CATALYTIC OXIDATION OF METHANOL
4213
The Catalytic Oxidation of Methanol in the Presence
of Small Amounts of Water
by Thomas C. Franklin and M o b Kawamata Chemistry Department, Baylor University, Waco, Tern8 (Received February 17, 1967)
It has been shown that the coulometrically determined amount of adsorbed hydrogen in 2 N sulfuric acid and in methanolic hydrogen chloride solutions is a measure of the catalytic activity of the electrode for the oxidation of methanol in nearly anhydrous methanol solutions. The rate of oxidation of methanol has been shown to follow the rate law rate a (MeOH) "(area)(Po,)*/'(HCl). From this an abbreviated mechanism was suggested.
A number of studies have been made of the electrochemical oxidation of methanol in aqueous A recent investigation was made correlating the catalytic and electrolytic oxidation of methanol in the same aqueous solution^,^* as part of a series of studies of the possibility of using the coulometric determination of the amount of hydrogen adsorbed on metals as a measurement of the surface area available for various heterogeneous processes.16-18 I n the previous studies in this series, the areas have been measured, and the reactions have been studied and correlated only in aqueous solutions of sulfuric acid and sodium hydroxide. The study of the catalytic oxidation of methanol on platinized platinum has been extended to methanolic solutions containing only small amounts of water in order to evaluate the use of the coulometric method of measuring catalyst areas in nonaqueous solutions. At the same time, it was expected that the study would give information on the effect of water on the rate of oxidation of methanol. Experimental Materials and Methods The platinum catalysts were prepared by sealing lengths of 0.8 mm in diameter wire in 6-mm glass tubing. The exposed wire was 1-3 cm long and electrical contact was made through mercury. The wires were immersed in aqua regia, and then cathodically and anodically charged several times in 2 N sulfuric acid. The cleanness of the catalyst was checked by observing the uniformity of hydrogen evolution. The catalysts were then rinsed with distilled water and plated with platinum black.l9 A 3% solution of chloroplatinic acid containing 0.06%lead acetate was used as the plating solu-
tion.20 Plating was performed at 5-50 ma for 60-300 sec, depending on the surface area of the substrate platinum and the thickness of the deposit desired. The
(1) E. Muller and A. R. Y. Miro, Z . Ekktrochem., 27, 54 (1921). (2) E. Muller, &id., 27, 558 (1921). (3) E. Muller and 5. Takegami, ibid., 34, 704 (1928). (4) H. Hoyer, Z . Naturjorech., 4a, 335 (1949). (5) A. Kutschker and W. Vielstich, Electrochim. Acta, 8,985 (1963). (6) J. E. Oxley, G. K.Johnson, and B. T. Buzalski, ibid. 9,897 (1964). (7) G. A. Bogdanovski and A. I. Shlygin, Zh. Fit. Khim., 34, 57 ( 1960). (8) R. P. Buck and L. R. Griffith, J. Electrochem. SOC.,109, 1005 (1962). (9) G. A. Martinyuk and A. I. Shlygin, Zh. Fit. Khim., 32, 164 (1958). (10) T. 0. Pavela, Ann. A d . Sei. Fennicae Ser. A 11, No. 69 (1954). (11) T. I. Borisova and V. I. Veselovskii, Zh. Fit. Khim., 27, 1195 ( 1953). (12) C. Liang and T. C. Franklin, Electrochim. Acta, 9, 517 (1964). (13) V. S. Bagotzky and Yr. B. Vasilyev, aid., 9, 869 (1964). (14) T. C. Franklin, et al., Final Report Contract No. AF 19(604) 8414, AFCRL-64-517, 1904. (15) T. C. Franklin and 9. L. Cooke, Jr., J . Electrochem SOC.,107, 556 (1960). (16) T. C. Franklin and D. H. McClelland, J. Phys. Chem., 67, 2436 (1963). (17) T. C. Franklin and J. Goodwyn. J. EZeetrochem. Soc., 109, 288 (1962). (18) F. Matsuda and T. C. Franklin, ibid., 112, 767 (1965). (19) W. M. Clark, "The Determination of Hydrogen Ions," Williams and Wilkens, Baltimore, Md., 1932. (20) W. M. MacNevin and M. Levitsky, Anal. Chem., 24, 973 (1952).
