The effect of temperature on ultraviolet absorption spectra and its relation to hydrogen bonding

JOURNAL

OF

MOLECULAR

4, 106-124 (1960)

SPECTROSCOPY

The Effect of Temperature and

on Ultraviolet

Its Relation

to Hydrogen

MITSUO Department

of Chemistry,

Faculty

of Science,

Absorption

Spectra

Bonding

ITO Kyushu

University,

Fukuoka,

Japan

The temperature effect on the absorption spectra of several compounds was studied over a wide range of temperature. It was found that all the absorption spectra characterized by P-K* transition exhibit a large red shift and an increase in total absorption intensity with decreasing temperature when hydrogen bonds are formed between the solute and solvent molecules. The substances which showed this type of temperature effect were phenol, aniline, and diphenylamine. The absorption band of acetone at 280 mp, which is characterized by nlr* transition, showed a different temperature effect. Its total absorption intensity decreased progressively with decreasing temperature even when no special interaction existed between acetone and solvent molecules. When hydrogen bonds were formed between these molecules, a blue shift in the absorption maximum is observed in addition to a decrease in the total absorption intensity. The effect of t,emperature was also studied on the n-T* absorption spectra of pyrazine and crotonaldehyde. INTRODUCTION It, is well known that the spectral features of ultraviolet solution spectra change with temperature. As causes of temperature-dependent change in the absorption spectra, the following three important factors may be considered: temperature effect on a refractive index of a solvent, temperature effect on the Boltzmann distribution among vibrational and rotational energy levels of a solute molecule, and temperature-dependent, interactions among molecules in the system. Contributions of the first and the second fact,ors to the spectral change were studied fairly well (I-J), while the contribution of the third factor is not yet well recognized. Among many kinds of temperature-dependent interactions, an interaction due to hydrogen bonding is the one most often encountered in usual absorption experiments, and our attention has been concentrated in this interaction. Since the formation of the hydrogen bond proceeds with decreasing temperature, the effect of the hydrogen bonding on absorption spectra is expected to be most easily recognized by measurements of the absorption spectra over a wide range of temperature. 106

TEMPERATURE

EFFECT

107

In this paper, temperature dependence of the absorption spect,m of sr~rr:~l aronlatic cornp~)uIids ~ontai~~ing proton donating and accepting groups was studied in various solvents over a wide range of temper&n-e and the results obtained were interpreted in term of hydrogen bonding.

Aln apparatus for quantit,ative measuremeMs of ultraviolet absorption spectr:l at low temperatures was con&ucted after t)he model reported by Bevle and 130~ (G), but, several modifications were made. It was essent~ially a metal &w:tr vessel with absorptiox~ cells suspended in a vacuum chamber. The details of t.hr apparatus are shown in I’ig. 1. It consists of absorption cells, an inner VCSSPI, and an out,er vessel. A cross section of the absorption cells is also shown in the figure. il. re~~tal~g~~l~~1 brass block A had two cylindrical holes C which were screw-threaded inside. Cylindrical brass tubes I threaded outside were screwed into the holes C and soldered to the brass block A at the position of equal projection on each side. The ends of I must be ax plane as possible, otherwise a sample solution might leak when t,he cells were assembled. Holes J drilled obliquely through the edge of the block A met at the tops of I, so that, bubbles formed by contraction of the sample on cooling process could escape into J. Outer side tubes K, which

