PHYSICAL CHEMISTRY

K. N. MOCHALOV, V. S. KHAIN, G. G. GILMANSHIN

GENERALIZED SCHEME FOR THE HYDROLYSIS OF THE BOROHYDRIDE ION AND DIBORANE

(Presented by Academician V. N. Kondrat’ev, 11 XI 1964)

The most important reaction of the borohydride ion—its interaction with water (hydrolysis)—is usually represented by the gross equation
$\mathrm{BH_4^- + 2H_2O = BO_2^- + 4H_2}$, which reflects only the stoichiometric relations and the final result of the process, but gives no idea of the mechanism and intermediate stages of this complex reaction. Above all, it does not take into account the participation of hydrogen ions in the process, whose role, as experiment shows, is very large.

Studying the hydrolysis of $\mathrm{LiBH_4}$, $\mathrm{NaBH_4}$, $\mathrm{KBH_4}$, and $\mathrm{CsBH_4}$ in borate buffer solutions at various pH values, we found that the reaction rate is proportional to the concentrations of borohydride and hydrogen ions and is practically independent of the nature of the cation ($^1$). The plots we obtained for the dependence of the first-order constant with respect to $\mathrm{BH_4^-}$ on the concentration of $\mathrm{H^+}$ pass through the origin, which means the absence of hydrolysis at $[\mathrm{H^+}] = 0$.

In studying the decomposition of $\mathrm{LiBH_4}$ in a strongly acidic medium ($^2$), it was established that diborane, which is a dimer of borine, is evolved in the first stage. Therefore the latter must necessarily arise in the course of hydrolysis, which was also proved experimentally ($^3$).

Thus, the hydrolysis of the borohydride ion proves to be closely connected with the hydrolysis of diborane (borine) ($^{4–6}$).

The currently accepted scheme for diborane hydrolysis, due to Weiss and Shapiro ($^{7–9}$), includes stages involving the formation of neutral intermediate products $\mathrm{BH_2OH}$ and $\mathrm{BH(OH)_2}$ and refers to a reaction in the gas phase. It is not applicable to hydrolysis conditions in an aqueous alkaline medium, since in this case the ions $\mathrm{BH_3OH^-}$ and $\mathrm{BH_2(OH)_2^-}$ arise, which, according to V. I. Mikheeva and V. Yu. Surs ($^{10}$), may be regarded as products of individual stages of borohydride hydrolysis.

Subjecting a hydrolyzing $\mathrm{NaBH_4}$ solution to chromatographic separation on paper under strongly alkaline conditions, we succeeded in confirming the formation in it of the indicated hydroxy derivatives ($^{11}$). On the basis of the available data one can construct a unified, more general scheme encompassing the hydrolysis of the borohydride ion and diborane in neutral, acidic, and alkaline media.

The generalized scheme for the hydrolysis of the borohydride ion and diborane is:

\[ \mathrm{BH_4^- + H_3O^+ \rightleftarrows \left[ \begin{array}{c} \mathrm{H^+ \cdot BH_4^-} \\ \mathrm{H_2O} \end{array} \right]^* \to BH_3 + H_2O + H_2.} \tag{1} \]

\[ \begin{aligned} \text{a)}\quad &\mathrm{BH_3 + H_2O \rightleftarrows [BH_3 \cdot H_2O]^* \to BH_3OH^- + H^+;}\\ \text{b)}\quad &\mathrm{BH_3 + OH^- \rightleftarrows [BH_3 \cdot OH^-]^* \to BH_3OH^-;}\\ \text{c)}\quad &\mathrm{2BH_3 \rightleftarrows B_2H_6.} \end{aligned} \tag{2} \]

\[ \mathrm{BH_3OH^- + H_3O^+ \rightleftarrows \left[ \begin{array}{c} \mathrm{H^+ \cdot BH_3OH^-}\\ \mathrm{H_2O} \end{array} \right]^* \to BH_2OH + H_2O + H_2.} \tag{3} \]

\[ \begin{aligned} \text{a) }& \mathrm{BH_2OH + H_2O \rightleftarrows [BH_2OH \cdot H_2O]^* \to BH_2(OH)_2^- + H^+;}\\ \text{b) }& \mathrm{BH_2OH + OH^- \rightleftarrows [BH_2OH \cdot OH^-]^* \to BH_2(OH)_2^-.} \end{aligned} \tag{4} \]

\[ \mathrm{BH_2(OH)_2^- + H_3O^+ \rightleftarrows \left[ \begin{array}{c} \mathrm{H^+ \cdot BH_2(OH)_2^-}\\ \mathrm{H_2O} \end{array} \right]^* \to BH(OH)_2 + H_2O + H_2.} \tag{5} \]

\[ \begin{aligned} \text{a) }& \mathrm{BH(OH)_2 + H_2O \rightleftarrows [BH(OH)_2 \cdot H_2O]^* \to BH(OH)_3^- + H^+;}\\ \text{b) }& \mathrm{BH(OH)_2 + OH^- \rightleftarrows [BH(OH)_2 \cdot OH^-]^* \to BH(OH)_3^-.} \end{aligned} \tag{6} \]