Volume 71. Number 13 December 1967
4214
THOMAS C. FRANKLIN AND MOTOO KAWAMATA
catalysts were charged anodically and cathodically several times in 2 N sulfuric acid to clean the surface. The prepared catalyst was rinsed and placed in a cell containing 2 N sulfuric acid and a large platinum gauze electrode to be used as a nonpolarizable electrode. The gauze electrode was platinized in some experiments, and in others it was not. A t no time did this seem to make any difference in the results. Hydrogen was bubbled through the solution until the potential between the catalyst and the platinum gauze became constant. This potential difference was approximately zero. Potentials when measured were compared with another hydrogen electrode in the same solution. The flow of hydrogen was stopped and, after allowing the system to stand for a few minutes, a polarogram was obtained for the oxidation of the hydrogen on the platinum catalyst. The sweep rate of the potential was about 3.3 mv/sec. Polarograms were taken repeatedly until reproducible results were obtained. If the current-voltage curves decreased after each run, the apparatus was disassembled and recleaned. The platinum catalyst was rinsed with redistilled water, freshly distilled methanol, and finally with a methanol solution of the same composition as that t o be studied. It was then placed in a cell containing the appropriate methanol solution with hydrogen chloride added as a supporting electrolyte. Polarograms for the oxidation of adsorbed hydrogen were obtained in the same manner described above. The area under the hydrogen polarograms was determined and used as a measure of the amount of hydrogen adsorbed on the metal surface.*b-** Figure 1 shows typical current-voltage curves in the hydrogen region, taken in aqueous sulfuric acid and methanolic hydrogen chloride solutions. The shaded area is the number of coulombs that was used as a measure of the catalytic activity of the metal. The residual current used to subtract graphically the amount of solution hydrogen that diffuses to and reacts at the electrode is illustrated in both cases. This residual current was determined by immediately rerunning a polarogram in the same hydrogen-saturated solution. This gave the potential for the start of oxidation of the solution hydrogen. Recognizing that the concentration of the hydrogen adjacent to the electrode had been depleted somewhat, the actual residual was drawn from this potential to the upper curve. Results obtained by this method give areas that agree within experimental error (+5%) with oxygen coverage determinations in sulfuric acid. The results in methanol could not be checked against the oxygen COVerage because of oxidation of methanol. It is obvious from the figures that the determination Of areas in The Journal
of
Phyeical Chemistpy
Potential.
Figure 1. Typical current voltage curves for the oxidation of hydrogen in methanolic hydrogen chloride and in aqueous area measured in methanolic sulfuric acid solutions: I, hydrogen chloride solutions; and m, area measured in aqueous 2 N sulfuric acid solutions.
methanol is subject to much more error than the determinations in sulfuric acid. Oxygen was then bubbled over the platinum catalyst and the rate of oxidation of methanol was followed by measuring the formaldehyde concentration as a function of time. The formaldehyde concentration was measured by the chromotropic acid method.21 Blank runs were made both without the platinum and with platinum using nitrogen or hydrogen in the place of oxygen. This established the fact that in the time studied, the catalytic oxidation was the only significant reaction. The methanol was purified by fractional distillation over magnesium.22 The water content in the methanol solution was determined by the Karl Fischer method.23 All water was redistilled from potassium permanganate. All other chemicals were the purest grade available commercially. All experiments were run in a water bath maintained at 25.0 f 0.1”.
Experimental Results Qualitatively, it can be pointed out that the area under the portion of the curve attributed to the oxidation of adsorbed hydrogen is much larger in methanolic hydrogen chloride solution than in aqueous sulfuric acid solutions, but the rate of growth of these maxima is much slower in the methanolic solutions. The slowness of the growth of this region of the curve is probably due to the presence of the chloride ion, since slow adsorption was also noted in aqueous hydrogen chloride solutions. The hydrogen area measured in 2 N sulfuric acid is pro(21) C.E.Brickes and H. R . Johnson, Ind. Eng. Chem. Anal. Ed. 17, 400 (1945). (22) N. Bjerrum and L. Zechmeister, Ber. Deut. Chem. Ges., 56,894 (1923). (23) “1962 Supplement,” ASTM Standards, American Society for Testing Materials, Philadelphia, Pa., Part 7, p 259.
