The adsorption and reaction of simple molecules on metal surfaces

Surface Science 89 (1979) 540-553 © North-Holland Publishing Company

THE ADgOPO' riON AND REACTION OF SIMPLE MOLECULES ON METAL SURFACFS Robert J. MADIX

Department of Chemical Engineering, Stanford University, Stanford, California 94305, USA Receivea 3 May 1979

The ~es,,Rs of experiments are discussed which exemplify surface science studies of heterogeneou~ reactivity. For simplicity throughout, the reactions of formic acid, H-C ~D , are used \OH as a mede! system. Direct comparison of reactions of formic acid on metal surfaces at high and H

H

low pressures show striking similarities. The existence of the intermediate

\ C / is demon0 / \O strated at 200 K and the relevance of such results to higher temperature reaction studies is discussed. Finally the manner in which temperature programmed reaction spectroscopy and electron spectroseopies compliment one another is discussed briefly.

1. Introduction Some of the previous speakers in this Cengress have discussed the importance of surface science in catalysis. There are a few points ihat I would like briefly to discuss in co~mection with this topic. In studying the interactions of molecules with macroscopic single crystal metal surfaces we are hoping to model the conditions existent on support:d metal catalysts, at least as far as these studies relate to catalysis. Generally, in metallic catalysts small metal particles are supported on high surface area inorganic materials or minerals such as alumina, silica or faujacites. These particles, which.may range from a few A to several hundred A in dimension are typically located along tortuous pores or channels in the inorganic matrix. Our idealization of these surfaces by the use of extended single crystal surfaces is an attempt to gain further understanding about the influence of surface structure and composition on reactiv~iy via the better surface definition realizable through surface science, but it must be realized that only in a few cases are metallic materials of macroscopic dimensions used for catalytic application. In general, the surface catalyzes the reaction under consideration by providing 540

K.a. Maatx / Adsorption o f simple molecules on metal surfaces

541

an alternative sequence of molecular bond rearrangements to that available if the reaction were to occur in the gas phase alone. The main point to be made here is that the chemical nature of the reaction intermediates and their bonding to the surface is crucial to the understanding of the reaction pathways available on the surface. As an example of this consider the oxidation of methanol on a silver surface to yield formaldehyde 2 CHaOD + ~ 02 --* 2 H2CO + H2 + D 2 0 . This reaction occurs via the steps [ 1]

O2(g) --* 2 0 a

dissociative adsorption of 0 2 ,

CH3OD(g) ~ CH30(a) + D(a)

does not occur,

CH3OD(g) + O(a ) -* CH30(a) + OD(a)

dissociative adsorption of CH3OD,

CH3OD(g) + OD(a ) -'* CH30(a ) + D20(g)

dissociative adsorption of CH3OD,

CH3Ofa) ~ H2CO(g) + H(a)

surface decomposition of CH30,

2 H(a) ~ H2(g)

atom recombination.

This sequence of steps clearly delineates which bonds are broken and formed in the molecular transformation (the subscript (a) denotes an adsorbed species). The two salient features here are: (1) the dissociative adsorption of CH3OD requires surface oxygen to be present, and (2) the formation of H2CO is governed by the decomposition of a relatively stable surface species, CH30(a). The basic understanding of both of these phenomena lies within the scope of surface science. Much has also been said of the so-called "pressure gap" which exists between studies of surface reactivity under ultrahigh vacuum conditions and catalytic studies above one atmosphere. It is important to clarify what different effects may occur under these extreme conditions; this "problem" is probably overstated. In general, the rate of conversion of molecular species is a function of the concentration of the species on the surface and the temperature Rmolecular conversion = f(ol, 02, 03, ..., o n, 7'). The surface concentrations, On, are indeed affected by the gas pressure, but, nonetheless, the reactivity characteristics depend on the surface concentrations. It is important to note that with "low pressure" methods surface coverages from 10 -3 to 1.0 monolayer can be effected. Thus, as long as the actual state of the surface being studied simulates that of the surface under conditions characteristic of the ,.,l~.,aLu~ t . a t a t y n t w t t J t respect to composition and structure, rate studies made employing UHV techniques should simulate the real conditions adequalely and help to provide the scientific basis for heterogeneous catalysis. The so-called pressure gap is, in reality, a coverage continuum. At very high pressures, of course, compression layers may form which strongly