/ 10 cm

,

,

5cm

OF CELL UNIT

PART SECTION & ELEVATION OF APPARATUS

FIG. 1 Low temperature

I

SIDE ELEVATION

ahsorption

cell

108

MITSUO

IT0

had small screw-in stoppers L carrying “Teflon” gaskets, were inserted into J and soldered. Circular quartz windows E closing the absorption cells had “Teflon” gasket on each face of them and were held in position by threaded ring-caps D screwed on to I. The absorption cells thus constructed were separable from the imler vessel M by mean of threaded cylindrical projection B for sample changing or for cleaning. The inner vessel M, in which the refrigerant was stored, was soldered at the top of it to a thin iron tube N of 1 mm in thickness, which acts as a thermal insulator. An annular iron ring 0 was soldered to N and to an iron tube P as shown in Fig. 1. The inner vessel carried also a brass flange Q, which was held by four sets of bolts and nuts R, and a rubber gasket was placed between P and Q. The outer cylindrical brass vessel S was equipped with four quartz windows which were aligned to the axes of the absorption cells. Connection to a vacuum pump was made through a side arm T, Cell temperature was measured at the midpoint between the two absorption cells by means of a thermocouple and galvanometer. First, solution and solvent to be investigated were introduced through J into the upper and the lower absorption cells, respectively, and sealed with the stoppers L. It is necessary to fill the cells with an excess of the solution and the solvent to allow for the great contraction of volume on cooling. The absorption cells filled with the solution and the solvent were screwed into the inner vessel, and then suspended in the outer vessel. After thus assembling the apparatus it was evacuated. About 0.1 mm Hg pressure in the Dewar vessel was sufficient to prevent condensation of moisture on the outer vessel, although pressures as high as I mm was still good enough. After evacuation the inner vessel was filled with refrigerant. When it was necessary to keep the cells at low temperature for a long time, the quartz windows of the outer vessel were warmed electrically with Xchrome wire wound around them in order to prevent frost formation. Since the apparatus described above had a great heat capacity and fluctuation in temperature could be suppressed within a small range, it could be favorably used in combination with a point-by-point reading spectrophotometer such as Beckman model DU or analogous quartz spectrophotometer, which usually requires a long time for the absorption measurements. For the present investigat,ion a Shimadzu model &B-50 quartz spectrophotometer was available, so that the absorption apparatus was designed for the easy adaptation to it. Figure 2 shows a side view of the arrangement. A is the apparatus, which rests on a base B. B can be moved vertically by means of sliding rod C, guide tube D and lever handle E, range of the vertical motion of B was limited by a slot hole and a pin.1 When E was down, the lower absorption cell was in the light beam and when E 1This idea was suggested by Dr. Y. Kanda of this laboratory.

T~~~~~RATURE

EFFECT

to9

Although this apparatus could be operated for the temperature range between + 1.00” C and - 190” C, it was used in the present, investigation only within the rsng~ in which solvents used remained fluid. Solvents used were ethanol, ethyl &her, and n-hexane. They remwixxt:d licluid in temperature ranges between +70” and - 110”, f30” and - 1 IO*, and +GO” and - 100” C, respectivcly. Temperatures higher t,han the room t2emperature were obtained by circulating

fSvcr the t,emperature range invckigated reproducibility of optical density valuers was very good when ethanol or et,hyl ether was used a~ a xtrlvent, while it, WURrather poor wkh n-hexane at low temperature, especially below -70” c‘. 1x1 the lat,ter case, t.he varianccr between successive optical den&y values at, - 70” C was less than 0.01 optical density units in the spectral region Rbove 2300 A, whiie it increased to about 0.X units at 3100 ,4. Therefore, molecular extinction eoefEcients for t.he tz-hexsne sdution at low temperature is believed t.o he correct wit.hin a small error in the spectra.1 region &ove E&l0 A, 1~12:itt the W-