\[ \mathrm{BH(OH)_3^- + H_3O^+ \rightleftarrows \left[ \begin{array}{c} \mathrm{H^+ \cdot BH(OH)_3^-}\\ \mathrm{H_2O} \end{array} \right]^* \to B(OH)_3 + H_2O + H_2.} \tag{7} \]

\[ \begin{aligned} \text{a) }& \mathrm{B(OH)_3 + H_2O \rightleftarrows [B(OH)_3 \cdot H_2O]^* \to B(OH)_4^- + H^+ \to BO_2^- + 2H_2O + H^+;}\\ \text{b) }& \mathrm{B(OH)_3 + OH^- \rightleftarrows [B(OH)_3 \cdot OH^-]^* \to B(OH)_4^- \to BO_2^- + 2H_2O.} \end{aligned} \tag{8} \]

According to \(^{(12)}\), the borate ion in aqueous medium exists in the form \(\mathrm{B(OH)_4^-}\). It is easy to see that, first, summation of the partial equations (1), (2a), (3), (4a), (5), (6a), (7), (8a), and also (1), (2b), (3), (4b), (5), (6b), (7), (8b) leads to the usual gross equation for hydrolysis of the \(\mathrm{BH_4^-}\) ion and, second, pairwise summation of equations (2a) and (3), (4a) and (5), (6a) and (7) gives the scheme of hydrolysis of borane (diborane) proposed by Weiss and Shapiro. Finally, pairwise summation of (1) and (2a), (3) and (4a), (5) and (6a), (7) and (8a) gives the scheme of V. I. Mikheeva and V. Yu. Surs.

The character of the medium in which hydrolysis proceeds affects the course of the process as follows: the borane \(\mathrm{BH_3}\) initially formed (according to equation (1)) in alkaline medium reacts as shown by equation (2b). At a low concentration of hydroxyl ions—in a neutral and not too acidic medium—it may interact according to the equation \(\mathrm{BH_3 + H_2O \to BH_2OH + H_2}\) (summation of (2a) and (3)). Finally, at very low concentrations of hydroxyl ions and water (in a strongly acidic medium), the only possibility left for borane is dimerization, and hydrolysis is limited to reactions (1) and (2b).

To confirm the participation of hydrogen ions not only in the initial stage of the process but also in the subsequent ones (equations (3), (5), (7)), we studied the kinetics of hydrolysis of the ions \(\mathrm{BH_4^-}\), \(\mathrm{BH_3OH^-}\), \(\mathrm{BH_2(OH)_2^-}\), and \(\mathrm{BH(OH)_3^-}\) as a function of the pH of the medium. Mono- and dihydroxy compounds were obtained for this purpose according to \(^{(10)}\) by passing tetraborane into 30% and, respectively, 20% solutions of caustic alkali at a temperature close to \(0^\circ\). Their hydrolysis was carried out in borate buffer solutions with various pH values at 15, 25, and \(35^\circ\). The pH measurements were performed with an accuracy of up to 0.001 on a pH-Meter Radiometer 72 Emdrupvej—Copenhagen Denmark instrument with a G-200B glass electrode. The concentration of hydroxy compounds was determined during hydrolysis by argentometric titration \(^{(13)}\) and was calculated from the content of “active” hydrogen (bound directly to boron). In both cases hydrolysis proved to be a first-order reaction with respect to the initial substance (Table 1).

The dependence of \(k_1\) on the concentration of hydrogen ions is expressed by straight lines passing through the origin. The hydrolysis of the \(\mathrm{BH(OH)_3^-}\) ion was studied by us by the polarographic method \(^{(14,15)}\). Its rate (\(k_2\)) proved to be three orders of magnitude greater than the rate of hydrolysis of the \(\mathrm{BH_4^-}\) ion (Table 1). The activation energy for the \(\mathrm{BH_4^-}\), \(\mathrm{BH_3OH^-}\), \(\mathrm{BH_2(OH)_2^-}\), and \(\mathrm{BH(OH)_3^-}\) ions has, respectively, the following values: 12 337, 10 928, 10 575, 10 575 cal/mole. In order to judge the relative rate of the individual consecutive reactions of this scheme, we studied the kinetics of the transformations \(\mathrm{BH_4^- \to BH_3OH^-}\), \(\mathrm{BH_3OH^- \to BH_2(OH)_2^-}\), and \(\mathrm{BH_2(OH)_2^- \to BH(OH)_3^-}\).

It turned out that potassium ferricyanide, under certain conditions (in a strongly alkaline medium), reacts only with \(\mathrm{BH_3OH^-}\) ions. Since this reaction proceeds 250 times faster than the interaction of \(\mathrm{BH_3OH^-}\) with hydrogen ions, the rate of disappearance of ferricyanide in a borohydride solution can be used to judge the rate of conversion of the latter into \(\mathrm{BH_3OH^-}\). The reaction was carried out in borate buffer solutions at various pH values and temperatures. Changes in ferricyanide concentrations were determined spectrophotometrically from the absorption at \(\lambda = 415\ \mathrm{m\mu}\) on an SF-5 instrument. The reduction of ferricyanide proved to be a first-order reaction with respect to the ions \(\mathrm{Fe(CN)_6^{3-}}\), \(\mathrm{BH_4^-}\), and \(\mathrm{H^+}\).