542

R.J. "ffadix / Adsorption o f simple molecules on metal surfaces

pert:':b the bonding t,f adsorbed species to the surface. Results presented in this meeting sho~ frequency shifts of 100 cm -~ for C=---Ostretching frequencies as coverage increases to fuli compression. This shift represents a change in the force cons,'ant of the C=O bond of about 10%. It remains to be answered whether such changes are of much significance in terms o f the reactivity of the C'~O bond. There should then be a class of reactions, however, which involves species whose binding energi,es are much larger than the lateral interaction energies and which are affected negligitqy by high coverages. If on the other hand the binding energy is nearly equal to the interaction energies, significant changes in qualitative behavior may occur betweer, 2¢,w and high surface coverages. In these cases an effect that must be studied ~:.; the two-dimensional phase behavior of mixed adsorbate systems. Phase segrega" ion of CO and O, for example, would have marked effects on the rate of CO oxidation. This effect will obviously be concentration dependent. In ,:rder to reach high surface concentrations at the pressures employed in UHV experiments, temperatures lower than those used at higher pressures must be utilized. The use of lower temperatures may be a more important consideration for comparison with catalytic conditions lhan the use of lower pressures, since at low temperatures reactions having the lowest activation energies will be favored. At high temperatures, the rates of reactions tend to equalize, and entropy factors become more important. On the other hand, the use of low temperatures allows us to identify "short-lived" intermediates which may be of importance in catalytic reactions at higher temperatures. For example, the surface lifetime of a species, r = r o e ~/Rr. If 7"0 = 10 -la see and E = 60 kJ mol -~ , 7"= 10 -7 secat 500 K, a n d r = 10 sec at 200 K. At 200 K this species is readily observable with TPRS (temperature programmed reaction spectroscopy), whereas at 500 K it is a "transient" or short-lived intermediate. In summary, then, the techniques of surface s'ience can be applied to problems of significance in heterogeneous catalysis. One should, of course, not overstate their relevance, but as a guide to the development of the scientific basis of the subject they will prove most useful. In some cases, specific knowledge can be gleaned which bears directly on catalytic reactivity. Examples of some of the points raised here are presented below.

2. Experimental ".,,,e,,.'y for our s~uu~es is shown in fig 1. it is a conventional UHV system equipped with LEED, XPS, UPS, AES (all using a double-pass CMA), and a quadrupole mass spectrometer. The state of the surface is prepared by standard argon bombardment, annealing, etc. procedures. If desired, structural overlayers of carbon, oxygen, sulfur or potassium can be deposited with C2H4, 02, H:S or K + sources, respectivly, located around the periphery of the chamber. Reactant gas can be adsor0ed onto the surface down to 150 K, and the sample can

R.J. Madix /Adsorption of simple molecules on metal surfaces

ION BOMBARDMENT GUN

543

,/x ..~/'/

X- RAY SOURCE / J ~

QUADRUPOLEMAsS ,a. A " OPTICS !L / SPECTROMETER~ .

DOSING SYRINGES '

E

-

Fig. 1. UHV apparatus showing various appendages.

be linearly heated at rates of 0.5 to 50 K s -1 . With the sample positioned in front of the ionizer of the mass spectrometer, species evolved from the surface in 10 -3 monolayer quantities are easily observed.

