Adsorption, Desorption, and Displacement Kinetics of H2O and CO2

Publication Date (Web): September 22, 2014. Copyright © 2014 American Chemical Society ... [email protected]., *Phone: (509) 371-6143. E-mail: Bruce...
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Adsorption, Desorption, and Displacement Kinetics of H2O and CO2 on Forsterite, Mg2SiO4(011) R. Scott Smith,* Zhenjun Li,† Zdenek Dohnálek, and Bruce D. Kay* Fundamental and Computational Sciences Directorate, Pacific Northwest National Laboratory, Richland, Washington 99352, United States S Supporting Information *

ABSTRACT: We have examined the adsorbate−substrate interaction kinetics of CO2 and H2O on a natural forsterite crystal surface, Mg2SiO4(011), with 10−15% Fe2+ substituted for Mg2+. We used temperature-programmed desorption and molecular beam techniques to determine the adsorption, desorption, and displacement kinetics for H2O and CO2. Neither CO2 nor H2O has distinct submonolayer desorption peaks, but instead both have a broad continuous desorption feature that evolves smoothly into multilayer desorption. Inversion of the monolayer coverage spectra for both molecules reveals that the corresponding binding energies for H2O are greater than those for CO2 on all sites. The relative strength of these interactions is the dominant factor in the competitive adsorption and displacement kinetics. In experiments in which the two adsorbates are codosed, H2O preferentially binds to the highest-energy binding sites available and displaces CO2. The onset of significant CO2 displacement by H2O occurs between 65 and 75 K.

I. INTRODUCTION There is great interest in reducing and capturing atmospheric CO2 in order to mitigate its effects on climate change. One proposed strategy is to pump supercritical CO2 into geological formations wherein it can react to form carbonates and remain trapped.1−3 Olivines are a class of minerals that are believed to have the best potential for carbonation.4−6 Olivines are orthosilicates with the formula X2SiO4 where X is typically composed of Mg2+, Fe2+, or Ca2+, either entirely of one cation or in various mixed fractions. For example, forsterite, Mg2SiO4, is an olivine where X is entirely Mg. Recently, forsterite has been used as a model system to study carbonation reactions in the laboratory.5−11 To understand these reactions at a fundamental level, calculations on the interaction energetics of CO2 and H2O with forsterite and other olivines have been performed.12,13 In a recent study we examined the adsorbate−substrate interaction kinetics of CO2 and H2O on a well-characterized TiO 2 (110) model system to provide a baseline for interpretation of the experimental results on more complex, real-world substrates.14 In the current paper we extend that work to a more complicated system to study the interactions of CO2 and H2O with a natural mineral crystal, forsterite, Mg2SiO4(011). We use temperature-programmed desorption (TPD) and molecular beam techniques to determine the adsorption, desorption, and displacement kinetics for H2O and CO2. Neither CO2 nor H2O has distinct submonolayer desorption peaks but instead both have a broad continuous desorption feature that evolves smoothly to multilayer desorption. Inversion of the monolayer coverage spectra for both molecules reveals that the corresponding binding energies for H2O are greater than that for CO2 on all sites. The relative © 2014 American Chemical Society

strength of these interactions is the dominant factor in the competitive adsorption and displacement kinetics. In experiments in which the two adsorbates are codosed, H2O preferentially binds to the highest-energy binding sites available by displacing CO2. The onset of significant CO2 displacement by H2O occurs between 65 and 75 K.

II. EXPERIMENTAL SECTION The substrate was cut from a natural forsterite crystal (peridot crystal; origin is Sapat, Kohisthan, NW Frontier, Pakistan region). An olivine is a silicate mineral that has the general formula (Mg,Fe)2SiO4. Strictly speaking, forsterite is an olivine that has the formula Mg2SiO4; however, natural forsterite samples often contain various amounts of Fe. An X-ray photoelectron spectroscopy (XPS) analysis of our sample revealed that ∼10−15% of the cation content was iron, i.e., the molecular formula was approximately Mg1.8Fe0.2SiO4. The XPS spectra are shown in the Supporting Information (Figure S1). The sample was cut from the mineral and polished. The crystal surface orientation was determined by X-ray diffraction analysis to be (011) (Figure S2 in the Supporting Information), although a clear and distinct low-energy electron diffraction (LEED) pattern could not be obtained. The lack of a LEED pattern was likely due to surface charging by the incident electron beam. All experiments were conducted utilizing an ultrahigh vacuum system (UHV) with a base pressure of 18 MΩ cm) adsorbates at normal incidence. The backing pressure for both species was 1 Torr. The beam fluxes were calibrated using a quartz crystal microbalance using a previously described procedure.16 The absolute fluxes were 4.3 × 1013 CO2/cm2 s and 1.1 × 1014 H2O/cm2 s. As we show below neither CO2 nor H2O has a distinct monolayer desorption feature on the forsterite substrate. Therefore, we use the saturation monolayer coverages for CO2 (0.41 × 1015 CO2/cm2)18 and H2O (1.1 × 1015 H2O/cm2)19 on Pt(111) to define the respective ML values. This gives a relative beam flux of 0.11 ML/s for CO2 and 0.10 ML/s for H2O. In this definition “1 ML” does not correspond to the same number of molecules for CO2 and H2O. All TPD spectra were acquired using a quadrupole mass spectrometer (UTI 100C) in a line-ofsight geometry and at a constant heating rate of 1 K/s.

