Hydrogen bonds between co-ordinated BH3 and BH2 groups and OH groups. Thermodynamics of formation by infrared spectroscopy

Sl~ectrochhnica Acta, Vol. 80A, pp. 1125 to 1131. Pergamon P r e ~ 1974. Printed in Northern Ireland

Hydrogen bonds between co-ordinated BH3 and BH2 groups and OH groups. Thermodynamics of formation by inc-eared spectroscopy M. P. BROWNand P. J. WAT.KER Donnan Laboratories, University of Liverpool, P.O. Box 147, Liverpool L69 3BX (Received 5 September 1973)

A1~;ract--The role of co-ordinated BH~ and BH 8 groups as proton acceptors in hydrogen bonding has been studied by examining complexes between phenol, as an example of a proton donor, a n d t h e co-ordinated boranes MeaN.BH~X ( X = H , C1, Br and I), py.BH s and EtaP.BH s. Some complexes of p-fluorophenol and t-butanol have also been examined. The enthalpies of formation have been determined by measuring the absorbance of the free-OH stretching band in the infrared over a range of temperature. The values of --AH fall in the range 1.7-3.5 kcal/mol. For phenol as the proton donor, the strengths of hydrogen bonds formed by MeaN-BH~,Y deorease in the order H > C1 > Br > I. INTRODUCTION I ~ P~EVmUS publications [1, 2] we r e p o r t e d the f o r m a t i o n o f a novel t y p e o f h y d r o g e n b o n d involving co-ordinated B H s or B H 2 groups as p r o t o n acceptors. T h e coo r d i n a t e d boranes, Me3N'BHs a n d MeaN.BH2X (X ---- CI, B r a n d I), for example, were shown t o f o r m h y d r o g e n - b o n d e d complexes w i t h O H c o m p o u n d s such as p h e n o l a n d methanol. Me3N.BIt 3 ~- R O H ~- Me3N'BH3 . . . .

H-O-R

T h e f o r m a t i o n of these complexes was studied, m a i n l y in c a r b o n t e t r a c h l o r i d e solution, b y observing t h e lowering (Av) of the O H stretching f r e q u e n c y in t h e infral red. T h e values of Av o f a b o u t 80-160 cm -1 are c o m p a r a b l e to those o f m a n y o t h e r systems where t h e h y d r o g e n b o n d s are k n o w n to be of a m o d e r a t e s t r e n g t h , viz. 2-3 kcal/mol. T h e relationship b e t w e e n t h e f r e q u e n c y shift a n d h y d r o g e n b o n d s t r e n g t h is specific to a given t y p e o f h y d r o g e n b o n d a n d so for a new t y p e o f h y d r o gen b o n d o n l y a rough idea o f its s t r e n g t h m a y be o b t a i n e d f r o m the value o f A~. T h e systems h a v e therefore been f u r t h e r studied and here we r e p o r t the free energies a n d enthalpies of f o r m a t i o n as o b t a i n e d b y measuring the absorbance o f the free-OH stretching b a n d as a function o f t e m p e r a t u r e a n d concentration. Phenol, p-fluorophenol a n d t-butanol are used as t y p i c a l p r o t o n donors. EXPERIMENTAL T h e m e t h o d was essentially similar t o t h a t described b y LOPES a n d T~OMPSON

[3]. (a) A:p:paratus A P e r k i n - E l m e r 125 s p e c t r o p h o t o m e t e r was used. This i n s t r u m e n t was flushed [1] M. P. BROWN and R. W. HESELTrS~, Ghem. Commun., 1551 (1968). [2] M. P. B~ow~, R. W. I~S~.T,TINE, PAmeLA A. SMI~ and P. J. W ~ , J. Chem. Sac. (A), 410 (1970). [3] MAR~ C. SousA LoPEs and H. W. THOMPSON,~pe.cSrochi~..A.Cf,a 24A, 1367 (1968). 1125

