The interactions of Co, Mn and water with calcite surfaces

..(.. .,.,.,.,. ,.,~ ..........:.l ~~ .:.:.

.‘:.:.:.:.:.:‘..:.‘.:.s: . . .. . . . . . .:.:.:.: . .. . . . . . .. . . . .A...... ..>: . .. . . . . ..i..... :.: . . . .. . . . :.:(.:

.::.:j:::::~::::::::jj:::::‘:::;;.:.: BI:$.: i.....” ‘.

Surface Science 276 (1992) 27-39 North-Holland

surface

.

.

.

. .

.

.

.

.

. .

.

.

.

~~

.. ...........i. :.: .......~.:.:.: .,:.:.:.:~:.:.: .’ .. .... ....~::.::::::

science

i::.>.; ... ‘.‘.““...‘.A :...: :::::!:I ““““.‘.‘. .,.,.,.( .:‘..:.:.:.:..: ::::;:::~:~:: :::.:.. ‘:‘:“‘::: .G’.........i.........: ...........1.,.. .,., ,,(, “‘i.‘.‘;.:‘::.:.:.:.:. ......:.~ i,., ,,,,,, “‘:‘:‘.:.:x:::l:~::ii~~~~:~:~~~~~~:~:~:~: :.:.:. j::. : ,.,,., ‘.

The interactions

of Co, Mn and water with calcite surfaces

David L. Blanchard, Jr. and D.R. Baer Molecular Sciences Research Center, Pacific Northwest Laboratory *, Richland, WA 99352, USA

Received 21 February 1992; accepted for publication 30 April 1992

The interactions of Co, Mn and water with surfaces of single crystal calcite (CaCO,) have been examined using surface analysis techniques (primarily X-ray photoelectron spectroscopy, XPS) and a specimen transfer system. The metals were deposited on the calcite by evaporation to give submonolayer concentrations. For coverages of Co as low as 0.16 ML, the metal clearly appears in both metallic and oxidized states. Immersion in deionized water completely converts all Co to a compound most consistent with the hydroxide. As previously observed .for Mn, some of the Co is removed from the surface during immersion. The rate of removal of Co is two times faster than that of Mn. Further, these rates are about three orders of magnitude slower than that reported for the dissolution of bare calcite. Recently reported solution experiments suggest that Mn cations are adsorbed substitutionally from aqueous solution while most other metals, including Co, are adsorbed in a hydrated form. The significance of our results in light of the conclusions of these solution experiments is discussed.

1. Introduction Oxide and carbonate compounds in their various forms are by far the major components of soils. Many ions, organics and other species are naturally present in groundwater percolating or flowing through soils. Additional species such as organics and/or heavy metal ions may enter an aquifer from a variety of artificial (man-made) sources. Both natural and man-made components of the groundwater solution may be adsorbed onto (or exchange with ions in) the adjacent mineral surfaces. Because of this interaction, these surfaces help determine the rate and direction these species move through the earth. Clearly then, knowledge of various characteristics of the oxide/ solution interaction is necessary to predict the fate of contaminants released into the earth. An accurate picture of this interface will require the elucidation of rates of adsorption, desorption, precipitation and dissolution under many different conditions, and detailed descriptions of spe* Operated for the US Department of Energy by Battelle Memorial Institute under contract DE-AC06-76RL0 1830. 0039-6028/92/$05.00

ties and surface sites involved. Several reviews of mineral surface chemistry have been published [l-3]. Most of the information about the surface chemistry of minerals has been gathered from solution studies. Roughly one sixth of the global sedimentary mass is composed of carbonate minerals [4,51. Calcite, the rhombohedral form of calcium carbonate, is a principle contributor to this portion of the Earth’s solid material. Because of this abundance in soils and sediments, and because of the reasons for the interest in the mineral/ solution interface outlined above, calcite in particular has been the subject of a number of studies. Much of a review (1983) of the mineralogy and geochemistry of carbonates [6] was devoted to calcite. Since then, there have been a number of studies of the interaction of calcite with aqueous solutions of metal ions [7-171 and organics [lS]. Among these studies is one [16] which shows that at constant total ionic strength, divalent metal cations adsorb on the calcite surface below saturation with the metal carbonate. The extent of adsorption depends on pH. The desorption of the same metal ions from the surface shows quite

0 1992 - Elsevier Science Publishers B.V. All rights reserved

28

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

interesting behavior. One group of metals exhibits almost complete desorption (i.e., - 80% of the adsorbed metal readily desorbs) while another group appears to remain incorporated in the lattice to a large extent (only N 20% desorption). The authors suggest that ions in the former group form surface hydration complexes, while those in the latter are actually incorporated into the calcite lattice. A basic molecular level understanding of the reactions and species involved is lacking. To begin to fill this gap, we have undertaken a study of various calcite surfaces and their interactions with water and several metals using UHV techniques. The natural occurrence of calcite in single crystal form makes it ideal for such UHV studies. We have previously examined the bare calcite surface prepared in several different ways, including fractured in air and UHV, sputtered, exposed to water vapor and immersed in water [19]. (Stipp and Hochella [20] have also recently examined bare calcite, with generally similar results.) A significant result of our previous work is the removal of CaO (formed at the surface by sputtering) when the sample is dipped in deionized (DI) water. The XPS spectrum of the resulting surface is almost indistinguishable from that of a cleaved one. In that study we also examined the interaction of the surface with different amounts of evaporated Mn, and the interaction of this Mn coated surface with water. We observed the formation of an oxidized Mn complex upon water immersion. The photoline peak shapes and positions were consistent with the formation of MnCO,. Zachara et al. [16] found manganese to be a member of the group of metals that does not readily desorb from the calcite surface following sorption from solution. These authors suggest that a CaCO,-MnCO, solid solution is formed at the surface upon adsorption of the Mn. The results of our UHV study are consistent with these conclusions from solution studies and demonstrate the utility of combining macroscopic studies with molecular level ones. In this paper we examine the interaction of Co metal with the bare calcite surface. We also compare the rates of dissolution of Mn and Co into water upon immersion of a calcite surface with a