Figure 1. Space-filling model of a natural forsterite, Mg2SiO4(011). The coordinates are from an X-ray diffraction study on a natural forsterite sample with 10% Fe impurities.21 In the model, Si atoms are blue, O atoms red, and Mg atoms yellow. (Fe is not shown.)

of the sample. The main point is that the forsterite (011) surface is likely to be more complex than the TiO2(110) surface we recently studied.14 The model shows that O, Mg, and Si adsorption sites are all available on the (011) surface. The experimental XPS data (Figure S1 in the Supporting Information) show the existence of O, Mg, Si, and Fe on the surface, but for simplicity the Fe that is present on the surface is not shown. It is believed that Fe is randomly substituted for Mg in the sample. This means that in the model surface displayed in Figure 1, approximately one of the Mg surface atoms would be replaced by an Fe atom. Because of the presence of all four atomic species, the real surface will likely have a range of binding sites that will give rise to a broad distribution of adsorbate binding energies. Figure 2 displays TPD spectra for CO2 and H2O deposited at 50 K on forsterite and then heated at 1 K/s. The coverages for CO2 in Figure 2a range from 0.09 to 3.30 ML and for H2O in Figure 2b from 0.18 to 3.24 ML. In both cases the curves fill from high to low temperature indicating that the adsorbates have sufficient surface mobility to find the highest-energy binding sites prior to desorbing. While the TPD spectra and subsequent analysis (below) show that there are distributions of binding sites and energies, neither the adsorbate orientation nor the specific adsorption site on the surface is known. For both CO2 and H2O, the monolayer desorption spectrum (red curve in each panel) spans a wide temperature range and smoothly continues to grow as multilayer coverages develop on the surface. An important point is that the desorption temperatures of the two adsorbates are much different with the onset of CO2 desorption beginning at ∼70 K and the onset of H2O desorption at ∼140 K. This point is more clearly seen in the inset in Figure 2b, which displays the 1 ML desorption spectra for both adsorbates on the same temperature scale. These spectra show that most of the 1 ML CO2 dose has desorbed prior to the onset of H2O desorption at 140 K. Note that the H2O films were heated to 650 K, and we did not observe any high-temperature desorption peaks. The existence of such peaks could be indicative of recombinative desorption of dissociated H2O.14 The monolayer desorption spectra (red curves) in Figure 2 were analyzed using a TPD inversion procedure described in

III. RESULTS AND DISCUSSION A. Desorption Kinetics of CO 2 and H 2 O from Forsterite(011). Figure 1 displays a space-filling model view of the (011) surface of a Mg2SiO4 crystal.20 The atomic positions for the model were obtained from a published X-ray diffraction pattern on a sample in which 10% of the cation was Fe (Mg1.8Fe0.2).21 This amount of Fe is close to the amount in our sample as determined by XPS analysis. Figure 1 is intended only to be illustrative of a possible (011) surface and no information is available to create an exact surface structure as we were unable to obtain a LEED pattern because of charging 29092

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Figure 2. TPD desorption spectra for CO2 and H2O deposited on forsterite at 50 K and heated at 1 K/s. (a) TPD spectra for CO2 coverages of 0.09, 0.30, 0.50, 0.71, and 0.88 ML (blue curves); 1.0 ML (red curve); and 1.27, 1.60, 2.74, and 3.3 ML (black curves). (b) TPD spectra for H2O coverages of 0.18, 0.38, 0.60, and 0.81 ML (blue curves); 1.0 ML (red curve); and 1.25, 1.46, 1.83, 2.67, and 3.24 ML (black curves). Inset: Plot of the 1 ML spectra for both CO2 and H2O on the same temperature scale.

Figure 3. Coverage-dependent binding energy curves for CO2 and H2O obtained by inversion of the respective 1 ML spectra in Figure 2. (a) Binding energy versus coverage, E(θ), curves for CO2 obtained by inversion using the prefactor that gives the best fit to the experimental data, vBest = 1.0 × 1019 s−1 (solid line) and prefactors 10 ±2 times vBest (dashed lines, i.e., 1.0 × 1017 s−1 and 1.0 × 1021 s−1). Inset: CO2 desorption energy probability distribution, P(E) = −dθ/dE obtained by differentiating the E(θ) curve. (b) Binding energy versus coverage, E(θ), curves for H2O obtained by inversion using the prefactor that gives the best fit to the experimental data, vBest = 2.3 × 1015 s−1 (solid line) and prefactors 10 ±2 times vBest (dashed lines, i.e., 2.3 × 1013 s−1 and 2.3 × 1017 s−1). Also shown in the plot are calorimetric measurements (solid circles) of the heat of adsorption for H2O on nanocrystalline forsterite powders.7 Inset: H2O desorption energy probability distribution, P(E) = −dθ/dE, obtained by differentiating the E(θ) curve.