1126

M.P. B~ow~ and P. J. WAT.~r~

with dry air, the transmittance was checked with standard rotating sectors, a narrow effective slit width of about 1.5 cm -1 was used and the wavenumber calibration was checked with the standard lines of water vapour [4]. The housing of the variable temperature cell was constructed from a hollowed out, thermally insulated brass block fitted with two outer, large diameter sodium chloride plates. The cell insert was from a standard R I I C variable temperature cell and consisted of two stainless steel plates and two sodium chloride windows and it fitted snugly into a cylindrical hole in the brass block. The temperature of the cell was controlled by passing an aqueous antifreeze solution from a thermostat through the hollow block. The air spaces between the windows of the cell insert and the outer windows of the brass block were flushed with a separate dry air supply which was heated or cooled to approximately the same temperature as t h a t of the antifreeze solution. This precaution was taken because of the rather large temperature difference which existed between the instrument cell well (normally 45-50 °) and the cell itself when the latter was at its lowest temperature of 0% The temperature of the cell was measured, to better than 4-0.2 °; b y means of a calibrated copper-constantan thermocouple using a null method and a Pye Universal Decade Potentiometer. The thermocouple was inserted into a small hole drilled in one of the inner sodium chloride plates. An upper limit of temperature of abouS 50 ° was set by a slow and irreversible reaction between amine-borane and carbon tetrachloride which set in at temperatures higher t h a n this. I t was therefore necessary to set the lower temperature limit below room temperature in order to obtain a reasonable range. This was conveniently achieved by fitting the thermostat t a n k with a dip cooler (Tecam R U 5) in addition to a heater and by using a commercial glycol antifreeze mixture. A steady cell temperature of 0 ° was obtained when the circulating solution was held at about --4.5 ° in the thermostat tank. The path length of the cell insert was normally about 2 mm (5 mm for t-butanol complexes). Amalgamated lead rather t h a n teflon spacers were used because of the tendency of teflon-spaced cells to leak. Cell path lengths were measured to an accuracy of 4-0.01 mm with a travelling microscope fitted with a Mercer dial gauge. (b) Materials and procedure The amine-boranes and related compounds were prepared as previously described [2] and were freshly sublimed before use. Phenol (m.p. 41 °) and p-fluorophenol (m.p. 48 °) were resublimed before use and t-butanol was distilled and a middle cut taken. Carbon tetrachloride was dried and distilled from 1)205 and stored in the dark. Solutions containing amine-boranes were always freshly prepared as they sometimes slowly decomposed on standing. A matched cell similar to the sample cell, but unheated, and containing pure CC14 was placed in the reference beam. Spectra were recorded, aS each temperature, once on a moderately fast scan mode from 4000-3000 cm -1 to record both the freeand bonded-0H bands and to determine I 0 (base line percentage transmission). Although the separation of free and bonded bands is good and compares favourably [4] I.U.P.A.C. Report, Tables of wavenumbers for the calibration of infrared spoetromoters, Butterworths, London (1961).

Hydrogen bonds between co-ordinated BH 3 and BH 2 groups and OH groups

1127

with most other systems of similar hydrogen-bonded complexes, there is a small overlap between the two bands. This m a y be seen in the previously published spectra [2]. However, simple extrapolation suggests t h a t at the frequency of the free-OH band maximum, overlap by the bonded band is slight or even insignificant. Consequently the most reliable (and least subjective) estimate of I 0 was considered to be the base line on the higher frequency side of the free-OH band. The free-OH band was finally recorded twice on the maximum expansion mode of the wavenumber scale to obtain I (percentage transmission at the band maximum) and the peak frequency. Spectra were recorded at 10 ° intervals usually over the range 0-50 ° b u t sometimes 10-50 ° where solubility problems arose. To check for a n y cell leakage or decomposition of the solutions, a recording a t 20 ° or 30 ° was first made, the temperature lowered to 0 ° and then raised stepwise after each recording. At the completion of the run a repeat recording at 20 ° or 30 ° was made. At each temperature, 30 rain was allowed for the cell to reach thermal equilibrium. (c) Determination of eq~tilibri~tm constants The 1:1 stoichiometry of the complexes has previously been established [2]. l%om the known initial concentrations of boron compound and OH compound and

