Atomic‐Scale Studies of Fe3O4(001) and TiO2(110) Surfaces Following Immersion in CO2‐Acidified Water

Abstract Difficulties associated with the integration of liquids into a UHV environment make surface‐science style studies of mineral dissolution particularly challenging. Recently, we developed a novel experimental setup for the UHV‐compatible dosing of ultrapure liquid water and studied its interaction with TiO2 and Fe3O4 surfaces. Herein, we describe a simple approach to vary the pH through the partial pressure of CO2 (pCO2 ) in the surrounding vacuum chamber and use this to study how these surfaces react to an acidic solution. The TiO2(110) surface is unaffected by the acidic solution, except for a small amount of carbonaceous contamination. The Fe3O4(001)‐(2 ×2 )R45° surface begins to dissolve at a pH 4.0–3.9 (pCO2 =0.8–1 bar) and, although it is significantly roughened, the atomic‐scale structure of the Fe3O4(001) surface layer remains visible in scanning tunneling microscopy (STM) images. X‐ray photoelectron spectroscopy (XPS) reveals that the surface is chemically reduced and contains a significant accumulation of bicarbonate (HCO3 −) species. These observations are consistent with Fe(II) being extracted by bicarbonate ions, leading to dissolved iron bicarbonate complexes (Fe(HCO3)2), which precipitate onto the surface when the water evaporates.


Supplemental Information
Equilibria involved in the CO2 dissolution in water, acidic constants and pH values.
CO2 dissolves in liquid water and then forms H2CO3 according to the equilibria (1) and (2) respectively: The reaction constant Kr [1] of equation (2) indicates that the forward reaction is slow with respect to carbonic acid decomposition [2]. At equilibrium, only a small fraction of the dissolved CO2 is converted into carbonic acid, while most of the CO2 remains as molecular CO2.
Carbonic acid is a weak acid, which dissociates in two steps to bicarbonate (HCO3 -) and carbonate (CO3 2-), respectively, as described by the equilibria (3) and (4): 3 − + 2 ↔ 3 2− + 3 + 2 = 4.7 · 10 −11 A weak acid only partially dissociates into its ions in aqueous solutions, and the corresponding acidic constant can be considered as a quantitative measure of the strength of an acidic solution. In the case of carbonic acid, the two acidic constants [3], related to the two dissociation steps, are indicated in (3) and (4) as Ka1 and Ka2. Given that Ka2 << 1, we can assume that the concentration of CO3 2is always negligible compared to HCO3 -. Consequently, the equilibrium (4) plays no quantitative role in the pH of the solution.
Considering, then, only the sum of the equilibria (2) and (3), we will have a new equilibrium (5) with the corresponding acidic constant necessary for the pH calculation of our solution (6), as described below: By increasing the partial pressure of CO2 surrounding the water drop, we are able to shift the described equilibria more to the right, resulting in the acidification of the water. We can use Henry´s law [4] (equation (7)), which describes the equilibrium between the partial pressure of a gas above the liquid and the concentration of this gas dissolved in the liquid, to tune the pH of the water drop to the desired value: where is Henry´s constant (29.4 · 10 3 · L ⁄ at 25 °), 2 is the partial pressure of the CO2, and [ 2( ) ] is the concentration of the species in the aqueous phase.
The pH can be written as a function of the CO2 concentration as shown in equation (8): Table 1 shows the concentration values calculated using the equations described above, for the dissolved species in water as a function of the 2 surrounding the water drop. While the bicarbonate concentration increases as the pH decreases, the carbonate concentration can be considered constant around the value of 5.6 × 10 -11 mol/L.   At pH 4.8, the surface OH corresponds to 19.9% of the overall O 1s area, and decreases to 15.9% and to 11.1%, at pH 4 and 3.9, respectively. The HCO3contribution corresponds to the 10.2% of the overall O 1s area when the pH of the acidic solution is 4.8, and increases to 36.6% and to 61.1%, as the pH decreases to 4 and to 3.9, respectively. These results follow the same trend of the surface bicarbonate concentrations measured based on the C 1s integral area reported in Fig. 3.