1120
William Hertl and Angela Maria Cuenca
ber of negative charges on S increase. When an additional aromatic ring is available for 0- attack, as in the case of diphenyl-4-carboxylic acid, the rate is approximately twice that of benzoic acid (k(0H + S ) / k ( O S) 2 98). However, with the diphenic acids which provide two aromatic rings und a doubly negative anion, the rate is established by both effects. The ortho or para nature of the diphenic acids does not have any apparent effect on the 0 rate. The most interesting trend in the k ( 0 S) values is exhibited by the singly charged anions of benzoic, naphthoic, and anthroic acids for which the rate of the reaction increases by over a factor of 10 in going from the simplest molecule to the most complex. If the rate of addition of 0- to the aromatic ring were governed simply by the number of sites available for bonding, then one would expect the values of k ( 0 - + S) to be in the ratio of 6:10:14, which they are not. Furthermore, one would expect that some sites on the aromatic nucleus (for example, the bridgehead carbon atoms) would be sterically unaccessable for 0- additiofi rendering the predicted range of rate constants even smaller. Rather, the value of k ( 0 + S)for these compounds appears to be related to the extent of delocalization of the aromatic A-electron system for which benzene < naphthalene < anthracene. Apparently, the rate of free-radical addition to aromatic hydrocarbons is governed somewhat by the energy required to isolate a single 7r electron a t a carbon atom to form a " u complex."18 It is unfortunate that the solubility of these aromatic hydrocarbons (even as the carboxylate deriva-
+
+
tive) is so low in aqueous solution which precludes a further test of this somewhat crude model. Mechanistically, it is not yet clear whether the 0- (or OH) radical adds directly to a specific site or first forms a molecular complex with the aromatic A system which rapidly undergoes reorganization to give the various isomers. The 0- radical is clearly a weaker electrophilic reagent than is OH in its addition reactions with the aromatic ring as shown by its less than diffusion-controlled reaction rates. It is difficult to assess the reasons for the difference between this acid-base pair but it should be noted that a similar behavior has been observed for NH3+-NH2 radicals. Specifically, NH3+ adds to benzene to give an intermediate observable by the pulse radiolysis technique; NH2 does not give a detectable transient.lg The lowestlying molecular orbital available for bonding in OH is virtually identical with the corresponding orbital in 0 thereby ruling out considerations on symmetry and energy grounds. At the present time we can only observe that the charge distribution and polarizability differences in the radicals may account for the observations reported here.
Acknowledgment. We thank Dr. E. Hayon of the U. S. Army Natick Laboratories for the use of the pulse radiolysis apparatus and his interest in this work. (17) G. Hughes and H. A . Makada, Trans. Faraday SOC., 64, 3276 (1968). (18) J. N. Murre% S. F. A. Kettle, and J. M. Tedder, "Valence Theory," 2nd ed, Wiley, New York, N. Y., 1970. (19) M. Simic and E. Hayon. J. Amer. Chem. SOC.,93, 5982 (1971).
Infrared Kinetic Study of Reactions of Alcohols on the Surface of Alumina William Herti*' and Angela Maria Cuenca Departamenlo de Quimica, Facultad de Ciencias, Universidad de /os Andes, Merida, Venezqela
(Received December 6 . 1972)
Methanol, ethanol, propanol, and butanol adsorb on alumina as stable alkoxyl groups. Above 150" in the presence of alcohol vapor or air these alkoxyl groups are converted to carboxyl groups. A kinetic analysis of this reaction using infrared spectroscopy showed that the surface reaction is pseudo-zero order. The presence of a Lewis acid site (Al+) is necessary in order that the gaseous alcohol or oxygen may adsorb and produce an oxygen-containing species. This species is mobile on the surface and reacts with the bonded alkoxyl groups to produce a bonded carboxyl group. Poisoning of the Lewis acid sites with pyridine strongly inhibits the reaction. Based upon the data a reaction mechanism is proposed.