3. Results and discussion For the sake of simplicity illustrations of the application of surface science to catalytic reaction studies will be taken from our work on the decomposition of formic acid. The interested reader may wish to consult more extensive reviews of the subject [2]. This simple molecule, 0 II H-C \OH decomposes to form H~, CO2, H 2 0 and CO, primarily via the formation of an Hi adsorbed tC\o species (with the exception of nickel surfaces near room O temperature). A typical temperature programmed reaction spectrum showing the decomposition rate versus temperature is shown in fig. 2. Formic acid (a mixture of DCOOH and HCOOD) was aasorbed on Cu(110) at 200 K. Clearly, there are two widely separated reaction events shown by the 273 K hydrogen peaks and the 473 K COs and hydrogen peaks, renpectively. These two channels result from the

R.J. Madix

544

/Adsorption o f simple molecules o~,,~metal surfaces TEMPERATURE (*C)

-I00

- 50

0

i

I

50

moo

150

200

250

t

t

I

I

I

'

J

(4. Z. t~ ol

PROD/HCOOD ( 2 0 0 K)

E t.l.I I'-'ILl

0 11.I-. o ~,~ v?

D2 (x2) HDlx21 173

1

I

223

273

I

""--I

....

I,

323 373 423 TEMPERATURE ( K )

I,

I

473

523

Fig. 2. TPRS for the products evolved from C u ( l l 0 ) subsequent to HCOOD adsorption at 200 K. Some HCOOH was present.

molecular rearrangements [3] HCOOD(a ) -~ HCOO(a ) + D(a) ,

2 D(a ) ~ D2(g ) ,

273 K ,

2 H(a) ~ H2(g) ,

473 K .

and HCOO(a) ~ H(a) + CO2(g ) ,

The use of pure HCOOD readily allows positive identification of these two processes, as only D2 was evolved at 273 K. The resolution inherent in the technique is evident from this result. A quantitative evaluation of the amounts of H2 and CO2 evolved at 473 K verified the H2/CO2 stoichiometry of 1/2 expected from decomposition of HCOO(a). Fig. 3 shows that the rates of H2 and CO2 evolution were identical (when normalized to the peak maximum) for all temperatures in the peak. This signifies that they were formed simultaneously from a common ratelimiting step slow

HCOOad s ~

H a + CO2(g ) ,

to!lowed by the rapid recombination of the hydrogen atoms to form H2 in a step too fast to be seen as a difference in peak shape. This peak position can now be taken as a characteristic "fingerprint" of the presence of the presence of HCOO(a) on Cu(110); it will be extremely useful to ascertain if this intermediate is formed from other reactants on the same surface.

R.J. Madix / ~_dsorption o f simple molecules on metal surfaces

373

398 ,.

423

I

I

TEMPERATURE ( K ) 448 473 498 523 I

I

I

545

548

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I

1.0 0.

z z tlJ n,. ILl N ..J cir. o z

i

0.8

0.6-

0.40.20

I I

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1

,

125

~ _

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I

,I

, I

,

,,.I

....

150

175 200 225 250 275 300 TEMPERATURE (*C~ Fig. 3. Normalized product spectra for H2/HCOOD, CO2/HCOOD(200).

In this regard, fig. 4 shows the TPRS resulting subsequent to H2CO adsorption on a Cu(110) surface which was predosed with isotopic molecular oxygen to form adsorbed =sO. Two reaction pathways for the formation of I-t2 are immediately evident. The 8-state at 470 K (CO2 evolution also observed) is identical to the result

I" .......

i

-I

"

I

I'

t

I ....

I

"1

I

HzlH2CI6 0

I

""

1

/"~

if)

a I

250

500

350

I

~

I

400

I

I

450

I

i

l

500

TEMPERATURE ( K} Fig;4. Hydrogen product spectrum subsequent to H2CO adsorption on preoxygenated Cu(110). H2 desorption-limited and HCOO surface-reaction-limited peaks are present at 350 and 475 K, respectively.

