J . Phys. Chem. 1984,88, 4439-4444
4439
Influence of Adsorbate Interactions on Heterogeneous Reaction Kinetics. Formic Acid Decompositlon on Nickel Jay B. Benziger* and Gregory R. Schoofs Department of Chemical Engineering, Princeton University, Princeton, New Jersey 08544 (Received: February 23, 1984)
The catalytic decomposition of formic acid on Ni catalysts has been studied as a model system to show the influences of adsorbate interactions on heterogeneous reaction kinetics. Formic acid decomposition on a Ni( 1 1 1) surface was studied under ultrahigh-vacuum (UHV) conditions. It was found that when formic acid adsorbs as a dimer it dehydrates to leave CO and formate adsorbed on the surface. Formic acid monomer dehydrogenates on Ni(ll1) to an adsorbed formate. Adsorbate interactions strongly influenced the decomposition kinetics of adsorbed formate. Formate-formate interactions were found to be attractive, reducing the reaction rate with increasing formate coverage. F o r m a t 4 0 interactions were strongly repulsive, causing a significant acceleration of the reaction rate with increasing CO coverage. Steady-state reaction kinetics of formic acid decomposition on a Ni wire were well represented by the UHV kinetics. The attractive formate-formate interactions caused a maximum in the isothermal decomposition rate as a function of formic acid pressure. A significant enhancement of the steady-state formic acid decomposition rate was observed upon addition of small amounts of CO. It is shown that CO acts as a cocatalyst for heterogeneous formic acid decomposition by causing phase separation between adsorbed CO and adsorbed formate on the nickel surface.
Introduction The rates of heterogeneously catalyzed reactions have typically been analyzed in terms of Hougen-Watson type rate expressions.’*2 These expressions take into account the following sequence of steps: (i) adsorption of reactants, (ii) surface reaction, and (iii) desorption of products. In their analysis Hougen and Watson assumed adsorption on an ideal surface where adsorbed species do not interact with one another. Surface concentrations may then be expressed in terms of Langmuir adsorption isotherms and substituted into appropriate rate expressions. Deviations from ideal, Langmuir adsorption isotherms are well-known. Numerous investigators have considered the influence of a nonuniform surface on adsorption and reaction kinetics. Adsorption isotherms may be developed by summing isotherms over all sites?” Temkin developed an isotherm based on a linear variation of adsorption energy with coverage.6 He found that, for bimolecular reactions on such a nonuniform surface, the rate expression closely resembles that for a uniform ~ u r f a c e . ~ Deviations from ideal, Langmuir adsorption also result from adsorbate interactions.*g9 Although adsorbate interactions affect adsorption behavior, the influences of adsorbate interactions on surface reactions has not received much attention. The lack of interest is due to limited data, which have become available only recently. Experiments on well-defined single-crystal surfaces have shown that adsorbates interact routinely. Low-energy electron diffraction has shown that ordered phases almost always result from adsorption of simple molecules on metal surfaces.I0 Large decreases in isosteric heats of adsorption resulting from repulsive
interactions have been Surface-phase condensation due to attractive interactions has been observed for many systems, including hydrogen on Ni(1 lO),I4 and water on ZnO, SnOz, and c r ~ 0 3 . ’ Separate ~ phases for two adsorbents have been reported for CO and O2adsorption on Pt,I6 H2and O2adsorption on Pt,” and C O and NO adsorption on Pt.18 The consequences of adsorbate interactions for surface reactions have been examined in the case of simple d e s o r p t i ~ n . ’ ~ The -~~ influence of adsorbate interactions on temperature-programmed desorption results of simple reactions has also been e ~ a m i n e d . * ~ - ~ ~ However, these studies have not taken full account of the consequences of adsorbate interactions on the entire reaction sequence involving both adsorption and reaction. In this paper, the decomposition of formic acid on nickel was chosen as a test system to examine the effects of adsorbate interactions on the entire reaction sequence noted above. The reaction displays complexities that cannot be accounted for with simple Hougen-Watson rate expressions. The reaction kinetics determined from ultrahigh-vacuum experiments agree with steady-state kinetics at moderate pressures when proper accounting is made for adsorbate interactions and formic acid dimerization.
Experimental Section The decomposition of formic acid on a Ni( 1 1 1) surface was studied by temperature-programmed desorption (TPD) in a stainless steel ultrahigh-vacuum (UHV) system. A constant heating rate of 20 K/s was used in all of the experiments. The Ni( 111) crystal was characterized by low-energy electron diffraction (LEED) and Auger electron spectroscopy ( A B ) . Formic
(1) 0. A. Hougen and K. M. Watson, Ind. Eng. Chem., 35, 529 (1943).
