the energy of adsorbed hydroxyl on pt(111) by microcalorimetrydepts.washington.edu/campbelc/pdf/oh...

Published: May 20, 2011

r 2011 American Chemical Society 11586 dx.doi.org/10.1021/jp201632t | J. Phys. Chem. C 2011, 115, 11586–11594

ARTICLE

pubs.acs.org/JPCC

The Energy of Adsorbed Hydroxyl on Pt(111) by MicrocalorimetryWanda Lew,† Matthew C. Crowe,† Eric Karp, Ole Lytken, Jason A. Farmer, Líney �Arnad�ottir,Carolyn Schoenbaum, and Charles T. Campbell*

Department of Chemistry, University of Washington, Box 351700, Seattle Washington 98195-1700, United States

’ INTRODUCTION

Surface hydroxyl groups are reaction intermediates in a largenumber of important catalytic and electrocatalytic reactions onPt and other late transitionmetal surfaces. These reactions includethe combustion and partial oxidation reactions of practically anyorganic molecule, the water-gas shift reaction, steam reformingof oxygenates or other organic molecules, the three-way catalystfor automotive exhaust cleanup, fuel cells for generating electricityfrom H2 or any other hydrogen-containing fuel, and watersplitting. Platinum has been identified as the most active catalystfor many of these reactions,1 and therefore, the enthalpy offormation of adsorbedOHon Pt surfaces is particularly important.

Because of its importance, much work has been devoted tounderstanding the interaction between water on clean and oxygenprecovered transition metal surfaces, as has been summarizedrecently in a review.2 Water adsorbs intact on clean Pt(111), anddesorption from the multilayer and monolayer happens at ∼150and 165K, respectively. In the temperature range from 130 to 185K,water is known to react with preadsorbed oxygen adatoms onPt(111) to produce adsorbed OH groups in a coadsorbed,hydrogen-bonded H2O 3 3 3OH complex.3�6 This complex canbe stable up to 205 K, with its stability dependent on the H2O:OHsurface stoichiometry. At∼205 K, water desorbs from the surface,restoring the initial oxygen-covered surface. Carefulmeasurementsby Clay et al.5 of the amount of water adsorbed per oxygen adatomover a wide range of oxygen precoverages at 163 K (where wateron oxygen-free Pt(111) is no longer stable) indicate that the com-plex formation can be described by the following reaction:

3H2Og þOad f 2ðH2O 3 3 3OHÞad ð1Þ

Consistent with this is the observation that themixedH2O 3 3 3OHphase is most stable to decomposition when formed by a reactionratio of 3H2O:1Oad to give a total local coverage of 0.67 ML Oatoms.5�7 At saturation, this corresponds to 1/2MLH2O reactingwith 1/6 ML of oxygen adatoms, but the same ratio at lowercoverages gives islands of very similar stability.8 Earlier experi-ments of the 18O fraction (0.31) in the desorbing water evolvedupon thermal decomposition of the surface hydroxyl complexformed using predosed 18O adatoms by Creighton and Whitesuggested a different reaction stoichiometry, for which the follow-ing reaction was proposed:4

2H2Og þOad f 3OHad þHad ð2Þ

Later experiments byBedurftig et al.7 repeated this isotope fractionmeasurement and obtained 0.29 instead, based upon which theyagreed with reaction 2 above proposed by Creighton and White.However, 0.29 actually agrees just as well with the fraction 0.25expected for reaction 1 as with the value 0.333 expected for reaction2, and both are probably within its experimental error. Whetherthe stable species formed is purely OHad or the (H2O 3 3 3OH)adcomplex, there seems to be general agreement from low energyelectron diffraction (LEED) and scanning tunneling micro-scopy (STM) measurements that it forms islands of either a(√3�√

3)R30� structure or a (3�3) structure, in either casewith a local coverage of 2/3 ML of O atoms.2,5�7,9 The

Received: February 18, 2011Revised: April 23, 2011

ABSTRACT:The energy of adsorbed hydroxyl on Pt(111) wasmeasured by dosing D2O gas onto oxygen precovered Pt(111)at 150 K while following the heat of reaction with single-crystaladsorption calorimetry. The adsorption of D2O on oxygenprecovered Pt(111) is known to produce surface OD(hydroxyl) coadsorbed with molecular D2O in a well-definedstructure. The heat of reaction and sticking probability of D2Oon Pt(111) were measured as a function of oxygen precoverageand D2O dose. With 0.25 monolayers (ML) of oxygen atoms, the differential heat of adsorption is nearly constant at 61.3 kJ/mol forthe first ∼1/3 ML but drops to 57.9 kJ/mol by 0.50 ML and 50.5 kJ/mol by saturation (0.62 ML). Similar experiments with Oad

precoverages of 0.18 and 0.07ML gave lower saturation D2O coverages (0.55 and 0.22ML, respectively) and lower heats of reactionby∼3.4 and∼4.6 kJ/mol, respectively, except at very low D2O coverage where step sites may play a role. From the integral heat ofD2O adsorption, the standard enthalpy of formation of the (D2O 3 3 3OD)ad complex is estimated to be�527 kJ/mol. Assuming thatthe adsorbedD2Omolecules in this (D2O 3 3 3OD)ad complex have the same enthalpy as when adsorbed as a pureD2O film onPt(111),the enthalpy of formation of adsorbed OD is found to be �226 kJ/mol, and the Pt�OD bond enthalpy to be 263�274 kJ/mol.From these values, we estimate that the reaction D2OadfODadþDad is endothermic by 29�39 kJ/mol. At D2O coverages below0.4 ML, the sticking probability of D2O on Pt(111) at 150 K is ∼0.9 for oxygen precoverages between 0.14 and 0.25 ML.