546

R.Z Madix / Adsorption of simple molecules on metal surfaces

of HCOO(a ) decomposition, as noted above. Obviously, the overall reaction H2CO + i aOa,ts ~ HCG 1SOars + Hads occurred, iollowed by decomposition of the formate upon heating. The adsorbed hydrogen formed in this reaction desorbed near 325 K in a desorption limited step which is clearly shown by the downward shift in the peak temperature with incre~Lsing amounts of H2CO adsorbed (curves (a)-(e)) [4]. Unfortunately, it is not po~il~le to determine with certainty if the species H

fl

\C / 0 / \,;~ \Cu / actuai~.y ~tably exists on the surface from these measurements, since the low temperature Hz peak is desorption-limited; i.e., only direct evidence for the existence of the stable species H I kCu , /

H I Cu

was observed. The presense of the intermediate H2CO2 could 0nly be inferred as a transient species. Nonetheless a comparison of figs. 2 and 3 is clear evidence that HCOO(a) formed in both cases. Similar evidence is provided by UPS, and a comparison of the results serves to show the comparative power of the two techniques for the identification of intermediates. The UPS spectra for HCOO(a) (or DCOO(a)) on Cu(110) for the cases DCOOH(g) ~ DCOO(a ) + ! H2(g) and ' H2CO(g) + O(a) -~ HCOO(a) + 21H2(g) are shown in fig. 5 [5]. The curves obtained at 400 K show essentially identical spectra, indicating that the same species results from both reactions. Without TPRS, however, the actual identity of this intermediate would obviously remain a mystery, as a comparison of the two spectra for DCOOH shows. It is interesting to compare the results of studies of the formic acid decomposition on a number of metals l:rLd~~UHV conditions [3-10] with those from high pressure conditions [11 ], ~.s shown in fig. 6. In or&:r to do this the temperatures necessary to produce equa~ rates of formic acid decomposition on metal catalysts, TR, were plotted as the c,pen squares against the heat of formation of the corresponding metal formate [ 11 ], and against the estimated M - O surface bond energy. The peak temperatures, Tp for CO2 peak from the TPRS results at low pressures are plotted as the dark circles. The strong similarity of the results indicate that pres-

R.J. Madix /Adsorption o f simple molecules on metal surfaces

547

150K

I -DC:OH,4OOK *0 H2CO,4OOK

-H2COJO,: CLEAN

i

0

5

i

!

IO E. (e,/)

15

Fig. 5. Hell UPS spectra for formate formation on Cu(110) from DCOOH and H2CO + O(a ), respectively. The two 400 K spectra show nearly identical features.

200 @Pt(tlO)

I"1 HIGH PRESSURE, SUPPORTED METe.. @LOW PRESSURE, SINGLE CRYSTALS

300

400

/

i-IRu\

Cu(llO}O

g

L -~)

~l-INi

~

500

60

ONi(llO)

®a~mo) e~u(ioro) cur..1~cuI~O'o)

i

80

t

90

Ft(lOO)

1

I00

,, L,

120

AHf Fig. 6. Comparative high/low pressure results for the decomposition of formic acid (see text).

R.J. Madix / Adsorption of simple molecules on metal s u r f a c e s

548

I I I I I I I

"

!

, !

,, ,, I ! ! I

I.__ i

0

I

1

1

5

I0

15

I

I

•

I

I

0 5 I0 TIME(s) Fig. 7. "E,~e rate of water e,,.~iution at (a) 246 K and (b) 335 K following the sudden exposure of preox?'genated Ag(i 1O) to HCOOH.