(2) 0. A. Hougen and y.M. Watson, “Chemical Process Principles 111. Kinetics and Catalysis”, Wiley, New York, 1947. (3) A. Clark, ‘The Theory of Adsorption and Catalysis”, Academic Press, New York, 1970. (4) L. A. Rudnitsky and A. M. Alexeyev, J. Catal., 37, 232 (1975). (5) H . Freundlich, “Kapillarchemie. Eine Darstellung der Chemie der Kolloide und verwandter Gebiete”, Akademische verlagsgesellshaft, Leipzig, 1909. (6) M. I. Temkin, Kiner. Katal., 8, I005 (1967). (7) M. Boudart in “Physical Chemistry, An Advanced Treatise”, Vol. VII, H. Eyring, Ed., Academic Press, New York, 1975. (8) T. L. Hill, “An Introduction to Statistical Thermodynamics”, Addison-Wesley, Reading, MA, 1960. (9) K. J. Laidler in “Catalysis”, Vol. I., P. H. Emmett, Ed., Reinhold, New York, 1956. (IO) G. A. Somorjai and H. H. Farrell, Adu. Chem. Phys., 20,215 (1971). (11) J. C. Tracy and P. W. Palmberg, J. Chem. Phys., 51, 4852 (1969). (12) J. C. Tracy, J. Chem. Phys., 56, 2736 (1972). (13) G. Ertl, M. Neumann, and K. M. Streit, Surf.Sci., 64, 393 (1977).
0022-3654/84/2088-4439$01.50/0
(14) K. Christmann, 0. Schober, G. Ertl, and M. Neumann, J. Chem. Phys., 60, 4528 (1974). (15) K. Morishige, S. Kittaka, and T. Morimoto, Surf.Sci., 109, 291 (1981). (16) R. A. Shigeishi and D. A. King, Surf.Sci., 75, L397 (1978). (17) G. B. Fisher, J. L. Gland, and S . J. Schmieg, J. Vac. Sci. Techno/., 20, 518 (1982). (18) R. J. Gorte and L. D. Schmidt, Surf.Sci., 111, 260 (1981). (19) C. Goymour and D. A. King, J. Chem. SOC.,Faraday Trans. I , 69, 749 (1973). (20) D. A. King and M. G. Wells, Proc. R. SOC.London, Ser. A , 339, 245 (1974). (21) D. L. Adams, Surf.Sci., 42, 12 (1974). (22) D. A. King, Surf.Sci., 47, 384 (1975). (23) J. B. Benziger, J. Chem. Soc., Faraday Trans. I , 76, 49 (1980). (24) V. P. Zhdanov, Surf.Sci., 102, L35 (1981). (25) V. P. Zhdanov, Surf, Sci., 111, 63 (1981). (26) V. P. Zhdanov, Surf. Sci., 123, 106 (1982).
0 1984 American Chemical Society
4440
The Journal of Physical Chemistry, Vol. 88, No. 19, 1984
Benziger and Schoofs
TABLE I: Summary of Product Yields from Formic Acid DecomDosition on Ni( 11 1)
+
pdlmcr/ 10-14(co co2); PmOnOmC~ molecules cm-2
I0-I4(adsorbed oxygen product), CO/CO,E molecules cm-2
A
L
i a W a
8.4 9.1
1.2
1.6
0.5
0.7
2.4
V T3
0.03
8.7
0.3
3.7
W K
3.6
In manifold. *Total CO + C 0 2 product. Product ratio.