L'BH2X + R O H ~- [L'BH2X . . . . HOR] (L = Me~N, py, or Et3P; X = H, C1, Br or I) the concentrations of free-OH compound as determined directly by the infrared measurements, the necessary concentrations of complex and of uncomplexed boron compound can be readily derived and the equilibrium constants calculated. The free donor concentration was calculated from the following relationship:

1 Io = ~ log T where e is the absorption coefficient, l the cell path length and c the concentration. The applicability of the relationship was verified using solutions of donor compound of several different but known concentrations at a constant temperature (20°). Some typical results for phenol are shown in Table 1. The concentrations were such t h a t the absorbances of the free-OH band more t h a n covered the range of values shown by solutions of the complexes. The molar absorbance values can be seen to be in satisfactory agreement with one another. Table 1. Molar absorption coefficients (e) for phenol in CC14from 0.004 to 0.010 ~ (Temperature, 30°; cell path length, 0-1954 cm). Concentration (~s)

I°

I

~(1 reel-1 cm-1)

0.010 0.008 0.006 0.004

96.8 96-8 97.0 97.0

33.5 40.7 50-0 62.7

235.8 240.7 245.6 242.5

1128

M . P . BRow~ and P. J. WA~.w~.~

Since s varies with temperature, other calibration runs were carried out at fixed concentration to determine s values at the appropriate temperatures. The average values obtained for each proton donor are given in Table 2. For these determinations Table 2. Molar absorbance values (e) as a function of temperature e(1 reel-1 cm-1) Temperature

phenol

p-fluorophenol

t-butanol

0° 10° 20° 30° 40°

263"3 256"7 250"8 241.5 231.1 221.6

248"8 239"8 233"8 226.4 216-0 205.6

78"2 74"1 70"3 65-3 61.0 57.3

50 °

the concentrations of OH compound were slightly less than the initial concentrations normally taken in the solutions containing complexes. Therefore, allowing for some complexing, the actual absorbance values are very similar and possible errors mlnimised. The values for phenol compare excellently with results (also in CC14 solution) reported by BAKER et al. [5]. They report an "optical density", E~s of 0.562 1 mmo1-1 cm -1 which corresponds to a molar decadic absorption coefficient e~5 of 244.0 1 mo1-1 cm -1. Solutions of complexes were made up using the following concentrations of proton donors: phenol and p-fluorophenol' 0-010-0.015 M; t-butanol, 0-008-0.015 M. The values resulted from a compromise between the need to have a high enough concentration to allow use of a cell having a reasonably short path length and yet low enough to avoid interference from self-association of the donor. Phenol at 0.02 •, for example, undergoes self-association only to the extent of about 1 ~ . Convenient concentrations of boron compounds were of the order of 0.1-0.3 M, allowing a reasonable degree of eomplexing. I n each system examined, three runs were carried out on separate solutions and the appropriate K vaIues determined. AH values were determined in the usual way from these sets of K values. The usual graphical method and a least squares computer fit gave results in close agreement and the values reported here are those obtained by the latter method. A typical set of results is t h a t for the MeaN.BHa-PhOH system given in Table 3. The results for the three separate runs m a y be compared and lead to - - A H values of 3.26, 3.21 and 3.35 respectively and an average of 3.27 kcal/mol. As m a y be seen from Table 3, the unweighted least squares programme gives standard deviations of the order of ~:0.1 kcal/mol, the individual values being shown in the table. Other results with phenol and with the other proton donors were generally similar to the example given in the table with standard deviations of about ±0.1 kcal/mol for each run. The three --AH values obtained for each system had an overall spread of no more than 0.17 kcal/mol except in one instance, the MeaN.BH~Br--PhOH system, where the individual values were 2.81, 2-88 and 3.04. [5] ilK.N. BA3g.ER,H. 0. KEI%LII~TGERand A. T. S~ULGI~,Spectrochim. Acta 20A, 1467 (1964).