11587 dx.doi.org/10.1021/jp201632t |J. Phys. Chem. C 2011, 115, 11586–11594

The Journal of Physical Chemistry C ARTICLE

(H2O 3 3 3OH)ad product via reaction (1) seems to explain the priorexperimental data2�4,6,7,9�13 and theoretical calculations8,14�19

for this system much better than reaction 2. The (3�3)structure is more stable to decomposition than the(√3�√

3)R30� structure.7 Density functional theory (DFT)calculations by Karlberg et al.8 based on reaction 1 also showthat the (3�3) structure is marginally (1.8 kJ/mol H2O dosed)lower in energy than the (

√3�√

3)R30� structure. However,these two structures are very closely related, as shown by theirtwo structures in Figure 1 that were proposed based on DFTcalculations by Karlberg et al.17 Similar structures were pro-posed by others.5�7,9�11,14,15,17,20,21 The (

√3�√

3)R30� struc-ture seems to be more commonly observed, since it requires lesslong-range order.6

Traditionally heats of adsorption are measured in ultrahighvacuum (UHV) using temperature programmed desorption(TPD); however, this method is limited to species adsorbedreversibly and with insignificant excess activation barrier beyondthe net reaction energy for desorption. Since the excess activationbarrier could be large for the decomposition of (H2O 3 3 3OH)ad(due to the barrier to break theO�Hbond inwater in the reversereaction of H2Og with Oad), TPD cannot give the net reactionenergy in a reliable way. Here, we report direct microcalorimetrymeasurements of the heat of adsorption during formation of the(H2O 3 3 3OH)ad complex as a function of oxygen precoverageand water coverage. The heat of adsorption and sticking prob-ability of water on clean and oxygen-predosed Pt(111) surfaceswere measured as a function of water coverage at sampletemperatures ranging from 120 to 150 K. For reasons to bedescribed, the experiments are done with D2O rather than H2O.Clay et al.5 studied the water�hydroxyl complex formationusing both D2O and H2O but saw no discernible differences.These measurements provide the heat of formation of the(D2O 3 3 3OD)ad complex, from which we estimate the heat offormation of surface ODad and the Pt�OD bond energy. Theseare important values for understanding the energetics of anycatalytic reaction mechanism on Pt involving adsorbed OH.They also provide important benchmark systems with knownenergies against which to compare new computational methodsbeing developed to improve upon the energy accuracy of densityfunctional theory, which is known to have some problems in thatrespect22 but is so widely used that its improvement would havea major impact. To our knowledge, the heat of formation ofadsorbed hydroxyl groups in a well-defined geometry has never

been measured on any surface previously, except for our pre-liminary presentation of a very small part of these results (twospecific oxygen and water coverage combinations).23 There, wealso present new density functional theory calculations, whichaccount for van derWaals interactions and zero point energies, ofthe energies of the two structures produced at those two specificcoverage combinations which agree with these measured en-ergies to within 1 and 15 kJ/mol of D2O, and thus give moreinsight into the details of the structure of the (D2O 3 3 3OD)adcomplex.

’EXPERIMENTAL SECTION

Experiments were performed in an ultrahigh vacuum (UHV)chamber (base pressure <2 � 10�10 mbar) with capabilities forsingle crystal adsorption microcalorimetry and surface analysis,as described previously.24 Briefly, the chamber is equipped forX-ray photoelectron spectroscopy (XPS), Auger electron spec-troscopy (AES), low energy ion scattering spectroscopy (LEIS),and LEED. The sample used was a 1 μm thick Pt(111) singlecrystal foil, supplied by Jacques Chevallier at Aarhus University inDenmark. The Pt(111) sample was exposed to a pulsed molec-ular beam of D2O; each pulse was 102 ms long and repeatedevery 2 s. D2O (Cambridge Isotope Laboratories, purity >98% ,with impurities being mainly HDO and H2O) was outgassedafter putting into its reservoir on the vacuum chamber by fivefreeze-pump-thaw cycles. Its purity was checked with a massspectrometer and found to be consistent with the manufacturer'sclaim. The beam was created by expanding ∼2.3 mbar of D2Othrough a microchannel array and then collimated through aseries of five liquid nitrogen cooled orifices. Coverages arereported here inmonolayers (ML) and are defined as the numberof D2O molecules which adsorb onto the surface irreversibly,independent of the actual species they form on the surfacefollowing adsorption. We define one monolayer (ML) of ad-sorbate coverage as being equal to the density of Pt atoms inthe Pt(111) surface (1.50 � 1019 m�2). A typical D2O dose is0.02 ML (∼4 � 1012 molecules) per pulse. A more detaileddescription of the experimental principles and implementation ofthe molecular beam can be found elsewhere.25

The flux of D2O from the molecular beam is measured byimpinging the beam onto a liquid nitrogen cooled quartz crystalmicrobalance (QCM). A sticking probability of D2O onto theQCM of 1.00 was ensured by precovering the QCM with

Figure 1. Structures of the coadsorbed (H2O 3 3 3OH)ad complex on Pt(111) in its (3 � 3) structure (a) and its closely related (√3�√

3)R30�structure (b), as proposed by Karlberg et al. based on DFT calculations and comparisons to experiments. Reproduced with permission from ref 17.



multilayers of D2O (see the Results section). Calibration of theQCM has been described previously.26 The heat released fromthe adsorption of oneD2O pulse to the sample surface is measuredwith a pyroelectric polymer ribbon gently pressed against theback side of the platinum sample, as described previously.24 Theindicated references provide amore in-depth discussion of thermalconduction between sample and ribbon.24,27

The sensitivity of the pyroelectric detector was calibrated aftereach experiment by depositing a known amount of energyinto the sample using a HeNe (632.8 nm) laser, as describedpreviously.24 The absolute accuracy of the calorimetric heats (afteraveraging >5 runs) is estimated to be better than 3% for systemslike those studied here which have sticking probabilities above 0.8,based on comparisons of our measurements to literature valueswhen forming multilayers with known energies for four differentmolecules.28 (This error is due to possible errors in the absoluteflux and heat signal calibrations and in the literature values.) Theerror is estimated to increase to ∼5% when making hydroxylgroups here, due to errors associated with reproducing oxygenadatom precoverages. Relative measurements (e.g., differences inheat with increasing coverage) are muchmore accurate, due to thehigh precision in heat measurements within a run (see below).