Introduction The catalytic dehydroxylation of alcohols Over alumina has been studied for many years.2a Generally, lower ternperatures produce ethers and higher temperatures produce alkenes, although methanol produces ethers, CzH4, C2H6, at higher temperatures' The genera' pattern for these catalytic reactions is the removal of a hydroxyl group and of a proton. There have been several infrared spectroscopic studies of alcohol adsorption on alumina, commencing with BaThe Journal of Physical Chemistry, Vol. 77, No. 9 , 1973
bushkin and Uvarovzb who found that, a t low temperatures, ethanol Produced surface -OH, -CHZCH2, and ethoxy floups. Greenler3 carried out a detailed study of methanol and ethanol adsorption on alumina, which included the use of ?3C and deuterated compounds for iden(1) Present address: Technical Staffs Division, Corning Glass Works, Corning, N. Y. 14830. (2) (a) M. E. Winfield in "Catalysis," Vol. VII, P. H. Emmett, Ed., Reinhold, New York, N. Y., 1960, p 93 f f ; ( b ) A . A . Babushkin and A. V. Uvarov, Dokl. Akad. Nauk SSSR, 110, 581 (1956). (3) R . G. Greenler, J. Chem. Phys., 37, 2094 (1962).
1121
Reactions of Alcohols on the Surface of Alumina tifying the various bands which appeared. Three surface entities were found, uiz., (1) physically adsorbed alcohol removable by evacuation at 35", (2) adsorbed alkoxyl groups, and (3) surface formate ions (from methanol) or acetate ions (from ethanol) when the alumina was heated to 170" in the presence of alcohol vapor. These carboxyl groups are stable to >400". All subsequent work has confirmed his observations and band identifications. Greenler suggested that the alcohol reacts with an A1 atom to produce alkoxyl groups with the elimination of one hydrogen, and with an oxide atom to produce carboxyl groups with the elimination of three hydrogens. In reviewing this work, both Little4 and Hair5 considered the reaction between the alcohol and a surface hydroxyl group as more likely to form alkoxyl groups. Kagel6 studied C1 through Cd normal alcohols, and concluded that the weakly adsorbed alcohols were H bonded to the surface. The chemisorbed alkoxyl groups were formed by adsorption of alcohol on oxygen atoms to form alkoxyl and hydroxyl groups. He proposed that the chemisorbed alkoxide reacts with an adjacent A1-OH group to form a bridged carboxylate. Deo and Dalla Lana, in a study of 1-propanol7 and secondary alcohol* adsorption on alumina, surmised that the alkoxylate forms by adsorbing alcohol on A13+ which then splits off H+, and that the carboxylate forms by reaction of an alkoxylate with an adjacent OH group. On the basis of NaOH doped alumina, they concluded that the highfrequency O H groups (-3785 cm-l) are the principal cause of dehydration of the alkoxyl group, and that by removing these groups a reaction involving A13+ ions becomes dominant. Although the experimental observations between these different workers are generally in agreement, the mechanisms proposed for both the alkoxide and the carboxylate formation are rather divergent. Most of the experiments involved heating the alumina in alcohol vapor and recording the spectra at room temperature. All the mechanistic conclusions were drawn on the basis of these spectra and gas analyses. This study describes a kinetic evaluation of the formation of surface carboxylates using methanol, ethanol, propanol, and butanol. A study of this type yields much more information about the course of the reactions than does observation of only initial and final states. In so far as the experiments carried out here overlapped those reported above, they are in broad agreement. In this study, all spectra were recorded at reaction temperature, and reaction curves were obtained for the formation of carboxyl groups on the surface. The reactions were carried out at various temperatures, using various pressures of alcohol vapor or air. Experiments were also carried out in which the alumina surface was treated with pyridine.
Experimental Section The details of measuring surface kinetics have been previously d e s ~ r i b e d .Briefly, ~ a self-supporting disk of alumina was mounted in a cylindrical furnace cell placed in a Perkin-Elmer Model 621 infrared spectrophotometer. The furnace cell had water-cooled end plates fitted with Irtan-2 windows, and was connected to a conventional vacuum rack. Gas or vapor at any desired pressure was admitted to the cell for a given time and then evacuated. Spectra were taken at each step of the process. By measuring the intensities of the absorption bands a t various times one can construct a reaction curve. Since the cell was frequently evacuated, the gas-phase concentration
was essentially constant during the course of any given experiment. All spectra were recorded at reaction temperature. The Alon C alumina disks were prepared by pressing approximately 50 mg in a 1-in, die at 24,000 psi between two sheets of tissue paper moistened with acetone. This "sandwich" was heated at 900" in air for 1 hr and transferred to the reaction cell, The cell with the alumina disk was then heated to 400" in uucuo for 15 min, after which the temperature was lowered to the desired reaction temperature and the reaction carried out. Analytical grade reagents were used throughout.