vapor deposit of Co and/or Mn. Cobalt is a member of the group of metals which were found to reversibly adsorb on the calcite surface. Because of this ease of desorption, it was suggested [16] that it forms a hydrated complex at the surface, instead of a solid solution. For this reason the two metals are expected to show significantly different surface complexes and dissolution kinetics. Surface sensitive methods offer the possibility of obtaining molecular level information about the chemical-geochemical interactions that occur at the mineral-solution interface. The specific goals of this paper are to determine if Co and Mn produce different types of surface complexes, to identify the likely composition of the Co surface complex, and to measure rates of dissolution of the metals from the surface into water. The application of vacuum based analysis methods in combination with specimen transfer capability allows complex geochemical interactions to be broken into components for detailed examination and verification. Information learned at the current level of experiments can both relate to work at more complex levels and suggest other detailed studies to answer specific questions. Results from the current experiments provide, therefore, both a somewhat more detailed understanding of complex measurements and indications of specific additional experiments that can further clarify reaction mechanisms or processes.

2. Experimental 2.1. Instrumental parameters All XPS spectra of Co or Mn on calcite were taken using a Perkin Elmer Physical Electronics (Eden Prairie, MN) 560 XPS/AES system, which uses a model 25-270 double pass cylindrical mirror analyzer (CMA). Cobalt carbonate and cobalt hydroxide standards (powders) were run on a 550 XPS/AES system. This instrument, similar to the 560, is also made by Perkin Elmer and also uses a double pass CMA (model 25-260). For both instruments, the incident X-ray beam and CMA axis are perpendicular. The calcite sample bars

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

(next section) were mounted such that the sample axis was roughly 30” away from the CMA axis toward the X-ray source. Due to the cleavage geometry of calcite, the surface normal was rotated 15.03” away from the specimen axis. This put the surface normal approximately 42” from the CMA axis toward the X-ray source and lo” to the side [21]. All data reported here were taken in the angle integrated mode. The lowest system pressure recorded for the 560 system during the course of these experiments was 3.7 X lo-” Torr. Normal operating pressures were typically 1 X lop9 Torr, with a maximum of 1 X lo-’ Torr during metal evaporation. A transfer system connects the 560 main analysis chamber to an electrochemical cell. This permits movement of samples to and from UHV for dipping in aqueous solutions. (In all cases the water for dipping was sparged with nitrogen gas to remove solution CO,.) Gas exposure is limited to nitrogen and water vapor. Non-monochromatized Mg Ka Xrays were used for measurments in both systems. Energy scales were calibrated to the Cu 2p3,2 and Au 4f 7,2 photolines at 932.67 and 84.00 eV, respectively. Survey spectra were collected at 100 eV pass energy, and multiplex data at 25 eV. The insulating nature of the samples required charge correction of the energy scale. For the calcite samples, this correction was made by assuming that the CO, carbon peak is located at 289.5 eV, as generally reported in the literature (ref. [19], and references therein). For powders an adventitious carbon peak at 284.8 eV was used for the correction. The pressure during the powder experiments in the 550 system was roughly 5 x lo-’ Torr. 2.2. Materials and methods All water used for dipping was deionized with a MilliQ system (Millipore Corp., Bedford, MA). The resulting water displayed a resistivity of 15 MR/cm or greater. An optical grade single crystal of geologic calcite was obtained from Commercial Crystal Laboratories, Inc. (Naples, FL) for this study. Bars of nominal dimension 7 x 4 x 4 mm3 were cleaved from the main crystal for mounting in vacuum. Calcite cleaves along the

29

{lOi4} faces, as referenced to the hexagonal cell. The surface was prepared for vapor deposition of the metal in the following manner. The face was sputtered with Ar ions at 1 kV energy for 5-10 min to remove surface contamination. This was followed by immersion in DI water for 2 min to remove sputter damage. The resulting surface gives an ESCA spectrum nearly indistinguishable from the cleaved surface - peak positions, shapes and area ratios are all nearly identical, as previously reported [19]. This surface was used to determine an area sensitivity factor for our instrument for the Ca2p peak pair. (Following the work of Wagner et al. [221, the Ca to CO, carbon ratio was assumed to be 1 and the area sensitivity factor for C was taken to be 0.25.) Cobalt was evaporated in vacuum from a crumpled 0.003” sheet (99.9975% purity) from a crucible evaporator as previously described [19]. The purity analysis provided with the Co sheet listed iron and silicon as the major contaminants. Neither was detected in any of the deposited films with XPS or AES throughout the experiments. However, in all cases of Co deposition on the surface, Mn was also present at slightly less than half the Co coverage. This was due to previous use of the evaporator for deposition of Mn. As higher temperatures are required for Co evaporation, some Mn was removed from non-source areas of the evaporation system and deposited with the cobalt. Because of our interest in comparing the behavior of the two metals, we investigated further. We found that the Mn XPS peak shapes and positions for a submonolayer coverage of just Mn on a calcite surface [19] are in reasonable agreement with those observed in this study in which Mn and Co are present simultaneously, and the dissolution rate of Mn from the calcite surface into water is the same in both cases. These two observations lead us to believe that the two metals behave independantly under the conditions of low coverage in this study. Although initially deposited together on the calcite unintentionally, this provides an excellent opportunity to compare the behavior of the two metals under precisely identical conditions. Sputter cleaned copper was used to characterize the vapor deposition as described in section 3.