110

MITSUO IT0

gion below it a great confidence cannot be given to the data obtained. Although a similar phenomenon was also reported by Potts (6), the reason for this has never been understood well. Lengths of the absorption cells (about 3 cm) were determined accurately by comparing the optical density of a standard solution with that measured with a cell of known length for the same concentration and temperature. Concentrations of solution at various temperatures were corrected by use of contraction or expansion data for the solvent used (7). In the absorption measurements, spectrometer slit width was kept constant for the spectra to be compared, i.e., for one solution in a series of temperatures. n-Hexane used as solvent was repeatedly shaken with concentrated sulfuric acid and washed with water, sodium hydroxide, and water. After it was passed through a silica gel column, it was refluxed over sodium and fractionated. Ethanol of a commercial grade containing a trace of benzene was refluxed over calcium oxide and distilled, and it was refluxed with silver nitrate and potassium hydroxide, and then fractionated. Ethyl ether was dried over sodium sulphate and distilled. The compounds studied here were the best obtainable commercially and. when necessary, they were further purified by recrystallization or distillat,ion. RESULT AND DISCUSSION GENERAL EFFECT OF TEMPERATURE It is generally accepted that an ultraviolet solution spectrum is temperature dependent even in the case where no special interaction, such as hydrogen bonding, exists in the system. This is illustrated with the near ultraviolet absorption spectrum of chlorobenzene in n-hexane (Fig. 3). A decrease in temperature leads to an increase in intensity of band maxima and a decrease in that of band minima, that is, individual vibrational bands increase in sharpness with decreasing temperature. It is also seen that positions of the band maxima do not shift very much with decreasing temperature but tend to move a very little towards the longer wavelength. A similar temperature dependence has been reported by several authors for the solution spectra of several aromatic compounds (1-J) and for the vapor spectra of diatomic molecules (8, 9). Sulzer and Wieland (10) have recently observed the absorptions of Cl2 , Brz , and 12 vapors over the wide range of temperature, and have found sharpening as well as red shifts of maxima of the ultraviolet absorption bands of these molecules by decreasing the temperature. Moreover, they noticed the fact that the total absorption intensit,y, i.e., area below the absorption curve, was nearly independent of temperature. This latter finding seems to apply also for the solution spectra, as illustrated in Fig. 3 and in Figs. 1 and 2 of Ref. 3.

T~~~~RATURE

EFFECT

Ill

&

0

36

-3-k”--

Frc. 3. The near ultraviolet absorption spectrum of chlorobenxene in 15q --- -I - 70°C.

__)

Thus, the fo5lowing three criteritt, may he given as general effect of temperature on solution spectra.: A. Ba,nd shmpness incremes with decreasing temperttfwet B. Position of absorption maximum does not move or moves very litt.Ie towards the longer wavelength side, with decreasing temperature, C, The t,otatl absorption int’ensity is approximately independent of the temperature. As these empirical criteria have been fully discussed already by Grubb and Kistiakowsky (3), and Sulzer and Wieland (IO), further explanations will not be given here. With these criteria the temperature effect on the absorpt,ion spect’ra for the systems involving hydrogen bonding will be described in the following section.

112

MITSUO TEMPERATURE

ITO

EFFECT FOR SYSTEMS HYDROGEN BONDING

INVOLVING

1. PHENOL Temperature effect on the absorption band of phenol near 270 rnp was first studied by Coggeshall and Lang (II), who observed the absorption spectrum of phenol at various temperatures within the range between 24” and 75” C for the solution of about 1O-4M in a solvent containing 0.2 M ethanol in iso-octane. However, they failed to find a significant change in the absorption spectrum with variation of temperature. For the solution of 1.5 X lop4 M in a solvent comprising 2 X 10e2 M ethanol in n-hexane, the experiments of temperature variation were repeated here over a wider range of temperature, and the results obtained are shown in Fig. 4. A

h 241

260 ,

I

(m(l) 270 I

FIG. 4. Effect of temperature on the absorption taining 2 X 10-Z M ethanol in n-hexane. Key: --, -70°C.