The data obtained show that the rate of interaction of ferricyanide with a borohydride solution is determined by the rate of hydrolysis of the latter, and this, in turn, coincides with the rate of conversion of \(\mathrm{BH_4^-}\) into the \(\mathrm{BH_3OH^-}\) ion (Table 2).

Consequently, the conversion

\[ \mathrm{BH_4^- \rightarrow BH_3OH^-} \]

(through the intermediate complex \(\mathrm{H^+BH_4^-}\)) is the limiting stage of the consecutive hydrolysis reaction of borohydride.

As for the conversion

\[ \mathrm{BH_2(OH)_2^- \rightarrow BH(OH)_3^-}, \]

it is the first stage of the hydrolysis of \(\mathrm{BH_2(OH)_2^-}\) to borate; the values of the rate constants for this process are given in Table 1. It is seen from it that the second stage of this process,

\[ \mathrm{BH(OH)_3^- \rightarrow B(OH)_4^-}, \]

proceeds about 500 times faster than the first. Therefore, the rate of the conversion of interest to us may be taken as equal to the overall rate of hydrolysis of \(\mathrm{BH_2(OH)_2^-}\).

Table 1

Hydrolysis of \(\mathrm{NaBH_4}\) and its hydroxy derivatives in borate buffers \((\mu \cong 0.4)\)

For the convenience of comparison, the kinetic data obtained by us are presented below.

Along with the “acid” stages, the generalized scheme also contains reactions of addition of hydroxyl ions. The participation of \( \mathrm{OH^-} \) ions in the hydrolysis process begins from the stage of formation of borine, which, being a strong Lewis acid, readily adds \( \mathrm{OH^-} \) with formation of the more stable species \( \mathrm{BH_3OH^-} \). The actual effectiveness of this is confirmed by investigation of the anodic behavior of \( \mathrm{BH_4^-} \) (17).

\[ \mathrm{BH_4^-} \xrightarrow{k_1=5.31\cdot 10^7} \mathrm{BH_3OH^-} \xrightarrow{k_2=3.57\cdot 10^{11}} \mathrm{BH_2(OH)_2^-} \xrightarrow{k_3=10.06\cdot 10^7} \mathrm{BH(OH)_3^-} \xrightarrow{k_4=5.68\cdot 10^{10}} \mathrm{B(OH)_4^-} \]

\[ \mathrm{BH_4^-} \xrightarrow{k'_1=5.15\cdot 10^7} \mathrm{B(OH)_4^-} \]

\[ \mathrm{BH_3OH^-} \xrightarrow{k'_2=9.44\cdot 10^7} \mathrm{B(OH)_4^-} \]

\[ \mathrm{BH_2(OH)_2^-} \xrightarrow{k'_3=10.06\cdot 10^7} \mathrm{B(OH)_4^-} \]

Table 2

| \(T\), °C | \multicolumn{2}{c}{\(k_2\), mole\(^{-1}\)·min\(^{-1}\)} |
|---:|---:|---:|
| | \( \mathrm{BH_4^- \to B(OH)_4^-} \) | \( \mathrm{BH_4^- \to BH_3OH^-} \) |
| 15 | \(5.15\cdot10^7 \pm 0.32\cdot10^7\) | \(5.31\cdot10^7 \pm 0.23\cdot10^7\) |
| 25 | \(11.14\cdot10^7 \pm 0.40\cdot10^7\) | \(10.89\cdot10^7 \pm 0.42\cdot10^7\) |
| 35 | \(20.92\cdot10^7 \pm 0.56\cdot10^7\) | \(21.58\cdot10^7 \pm 0.41\cdot10^7\) |

Although, owing to experimental difficulties, we did not study the kinetics of these processes, it may be confidently assumed that these reactions must be very fast, since the species \( \mathrm{BH_3} \), \( \mathrm{BH_2OH} \), and \( \mathrm{BH(OH)_2} \) are extremely unstable and highly reactive.

Table 3

Transition of \( \mathrm{BH_3OH^-} \) to \( \mathrm{BH_2(OH)_2^-} \)

From the foregoing it is evident that an aqueous borohydride solution is a very complex system of many components. We found their concentrations by the method of N. M. Rodigin and E. N. Rodigina (18) and determined the quantitative composition of the system at different moments of time. The concentrations of \( \mathrm{BH_3OH^-} \) and \( \mathrm{BH(OH)_3^-} \) ions remain small throughout the entire course of the process (0.03–0.05 mole %), whereas the concentration of the \( \mathrm{BH_2(OH)_2^-} \) ion reaches 25.7 mole %.

Kazan Chemical-Technological Institute
named after S. M. Kirov

Received
6 XI 1964

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