acid obtained from Alfa Products (95%) was further purified by quadruple distillation prior to its use. Formic acid monomer and dimer were prepared by controlling the temperature and pressure of formic acid in the dosing manifold. The equilibrium constant of formic acid dimerization, K,is 27 K = p m z / p d3 10(10.755-3090/T) torr (1)
tW
1 0
a
t-
V W
a
v)
In (0
a I
I 300
where P, = partial pressure of formic acid monomer (torr), Pd = partial pressure of formic acid dimer (torr), and T = absolute temperature (kelvin). In practice, formic acid monomer was prepared at a pressure of 100 mtorr and a temperature of 300 K, where the ratio of dimer to monomer is 0.03. Formic acid dimer was prepared at a pressure of 100 mtorr and a temperature of 328 K, where the ratio of dimer to monomer is nearly 4. The formic acid was admitted into the vacuum system through a 0.2 mm i.d. tube under conditions of a free jet expansion. The cooling in this expansion should result in further dimerization. We have assumed that the rate of dimerization is small compared to the transit time of the gas jet so that the dimer concentration in the manifold is a good measure of the dimer concentration impinging on the Ni( 111) surface in the vacuum system. Nearly all previously reported UHV experiments of formic acid decomposition on metals have occurred under conditions where the dimer concentration is much higher than the monomer concentration. The possibility of formic acid dimerization was neglected in these studies, however. Experiments of formic acid decomposition on metals at moderate pressures have generally occurred under conditions where formic acid exists only as a monomer. This dichotomy has hindered comparisons of the reaction mechanisms and kinetics observed in the two regimes and is reconciled below. The steady-state decomposition of formic acid monomer on polycrystalline nickel wire was studied in an isothermal differential reactor. Helium was bubbled through formic acid at 20 O C and diluted with additional He to obtain the desired formic acid pressure. Provision was also made to add CO and Hzto this gas stream to obtain other desired gas compositions. The total pressure in the reactor was 1 atm. Nickel wire (Alfa Products, 99.97% purity) was first reduced in flowing H2 at 548 K for 1 h. The reaction gas mixture was then introduced to the reactor and the effluent was run through a Beckman IR-10 infrared spectrometer. Steady-state conversion of the formic acid was determined by the change in absorption of the effluent gas at 1755 cm-I, and reaction rates per unit mass of wire were determined. Measurable rates could be obtained only under conditions where the ratio of dimer to monomer was less than 0.01. Under these conditions the contribution of the vapor-phase dimer may be neglected. Results Ultrahigh-Vucuum Results. Figure 1 shows the TPD spectra following adsorption of formic acid with different concentrations of dimer on a clean Ni( 111) surface at 250 K. The Auger spectra taken after each TPD experiment indicated the presence of oxygen on the surface. Table I summarizes the amount of adsorbed oxygen product and the CO/CO2product ratio from these reactions. The results in Figure 1A represent conditions where the dimer is the primary species impinging on the Ni(ll1) surface. These results are virtually identical with the autocatalytic decomposition (27) A. S. Coolidge, J. Am. Chem. SOC.,50, 2166 (1928).
350
TEMPERATURE
400
450
(OK)
B
t-
z W
a
a 0 2 W K
I-
#
a
t; W a v)
In v)
a
I I 300
350
TEMPERATURE
400
450
(OK)
C
z t-
v
U u 3
a W
w t-
aa
tV W
Ia n 0) v)
a
I I 300
350
TEMPERATURE
400
450
(OK)
Figure 1. Product desorption spectra for formic acid decomposition on N i ( l l 1 ) : (a) manifold P = 100 mtorr, T = 238 K, P,,,,,,/P,,,,,,, = 3.6; (b) manifold P = 200 mtorr, T = 273 K, Pd,mer/Pmonomer = 0.5; (c) manifold P = 100 mtorr, T = 300 K, Pd,mer/Pmonomcr = 0.03.
of formic acid observed on Ni( 110)28329 and Ni(100).30 Hydrogen and C 0 2 desorbed in a very narrow desorption peak at 362 K; subsequently CO desorbed at 450 K by a desorption-limited process. The narrow desorption peak did not vary with surface (28) J. McCarty, J. L. Falconer, and R. J. Madix, J . Catal., 30, 235 (1973). (29) J. L. Falconer and R. J. Madix, Surf. Sci., 46, 476 (1974). (30) J. B. Benziger and R. J. Madix, Surf. Sci., 79, 394 (1979).