CATALYTIC OXIDATIONOF METHANOL
4215
0
5
12.0
X h
1 0
0
-2 X
.-D
-E
8.0
\
s
*a
e
::
. I
e
9 0
-2.0
.* 4.0 ;d
- 1.0
Log (mole fraction of water)
.* 51 3
Figure 3. The effect of water on the amount of hydrogen adsorbed on platinum and on the rate of oxidation of methanol in methanolic solutions: A, amount of adsorbed hydrogen; and 0, rate of oxidation of methanol. 0.0
0
80 Surface area, m*/g.
160
Figure 2. The effect of the area of the catalyst on the rate of oxidation of methanol: 0, area measured in 2 N sulfuric acid; and 0, area measured in 0.03 N hydrochloric acid in methanol.
portional to, but not equal to that measured in the methanolic solution. That the hydrogen area is proportional to the active surface area of the catalyst can be seen in Figure 2. It can be seen that the rate of oxidation of methanol is proportional to the area under the curve in the hydrogen region, whether the area was measured on electrodes immersed in aqueous or methanolic systems. I n order to compare with previous studies in aqueous systems and because of greater uncertainty in drawing the residual in methanolic solutions, all results were reduced to a unit area as measured in aqueous 2 N sulfuric acid. The introduction of small amounts of water into the methanol solution caused a decrease in the measured amount of adsorbed hydrogen on the catalyst (Figure 3), and correspondingly, there was a decrease in the rate of catalytic oxidation of methanol. Although there are slight differences in the two results, it can be concluded that the change in rate of oxidation of methanol upon changing the water content is essentially due to changes in the available area of the catalyst. This point is further emphasized if it is realized that over the range of water concentration running from the minimum to the high methanol concentration end of the solid line, the
rate of oxidation per unit area measured in sulfuric acid changed ,by a factor of 4.75. On the other hand, the rate of oxidation per unit area measured in the methanol solutions was constant at a value of 0.55 X lo-&mole/l. min mcoulomb with an average deviation for the 13 runs of h0.06 X Figure 4 shows the effect of the pressure of oxygen on the rate of oxidation of methanol. The slope of this curve is 0.5, indicating that the reaction is half order with respect to oxygen. As the reaction proceeded and formaldehyde built up in the solution, eventually a steady state was reached in which the rate of oxidation of formaldehyde became equal to the rate of formation of formaldehyde (Figure 5 ) . It can be seen from Figure 5 that the addition of water decreases the steady-state concentration of formaldehyde, while the addition of hydrochloric acid caused an increase in the rate of the reaction and an increase in the steady-state concentration of formaldehyde. This is shown in another form in Figures 6 and 7. Figure 7 was obtained by subtracting the effect of water shown in Figure 6, so that the results are expressed in terms of the effect of hydrochloric acid on a system with a constant water content (mole fraction of 0.028).
Discussion of Results It has been commonly assumed, since the number of sites calculated from oxygen adsorption agree with the number of sites calculated from hydrogen adsorption, that the platinized platinum electrode is completely covered with hydrogen at a hydrogen pressure of 1 Volume 7 1 . Numbm 1.9 December 1967
THOMAS C. FRANKLIN AND MOTOO KAWAMATA
4216
- 1.0
-0.5 0 0.5 Log (partial pressure of oxygen atm).
Figure 4. Initial oxidation rate a3. a function of oxygen partial pressure a t a constant water content (water content, 5-6 mg/ml of solution).
- 1.5
- 1.0
-0.6
Log (mole fraction of water).
Figure 6. The effect of water on the steady-state formaldehyde concentration.
f
0.20
i
g I
:
6 0.10
10
20
Time, min.
Figure 5. Change of formaldehyde concentration during the oxidation of methanol containing different amounts of water and hydrogen chloride. Mole fraction of water, normality of hydrogen chloride, and catalyst surface area (millicoulombs) are as follows, respectively: V, 0.078, 0.24, 58.9; . , 0.040,0.12, 51.0; 0,0.028, 0.03, 42.3; and 0, 0.169, 0.03, 58.9.
atm over an aqueous sulfuric acid solution.24 The fact that the area under the curve in the hydrogen region is larger in methanol solutions than aqueous soluThe Journal of Phyaical Chemistry
1*
0.00
3.5
5.5
7.5
Formaldehyde concn/unit area X a correction for the effect of water. mole/l. mcoulomb X 10s.
