A microcalorimetric estimation of the surface acidity of NiSO4·xH2O and catalytically active Al2O3

JOURNAL OF COLLOID SCIENCE 20, 8 3 8 - 8 ~ 5

(1965)

A MICROCALORIMETRIC ESTIMATION OF THE SURFACE ACIDITY OF NiSO4.xH20 AND CATALYTICALLY ACTIVE Al2Oa William H. Wade, Shiichiro Teranishi, and Jack L. Durham The University of Texas, Department of Chemistry, Austin 12, Texas Received March 3, i965 ABSTRACT The relative surface acidity of several alumina powders with a variety of specific surface areas has been measured by a microcalorimetric technique. The alumina surfaces studied here are those for which the catalytic a c t i v i t y for alcohol dehydration has been measured. The technique used was to measure the heats of immersion (hill) in pyridine and 2,6-dimethyl pyridine. A maximum in AH~ for both liquids was observed as a function of a particle size corresponding to a maximum in the catalytic activity. The technique was also applied to NiSO4.xH~O samples, and the results were compared to other surface acidity measurements for these samples. INTRODUCTION

In recent studies in this laboratory the catalyzed dehydration of ethanol (1) and methanol and isopropanol (2) has been investigated. The catalysts were four samples of Al~03 the specific surface area (~) of which varied from 2.72 to 221 m.2/g. The results uniformly showed the catalytic activity to be a marked function of alumina particle size with maximal activity being observed for samples of intermediate ~. This behavior is consistent with excessively weak bonding of the alcohols on the highest area alumina and excessively strong bonding of water on the lowest Z sample (1, 2). Moreover, this is consistent with the room temperature heats of immersion of these aluminas in H20 (3), where AH¢ uniformly decreased with increased ~. On the basis of this latter study and several others (4-6), it was postulated that this common decrease in AHi with increased ~ is due to a gradual crystalline to amorphous surface transition. Exactly how this affects the chemical properties of the surface is still open to question. It is well established (7) that A1203 surfaces are highly populated by surface hydroxyl groups and ionization of the - O H bond is responsible for the usually observed (8, 9) acidic character of these surfaces. Chessick and Zettlemoyer (10) have rated the surface acidity of silica- alumina catalysts by means of immersional heats measurements in BuNHs. Since regular trends in catalytic activity have been noted (1, 2) it appeared worth while 838

SURFACE ACIDITY OF NISO4.XH20 AND CATALYTICALLYACTIVE A]2o3 839 to explore the existence of any correlation between catalytic activity and surface acidity for the same alumina samples. To establish further the validity of surface acidity rating by heat of adsorption of an amine an auxiliary experiment was performed. Tenabe (11) prepared various samples of NiSO~-xH20 by calcining NiSO4-6H20 at various temperatures. By titrametric techniques he noted a maximum in acidity for samples heated to approximately 350°C. Similar samples should be easily generatable in any laboratory since a relatively large surface area of maximal purity is generated during the dehydration process itself. Similarly heats of immersion were measured as for the four alumina samples. EXPERIMENTAL

Samples Four samples of alumina with various specific surface areas and six samples of NiSO4-xH20 were investigated and are listed in Table I. The four samples of alumina are the same as used in the dehydration of methanol, ethanol, and isopropanol which have been reported previously (1, 2). The six samples of NiSO4.xH~O were prepared by the calcination of C.P. NiSO4.6H~O in air under the conditions listed in Table I. The B.E.T. surface areas were obtained from Kr adsorption isotherms for all alumina samples and Ar isotherms for all NiSO4.xH20 samples. All surface areas were independent of outgassing temperature above 120°C. The samples prior to immersion were outgassed at 10-5 mm. Hg for 48 hr. in Pyrex bulbs and were sealed off on the outgassing apparatus.

Solutions Pure grade n-decane from Phillips Petroleum Company was employed as solvent for pyridine, 2,6-dimethyl pyridine, 2,4-dimethyl pyridine, and TABLE I Sample

Area (m?/g.)

A12Oa(A)

221

AI~O3(B) Ah03(C) AI~O3(D) NiSO4"xH~O NiSO4"xH20 NiSO4"xH20 NiSO4. xH20 NiSO4.xH~O NiSO~'xH20

104 65 2.72 22.9 27.2 22.9 19.8 15.3 13.1

Ref. Amorph, supplied by Aluminum Company of America vA120~, supplied by Gulton Ind., Inc. ~,Al~O3, supplied by Gulton Ind., Inc. aA120~, supplied by Godfrey L. Cabot, Inc. Calcined at 250°C. for 4 hours in air Calcined at 300°C. for 4 hours in air Calcined at 350°C. for 4 hours in air Calcined at 400°C. for 4 hours in air Calcined at 450°C. for 4 hours in air Calcined at 500°C. for 4 hours in air

WADE,

TERANISHI,

AND

DURHAM .