I

260 I

spectrum of phenol in the solution 10°C; ----, - 10°C; ~.-.-. , -40°C;

con-

. .,

remarkable change in t,he absorption spectrum was observed which was overlooked by Coggeshall and Lang. At IO” C, the spectral fe&ure was complet,ely identical with that of the pure n-hexane solution, but it gradually varied with decreasing temperature. The band maxima shifted to the red by about 250 cm-’ with decreasing temperat~lre from 10” C to -70” C. ~u~her~n~~r~, at intermediate temperatures, the absorpt,ion exhibited shoulders near at. 274 and 271 ml*. On the other hand, such a remarkable change could not be observed for the spect,run~ of the pure n-hexane solution, where the spectral shift i~~du~ed w&h decreasing t,~~rn~rature ws,s quite small and was less than 100 em-1 in the region between 60” C and -?O” C (Fig. 5). This indicates that t
241

280 ,

I

I

36

I

I

260

I

t

I

1

38

1

FIG. 5. The absorption spectrum of phenol in n-bexane at various temperatures

114

MITSUO

IT0

The following three forms may be assumed for the hydrogen-bonded of phenol and ethanol (I-III).

complex

C2b I AH “‘;a/

In I and II a phenol molecule acts as a proton donor and a proton acceptor, respectively, in III both characters are shown by the molecule. The hydrogenbonded complex analogous to I can be materialized in the complex formed between phenol and ethyl ether. According to the investigation of Xagakura and Baba (12), the absorption bands of phenol shifted to the red due to the complex formation with ethyl ether. This fact and the observed red shift with decreasing temperature enable us to expect that the hydrogen-bonded complex between phenol and ethanol has the form I. On the other hand, if the complex was of the form II or III, the blue shift of the absorption band, contrary to the actual observation, is expected from the reasons which will be described in the subsequent paper (13). The greater possibility of I than that of II or III implies that proton donating power of phenol is much greater than that of ethanol. 2. ANILINE The temperature effect on the absorption band of aniline near 280 rnp was studied in n-hexane and in ethyl ether. The results obtained are given in Figs. 6 and 7. For the n-hexane solution the observed temperature effects are readily classified by the general effects since all the criteria listed in the previous section are satisfied. On the other hand, the absorption band of aniline in ethyl ether exhibits a large red shift of about 500 cm-’ with decreasing temperature from 15” C to -70” C and an increase in the total absorption intensity. These deviations from the criteria for the general effects suggest that there is a considerable contribution of hydrogen bonding to the spectral change observed with the ethyl ether solution.

MITSUO

116

IT0

The hydrogen-bonded complex of aniline and ethyl ether possibly takes the two forms shown (IV, V).

and

V IV is the 1: 1 complex and V the 1: 2 complex. The spectral shift induced by formation of V may be larger than that of IV since the number of hydrogen bonds is larger for the former. The observed large spectral shift by lowering the temperature seems to indicate that in addition to the 1: 1 complex the 1:2 complex is also formed at lower temperatures. The existence of such a complex was suggested by Vidale and Taylor (14) in the hydrogen-bonded complex of a molecule having several H-X groups and ethyl ether. 3. DIPHENYLAMINE Temperature dependence of the absorption band of diphenylamine was observed for ethanol and n-hexane solutions. The results obtained are shown in Figs. 8 and 9. In the case of the ethanol solution the absorption maximum shifted to the red by about 400 cm-l with temperature decrease from 10” C to -70” C. This amount of the spectral shift was about three times as large as the corresponding shift for the n-hexane solution (120 cm-l), indicating again the formation of hydrogen-bonded complex of diphenylamine and ethanol. It is possible for the complex to take the two forms shown (VI, VII).

and

/

TEMPERATURE

EFFECT

117

MITSUO ITO

1.18

If’ the complex took the form VII, it would be expected that the corresponding absorption band of N-methyldiphenylamine also exhibits a large spectral shift in the ethanol solution since this molecule has a possibility to form a complex analogous to VII. However, in the actual observation (Fig. lo} only a small shift was observed. This shows a greater possibility of VI than that of VII for the h~drogcn-bonded complex of diphenylamine and ethanol. It is to be observed in Figs. 8 and 9 that the total absorption intensity of diphe~ylamine remar~bly increases with decreasing temperature. In the ease of the ethanol solution, this increase in the total absorption intensity may be interpreted as a result of the hydrogen bond formation between the solute and solvent molecules as mentioned above. However, the same explanation seems to be dishcult) to apply for the n-hexane solution, where the solute-solvent interaction is generally small. The fact that the corresponding absorption band of N-methyldiphenylamine in n-hexane did not exhibit any remarkable change in the total