detail elsewhere.14,17,22−24 Briefly, the Polanyi−Wigner rate equation, dθ/dt = −νθ n exp(−E/RT), where θ is the coverage, T the temperature, E the desorption activation energy, R the gas constant, ν the prefactor, and n the desorption order, is rearranged to give, E(θ) = −RT ln[(−dθ/dt)/vθ n]. The equation is solved using the monolayer TPD spectrum for the rate, n = 1 (first-order desorption) and an assumed constant prefactor. The obtained E(θ) curve is then used to numerically integrate the Polanyi−Wigner equation to simulate the entire set of TPD spectra, and this set is compared to the experimental spectra. The process is repeated using another prefactor until the value that best fits the experimental data is determined. Representative sets of simulated TPD spectra and error analysis data are shown in the Supporting Information (Figure S3 for CO2 and Figure S4 for H2O). Figure 3 displays the coverage-dependent binding energy curves (solid lines) for CO2 (Figure 3a) and H2O (Figure 3b) obtained from the inversion procedure using a prefactor that minimizes the error between the experiment and simulation. For CO2, the optimum (best fit) prefactor was vBest = 1.0 × 1019 s−1, and for H2O, the optimum (best fit) prefactor was vBest = 2.3 × 1015 s−1. As we have observed previously,14,15 in cases in which the monolayer desorption spectrum spans a wide

temperature range, there is a wide range in the magnitude of the prefactor that can give a reasonable fit to the experimental data. For that reason, we display E(θ) curves (dashed lines) obtained with prefactors that are 10±2 times the best fit prefactor. In Figure 3a, the E(θ) curves steadily increase with decreasing coverage. In fact, the binding energy for CO2 increases by more than a factor of 2 going from a coverage of 1 to 0 ML. The inset displays the site distribution of binding energies calculated by differentiating the E(θ) curves, P(E) = −dθ/dE(θ). The distribution has a single peak centered at ∼33 kJ/mol. A similar trend is observed for the H2O E(θ) curves displayed in Figure 3b, namely, there is a steady increase in binding energy with decreasing coverage and the increase is 29093

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more than a factor of 2 going from a coverage of 1 to 0 ML. The inset shows that the distribution of binding sites is a single peak centered at ∼51 kJ/mol. Also shown in the plot are calorimetric measurements (solid circles) of the heat of adsorption for H2O on nanocrystalline forsterite powders.7 The data have the same trend as the TPD results, showing the same magnitude increase in the heat of adsorption with decreasing coverage. Because the powders present many different surface terminations, quantitative agreement is not expected; however, qualitatively the agreement is reasonable. The E(θ) curves from Figure 3 were used to numerically simulate the experimental TPD spectra. Figure 4 compares the

Figure 5. Comparison of experiments and simulations of the H2O desorption spectra for coverages of 0.07, 0.18, 0.24, 0.32, 0.40, 0.48, 0.55, 0.64, 0.76, and 1.0 ML using the E(θ) curves in Figure 3b. The experimental data are open circles. Simulated TPD spectra use the E(θ) curve in Figure 3b.

the surface and the net entropy change could possibly be greater than vmax. Heuristically, one might expect that such adsorbate-induced changes in the surface phonons are more likely in chemically complex materials like forsterite in which the silicate anion moiety (SiO44−) is “molecular” rather than “atomic” and as such will have numerous low-frequency vibrational modes.26−29 Thus, to fully account for the total entropy change, one would need to take into account the adsorbate-induced changes in the frequencies of these modes. This is an interesting idea, but further exploration of it is beyond the scope of the present paper. The best-fit prefactor for H2O is very close to the calculated value of νMax (2.3 × 1015 s−1 compared to 5.6 × 1015 s−1). However, as described above, the range of prefactors that provide reasonable fits to the experimental results does extend above νMax. The reasonably good agreement of the simulations with the experimental TPD confirms that the inversion results provide a good description of the coverage-dependent binding energies for the two adsorbates. The main result is that the binding energy for H2O (∼51 kJ/mol at the distribution peak) is stronger than that for CO2 (∼33 kJ/mol at the distribution peak). In the next section we investigate how this result affects the competitive adsorption kinetics of the two adsorbates. B. Competitive Adsorption and Displacement Kinetics for CO2 and H2O on Forsterite. Here we study the competitive adsorption and displacement kinetics for CO2 and H2O on forsterite by sequential codosing of the two adsorbates at 50 K. Figure 6a displays a series of CO2 TPD spectra in which 0.2 ML of CO2 is deposited after the deposition of various amounts of H2O (0 to 0.8 ML). With increasing amounts of predosed H2O, the CO2 desorption spectra shift to lower and lower temperature. Interpretation of these results is aided by considering two limiting cases for the adsorption of the predosed H2O: preferential adsorption on the highest binding energy sites and random adsorption. If one assumes that the preadsorbed H2O preferentially binds to the highest binding energy sites, the subsequently adsorbed CO2 molecules will not interact with those sites. For example, when 0.2 ML of H2O is preadsorbed, the binding sites where the 0.2 ML of CO2 can adsorb will have binding energies in the 0.2 to 0.4 ML coverage range of the E(θ) curve in Figure 3a. In a