The time scale of our heat measurement is comparable to thebeam pulse duration (∼100ms). If the gas molecules in the pulsecome to their final, most stable surface structures on a time scaleof∼10 ms or less, the heat signal has the same line shape as seenduring the calibration using the (instantaneous) laser heating, sothat calibration is valid. If the molecule’s relaxation time into itsmost stable structure is longer than ∼10 ms, it affects the heatsignal line shape. The contribution to the heat signal can even becompletely lost, if it is much slower than 100 ms. In themeasurements described below of the heat of adsorption ofD2O onto oxygen-predosed Pt(111) at 150 K, we observed thatthe heat signal line shape was ∼6% broader in full width at half-maximum than that for the laser calibration, proving that there issome small time delay in achieving the most stable structure(here, the coadsorbed D2O 3 3 3OD complex). This line shapebroadening was well-fitted by convoluting the instrument responsefunction (i.e., the laser heat signal line shape) with a first-orderreaction (i.e., an exponential decay from the initially populatedadsorbed state into its final, most stable structure) with a timeconstant of 30 ms. This convolution resulted in a decrease in heatsignal by 10.% To correct for this small amount of demodulation,the heat signals measured here were multiplied by 1.10 beforescaling with the laser calibration factor (details given elsewhere28).

The Pt(111) sample was heated radiatively by placing a hottungsten filament (wound into a flat, square-shaped coil)∼1mmbehind its rear face. The temperature of the sample duringannealing was measured using an optical pyrometer. The samplewas cleaned by 1.25 kV Arþ-ion sputtering, annealed at 1073 Kfor 2 min, and exposed to 1� 10�8 mbar O2 at 723 K for 1 min.After the oxygen treatment, the sample was allowed to cool downfor 1 min in the presence of 1� 10�8 mbar O2, directly dosed tothe sample surface, to cover the surface with oxygen (intended tohelp prevent impurities from adsorbing or react them away later).After thermal equilibrium had been re-established throughoutthe calorimeter, the sample was flash heated to 1073 K (<1 s)to remove oxygen, producing a clean Pt(111) surface, andbrought back into contact with the pyroelectric detector. Ther-mal equilibrium was usually re-established 5�10 min after flashheating the sample, after which an experiment was started. After

this treatment, impurities were below the detection limit of Augerelectron spectroscopy across the area of the platinum sample.

The calorimeter was mounted on top of a large thermalreservoir cooled by passing two separate streams of dry nitrogengas first through coils of Cu tubing submersed in two differentlow-temperature baths (i.e., a mixture of ice and water (273 K)and liquid nitrogen (77 K)). The two different-temperature gasstreams were mixed and passed through the reservoir, which is inthermal contact with the sample, thereby cooling the sampleto the desired temperature. We were not able to directly measurethe temperature of the thin platinum sample because mountingthermocouples directly to the sample or sample holder would beimpractical with our instrument design. Instead, we measuredboth temperature A, that of the holder of the pyroelectric ribbon,which is in thermal contact with the backside of the sampleholder, and temperature B, that of the directly cooled thermalreservoir in thermal contact with the front side of the sampleholder. The average of these two temperatures was used to esti-mate the sample temperature. The typical temperature differencebetween A and B is about 20 K. The temperature was monitoredby alumel/chromel thermocouples spot-welded to the holder ofthe pyroelectric ribbon and to the thermal reservoir. The sampletemperatures used in this study were 90, 120, and 150 K.

We used D2O in the molecular beam, as opposed to H2Owhich is present in the background of the vacuum chamber, tomaximize the signal-to-noise ratio in mass spectrometry detec-tion of pulsed molecular beam signals used to determine stickingprobability. (Note, however, that the H2O background pressurewas still so low that <0.01 ML of H2O could adsorb during thetime of our measurement.) The (D2O�OD)ad complex wasprepared by dosing D2O onto the Pt(111) surface precoveredwith oxygen adatoms, under conditions which are known toproduce the analogous (H2O�OH)ad complex upon dosingH2O at 130�165 K.2,3,5,8,13,29 Atomic oxygen is known to forma p(2 � 2) structure on Pt(111) with a saturation coverage of0.25MLOad, which can be formed by dosing O2 onto the surfaceat temperatures above 150 K, where O2 either dissociates ordesorbs.12,30�32 The Oad layers were prepared using a directedO2 doser translated 1 mm away from the front face of the sample.The sample was flash heated to 1100 K immediately prior todosing O2 so that the sample was still above 150 K during dosingand the O2 dissociated. The directed doser resulted in anenhancement factor of ∼10� compared to dosing oxygen byonly backfilling the chamber. The coverage of atomic oxygen wasdetermined by the O(510 eV)/Pt(238 eV) ratio observed withAuger electron spectroscopy, calibrated against the saturatedp(2� 2)-Oad layer (which gave the expected sharp LEEDpattern).The Pt(111) surface with different amounts of adsorbed atomicoxygen (Oad) was exposed to a choppedmolecular beam of D2O,and the sticking probabilities and the heats of adsorption weremeasured as a function of D2O coverage at 90, 120, and 150 K.Care has to be taken when working with oxygen on Pt(111), sinceoxygen is readily removed from the surface by backgroundCOandH2. It is therefore important to keep their background pressureslow and to dose water as soon as possible after dosing oxygen.

Sticking probabilities were measured simultaneously withcalorimetric measurements, using the King and Wells method.33

A mass spectrometer, without line-of-sight to the sample, mea-sured the background pressure increase of D2O (m/z = 20) in thechamber. A gold flag was positioned in front of the sample andused to determine the mass spectrometry signal correspondingto full reflection of D2O. A gold flag was used because D2O does



not stick to gold at room temperature.34 The sticking probabilityof D2O is calculated by integrating the mass spectrometer signalmeasured from the increase in D2O partial pressure above back-ground when the molecular beam is pulsed onto the samplesurface in comparison with the increase in D2O partial pressureresulting whenD2O is pulsed onto the inert gold flag. Wemeasuretwo types of sticking probabilities:25 the long-time sticking prob-ability, which is defined as the probability that molecules in a gaspulse stick on the surface until the next pulse arrives 2.0 s later, andthe short-time sticking probability, which is defined as the prob-ability that molecules in a gas pulse stick on the surface until theend of the ∼100 ms time period used to measure the heat signal.The long-time sticking probability is used to calculate the coverageat the start of the next pulse, and the short-time sticking probabilityis used to calculate adsorption energies per mole adsorbed. Whenthere is no desorption between pulses, they are the same.