Results In Figure 1 are given some typical spectra in the region 1700-1300 cm-1 taken during the course of a kinetic ex. periment. The bands in the vicinity of 1460 and 1560 cm-1 are due to the symmetric and asymmetric OCO group vibrations of the surface carboxylate. The precise frequencies, as well as the relative band intensities, for each alcohol are given in Table I. With the exception of methanol, the two observed carbonyl bands have integrat,. ed intensities which differ by no more than about 15%. With methanol, the symmetric band is only about I/3 ai1 intense as the asymmetric band. The intensities of these C=O bands during the course of an experiment were used to construct the kinetic reaction curves. Some typical reaction curves are given in Figure 2. These C=O bands increased in intensity only in the presence of gaseous alcohol or added air. No changes were observed in these intensities when the sample was allowed to stand in uucuo at reaction temperature for periods up to several hours. Plots of the lower frequency C=O band against the higher frequency C=O band were linear for any given alcohol, demonstrating that both bands arise from the same functional group. Bands at 2800-3000 cm-1 and near 1385 cm-l, due to C-H stretching and bending vibrations, appeared initially. These C-H bands rapidly reached their final intensity. Thereafter they remain nearly constant during the course of a reaction (cf. Figures 1 and 3). During this initial period the hydroxyl group bands at 3600-3760 cm-I disappeared (Figure 3) and the OH group bands between 3460 and 3560 cm-1 increased in intensity. During the course of the reaction there is a small increase in the region 3600-3100 cm-l. When the alcohol reacted with the surface at a lower temperature (150") bands also appeared near 1070, 1100, and 1150 cm-l. Greenler3 observed bands in this vicinity and assigned them to a surface alkoxide group. Thus, the first reaction is that between the alcohol and some of the surface hydroxyl groups to produce surface alkoxyl groups. In the presence of gaseous alcohol these alkoxyl groups are then oxidized to carboxyl groups. Figure 2 shows that the reaction curves for the formation of carboxylate, using ethanol, propanol, and butanol, are linear throughout except for an initial fast reaction. With methanol, the reaction curves showed a sharp break, but the two portions were linear. In these experiments the (4) L. 'H. Little, "Infrared Spectra of Adsorbed Species," Academic Press, London, 1966, p 178. (5) M. L. Hair, "Infrared Spectroscopy in Surface Chemistry," Marcel Dekker, New York, N. Y., 1967, p 157. (6) R. 0. Kagel, J. Phys. Chem., 71,844 (1967). (7) A. V . Deo and I . G. Dalia Lana,J. Phys. Chem., 73,716 (1969). (8) A. V. Deo, T. T. Chuang, and I . G . Dalla Lana, J. Phys. Chem., 75, 234 (1971). (9) W. Hertl, J. Phys. Chem., 72, 1248 (1968). The Journal of Physical Chemistry, Voi. 77, No. 9, 1973
1122
William Hertl and Angela Maria Cuenca r
1
I 1700
I
I
I500 FREOUENCY (cm-')
I
I
..- - -
Figure 3. Spectra of alumina at 275': initially; - - - - - after sec of reaction with propanol; --- after 145 min of reaction with propanol.
30
J 1300
Figure 1. Spectra of alumina at various times showing growth of carboxylate bands due to reaction of 10 Torr of ethanol at 275". The initial spectrum is at the bottom; the top spectrum is after 8 4 min of reaction.
TABLE II: Gas-Phase Products from Reaction of Alcohols on Alumina; Reactant
Products
Methanol Ethanol Propanol
Mixture of ether and alkanes Ethylene Propylene
The alumina was heated at 400' in vacuo prior to the reaction with the alcohol.