30

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, h4n and water with calcite surfaces

Cobalt carbonate and cobalt hydroxide powders were run for comparison, also described below. In addition, the CoCO, was used to determine the XPS area sensitivity factor for the Co2p,,, peak in the same manner given above for Ca. This is particularly important because the sensitivity factor for this cobalt photopeak is reported to vary greatly with chemical state [23,24]. For the dissolution experiment, a calcite surface with either pure Mn or a combination of Mn and Co was immersed in DI water for a specified time interval, then returned to vacuum. 2.3. Data analysis Peak positions were determined after charge correction and smoothing (seven point SavitskyGolay). Peak areas were then determined after a Shirley inelastic background subtraction [25]. The metal overlayer coverages were calculated using these areas and assuming two adsorption sites per CaCO, unit 1191. The Co (or Mn) 2p3,2 and Ca2p photopeaks were used for the calculation. The derivation of Seah 1261 was followed to determine the appropriate expression for a single component deposit on a two component homogeneous substrate. (The CO, polyanion is taken to be one component.) For submonolayer coverages I,,/&, = $ASFc,(I

- exp[ -acO/Ac,(&J

cos 01)

x O.SASFc,{ 1 - $J + $J Xexp[ -acO/AcO(&J

cos

f4},

(1)

where 1, is the area (in counts per second) of the peak of interest of element x, ASF, is the area sensitivity factor for that peak, 4 is the fractional coverage, a,, is the thickness of a monolayer of Co (taken to be the atomic diameter of Co as calculated from the density and molar mass of the pure crystalline metal), and h,(E,) cos f3 is the escape depth [26] in medium x for electrons of kinetic energy E, (which is characteristic of the Auger transition of interest of element y). The 0.5 factor is introduced in the denominator because the calcium ions account for only half of the possible surface sites. This equation may be

solved for 4; all other quantities are known or can be calculated. The coverages calculated using this equation are surprisingly linear in I&Z,,. For coverages greater than one monolayer, the following expression was used:

=

ASFc,{l-

exp[ -dc,/&,(

/O.SASFc, exp[ -d,,/h,,(

Ec,) cos e]} E,-,,) cos 131. (2)

Constants and variables are as defined for eq. (1) above, and do, is the thickness of the overlayer in nanometers. This equation does not have an analytic solution - it must be solved iteratively. A readily available spreadsheet program was used on an IBM PC (Boca Raton, FL) to calculate the coverage using eq. (1). If the result was greater than 1.0 ML, eq. (2) was used instead. Using hypothetical peak areas, the two equations agree to within 1% for coverages from 0.98 to 1.02 ML. The mean free path, A, of electrons with energy E, in a metal x was calculated using the expression [26]: A,( EY) = 538a,/(

E,)’

+ 0.41( ax)3’2( Ey)“2. (3)

We estimate the uncertainty for the low coverages of this study to be approximately 20%. No attempt was made to determine simultaneously the coverage of Co and Mn when both were present on the surface. The implicit assumption is that the coverages are so low that the Ca signal is not significantly affected. Although this is not strictly the case, the procedure may be justified by the following observation. As noted, a plot of 4 co versus Ic,/Z,. gives a straight line. Because the coverage of Mn was always lower than that of Co, it was completely removed from the surface first during the dissolution experiments. No significant change in slope is observed for 4co versus Z,-.,/Ic, for those points before and after the removal of Mn. This suggests that any error in coverage due to neglect of other species present at the surface is within the experimental error, at least for the coverages considered here.

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

A

a)Ola

0.30

-537

-535

-533 Binding

-531 Energy,

-529

ML GY

-527

-525

-794

-791

-788

eV

-785

-782

Binding

Energy,

-779

-776

-773

eV

Fig. 1. (a) 0 1s photopeak for bare calcite and calcite with two Co coverages. The surface was prepared in all three cases by sputtering and dipping as described in the text. (b) Co2p,,, p hotopeak for a cobalt deposit on copper and for the same two deposits on calcite as in (a).

was used. For the low concentrations of Mn in that run, we see no significant difference between the two calculations.