The Journal of Physical Chemistry, Vol. 88, No. 19, 1984 4441
Formic Acid Decomposition on Nickel
.O
0
a
1
b
LT W
0
0
1
2
3
4
5
6
7
8
coverage, except at very low coverages where it broadened and shifted to higher temperatures. Water formed from the acid hydrogens also desorbs from formic acid dimer at temperatures lower than those which we could ~ b t a i n . ~ ' The - ~ ~fact that the decomposition of formic acid dimer left less adsorbed oxygen on the surface than formic acid monomer did supports this hypothesis. Figure 1C shows the TPD spectra under conditions where formic acid monomer was the dominant species impinging on the N i ( l l 1 ) surface. Hydrogen and C 0 2 desorbed at 404 K. The CO, desorption peak width was 22 K at half-maximum, typical of that expected for a first-order decomposition. The CO, and H, desorption peaks shifted to lower temperatures and broadened slightly with decreasing surface coverage. A second hydrogen peak was also observed at 340 K. This low-temperature peak is typical of hydrogen desorption following hydrogen adsorption on a clean Ni(ll1) s ~ r f a c e . ' ~The , ~ ~ratio of the areas of the two hydrogen peaks was approximately 1 . Table I indicates that the CO/CO, product ratio from the monomer was substantially less than that observed from the dimer, whereas the monomer left more adsorbed oxygen on the surface than the dimer did. Figure 1B shows the TPD spectra of formic acid decomposition under conditions where roughly equal fractions of monomer and dimer impinged on the surface. In this case, the TPD results indicate a superposition of the monomer and dimer decomposition results. Moderate Pressure Differential Reactor ResuIts. The steady-state rate of formic acid monomer decomposition on a nickel wire was studied as functions of temperature and formic acid pressure. Figure 2 shows that at low pressure the reaction rate is first order in formic acid pressure. At higher pressures the reaction rate becomes zero order in formic acid pressure. These results are typical of a reaction sequence consisting of equilibrium adsorption of reactants followed by a rate-limiting unimolecular surface reaction. The unusual feature found in this study is that the isothermal reaction rate passed through a maximum as a function of formic acid pressure. The rate of formic acid decomposition on nickel wire at 549 K was also measured as a function of CO pressure at a fixed pressure of formic acid. The results shown in Figure 3, a and b, correspond to partial pressures of formic acid in the zero-order and first-order reaction regimes. For the higher formic acid pressure, where the surface coverage should be close to a monolayer, the rate of formic acid decomposition showed a dramatic maximum at a C O pressure of about 2 torr (Figure 3a). At low formic acid pressures the addition of C O had little effect on the rate of decomposition (Figure 3b). We also investigated the influence of H, on the rate of formic acid decomposition. Figure 4 shows that increasing the H2 pressure while holding the temperature and the formic acid partial pressure constant caused a gradual increase in the rate of reaction. (31) J. Lapujoulade and K. S . Neil, J . Chem. Phys., 57, 3535 (1972).
20
CARBON MONOXIDE PRESSURE ( t o r r )
FORMIC ACID PRESSURE (torr)
Figure 2. Rate of formic acid decomposition on N i wire: (a) T = 563 K, (b) T = 548 K, (c) T = 533 K.
15
IO
5
9
Figure 3. Influence of CO on formic acid decomposition at 549 K: (a) PHCooH = 7.1 torr, (b) PHCOOH = 0.8 torr.
I
I
$1
I
P
E! z -
i t , , LT
0
5
IO
,
.
,
.
15
20
25
30
, I 35
40
HYDROGEN PRESSURE (torr)
Figure 4. Influence of PHCooH = 6.9 torr.
H2on formic acid decomposition at 548 K;
Discussion Mechanism of Formic Acid Decomposition. Formic acid decomposition on nickel catalysts has been extensively studied, both at moderate pressures and under UHV conditions. However, a connection between the two regimes has not yet been established. Iglesia and Boudart recently reviewed the literature of formic acid decomposition on nickel and concluded that a connection between the two pressure regimes could not be made because of differences in the proposed reaction mechanisms in the two pressure regime^.^' Results presented here indicate that when the extent of gas-phase formic acid dimerization is taken into account, one can relate the mechanisms observed in the two pressure regimes. Previous TPD experiments of formic acid dimer decomposition on Ni(1 10)28329 and Ni(100)30found a C O / C 0 2 product ratio of approximately 1. In addition, a low-temperature dehydration step occurred in which the acid hydrogens were lost in the formation of water. These results suggested the formation of a formic anhydride surface intermediate, which then decomposed by unusual autocatalytic reaction kinetics. Recent vibrational spectroscopy results have shown that formic acid adsorbed on Ni(1 and R U ( O O ~decomposed )~~ to adsorbed CO and adsorbed formate below room temperature, discounting the existence of a surface anhydride. However, there was no explanation why formic acid decomposed to adsorbed C O below room temperature. The dehydration of formic acid and formation of CO may be understood in terms of the reaction of the adsorbed dimer. In the hydrogen-bonded dimer, a proton transfer occurs, leading to (32) E. Iglesia and M. Boudart, J . Cutuf.,81, 224 (1983). (33) R. J. Madix, J. L. Gland, G. E. Mitchell, and B. A. Sexton, Surf Sci., 125, 481 (1983). (34) L. A. Larsen and J. T. Dickinson, Surf Sci., 84, 17 (1979). (35) N. R. Awry, B. H. Toby, A. B. Anton, and W. H. Weinberg, Surf. Sci., 122, L574 (1982).