Figure 4. Comparison of experiments and simulations of the CO2 desorption spectra for coverages of 0.11, 0.16, 0.25, 0.33, 0.40, 0.48, 0.57, 0.69, 0.80, and 1.0 ML using the E(θ) curves in Figure 3a. The experimental data are open circles. Simulated TPD spectra use the E(θ) curve in Figure 3a.

experimental (open circles) and the simulated spectra (solid lines) for CO2 using the E(θ) curve calculated with the best-fit prefactor (vBest = 1.0 × 1019 s−1). Of course, simulations of the 1 ML spectrum reproduce the experiment exactly because these data were used in the inversion procedure. The submonolayer simulation results provide a good description of the experimental TPD. Figure 5 displays the analogous results for H2O using the best-fit prefactor (vBest = 2.3 × 1015 s−1). Again, the agreement between the experiment and simulation results is excellent. It is interesting to compare the best-fit prefactors obtained above with the values one might expect considering two extreme case scenarios.23,25 The maximum expected prefactor, vMax, can be calculated assuming that the adsorbate is fully hindered in all translational and rotational modes. This results in the maximum entropy change going from the surface to the gas phase. The minimum prefactor, v2Dgas, is calculated assuming that the adsorbate has the translational and rotational degrees of freedom of a two-dimensional gas on the surface. This results in the minimum change in entropy going from the surface to the gas phase. For CO2, the calculated prefactors were v2Dgas = 1.1 × 1013 s−1 and vMax = 1.7 × 1016 s−1 (calculated at T = 86 K), and for H2O, the prefactors were v2Dgas = 1.7 × 1013 s−1 and vMax = 5.6 × 1015 s−1 (calculated at T = 162 K). The best-fit prefactor for CO2 is ∼600 times greater than the calculated vMax. The calculation of v2Dgas and vMax take into consideration only changes in the adsorbate itself. One could imagine a case in which an adsorbate induces some change in 29094

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accounts for the reduced desorption area as the water dose increases. Simulations in which the preadsorbed H2O was assumed to be randomly deposited on the surface are displayed in Figure 6c. Of the two limiting cases described above, the model in which H2O is assumed to preferentially adsorb on the highestenergy binding sites (Figure 6b) best captures the overall qualitative behavior of the experimental TPD spectra. The reason that the simulation results are not a quantitative match to the experiment may be because the two models represent the extreme cases and there is the possibility that the preadsorbed H2O does not block all of the highest-energy binding sites. Another possible explanation for the nonquantitative agreement is that the E(θ) curve used in the simulations was obtained from a TPD spectrum for the desorption of pure CO2. However, in these experiments, CO2 is coadsorbed with H2O. The interactions of CO2 with H2O will likely affect the “real” E(θ) curve; therefore, quantitative agreement with a model that does not include adsorbate−adsorbate interactions is not expected. Nonetheless, the simulation results in Figure 6 qualitatively capture the overall experimental behavior, and we conclude that the H2O has occupied most of the high-energy binding sites prior to CO2 desorption. Additional insight into whether the preadsorption of H2O occurs preferentially on high-energy binding sites or randomly is displayed in Figure 7. In Figure 7, the experimental CO2

Figure 6. (a) Series of TPD spectra in which 0.2 ML of CO2 is deposited after the predosing of various amounts of H2O (0.0, 0.2, 0.4, 0.6, and 0.8 ML). All depositions were at 50 K, and the heating rate was 1 K/s. (b) Simulations using the CO2 E(θ) curve obtained in Figure 3a and a model that assumes that highest-energy binding sites equal to the amount of predosed H2O are blocked. (c) Simulations using the CO2 E(θ) curve obtained in Figure 3a and a model assumes that the predosed H2O occupies random binding sites.