1

32

I

I

34

J

I

36

I

I

36

Y x I o-3hli’)

FIG. 10. Absorption spectrum of N-methyldiphenyl&mine in ethanol

m

2 p

*3

120

MITSUO

IT0

affected by temperature change, while, in the ethanol solution, it shifted to the blue with decreasing temperature as shown in Fig. 13. This blue shift may be explained as a result of hydrogen bond formation with ethanol at low temperatures, and this interpretation is quite in agreement with the fact that hydrogen bonding has the effect of causing a blue shift for transitions designated as n-?r* (16). It is seen in Figs. 12 and 13 that total absorption intensity gradually decreases with decreasing temperature. This type of temperature effect is a type different from that encountered in the previous subsections. According to the theoretical considerations made by Holmes and McClure (17), this decrease in the total absorption intensity implies that the electronic transition associated with the absorption is forbidden by symmetry. If one assumes that the symmetry of the acetone molecule is CzV like formaldehyde, the forbidden transition must be an AZ + A1 transition, which is in agreement with the assignment for the corresponding absorption of formaldehyde (18). A(mp) 320

300 I

I

280 I

I

260 I

240 I

I

15 -

+15” -25” -85”

I

I

38

FIG. 12. Absorption

spectrum

of acetone

I

I

I

40

in n-hexane

TEMPERATURE

121

EFFECT

X (mv) 320

300

280 I

260

I

FIG.13.Ahorptionspectrum. of acetone in ethanol

Figure 14 shows the absorption band at 320 mi.t of pyrazine, which is known as due tan the ?1--?r*tral~s~t~i(~~~(19). It is noted here that at 15” G the absorption spectrum of pyrazine in an ethanol solution exhibited only a sin& broad band with no vibrational structure, whereas at -RI!” G several inflections occurred in the longer wavelength region of the band maximum. Moreover, t>heir spectral positions were completely identical with those of the vibrational bands shown in the spectrum of the n-hexane solution. A similar phenomenon was observed also for the 320 rnp band of crotonaldehyde as shown in Fig. 15. Bresley and Kasha (20) have recently found in their careful measurements of the absorption band at 3:30 r~ of p~r~da~~ne that the vibrational bands of the 3t-hexane solution were preserved afso in the spectrum of the solution containing a large amount of ethanol in spite of a considerable blue shift of the absorpt.ion maximum. But., under t.heir ex~eril~~~~ta~ ~ond~t~o~~s at 25” C, the ~reser~~~,~on

123

B A

t

I

26 FIG. 15, ~b~~rp~~ion spectrum ethanol at - 70°C; C, in n-hesane

I

I

t 3t, t/ x lo-3hK’~

i

28

of ~~otonaldeh~de, at - 70°C.

TABLE

Key:

1 s k, in ethanol

15°C; B,

I

SPECTR.M, SHIFTS AND CHANGES IN THE TOTAL ABSORPTION IXTENS~TY WETH DECREAEX-JG TEMP~~AT~~~~

AdO* and

Compt’ulads

I.- .--.

-TO”)

..-~--_-.-”

Ch~~r~ben~~~~~ N-*-lethgldiph~nplamine wlenni Aniline

1Xphen$amin~ Act?t~one * Thr

values

obt.ained

in 5x1 ethyl

in cm-~

..,----.. __---

ji-w&l f(io”t

- ____. .._

-80

1

-120

1

-30

1

-40

1

- 120

1.2

0

0.9

ether solution.