’RESULTS

Sticking Probability. Figure 2 shows the long-time stickingprobability of D2O on Pt(111) at 150 K precovered with 0.05,0.09, 0.14, 0.21, and 0.25 ML of oxygen adatoms. Oxygencoverages above 0.13 were estimated from the O2 dose using acalibration plot of coverage by AES versus O2 dose. Unfortu-nately, this method had large error bars at low coverages andtherefore lower oxygen coverages were determined from themeasured saturation water coverage, assuming a 4:1 ratio ofwater to oxygen at saturation as reported by Clay et al. at 163 K.5

The resulting saturation water coverage versus Oad precoverageagreed very well with Clay et al. at all coverages up to 0.25 MLO,so it is not reproduced here. The sticking probability of D2Oon >0.14 ML Oad on Pt(111) at 150 K is near unity (>0.90) atD2O coverages below 0.4ML (see Figure 2), indicating a precursor-mediated adsorption mechanism. At high Oad precoverages, theuptake of D2O saturates at 0.55�0.65ML. Even at the lowest Oad

precoverage, D2O is stabilized at this temperature, whereas D2Odesorbs again on this time scale from oxygen-free Pt(111).28 Thesurface saturates at 0.2 ML of D2O at ∼0.05 ML Oad coverage.This behavior is consistent with LEED results which show a small

amount of oxygen is enough to pin D2O to a different structurethan it adopts on clean Pt(111) and with TPD results which showa higher D2O desorption temperature even for low Oad coverage.

5

The short-time sticking probability at 150 K was 0.90 ( 0.04below 0.5 ML D2O for all the oxygen precoverages of Figure 2and dropped to ∼0.8 by 0.6 ML D2O coverage.Adsorbed D2O did not form multilayers at temperatures

greater than 120 K with or without oxygen precoverage. At120 K and below, D2O forms multilayers with or without Oad.The initial sticking probability of D2O on both clean and oxygen-covered Pt(111) at 120 K was 0.91 ( 0.01, independent of Oad

coverage up to 0.25 ML. The sticking probability increased to0.97 by 0.5 ML D2O coverage and to 1.00 by 1.0 ML D2O, asD2O forms multilayers.28

Heat of Adsorption. Throughout this paper, we will use theterm heat of adsorption, which we define here as the negative ofthe differential standardmolar enthalpy of adsorption (�ΔH�/nads),with the gas and surface both at the temperature of the Pt(111)sample. Since the gas starts out at a different temperaturethan the Pt(111) in the real experiment, this requires a smallcorrection to the heat actually measured, which is the differencein internal energy between the initial and final states of thesystem, as described elsewhere.25

The heat of adsorption of D2O on oxygen-predosed Pt(111)at 150 K versus D2O coverage is shown in Figure 3. At 150 K,water is known to react withOad on Pt(111), in a ratio of betweentwo and four water molecules per Oad (depending on the amountof oxygen preadsorbed), to produce a coadsorbed H2O 3 3 3OHcomplex.2,3,5,8,29 Therefore, it is safe to assume we are making theanalogous (D2O 3 3 3OD)ad complex here, except at the highestD2O coverages (see below). The different curves shown here arefor different coverages of predosed oxygen adatoms (∼0.07,∼0.18, and 0.25 ML) and are the averages of several runswith similar Oad coverages in the ranges specified. At each Oad

precoverage, the highest D2O coverage point shown here is thesaturation coverage achieved at 150 K for the D2O dose rate used(except for the transiently adsorbed ∼0.02 ML dose per pulsethat is pumped away between pulses of the D2O beam atsaturation). These agree with the results of Clay et al. at 170 K.5

Figure 2. Average sticking probability of D2O versus D2O coverage onPt(111) at 150 K with different Oad coverages (0.05, 0.09, 0.14, 0.21,and 0.25 MLOad). The gas here is a collimated molecular beam from aneffusive source held at 300 K. Precursor mediated behavior is observedfor oxygen coverages greater than 0.05 O ML.

Figure 3. Heat of adsorption of D2O versus D2O coverage on Pt(111)at 150 K that had been precovered with different amounts of oxygenadatoms (0.07, 0.18, and 0.25 ML Oad). Also shown is the integral heatof adsorption for 0.25 ML of Oad. Each curve is the average of severalruns with similar coverages of Oad in the range given.



Returning to Figure 3, with 0.25 monolayers (ML) of oxygenatoms, the differential heat of adsorption is nearly constant at61.3 kJ/mol for the first ∼1/3 ML but drops to 57.9 kJ/mol by0.50ML and 50.5 kJ/mol by saturation (0.62ML). The top curveis the corresponding integral heat of adsorption of D2O on 0.25ML Oad precovered Pt(111). Similar experiments but with Oad

precoverages of 0.18 and 0.07 ML gave lower saturation D2Ocoverages (0.55 and 0.22 ML, respectively) and lower heats ofD2O adsorption by∼3.4 and∼4.6 kJ/mol, respectively, except atvery low water coverage (∼0.02 ML), at which a larger differ-ential heat of adsorption of ∼62 kJ/mol was observed, nearlyindependent of oxygen precoverage.We made the above measurements at 150 K because we were

unable to make calorimetric measurements at temperatureshigher than 150 K (but below the ∼180 K decompositiontemperature of the surface hydroxyls5,6,13,29) without significanttransient desorption of D2O during the 2 s period of the pulsedD2O molecular beam, which cools the sample and gives a signalthat complicates heat determination. For calorimetry, it is alwaysbest to make the measurements at the highest temperatureknown to produce the species of interest (surface hydroxyl here),since they come to their final stable structuremost rapidly. (If thisoccurs on a time scale similar to or slower than the pulse duration(100 ms), it changes the heat signal linehape and thus compli-cates heat signal analysis.) The transient mass spectrometersignal for D2O (used mainly for measuring the sticking probability)showed that, at >150K, there are two types of water binding to thesurface: one which is populated only at high water coverage, isweakly bound and leaves the surface within a few hundredms afterthe 100 ms long water pulse, and another, more strongly boundwater, which remains bound to the surface over the entire 2 speriod of each pulse. Desorption of the weakly bound speciesaffects the heat signal pulse shape at high D2O coverage (when thesticking probability is strongly decreasing) but not enough tochange the heatmeasurement significantly, which is taken from theslope of the initial sharp signal rise.25