Figure 7. Plot of the coverage-dependent desorption energy for the experimental CO2 TPD spectra from Figure 6a (red lines) obtained using the inversion procedure with a prefactor of 1.0 × 1019 s−1. These curves are truncated below a coverage of 0.02 ML, which corresponds to desorption in the low signal−high noise far end of the hightemperature tails of the TPD spectra in Figure 6a. Also shown is the E(θ) curve for CO2 desorption from forsterite (blue line) from Figure 3a. Each red curve is from the desorption of 0.2 ML of CO2 after the adsorption of various amounts of H2O (see labels). The black scale bar shows a span of 15 kJ/mol in 5 kJ/mol increments.

sense, the E(θ) curve is shifted to the left where the new zerocoverage limit for the binding energy of CO2 begins at the amount of predosed H2O. Simulations in which the highestenergy binding sites in the E(θ) curve (the vMax curve from Figure 3a was used) were blocked by various amounts of preadsorbed H2O are displayed in Figure 6b. The other extreme is if the preadsorbed H2O adsorbs on random sites on the surface. In this case, the entire energy range of binding sites is available for CO2 adsorption but there are fewer sites. This changes the effective coverage of CO2. For example, when 0.2 ML of H2O is preadsorbed, the effective coverage of 0.2 ML of CO2 becomes 0.25 ML. The effective coverage is given by the formula, θEffective = θCO2/(1 − θH2O). In a sense, the adsorbed CO2 would sample a wider range (i.e., higher coverages) of the E(θ) curve than it would without the preadsorbed H2O. In order that the simulation TPD spectra have the same integrated area (the CO2 dose is the same in all cases), the final rate has to be multiplied by θ/θEffective, which

desorption spectra in Figure 6a are inverted using a prefactor of 1.0 × 1019 s−1 and plotted versus the total adsorbate coverage, i.e., the sum of the amount of preadsorbed H2O and the CO2 dose (0.2 ML). The E(θ) curves for each experiment (red curves) are shown along with the E(θ) curve for CO2 desorption from forsterite (blue curve) in Figure 3a. As should be the case, the 0 ML H2O dose experimental curve aligns well with the blue curve in the 0−0.2 ML total coverage region. (The small differences are due to typical day-to-day 29095

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cases CO2 is not desorbing after the H2O overlayer which begins to desorb above ∼140 K. When the desorption of CO2 from H2O (Figure S5 in the Supporting Information) and forsterite (Figure 4) are compared, one sees that the saturation coverage spectrum for CO2 on forsterite is much broader than the one for desorption from the H2O layer. (A direct comparison of these two spectra is shown in Figure 10.) A comparison of the binding energies of CO2 on forsterite and on 1 ML of H2O is shown in the Supporting Information (Figure S6). These results show that the binding energy of CO2 on H2O is lower than that on forsterite; however, because of the compensating nature of the prefactor and desorption energy, the desorption peak for 1 ML of CO2 on either surface occurs at about the same temperature (see Figure 10). The E(θ) curves in Figure S6 show that at lower coverages the CO2 binding energy on forsterite increases more rapidly than on H2O, which accounts for its broader TPD spectrum and its desorption trailing edge that extends to higher temperature. The shift to lower temperature and the narrowing of the desorption peak in Figure 8 is consistent with more of the CO2 desorbing from on top of H2O with increasing H2O dose. Finally, one can invert the pairs (CO2 first vs CO2 second) of CO2 TPD spectra in Figure 8 and show that the resulting coverage-dependent binding energy (E(θ)) and site probability (P(E)) curves are very similar. We do this for the 1 ML H2O dose spectra in Figure 8, and the results are given in the Supporting Information (Figure S7). The corresponding H2O TPD for the experiments in Figure 8 (shown in Figure S8 in the Supporting Information) are unaffected by the dose order and are similar to those in Figure 5, indicating that complete CO2 displacement−desorption precedes the onset of H2O desorption. Despite the small difference in a few of the peak lineshapes, the desorption of CO2 is largely independent of dose order. The results in Figure 8 demonstrate that an incoming CO2 will not displace a preadsorbed H2O but an incoming H2O will displace a preadsorbed CO2. If this were not the case we would have expected the CO2 to desorb only after the desorption of the H2O overlayer at higher temperatures (>140 K), which is not the case. The CO2 desorption in Figure 8 does not begin until above 75 K; therefore, these experiments are blind to the onset of the displacement kinetics. In the next section we explore the onset temperature for the displacement kinetics. C. Dynamic Displacement Kinetics for CO2 and H2O on Forsterite. Molecular beam experiments were conducted to determine the temperature for the onset of H2O/CO2 displacement. In these experiments, the forsterite substrate was exposed to a 1 ML equivalent dose of CO2 at a given isothermal temperature. The substrate was then exposed to a molecular beam of H2O and the amount of desorbed CO2 was measured. Figure 9 displays the results for a series of isothermal measurements at temperatures from 52 to 140 K. The H2O beam was started at 10 s (indicated by vertical dashed arrow) and kept on for 60 s at a flux of 0.10 ML/s. The observations are best described by considering three temperature regimes. At low temperature, from 52 to 77.5 K (blue lines), very little to no CO2 desorption is observed. In this regime even if CO2 is displaced to lower-energy binding sites on forsterite and/or moved on top of H2O, the temperature is too low for desorption. In the temperature range from 82.5 to 100 K (black lines), desorption begins just after exposure to the H2O beam with a large fraction of the initial CO2 dose being displaced. In this case the temperature is high enough for the CO2 displaced

