India ,Journal 01CbemlstryVol. 14A, February 1976, pp. 104-186

Electronic Absorption Spectra & Hydrogen Bonding: Phenols withDioxane, Ethyl Acetate & '3 Few Tertiary Amines

S. B. SHAH & A. S. N. MURTHY·

Department of Chemistry, Birla Institute of Technology & Science, Pilani 333031

Received 14 October 1974; accepted 4 April 1975

The hydrogen bonding between phenol as electron acceptor and dioxane, ethyl acetate, and anumber of tertiary amines as electron donors has been studied in cyclohexane by electronicabsorption spectroscopy. The thermodynamic and spectroscopic parameters have beenevaluated and their significance discussed. Hydrogen bonding interaction of para-nitrophenoland sym-tribromophenol with these donors has also been investigated. Electronic absorptionspectroscopy has been found to be a useful technique to Investtgate hydrogen bonding pheno-menon.

BABA and Suzuki! have investigated thehydrogen bonding of phenol, oc-naphtholand ~-naphthol with dioxane by electronic

absorption spectroscopy. Besides obtaining theequilibrium constants and enthalpies for hydrogenbond formation, the nature of electronic transitionsand their behaviour in hydrogen bond formationhave been discussed. The hydrogen bond energiesfor ground and excited state and the factorsresponsible for hydrogen bond effects were discussedin detail. Kubota- has carried out a detailed studyon the hydrogen bonding of phenol and naphtholswith trimethylamine N-oxide. Singh et a/.3 deter-mined the equilibrium constants of a few systemsusing electronic absorption spectroscopy and foundgood agreement WIth those obtained either byinfrared or NMR spectroscopy. .

In the present study, the hydrogen bonding ofphenol, p-nitrophenol and tribromophenol withvarious tertiary arnines has been investigated em-ploying electronic absorption spectroscopy and theequilibrium constants (K) and enthalpies (-AHO)for the 1: 1 complex formation have been deter-mined. Spectroscopic properties such as hydrogenbond frequency shifts, molecular extinction coeffi-cients, oscillator strengths and transition momentshave also been evaluated. A comparative studyon the electron donating ability of tertiary amineswith varying substituents to form hydrogen bondshas not been reported in the literature so far.Besides, the effect of the acidity of proton donor~m the. extent and strengt~ of hydrogen bondinginteraction would be of interest. The objective~as to estimate the reliability of electronic absorp-tion spectroscopy in comparison with infrared andNMR spectroscopy in investigating hydrogenbonding,

Materials and MethodsPhenols were purified by distillation and fractional

crystallization to constant melting points. Tri-ethylamine, tri-isopropylamine, tri-s-butylarnine,.--------_._--

-To whom all the correspondence should be addressed.

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dioxane, ethyl acetate and the solvent cyclohexanewere purified by standard methcds+, The electronicabsorption spectra were recorded on a Perkin-ElmerHitachi 139 spectrophotometer equipped with atemperature regulated cell holder. The tempe-rature was kept constant within ± 0'5°. Freshlyprepared stock solutions were used and kept indark except during sampling. Matched quartzcells (stoppered) of 1 em path length were used.The concentration of electron donor was alwayskept greater than the concentration of electronacceptor. The concentration of any of the phenols(electron acceptors) was kept as low as possible(--4 X to·tM) to prevent self-association. The refer-ence solutions were of the same concentration ofelectron donors used in complexing with phenols.

The increase in absorbance at the shifted bandwith increase in donor concentration was employedfor the calculation of equilibrium constants. Theequilibrium constants were calculated from therelation (1) given by Baba and Suzuki-,

(£ 1fif) =[1/K(£b-Ef)](~)+[I/(fib-fi/)J ••. (1)

where <t is the molar extinction coefficient of thenon-hydrogen bonded or free solute molecule andEb is that of the hydrogen bonded molecule. £ isthe molar extinction coefficient as observed fora solution in which the initial concentration ofelectron donor is C. From the equation a plot ofl/(fi-fi/) vs l/C is expected to be linear and equi-librium constant can be calculated from the slopeof the linear plot. fi'S in above equation may bereplaced by the corresponding absorbances A's.Thus 1/(A -At) was plotted against IIC for theshifted bands. These plots were linear and theequilibrium constants (K) were calculated fromthe slopes of these plots (Fig. 1).

