizes at zooo. We have succeeded in splitting up three types of thew compounds into optical antipodes, viz. the anions:
by precipitating them from aqueous solutions with optically active bases. The spirane configuration of these compounds and the presence of these complexes also in aqueous solutions can, therefore, hardly be called in question. As we have exactly proved the existence of the anions BR-2 (as well as of BR-) in aqueous solutions, as will be set forth presently we will pass by Bancroft and Davis's other considerations in silence. We will only point out that the tartaric acids form a complicated case, because compounds of boric acid are possible both with the two a-hydroxy-
5
Doctor thesis, Delft, pp. 35-36 (1924). Meulenhoff: 2. anorg. allgem. Chem , 142, 377-379 (1925). Boeseken, Muller and J a p Hong Jouw: Rec. Trav. chim., 45, 919 (1925). Boeseken and Mijs: Rec. Trav. chim., 44,750 (1925).
COMPOSITION O F ACID B O R I C ACID-DIOL COMPOUNDS
I479
carbonic parts of the molecule, and with the middle glycol part. This is one of the reasons why the esters and the amides of the tartaric acid have been examined, where we have only to do with the glycol part of the acids. Here, too, it was found that the active compounds had a greater influence on the conductivity than the inactive ones, which can be explained by the stronger mutual repulsion of the COOR (CONH,) groups than that of the OH-groups, owing to which these latter will lie closer together in the active acids than in the inactive acids. We may add to this that in the imides, where in consequence of the formation of the 5-atomic ring, the position of the OH-groups is fixed, the very reverse is to be expected, and has also been found by J. coop^.^
OH
011
\
\
OH Anti-tartaric-imide
act-tartaric-imide
The anti-tartaric acid imide causes a considerable increase of the conductivity of the boric acid, the act-tartaric acid imide absolutely none. We do not see very well, how this difference can be accounted for in another way than by the formation or non-formation of compounds with boric acid. If Bancroft and Davis should rest satisfied with “the fact that the boric acid dissolves preferentially in the activating substance and is therein more highly dissociated than in water” this purely physical formulation does not seem satisfactory to us, and the assumption of a compound with the favourably situated hydroxyl groups of strongly acid character in the first case, the formation of which is impossible in the second case, seems very much more rational. As regards the tartaric acids themselves, the separation of the potassium boro tartrate KBT2 with two tartrate groups by Lowry7 is so important, because it shows so convincingly how easily, with a favourable position of the hydroxyl groups, a compound of the spirane type is formed; accordingly we readily leave it to the readers to decide whether they will accept this formula given by Lowry, or that given by Bancroft and Davis on p. z j o z . I n conclusion we will point out, that the summary given on pp. 2495-2496 of investigations of one of us and his pupils, is very incomplete for which reason we may lay aside their opinion on the small value of our investigations with a view to configuration determinations. Let us now proceed to the closer examination of aqueous solutions of boric acid and polyols. Bancroft and Davis hold the view that “the activation of boric acid in the presence of sufficient glycerine or mannite or certain other substances, is not Doctor thesis, Delft, pp. 62-67 (1924).
’ J. Chem. Soc., 1929 11, 2853.
1480
J. BOESEKEN AND N. VERMAAS
due to compound formation, but to the fact that the boric acid dissolves preferentially in the activating substance, and is therein more highly dissociated than in water”. Among other things the writers also make the following remarks: “The phenomena studied by Boeseken do not require such a ring structure as he adopted, but may more probably be due to a kind of complex formation involving only one hydroxyl group, the adjacent carboxyl, hydroxyl, or keto groups being required to activate the one reacting with the boric acid. As evidence against the hypothesis of compound formation in these systems is the fact that there is no indication of the existence of a solid compound in the phase study of their aqueous solutions. This is true of the mannite system as well as of the tartaric acid systems, although mannite is one of the most effective activators of boric acid, and should give a compound if one exists.” It will now first be briefly discussed to what the acidity of boric acid diol solutions is to be attributed, after which in the conclusion Bancroft and Davis’s views will be tested by the obtained results.