Figure 7. The effect of hydrochloric acid on the steady-state formaldehyde concentration.
tions leads one to raise the question as to what is being measured in the area determinations in methanol solu(24) M. Breiter, “Transactions of the Symposium on Electrode Processes,” Philadelphia, Pa., 1959, John Wiley and Sons, Inc., New York, N. Y., 1961, p 307 ff.
CATALYTIC OXIDATION OF METHANOL
tions. There are three possible explanations for the results. (1) The larger area in the nonaqueous media is due to an experimental error. It has been pointed out in the discussion of the method of drawing the residual that there could be an appreciable error in the areas determined in methanolic hydrogen solutions. No major error can be visualized in the determination in sulfuric acid. This is especially true since this area agreed with areas obtained by oxygen coverage method. However, the oxidation rate curve in Figure 3 is almost identical in shape with the hydrogen adsorption curve. The oxidation rate was measured per unit area in sulfuric acid while the hydrogen adsorption curve involved areas measured in methanol. I n other words, the oxidation rate per unit area as measured in the methanol solutions was constant. It is difficult to see how any error in the determination of the amount of hydrogen adsorbed on platinized platinum in methanol or mixed solvents can be reflected in the rate measurements. (2) The larger area in methanol is caused by another electrolytic reaction. The electrochemical oxidation of methanol in the hydrogen potential region has been shown to occur.25 However, the conclusions of this study were that the oxidation proceeded by a catalytic decomposition and that the material oxidized in the hydrogen potential region was hydrogen obtained from the decomposition of methanol. If the electrode were completely covered with hydrogen in 2 N sulfuric acid, it is difficult to visualize more hydrogen on the surface whether the hydrogen comes from dissolved hydrogen or from decomposition. Also, it is again difficult to see an explanation of the data in Figure 3 in terms of this error. (3) The larger area in the methanol is observed because the electrode is not completely covered by hydrogen in sulfuric acid solutions. A competitive equilibrium exists between the solvent and hydrogen. The measured amount of hydrogen is proportional to, but not equal to, the total number of sites. This equilibrium could be shifted to put more hydrogen on the electrode in methanol because the solubility of hydrogen is greater in methanol than in water. It is difficult to reconcile this assumption with the experimental observation that the number of sites measured by oxygen coverage equals the number measured by hydrogen coverage in sulfuric acid. This, however, seems to be the only explanation that agrees with the effect of water on the kinetic data. No matter which explanation is accepted, it seems clear from Figure 3 that water does not enter into the
4217
methanol reaction, but instead changes the catalytic activity of the metal. The kinetic data can be summarized in the following rate law. rate of oxidn of MeOH
a:
(3feOH)O (area) (PoJ”’(HCI)
The first-order dependency of the area of the catalyst is shown in Figure 2, the half-order dependency of the oxygen is shown in Figure 4, and the first-order dependency of the hydrochloric acid is shown in Figure 7. The zero-order dependency of the methanol is probably caused by its high concentration, since in aqueous systems it has been shown that the reaction is first order up to approx 1 M . At approximately 1 M , the order drops towards zero. The above rate law can be interpreted in terms of the following activated complex
Pt-0
H
H
/\
I
+ 0-C-H
\/ H
I
H
This activated complex, in order to agree with the rate law, involves one platinum site, one oxygen atom, one proton, and a methanol molecule. Very probably, a protonated methanol molecule, either on the surface or in the solution, is attached to a surface oxide, PtO.
H+
H
I
/\
CH3-OH
H
I I
+ PtO +PtO +
0-C-H
H
H
\/
This surface oxide can then undergo reaction with a methanol molecule probably in two rapid steps to produce formaldehyde. H
/\ 0-C-HD \/
PtO
+PtOH2
+ OCHa’
H+
CH30H
+ OCH3+
4formaldehyde
+ CHr+OH2
Acknowledgment. We wish to thank the Robert A. Welch Foundation of Houston for their support of this study. (25) 0. A. Petry. B. I. Podlovchenko, A. N. Frumkin, and H. Lal, J. Elcctroanal. Chem., 10, 253 (1965).
Volume 71,Number 13 December 1867