The enthalpies (-.lHO) w~re calculated fromthe equilibrium constants at different temperatures.The uncertainty in K is less than 10% and theuncertainty in -AHo is ± 0-5 kcal mole=. Theentropy of activation, -ASo was evaluated fromthe relation AGo= AHo-TASo. The oscillator

8r----r---,----.----r----.-~~--~SHAH & MURTHY: HYDROGEN BONDING BETWEEN PHENOLS & AMINES

6

10 20 30 40 50 60

40 120 200

1

C

Fig. 1 - Typical plots for calculation of equilibrium con-stants by Baba and Suzuki's equation for the system incyclohexane [Triethylamine with (A) phenol at differentwavelengths (28°), (B) p-nitrophenol at 307 nm (25°),

(C) tribromophenol at 320 nm (26°)1

strengths (f) for the shifted bands and for pureacceptors were calculated from the relation

f = 4·319x 1O-8(EM'~Vl/2)where EM is molar extinction coefficient at the maximaof the shifted band and t::..Vl/2 is the width in crrr+of the shifted band at half intensity. This is anapproximate method for calculating the oscillatorstrength, and considerable uncertainty is likelyto be involved in its measurement in view of thedifficulty in resolving the bands. The transitionmoments (D) were calculated from the relation

[

A J1/"D = 0.0958 EM'~:1/2 -

where EM and t::..Vl/2 are same as above. v", is thefrequency of the shifted band maxima in crn ".

Results and DiscussionBaba and Suzuki- obtained for phenol-dioxane

system an equilibrium constant of 16·4 litre mole?and enthalpy 5·5 kcal mole! in isooctane solvent.West and coworkers- obtained a value of 16·3 litremole! and 5·25 kcal mole! respectively in carbontetrachloride solvent employing infrared spectro-scopy. We have repeated the same system incyclohexane as solvent and obtained a value of14·3 litre mole= and 4·8 kcal mole! for K and-t::..Ho respectively. With ethyl acetate as electrondonor we have obtained a value of 11·0 litre mole!

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for K and 4·8 kcal mole! for -t::..Ho in cyclohexanewhich compares well with the values 9·8 litre mole?and 4·2 kcal mole <, obtained by Grarnstad" by IRspectroscopy.

All our studies were confined to the benzenoidabsorption region in the range 250-290 nm. Theabsorption spectra of the system under investi-gation show the presence of isobestic points whichindicate the existence of 1: 1 hydrogen bondedcomplexes between OH group of phenols and thelone pair of electrons on nitrogen atom. Similarspectral changes were observed for the interactionof all the other phenol'> with tri-isopropylamine,trion-butylamine and tribenzylarnine. For the cal-culation of thermodynamic data, the perturbationson the peak around 290 nm was chosen for con-venience. The results are given in Table 1.

As can be seen from Table 1 the equilibriumconstants for the hydrogen bond formation betweenvarious phenols and tertiary amines vary in therange 305-4·8 litre mole! and the enthalpies varyin the range 9·0-1·7 kcal mole ", The entropies,_t::..So are in the range 5·8-23·8 e.u. Since thetemperatures at which our measurements havebeen carried out are a little different from thosereported in literature, a strict comparison is notvalid. However, the trend is in the expected di-rection. Notwithstanding these limitations, elec-tronic absorption spectroscopy is as reliable asinfrared or NMR spectroscopy in studying hydrogenbond equilibria. With p-nitrophenol the equi-librium constants are quite large. This is becauseof the high acidity of p-nitrophenol which makesit a better electron acceptor (proton donor) thanphenol. The equilibrium constants with tribromo-phenol are, however, intermediate. Tribromophenolis intramolecularly hydrogen bonded and anyelectron donor has to break its intramolecularhydrogen bond before forming intermolecular hy-drogen bond. For phenol the equilibrium constantwith amines vary in the order triethylamineo-tri-isopropylamine > tri -n- butylamine> tribenzylamine,cons istent with the bulk of the alkyl group. Forp-nitrophenol and tribromophenol, the equilibriumconstants vary in the order triethylamineo-triiso-propylamine>tri-n-butylamine. The _t::..Ho valuesfor the hydrogen bonds formed by triethylamineand trion-butylamine with p-nitrophenol are foundto be lower than the corresponding values for phenol.This is rather surprising and the reasons are notimmediatelv clear. One possible reason could bethe formation of the anion of p-nitrophenol, butwe could not find any experimental evidence insupport of this. The equilibrium constants andenthalpy values obtained presently suggest thatthe hydrogen bonds formed in the present systemsare quite strong. Tertiary amines were indeedfound to be strongest donors towards iodine asacceptor", This similarity probably indicates th~tcharge transfer forces play an important role III

phenol-amine interactions. Considerable theore-tical studies using semi-empirical quantum mecha-nical methods have shown that charge transfer isquite important in hydrogen bonding interactionss-".

Spectra of hydrogen bonded phenols - The spectralparameters such as the hydrogen bond shift, t::..v

280

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INDIAN J. CHEM., VOL. 14A, FEBRUARY 1976

TABLE 1 -- THERMODYNAMICDATA ON PHENOL-DoNOR SYSTEM IN CYCLOHEXANE

Electron donor Concentration K (litre mole'<) at 290 nm -t:.Horange [kcal mole'<)(M) Present Lit.