General Remarks on Boric Acid Diol Complexes When a “favourable” diol (a-oxy-acids etc. included) is added to a solution of boric acid in water, the ring compound
R /‘>OH
is primarily
\o
formed, which, on account of the tendency of boric acid to assume the penta-
[
valent condition, can pass into resp. the monobasic acids H R B(OW2]
with water or with a new molecule of diol. These acids are comparatively strong, and as regards structure perfectly analogous to the acid HIB(OH)a] to which at present the acidity of a boric acid solution is a s ~ r i b e d . * ~ ~ With less favourable diols the position in space of the OH-groups will render the transition of boron to the tetrahedric valence arrangement almost impossible; hence the complex formation will go no further than the
ROH,
and no increase of acidity will be observed
tane -2-4 diol boric acid.)* 8Hermans: 2.anorg. allgem. Chem., 142, 83 (1925). 9 Boeseken and COOPS: Rec. Trav. chim., 45, 407 (1925).
(2-4 dimethyl
pen-
COMPOSITION OF ACID BORIC ACID-DIOL COMPOUNDS
1481
K i t h absolutely unfavourable diols no complex formation will take place at all. Combining all the undissociated complex compounds under HBR and HBR,, the course of the reactions may be represented as follows: H B (boric acid) HBR+ HB
+ R (diol) s HBR
H+
+ BR-
R a o r
+ 2R e HBRz
H+
+ BR-2.
On the assumption that H B R and HBR, are strong acids, HBR and HBRz may be neglected with respect to their ions in the case of favourable diols in diluted solution. The condition of the solution is then determined by the following equilibrium constants:
These constants are analogous to those of carbonic acid and ammonia. If for less favourable diols H B R and HBRz may possibly not be neglected, w r y certainly these equations cannot be used for more concentrated solutions, because then on one side tetra-boric acidlo on the other side the undissociated HRR and HRRz may begin to play a part also for the favourable diols. By division of the equilibrium constants we get: I)
..................
If now K ~ / K Iis = i 1 0 3 and the diol concentration is f 0 . 1 mol., then (BR-l) may be neglected with regard to (BR-,) and the acidity of such a boric acid diol solution is governed by the di-complex.
Poly-alcohols For poly-alcohols this has been proved in the Rec., 49, 711, to which we refer here, as also to the theoretical derivations. Inversely it has appeared from this, that actually Kz/K1 = + 1 0 3 or > 103. To the measurements mentioned in Rec., 49, 7 I I only some results are added calculated from measurements of others: I. Rimbsch and Ley1' measured the hydrogen-ion concentration of boric acid with different quantities of mannitol. From these data the following values were calculated by the methods described in Rec., 49, 7 1 2 (1930). Boric acid = 0.1mol. lo l1
Kolthoff: Rec. Trav. chim., 45, 501 (1925). Z. physik. Chem., 100, 393.
J. BOESEKEN AND N . VERMAAS
1482
n
PH
Boric acid: polyol
r : 6 1 : 1
3.86. 3'08\ 3.86.
1 : s
33 .. 11 11
1 : 4
3.27)
2
I
.o
2.1
0.59
2
0.46
1.9
.o
1 : 1
3.40
1:3
Boric acid =
0.2 j
mol.
Boric acid: polyol 1 : 2
PH 2
'
93)
1 : 6 I : I
3.18
There is, therefore, no doubt but n is = 2 for these solutions but on the other hand it is easy to understand that for more concentrated solutions, where, moreover tetraboric acid appears, also the remaining OH-groups of the mannitol play a part through which the difficulties that arose in the separation of these acids, find a natural explanation.12 11. Years ago van Liempt (Rec., 39, 3 58) titrated boric acid diol solutions potentiometrically with alkali. One of his measurements (on fructose) can be used for the calculation mentioned in Rec., 49, 7 I I on account of the proportion boric acid : diol = I : j. I n view of the irregular intervals of the pH his measurements appeared to be not very accurate, among other things for the reason that he made use of a very concentrated boric acid and fructose solution, which he greatly diluted with sodium hydroxide during the titration. According to Rec., 4 9 , j 14.
[H+l X [BD%-l Ks = [HB] X [D]' hence if the initial concentration of boric acid = a, and of the fructose = sa, and if xa mol of alkalies have been added, then:
The final point of the titration being reached a t 1 5 cc, x is = y/rg if y is = the number of added cc's of NaOH, hence: '*Hermans: Z. snorg. allgem. Chem., 142, 1 x 1 (1925);Fox and Gauge: J. Chem. Soc., 99, 1075
(1911).