sure, per 3e, has no particular effect on the rate in this case. For this particular reaction the high pressure rate may well be determined by the rate of decomposition of the surface formate, and the results of the UHV experiments can be used to account for the higher pressure behavior. Another interesting example of surface reactivity is illustrated by the interaction of formic acid with Ag(110) surfaces which as noted above are rather inert by themselves. When clean Ag(110) is exposed to formic acid, no dissociative adsorption occurs. When the surface is predosed to form submonolayers (10 -a to 5 X 10 -2 ML) of atomic oxygen, however, the acid adsorbs with the concq~ ~itant release of water. This evolution can be observed directly by suddenly rotatin£ he preoxygenated surface into a steady beam ef DCOOH and observing the water evolved with the mass spectrometer due to the reaction [5] 2 DCOOH(g) • O(a ) -~ 2 DCOO(a) + H20(g). Fig. 7 shows the rate of water evolutio~l with time after turning the sample into the beam. The characteristic time for water formation decreased with increadng temperature, as the preadsorbed oxygen was selectively reacted to form water and the formate species. The amount of formate adsorbed per surface oxygen was quantitatively determined by measurement of the relative amounts of CO2 and H2 desorbed for a given oxygen dose. As 'shown in Fig. 8, two fomlate species form per surface oxygen atom. Results similar to those discussed above for Cu(110) were also obtained with (Ag(110) for the reaction H2CO(g) + 180(a ) ~ HCOO(a) + H(a ). in • his "'~"" ~. . . . . . . . . . evidence for the intermediate H

H \C /

0 /

\0

\Ag /

R.J. Madix ~Adsorption of simple molecules on metal surfaces 6

I

I

549

/i

/

0

/

/

/

'sLo =2 Q

O

/ /

IJ

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C02~

3

sJ

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/

/ E

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6]3--

/

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H2

0/

0~./,. I

I

I

1

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2

OXYGEN

PRECOVERAGEIatoms

3 c m - 2 , 1 0 "13]

"

Fig. 8. The amount of CO2 and H2 formed in the high temperature (HCOO) TPRS peaks plotted against oxygen predose.

1

v

l

i

~ - - ~ .... t " r

t

PRODUCT DISTRIBUTION HzCO/Ag(IIO)--O

I

Z"

_o Q. (/3 UJ a

200

2

TEMPERATURE(K)

400

Fig. 9. The TPRS spectrum for the products [brined following H2CO exposure to the preoxygenated Ag(110) surface. Two reaction-limited H2 peaks are observed at 240 and 390 K, respectively, indicating both H2CO2(a) and HCOO(a ) intermediates.

550.

R.J. Madix / Adsorption of simple molecules on metal surfaces

was clearly obtained. As shown in fig. 9, hydrogen was evolved at 240 and 400 K. The high temperature pc,& was due to formate decomposition. The low temperature peak exhibited first order kinetics and was therefore not rate limited by hydrogen atom recombination. In fact, hydrogen atom desorption would occur below 200 K, and the peak at 240 K must be attributed to the decomposition reaction

F

H I

I-{

H I

"C / O/

+ Ag

N O .+ o / C \ o

Ag Ag Ag

--"Ag/

followed :~y rapid hydrogen atom recombination. This intermediate is an example of one that would be extremely short-lived at 500 K for the high temperature oxidation of H2CO to CO2 and H20 and which was readily detected under these experimental conditions. The reactive properties that oxygen adsorbed on silver exhibits are unique. Apparently oxygen on this surface is such a strong electron donor that it selectively attackes electron deficient centers of the adsorbing molecules. Examples of this behavior are shown below (R designates CHa- or H - ) [ 1,5,17]

R

ON Ag

H \CII/ + A g /

H\C/R 0 / "0

-*

0

Ag

R

Ag

R

I

C

+

O-H

Ag"

Ag

H

I

/0\t8

( 1)

/

-+ 0 + ~ o XAg

Ag

(2)

Ag

,c_ O:C=O

.t..