experimental variations in the TPD spectra.) For the experiments with preadsorbed H2O, in all cases the desorption energy is initially close to the blue curve (within 10%) and then increases as the CO2 coverage decreases (the maximum deviation is ∼30% for the 0.2 ML H2O dose experiment). The upturn is likely due to increasingly attractive CO2/H2O interactions when the ratio of H2O to CO2 becomes large at low CO2 coverages. Despite this, the envelope of the initial CO2 coverage point (0.2 ML) on the red curves closely follows the blue curve and demonstrates that the preadsorbed H2O occupies the most favorable binding sites on the forsterite surface. The fact that the red curves extracted from the inversion of the TPD curves in Figure 6a lie slightly above the blue curve indicates that there is a measurable attractive interaction between the CO2 and the preadsorbed H2O. A further test of the preferential binding of H2O to the highest-energy binding sites is displayed in Figure 8. Displayed

Figure 8. TPD spectra for 1 ML of CO2 deposited before (red lines) and after (black lines) the deposition of various amounts of H2O (0.0, 0.2, 0.4, 0.6, 0.8, and 1.0 ML). All depositions were at 50 K, and the heating rate was 1 ML/s.

are CO2 TPD spectra in which 1 ML of CO2 was deposited either before (red lines) or after (black lines) the deposition of various amounts of H2O (0 to 1.0 ML) at 50 K. For a given H2O dose, the CO2 TPD spectra for the two cases are similar, i.e., the spectra are largely independent of the dose order. The small difference in the TPD spectra for H2O doses of 0.4 ML and less are due to the run-to-run experimental variation. For H2O doses of 0.6 ML and above, there are some differences in the line shape which we discuss below. Despite this, the overall picture is that with increasing H2O dose, there is a sharpening of the TPD line shape and a small shift to lower temperature. For reference, the TPD spectra for various coverages of CO2 on 1 ML of H2O deposited on forsterite are shown in the Supporting Information (Figure S5). Those results show that for a saturation coverage of CO2 (0.90 ML), there is a desorption peak at ∼79 K, which is slightly lower than the ∼85 K peak observed for CO2 on forsterite (see Figure 4). At higher coverages a peak at ∼85 K emerges, which we attribute to multilayer CO2 desorption. As mentioned above, there is a small difference in the peak shape for the two cases (CO2 first vs CO2 second) that develops with increasing H2O dose. The small difference in the peak shape may indicate a slight preference for CO2 to desorb from on top of H2O in the case when it is dosed after H2O. However, it is clear that in both 29096

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surface is narrower. The TPD spectrum corresponding to the 52 K beam experiment has a broad desorption peak that is centered near 125 K and extends out to ∼160 K. There is a small peak near 85 K that is consistent with CO2 desorption from on top of H2O. The higher temperature desorption (>100 K) is due to CO2 that is trapped under a H2O layer.30−32 The trapped CO2 desorbs when the water overlayer starts to desorb. The 65 K TPD is similar to the 52 K TPD, but with more desorption at 85 K. The desorption of CO2 in the 80−90 K range is consistent with desorption from on top of H2O (see Figure S5 in the Supporting Information). This trend continues for the 72.5 and 77.5 K spectra which show that most of the CO2 desorption occurs from on top of H2O. For the 82.5 K spectrum and higher there is very little CO2 desorption, meaning that nearly all of the CO2 desorbed during the beam exposure experiments in Figure 9. These results show that a significant amount of displacement of CO2 by H2O begins around 65 K where the desorption of both CO2 from weak binding sites on forsterite and on top of H2O is observed. Above 80 K nearly all the CO2 is displaced by H2O during exposure to the H2O beam flux. Figure 11 is a summary of the isothermal beam exposure experiments in Figure 9 and the postbeam TPD experiments in

Figure 9. Series of CO2 desorption rate versus time curves at various isothermal temperatures (72 to 140 K). At a given isothermal temperature, the substrate was exposed to a 1.0 ML equivalent dose of CO2 followed by exposure to a H2O beam (0.10 ML/s) for 60 s. The vertical dashed arrow marks the onset of the H2O beam exposure at 10 s.

by the H2O to desorb. From 110 to 140 K (red lines), the amount of beam-induced CO2 that desorbs decreases with increasing temperature. At these temperatures the displaced CO2 can desorb but the desorbed amount is less because the initial amount of CO2 that can be deposited at these higher temperature is less. The beam results in Figure 9 show that at 82.5 K and above, H2O displaces CO2, but below 77.5 K we do not know if the displacement has occurred because the temperature is too low to observe CO2 desorption. To address this question, postbeam TPD experiments were conducted. Figure 10 displays TPD spectra (1 K/s) obtained after the beam exposure experiments in Figure 9 for several of the isothermal experiments. For reference, the bottom two TPD spectra are for 1 ML of CO2 from H2O free forsterite (green dashed line) and from a 1 ML H2O film on forsterite (green solid line). As mentioned above, the desorption peaks are close in temperature, but the spectrum from the H2O

Figure 11. Amount of CO2 desorbed during the beam exposure experiments in Figure 9 (red squares) and the post-beam TPD experiments in Figure 10 (blue circles). There are more temperatures than shown in Figures 9 and 10. The sum of these two quantities is also shown (black diamonds). Note that CO2 displacement or desorption is complete prior to the onset of H2O desorption at 140 K.