COMPOSITION O F ACID BORIC ACID-DIOL COMPOUNDS
[H+l
(IS
- y)
a2KI 3' -= constant = (75 - 2y)Z 225
C.
For the middle part of van Liempt's titration curve, where the error in ( I j - y) and in y is slight, and the electrolyte content not too great, and where also the dilution does not vary too much, the following values were calculated: cc's NaOH
PH
5 6 7 8 9
3.21 3.42 3.56 3.69 3.83 4.09
c
x
I08
7.3 6.3 6.5 6.7 6.8
5.5 Accordingly the K 3 appeared to be constant even under these comparatively nnfavourable circumstances. a-oxy-acids IO
I. I n the first place a small orientating calculation will be given in connection with Rimbach and Ley's measurements" on a-oxyacids. They measured the hydrogen-ion concentrations in oxyacid solutions of different concentrations with and without, boric-acid. If it is assumed that, just as it has appeared with the poly-alcohols, the mono-complex acid may be neglected with respect to the bi-complex acid in these solutions, the following calculation may be made: Let a be the (H+)-conc. of n X 0.1 mol. of oxy-acid solution without boric acid, and b the (H+)-conc. of the same solution, but with 0.1 mol. of boric acid, then, if the variation in the conc. of the undissociated acids is neglected, the following equations hold:
a2 -b X P n X 0.1 n X 0.1
p and q being resp. the (Z-) and the (BZ2-) concentrations of the boric oxyacid mixture. Hence : n2 x I O - ~ K ~ n X 0.1 XI , and n x 0.1K1 = a* P = b ' q = b
Korv p
n2 X 1o-3K2 + q = b = a2 - + b b
b2 - a? Hence
-~
n2 Lactic acid.
=
c must be constant for different values of
D
J. ROESEKEN
I484
AND
N. VERMAAS
H+ conc.
H+ conc. without boric acid
Conc. oxy. acid
with o. I mol of boric acid
6 . 0 ; X IO-^
0.1
0.2
8.36 X
9.06
IO+
c
x
10'
0.33
17.1
0.j2
11.5
21.7
0.37
'3.5
27.6
0.36
0.5
14.6
0.6
15.8
35.1 38.0
0.33
0.7
17.1
41.2
0.29
0.3 0.4
0.40
b2 - a2 For the mono- and the tri-complex acid, the values of resp. and n b2 - a2 would have to be constant, which is, of course not the case in the above n3
observations, where n varies from 1-7, A similar result, though not so beautiful, is also yielded by the glycollic acid measured by Rimbach and Ley. 11. For the determination of the dissociation constants of the acid complexes of boric acid and a-oxy acids, the pH-curve was determined for the titration of I O cc 0.1 mol. of oxy acid with 1.0cc 1.0mol. KBO2-solution. Assuming the 1-2 complex acid to constitute the last stage of the tendency to complexity of boric acid, we have, therefore, here four acids side by side, with the following dissociation constants:
KS
=
[H+l X [Z-I [HZI
Kq =
E-I w31
[H+l X
Further, when b mol. of KROZ are added:
+ + + + + + +
initial conc. of oxy acid a = ~[BZZ-] [BZ-I [Z-I WI. all the negative ions b [H+] = [BZz-] [BZ-I EZ-1 4- [B-I. [BZ-I [HB] [B-I. all the boric acid b = [BZz-]
+
+
The whole titration curve lying in an acid region, we may neglect the (B-) conc. with respect to (HB), so that remains: Kz X [HZ]' X [HB] = [H+l X [BZz-l Ki X [HZ] X [HB] = [H+] X [BZ-I Ka X [HZI = [H+l X [Z-I a = ~ [ B Z Z - ] [BZ-I [Z-I [HZI b [H+] = [BZz-I [BZ-I [Z-I b = [BZ*-] [BZ-I [HB]
+
+
+
+ + +
+ +
COMPOSITION O F ACID B O R I C ACID-DIOL COMPOCNDS
1485
For a pretty well equal electrolyte concentration, K3 may be measured by means of a simple potentiometric titration of the oxy-acid with alkali, so for a definite point of the titration curve K,, a, b, (HT) are known; unknown are: KP,XI, (BZz-), (BZ-), (Z-), (HZ), and (HB). Hence there are 6 equations, and 7 unknowns. By the application of a neglection, K, can be calculated for a portion of the titration curve, in consequence of which 6 equations and 6 unknowns remain for the general case. After elimination of the unknown quantities, two equations remain in the end from which the (HZ) cannot be eliminated any more 2)
.
