NMR studies of hydrogen bonding of phenol with several proton acceptors1

Joumol of Molecular Structure Elsevicr Publishing Company, Amsterdam.

NMR STUDIES OF HYDROGEN PROTON ACCEPTORS*

253

Printed in the Netherlands

BONDING

OF PHENOL

WlTH SEVERAL

THOR GRAMS-I-AD+’ AND EDWIN D. BECKER National Institute of Arthritis Md. 20024 (U.S.A.) (Received

and Metabolic

Diseases,

National

Institutes of Health, Bethesda,

July 18th. 1969)

ASTRACT

Proton magnetic resonance studies are reported of self-association of phenol in carbon tetrachloride and in cyclohexane, and of hydrogen bonding between

phenol and two carbonyl and several phosphoryl compounds. The NMR chemical shift, S,, of a proton in a hydrogen bond has been found to be dependent on both the temperature and the concentration of the proton acceptor. Furthermore, the chemical shift data at low concentration of proton acceptor suggest the presence of a 2: 1 proton donor-acceptor complex in addition to a 1: 1 complex. A linear relationship exists between the change in chemical shift and O-H stretching frequency shift upon complexation of phenol with phosphoryl compounds.

lNTRODUCTION

Proton NMR has been applied extensively to the study of hydrogen bonding. For example, GrBnacher* has shown that there exists for binary solutions of phenol in different solvents a linear correlation between the phenol hydroxyl proton chemical shifts and the corresponding shifts of the infrared hydroxyl stretching frequency, Av(OH). A similar correlation has been found for intramoIecular hydrogen bonding by Reeves ct aL3 -‘, who determined the equilibrium constants and AH values for the cis-frans conversion in orrho-halophenols. Mathur et aL6,

have reported NMR studies of hydrogen bonding bctwccn the S-H proton of benzenethiol and various proton acceptors. By treating the variation of the chemical shift of the S-H proton as a function of the acceptor concentration, they obtained both the chemical shift of the complex, 6,. and the equilibrium constant. The same method has been utilized by Takahashi and Li’ to study the association of 2propanol and chloroform with various proton acceptors. Eyman and Drago’ have l

l Paper XVII of a serieson hydrogen bonding, for Part XVI see ref. I. * Visiting Scientist, 1966, permanent address; Kjcmisk Instilutt, Univcrsitctct

Norway.

1. Mol. Structure,

i Bcrgcn, Bergen.

5 (1970) 253-261

T. GRAMSTAD,

254

E

D. BECKER

demonstrated a linear correlation between chemical shifts, a,, and Av(OH). (Previously Joeston and Dragog had reported an approximate linear relation between Av(OH) and AH for a number of hydrogen bonding systems.) Eyman and Drago also suggested a method for evaluating anisotropic contributions in proton acceptors. The main purposes of the present work were (a) to study the temperature dependence of the NMR spectrum of a proton in a hydrogen bond, and (b) to investigate the effect of solvent interaction on the hydrogen bonded complex.

EXPERIMENTAL

Phenol was purified by repcatcd recrystallization from petroleum ether and the white needles obtained were dried over phosphorus pentoxide in a desiccator. Carbon tctrachloridc and cyclohexanc wcrc purified chromatographically by using Molecular Sieve Type 4A and basic aluminum oxide. The proton acceptors (Table 2) were redistilled or recrystallized and liquids were chromatographed through basic aluminum oxide just before use. All proton NMR spectra were obtained on a Varian A-60 spectrometer with temperature regulating accessories. The temperatures were determined using a methanol or an ethylene glycol sample and the calibration data furnished by Varian Associates. Temperature accuracy is estimated to be f2” *. Chemical shifts of the O-H signal in phenol were measured with respect to tetramethylsilane (TMS, 0.5 OA). All lines were downfield from TMS. Chemical shifts were measured by interpolation between audio side-bands with a 50 Hz sweep width. The reported shifts (in p.p.m.) arc accurate to at least 0.003 p.p.m. The solutions were made up by weighing, and the molarity of each solution was calculated at various temperatures from the change in density with temperature of carbon tetrachloride and cyclohexane. Since the concentration of phenol was normally below 0.05 M it was sufficient for most purposes to assume that the density of the solution was the same as that of the solvent.