+

Ag

(3) "~Ag

Ag Ag

0

H

II HCOCH3 +

18o

O Ag

II

RCOH+

O

Ag /

NAg

I

/C \

--9,

Ag/ O

CH3

I

+O

018 Ag

I

Ag

R

H

/C \

+O i Ag

I

0 O Ag Ag Ag

I

(4)

Ag

(5)

R.J. Madix /Adsorption o f simple molecules on metal surfaces

551

H H I I H H 0 -+ 0 0 ! I \ 0 / + A g / NAg AgAg

(6)

In some cases the use of other spectroscopic tools may be needed to supplant TPRS. The decomposition of f o ~ i c acid on Fe(100) illustrates the combined use of TPRS and XPS to elucidate a reaction mechanism. The TPRS shown in fig. 10 indicates that H2 and CO desorbed in desorption.limited peaks at 350 and 800 K, respectively. The CO, H= and CO= peaks at 490 K were indicative of HCOO decomposition, but the desorption spectrum alone left some doubt as to whether CO formed at a lower temperature, at first inspection. Examination of the progressive changes in the C(1 s) spectrum following adsorption of HCOOH at 200 K and progressive heating were extremely helpful in resolving this question. The C(1 s) binding energy for both molecularly bound HCOOH and HCOO was nearly identical and significantly higher than that expected for adsorbed CO as shown in fig. 11. The O(ls) spectrum, fig. 12, also shows conversion of two forms of oxygen to one between 200 and 360 K with formation of O(a) above 500 K. The inescapable bonclusion from this information is that HCOOH dissociates to HCOO between 200 and 360 K with the HCOO decomposing to H2, CO2 and CO (both molecular and dissociative) near 500 K. Though, TPRS itself would actually have led to the same conclusion, the XPS results lend substantial support to the result.

Formic Ac,d Decon3Dosi|ion on

Fe(lO0)

5=10"9A

+

JL_ co

U IJd O3

&

200

I

l

__

A

~

400 600 TEMPERATURE {g ~

ml

"

~0

Fig. 10. The TPRS spectrum of HCOOH on Fe(100). The peaks at 490 K are surface-reactionlimited peaks.

552

R.Y. MrMix / Adsorption of simple molecules on metal retraces

C(Is) ×PS o! HCOOH

e(

Adsorbed on

Q

m

I'l

b-

.Jk il 2 70

,, I

I ,, , I 280 BINDING ENERGY

* 290 (eV)

'"

Fig. ] ]. The C(]s) XP$ for HCOOH adsorbed at 200 K and heated to (a) 200 K, (b) 360 K, (c) 500 K, and (d) prior to adsorption. The spectrum clearly shows that molecular CO does not form at 360 E where H2 is evolved in forming the formate.

O(Is) XPS of HCOOH Adsocbed on ~Fe{lO0)

,

520

ml

/

~

BiNDiNG

~

530,

~

T

ENERGY

~ . . . .

~

.

540

__

I

(eV)

Fig. 12. The O(ls) XPS for HCOOH adsorbed at 200 K and heated to the same temperatures in fig. 11. The desorption and conversion of HCOOH(a ) is evident from (a) and (b).

R.J. Madix /Adsorption o f simple molecules on metal surfaces

553

Acknowledgment The author gratefully acknowledges the support of the National Science F o u n d a tion (NSF GK 12964). %

References [ll I. Wachs and RJ. Madix, Surface Sci. 76 (1978) 531. [2] R.J. Madix and J. Benziger, Ann. Rev. Phys. Chem. 29 (1978) 285. [31 D.H.S. Ying and R.J. Madix, J. Catalysis, to be published. [4] I.E. Wachs and R.J. Madix, Surface Sci. 84 (1979) 375. [51 M. Bowker, M. Barteau and R.J. Madix, to be published. [6l D.H.S. Ying and R.J. Madix, J. Inorg. Chem. 17 (1978) 1103. [7] J. Benziger and R.J. Madix, to be published. [81 T.J. Dickinson, to be published. [9] N. Abbas and R.J. Madix, to be published. [10] J. Falconer and R.J. Madix, Surface Sci. 46 (1974) 473. [11] J. Fahrenfort, in: The Mechanism of Heterogeneous Catalysis, Ed. J.H. de Beer (Elsevier, Amsterdam, 1960) p.23. [121 M. Barteau and R J . Madix, to be published.