One set of data for the dissolution of Mn was initially presented in tabular form in a previous publication [19]. A simpler coverage calculation

Table 1 Binding energies (eV) of samples and reference compounds Compound/sample

Cls

Ca 2~s~~

01s

CaCO,

289.4 289.4 289.6

346.7 346.8 347.1

531.2 531.3 531.4

289.5

347.0

531.4 529.6 530.2 530.2 529.6 531.2 531.2

coo

Co(OH),

coca, 0.16 ML Co on CaCo,

289.4 289.5

346.8

Previous + 30 s dip (0.07 ML remaining) 0.30 ML Co on CaCO,

289.5 289.5

347.0 346.8

Previous + 30 s dip (0.26 ML remaining) 0.36 ML Co on Cu

289.5

346.9

530.9 531.3 531.3; 528.9 531.4 531.4; 530.1; 528.9 531.4

Co2p,,,

Source ESCA DB ESCA DB Ref. A

781.3 780.8 781.3 780.5

Current Ref. C ESCA DB ESCA DB Ref. B Ref. C ESCA DB Ref. L Current Current Current

780.9 780.3

Current Current

780.9 778.1

Current Current

780.0 780.4 780.4 780.6 781.0 _

32

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

Although Mg Ka X-rays gave less peak overlap in the spectra than AlKa, there is a small peak in the same region as the Co2p,,, peak. This is part of the O(KVV) Auger transitions. To remove this source of error, it was assumed that this peak area scaled with the 0 Is peak area. The bare CaCO, was used as a reference to estimate the intensity of this peak, which was then subtracted from the Co peak.

3. Results 3.1. Vacuum surfaces The calcite crystal was prepared by sputtering and dipping as described above. Co was then deposited on the surface by evaporation. The metal clearly interacts with surface oxygen as indicated by the appearance of an extra shoulder on the low binding energy side of the 0 1s peak (fig. la). (All XPS peaks shown have been normalized to the peak maximum.) The curves, for different coverages as noted in the figure, have been aligned to remove shifts due to charging. The true charge corrected binding energies are shown in table 1. The shoulder for the low coverage (0.16 ML) curve is well accounted for by a single extra peak at 528.9 eV. Assuming that this extra peak arises from oxygen interacting with Co, it must be very near the surface. A coverage may be calculated if we further assume that it

Table 2 Coverages

(ML) of surface

species

(see experimental

forms an overlayer of 02- anions (table 2). The larger shoulder on the higher coverage (0.30 ML) oxygen peak required two extra peaks for a fit. One of these is also located at 528.9 eV, as noted in table 1. The corresponding coverages for these two peaks are also provided in table 2. When using more than two curves to fit a peak, the uniqueness of the solution must be considered. In this case, the bare calcite 0 1s curve was fitted, then this fit was scaled and subtracted from the 0 1s peak when 0.30 ML of Co was present. The residual was quite asymetric and clearly composed of several peaks. This was used as a guide for the locations and intensities of the extra peaks when the complete curve was refit. As a check, the peak locations and intensities were varied and the curve was fit again. Several variations and refits made it clear that the original fit was the best, with the bare surface peaks as the dominant contribution. As mentioned above, a single peak added to the scaled bare surface fit was sufficient for a good reproduction of the lower coverage data. This fit was done without using the 0.30 ML peak fit as a guide, and so the match in peak location at 528.9 eV gives some measure of credence to the fitting accuracy. Fig. lb shows the Co2p,,, peaks for the same two coverages on calcite, and for a comparably low coverage on sputter cleaned copper. (In this case the peaks are at their true charge corrected binding energies.) For as little as 0.16 ML of Co

for calculational

details)

Surface/treatment

0(0x)

Co(ox) h’

Cotme)

0.16 ML Previous Previous 0.30 ML Previous Previous Previous 0.28 ML

0.15 _ _

0.15 _ _

0.01 _ _

0.30; 0.27 _ _ _ _

0.24 _

0.06 _ _ _

‘) ‘) ‘) d,

Co on CaCO, + 30 s dip + 30 s more Co on CaCO, + 30 s dip + 30 s dip + 30 s dip Co on Cu

a)

From shoulder on carbonate 0 1s peak observed Oxidized Co (before dipping) Metallic Co (before dipping). Oxidized Co (after dipping).

_ _ after Co deposition.

0.28

‘)

Cothyd)

‘)

Mn(all)

0.07 0.03

0.07 0.02 _

0.26 0.12 0.04 _

0.13 0.10 0.03 _ _

33

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

on calcite, the overlayer appears in both the 2 + state and the 0 (metallic) state. This is evident from the shoulder at roughly 778 eV on a much larger peak pair with maxima at approximately 780.5 and 785.9 eV. The latter is characteristic of high spin cobalt(I1) compounds, such as CoCO,, Co0 or Co(OH), [28,291. Other possible cobalt compounds may be excluded on this basis. The observed peak binding energy is closest to those of the oxide and hydroxide (table 1). The shoulder arises from metallic Co. Other possibilities for the metallic peak can be discarded on the basis of peak location, as is evident by the overlay of the Co2p 3,2 photoline due to roughly 0.28 ML of Co on a copper substrate. Although the amount of the Co oxide increases with Co coverage, the amount of metallic Co increases more quickly. This can be seen by the relative increase in the intensity of the shoulder with increasing Co deposition. The possible reasons for this are discussed in section 4. The distinctly metallic shape and position of the 0.28 ML deposit of Co on copper is important because it indicates that the vapor deposition is clean; only metallic Co is being deposited. Therefore the Co oxide seen on the calcite is not due to accidental deposition of this species directly from the vapor at the start of the experiment. (The