RESULTS

AND

DISCUSSION

Self-association

of phenol

The results of the dilution

study of phenol

in carbon

tetrachloride

and in

+ Recent work in this laboratory (R. R. Shoup, unpublished) and clscwhcrc’” has shown that the Varian calibration curves arc somewhat in error. In the temperature range S-50 “C employed here the true temperatures arc about 1” to 2.5” below those reported. This systematic error is of little consequence in this work; in particular the use of equilibrium constants for the supposed tcmpcraturcs instead of the actual temperatures can bc shown to cause negligible changes in the results. J. Mol. Slructrue,

5 (1970) 253-261

NMR

STUDIES

OF HYDROGEN

BONDING

OF

PHENOL

255

cyclohexane at various temperatures is presented in Fig. 1. As can be seen, the resonance of the phenol 0-H is, at low concentration, shifted further downfield in carbon tetrachloride than in cyclohexane, whereas the opposite is the case with increasing concentration. The chemical shifts, 6 r, found by extrapolation to zero concentration of phenol in carbon tetrachloride and cyclohexanc, are given in Table 1. Since the temperature coefficient in cyclohexanc (0.0022 p.p.m./“C) is almost as large as that in Ccl.+ (0.0027 p.p.m_/“C), hydrogen bonding of phenol to

0.05

r

0.1

I-

i--

0.2

0.15

0.25

PnENOL

Fig. I. Chemical TABLE

shift of phenol OH proton

0.35 (molar/l

0.4

0.45

INFINITE

DlLUTION

1

in CC14and in cyclohcxane.

1

TEMPERATURE

AND

SOLVENT

Chemical Solcent

EFFEECT

shift.

ON

6,.

THE

O-H

RESONANCE

OF

PItENOL

AT

in p.p.m.

5”

20”

35”

50”

4.327

4.292 4.070

4.258 4.040

4.223 4.017

__.. ccl, CsH,,

0.3

CONCENTRATION

1. Mol. Strucrurc. 5 (1970) 253-261

256

T. GRAMSTAD,

E. D. BECKER

the solvent is ruled out as the principal cause of the temperature variation. Other, less specific solvent effects, such as van der Waals forces, arc probably responsible*. Complexing

of phenol

with phosphoryl

compounds

In general, the observed NMR chemical shift, &_, equilibrium with a donor-acceptor complex is

of a proton

donor

in

(1)

6ObY

where 6, is the chemical shift of the non-bonded donor, 6, is that of the donor proton in the complex, and Cr, C, and C,, are, respectively, the free, complexed and total donor concentrations. When the association constant, K, is rather small, it is impossible to obtain essentially 100 oA complex; hcncc 6, cannot bc obscrvcd directly, and indirect methods of finding both S, and K are used (e.g., the graphical procedure of Mathur et a1.6, or a computer-aided iterative procedure). For some

. 8.567p.p.m.

it

o dxob= = 8.5OOp.pm. = 8.433 l

02

oI2

ot 20°

p.pm.

ot 35O

JxobSt 8.333p.p.m.

al4

ot 50°

03

1 shift of phenol OH proton for the system (CIHSO)sPO/CcH,OH (C&O),

Fig. 2. Chemical

05

5’

PO

CONCENTRATION

0!6

(moles/l

in Ccl..

l The depcndcncc of the OH chemical shift of phenol on solvent was also studied at room tcmpcrature by using a mixed solvent of Ccl4 and cyclohexane. The OH resonance frequency varied fineurfy with mole fraction Ccl.; hence a treatment of 1: 1 complex formation between phenol and CCL (for example, by the method of Mathur et aL6) is precluded.

f. Mol. Structure, 5 (1970) 253-261

NMR STUDIES

OF HYDROGEN

BONDING

257

OF PHENOL

of the present phosphoryl acceptors, however, the association with phenol is so strong that > 99 oAof the phenol is in the complexed form with moderate (C 1 M) concentrations of phosphoryl acceptor and low concentrations of phenol (0.05 M). In these cases 6, can be found directly and the fraction of phenol complexed at each acceptor concentration determined from cqn. (1) and the measured values of &,,, 6r and C,, the acceptor concentration. Fig. 2 shows plots at different temperatures of the OH chemical shift of phenol as a function of concentration of the acceptor tricthyl phosphate in Ccl.,. At all temperatures the plots are characterized by a rapid change in Sot,, at low concentration with a much smaller variation above the concentration of acceptor of about

0.2

M.

