Vanadium As a Potential Membrane Material for Carbon Capture

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Vanadium as a potential membrane material for carbon capture: Effects of minor flue gas species Mengyao Yuan, Simona Liguori, Kyoungjin Lee, Douglas G. Van Campen, Michael F. Toney, and Jennifer Wilcox Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b02974 • Publication Date (Web): 14 Sep 2017 Downloaded from http://pubs.acs.org on September 18, 2017

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Environmental Science & Technology

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Vanadium as a potential membrane material for

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carbon capture: Effects of minor flue gas species

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Mengyao Yuan,† Simona Liguori,†,‡ Kyoungjin Lee,†,§ Douglas G. Van Campen,ǁ Michael F.

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Toney,ǁ Jennifer Wilcox*,†,‡

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†Department of Energy Resources Engineering, Stanford University, 367 Panama Street,

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Stanford, California 94305, United States

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‡Department of Chemical and Biological Engineering, Colorado School of Mines, 1613 Illinois

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Street, Golden, Colorado 80401, United States

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§Applied Materials, 974 E. Arques Avenue, Sunnyvale, California 94085, United States

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ǁStanford Synchrotron Radiation Lightsource, SLAC National Accelerator Laboratory, Menlo

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Park, California 94025, United States

ACS Paragon Plus Environment

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ABSTRACT

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Vanadium and its surface oxides were studied as a potential nitrogen-selective membrane

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material for indirect carbon capture from coal or natural gas power plants. The effects of minor

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flue gas components (SO2, NO, NO2, H2O, and O2) on vanadium at 500–600 °C were

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investigated by thermochemical exposure in combination with x-ray photoelectron spectroscopy

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(XPS), scanning electron microscopy (SEM), and in situ x-ray diffraction (XRD). The results

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showed that SO2, NO, and NO2 are unlikely to have adsorbed on the surface vanadium oxides at

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600 °C after exposure for up to 10 hours, although NO and NO2 may have exhibited oxidizing

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effects (e.g., exposure to 250 ppmv NO/N2 resulted in an 2.4 times increase in surface V2O5

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compared to exposure to just N2). We hypothesize that decomposition of surface vanadium

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oxides and diffusion of surface oxygen into the metal bulk are both important mechanisms

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affecting the composition and morphology of the vanadium membrane. The results and

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hypothesis suggest that the carbon capture performance of the vanadium membrane can

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potentially be strengthened by material and process improvements such as alloying, operating

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temperature reduction, and flue gas treatment.

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1. Introduction


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Vanadium and vanadium oxides are widely studied for a suite of environmental and energy

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applications, e.g., selective catalytic reaction (SCR) catalysts, electrodes for lithium-ion batteries,

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hydrogen storage, and fusion systems.1-3 In recent years, vanadium has been investigated as a

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potential membrane material for carbon capture,4-8 the process of separating carbon dioxide

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(CO2) from the flue gas of fossil fuel power generation as a climate change mitigation method. It

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is therefore important to understand the effects of flue gas pollutants and minor components on

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vanadium at realistic power plant operating conditions. In this paper, we aim to investigate the

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effects of selected flue gas species, including sulfur dioxide (SO2), nitric oxide (NO), nitrogen

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dioxide (NO2), water vapor (H2O), and oxygen (O2) on vanadium and its surface oxides using x-

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ray photoelectron spectroscopy (XPS), scanning electron microscopy (SEM), and in situ x-ray

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diffraction (XRD).

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Membrane separation is considered a competitive separation technology along with solvent-

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based absorption and physical adsorption. Recent modeling studies have shown that membrane

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separation can achieve comparable or lower energy consumption than amine absorption,9 one of

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the more mature technologies proposed for carbon capture. In addition, membrane separation has

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unique advantages such as ease of scale-up, smaller footprints, and requiring little chemicals in

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the separation process.10-11 Currently, the most-studied carbon capture membranes are polymeric

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membranes that selectively separate CO2 from flue gas (known as “CO2-selective” membranes).

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Recently, there has been an emerging research interest in the development of nitrogen-selective

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membranes, which is primarily motivated by their potential for reducing the energy use of CO2

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separation.7 In a typical flue gas from coal or natural gas combustion, nitrogen (N2) is the

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dominant species, with concentrations above 70 vol%. The high N2 concentration in flue gas is

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able to provide a higher driving force for separation (i.e., higher partial pressure difference


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across the membrane), whereas the low CO2 concentration (as low as 3–5 vol% from a natural

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gas power plant12) can pose design limitations to CO2-selective membranes. For example, it has

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been shown a CO2-selective membrane with experimentally achievable material properties

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cannot simultaneously achieve 90% capture and 95% CO2 purity, a common separation target for

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carbon capture, within single-stage separation for an inlet CO2 concentration typical of a coal-

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fired flue gas (e.g., 10–12 vol% CO2).13-14 Unlike in CO2-selective membranes, concentrated

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CO2 product stream exits on the high-pressure side of a nitrogen-selective membrane, which

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could reduce the energy use for compression to high pressures (e.g., 110 bar12) for long-term

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geologic storage.