uncertainty in the Co coverage on Cu, approximately f0.20 ML, is a bit large due to uncertainty in the sensitivity factors (SF’s) for the metallic species. However, all reasonable guesses for the SF’s give submonolayer coverages, and so the Co peak location and the conclusion of a clean evaporate are unaffected by this uncertainty.) 3.2. Dipped surfaces The Co-on-calcite surface (i.e., with a Co deposit) changes significantly when dipped (fig. 2). The first panel shows the calcite 0 Is peak for bare calcite, calcite with 0.30 ML of Co, and the latter surface after a 30 s immersion in DI water. After dipping, the peak looks nearly identical to that of the bare surface, although, as noted in the figure, there is still 0.26 ML of Co left at the surface. Fig. 2b is a plot of the Co2p,,, peak for the 0.30 ML Co-on-calcite surface after dipping. Also shown for comparison are the Co peaks for the 0.16 and 0.30 ML coverages before dipping. (The peak maxima have been aligned for ease of peak shape comparison.) It is immediately apparent that there is no longer metallic Co on the surface. The water immersion has completely oxidized it. The peak maximum has changed posi-

b) Co 2pm

w iii iz

0.30 (0.26

ML Co + 30s

Dip

w c P

ML Remaining)

I” -537

-535

-532 Binding

-530 Energy,

-527 eV

-525

-794

I

-791

‘.

I

-708

‘-

I

*

‘1

-785

-782

Binding

Energy,

‘.

I

-779

-

’

I

-776

“I -773

eV

Fig, 2. (a) 0 1s photopeak for bare calcite, calcite with a 0.30 ML deposit of Co, and the latter surface after water immersion for 30 s. The peak for the dipped surface is almost identical to that of the bare surface, though 0.26 ML of Co remains. (b) Co2p,,, photopeak for two coverages of cobalt on calcite and for the 0.30 ML coverage after dipping (0.26 ML Co remaining).

34

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

at their true charge corrected binding energies.) These two compounds are the most likely suspects for the Co compound at the calcite surface after dipping. Although the powders give broad lines, Co(OH), has the most similar peak shape and binding energy. However, the presence of CoCO, cannot be excluded. In addition to the chemical state changes, dipping alters the elemental composition at the surface - some of the Co is lost to the solution (table 2). This is also the case for Mn deposited on calcite and then dipped in pure water, as shown previously [19]. Fig. 4 displays a plot of the amount of remaining metal versus total time in contact with water for both metals. Dips were consecutive - the metal remaining after the first dip was used for the second and so on. Three runs with Mn and two runs with Co are shown. Although the data set for any one run is small, the slopes for different runs for the same metal agree very well. (The assumption of linearity is not a priori justified, since the first dip involves the ionization of the metal atoms in addition to hydration and dissolution, and the form expected for the dissolution of a thin film into solution is

ct! G i

-794

-791

-708

-785

-782

Binding

Energy,

-779

-776

-773

eV

Fig. 3. Co 2p,,, photopeak for powder samples of CoCO, and Co(OH),, and for a 0.30 ML coverage of Co on calcite after dipping. The latter is most similar to the Co photopeak for the hydroxide.

tion as well. The new position, at 780.9 eV, is very close to that expected for Co(OH), [28,29] (table 1). This peak is compared with the Co 2p,,, peak observed for powder standards of Co(OH), and CoCO, in fig. 3. (All peaks in this figure are

0.3

0.2

0.1

.

0.0

L

0

60

120 Dip

la0 Time

240

300

? 30

(s)

Fig. 4. Plot of the cobalt coverage on the calcite surface versus dipping time. The negative slopes for the different runs reflect the loss of Co from the surface into solution. The two runs for Co give an average slope (i.e., rate of removal) of - 0.0030 ML/s and the three runs for Mn give an average of - 0.0015 ML/s.

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

not known. However, the agreement indicates that the slopes are nearly linear, at least over the range of initial coverages investigated.) Most interesting, however, is the significantly slower average rate of removal for Mn ions. The rate of removal of the Mn is calculated to be 0.0015 + 0.0003 ML/s, and that of the Co is 0.0030 + 0.0006 ML/s, twice as fast. Note that the data for both metals span a coverage range from 0.30 to less than 0.05 ML, so the difference is not merely due to differences in coverage.

4. Discussion 4.1. Co on calcite in vacuum We consider first the interpretation of the spectra of the Co-on-calcite surface. There are clearly two states of Co as deposited on the calcite surface in vacuum. Several observations suggest that a species similar to cobalt oxide is the origin of one of these. The satellite structure p eak indicates either the oxide, of the Co2p,,, hydroxide or carbonate, as noted. The position is closest to that of Co0 (table 1). The shoulder on the 0 1s peak is a clear indication of an interaction of Co with 0. The assumption that an overlayer of 0 anions interacting with Co gives rise to this shoulder is very simplistic. Nevertheless, the amount of 0 calculated in this approximation is very close to the amount of Co observed in the oxidized state (table 2). This 1: 1 ratio is expected for COO. The position of the Co photopeak matches that seen for COO, and the binding energy for the 0 Is “oxide” shoulder, although not a perfect match, is close to that expected for COO. Further, it is at the correct end of the range of 0 1s peaks of possible oxides (table 1). Considering the difference between a bulk sample and a thin film on a calcite crystal, a shift is not surprising. There are several possible explanations for the coexistence of metallic and oxidized Co on deposition - oxidation only at defect sites, formation of islands, or an inequivalence of normal surface sites. We have obtained scanning force microscopy (SFM) data for air cleaved calcite that