Figure 4 compares the results of Figure 3 for the heat ofadsorption of D2O at 0.25 ML Oad and 150 K with controlexperiments at 120 K for D2O on clean Pt(111) and on Pt(111)predosed with O.25 ML Oad, all plotted versus coverage of D2O.The data at 120 K are reproduced from our earlier work.28 At120 K, water does not dissociate to produce hydroxyl groups oneither clean or O-predosed Pt(111).5,6,29,35 The coadsorbed(H2O 3 3 3OH)ad complex is known to only form above 130 Kon oxygen precovered Pt(111).5,6 The heat of adsorption for thefirst 1/2ML of D2O seen here at 120 K (55 kJ/mol) is lower thanthat at 150 K (61 kJ/mol). This is consistent with the reaction notforming this more stable (D2O 3 3 3OD)ad complex at 120 K,presumably because it cannot scale the activation energy asso-ciated with breaking the O�D bond but instead simply adsorb-ing molecularly. This heat is still higher than the heat for D2O onclean Pt(111) at 120 K, 51.3 kJ/mol.28 Our measured heat ofadsorption for multilayer D2O on clean or oxygen-predosedPt(111) at 120 K is 47.2( 1.4 kJ/mol. This is within error bars ofthe literature value from TPD experiments, in which the heat ofsublimation of deuterated amorphous ice multilayers on Ru-(0001) was determined to be 49.11 kJ/mol ( 2% in thetemperature range 145�155 K.36 (Our measurement is theenthalpy of adsorption, to which one must also subtract 1/2RT for direct comparison to an Arrhenius activation energy,giving an activation energy of 46.6( 1.4 kJ/mol at the desorptionrate measurement's average temperature.) This confirms the

accuracies of our heat and water beam flux measurements andthe calibrations upon which they are based.

’DISCUSSION

On the basis of our understanding of the literature assummarized in the Introduction, the best model to explain theprior literature is that water reacts with predosed oxygen adatomsunder the conditions present in our data of Figure 3 according toreaction 1, rewritten here with D instead of H:

3D2Og þOad f 2ðD2O 3 3 3ODÞad ð1ÞHowever, as noted in the Introduction, some have proposed thatthe reaction instead is

2D2Og þOad f 3ODad þDad ð2ÞWenow interpret the heat data of Figure 3 according to these twopossible reaction mechanisms to determine the heat of formationof the resulting product ((D2O 3 3 3OD)ad or ODad), and furthershow from these values that both mechanisms result in a verysimilar value estimate of the Pt�OD bond enthalpy (between263 and 274 kJ/mol). As noted above, the literature much morestrongly supports reaction 1, but this comparison to reaction 2 isnevertheless interesting in that it gives such a similar Pt�ODbond energy (see below).

Figure 5 shows the thermodynamic cycle for the heat ofreaction, assuming the reaction proceeds via reaction 1 to formthe (D2O 3 3 3OD)ad complex. This cycle is for the highest oxygenprecoverages that allow this complex to form cleanly by thismechanism, which is for the case of 0.167 ML Oad plus 0.50 MLD2Oad, giving the total O coverage of 0.67 ML. This 0.67 MLvalue is the known coverage in either the (

√3�√

3)R30� or(3�3) structure,5 although under the present conditions the

Figure 4. Differential heat of adsorption of D2O on Pt(111) predosedwith 0.25 ML Oad at 150 K (black) and at 120 K (red). For comparison,the curve for clean Pt(111) (reproduced from ref 28) is also shown(blue). Multilayers of D2O form after ∼0.8 ML of water at 120 K. Thelines drawn represent the regions used to estimate the monolayer andmultilayer energies. The error indicated represents run-to-run variationsin the absolute heats, reflecting the statistical errors in the calibration ofcombined heat signal and gas flux measurements. These are the averagesof 4, 10, and 5 runs, respectively. We assumed here that the multilayerheat at 120 K is the same with or without the Oad, so we scaled the heatsslightly (2�4%) so that their multilayer heat is the same as the overallaverage for all 15 runs (47.2 ( 1.4 kJ/mol).



(√3�√

3)R30� structure is known to be produced.5 This cyclegives a heat of formation of the (D2O 3 3 3OD)ad complexof �527 kJ/mol and a Pt�OD bond enthalpy of 263 kJ/molusing heat data from Figure 3. Higher Oad precoverages cannotcompletely react via reaction 1, since this would lead to localcoverages exceeding the known total O coverage of 0.67 ML inthese (

√3�√

3)R30� or (3�3) structures.Starting with the elements in their standard states on the left

side of Figure 5 and following the cycle pathway across thebottom, the first step’s enthalpy is the sum of 3 times the standardenthalpy of formation of D2Og