4442
The Journal of Physical Chemistry, Vol. 88, No. 19, 1984
an adsorbed formate and a protonated hydroxyl group
Benziger and Schoofs feature of a nonrandom configuration is the quasi-chemical approximation developed by G ~ g g e n h e i m . To ~ ~ model the kinetics of formic acid decomposition, it is necessary to consider adsorpion equilibria for two species, formate and CO. Assuming competitive adsorption on a single type of site with Z nearest neighbors, the quasi-chemical approximation yields the following six equations that describe the surface configuration:
HCOO(ad)
-
H(ad)
2H(ad) CO,(ad)
-+
(1
+ -
- x)CO,(g) CO(ad)
ex~(-W.A/kr)
(2)
~ B B ~ O O / ~= B O Y4 ~ex~(-WBB/kr)
(3)
~ A A ~ O O / ~ A= O 'Y4
that subsequently splits off water while CO and H remain adsorbed on the surface. This mechanism has also been observed on Ru34,35 and Mo36surfaces, though the role of the formic acid dimer was not previously understood. In all cases, the adsorbed formate then decomposed at higher temperature into primarily C 0 2 and H2
" A B ~ O ~ / ( ~ A O ~ B= O )y2
+ CO,(ad)
H2(g) xCO(ad)
+ xO(ad)
CO(g)
The unusual kinetics of the formate decomposition are due to adsorbate interactions, which are discussed in the next section. The results presented in Figure 1C and Table I suggest that formic acid monomer dehydrogenates upon adsorption into adsorbed formate and adsorbed hydrogen. The adsorbed formate then decomposes to C 0 2 and H2 at higher temperatures, as indicated above. This mechanism is consistent with the results of formic acid monomer decomposition on nickel at moderate pressures. Infrared spectroscopy studies identified the decomposition of surface formate as the rate-limiting The proposed mechanism was that formic acid monomer adsorbed dissociatively on nickel to form adsorbed hydrogen and a surface formate, and subsequently the surface formate decomposed to primarily C 0 2 and HZ. The CO/CO, product ratio observed here for formic acid monomer decomposition in the TPD experiments is typical of the CO/CO2 product ratio observed in many previous experiments at moderate pressure^.^^^^^^^ It is important to note that the CO product from formic acid dimer occurs primarily from the dehydration reaction, whereas the CO product from the monomer occurs primarily from the decomposition of the product CO, to CO and adsorbed oxygen. Table I clearly illustrates this point, as the CO/CO2 ratio varies with the extent of dimerization and inversely with the extent of surface oxidation. Reaction Kinetics. As noted above, differences in the proposed reaction mechanisms have inhibited the comparison of formic acid decomposition in UHV studies with experiments at moderate pressures. The perceived "pressure gap" between the two regimes has further hindered comparisons of the reaction kinetics. Results presented here indicate that the reaction kinetics in the two regimes may be related to one another provided one translates between pressure regimes properly and includes the effects of adsorbate interactions. The following model is based on the belief that the appropriate way to view the difference between the two regimes is not a pressure gap, but rather a coverage continuum. In order to model the effect of adsorbate interactions on heterogeneous reaction kinetics, we shall treat the adsorbed molecules as a lattice gas and allow for nearest-neighbor interactions. The most tractable solution to this problem that retains the essential (36) S. L. Miles, S. L. Bernasek, and J. L. Gland, Surf. Sci., 127, 271 (1983). (37) J. K. A. Clarke and A. D. E. Pullin, Trans. Faraday SOC.,56, 534 ( 1960). (38) J. Fahrenfort, L. L. van Reijen, and W. M. H. Sachtler, Z . Elekfrochem., 64, 216 (1960). (39) W. M. H. Sachtler and J. Fahrenfort in "Actes Deuxieme Congres International de Catalyse", Technip, Paris, 1961, p 838. (40) D. K. Walton and F. H. Verhoek, Adu. Catal., 9, 682 (1957). (41) H. S. Inglis and D. Taylor, J . Chem. SOC.A , 2985 (1969). (42) M. S. Platonov and V. I. Tomilov, Zh. Obshch. Khim., 8, 346 (1938).
"00
exp(-WAB/ k r )
(4)
aAA
=
7Z8A
- I/2ffA0
-
%NAB
(5)
aBB
=
Y28B
-
- l/ZaAB
(6)
Y2/2(l
- 8A
-
=
l/ZaBO
- 1/2ffAO -
Y2aBO
(7)
where 4 is the fractional coverage of species i, a , is the fraction of ij site pairs, and W, is the ij nearest-neighbor interaction energy (Wi, > 0 for repulsion). The adsorption isotherms are given implicitly by the expressions
where plo is the standard-state chemical potential of species i at temperature T, and q1 is the partition function of an isolated adsorbed molecule. The rate of surface reaction is given byz4 z
r = k,(
z
c
PA,nA,mBe(nWIAA+mWA,)l~~~A
(10)
n=O m=O
where k, is assumed to have Arrhenius form, and PA,nA,mB is the probability that an A molecule (formate) has n type A nearest neighbors and m type B nearest neighbors. These probabilities are given by a multinomial distribution PA,nA,mB
Z! - ( Z - n - m)!n!m!