35

appears to rule out selective oxidation at large defect sites. For these scans the microscope was used essentially as a profilometer to directly image the surface. A simple calculation from a 3 X 3 pm2 SFM image of the surface (not shown) indicates a maximum of 0.02 ML of step edge defect sites. (For this scan the resolution is roughly 5 A perpendicular to the surface and 40 A parallel to the surface.) The observed step density is consistent with that reported by others 1301. All other defects visible in the image (predominantly ridges that appear to be remnants of the opposite cleavage surface) contribute no more than 0.04 ML. (The SFM data will be presented and discussed in greater detail in an upcoming publication by Blanchard and Baer.) For the 0.30 ML Co coverage, there is roughly 0.24 ML of the oxide. Assuming that an air cleave is not significantly different from a vacuum cleave, this suggests that the oxide is not formed at large defect sites, such as ridges and steps. This of course does not rule out adsorption at smaller defects such as vacancies, ad-ions or -atoms, step kinks, etc. Because of the predominance of the oxide over the metal, we believe that the majority of those Co in contact with the calcite surface interact with one or more carbonate oxygens, resulting in an oxidized surface complex that looks very much like COO. The chemical state characteristic of metallic Co probably arises from atoms that adsorb on top of those that have already arrived (i.e., island formation). Islands could also form due to deposition of metal clusters on the surface. Prelimina~ analysis of the inelastic background [31] around the Co 2p photopeaks is consistent with island formation (R. Williford, private communication). The last possibility is that there are two distinct kinds of calcite surface sites for the metal, one at which the Co atoms are oxidized, and one at which they remain fairly metallic. The increase in Co metal relative to oxidized cobalt with increasing deposition is consistent with either of the latter. More data (such as additional topographic SFM images at higher resolution and with Co present, mobility of surface species, barrier to metal - oxide interconversion, etc.) are needed to say with certainty why only part of the sub-monolayer cobalt deposit is oxidized.

36

D.L. Blanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

4.2. Co on calcite after dipping

The changes due to dipping the Co-on-calcite surface are quite interesting. After dipping, the shape and position of the Co2p,,, peak indicate the presence of either cobalt hydroxide or cobalt carbonate, both oxygen containing species. Therefore, an 0 Is peak from this compound, in addition to that due to CaCO,, is expected in the spectrum. However, only one oxygen peak is apparent, so the calcium carbonate oxygen and the cobalt compound 0 1s peaks must appear very close together at N 531.2 eV. In fact, the absence of appreciable broadening of the dipped 0 Is peak relative to the bare calcite peak indicates that the peak shapes and positions for the two must be almost identical. The binding energy expected for the 0 1s peak of Co(OH), [29] is very close to the 0 1s binding energy for CaCO, reported by others [32] and observed in this lab previously [19] and in the present study (table 1). Therefore, the disappearance of the shoulder on the oxygen peak is consistent with the formation of cobalt hydroxide. As noted above, the shape and position of the Co2p,,, peak after dipping are closest to those expected for Co(OH1, as observed previously (table 1, and references therein) and in this study (table 1 and fig. 3). Consideration then of both the changes in the 0 1s peak and the Co 2p 3,2 peak upon immersion suggest the formation of Co(OH),. The very close agreement between the peak positions for bulk Co(OH), and the observed Co and 0 peaks after dipping may reveal the extent of interaction between the cobalt complex and the calcite surface. Although the match could be somewhat fortuitous, it may instead point to a lack of interaction between cobalt hydroxide overlayer “molecules” and the calcite substrate. Since both are insulators with fairly localized charge densities, they would be expected to perturb one another less than, for instance, a metal on a semiconductor. This is especially true if the hydroxide remains at the surface without mixing appreciably with the calcium carbonate. In contrast, the Mn complex seen after dipping does not display peak maxima in the exact positions expected for manganese carbonate or other reason-

able manganese compounds (ref. [19], and references therein). This suggests that in this case there is a greater interaction, one strong enough to shift the core levels. This shift must be on the order of 0.4-0.5 eV, and possibly reflects the incorporation of the Mn into the calcium carbonate lattice. This difference in strength of interaction of the Co(OH), and MnCO, is of course consistent with the solution studies previously discussed. We have also noted several other changes with dipping. When a large amount of metal (- 2.5 ML) was deposited directly on a sputtered surface (i.e., no dip to remove sputter damage) a subsequent 30 s dip was sufficient to convert all the metallic Co to the hydroxide. Most of the sputter damage is simultaneously removed as well. This is comparable to the rate of removal of sputter damage from bare calcite. The implication is that in spite of the metal overlayers, the water has reasonably good access to calcite at the interface. This supports the case for island formation made above to account for the appearance of both an oxidized and metallic Co as deposited from the vapor. The high surface-to-volume ratio of islands would presumably allow rapid ionization and solvation of the metal atoms upon dipping. This seems to be the case, as both oxide and metal were completely converted to hydroxide for relatively short dips (30 s> with large coverages of metal, and before the complete removal of the sputter damage. Alternatively, the metal may quickly hydrate and move into the double layer, giving a similar effect. 4.3. Comparison to solution studies The interaction of hydrated metal ions with the calcium carbonate surface is of particular interest for reasons which were outlined in the introduction to this paper. As previously discussed, measurements of both the sorption and desorption of a series of metal ions onto and from calcite in solution suggest two different mechanisms of aqueous deposition of metal ions at the surface [16]. Mn and Cd ions appear to dehydrate rapidly after being adsorbed, forming a surface CaCO,(s)-MeCO,(s) solid solution (a

