780
The Journal of Physical Chemistry, Vol. 82, No. 7, 1978
T. Sakurai and A. Takahashi
Behavior of Ruthenium in Fluoride-Volatility Processes. 3. Thermal Decomposition of RuOF, Tsutomu Sakural" and Akira Takahashi Fluorine Chemistry Laboratory, Japan Atomic Energy Research Institute, Tokai-mura, Naka-gun, Ibaraki-ken (3 19- 1 I), Japan (Recelved Aprll 25, 1977; Revised Manuscript Received November 28, 1977) Publication costs assisted by the Japan Atomic Energy Research Institute
Solid RuOF4 releases its oxygen even at room temperature. The rate of this thermal decomposition, expressed in mol of O,/"C, reaches a maximum at about 70 "C and then decreases markedly with temperature; RuOF4 is an unstable material contrary to the description of earlier literature. The vapor pressure of RuOF, was checked by transpiration method; it is expressed approximately by the following equation in a temperature range 0-70 "C: In P,, = 8.48 - 2.74 x 103(l/T). Since the bulk of solid RuOF4 decomposes at about 70 "C, the vapor pressure measurement is meaningless over 70 "C. Infrared analysis of the solid indicates that the Ru=O bonds change into single bonds upon condensation; this may be related to the instability of solid RuOF4.
Introduction Although the product of Ru02-F2 reaction had been suggested to be RuF, by many other workers,1,2we reported in the previous work that it was RuOF,.~ The heaviest species in the mass spectrum of the product was RuOF4+;chemical analysis gave a value 4.0 f 0.1 for the F / R u mole ratio in the product, and mass spectrometry confirmed the release of oxygen from the product. The papers so far published on RuOF, are very few; the only paper may be by Holloway and Peacock. They reported in 1963, that they had obtained RuOF, by the reaction of ruthenium metal with a mixture of bromine trifluoride and bromine a t 25 "C. Its magnetic moment and vapor pressure data were presented; they used a static method in measuring the vapor pressure. According t,o our experiments, however, RuOF, is a rather unstable material; it releases oxygen even at room t e m p e r a t ~ r e .Therefore, ~ it appears that the static method is not suitable for measuring RuOF4 vapor pressure, therefore the earlier data are obscure. On the other hand, the property of releasing oxygen seems to be peculiar to RuOF,, because no such data have been reported for other oxide tetrafluorides such as MoOF,, WOF,, and ReOF4.5 The question was raised, why RuOF, was so unstable. These uncertainties on the properties of RuOF, urged us to study its properties in further detail. In the present work, the temperature dependence of the thermal decomposition is studied with attention to the oxygen quantity released from the solid, the vapor pressure is measured by a transpiration method, and the instability is discussed on the basis of the infrared spectra obtained. Experimental Section Materials. Commercial grade Ru02.rzH20 of purity higher than 99.9% was heated in air a t 700 "C; the x-ray diffraction pattern of the resulting substance agreed with that of R U O ~ Infrared .~ analysis showed no presence of water in the substance. Its specific surface area, determined by the BET method using krypton, was 6.3 m2/g. Fluorine from Matheson Co. was passed through a chemical trap charged with sodium fluoride for removal of hydrogen fluoride. Helium of purity higher than 99.999% served as the carrier of fluorine and RuOF4, after passage through a cold trap at -196 O C . Apparatus and Procedure. The apparatuses used were made of corrosion-resistent materials such as Monel, nickel, 0022-3654/78/2082-0780$0 1.OO/O
copper, and Teflon. Monel valves of Autoclave Engineers and Hoke Inc. were used for the connecting lines through which fluorine and RuOF, were passed. Leakage of the apparatuses was checked with a helium leak detector, which was kept below atm cm3/s, RuOF, was prepared by the reaction of RuOz with f l ~ o r i n e .RuOz ~ (1-2 g) was placed in a "boat" reactor, heated under high vacuum to remove moisture, and then fluorinated with 20% F, in a He stream a t 300-400 "C. The volatile products from the reactor were collected in a Monel container cooled to -196 "C. At the end of the fluorination, the container was evacuated a t 0 "C by pumping to remove HF and trace amount of R u 0 4 (byproduct). The container was then separated from the fluorinating apparatus and positioned in another apparatus shown in Figure 1 for the measurements of the vapor pressure and thermal decomposition of RuOF,. The transpiration method was employed for vapor pressure measurement in order to avoid the influence of the oxygen released from RuOF,. The RuOF, container was thermostated at an operating temperature between 0 and 100 "C and then helium was allowed to pass through it at a constant flow rate between 2 