37 plus the integral heat of adsorp-tion of 1/2 O2,g to make Oad at 0.17 ML coverage.38 The nextstep’s enthalpy is the calorimetrically measured heat of adsorptionat these coverages, from Figure 3. Now following the cycle pathwayover the top, the first step’s enthalpy is simply twice the standard

enthalpies of formation ofD2Og andODg, from ref 37. The secondstep’s enthalpy is determined to be �629 kJ/mol by forcing thecycle to sum to 0 kJ/mol. If we assume that the D2Oad in the(D2O 3 3 3OD)ad complex has the same enthalpy as D2Oad onO-free Pt(111), this step’s enthalpy is simply twice the sum of theheat of adsorption of D2O on clean Pt(111) measured here inFigure 4 (51.3 kJ/mol) plus the heat of adsorption of ODg. Fromthis, we estimate the heat of adsorption of ODg, or the Pt�ODbond enthalpy, to be 263 kJ/mol, for OD adsorbing into pre-adsorbed D2O to make the stable (D2O 3 3 3OD)ad complex.Combining this with the standard enthalpy of formation of ODg

(þ37.23 kJ/mol) gives the enthalpy of formation of ODad

to be �226 kJ/mol, assuming that the coadsorbed D2O in thiscomplex with OD has the same stability as in a pure D2O adlayeron Pt(111). The central pathway from left to right in Figure 5 isdetermined to be �1053 kJ/mol by summing either the upperor lower pathway. This is, by definition, twice the standardenthalpy of formation of this (D2O 3 3 3OD)ad complex, which isthen �527 kJ/mol.

Figure 6 shows the thermodynamic cycle for the heat ofreaction, assuming the reaction proceeds via reaction 1 to formthe (D2O 3 3 3OD)ad complex. This cycle is for the case of 0.25ML Oad þ 0.375 ML D2Oad, which results in 75% of the surfacearea covered by the pure (D2O 3 3 3OD)ad complex with itslocal coverage of 0.67 ML total O. The remaining 0.125 ML ofunreacted Oad is assumed here to be compressed into theremaining 25% of the surface with a local coverage of 0.50 ML(referred to here as Oad(comp)). The enthalpy cost of thiscompression, ΔHcomp(Oad), was estimated to be 17 kJ/molOad using the desorption activation energy versus coveragemeasured by Koel’s group.39 (We used here their result thatthe average desorption energy from 0.25 to 0.5 ML is 34 kJ/molO larger than that from 0 to 0.25 ML, which means the integralheat of adsorption drops by 17 kJ/mol from 0.25 to 0.5 ML.)This cycle gives a heat of formation of the (D2O 3 3 3OD)adcomplex of �537 kJ/mol and a Pt�OD bond enthalpy of274 kJ/mol. The other literature enthalpy values needed tocompute this cycle and the cycle below were obtained asdescribed above for Figure 5.

Figure 5. Thermodynamic cycle for the heat of reaction, assuming thereaction proceeds with a stoichiometry of 3 D2O to 1 Oad via reaction 1to form the (D2O 3 3 3OD)ad complex. This cycle is for the case of 1/2MLD2Oadþ 1/6MLOad, which results in a total O coverage of 2/3ML,covering the entire Pt(111) surface in the (

√3�√

3)R30� structure ofpure (D2O 3 3 3OD)ad complex.5 It uses the data from the 0.18 ML Oad

curve in Figure 3 (which is the closest measurement to 1/6MLOad) andat 1/2 ML D2O. This cycle gives a heat of formation of the(D2O 3 3 3OD)ad complex of�527 kJ/mol and a Pt�OD bond enthalpyof 263 kJ/mol. Enthalpies measured here are shown in bold font. See textfor details of other enthalpies.

Figure 6. Thermodynamic cycle for the heat of reaction, assuming the reaction proceeds with a stoichiometry of 3 D2O to 1 Oad via reaction 1 to formthe (D2O 3 3 3OD)ad complex. This cycle is for the case of 3/8 ML D2Oad þ 1/4 ML Oad, which results in 75% of the surface area covered by the pure(D2O 3 3 3OD)ad complex with a local total O coverage of 2/3 in either the (

√3�√

3)R30� or (3�3) structure.5 The remaining 1/8 ML of unreactedOad is assumed here to be compressed into the remaining 25% of the surface (so that its local coverage is 1/2ML). The enthalpy cost of this compression,ΔHcomp(Oad), is estimated to be 17 kJ/mol O as described in text. This cycle gives a heat of formation of the (D2O 3 3 3OD)ad complex of�537 kJ/moland a Pt�OD bond enthalpy of 273 kJ/mol. See text for details.



Starting with the elements in their standard states plus Oad atits initial coverage (0.25 ML) on the left side of Figure 6, theenthalpies for the two steps in the cycle pathway across thebottom are calculated in exactly the same way as in Figure 5,except that the coverages are different here. Following the cyclepathway over the top, the first step’s enthalpy is also the same asin Figure 5. The second step is the compression of the oxygenadlayer, whose enthalpy was given above. The third step’senthalpy is determined to be �649 kJ/mol by forcing the cycleto sum to 0 kJ/mol. Analyzing this last step’s enthalpy in the sameway as in Figure 5 gives the heat of adsorption of ODg, or thePt�OD bond enthalpy, to be 273 kJ/mol, for OD adsorbing intopreadsorbed D2O to make the most stable (D2O 3 3 3OD)adcomplex. The corresponding standard enthalpy of formation ofODad (coadsorbed with D2O) is �236 kJ/mol. The centralpathway from left to right in Figure 5 is determined to be�1057kJ/mol by summing either the upper or lower pathway. This is,by definition, twice the standard enthalpy of formation of thiscomplex plus the enthalpy to compress the unreacted Oad into ahigher local coverage, mentioned above to be 17 kJ/mol O. Thisthen gives �537 kJ/mol for the standard enthalpy of formationof the (D2O 3 3 3OD)ad complex, very similar to the�527 kJ/moldetermined from the lower oxygen coverage data in Figure 5. The10 kJ/mol greater stability here arises from the∼3 kJ/mol higherheat of adsorption in Figure 3. The higher heat of adsorption issurprising given the energy cost here of compressing theunreacted oxygen adatoms into more dense domains. At present,we do not have an explanation for this. Perhaps there is an error inour assumption that the reaction here forms the (D2O 3 3 3OD)adcomplex with its local coverage of 0.67 ML total O and that theremaining 0.125 ML of unreacted Oad is compressed to a localcoverage of 0.50 ML. However, this seems unlikely given thatprior results show that similar exposures with nondeuteratedwater produce the (H2O 3 3 3OH)ad complex in domains of its(√3�√

3)R30� structure.2,5�7,9 The extra Oad would seem tohave no other choice but to compress as assumed.