(~AA)~(~AB)~(~AO)'-~~ (11) (2au (YAB aAo)Z
+
+
If we assume an ideal gas and neglect rotational and vibrational contributions to the gas-phase partition function, the standard-state chemical potential may be approximated as
wlo = -kT In [ ( 2 ~ M , k T / h ~ ) ~ / ~ k T ]
(1 2)
where Mi = the mass of molecule i. The partition function for the adsorbed molecule, q,, may be approximated as q, = e-AHalkT
(1 3) where AH, = the enthalpy of adsorption. This approximation assumes a tightly bound species with negligible vibrational energy. The adsorption isotherms given implicitly by eq 8 and 9 may be used in conjuction with eq 10 to find the steady-state reaction rate as a function of temperature and pressure. TPD results may be simulated by recognizing that r = -(deA/dt) (14) while assuming surface equilibrium so that eq 2-7 still hold, and using an appropriate equation for the heating rate. Kinetic parameters from UHV experiments could then be used to predict (43) E. A. Guggenheim, "Mixtures", Oxford University Press, London, 1952.
The Journal of Physical Chemistry, Vol. 88, No. 19, 1984 4443
Formic Acid Decomposition on Nickel behavior at higher pressures or vice versa. We chose the former route in this study. The application of the expressions given above to model formic acid decomposition requires the specification of several parameters, namely the enthalpies of adsorption for formic acid, CO, and hydrogen, the activation energy and preexponential factor for the decomposition of an isolated adsorbed formate, and the interaction energies between pairs of molecules. All of these parameters may be found from UHV experiments. In comparing the UHV results to moderate-pressure experiments on nickel wire we will use parameters for Ni( l l l ) and Ni(100) surfaces as they are the thermodynamically most stable planes and are expected to predominate on polycrystalline nickel wires. The isosteric adsorption enthalpy of CO on Ni(100) has been measured to be -122 kJ/mol with repulsive interactions of +2 kJ/mol; l 2 similar results have been obtained for CO on Ni( 11 1).44 The enthalpy of adsorption of hydrogen on Ni( 100) is -96 kJ/mol with repulsive interactions of 1.6 kJ/mol; 14*45 identical results have been obtained for hydrogen on Ni( 11I ) . l 4 The heat of adsorption of formic acid on nickel may be approximated from the heat of formation of nickel formate, which has been determined to be -413 k J / m 0 1 . ~ ~For the reaction HCOOH(g)- HCOO(ad) H(ad)
+
+
the heat of reaction is -105 kJ/mol, which is the value we will use for the enthalpy of adsorption of formic acid on nickel. The formate-formate interactions are slightly attractive. Increasing the initial coverage of formic acid monomer decreased the reaction rate, as indicated by higher decomposition peak temperatures. From the COz peak shift with formate coverage the formate-formate interaction energy was found to be -3 kJ/mol. Adsorbed CO-formate interactions are strongly repulsive, destabilizing the formate and increasing the reaction rate with increasing C O coverage. A repulsive interaction energy of +6 kJ/mol can account for the change in the reaction rate shown in Figure 1. Values of the activation energy and preexponential factor have been determined from heating rate variation TPD experiments to be 101 kJ/mol and lOI5 s-I, respectively, for Ni(l11) and other planes of Ni.2s-30 The underlying physics of the interactions are not clear. Since adsorbed C O and formate have opposing dipole moments, one might expect the formate-formate interactions to be repulsive and the CO-formate interactions to be attractive if both the CO and the formate bond perpendicular to the surface. Since the opposite type of interaction is observed, we suggest that the molecules do not adsorb perpendicular to the surface and/or that factors other than dipole moments contribute to the total interaction energy. The TPD experiments are reproduced quite well with the model accounting for adsorbate interactions. The autocatalytic kinetics are displayed as well as the destabilizing influence of adsorbed C O on the surface formate. The simulated TPD results do not show two separate desorption peaks as the experimental results in Figure 1B do, however. The reason for the discrepancy is that the simulated results assume an equilibrium surface configuration. Because of the nonequilibrium nature of TPD experiments this assumption is probably in error. The simulated results do accurately reflect the peak shift in going from no adsorbed CO (Figure 1C) to a 50/50 mixture of adsorbed CO and formate (Figure 1A), which is the main feature of interest. The model presented above was also used to calculate isothermal reaction rates for formic acid decomposition on nickel at moderate pressures. These results are in excellent qualitative agreement with the experimental results in Figure 2, though quantitatively the pressures at which the maximum rate occurs differ by a factor of 2. We found that by adjusting the adsorption enthalpy of formic acid by about 5 kJ/mol we could obtain quantitative agreement. This adjustment is well within the experimental error, indicating (44)K.Christmann, 0. Schober, and G. Ertl, J . Chem. Phys., 60,4719 (1974). (45)J. Lapujoulade and K. S . Neil, Surf. SA., 35, 288 (1973). (46)G.M. Schwab and E. Schwab-Agallidis, Ber. Dtsch Chem. Ges. B, 76, 1228 (1943).