and 25 cm3/min for a specific period. The effluent gas from the container was passed through a connecting line kept at a temperature slightly higher than that of the container and then introduced into a trap cooled to -196 "C, where the RuOF, vapor in the gas stream condensed. The temperature of the container was kept constant to within h0.2 "C and the flow rate of helium was regulated with bellows valves from Nuclear Products Co. with an error smaller than f0.2 cm3/min. A mass flow meter of Hasting Raydist a Teledyne Co. Model LF-20 measured the flow rate of He. After completion of each run, the trap was evacuated, warmed to room temperature, and then allowed to introduce about 60 mL of 2 N NaOH (containing 1.5% K2S208)through a Teflon stopcock with the aid of vacuum. Immediately after separation of the trap from the apparatus, a small quantity of 2 N NaOH was admitted to the sidearm of the trap for collecting the RuOF, deposited therein. Chemical analysis of ruthenium and fluorine followed this hydrolysis of the RuOF4 collected. The ALC photometric method was used for the determination of fluorine7 and spectrophotometry with 8-quinolinol was used for the determination of ruthenium.8 It was experienced in the previous3 and present works that on hy@ 1978 American Chemical Society
The Journal of Physical Chemistry, Vol. 82, No. 7, 1978 781
Thermal Decomposition of RuOF, 2N- NaOH (containing Mass flow meter
\
T? air
Trap for RuOF4 vapor
L
- - - _-
2
Thermosta t \Container
of
solid R u O F ~
Figure 1. The apparatus used for measuring the vapor pressure and thermal decomposition of RuOF,. The RuOF, container was thermostated at the desired temperature and He was allowed to pass through it at a constant flow rate. RuOF, vapor in the He stream was collected in the cold trap and, at the end of the run, it was hydrolyzed with a NaOH solution for calculation of its partial pressure in the gas stream. The effluent gas from the trap was introduced into a gas chromatograph for the determination of the oxygen liberated from RuOF, solid.
drolysis about 6% of the ruthenium unavoidably precipitated as hydroxide; so a step of dissolving it was necessary before the determination of ruthenium. This operation increased greatly the complexity of the analysis of ruthenium. In the present work, ruthenium analysis was made with some typical samples only to confirm that the fluorine-to-ruthenium mole ratio in solution was close to the theoretical value (= 4). The RuOF, partial pressure in the helium stream was calculated from the analytical values of fluorine. The apparatus in Figure 1 was also used in measuring the quantity of the oxygen produced by the thermal decomposition of RuOF,. In this case, the effluent gas from the trap was fed to Hitachi Model 063 gas chromatograph. The flow rate of helium was kept constant a t 20 cm3/min according to specifications of the instrument. A t a temperature of the container between room temperature and 400 "C, helium was allowed to passed through the system until oxygen was no longer detectable in the effluent gas from the trap. The helium flow was then turned off, and temperature of the container was raised for the next higher temperature measurement. Measurements were thus made stepwise: the temperature was raised by 5 or 10 "C in each step up to 100 "C, and then 100 "C in each step beyond 100 O C . Finally, chemical analysis of the residue in the container was carried out in the same manner as in the previous worka3 The infrared absorption spectrum of solid RuOF, was obtained in a manner similar to that in the previous work.3 An absorption cell flanged a t one end was connected directly to the "boat" reactor so as to enable gaseous RuOF, to diffuse smoothly into the cell from the reactor. R u 0 2 (200-300 mg) was placed in the reactor and heated to 300-350 "C, then fluorine was admitted to the system to a pressure of 200 mmHg. Gaseous RuOF, diffusing from the reactor was allowed to condense on the silver chloride windows kept a t room temperature. Infrared spectra were scanned with Japan Spectroscopic Co. Model IRA-2 recording spectrophotometer (4OC-4000 cm-l). The windows were not corroded by the solid; therefore they were used repeatedly after wiping with gauze. Results and Discussion Thermal Decomposition. Mass spectrometry confirmed in the previous work3 that the gas produced by the thermal decomposition was only oxygen; no free fluorine and oxygen fluorides were detected. In the present work, the
Tempernture
in
"C
Figure 2. The amount of the oxygen produced by the thermal decomposition of solid RuOF,. The histogram shows the average per degree of temperature rise, and the broken line shows the integrated amount of oxygen in its ratio to the theoretical value calculated from the initial RuOF, quantity 1.786 g (9.25 X mol).