Figure 7 shows the thermodynamic cycle for the heat ofreaction assuming the reaction proceeds instead via reaction 2to form 3 ODad plus Dad, for the case of 0.18 MLOadþ 0.36 MLD2Oad, resulting in a total O coverage of 0.52 ML. It gives a heatof formation of the ODad of �237 kJ/mol and a Pt�OD bondenthalpy of 274 kJ/mol. The enthalpies of the individual steps are

calculated in essentially the same way as in Figure 5 except thatthe second step across the top path does not produce thecomplex but instead only makes 3 ODad plus Dad. Its enthalpyof �859 kJ/mol, determined by summing the cycle to zero netenthalpy, is then the sum of 3 times the Pt�OD bond enthalpyplus the integral heat of adsorption of 1/2 H2,g to make Had

(�36 kJ/mol, from refs 40�46), which we assume is the same forthe D isotope. This gives the heat of adsorption of ODg, or thePt�OD bond enthalpy, to be 274 kJ/mol.

Note that all three cycles give a very similar Pt�OD bondenthalpy of between 263 and 274 kJ/mol, so this range of valuesis independent of choosing whether reaction 1 or 2 best describesthe real situation. As noted above, reaction 1 is more consistentwith the literature data and calculations summarized in theIntroduction.

From previous experimental results of Clay et al.,5 theexpected water uptake for a Pt(111) surface precovered with0.25 ML Oad at 163 K is 0.58( 0.03 ML, resulting in 0.80�0.86ML adsorbed oxygen-containing species total. At 150 K, weobserve an average water uptake of 0.62 ( 0.02 ML of wateradsorbing with 0.25 ML Oad on Pt(111), resulting in 0.85�0.89ML adsorbed oxygen-containing species total. This is slightlylarger than the results of Clay et al.,5 although within experi-mental errors. A slightly greater amount of total Oad species isexpected here, due to the calorimetric measurements being madeat a slightly lower temperature (150 K versus 163 K) whichwould allow more water to adsorb onto the surface. This isconsistent with the observation of Schiros et al.3 that it wasnecessary to heat their Pt(111) sample to 156 K to desorbedweakly bound second-layer water so that only the monolayerremained. This extra amount of water that has adsorbed at 150 Kwould have adsorbed with a lower heat and may not haveadsorbed in the 163 K experiment of Clay et al. This is likely,since we have observed that an additional ∼0.02 ML of water istransiently adsorbed with every gas pulse beyond saturation at150 K, but it desorbs again in <2 s. We made our measurementsat 150 K instead of 163 K for the reasons outlined above.However, we noted elsewhere28 that 150 K here correspondsto 160 K in the papers by Clay et al.5,6 due to systematic errors inthermocouples.

The dissociation of adsorbed water

D2Oad f ODad þDad ð3Þ

is a key elementary step in the water-gas shift reaction and thesteam reforming or oxidation of practically any organic moleculeon Pt catalysts. We can estimate the enthalpy for this reaction onPt(111) using: (1) the heat of formation of ODad of�226 kJ/molfrom the thermodynamic cycle in Figure 5 (see above), (2) theintegral heat of formation for Had (�36 kJ/mol at low coveragewhere it is most stable, from refs 40�46), and (3) the heat offormation of D2Oad (standard enthalpy of formation of D2Og

�249.2 kJ/mol, from ref 37, plus the measured integral heat ofadsorption of D2Og on clean Pt(111) at 0.5 ML coverage of�51.3 kJ/mol).28 The enthalpy of this dissociation reaction isthus estimated to be þ39 kJ/mol of D2O (endothermic). Usinginstead the cycle in Figure 6 for a higher oxygen precoverage(which gives ODad to be 10 kJ/mol more stable) gives thisenthalpy reaction to be þ29 kJ/mol D2O (endothermic).Reaction 3 is even more endothermic than this 29�39 kJ/molestimate when the ODad product is not stabilized by coadsorbedwater, or when the Dad product is at higher coverage. Both these

Figure 7. Thermodynamic cycle for the heat of reaction assuming thereaction proceeds with a stoichiometry of 2 D2O to 1 Oad via reaction 2to form 3 ODad plus Dad. This cycle is for the case of 0.36 ML D2Oadþ0.18 ML Oad, which results in a total O coverage of 0.52 ML. It gives aheat of formation of the ODad of �237 kJ/mol and a Pt�OD bondenthalpy of 274 kJ/mol. See text for details.



values are about half the energy barrier to dissociation predictedby DFT, 65 kJ/mol.14 This is only partially due to the stabiliza-tion of the product ODad by coadsorbed D2O here, since thereactant D2Oad is also stabilized by coadsorbed D2O.

The larger differential heat of adsorption of ∼62 kJ/molobserved at the lowest water coverage (∼0.02 ML) (which isnearly independent of oxygen precoverage above 0.06 ML O) isprobably due to hydroxyl complex formation at defects. DFTcalculations of water monomers, dimers, and trimers have beenshown to have higher binding on steps and kinks than onterraces.47 More importantly here, adsorbed OH is more stableat step edges, as shown by our previously unpublished DFTcalculations presented in Table 1.