Figure 2. Kinetic plots obtained at 250" for: 0 ,20 Torr of methanol; A , 20 Torr of ethanol; 0 ,10 Torr of propanol; 0 , 5 Torr of butanol. TABLE I: Observed Frequencies (in cm- l ) and Intensities of Carboxyl Bands Produced by Reaction of Alcohols on Alumina Intensity symmetric
oco/
Alcohol
Methanol
Ethanol Propanol Butanol
Asymmetric OCO 1587 1572 1564 1556 1555 1567
Symmetric
OCO 1378 1455 1444 1469 1455 1463
CH bending
intensity asymmetric
OCO
1390 1388 1384
33% 89% 85%
1380
110%
alcohol vapor was evacuated a t short intervals and fresh vapor added, so that the gas-phase composition was essentially constant during an experiment. Under these conditions the reaction curves describe a pseudo-zero-order process. The Journal of Physical Chemistry, VoI. 77, No. 9, 7973
In order to determine the gas-phase products, some alumina was placed on the floor of the furnace a t 250". Alcohol vapor was added for a time sufficient to ensure that the initial bonding reaction to form the alkoxide was complete. Fresh alcohol vapor was then added and allowed to stand for 45 min. Spectra were taken at various times. Initially, unreacted alcohol was observed and at later times only the products listed in Table I1 were found. The rate of carboxylate formation a t various alcohol pressures was measured by carrying out the reaction a t a fixed datum pressure and measuring the rate; then, using a different pressure, the rate was again measured. Thus all the rates for a given alcohol are referred to a datum pressure. This method of normalization is necessary due to slight differences in thickness between different alumina disks. In Figure 4 are given the measured pressure dependencies of the reactioh rates. These curves should reflect the shape of the adsorption isotherm of the alcohol which reacts with the bonded species, and appear to be type I isotherms. The initial fast reaction which appears as an intercept on the kinetic plots accounts for about 10% or less of the total reaction and is completed within several minutes, compared to hundreds of minutes for the complete reaction. This fast reaction takes place during the time when the CH bands are building up to their final intensity and the OH bands above 3600 cm-' are disappearing. By carrying out the reaction at 150" one can readily follow this initial process. At this low temperature one observes a small growth in the C=O band a t 1560 cm-I and a large growth in the C=O band a t 1462 cm-l during the time when the CH bands at 1380 cm-1 are building up to their final intensity, after which there is no further change. On raising the temperature to 250" with evacuation about 75% of the large C=O band and the 1380-cm-' CH band disappear, leaving small C=O bands at 1560 and 1462 cm-1 and the CH band a t 1380 cm-l. Thus, this initial fast oxidation reaction is sufficiently rapid to be
1123'
Reactions of Alcohols on the Surface of Alumina
r
t
c t-/ 1400
1600
le00
FREQUENCY
(CHI)
W
$
I
I
0
a0
I
I
30
40
P (tar)
Figure 4. Plots of carboxylate reaction rate at 250" against pressure of alcohol vapor. The rate at a fixed reference pressure was determined for each disk. A different pressure was then used. All the rates are normalized with respect to the reference pressure. Reference pressures used were methanol, 20 Torr; ethanol, 20 Torr: and propanol, 10 Torr. observable even a t 150". The experiment is somewhat obscured due to the oxidation of loosely bound alcohol, which desorbs on raising the temperature. This loosely bound alcohol would not be present a t the usual reaction temperatures. These intensity changes are not due to the temperature dependence of the infrared absorption bands themselves. Over this temperature range the peak intensities normally decreased about 25%, whereas the large C=O band as well as the CH bands decreased about 75% on raising the temperature. The temperature dependencies were also measured. The reaction rate variation with temperature will depend on at least two factors: the true activation energy of the oxidation reaction and the concentration of the intermediate adsorbed oxidant. In these experiments it is not possible to isolate the two steps. Although this experimental activitation energy has no fundamental significance it is a convenient way to describe an experimental temperature dependence. The results obtained were (in kcal/mol) butanol 16, propanol 16, ethanol 17, and methanol (first linear part of reaction) 21. The kinetic behavior of the methanol is quantitatively somewhat different. The kinetic curves show the rapid initial formation of C=O bands in parallel with the buildup of the CH bands in the 2800-3000 cm-' region for the first few minutes of reaction, as with the other alcohols. After this initial fast reaction the curve for C=O formation is linear for about 50% of the reaction. After this point the reaction curve is also linear, with a slope indicating a reaction rate only about half as fast as during the first half of the reaction. There are also small increases in the CH bands a t 2800-3000 and 1390 cm-l. The gaseous reaction products obtained with methanol are also different from those obtained with the other alcohols. When the partially reacted alumina is heated to 400" in uucuo the C=O bands are stable in frequency and intensity, but the CH band intensities decrease. This is probably due to desorption of the bonded alkoxy1 groups. On heating the alumina at 400" in air, after substantial reaction with methanol, no changes occur in the intensities or frequencies of the C=O bands and the CH band a t 1390 cm-l decreases only slightly. The CH bands near 2900
Figure 5. Spectra of alumina at various times showing the initial fast reaction at 150" using 5 Torr of butanol. The initial spectrum is at the bottom; the top spectrum was taken after 90 sec of reaction. No further changes were observed at this temperature.