temperature dependence of the thermal decomposition was studied. The amount of released oxygen depended on both the quantity and temperature of the solid. Moreover, it was found that the solid liberated oxygen only when the temperature was being raised. After the temperature became constant, the liberation of oxygen slowed gradually and ceased within 1-2 h. Further increase in temperature resulted in an additional release of oxygen. This finding means that thermal decomposition takes place partially depending upon temperature. The measurement of oxygen was carried out in each 5, 10, or 100 "C rise in temperature up to 400 "C, and the average per degree of the rise was taken, Le., mol of 02/"C. The results up to 300 "C are shown in Figure 2 as a histogram. The release of oxygen reaches a maximum at about 70 "C and then decreases sharply with temperature; however, a small amount of oxygen is liberated even at 400 "C. The broken line in the figure shows the integrated amount of the oxygen in its ratio to the theoretical value calculated from the initial amount of solid RuOF4. The oxygen quantity amounted to about 80% of the theoretical value in heating to 400 "C. The remaining oxygen (-20%) is partly present in the form of RuOF, in the cold trap and is partly confined still in the residual solid in the container. The value of 80% obtained for the oxygen quantity substantiates the previous conclusion3that the product of the Ru02-F2 reaction is RuOF,. Also, it is evident from the present experiments that solid RuOF, is so unstable as to decompose below 100 "C. Since chemical analysis gave values of 3.8-4.0 for the F / R u mole ratio in the residual solid, it is considered that the thermal decomposition proceeds to produce RuF, and 02: RuOF,(s)
+
RuF,(s)
+
'/2
O,(g)
The residual solid was black and sensitive to moisture. We have not succeeded in x-ray analysis of RuF,. The only paper so far published on RuF4 may be by Holloway and P e a c ~ c k .They ~ prepared RuF, by the reduction of RuF, with I2 and attempted x-ray analysis; however, the cyrstal structure data were not presented. In other experiments, the interaction of gaseous RuOF, with the surface of a Monel tube was checked by the use of radioactive 103R~02.10 The Io3Ru02was obtained by irradiating R u 0 2 in JRR-2 for 180 h in a neutron flux of 4.3 X 10I2 n/cm2 s.ll About 30 mg of it was fluorinated in the same manner as described already, and the resulting
782
The Journal of Physical Chemisfry, Vol. 82, No. 7, 1978
T. Sakurai and A. Takahashi
TABLE I: Dependence of RuOF, Partial Pressure on H e l i u m Flow Rate a t 50 C 50
He f l o w F l u o r i n e quan- RuOF, parrate, e m 3 - D u r a t i o n o f tity in soh- t i a l pressure: ( N T P ) / m i n gas flow, h tion, m g mmHg
9.7 9.7 5.0 19.1
4.0 4.0
5.1
4.0
3.0 2.6
14.4
1.3 2.0
2.5 19.6
8.0 2.0
2.1 2.1
8.0 8.0
4.4 2.8 4.0 3.0 4.1 4.3
\
20-
0.48 0.42
0.55 0.34
0.35 0.75 0.28 0.81 0.82
Calculated from the fluorine q u a n t i t y , H e flow rate, and d u r a t i o n of gas flow, a
gaseous mixture was passed through the tube heated to 100-200 "C. Then, a scintillation detector was scanned along the length of the tube. The decomposition of gaseous RuOF4 in small quantity was observed through the increase of radioactivity. However, once coated with ruthenium, the surface no longer decomposed gaseous RuOF,. In order to observe RuOF,, the gas in the RuOF4 container was expanded into an evacuated Pyrex container cooled to -196 "C. A small amount of yellowish green RuOF, condensed in the Pyrex container. However, it turned black gradually when warmed to 30 "C; the thermal decomposition of solid RuOF4 takes place even in a Pyrex container. Thus, RuOF4 turned out to be a rather unstable material. Vapor Pressure. Since the earlier vapor pressure study used a static method: it may include an error arising from the oxygen released. We tried to check the vapor pressure by means of a transpiration method. The instability of RuOF4 forced us to use the product container of the fluorinating apparatus as the transpiration chamber. This greatly restricted the favorable structure of the transpiration chamber, e.g., the heat ~ h i e l d ' ~was J ~ not available because a small opening became pluged during the fluorination. The present experiments were carried out under such a disadvantageous condition. Bell, Merten, and Tagami have obtained saturation at flow rates between 1and 5 cm3/min.13 So, in the present experiments, the helium flow