Figure 10 (plus additional data points for experiments not shown in those figures). The amount of beam-desorbed CO2 (red squares) and post-beam TPD CO2 (blue circles) are plotted versus isothermal temperature. Also plotted is the total amount of desorbed CO2 (sum of beam-desorbed and postbeam TPD). Below ∼75 K no CO2 desorbs during the beam experiment; however, as described above, near 75 K CO2 can be displaced by H2O. Above 75 K, an increasing fraction of the CO2 is displaced by H2O (see Figure 10) and a fraction of the CO2 on top of H2O can desorb (see Figure 9). As the temperature increases above 80 K, nearly all of the displaced CO2 can desorb during the H2O beam exposure; thus, the amount in the post-beam TPD decreases (rapidly above 75 K, as shown in Figure 11). The total amount of CO2 decreases above 75 K because the amount of CO2 that can be adsorbed decreases with increasing temperature, consistent with the TPD spectra displayed in Figure 4.

Figure 10. Postbeam CO2 TPD spectra conducted after the isothermal beam exposure experiments in Figure 9 for selected isothermal temperatures. After the 60 s H2O beam exposure (see Figure 9), the sample was cooled to 50 K and a TPD spectrum at a heating rate of 1 K/s was obtained. The bottom two TPD spectra are for 1 ML of CO2 from forsterite (green dashed line) and from a 1 ML H2O film (green solid line) deposited on forsterite at 50 K and are shown for reference. 29097

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The beam-induced CO2 desorption observed in Figure 9 could in principle be the result of either direct desorption into the vacuum or CO2 that desorbs thermally after being displaced to a lower-energy binding site. Figure 12 displays a series of

Figure 13. Amount of adsorbed H2O (blue circles) at the onset of CO2 desorption versus the initial CO2 dose for the beam experiments at T = 85 K in Figure 12. The H2O amount was calculated using the water beam flux (0.10 ML/s) and the onset delay time. A linear fit to the data (solid line) yields a slope of −1.5 ML of H2O per ML of CO2. Also plotted are the results (red circles) for beam experiments at T = 87.5 K (not shown). These data have a slope of −1.4 ML of H2O per ML of CO2.

Figure 12. Series of CO2 desorption rate versus time curves for various doses of CO2 deposited at 85 K and then exposed to a H2O beam (0.10 ML/s) for 60 s. The initial CO2 coverages (from bottom to top) were 0.06, 0.12, 0.23, 0.34, 0.45, 0.54, 0.59, 0.61, and 0.63 ML. The vertical blue arrow marks the start of the H2O exposure at t = 10 s.

are occupied, the CO2 is pushed into lower-energy binding sites from which the onset of desorption is observed. The slope gives the coverage of H2O per CO2 needed to replace the CO2 and push it into a site from where it can desorb. However, given that we have defined 1 ML for both CO2 and H2O based on their saturation coverages on Pt(111), we would expect the slope to be one. The higher value of ∼1.5 suggests that the CO2 can be compressed in the surface layer and perhaps can adopt a different surface configuration to allow it to remain in the first layer as H2O is added. For example, if upon H2O adsorption CO2 underwent a structural change from lying flat to being upright, the amount of CO2 that could remain on the surface would increase. Such a change in the CO2 adsorption geometry could potentially be probed using vibrational spectroscopy. While these issues clearly warrant further investigation, they are beyond the scope of the present study because they will require detailed structural information not directly provided by desorption and/or scattering alone.

CO2 desorption versus time spectra in which various initial amounts of CO2 (0.06 to 0.63 ML) are deposited at 85 K and then exposed to a H2O beam having a flux of 0.10 ML/s for 60 s. At this temperature, CO2 displaced into another forsterite site will remain surface-bound whereas CO2 displaced into the second layer (either on top of H2O or another CO2) will desorb promptly. The vertical dashed arrow marks the onset of the H2O beam exposure at 10 s. The bottom spectrum has an initial 0.06 ML dose of CO2 and shows that the onset of CO2 desorption is delayed more that 10 s after the start of the H2O beam exposure. The delay time before the onset of CO2 desorption decreases with an increase in the initial CO2 dose. The observed delay rules out a direct desorption mechanism, otherwise we would expect to see CO2 desorption at the onset of the H2O beam. Instead, these results support a desorption mechanism in which the incident H2O displaces a CO2 from a high-energy binding site but the CO2 molecule remains bound to forsterite on a lower-energy binding site. This process continues until a sufficient amount of H2O has adsorbed to push the CO2 into a less strongly bound forsterite site or into the second layer whereupon it thermally desorbs at a rapid rate. The fact that most of the CO2 desorption occurs after more than 1 ML of H2O is deposited (above 20 s) suggests that some if not all of the CO2 is desorbing from on top of H2O in contact with forsterite. Figure 13 displays the amount of H2O adsorbed (beam flux × exposure (delay) time) at the onset of CO2 desorption versus the initial CO2 dose for the spectra in Figure 12 (blue circles). Also shown are additional results (beam experiments not shown) in which the substrate temperature was held at 87.5 K (red circles). In both cases the amount of H2O required to induce CO2 desorption decreases linearly with initial CO2 dose with a slope of −1.5 for the 85 K data and a slope of −1.4 for the 87.5 K data. These results suggest that at a given isothermal temperature there are a finite number of sites where CO2 can remain bound to the forsterite surface. When all of those sites