Present Lit.PHENOL

Dioxane 2 x 10-' to 4 X 10-1 14·3 (28°) 16·3 (25O)a 4'8 5·3Ethyl acetate 2xl0-2 to 2xl0-' 11·0 (32°) 9·8 (200)b 4·8 4·2Triethylamine 1 xl0-a to 1 xl0-' 52·4. (28°) 58·0 (25°)c 8·0 9·1TriisopropyJamine 2 X to-a to 2 X 10-1 30·0 (30°) 22·4 (200)d 3·8 5·9Tri-a-butylamine 2 x 10-' to 2 X 10-' 13·5 (30°) 29·2 (200)d 6·8 6·9Triben zylamine 1 x 10-' to 7xl0-1 4-8 (30°) 2.7 (200)d 1·7 1·6

p-NITROPHENOLDioxane 6xl0-1 to 2 X 10-1 16-6 (30°) 8·6Triethylamine 1 xl0-1 to 4 X 10-1 305'0 (25°) 2·0Triisopropylamine 5 X 10-1 to 2 X 10-1 113·0 (270) 5·8Tri-a-butylamine 2xl0-s to 2xl0-1 112·0 (30°) 2·5

TRIBROMOPHENOLDioxane 6xtO-. to 2xl0-1 12'5 (320) 7·3Triethylamine 1 xl0-1 to 4xl0-2 124·0 (26°) 4·0Triisopropylamine 5 X 10-' to 2 x 10-1 21·3 (26°) 5·2Tri-a-butylamine 2xl0-a to 2xl0-' 20'0 (30°) 9'0

·Average value for different wavelengths: 51'8 (275 nm); 49·9 (276 nm); 54-8 (280 nm); 50·0 (281·5 nm); 55·2(283 nm); 52·5 (285 nm).

(a) Ref. 5, (b) ref. 6, (e) ref. 3, (d) ref. 10.

TABLE 2 - SPECTROSCOPICDATA ON PHENOL-DoNOR SYSTEMSIN CYCLOHEXANE

Electron donor Av £M itree IH·bonded IH·bonded D(cnr't) (litre mole-l em-I) ---- (Debye)

Ifre.PHENOL

Dioxane 220 1221 4·4 X 10-1 5·8 X 10-s 1·32 0'59Ethyl acetate 110 1248 5·1 x 10's 1-15 0·55Triethylamine 480 1439 5·7 X 10-' 1·30 0'59Triisopropylamine 640 1074 4·6 x 10-' 1·05 0·53Tri-a-butylamine 640 1252 5·3 x 10-' 1·20 0'56

P-NJTROPHENOLDioxane 560 9995 1·8 X to-I l'8xl0-1 1-01 3-70Triethylamine 2130 11600 3·3 X 10-1 1'80 4·50Triisopropylamine 1650 13300 2·6 X 10-1 1·46 4-12Tri-a-butylamine 1170 11350 1·9 X to-I 1-06 2-82

TRIBROMOPHENOLDioxane 880 1482 1·5 X 10-1 i-s s io-s 1·01 0·96Triethylamine 2310 1340 1·6 X 10-2 1·05 0·96Triisopropylamine 2110 1926 l'9xl0-2 1·27 1-15Tri-a-butylamine 2020 1743 1·9 x 10-' 1'25 1-15

(cm+), mol extinction coefficient, £M, half band-width (AVI/2) and oscillator strength of free andhydrogen bonded peaks (j) and transition moments(D) are given in Table 2. On hydrogen bond for-mation all the bands show a red shift. The mostpronounced red shifts were found only in the caseof 290 nm band. The red shifts are quite largewhich indicates the aliphatic tertiary amines arestrong electron donors. A closer examination ofthe spectra reveals that there is a broadening ofthe band around 290 nm on hydrogen bond for-mation. This is probably due to the mixing ofthe ground state of the hydrogen bonded complexwith different electronic states resulting from thecharge transfer from the electron donor. Such amixing is expected to modify considerably thepotential energy surfaces of the ground and excitedstates and thus broaden the bands in the spectra.This seems to be a satisfactory explanation in thelight of ab initio and semi-empirical quantum

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mechanical calculation on hydrogen bonding carriedout in recent yearst-". It is also significant thatthe oscillator strengths of various phenols increaseupon hydrogen bond formation (Table 2).

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Faraday ssi., 62 (1966), 1056.4' Technique of organic chemistry, Vol. VII, Organic solvents,

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6. GRAMSTAD,T., Spectrochlm. Acta, 19 (1963),497.7. RAO, C. N. R., BHAT, S. N. & DWIVEDI, P. C., Appl.

spectrosc . Rev., 5 (1972), 1.8. MURTHY, A. S. N. & RAO, C. N. R., J. molec. Structure,

6 (1970), 253.9. KOLLMAN,P. A. & ALLEN, L. C., Chern. Rev., 72 (1972),

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