I
10
I 20
i
30
TIME (minl
Figure 6. Kinetic plot of growth of carboxyl bands obtained at 275" using 50 Torr of air. Prior to making this plot, the alumina was exposed to 30 Torr of methanol at 150 for 5 min. cm-1 are almost completely removed (15% remaining). The OH band a t 3720 cm-I reappears and the OH band near 3500 cm-I is removed. In Figure 6 is given a reaction curve taken under the following conditions. Methanol was added to the cell at 275" fos 5 min until the intensities of the observed CH bands were essentially constant, indicating that the alkoxyl bonding reaction was complete. Air (50 Torr) was introduced into the cell for a given period of time and then evacuated. This was repeated in exactly the same manner as in a kinetic experiment with alcohol vapor. The reaction curve is linear throughout. The frequencies of the C=O bands were identical with those observed when methanol vapor was used for the oxidation reaction. The same experiment was repeated with propoxyl groups bonded to the surface and then oxidized with air. A linear reaction curve was also obtained, as when propanol vapoI was used, but an additional C=O band was observed at 1580 cm-l, as well as the bands at 1560, 1470, and 1445 cm-l. This new C=O band appears near the C=O band observed with methanol (1587 cm-1). It seems likely that the hydrocarbon side chain had been partially oxidized, thus giving rise to this formate C=O band, as well as the propionate C=O bands observed with propanol vapor oxidation. The Journal of Physical Chemistry. Vol. 77, No. 9, 1973
1124
William Hertl and Angela Maria Cuenca
t 10
I 30
20
I 40
TIME (min)
Kinetic plot of growth of carboxyl bands obtained at 275" using 50 Torr of air from time 0 to 10 (e) using 50 Torr of air plus 5 Torr of pyridine from time 10 to 30 (0); using 50 Torr air only from time 30 to 46 (0).Prior to making this plot the alumina was exposed to 10 Torr of propanol at 150' for 6 Figure 7.
min.
Pyridine is often used as a diagnostic test for the presence of Lewis acid sites, since it preferentially adsorbs on the stronger sites. Some reactions were carried out for a time with pure alcohol vapor and then pyridine vapor was mixed with the added alcohol vapor. In the presence of pyridine the reaction rate was only about 25% as fast. When the reaction was continued without pyridine vapor added there was only a very slight increase in rate, and the reaction rate was still much less than prior to the pyridine addition. This demonstrates the importance of the Lewis acid sites. The kinetic plots of these experiments were similar to that given in Figure 7. An air oxidation experiment was done in which propoxyl groups were intially bonded to the surface. After partial reaction pyridine vapor was added to the air and the reaction was allowed to continue. Figure 7 shows that the rate is only about 15% as fast as initially. This demonstrates that with air oxidation the Lewis acid sites are also important for producing the intermediate species required for the oxidation of the alkoxyl group to the carboxyl group.