rate was varied over 2 cm3(NTP)/min and the dependence of RuOF, partial pressure on the helium flow rate was studied a t 0, 30, 50, 70, and 90 "C. Calculations12 made with the diffusion coefficients of various binary systems of helium14suggest that the diffusion effect12is negligible in the present case even a t 2 cm3/min. In measurements at 70 and 90 "C, the solid RuOF, sample was replaced with a fresh one after each run, for fear that the effect of the thermal decomposition at about 70 "C was superimposed on the flow rate dependence in the subsequent run. Table I shows the results of the experiments a t 50 "C. A t each temperature, an increase in the flow rate resulted in a decrease of the partial pressure. These results indicate that saturation is attained at flow rates less than 2 cm3/min. In the present work, however, it was difficult to find the point of saturation in such a small flow rate range as 0-2 cm3/min. Approximate values of RuOF, vapor pressure were obtained in the following way. Kubaschewski and Evans describe15that, theoretically, the saturation of carrier gas should be attained a t a zero flow rate; many workers once obtained vapor pressures by extrapolation to zero flow rate.16 In practice, however, the saturation is usually reached a t low yet finite flow rates; the extrapolation may result in an overestimation of vapor pressures. These facts suggest however that the RuOF, vapor pressure will be in
90
\
0
\ \
\
\,/
I Or
Earller literature
\
0 1 , -
I € 5.0E -
$
Temperature, "C 70 50 3 0
\
\ '\
\
2.0-
v)
0
&
I.0:
9
0.5:
-
0.2_ . I
2.6
2.8
3.0
3.2
3.4
3.6
Z 8
I./"K x io3 Figure 3. Vapor pressure of RuOF,. The full line shows the vapor pressure obtained in the present work by the use of the transpiration method. The thermal decomposition shown in Figure 2 made the measurement difficult over 70 "C. The broken line approximates the vapor pressure reported by the earlier worker^;^ they used static method.
the range between the partial pressure a t the least flow rate ( 2 cm3/min) and the value obtained by extrapolation of the plots of the partial pressure back to zero flow rate. By regarding this range as an uncertainty inherent in the present experiments, we estimated the vapor pressure a t 0, 30, 50, and 70 "C; the results are shown in Figure 3. (Table I gave a value of 0.89 f 0.08 mmHg as the vapor pressure a t 50 "C.) The full line approximating the plots is expressed by the following equation: In P,, = 8.48 - 2.74 X 103(l/T) where T i s temperature in K and P,, the vapor pressure is mmHg. The measurements were meaningless above 70 "C, because the bulk of solid RuOF, decomposed a t about 70 "C and the layer of the resulting RuF4 prevented the subsequent vaporization of the remaining RuOF4. The partial pressure was only 0.68 mmHg a t 90 "C and 2.1 cm3/min. Therefore, the apparent vapor pressure of RuOF, cannot increase beyond 70 "C. In this respect, the earlier data: which are approximated by a broken line in the figure, are erroneous. Although approximate, the present study has revealed the characteristic of RuOF4 vapor pressure. In the present study, it was assumed that there was no interaction between solid RuOF, and RuF4. This assumption is supported by the reproducibility of the partial pressure shown in Table I. The experimental runs in the table were carried out in that order. The obtained partial pressure depends on He flow rate but it is almost independent of the order of the run. This means that RuF4 does not affect RuOF, vapor pressure below 70 "C. Instability of RuOF,. The following two facts indicate that RuOF, is more stable in the gaseous state than in the solid state. (i) Gaseous RuOF, can be transfered from the fluorinating reactor to the product container through the connecting line heated to 150 "C. (ii) In the mass spectrum of R u O F ~the , ~ highest peak is for the fragment RuOF3+ and in all of the fragments expected for RuOF,, only the peak due to RuF4+does not appear. This means that the rupture of a Ru-F bond of
The Journal of Physical Chemktry, Vol. 82, No. 7, 1978 783
Thermal Decomposition of RuOF,
C
.0 c
E? in L)
a
,
,
,
1200 1100 1000 900
, 800
,
,
700 600 500
4
I
Wave numbers in cm" Figure 4. Infrared spectra of solid and gaseous RuOF,; the latter is from the previous work for c ~ m p a r i s o n .The ~ band at 1040 cm-' can be assigned to Ru=O bonds," which disappears upon condensation. The disappearance of this band may be related to the instability of solid RuOF,.