D. L. Hanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

surface precipitate) at aqueous concentrations significantly below that required for homogeneous MnCO, or CdCO, precipitation. Co, Zn and Ni ions, in contrast, seem to remain hydrated in surface complexes until r~c~stallization of the CaCO, incorporates them into the crystal structure. These metals have higher free energies of hydration (by about 50 kJ [331) than do Mn and Cd, and so are held more strongly to the surrounding cage of water molecules. This undoubtedly contributes to the difference in deposition mechanism. The X-ray photoemission spectra of a Mncoated calcite surface before and after dipping have previously been reported [19]. The manganese appeared to be more highly complexed after dipping, and the new photoline was consistent with the formation of MnCO,. The current study presents persuasive evidence for the formation of Co hydroxide when a Co-on-calcite surface is dipped. These two very different species, MnCO, and Co(OH),, seen after return of the surface from solution to vacuum, presumably reflect the nature of the metal ion complexes at the solid/ liquid interface. Therefore these results support the contention that Mn forms a surface CaCO~ts)-MnCO~(s) solid solution while Co remains in a hydrated complex. The rates of dissolution of the two metals (determined from fig. 4) also support this contention. The Mn at the solid/liquid interface is removed at a much slower rate than the Co. This is to be expected if Mn is incorporated in the crystal Iattice to form a solid solution at the surface while Co forms a less tightly bound hydrated species. The consistency of the results of the two experiments is somewhat surprising. The dissolution of a thin film into pure water and the desorption of metallic cations into a saturated calcite solution are not necessarily expected to demonstrate similar behavior. This suggests that the meta surface complexes formed are not particularly sensitive to the concentration of calcite in solution, at least for the low coverages we used. Perhaps the electrical double layer is providing a saturated solution in the vicinity of the surface. This possibility is supported by the observation that the rates of removal of metal ions from the calcite surface

37

measured in this study are three orders of magnitude lower than the reported rate of dissolution of calcite into pure water (2.55 ML/s) [34]. This would suggest (1) a very strong attraction between the surface and the metallic ions, in whatever form they assume at the solid/liquid interface; or (2) the presence of the metal greatly slows dissolution. The observation that a large metal deposit does not greatly affect the rate of removal of sputter induced damage (noted in the previous section) makes the first possibility most likely. A more accurate determination of the rate of calcite dissolution with and without a metal deposit at the surface would make obvious which of these possibilities is correct. Accordingly, experiments are being planned to measure these rates with our instrumental setup.

5. Conclusion When vapor deposited on a bare calcite surface, Co appears in both metallic and oxidized states. This occurs for coverages as low as 0.16 ML. When immersed in DI water, both states are completely converted to one with X-ray photoelectron spectral features most consistent with cobalt hydroxide. The rates of removal of Co and Mn with immersion have been calculated - Co moves into solution two times faster than Mn (0.0030 _t 0.0006 ML/s and 0.0015 k 0.0003 ML/s, respectively). Further, these rates are about three orders of magnitude slower than that reported for the dissolution of bare calcite (2.55 ML/s) [34]. The identification of the hydroxide and the difference in the rates of removal of the two metals are consistent with the interpretation of the results of recent solution experiments. The authors suggest that Mn is incorporated in the calcite lattice to form a solid solution, while Co forms a less tightly bound hydrated compIex [16]. The difference in dissolution rates of the metals and calcite qualitatively indicates the strength of attraction between the cation complexes and the calcite surface. The results presented here show a significant consistency between UHV and geochemical solution studies. However, even more important than

D.L. Hanchard, Jr., D.R. Baer / The interactions of Co, Mn and water with calcite surfaces

38

the corroboration of solution study results is the different information and new questions that UHV techniques bring to light. The current results point to specific follow up studies that will provide more information about the molecular scale interaction of metals with calcite in aqueous solution, and thus provide a basic understanding of processes important in sub-surface transport. Herein lies the strength of UHV techniques in probing the complexes formed at insulating oxide/ solution interfaces. Combining UHV and solution studies is fairly commonplace in electrochemical and corrosion studies [35], but has only recently been applied to environmentally related materials and problems [19,20,36,37], as in the study presented here. It is hoped that this paper will stimulate further research, and development and application of new techniques, in the area of fundamental processes and materials of significance to waste remediation and environmental restoration.

Acknowledgements

We are very grateful for help in data acquisition by Mark Engelhard and Matthew Porter, data analysis by Rick Williford, and for useful and stimulating discussions with John Zachara and Cal Ainsworth. We would also like to thank Park Scientific for providing SFM scans. This work was supported by the US Department of Energy under contract DE-AC06-76RLO 1830 with Battelle Memorial Institute. References [I] J.A. Davis and K.F. Hays, Eds., Geochemical Processes

[2] [3]

[4] [5]

at Mineral Surfaces, ACS Symposium Series (ACS, Washington, DC, 1986). W. Stumm, Ed., Aquatic Surface Chemistry (Wiley, New York, 1987). M.F. Hochella and A.F. White, Eds., Mineral-Water Interface Geochemistry, Reviews in Mineralogy, Vol. 23 (American Mineralogical Society, 1990). K.H. Wedepohl, in: Handbook of Geochemistry, Vol. 1, Ed. K.H. Wedepohl (Springer, Berlin, 1969) p. 250. R.M. Garrels and F.T. Mackenzie, Evolution of Sedimentary Rocks (Norton, New York, 1971) p. 397.