Discussion These reactions can be conveniently divided as follows: (1) alcohol bonding reaction to produce alkoxyl groups, (2) oxidation of the bonded alkoxyl groups to produce carboxyl groups, and (3) the initial fast reaction. Alkoxy1 Group Formation. Greenler suggested that the alcohol reacts directly with a surface A1 and eliminates hydrogen. In reviewing this work both Little and Hair consider a reaction between the alcohol and surface hydroxyl groups with the elimination of water more likely. This would be exactly analogous to the reactions between alcohols and silanol groups. This point might be resolved by observing spectral changes in the 3700-cm-l region and detecting water. Kagel's spectra are obscured in this region due to the large amounts of adsorbed water, but he did make a mass spectrometric examination of the gases immediately after the alkoxide formation from EtOD. The gases showed neither hydrogen nor deuterium and the presence of water was questionable, although adsorbed water was observed at room temperature on samples that had been heated to 320" in alcohol vapor. He proposed that the chemisorbed alkoxyl groups were formed by adsorption of alcohol on surface oxide ions to form alkoxyl and hydroxyl groups. There are difficulties of interpretation under these experimental conditions. In the temperaThe Journal of Physical Chemistry, Vol. 77, No. 9, 1973
ture region near 300" the catalytic dehydroxylation takes place which produces alkenes and water. Where the temperature was limited to about 170", where only the alkoxyl bonding reaction takes place, the spectra and gas analyses were taken a t room temperature. Alumina that has been heated readily adsorbs water on coolinglo and water might not appear in gas analyses. This adsorbed water (or H bonded OH groups produced by water adsorption) appears as the very broad, strong band from below 3000 to about 3800 cm-l. Deo and Dalla Lana's spectra show much less initially adsorbed water and the "free" OH bands are well resolved. They examined this region carefully and found that on exposure to alcohol up to 170" the band at 3758 cm-I disappeared completely and a broad band appeared near 3500 cm-l. This agrees with the results obtained here. Deo and Dalla Lana correctly point out that the bands appearing in this region can be difficult to interpret, but considered this to be indirect evidence of H bonding with the surface. From frequency shifts in the 1060-1250-~m-~region, they concluded that the alcohol molecules are loosely H bonded through the O of the alcohol to the surface OH. They surmised that the alcohol adsorbs via the oxygen of the alcohol to a AP+ site and forms alkoxylate by splitting off a hydrogen ion which is then free to migrate to oxide surface sites and form Hz. The spectra taken here at reaction temperatures, as well as those taken at 150" followed by raising the temperature to 275", showed that the alcohol reacts with the free OH groups to produce alkoxyl groups stable at 300" in vacuo. At this temperature coverage of OH groups via H bonding would probably not be important, so that the disappearance of these free OH groups must be due to chemical reaction. The broad band which appears near 3500 cm-1 could,then, be due to adsorbed water produced as a result of the reaction between the alcohol and a surface hydroxyl group. This experimental evidence points to the following reaction ROH
+
)A1-O-H
---+
)Al-O-R
4- H,O,,
or(ads)
Carboxyl Group Formation. Greenler assumed that the alcohol reacts directly with an oxide atom to form the carboxylate by splitting out three hydrogens. Kagel carried out the reaction at 320" until both propoxide and propionate were formed. He evacuated the system at EO", cooled it to room temperature, closed the cell, and heated it to 410". The propoxide disappeared and propionate formed. Gaseous hydrogen was the principal product. Similar results were obtained with butanol and propanol. On the basis of these experiments he concluded that the reaction i s R R I I
The experiments carried out here demonstrated that the presence of gaseous alcohol is necessary for the alkoxyl to carboxyl group reaction to take place. A reaction involving only alkoxy1 and hydroxyl groups cannot explain the mechanism of this part of the reaction. It was also ob(10) M. G. Neumann and W . Hertl, J . Chromatogr., 65,467 (i972)
1125