the molecule is brought about by electron impact in preference to that of the Ru=O bond. (RuF4+gives the highest peak in the mass spectrum of RuF,.17) Besides the relative instability of RuOF, to MoOF4 and WOF4, we must therefore discuss the lower stability of its solid. Instability of RuOF4 in comparison with MoOF4 and WOFl may be explained as follows. In these oxide tetrafluorides each metal is in the hexavalent state. Mo(V1) is stable because its electron configuration is the same as that of the ground state (4d5 5s1).lS W(V1) is obtained by promoting one 6s electron to the 5d orbital (6s25d4 6s1 5d5); this promotion can take place easily because there is not much difference in energy between the 6s and 5d orbitals. The presence of one electron in each of the five d orbitals also contributes to the stability of both Mo(V1) and W(VI).18 On the contrary, Ru(V1) is in an excited state and therefore unstable, because promotion of one 4d electron to the 5p orbital is necessary to produce the hexavalent state (5s1 4d' 5s1 4d6 5p1). The fact that Ru(V1) rarely forms its binary compounds supports the instability of Ru(V1); RuFs which may be the only binary compound of Ru(V1) is known to be ~ n s t a b 1 e . l ~The instability of RuOF, partly arises from Ru(V1) being in an excited state. The structure of solid RuOF4 is not known, because its x-ray diffraction pattern has not been obtained yet. However, infrared analysis provided a clue to the structure of solid RuOF,. Figure 4 shows the infrared absorption spectra of solid and gaseous RuOF,; the latter is from the previous work for c o m p a r i ~ o n .The ~ spectrum of the gas is similar to those of gaseous M O O F , ~and ~ WOF421 presented by other workers, both in position of the two bands and their relative intensities; the bands a t 1040 and 720 cm-l are assigned to the stretching vibrations of the Ru=OZ2 and Ru-F bonds, respectively. However, a difference is noted between the spectrum of solid RuOF, and those of MoOF, and WOF4 solids. In the case of MoOF,, the similarity of spectra of its solid and
-
-
vapor indicates that the metal-oxygen multiple bonds are still held by the solid. Edwards and Steventon concluded that the solid had a Mo-F-Mo bridge systemsz0Also, the spectra of solid and gaseous WOF, are similar to each othersz1 Edwards and Jones however suggests from the crystal structure that in the case of solid WOF4 the band at 1050 cm-l is due to a W-F stretching frequency and a part of the band system centered a t 650 cm-l relates to W-0-W bridge bonds. Their conclusion was that solid WOF4 consisted of tetramer molecular units with an oxygen bridge system. Contrary to these two studies, the spectrum of solid RuOF, involves no band around 1040 cm-l; instead, it has such a characteristic band at 630 cm-l as is not seen in the spectra of either MoOF, or WOF, solid (see Figure 4). This reveals that Ru=O bonds disappear on condensation. The instability of solid RuOF4 is related to the disappearance of Ru=O bonds. Comparison of the spectra of solid RuOF, and WOF4 does not permit assignment of the bands between 600 and 700 cm-l in the former to Ru-0-Ru bridge bonds, because there is no band around 1050 cm-l which would be expected from analogy to WOF, for the frequency of the Ru-F bonds trans to the oxygen bridge bonds. Moreover, the earlier study on [Cl6Ru-O-RuCl5I4- ions makes the presence of the oxygen bridge bonds in solid RuOF4 improbable, because it indicates that Ru-0-Ru bridge bonds are not so unstable as to decompose at temperatures below 100 0C.23 From the comparison with the absorptions of solid MoOF,, we assign the two sharp bands a t 700 and 650 cm-l to Ru-F stretching frequencies, the broad band centered a t 520 cm-l to the stretching vibration of Ru-F-Ru bridge bonds, and the characteristic band a t 630 cm-l to the stretching frequency of the Ru-0 single bonds