From

tion constants

infrared

spectra

Aksncs

and

of 100 to 500 in this temperature

Gramstad”

range.

have

found

associa-

Thus,

for C, > 0.2 to 1.0 M (depending on temperature), more than 99 % of the phenol is in the complcxed form, and 6, M do,,,_ The small change in fiob, in this high concentration region is probably due to the effect of variation of the medium surrounding the complex from essentially pure Ccl,, to a mixture of CCL, and (C2HSO)3PO*. From a direct plot of the data (see, for example, Fig. 2) we can find the value of 6, (hereafter called bxob’ to distinguish the value obtained in this way from a value for 6, that we shall obtain later). This simple treatment is, of course, applicable only to those systems that bond strongly, since, in the less strongly bound complexes, considerably less than 100 ‘A of the phenol is in the complexed form even at high acceptor concentration. Values of S,““l for the strongly bound complexes are given in Table 2. TABLE

2

CHEMICAL -

I II III IV V VI VII VIII

SHIFTS

(p.p.m.)

OF PHENOL OH .__~

PROTON IN IIYDROGEN --.--~

Proton ucceptor

6 cs1 (in CC/.) _=20” 35”

GHsO)IY(OKCI,

7.633 7.867 8.183 8.467’ 9.433 9.600 8.450 7.667

(C~H,O)lP(0)CIKI~ (C211s0)1P(0)CH2CI (C2~~5O)lPO (C6H,),PO

[(CIidzNld’O DC(O)N(CD,), CHI(CH2).C-0

_.50”

7.500 7.233 7.750 7.533 8.067 7.967 8.433 8.333 9.333 9.233 9.533 9.483 8.400 8.250 7.600 7.500

BONDED -

COMPLEXES -__.

blob’ (in CC/,) -.--20” 35”

50”

6 *’ (in CeHII)’ I. __ 20” 35” 5o”

8.500d 9.500 9.583 8.367 -

8.333 9.250 9.433 7.900 -

7.833 8.217 8.517 8.800 9.767 -

8.433 9.367 9.517 8.167 -

7.650 8.017 8.367 8.717 9.700 -

7.400 7.833 8.200 8.633 -9.617 -

’ See text for definitions of terms. b In cm-’ at 20" in CC&. a At 5”, 8.517 p.p.m. ’ At So, 8.567 p_p.m. l A rough calculation of the reaction field effcct*2, which should bc the dominant cffcct change in medium from non-polar to polar, indicates that the observed downfield shift of 0.08 p.p.m. is of the correct magnitude. (The calculation was actually carried out for phenol not the 1:I complex, in a medium changing from CC& to tri-o-cresyl phosphate. Clearly useful only as an order of magnitude calculation.)

J. Mol.

Structure,

upon about itself, this is

5 (1970) 253-261

dv(OH)b

244 270 305 330 415 460 280 235

258

T. GRAMSTAD,

Combined

use of IR and NMR

E. D. BECKER

data

Since all of the systems studied here have previously been subjected to careful In investigation’ t. r 3*I4 (using Ccl, solutions only), WC can use the values of K determined from the IR work in the treatment of the NMR data. For known values of K, C, and C,, WCcan find C, and Cr from eqn. (2) and then use these values in eqn. (l), K=

cX Cf

w.

-

(2) Cx)

together with the measured values of bob, and values of & obtained from Fig. 1, to calculate 6, (hereafter called 6:” ). The application of the equilibrium expression implied by eqn. (2) alone is strictly valid only if there is no self-association of phenol. By working at very low phenol concentrations we have minimized such self-association but have not eliminated it. To a high dcgrce of approximation we can, however, correct easily for this effect by using for a,, not the value extrapolated to C,, = 0, but the value of Sobs for phenol alone at the total concentration of phenol in the complex-forming system6. 13.0

12.0

I?

= 8.517p.pm.

11.5

8% -

11.0

o

d:%’

= i3.467p.pm.ot

8

qxt

=

8.433

.

dfxr

=

0333p.p.m.

ot 5O 200

pp.m.

ot 3S”

at 50°

9.0

1

. .