Y
e * X
a a Z
P
N
La a w 0
CARBON MONOXIDE PRESSURE (torr)
Figure 5. Simulated effect of CO on formic acid decomposition a t 549 K: (a) PHCooH = 7.1 torr, (b) PHCOOH = 0.8 torr.
that the model represents these data quite well. The isothermal reaction rate maximum is an odd feature for a unimolecular decomposition and is due to the attractive interactions between formate intermediates. Attractive interactions increase the magnitude of the enthalpy of adsorption with increasing coverage, causing the surface coverage to increase faster with pressure than in the absence of interactions. The higher coverage increases the number of reaction intermediates, which tends to accelerate the reaction rate. The attractive interactions also enhance the stability of the adsorbed intermediates, which raises the activation energy for reaction and tends to decelerate the reaction rate. The trade-off between these competing effects leads to the maximum in the isothermal reaction rate. Incidentally, the relative magnitude of the rate maximum is a function of the dimensionless interaction energy WAA/kT. As the attractive interactions become stronger, the relative magnitude of the maximum increases. With sufficiently large attractive interactions surface-phase condensation O C C U ~which S , ~ ~causes ~ an extremely prominent rate maximum. The TPD experiments clearly showed the destabilizing effect that CO had on adsorbed formate. The CO caused formates to coalesce into islands to reduce the CO-formate repulsions, and the CO destabilized adsorbed formates at the interface between a CO surface phase and a formate surface phase. In the steady-state experiments at moderate pressures, the CO influences the reaction in the same way if the formic acid pressure is sufficient to assure high formate coverage. The initial addition of C O destabilizes the adsorbed formate and causes the reaction to accelerate; however, as the CO partial pressure increases, CO becomes the dominant species on the surface and the reaction rate decreases. This result is clearly displayed experimentally and by the model in Figures 3a and 5a, respectively. The model predicts the CO partial pressure where the maximum occurs quite accurately. At low formic acid pressure where the formate coverage is low, the effect of CO is much less pronounced. Again this is reflected both experimentally and by the model calculations shown in Figures 3b and 5b, respectively. In performing our model calculations we found that multiple solutions existed for the reaction rate and surface configurations. This indicated phase separation, such that the adsorbed species formed a formate-rich phase and a CO-rich phase. The equilibrium between phases is described by equal chemical potentials for each component in the two phases, and equal spreading pressures for the two phases.47 Using the parameters evaluated from the TPD experiments we were able to construct the ternary diagram for the formate-CO-empty site adsorption system on nickel shown in Figure 6. From Figure 6 it is evident that the results in Figures 3a and 5a represent conditions inside the phase envelope, and hence give rise to the dramatic rate enhancement, ~~
(47)S.Sundaresan, presented at the Annual Meeting of the American Institute of Chemical Engineers, Los Angeles, CA, paper 59B, 1982.
J. Phys. Chem. 1984, 88.4444-4446
4444
an acceleration of the isothermal decomposition rate with decreasing coverage; they also were unable to account for this result. Many previous studies of formic acid on nickel have focused attention on dehydrogenation vs. dehydration as measured by the CO/CO2 product r a t i ~ . ~We ~ ~have ~ ' purposefully not emphasized that issue since the CO/CO2 product ratio can be skewed by secondary reactions. Nickel is known to catalyze C 0 2 decomposition into CO and adsorbed oxygen, C O and CO, hydrogenation, and the water gas shift reaction. Since the reaction kinetics more closely reflect the behavior of the primary reactions, it is our contention that reaction kinetics provide a more meaningful comparison between the UHV and moderate-pressure regimes.