+
KzK3[HZI3- KB[H+]*[HZ]* [H+]*[HZ!- [H+]*{ a - b - [H+]) = o
The value of (HZ) to be approximated from this is filled in, in the following equation, from which K1can be calculated [H+] [IILI
3 ) . . . .Ki = 7X
[H+1 { 2b - a
+
+
2[H+l [HZI) - K3[HZl K3[HZ] - [H+I2
K e now proceed to the computation of K B : By dirision of K B by K,, follows, compare: I )
If KB/K1has a value of i103or > 103,as has appeared to be true for the polyols, and as afterwards will appear to be true here too, (RZ-) may be neglected with respect to (BZ2-) over the portion of the titration curve where (HZ) > 0.02 mol. as the quantity of (BZ-) is then less than jqc of the quantity of (BZB-). These values may be neglected over the part of the titration curye from 0.10-0.40 cc KBO,. Hence we get for the calculation of Kz: K B X [HB] X [HZ]' = [H+l X [BZz-I K3 X [HZI = [H+l X [Z-I a = [HZ] [Z-] 2[BZZ-] (initial conc. of the oxy-acid)
+
+
+
b = [HB] [BZZ-] (total boric acid) [H+] = [Z-] [BZn-] (all negative ions) b
+
+
From these 5 equations (BZp-) (Z-), (HB), (HZ) may be eliminated after which we get:
1486
J. BOESEKEN 4 T D N . TERMBAS
Recapitulating: with formula 4) K1 can be calculated over the first part of the titration curve (0.10-0.40 cc KBOQ),with formulae 2 ) and 3) K, can be calculated over the remaining part, where the appearance of RZ- ions is very perceptible. The pH-curves are measured for a-oxy-iso-butyric acid, cyclo-pen tane-a-oxy-carbonic acid, cyclo-hexane-a-oxy-carbonic acid. Determination of K S at an electrolyte concentration varying between 0 . 0 2 j - 0 . 0 7 5 mol. by potentiometric titration of I O cc 0.1 mol. of oxy-acid with 1.0 cc 1.0 X. potassium hydroxide. cc's
a-oxy-iso-butyric acid
cyclo-pentane-a oxy-carbonic-acid
cyclo-hexane-a-oxy-carbonic-acid
KOH
K?
PH
x
10-4
x
10-4
0.2j4
3.387
I .4
0.52j
3.894
1.4
0.790
4.438
1.4
0,267
3 . j62
1.0
0.532
I
.o
0.792
4.066 4.581
I
.o
0.226
3,514
0 . 9 0 X 1o-I
0.537 0.690
4.104 4.383
0.91 0.92
ICow follow the different points of the determined pH curves for the titration of oxy-acid with KBOa with the values calculated for Kz and KI. A typical phenomenon, which has not yet been accounted for, was that the adjustments over the second parts of the titration curves required a good deal of time. a-oxy-iso-butyric acid CC'S
PH
K?
2.664
KBO?
0 . I94
2.7j8
2.3 2.3
0.240
2.869
2 . 2
0.288
2.2
0,336 0.384
2.99' 3.119 3.282
2.4
0 . '47
0.532
3.934
0.582
4.225
0.630
4.502
0.731
5.038
0.783 0.836
5
'
290
5'523
KI
2.4
3.I 3.1 3.3 3.2 3.3 3.6
x
:o-~
COMPOSITION O F ACID B O R I C ACID-DIOL COMPOUXDS
1487
cyclo-pentane-a-oxy carbonic acid PH
K2
0 . IO0
2.641
3.2
CC'S
KBOi
0.144
2.724
2 . ;
0.192
2.812
2.8
0.239 0.285
2.901
3.032
3.3 3.2
0.331 0.337 0.616 0.663
3'154
3.'
3.3'2 4.644
3.2
0.712
5 ' '99
0.762 0.808 0.854
K,
3 , 8 X IO-^
4.943
3.6 3.4 3.9 3.7
5.434 5.659 5.842
12
cyclohexane-a-oxpcarbonic acid PH
I