Reactions of Alcohols on t h e Surface of Alumina served here that a t 150°, the temperature at which Kagel evacuated his cell, some loosely bound alcohol is still present. Thus, in the absence of a leaky cell, this adsorbed alcohol desorbs on raising the temperature and readsorbs on the catalytically active sites, thereby leading to carboxylate formation. Deo and Dalla Lana heated their alumina in propanol vapor for 1 hr, between 170“ and 40O0,after which only carboxyl groups were observed on the surface. Infrared and mass spectra of the gases formed above 200” showed propylene to be a major product and hydrogen was detected. The gases produced at 400” showed mainly propylene and water. Based on gas analyses, and drawing an analogy with the adsorption of formic acid on nickel oxide, they proposed a bridged carboxylate species similar to Kagel’s. The precursor propoxide was considered to be chemisorbed on the electron-abstracting Lewis acid aluminum ion. The organic product of this surface reaction, an alkene, is the same as that obtained in the catalytic dehydroxylation of alcohols over alumina. The reactive intermediate could be common to both. Generally2a the alcohol is assumed to dissociate into carbonium ions and hydroxyl ions, but specific proof of an ion mechanism is difficult to provide. The catalytic dehydroxylation reactions are usually divided into two main categories: (1) an ether mechanism in which the elements of water are removed from two hydroxyl groups, and (2) an olefin mechanism in which the water molecule that is removed is made u p of one OH group together with a hydrogen atom withdrawn from either a CH2 or CH3 group. For the alcohols studied here, the methanol proceeds in part via the ether mechanism, and the other alcohols completely via the olefin mechanism (cf. Table LI). Any mechanism proposed has to be consistent with the following experimental observations. (1) Once the alkoxide is formed a t a given temperature, it is perfectly stable in the absence of alcohol vapor or air. Carboxyl groups form only on exposure to these oxidizing agents. (2) The oxidizing species must add (0). (3) Two (H) are eliminated from the bonded alkoxylate. Kagel and Dalla Lana showed that hydrogen is a product of the reaction. (4) The other gaseous product is (a) an alkene with the same number of carbon atoms as the alcohol, or (b) with methanol an ether and alkanes. (5) The reaction rate is dependent on the surface coverage of the reactive species (cf. Figure 4). (6) The oxidizing species must be mobile on the surface, since a pseudo-zero-order reaction is observed. If this species were not mobile, the reaction would be first order or greater with respect to the surface sites. (7) The oxidizing species is produced principally by Lewis acid sites (the Al+ atoms). Poisoning with pyridine inhibits the reaction. (8) With air as an oxidizing agent, the only possible intermediate species are oxygen molecules or atoms (the possibility of a charged species is not excluded). Poisoning with pyridine inhibits the reaction so that the oxygen molecules must interact with the Lewis acid sites and proceed via these sites. The general picture is that stable alkoxyl groups are bonded to the surface. Gaseous alcohol adsorbs on the Lewis acid sites and produces an oxygen-containing species, which is sufficiently mobile to migrate to the
bonded alkoxylate and react with this group. As a result of the reaction, two hydrogen atoms are eliminated from the alkoxylate and a carboxyl group results. The evidence for carbonium ions is not conclusive and one could write the reaction using any of the following species: (1) carbonium plus hydroxyl ions, (2) alkyl plus hydroxyl radicals, or (3) alkyl radical plus oxygen and hydrogen atoms. Although the experiments carried out here cannot differentiate between these possibilities, the experiments with added air to oxidize the alkoxyl group suggest that an oxygen atom is involved.
o2 + ’0
-’ -
Al+
Al+
‘0
0
‘0
o/
Al+
+ 2[0]
‘0
The reactions that take place with the alcohols as oxidants can also be written in a formal manner using oxygen atoms. For ethanol H,C. 3 CH,,CH,OH + AI+ C’‘ H o/ ‘ 0 H ’ \O’
-
”
AI
o/
+
‘0
+ LO] + 2[Hl +
CH,=CHz
o/
AIf
‘0
and
CHB
[O]
+ H-C-H I
I 0
-
FHI
C=O
I 0
2[H]
While not proving that the oxygen atom is the reactive species, it is consistent with all the experimental observations, including the mixture of products obtained when methanol is used. For methanol one can write
CH,OH
+
-
+
+
CH,. [O] €1. and/or CH,O* + H. and/or CH,. OH. Combinations of these radicals will lead to ether and alkanes, as is found with methanol. The break which occurs in the reaction curve for methanol could be due to a preponderance of one oxidizing species during the earlier part of the reaction and a different one during the latter part of the reaction. The reaction scheme proposed is also consistent with the products obtained from the catalytic dehydroxylation of alcohols. When all the surface alkoxyl groups have been oxidized to carboxyl groups, or concurrently with this reaction, the species produced by the Lewis acid site will AI+
0 ’
\o
u
3
o’A’o;
+
The Journal of Physical Chemistry, Vol. 77,No. 9, 7973
1126
Skolnik, Salesi, Russ, and Goodfriend