.

.

0.1 .

.

. -,

0.2

.

(c,HeO),

.

_

_ .,.

PO

.

0.3

CONCENTRATION

. _,

0.4

.

_

.

0%

0:s

hdr~ll)

Fig. 3. 6xu1c as a function of proton acceptor concentration for the system (CIH,O),PO/ CaH50H in CCL.

J. Mol. Structure,5 (1970) 253-261

NMR

STUDIES

OF

HYDROGEN

BONDING

259

OF PHENOL

A plot of &=“= vs. C,, is shown in Fig. 3 for phenol plus triethyl phosphate. It is apparent that bxCr’Cdepends upon both temperature and C,: it decreases linearly and rather gradually as C,, dccreascs and then increases rapidly at C, < 0.08 M. The linear variation is presumably due to the solvent effect discussed previously. By extrapolating bXc’Ic along the linear portion of the curve to C, = 0 we obtain SrCX’,which has thus been corrected insofar as possible for solvent effects. The values of ~5~“’ from Fig. 3, as well as those for other systems (obtained in the same way) arc given in Table 2. At low C,, bxc=‘= increases rapidly. Since C, becomes small in this region, the validity of the sharp rise in 6Xc*‘cwas open to question. However, calculations showed that any reasonable variation in Gobi, &, or K could not account for the observation. (Errors of 0.07-0.50 p.p.m_ in 6,, or 6, would be required to make sx==‘=linear down to C, = 0.001 M. Alternatively, errors of even 50 % in K could not begin to linearize the curve.) Consequently, we found it reasonable to assume that the deviation is due to the formation of 1:n proton acceptor-donor complexes. Since the concentration of donor and acceptor is relatively low we believe that the reaction mixture contains predominantly 1: 1 and I:2 acceptor-donor complexes. This assumption stems also to fit very well with infrared spectroscopic results. Fig. 4 shows the infrared absorption spectra in the O-H stretching region for the system (C2HS0)2P(0)CHC12/ C6HSOH at two concentration ratios. There is a considerable difference in the breadth of the hydrogen-bonded v(OH) band near 3350 cm- * as the ratio of phenol

01 4ooo

35bo FREOUENCY

=

2;oo

(cm-‘)

Fig. 4. Infrared spectra in the OH stretching region of the following solutions in Ccl,: 0.05 M phenol and 0.15 M (CIHSO)IP(0)CHCIz (II); (b) - - - - . 0.15 M phenol and (a) 0.05 M (II); (c) - - - - - 0.15 h4 phenol and 0.05 M (II) compensated with a reference cell containing 0.05 M phenol and 0.15 M (II). (a) shows a band near 3350 cm-’ ascribed to a 1: I complex, while in (b) the analogous band has a contribution from a 2:l acccptor+lonor complex, as well. In (c) the band due to the I:1 complex should be nearly cancelled. 1. Mol. Structure,

5 (1970)

253-261

260

T. GRAMSTAD,

to phosphoryl

compound

is varied. We assume that the broadening

due to the prcscnce of a 1:2 proton acceptor4onor

E. D. BECKER

in contour

is

complex in zddition to the 1: 1

complex. By cancellation of the 1: 1 complex by means of difference spectra (Fig. 4) there appear two new absorption bands I and 2 which may be assigned as v(OH) vibration bands in the 1:2 complex 2 p-0.

- - . H-0.

I - - - H-0.