containing an unpaired electron. The change of Ru=O bonds into single bonds may be explained as follows. On condensation, a fluorine atom of a neighboring RuOF4 molecule approaches the particular RuOF4 molecule considered from a direction trans to its Ru=O bond, forming a fluorine bridge bond. This raisesz4 the potential energy of the Ru=O bond through d-level splitting of the Ru atom and leads to breaking of the 7r bond, which causes the change of double bonds into single bonds. Thus, we concluded that the instability of solid RuOF4 was brought about by the change of Ru=O bonds into unstable single bonds in condensation. Solid RuOF, was found to decompose in large quantity a t about 70 "C. This thermal decomposition may be responsible for the complicated behavior of ruthenium encountered in the fluoride-volatility p r o c e s s e ~ . ~ ~ Acknowledgment. We thank Mr. S.Tsujimura, Dr. M. Iwasaki, and Mr. T. Tsujino for their valuable discussion. We are also indebted to Dr. H. Hashitani, Mr. T. Adachi, and Mr. H. Yoshida for their aid in chemical analysis, to Mr. A. Hoshino in gas chromatography, and to Dr. K. Ohwada and Mr. H. Shinohara in infrared analysis.
References and Notes (1) D. V. Steidl and R. L. Jarry, USAECRep.,ANL-7575, 55 (1969); R. C. Vogel et al., ibid., ANL-7550, 28 (1968). (2) M. A. Hepworth and P. L. Robinson, J . Inorg. Nucl. Chem., 4, 24 (1 957). (3) T. Sakurai and A. Takahashi, J . Inorg. Nucl. Chem., 39,427 (1977). (4) J. H. Holloway and R. D. Peacock, J . Chem. Soc., 527 (1963). (5) G. H. Cady and G. B. Hargreaves, J . Chem. Soc., 1568 (1961). ( 6 ) U.K.A.E.A. (Reactor group), A.S.T.M. card No. 18-1139. (7) H. Hashitani and H. Muto, Jpn. Anal., 14, 1114 (1965). (8) H. Hashitani, K. Katsuyama, and K. MotoJima, Talanta, 16, 1553 (1969). (9) J. H. Holloway and R. D. Peacock, J . Chem. Soc., 3892 (1963).
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The Journal of Physical Chemistry, Vol. 82, No. 7, 1978
(10) T. Sakurai and A. Takahashi, unpublished results. (11) T. Sakurai, A. Takahashi, and Y. Komaki, J . Nucl. Sci. Techno/., 11, 74 (1974). (12) U. Merten, J . Phys. Chem., 63, 443 (1959). (13) (a) W. E. Bell, U. Merten, and M. Tagami, J. Phys. Chem., 65, 510 (1961); (b) W. E. Bell and M. Tagami, bid., 87, 2432 (1963). (14) R. C. Rea and T. K. Sherwood, "The Properties of Liquids and Gases", McGraw-Hill, New York, N.Y., 1966, p 520. (15) 0. Kubaschewski and E. L. Evans, "Metallurgical Thermochemistry", Pergamon Press, New York, N.Y., 1958, p 151. (16) (a) N. G. Schmahl and P. Sieben, "National Physical Laboratory Symposium No. 9, The Physical Chemistry of Metallic Solutions and Intermetallic Compounds", Vol. 1, Her Majesty's Stationary Office, London, 1959, p 2k; (b) T. Kikuchi, T. Kurosawa, and T. Yanahashi,
Millero, A. Lo Surdo, and C. Shin
J . Jpn Inst. Met., 28, 497 (1964). (17) T. Sakurai and A. Takahashi, unpublished results. (18) E. Cartmell and G. W. A. Fowles, "Valence and Molecular Structures", Butterworths, London, 1961, p 49. (19) H. H. Claassen, H. Selig, J. G. Malm, C. L. Chernick, and E. Weinstock, J . Amer. Chem. Soc., 83, 2390 (1961). (20) A. J. Edwards and E. R. Steventon, J . Chem. Soc. A, 2503 (1968). (21) A. J . Edwards and G. R. Jones, J . Chem. SOC. A , 2074 (1968). (22) C. G. Barraclough, J. Lewis, and R. S. Nyholm, J . Chem. Soc. A, 3552 (1959). (23) A. M. Mathieson, D. P.Mellor, and N. C. Stephenson, Acta Crystalbgr., 5. 185 (19521 .(24) Reference 18, p 192. (25) G. Manevy et al., French Report CEA-N-1479, 243 (1971). \
The Apparent Molal Volumes and Adiabatic Compressibilities of Aqueous Amino Acids at 25 OC Frank J. Millero,*+ Antonio Lo Surdo,+ and Charles Shin$ Rosenstiel School of Marine and Atmospheric Science, University of Miami, Miami, Florida 33 149 (Received October 21, 1977) Publication costs assisted by the National Science Foundation