be an alkene, an oxygen atom, and two hydrogen atoms. On combination of the atoms the final products are an alkene and water. Initial Fast Reaction. The initial fast reaction appears as a positive intercept on the kinetic plots (or reaction curves) and occurs during the period when the alkoxide builds up to its final constant concentration. At 150" the alcohol bonding reaction takes several minutes. During this time there is a small buildup in the intensities of the C=O bands. Loosely adsorbed alcohol is also oxidized, but on raising the temperature this desorbs and would not be important at higher temperatures. After the first few minutes, no further changes are observed in the intensities of the C=O and CH bands in the presence of alcohol vapor. It is probably that this is the initial fast reaction observed at higher temperatures. Deo and Dalla Lana proposed that the high-frequency OH groups are catalytically active and that they are removed by NaOH doping, after which the aluminum ion sites assume the primary catalytic role. Under the condi-
tions used here, these high-frequency groups are removed during the initial part of the reaction. A possibility is that this could be the cause of the initial fast reaction. The equation for this reaction would be the same as that proposed by Kagel in his mechanism, Another possibility is that the water produced by the bonding reaction adsorbs on the strongest Lewis acid sites and deactivates them. Adding water vapor is a common method of partially deactivating alumina supports in gas chromatography.llJ2
Acknowledgment. The authors gratefully acknowledge the gift of the Alon C alumina by the Cabot Co., Boston. ( 1 1 ) C.G. Scott, J. Inst. PetroI., 45, 118 (1959). ( 1 2 ) Note: An anonymous reviewer suggested the use of a high-temperature H2 treatment to remove adsorbed 02.The experiments reported here show that added oxygen does indeed result in formation of the carboxyl groups from the bonded alkoxy1 groups, but that when the system is evacuated at reaction temperatures no reaction takes place. However, the presence of initially adsorbed oxygen could be the cause of the initial fast reaction observed.
Absorption and Flash Photolysis Kinetic Spectroscopy Studies on Difluoro-, Chlorodifluoro-, Dichlorofluoro-, and Tetrafluorophosphine Edward G. Skolnik,' Robert J. Salesi,' Charles R. Russ, and
P. L. G ~ o d f r i e n d * ~
Department of Chemistry, Universify of Maine. Orono, Maine 04473 (Received September 15, 1972)
Absorption studies and flash photolysis kinetic spectroscopy experiments were carried out on HPF2, P2F4, ClPF2, and FPC12. Absorption results were found to disagree with those of other investigators. Upper limits for two dissociation energies were obtained. No transient spectra that could be assigned to PF-containing fragments were observed.
Introduction The original motivation for the investigations reported here was the desire to detect, identify, and study the electronic absorption spectra of HPF, ClPF, and other unobserved PF-containing radicals by means of flash photolysis kinetic spectroscopy. Although this objective was not realized, information about the absorption spectra and photochemistry of HPF2, ClPF2, PFC12, and P2F4 was obtained. Of particular interest is the fact that results were obtained which disagree strongly with those of other investigators. Experimental Section A concentric photolysis flash sample cell arrangement was used in these studies. The concentric arrangement and the delay system are of the same form as previously reported as part of another apparatus.4 Delay times of from 0 to 300 psec were available, and photolysis flashes of energy between lo00 and 2000 J were used. Spectral plates were taken using a Jarrell-Ash F/6.3 75-000 plane grating spectrograph with a reciprocal dispersion of 11 The Journal of Physical Chemistry, Vol. 77, No. 9, 1973
A/".
Delay times were monitored using a photocell and oscilloscope. Sample cells were 65 cm in length except in the P2F4 studies where the length was 55 cm. The wavelengths available for photolysis were limited by the quality of the quartz used to the region above 1900
A. Wavelengths (in air) were measured on tracings from a Jarrell-Ash 23-100 recording microphotometer using an iron arc calibration spectrum. Sharp features could be measured to within 1A. Difluoroiodophosphine (PFzI) was prepared by the reaction of dimethylaminodifluorophosphine with hydrogen iodide as described by Rudolph, Morse, and Parry.6 Commercial HI was used. The dimethylaminodifluorophos( 1 ) Part of work done by E. G. Skolnik in partial fulfillment of the requirements of the Ph.D. degree at the University of Maine. (2) Part of work done by R. J. Salesi in partial fulfillment of the requlrements of the M.S. degree at the University of Maine. (3) To whom correspondence should be addressed. (4) P. L. Goodfriend and H. P. Woods, Rev. Sci. Instrum., 36, 10 (1965). ( 5 ) R. W. Rudolph, J. G. Morse, and R. W. Parry, Inorg. Chem., 5, 1464 (1966).