R

R

Temperature dependence of 6, Muller and Reiter” have shown theoretically that the observed variation with temperature of the NMR chemical shift of a proton involved in a hydrogen bond does not entirely arise from changes in the association equilibrium, but also to excitation of the hydrogen bond stretching vibration mode, i.e., from changes in the effective length of the OH * * -0 bond. Takahashi and Li’ have also reported data which led them to question the validity of the assumption that the NMR chemical shift, S,, of a proton in a hydrogen bond is temperature independent. For the systems we have studied dXeX’is clearly temperature dependent, as indicated in Fig. 3 and Table 2. It is interesting to note that the temperature dependence of the chemical shift for a proton in a hydrogen bond has been calculated for a general OH.... 0 bond by Muller and Reiter” to be within 0.002-0.008 p.p.m./“C, whereas we have found that the values lie between 0.001 and 0.018 p.p.m./“C (see SXCX,values in Table 2). Howcvcr, our data show no obvious correlation between the strength of the hydrogen bond and the temperature dependence of the chemical shift. Correlation between NMR and IR frequency shifrs Fig. 5 shows that there exists a linear relationship bctwccn the change in chemical shift, Aa,, and O-H stretching frequency shift, Av(OH), for the system organophosphorus compounds/phenol with maximum deviation of kO.13 p_p.m. Eyman and Drago’ reported a similar relationship which holds within f0.38 p.p.m. Since different solvents (Ccl, and CH,CI,) have been used, a direct comparison of our results with theirs is made difficult. (For example, they found the values of 6, at -60” and in CH2Cl, to be, 10.630, 10.630, 9.230 and 9.730 p.p.m. for the systems [(CH,),Nl~P(0)/C,H50H, (GH5)3P(0)/GH50H, (GH50)d’(0)/ C6H50H and HC(0)N(CHJ),/C6H50H, whereas we have calculated for the same systems at 20” and in CClc 9.600,9.433,8.467 and 8.450 p.p.m., respectively.) Fig. 5 also includes data for two carbonyl acceptors, cyclohcxanonc and N,N-dimethylformamide-cl,. While the point for cyclohexanone is nearly on the line defined by the phosphoryl data, that for DMF is rather far removed. Eyman and Drago found J. Mol. Strucmre, 5

(1970)

253-261

NMR STUDIES

OF HYDROGEN

I

0

250

.

.

..I

BONDING

..,

.,....,

300

.,..,....I

350 IR

SHIFT.

261

OF PHENOL

400 Al&,

450

500

km?

Fig. 5. Change in OH chemical shift of phenol OH on hydrogen bonding to various acceptors (in Ccl.) vs. change in OH stretching frequency on hydrogen bonding. A& = &car-& at 20”. Av(OH) from refs. 11.13 and 14. (0) Phosphoryl compounds (I)-(VI); 0) carbonyl compounds (VII) and (VIII). See Table 2 for identification of compounds.

a consistent deviation from the linear correlation with carbonyl acceptors and suggested that a constant difference for such acceptors is associated with the magnetic anisotropy of the C.-O group. Our limited results suggest that the nature of the acceptor and the strength of the hydrogen bond may be factors in determining the magnitude of the anisotropy contribution. Tables giving the observed NMR frequencies, as well as values of Cr and 8xca’c,are available from either of the authors on request. REFERENCES I T. GRAMTAD AND G. VAN BIN=. Specrrochim. Acta. 22 (1966) 1681. 2 1. GR~~NACHER, Helu. Phys. Acru, 34 (1961) 272. 3 L. W. REEVES, Can. J. Chem., 38 (1960) 736. 4 L. W. REEVES, E. A. ALLAN AND K. 0. STR~MME, Con. J. Gem., 38 (1960) 1249. 5 E. A. ALLAN AND L. W. m, J. Phys. Chem., 66 (1966) 613. 6 R. MATHUR, E. D. BECKER, R. B. BRADLEY AND N. C. Lr, J. Phys. C/rem., 67 (1963) 2190. 7 F. TAKAIMSHI AND N. C. LI, J. Phys. Chem., 68 (1964) 2136; 69 (1965) 2950. 8 D. P. EYMAN AND R. S. DRAGO. J. Am. Chem. Sot.. 88 (1966) 1617. 9 M. D. JOESTON AND R. S. DRAGO, J. Am. Chem. Sot.. 84 (1962) 3817. 10 A. L. VAN GEET, And. Chem., 40 (1968) 2227. 11 G. AK~NE~ AND T. GRAMSTAD, Acru Chem. Scund., 14 (1960) 1485. 12 A. D. BUCKINGHAM. T. SCHAEFER, AND W. G. SCHNEIDER, J. Chem. Phys., 32 (1960) 1227. 13 T. GUMSTAD, Spectruchim. Actu, 19 (1963) 497. 14 T. GRAMSTAD AND W. J. FUGLEVIK, Acfu Chem. Scan& 16 (1962) 1369. 15 N. MULLER AND R. C. REAR, J. Chem. Phys.. 42 (1965) 3265.

J. Mol.

Structure,

5 (1970) 253-261