F

O-

0 0. I N 0 0.5N 8 0.5N l 0.5N

.pyridine-decane pyridine-decane 2.6 dimethyl pyridine-decant 2.4 dimethyl pyridine-decana

O-

I-

O-

2.72

FIG.

1.

65

Heats of immersion

104

221

of alumina5 in the solution as noted at 150°C.

n-butylamine. The concentration of all solutes used in these experiments was either 0.1 or 0.5 M in n-decane solutions. Solutions were dehydrated using 5X molecular sieve. Spectroquality reagent grade pyridine and pure grade 2,6-dimethyl pyridine, 2,4-dimethyl pyridine, and n-butylamine were obtained from Matheson, Coleman, & Bell Co. Calorimder The calorimeter has been described previously (4). It is of t’he twin adiabatic type with thermistor temperature sensing elements. All measurements were made at 25” f 0.01%. Thin wall-spherical Pyrex bulbs which were shattered completely during immersion were used for al1 samples. Sample weights varied from 4 to 0.1 g. depending on the sample area. Total temperature changes during immersion varied from 1O-2 to 1O-3 degree. Differential base line temperature variations commonly were less than 10-W. over the 30 minute periods required for an individual measurement. RESULTS

AND

DISCUSSION

The heats of immersion (AHi) of four aluminas with surface areas from 2.72 to 221 m.“/g. at a fixed outgassing temperature are shown in Figs. 1

SURFACE ACIDITY OF NISu4"XH20 AND CATALYTICALLY ACTIVE A12O3 841

and 2 for pyridine-decane, 2,4-dimethyl pyridine-decane, and 2,6-dimethyl pyridine-decane solutions. Three trends are evident in the data: I. The alumina with the surface area of 65 m?/g. has the maximum value of AH~ above ,~175°C. outgassing temperature (Fig. 2), but a regular increase of AHi with increased particle size below ,~150°C. outgassing temperature is observed. 2. The concentration range of pyridine in decane used in these experiments (0.1 ~, 0.5N) yields almost constant values of AH~. 3. The heats of immersion of aluminas in pyridine-decane solution are much larger than in 2,4-dimethyl pyridine-decane solution. The effect of increased outgassing temperature on the heats of immersion of aluminas of various specific surface areas can better be shown in Figs. 3 and 4. The heats of immersion of all aluminas initially drop and then pass through a shallow minimum in the vicinity of 350°C. The microcalorimetric method was also applied to the various NiSO4. xH20 samples. The heats of immersion of pyridine-decane, 2,6-dimethyl pyridine-decane, and n-butylamine-decane are shown in Fig. 5. It is observed that the heats of immersion of NiSO4-xH~O have maximal

400.

-AHm ergs

cm-'~-J 50C

2001

o O,5N pyridine-declone • 0.5N 2.6 dimethyl pyridine-decone 2.72

65

104 221 ~.. (rnZ/grn)

~IG. 2. Heats of immersion of aluminas in the solution as noted at 250°C.

842

WADE,

TERANISHI,

AND

DURHAM

65me/g 400/104m*/g -AHi r. (5)

I)

-

300-

[*PI

m2/g

l

200

300 T,,

3. Dependence outgassing temperature. FIG.

of heats

of immersion

400 PC) of aluminas

in pyridine-decane

upon

.65 m2/g

300-

.

2.72 m2/g 2

(&;

i104m1g

Y /”

221 m2/g

I 200 FIG. decane

4. Dependence upon outgassing

of heats of immersion temperature.

300

400 of aluminas

in 2:Cdimethyl

pyridine-

SURFACE ACIDITY OF NISO4"XH20 AND CATALYTICALLY ACTIVE Al2o3

843

0.SN butylarnine-decane 600-

500 0.5N pyridine-decone .A HI (ergs~ -

400-

0.SN 2.6 dirnethyl pyridine 300 3(~)0

'

4()0 Te (°C)

'

560

Fio. 5. Heats of immersion of NiSO4.xH20 in the solutions as noted as a function of precalcination temperature. values between 350 and 400°C. for pyridine, 2,6-dimethyl pyridine, and BuNH~ solutions. However, the heats are somewhat greater for BuNH2 than for pyridine; the heats for pyridine in turn are considerably greater than for 2,6-dimethyl pyridine. The base strength of BuNH2 is greater than for pyridine and 2,6-dimethyl pyridine; this probably explains the enhanced AH~'s for BuNH~; however, the base strengths of the pyridines are quite similar. The low values of AHi for 2,6-dimethyl pyridine compared to the AH~ for pyridine is here postulated to be due to a considerable steric hindrance by the a, a' methyl groups. Alternately, it might be better to consider that the methyl groups prevent intimate contact between a surface hydroxyl group and the nitrogen with its lone pair of electrons. From these data alone it is not possible to determine if proton transfer actually occurs. If it does, the greater distance of separation of the resulting ion pair using