Hydrogen bonds between co-ordinated B H s a n d BH~ groups a n d OH groups

1129

Table 3. Detailed results for the Me~N.BHs--PhOH system a C o n c e n t r a t i o n ( × 10 - a ~ )

K T°

10

1

Free PhOH

Complex

Free MosN'BH s

(1 reel -1 )

4.51 5.22 5.85 6.39 6.88 7-38

7.99 7.28 6"66 6"51 5.62 5.12

192.0 192-7 193-4 193.5 194.4 194.9

9.22 7.34 5-89 5-26 4.26 3-55

5.29 6.38 6.70 7.15 7.63 8.14

7.22 6.12 5.81 5.35 4.87 4"36

142-8 143.6 144.2 144.7 145.1 145.6

9.56 7.25 6.01 5-17 4.40 3.68

5.51 6.19 6-99 7.72 8"35 8.94

9"49 8-81 8"01 7.28 6-65 6"06

190"5 191.2 192-0 192.7 193"4 193.9

9.04 7.42 5-97 4.90 4"12 3"49

--AR

--AS

(kcal.mo1-4)

(o.u.)

3-26 -4- 0"12e

7.55

3.21 :~: 0.12 e

7-37

3"35 -4- 0"03e

7.91

l ~ u n 1b 0 10 20 30 40 50

96.8 97.8 97.5 97"7 97.5 97.9

0 10 20 30 40 50

97.3 97.6 98"6 98.9 98.8 98.6

0 10 20 30 40 50

96.6 97-1 97.0 96.8 96-7 97"5

56.8 53.5 50-4 48.8 47.6 46.9 R u n 2c 52.0 47.1 45.3 44.5 43.7 42,8 S u n 3d 50.3 47-5 44.1 41.9 40'6 40-0

a I n CCI4 s o l u t i o n , b I n i t i a l c o n c e n t r a t i o n s : M % N . B H s , 0.2 M; P h O H , 0.0125 M. e I n i t i a l e e n e e n t r a t i o n s : M % N . B H 8, 0.15 M; P h O H , 0.0125 M. d I n i t i a l e o n o e n ~ r a t i o n s : M % N . B H a , 0-2 M; P h O H , 0.015 M. • S t a n d a r d deviations from least squares programme.

Reproducibility is therefore generally better than ±0.1 keal/mol and the results are c o n s i d e r e d t o b e a c c u r a t e w i t h i n a b o u t t h e s a m e l i m i t s . R e p r o d u c i b i l i t y i n - - A S

is generally about =L0.3 eu b u t the accuracy here is probably less and is likely to be q-1 eu. As a test of the experimental method, the previously well studied interaction between ethylaeetate and phenol was examined. The average value of - - A H found was 4-57 ±0.1 keal/mo! and this is in reasonable agreement with recently published values of 4.45 ± 0.05, b y POWELL and WEST [6] from a near infrared study, and 4-75 ± 0-08, b y ARNETT et al. [7], and 4.77 =k 0.1, b y EPLV,y and D~AOO [8] from calorimetric measurements. RESULTS

AND

DISCUSSION

The measured - - AH, -- AS and K20° values are given in Table 4 and are compared with A~(cm -1) values reported previously [2]. As can be seen, the - - A H values for interactions with phenol and p-fluorophenol lie within the rather narrow range of 2.7-3.5 kcal/mol. Values of this magnitude are roughly consistent with those expected for frequency shifts of about 150 cm -1 on the basis of comparisons with other hydrogen-bonded systems. [ 6 ] D . L . P O W E L L a n d R . W ~ . s T , Sl~ctrochim. Acta 20A, 9 8 3 ( 1 9 6 4 ) . [ 7 ] E . M . A R N E T T , L . J ' O R I S , E . I ~ I T C H E I ~ , T . S . S . R . M U R T Y , T . 1~. G O R R I E a n d

S c a L a r s , J. Am. Chem. See., 92, 2365 (1970). [8] T. D. EPLEY and R. S. D~Aoo, J . A m . Chem. See., 89, 5770 (1967).