Clay et al.5 did a TPD experiment to compare the thermaldecomposition stability of H2O�OH complexes formed fromdifferent water plus Oad compositions. They showed that, atlower Oad precoverage (0.14 ML), the TPD peak resulting fromH2O desorption from the (H2O 3 3 3OH)ad mixed layer occurs athigher temperature (∼6 K higher) than when the O precoverageis 0.25 ML, implying a higher desorption energy (assumingthe same prefactor).5 In contrast, at both 0.07 and 0.18 ML ofpreadsorbed O, we measure a slightly lower heat of D2Oadsorption (Figure 3). Specifically, the heat of adsorption on a0.18 ML Oad precovered surface remains constant up until 0.4ML of D2O (i.e., when the ratio of D2O to Oad is < 2) with anaverage heat value of 58.4 kJ/mol. The average heat of adsorptionon a 0.07 ML Oad precovered surface up to 0.22 ML of D2O is56.9 kJ/mol; there is no flat region at such a low precoverage ofoxygen because of the effect of defect sites mentioned above.There are two potential problems with the analysis of the TPDdata that could cause the observed 6 K peak shift to bemisleadingwith respect to adsorption energy: (1) a different prefactor or (2)

a small and different activation energy for forming the layer thatcontributes to the desorption activation energy but not theadsorption heat. These small differences in peak temperaturesin TPD can easily arise from these or other dynamical effects.Another possibility for the difference in trend is that the calori-meter measures heats over the 100 ms time scale of the gas pulse,so if the gas molecules takes >∼1000 ms to achieve their moststable structure after hitting the surface, it will miss some of theheat. (If this takes between ∼10 and ∼1000 ms, it causes a lineshape change in the heat signal, which is seen and corrected (seethe Experimental Section).) However, this effect would probablygo in the opposite direction, since the experiment at 0.18 ML ofOad already has the proper local coverage to optimize the(√3�√

3)R30� structure of the (D2O 3 3 3OD)ad complex,whereas the experiment at 0.25 ML of Oad has a local coveragethat is too high, and excess Oad must be squeezed out of thedomains of the (

√3�√

3)R30� (D2O 3 3 3OH)ad structure. Onewould therefore expect the latter process to take longer.

In principle, one can analyze TPD data following the decom-position of the (H2O 3 3 3OH)ad complex, such as presented byClay et al.5 and Creighton andWhite4 to get an activation energyfor the decomposition of the (H2O 3 3 3OH)ad complex. How-ever, large errors in activation energy result when analyzing thesedata, as has been discussed in detail.8 Also, any such activationenergy, even if determined accurately, would not give the netreaction energy we measured here, since it would include theexcess activation energy for adsorption, which could be large,since the DO�D bond is broken in making ODad. Nevertheless,assuming the adsorption process is not activated, we can use ourcalorimetric heats of adsorption for D2Oad adsorbing onto Oad

precovered Pt(111) to make the (D2O 3 3 3OD)ad complex as anestimate of the activation energy for its decomposition (Edec).We assume that this decomposition occurs by the reverse ofreaction 1: 2(D2O 3 3 3OD)ad f 3D2O þ Oad. From the TPDexperiments of Clay et al.5 with a heating rate (β) of 0.95 K/s at0.25 and 0.14 ML Oad precoverages, the decomposition peaktemperatures (Tp) are 199 and 205 K, respectively. We analyzethese by assuming a first-order decomposition process, forwhich the Redhead equation48 is Edec/(RTp

2) = (ν/β) exp-[�Edec/(RTp)], where ν is the decomposition reaction’s pre-exponential factor. Using our enthalpies of reaction at 0.25 MLOad coverage (61.2 kJ/mol) and 0.18MLOad coverage (57.4 kJ/mol)as the Edec values, we estimate the prefactors to be 2.0 � 1015/sand 6.6 � 1013/s, respectively. This lower prefactor at the lowerOad precoverage implies a more disordered (higher entropy)adlayer, which is not unreasonable. This difference in prefactorsfor dissociation could explain why the TPD experiments of Clayet al. implied that the complex formed at 0.25 ML Oad is lessstable to desorption (lower Tdes) than the complex formed at0.14 ML Oad precovered Pt(111). Using this same Edec value of61.2 kJ/mol with the corresponding Tp = 205 K value at 0.25 MLOad precoverage and β = 10 K/s from Creighton and White’sTPD experiment4 gives a slightly higher ν value of 6.9� 1015/s.

’CONCLUSIONS

Calorimetric measurements of the adsorption enthalpy ofD2O on Pt(111) at 150 K predosed with Oad give the heat offormation of the well-known (D2O 3 3 3OD)ad coadsorbed com-plex, fromwhich the enthalpy of adsorbedOD on Pt(111) can beextracted. With 0.25 monolayers (ML) of oxygen adatoms, thedifferential heat of reaction is nearly constant at 61.3 kJ/mol for

Table 1. DFT Calculations with Periodic Boundary Condi-tions of the Adsorption Energy (Ead) and O�Pt Distance(O�Pt) for OHad on Pt(111) and Its Defect Sitesa

a Energies are calculated for low coverages of OHad (g12 Pt surfaceatoms per OH) to minimize interactions between the periodic images.Computation parameters and surface structures have been publishedearlier.47 The adsorption energy is calculated as the energy differencebetween the energy of OH adsorbed on the Pt slab and the summation ofthe energies of the adsorbate-free slab and OH in the gas phase.



the first 0.4 ML but drops to 50.5 kJ/mol by saturation (0.62ML). Similar experiments withOad precoverages of 0.18 and 0.07ML gave lower heats of D2O adsorption by ∼3.4 and ∼4.6 kJ/mol, respectively. From these data, the heat of formation ofODad, stabilized by coadsorbed D2O in this (D2O 3 3 3OD)adcomplex, was estimated to be �226 to �236 kJ/mol; and thePt�OD bond enthalpy was estimated to be 263�274 kJ/mol.From these, we estimate that the dissociation of D2Oad to makeODad þ Dad is uphill in enthalpy by 29�39 kJ/mol. At very lowD2O coverage, the heat of reaction is slightly larger. New DFTcalculations which compare OH at step and terrace sites implythat this is due to step sites.

’AUTHOR INFORMATION

Corresponding Author*E-mail: [email protected].

Author Contributions†These two authors contributed equally to this paper.

’ACKNOWLEDGMENT

The authors acknowledge support for this work by theNational Science Foundation under CHE-1010287. L.A. wassupported by the Office of Naval Research.

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the energy of adsorbed hydroxyl on pt(111) by microcalorimetrydepts.washington.edu/campbelc/pdf/oh...

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