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A few studies on polymer-based nitrogen-selective membranes can be found in the

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literature.15-19 While these membranes have successfully demonstrated N2 selectivity over CO2,

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the reported N2/CO2 selectivities are often fairly low, e.g., below 20, despite a few outliners. In

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contrast, the focus of this study is a vanadium membrane that falls under the category of metallic

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nitrogen-selective membranes. Unlike its polymeric counterparts, the vanadium membrane

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provides nitrogen selectivity via a solution-diffusion mechanism similar to that found in

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palladium-based membranes for hydrogen separation. The key steps in this mechanism include

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catalytic dissociation of N2 into nitrogen atoms on the membrane surface and diffusion of

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nitrogen atoms across the membrane bulk.5 The solution-diffusion mechanism has been validated

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for vanadium by experiments as well as density functional theory (DFT) calculations4-5 and

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enables the membrane to have N2/CO2 selectivities orders-of-magnitude higher than those

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typically found in polymer-based membranes, and ongoing effort is dedicated to improving the

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N2 permeability of the membrane (e.g., currently, at 400 °C and a pressure difference of 2–5 bar,

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N2 fluxes on the order of 0.03–0.05 mol/m2-h have been measured4, 8) for large-scale


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applications. A recent process modeling and optimization study has shown that the metallic

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nitrogen-selective membrane concept based on vanadium has the potential of achieving energy

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consumption comparable to CO2-selective membranes while reducing process design

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complexities.7

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One common uncertainty in the application of membrane technologies for carbon capture is

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the effects of minor flue gas species on the membrane materials, yet this topic is little explored

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even for polymeric membranes.20 This paper represents the first attempt to the authors’

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knowledge that investigates the effects of SO2, NO, NO2, H2O, and O2 on vanadium as a

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candidate membrane material for carbon capture. Specifically, we focused on typical conditions

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found at the exit of the economizer and prior to the emissions controls in a coal-fired power

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plant, where the flue gas temperature is 370–650 °C,21 and the pollutant concentrations are

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higher than those found at the stack (see SI Section S1). Previous experimental studies have

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shown that a suitable operating temperature range for the vanadium nitrogen-selective membrane

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is 400–800 °C.4, 8 A potential location for installing the vanadium membrane, therefore, is at the

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exit of the economizer in a coal-fired power plant, where the high-temperature flue gas can be

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fed to the membrane with minimal heating.

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In this work, we performed thermochemical exposure experiments for vanadium foils using

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various gas mixtures and studied the chemical and morphological changes on the vanadium

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oxide surface using XPS and SEM. Surface studies were the focus of this work because they can

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reveal the first signs of potential chemical reactions or gas permeation and are indicative of bulk

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changes. In addition, we will briefly discuss a set of in situ transmission XRD experiments aimed

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at investigating potential phase changes in the bulk of vanadium during thermochemical

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exposure. While the purpose of this work is to facilitate the understanding of the chemical


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resistance and potential failure mechanism of vanadium in an actual flue gas environment, the

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methodology and results presented here have broader implications in other environmental and

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energy applications where vanadium and vanadium oxides play a key role.

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2. Experimental

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The experiments in this work can be divided into two parts. The first part consists of

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thermochemical exposure tests and ex situ XPS and SEM characterization. The second part is a

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series of in situ XRD measurements performed during gas exposure. The materials and

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experimental procedures are briefly described below.

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Materials. Commercial vanadium foils (Goodfellow Cambridge Limited, cold-rolled, 99.8%

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purity, 0.040 mm thickness) were used in all of the tests. These foils were cut into small pieces

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for gas exposure and characterization (approximately 25 mm2 for ex situ analyses and 85 mm2

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for in situ XRD measurements). Test gases were single gases or gas mixtures prepared from

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argon (Ar, 99.999%), N2 (99.998%), helium (He, 99.999%), SO2 (4.97 vol%, N2 balance), NO

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(1.97 vol%, N2 balance), NO2 (1018 ppmv, N2 balance), H2O (10 vol%, N2 balance), and O2 (10

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vol%, He balance). To simulate the flue gas concentrations in a coal-fired power plant, the

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following gas mixtures were tested: 2500 ppmv SO2/N2, 250 ppmv NO/N2, 50 ppmv NO2/N2, 10

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vol% H2O/N2, and 5 vol% O2/He. These gas concentrations were chosen based on the data in the

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Integrated Environmental Control Model (IECM) software program.22 The chosen gas

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concentrations are comparable to or higher than the gas concentrations exiting the economizer in

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a pulverized coal plant for typical coal types found in IECM (see SI Table S1).