VACANT SITE COVERAGE
co
FORMATE
Figure 6. Phase diagram for formate and CO adsorbed on Ni at 548 K. Tie lines correspond to the following gas-phase compositions: (a) PCO = 0.067 torr, PHcmH = 1.6 torr; (b) Pco = 0.10 torr, P H C ~=H 2.0 torr; (c) Pco = 0.29 torr, P H C m H = 4.0 torr; (d) PCO= 0.62 torr, P H C ~=H 7.1 torr. Plait point at PCo = 0.061 torr, PHCOOH= 1.5 torr.
whereas Figures 3b and 5b are in the single-phase region and little effect on the reaction rate is noted. The formate-formate interactions are not sufficiently attractive to induce surface-phase condensation in the two-component formate-empty site system; however, addition of a third species such as CO can induce phase separation. The enhancement of the reaction rate by the addition of hydrogen is due to a reduction in stability of the formate intermediates. Addition of hydrogen causes a slight reduction in the number of formate-formate nearest-neighbor pairs, which accelerates the rate. The hydrogen-formate interactions appear to be negligible, so that no phase separation occurs and no dramatic increase in the reaction rate is observed. Using the heat of adsorption of hydrogen noted above and no hydrogen-formate interactions, the model is able to reproduce the experimental results shown in Figure 4. There are several anomalous results in previous studies of formic acid decomposition on nickel at moderate pressures that can now be explained in terms of adsorbate interactions. Walton and Verhoek observed that as C O was added to formic acid vapor, the isothermal reaction rate initially increased and then decreased.40 They were unable to account for this result, though it is now apparent that it results from destabilization of the formate intermediates by adsorbed CO. It seems that adsorbed CO probably played an important role in the supercatalytic activity of nickel wires reported by Duell and Robertson4* and also by W i l l h ~ f t .Evidence ~~ for attractive interactions between formate intermediates is presented by Giner and R i s ~ m a nwho , ~ ~observed (48) M. J. Duell and A. J. B. Robertson, Trans. Faraday SOC.,57, 1416 (1961). (49) E. M. A. Willhoft, J. Chem. SOC.,Chem. Commun., 146 (1968).
Conclusions The results presented in this paper have elucidated the mechanism of formic acid decomposition observed in many TPD experiments. Formic acid dimers formed at low temperatures dehydrate on nickel surfaces, leaving adsorbed CO and formate on the surface. Formic acid monomers dehydrogenate to form formate intermediates. The TPD experiments indicated that adsorbate interactions caused the unusual kinetics. It was found that formate-formate interactions were attractive, while COformate interactions were strongly repulsive. Kinetic parameters determined from the TPD experiments were used to predict reaction kinetics at moderate pressurs successfully. Several important features that could be represented by the model include (1) a maximum in the isothermal reaction rate as a function of formic acid pressure, (2) a dramatic maximum in the isothermal reaction rate with the addition of carbon monoxide, and (3) enhancement of the isothermal rate of reaction with the addition of hydrogen. None of these results can be accounted for by simple Hougen-Watson type rate expressions, and they can be predicted only by considering adsorbate interactions. The enhancement of the reaction rate by the addition of a nonreactive component is an important finding. In this work CO had a dramatic effect on the rate of formate decomposition by destabilizing the formate. It was found the surface phases could be formed by adding CO, a nonreactive component. The formation of surface phases greatly affected the reaction rate. In this way CO acted as a cocatalyst, similar to a coenzyme in a biological system. These results suggest the possibility of adding nonreactive components to facilitate other heterogeneous chemical reactions. In cases of bimolecular reactions it may be possible to add a third component that would push the reactive components closer together to facilitate reaction. This concept clearly needs further investigation. Acknowledgment. We thank the Air Force Office of Scientific Research for their financial support of this work. Registry No. Ni, 7440-02-0; CO, 630-08-0; formic acid, 64-18-6; H,, 1333-74-0. (50) J. Giner and E. Rissman, J . Catal., 9, 115 (1967). (51) P. Mars, J. J. F. Scholten, and P. Zwietering, Adv. Caral., 14, 35 ( 1963).
COMMENTS Concernlng the Alleged Efficiency of Photoaquatlon in the Ultraviolet Photolysls of Bromo(pentammlne)cobalt( I I I ) S i c In a recent paper in this journal, Kirk et al.' have reported that ultraviolet irradiations of C O ( N H ~ ) ~ Bproduce ~ ~ + Co0022-3654/84/2088-4444$01.50/0
(NH,),OH?+ and Co2+in a wavelength-independent0.46: 1 ratio. This stands in contradiction to our earlier report^^,^ that we were unable to detect photoaquation products, and the contradiction has led us to reinvestigate the 254-nm photolysis of C O ( N H ~ ) ~ B ~ ~ + in acidic aqueous solution^.^ Owing to the uncertainties intrinsic to spectrophotometric determinations of small amounts of photolysis products in the 0 1984 American Chemical Society