The apparent molal volumes and adiabatic compressibilities of 15 amino acids in water have been determined at 25 "C from precise density and sound measurements. The electrostriction partial molal volume and compressibility of the amino acids were determined from p(meas) - P(int) and I?'(meas) - I?'(int). The values of P ( i n t ) have been estimated from P of the uncharged amide isomers and PCryst, the crystal molal volume. Both P(e1ect) and P(e1ect) yield values of 4.1 f 0.4 for the number of water molecules hydrated to the amino acids. Group contributions for the partial molal volume and adiabatic compressibilities have been determined from the amino acids. The volume results are in good agreement with the values calculated by other workers on soluble organic solutes. Introduction Recently there has been an increased interest in the state of water in the living cell. Since most biological macromolecules are physiologically active in aqueous solutions, a knowledge of water-protein interaction is necessary to understand the role of water solvated to soluble organics in the living cells. A better understanding of this type of interaction may be obtained from dipolar ions. Since amino acids are zwitterions in aqueous their volume and compressibility properties should reflect structural interactions with water molecules as in the case of electrolytes. Many experiments4on the properties of protein solutions indicate that ca. 0.3 g of water per gram of dry protein are bound or hydrated in aqueous solutions. Work on the volume properties of proteins and some amino acids have shown that they undergo a decrease in volume upon dissolving in ~ a t e r . ~ -This l ~ decrease in volume upon dissolution is similar to what occurs for electrolyte~l~ and can be attributed to electrostriction due to water-protein, and water-amino acid interactions. The number of water molecules determined from the electrostriction of proteins16 and amino a c i d ~ (using ~ ~ J ~methods that have been applied to electrolytes)lg agree very well with the values determined by other Although these results might indicate that proteins behave as electrolytes, the effect of pressure on the specific volumes12 indicate that the so-called "bound water" behaves more like water solvated to soluble organic solutes.16 To better
understand the hydration of proteins and amino acids, basic physico-chemical properties are needed. In this communication, we report on the volume and compressibility measurements made on amino acids in water at 25 "C. Experimental Section The amino acids used in these studies were obtained from ICN Pharmaceuticals, Inc., and were used without further purification. All solutions were made by weight with ion-exchanged (Millipore Super Q) water. All weights were vacuum corrected. The densities were measured a t 25 "C to f 3 X g cm-3 with a vibrating flow densimeter (Sodev, Inc.). The values of Ad = d - do, where d and do are the densities of the solution and water are listed in Table I.2oThe system was calibrated with N2gas and ion-exchanged water using the densities of Ke11.21 The accuracy ( f 3 ppm) of the system has been determined by measuring the densities of standard seawater solutions.22 The sound velocities were measured a t 2 MHz to a precision of f0.02 m s-l using a "sing around" sound velocimeter (Nusonic, Inc.). The system was calibrated with pure water using the sound velocities of Del Grosso and made^-.^^ The accuracy ( f 0 . 1 m s-l) has been determined by measuring the speed of sound in seawater solutions.24 The relative sound velocities (Au)of the aqueous solutions at 25 "C are listed in Table 120 and were determined from the frequency measurements using
TRosenstiel School of Marine and Atmospheric Science, Miami Fla.
33149.
* Chemistry Department, University of Miami, Coral Gables, Fla.
33100.
0022-3654/78/2082-0784$01 .OO/O
0 1978 American Chemical Society