844

WADE,:TERANISHI, AND DURHAM

2,6-dimethyl pyridine once again would dictate a decreased interaction energy. It has been or will be noted thai little or no consideration is given to simultaneous interaction between the alumina or NiSO4.xH20 surfaces and the other solution component, n-decane. These interactions are quite weak (5) compared to those of the highly polar amines as evidenced by the latters relatively large values of &Hi (a factor of 5 greater than usually observed for hydrocarbons). Further experimental evidence is provided by the similarity of AH~ in 0.1 and 0.5 M pyridine-decane solutions. As stated earlier, Tenabe (11) notes a maximum in surface acidity as measured titrametrically at approximately 350°C. The occurrence of similar maxima in AH~ for three different amines at the same temperature augurs for a common source of explanation. This most likely is a strong interaction between acidic hydroxyl groups and the amines. Still more striking is the agreement between the heats of immersion of the aluminas and their catalytic activity for dehydration of alcohols. The catalytic activity was found previously (1, 2) to have a maximum value for the 65 m.2/g. -y-AI~O3and with the present data for samples outgassed at temperatures where the A12Oais catalytically active (250°-350°C) (Fig. 2) once again augurs for a common source of explanation. This would indicate that catalysis perhaps involves the formation of species such as R-OH~+. Once again considerable steric hindrance is observed for 2,6-dimethyl pyridine as evidenced by the low values of AH~. It is interesting that the ratio of AH~ in 2,6-dimethyl pyridine to the &Hi in pyridine is 0.40-0.52, i.e., approximately constant, for all four alumina samples. As a final check on the steric hindrance effect as postulated the AH~ was obtained for the four alumina samples in 0.5 M 2,4-dimethyl pyridinedecane solution. In this latter there should be little steric hindrance, and this is evidenced by the AH~ differing but little from those in 0.5 M pyridine-decane solutions. Recently there has been an effort in the field of catalysis to separate "acid" sites into those with either Lewis or Bronsted character (12, 13). Perhaps the differing thermaI behavior between pyridine and 2,6-dimethyl pyridine partially reflects the presence of both Lewis and Bronsted sites which preferentially react with only one of the two species. In summary the hea~s of immersion in amines correlate well with the measured activity for Catalytic dehydration of alcohols by aluminas and likewise correlate well with the independently measured surface acidity of NiS04. xH20. ACKNOWLEDGMENT The authors express appreciation to the Robert A. Welch Foundation for their continued interest and support. Also, appreciation is expressed to Mr. James Gardner for assistance in the operation of the calorimeter.

SURFACE ACIDITY OF NISO4"XH20 AND CATALYTICALLY ACTIVE A12O3 845

REFERENCES WADE,W. H., TERANISHI, S., AND DURHAM, J. L., J. Phys. Chem. 69, xxx (1965). WADE,W. H., TERANISHI,S., AND DURHAM,J. L., in preparation. W~uv,, W. H., AND HACKE~MAN,N., J. Phys. Chem. 64, 1196 (1960). MAKRrDES,A. C., AND HACKERMAN,N., J. Phys. Chem. 63, 594 (1959). WADE, W. H., AND I'IACKERMAN,N., Advan. Chem. Ser. 43, 222 (1964). WHALEN, J. W., Advan. Chem. Set. 33, 281 (1961). PER1, J., J. Phys. Chem. 69, 211 (1965). JOHNSON, 0., J. Phys. Chem. 59, 827 (1955). BENESI, H. A., J. Am. Chem. Soc. 78, 5490 (1956). CHESSICK, J. J., AND ZV.TTLEMOV~R,A. C., Advan. Catalysis 11, 263 (1959); J. Phys. Chem. 62, 1217 (1958). 11. TENADE, K., J. Res. Inst. Catalysis, Hokkaido Univ. 10, 229 (1962). 12. HOLM, V. C. F., BAILEr, G. C., AND CLARK, A,, J. Phys. Chem. 63, 129 (1959). 13. LEFTIN, H. P., ANn HALL, W. K., Acres Congr. Intern. Catalyse, ~', Paris 1960, Vol. 1, 1353 (1960). 1. 2. 3. 4. 5. 6. 7. 8. 9. 10.

DISCUSSION H. Z. Friedlander (Dorr-Oliver, Inc., Stamford, Connecticut): What is the chemical nature of the acidic sites on nickel sulfate? W. H. Wade: For alumina, surface hydroxyl groups are the Bronsted sites and the exposed aluminum ions are usually considered to be the Lewis sites. For NiS04, the nature of acid sites must be rather more complex. The hexahydrate has largely been decomposed leaving a rather ill-defined surface structure. A. C. Zettlemoyer (Lehigh University): It would be valuable to have the surface capacity for butyl amine versus the pyridines. If the acid sites are nearest neighbors, then this situation would alone contribute considerably to the higher heats of immersion found with the butyl amine. W. H. Wade: The present work must be considered incomplete since we have yet to obtain the requisite isotherms permitting a molecular area evaluation.