161 See, e.g., R.J. Reeder, Ed., Carbonates: Mineralogy and Chemistry, Reviews in Mineralogy, Vol. 11 (American Mineralogical Society, 1983). f7] M.L. Franklin and J.W. Morse, Mar. Chem. 12 (1983) 241. lsl W.A. Kornicker, J.W. Morse and R.N. Damascenos, Chem. Geol. 53 (1985) 229. [91 N.E. Pingitore, Jr. and M.P. Eastman, Geochim. Cosmochim. Acta 50 (1986) 2195. l101J.A. Davis, C.C. Fuller and A.D. Cook, Geochim. Cosmochim. Acta 51 (1987) 1477. llll CC. Fuller and J.A. Davis, Geochim. Cosmochim. Acta 51 (1987) 1491. 1121R.N.J. Comans and J.J. Middelburg, Geochim. Cosmochim. Acta 51 (1987) 2587. 1131 J.M. Zachara, J.A. Kittrick and J.B. Harsh, Geochim. Cosmochim. Acta 52 (1988) 2281. 1141 J.M. Zachara, J.A. Kittrick, L.S. Dake and J.B. Harsh. Geochim. Cosmochim. Acta 53 (1989) 9. l151P. Wersin, L. Charlet, R. Karthein and W. Stumm, Geochim. Cosmochim. Acta 53 (1989) 2787. lt61 J.M. Zachara, C.E. Cowan and C.T. Resch, Geochim. Cosmochim. Acta 55 (1991) 1549. 1171 S.L. Stipp, M.F. Hochella, G.A. Parks and J.O. Leckie, Geochim. Cosmochim. Acta, 56 (19921, in press. 1181 J.J. Zullig and J.W. Morse, Geochim. Cosmochim. Acta 52 (1989) 1667. 1191 D.R. Baer, D.L. Blanchard, M.H. Engelhard and J.M. Zachara, Surf. Interface Anal. 17 (1991) 25. 1201S.L. Stipp and M.F. Hochella, Geochim. Cosmochim. Acta 55 (1991) 1723. l211For a complete account of the experimental geometry and a set of XPS survey and multiplex spectra for vacuum cleaved calcite, see: D.R. Baer, A.M. Marmorstein, R.E. Williford and D.L. Blanchard, Jr., Surf. Sci. Spectra 1 (1992) 80. [22] C.D. Wagner, LE. Davis, M.V. Zeller, J.A. Taylor, R.H. Raymond and L.H. Gale, Surf. Interface Anal. 3 (1981) 211. [23] D. Briggs and M.P. Seah, Eds. Practical Surface Analysis by Auger and X-ray Photoelectron Spectroscopy, app. 5 (Wiley, Chichester, 1983). 1241C.D. Wagner, W.M. Riggs, L.E. Davis, J.F. Moulder and G.E. Muilenberg, Handbook of X-ray Photoelectron Spectroscopy (Perkin Elmer. Eden Prairie, FL, 1979). [25] D.A. Shirley, Phys. Rev. B 5 (1972) 4709. 1261 M.P. Seah in ref. [23], chap. 5, and references therein. [27] C.V. Schenck, J.G. Dillard and J.W. Murray, J. Colloid Interface Sci. 95 (1983) 398. [28] T.J. Chuang, CR. Brundle and D.W. Rice, Surf. Sci. 59 (1976) 413. [29] N.S. McIntyre and M.G. Cook, Anal. Chem. 47 (1975) 2208. 1301 P.E. Hillner, S. Manne, A.J. Gratz and P.K. Hansma, Ultramicroscopy 42-44 (1992) 1387. 1311 S. Tougaard and H.S. Hansen, Surf. Interface Anal. 14 11989) 730.

D. L. Blanchard, Jr., D. R. Baer / The interactions of Co, Mn and water with calcite surfaces [32] NIST XPS Data Base, assembled

by C.D. Wagner and available from NIST, Gaithersburgh, MD 20899, USA. [33] J. Burgess, Metal Ions in Solution (Ellis Horwood, Chichester, UK, 1978) p. 481. [34] R.G. Compton and P.J. Daly, J. Colloid Interface Sci. 115 (1987) 493. [35] See, e.g., P.M.A. Sherwood, Chem. Sot. Rev. 14 (1985) 1; A.T. Hubbard, Chem. Rev. 88 (1988) 633; D.R. Baer, C.R. Clayton and G.D. Davis, Eds., The

39

Application of Surface Analysis Methods to Environmental/Material Interactions (The Electrochemical Society, Pennington, NJ, 1991). All have many further references on the subject as well. [36] M.F. Hochella in ref. [3]. [37] Environmental Materials and Interfaces, October 1991, PNL-SA-19675 (CONF-9009424) (Pacific Northwest Laboratory, Richland WA, USA, 1991).