P. v. R.

1130

M. P. BROWN and

P.

J . WALF~F.R

Table 4. Thermodynamic data for the co-ordinated borane complexes K20o

(1 reel-i )

--AH (keal.mo1-1)

--AS (o.u.)

Av~

(ore-1)

(a) With phenol MosN.BH s

5-96

3.27

7.61

EtsP.BHs Py.BHs M%N'BHsCI MosN.BHaBr M%N.BHsI (b) Withp-fluorophonol

7.32 7.69 5"88 4-70 3.48

3.32 3.30 3.10 2.91 2.72

7-37 7-30 7.08 6-86 6.81

142 b 141 148 142 153 149

8"57 11-93 6.44

3"44 3"51 2.94

7"48 7"13 6"36

148 146 158

1-07 1.23 1.19

1.71 1.90 1-70

5.68 6-06 5-44

--

MosN'BH s EtsP.BH s MosN.BHsBr

(c) With t-butanol MOsN.BH s EtsP.BH s Py.BH s

--

a Previouslypublisheddatas. b Experimentalerror in Av is about -4-3cm-1. p-Fluorophenol, being a stronger acid t h a n phenol, might be expected to form stronger hydrogen bonds t h a n phenol. The results however do not clearly demonstrate this. Although the t e nde nc y of p-fluorophenol to form complexes is much greater, as seen from comparison of the K values, t he hydrogen bond strengths, as measured b y - - A H values, are only marginally greater, differences being hardly significant in relation to estimated experimental error. Predictably, however, interactions with t-butanol are weaker on the basis of both K and - - A H values. E n t r o p y values for all three donors are r at her low but not uniquely so compared to other similar systems. F o r a given proton donor the spread of - - A H values is small. For example, the three B H 3 compounds with phenol give - - A H values which differ less t han the limits of experimental error. The same is true for interactions of two of these BH3 compounds with p-fluorophenol. Substitution of one of the H atoms of MesN.BH s b y halogen, however, does appear to lower the hydrogen-bonding ability, both in terms of K and --AH. For compounds Me3N.BH2X, the hydrogen-bonding strength decreases in the order H > C1 > Br > I. A possibility is t h a t the increasing inductive effect along this series progressively lowers electron density on the attached BHz group and results in progressively weaker hydrogen-bonds. In conclusion, this investigation demonstrates t h a t the hydrogen bonds between these co-ordinated boranes and OH compounds are of moderate strength. The -- AH values found, which range from 1.7 to 3.5 keal/mol, m a y be considered in relation to a range of about 1-10 kcal]mol within which most hydrogen bond strengths fall. Many examples are known of OH . . . . 0 donor-accept er hydrogen bonds, for example, with phenol as the proton donor and these usually fall in the range 3-7 kcal/ mol [9]. Typical examples are phenol complexes with acetone ( ~ 4 . 6 kcal/mol), [9] A. S. N. MurPHY and C. N. R. RAo, Spectroscopic studies of the hydrogen bond. In: Applied Spectroscopy Reviews. (Ed. E. G. B~aA~r~.),Vol. 2, pp. 69-191, Marcel Dekker, New York (1969).

Hydrogen bonds between co-ordinated BH 8 and BH 2 groups and OH groups

1131

diethylether ( ~ 5 . 1 kcal/mol) and ethylaeetate (~-~4.6keal/mol). Thus the coordinated boranes selected for this investigation appear to be only slightly weaker than divalent neutral oxygen compounds in proton acceptor strength. .4c/~owle~tgerr~n~-V~e are grateful to Professor H. W. Thompson for assistance with the experimental method.