IV. SUMMARY The adsorption, desorption, and displacement kinetics of CO2 and H2O on a natural forsterite crystal surface, Mg2SiO4(011) with 10−15% of Fe (substitutional in Mg lattice positions) have been studied using TPD and molecular beam techniques. In Section IIIA, Desorption Kinetics of CO2 and H2O from Forsterite(011), the desorption spectra show that both adsorbates preferentially bind to the highest-energy binding sites available. Neither adsorbate has a resolved monolayer or site-specific desorption peaks. Instead, analysis of the desorption spectra of both species shows a continuous coverage-dependent binding energy that smoothly increases with decreasing coverage. This is in contrast to our prior study on TiO2(110) in which specific adsorption sites had wellresolved desorption peaks.14 The results here show that the forsterite surface is more complex with numerous adsorption sites (Mg, Si, O, and Fe) and specific site energies cannot be resolved in the TPD spectra. These results are in qualitative agreement with calorimetric data on forsterite powders7 that 29098

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The Journal of Physical Chemistry C also show similar coverage-dependent behavior for H2O. The desorption studies and analysis show that the binding energy of H2O is greater than that for CO2 over the entire monolayer coverage range. In Section IIIB, Competitive Adsorption and Displacement Kinetics for CO2 and H2O on Forsterite, we used sequential coadsorption to study the competitive adsorption and displacement kinetics. We find that largely independent of dose order, H2O preferentially occupies the highest-energy binding sites and will readily displace CO2 from an occupied site. This is consistent with our previous work in which we found the same behavior for CO2 and H2O adsorption on TiO2(110).14 Simply put, the highest binding energy adsorbate displaces a lower binding energy adsorbate. We also present data that indicate the binding of CO2 to the forsterite surface is slightly enhanced by coadsorbed H 2 O because of attractive H 2 O/CO 2 interactions. In Section IIIC, Dynamic Displacement Kinetics for CO2 and H2O on Forsterite, beam experiments show that the onset of significant displacement occurs between 65 and 75 K. We also find that a displaced CO2 remains bound to the forsterite surface until a sufficient amount of H2O is adsorbed to “push” the CO2 to a lower-energy site from which it thermally desorbs in a facile manner. A coverage ratio of 1.5 H2O to CO2 is needed to fill the area of the missing CO2 before the CO2 is pushed into the lower-energy site for desorption. This may be the result of compression or reorientation of the CO2 in order to stay in contact with the forsterite surface. These results provide the fundamental adsorbate interaction energetics for CO2 and H2O on the complex forsterite, Mg1.8Fe 0.2SiO4(011), surface. We find no experimental evidence for the dissociation or reaction of either CO2 or H2O on forsterite. Of course, current experiments in which forsterite is used as a model substrate for carbonation to form MgCO3 are at elevated pressures (>100 atm) and temperatures (>300 K) that cannot be replicated in our ultrahigh vacuumbased experiments.5−11 However, theoretical studies that are trying to determine the fundamental energetics should be able to use our data as a benchmark for the accuracy of their calculations and adsorption models. Of particular interest is the energetics for determining the structure of the CO2 and H2O at the forsterite surface, and our results should be of value in this effort.

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ACKNOWLEDGMENTS

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REFERENCES

This work was supported by the U.S. Department of Energy (DOE), Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences, and Biosciences. The research was performed using Environmental Molecular Sciences Laboratory, a national scientific user facility sponsored by DOE’s Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory, which is operated by Battelle for the DOE under Contract DE-AC05-76RL01830. The authors thank Shawn Chatman and Kevin Rosso for providing the natural forsterite sample and Tamas Varga and Mark Engelhard for the x-ray diffraction and x-ray photoelectron characterization, respectively.

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ASSOCIATED CONTENT

S Supporting Information *

Eight additional figures that are not necessary for the overall understanding of the scientific arguments presented in the main paper but may be of interest to some readers: figures on the characterization spectra of the forsterite substrate, details on the inversion analysis presented in the main text, and additional TPD spectra and analysis. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*Phone: (509) 371-6156. E-mail: [email protected]. *Phone: (509) 371-6143. E-mail: [email protected]. Present Address †

Z.L.: Evans Analytical Group, Sunnyvale, CA 94086.

Notes

The authors declare no competing financial interest. 29099

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