-120 -200 - 4:w -370 1280

1

--. >l”

1.2 0.9

in

MITSUC)

121

ITO

the second and the fourth columns of the t~able, and the ratios of the t&d ahsorption intensities at these tsvo ~e~~peratur~§ are listed in t,he t,hird and the fifth columns. In this table the following characteristics may he found: 1. The magnitude of t,he spectral shift5 is usually small in the n-hexane soluCon, but it is large in the ethanol (or in et,hgl &her) solution espeeia~ll~~ for the substances involving the proton accepting or donating groups. 2. M the compounds, except acetone, exhihit the red shifb of the ahsorption bands with decreasing the temperature and these absorption bands are 911 characterized as 7~-?r* transitions. The absorption hand of acetone, which is due to an n-a* transition, shifts to the blue in the ethanol Aut~ion. 3. The ratio of the total shsorption inten&y is approximately eclual to I when hydrogen bonding is absent, from tShe system, hut. it is greater t,han 1 when it, assists. The small value for acetone may he interpret.ed :ts due to other causes as menGoned in the previous section.

The author wishes to express his sincere thanks to Professor for their constttnt ndviw and encouragement.

S. Imanishi

and t,a Dr. Y.

Em&

REFEREKCES 1. S. BRCWERSEN AND A. LANGSETH, Kyl. Dan&e Videnskab. S&k&, Xaf .-fys. Xedd. 26, 1 (1951). 2. R. PAs~l%~r;~ ARD I.G.It0SS,d. f%&'IF&. f%~&~, 1012 (19j-t). 3. W. T. GRIM AND G. B. X~ISTXAK~~~~~~,J.A~. Cf2em. Suc.72,-II9 ff950). 4. v. A.-k‘ARBOROUGR,J. E. &~SIUX, AND 1v.J. LAIMBDXN,d%%l. 6%5nr.86,15% 5. R. S. BEALE AR'D E. M, F. ROE, J. Sei. In&r. 28, IOQ (1951). 8. W. J. POTTS, J. Chew Phys. 21, 191 (1953).

7. “International 8. G. IL GIBSON,

(z?Mf.

Critical ‘t’ahles,” Vol. III. McGraw-Hill, Xew \iork, 19%. K. Rrw, AND Xi.6. BAYLISS, Phys. fZev.44,193 (1933).

0.

HERCZW AND IL W~~~~ND, H&J. Pkys. .teta 21, 936 (19-N); 22, 552 (1949); S~JLZER AND li. WIBXAN~, Heb. Phys. Acta 2% 591 (1949). SULZER AND K WIELMU, Hetv. Ph,ys. Aetu 26, 653 (1952). U. C~GFESHALL AM> E. M. LANG, J. Am. (‘hem. Xoc. 70,X&3 (KM). ii?. S. NAWKTIRA AND H. BABA, J. Arrr-. Chem. Sot. 74,569s (1952:.

9. A. P. IO. P. 11. N.

IS. &I. 1~0, this issue 13. Mol. ~~~(:~~Qs~~~~4, f25 (1QSQ)j. 1.4. G. IA VruaLE ANDR. C. TAYLOR, J. Aw,* Chem. see. 78,2Q4 (1956). 15. A. AFNARA, J. Chew& &c. Japan, P~tre &kf.. &xt. %%, -t:?$ (19%). Ifi.

17. 18. IS. 20.

M.

I(ABEIA,DZ‘~C. Farada!/ Sot. No. 9,14 (19.50); H. MCCONNELL, J. Clhem. Phys. 20, 700 (1952). 0. G. HOLMES AND I).8. MCCLURE, J. Chem. Phys. 26, 1686 (1957). T. ANKO AND A. SADO, J. (‘hem. Phys. 26, 1759 (1957). B’I. ITO, R.. SHIMADA, T. KCRMSHI, AND W. IMIZUSHIMA, J. fhetll. Phys. G. J. BREAI‘EY AND M. KMHA, f, r2nt. C%WL. Sot. 77, 4462 (19551.

26,

1508 (lQBT\.