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Thermochemical exposure experiments. Each vanadium sample was supported by quartz

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wool in a glass tube reactor, and the reactor was connected to continuous gas flow by Teflon

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fittings. The tube reactor was placed in a tube furnace (Carbolite) with a heating zone of

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approximately 30 cm. The total gas flow rate in all exposure tests was set to 120 ml/min by mass

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flow controllers. The test temperature was chosen to be 600 °C to represent the flue gas

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temperature exiting the economizer. Each sample was heated in Ar from room temperature

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(approximately 20 °C) to 600 °C at 3 °C/min. The temperature was then held constant at 600 °C

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for 5 or 10 hours, during which the sample was exposed to a gas mixture containing one of the

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minor flue gas species. At the end of the exposure, the reactor was quenched to room

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temperature at 2 °C/min while Ar was used to purge the system. In both the 5-hour and 10-hour

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exposure experiments, SO2/N2, NO/N2, and NO2/N2 were tested. In addition, N2 was tested in

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both 5-hour and 10-hour experiments to rule out possible effects of the balance gas, and H2O/N2

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was tested in 5-hour experiments. The samples are referred to in the following sections by the

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gas used in the test and exposure duration, e.g., “N2-5h” refers to the sample exposed to N2 for 5

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hours, and “NO-10h” refers to the sample exposed to NO/N2 for 10 hours.

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Ex situ XPS and SEM characterization. All XPS analyses were performed on a PHI 5000

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Versaprobe using a monochromatic Al Kα radiation source (1486.6 eV). The x-ray beam was

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200 µm in diameter with a beam power of 49.9 W and an operating potential of 16 kV. A

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vacuum of ~6 × 10−8 Pa was maintained in the main chamber. The angle between the analyzer

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and sample surface was 45°. Survey scans were recorded over a binding energy (BE) range of 0–

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600 eV using a pass energy of 117.4 eV. Region scans were recorded for V 2p, C 1s, O 1s, N 1s,

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and S 2p with a pass energy of 23.5 eV. All photoemission spectra were referenced to C 1s =

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284.8 eV. The spectra were fitted using CasaXPS 2.3.17 assuming Gaussian-Lorentzian


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lineshapes and a Shirley background. High-resolution SEM images were obtained using an FEI

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XL30 Sirion SEM with a beam intensity of 5 keV and a spot size of 3.

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Exposure and in situ XRD experiments. Thermochemical exposure experiments and in situ

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transmission XRD measurements were performed on Beamline 10-2 (BL10-2) at the Stanford

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Synchrotron Radiation Lightsource (SSRL). Gas exposure was conducted using a hermetic

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annealing chamber (see SI Section S2). The energy of the incident x-ray beam was 20 keV. The

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scattering intensity was detected by a two-dimensional image detector (Dectris Pilatus 300W)

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with a pixel size of 172 µm (1475 × 195 pixels). This combination of beam energy and detector

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gave a detectable q-space of up to 10 Å−1, or 10 lines for vanadium. SO2/N2, NO/N2, NO2/N2,

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H2O/N2, O2/He as well as N2 and He were tested. The total gas flow rate was 120 ml/min in all of

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the exposure tests. Several experimental conditions were different from those in the exposure

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experiments with ex situ characterization due to the setup at the beamline. A gas exposure

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temperature of 500 °C was used, which was set by the annealing chamber material but is still

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within the temperature range at the economizer exit. Each sample was heated from room

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temperature to 500 °C at 5 °C/min, during which the respective gas or gas mixture for testing

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was supplied to the chamber as purge gas. The temperature was held at 500 °C for 4 hours for

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each gas or gas mixture with the exception of O2/He, which was tested for 3.5 hours due to the

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availability of the beamline equipment. Cooling was not programmed. Transmission XRD data

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were collected with an exposure time (time during which the sample was exposed to x-ray) of 2

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seconds every 2 minutes during temperature ramp-up and every 5 minutes while the temperature

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was stabilized at 500 °C.

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3. Results and Discussion

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Since the results from the 5-hour and 10-hour exposure tests share many similarities, the

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discussion in this section will be focused on the latter, and the differences between the two sets

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of data will be highlighted. The results from the 5-hour exposure tests are shown in the

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Supporting Information (SI Section S4).

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Ex situ XPS characterization: Survey scans. Surface elemental compositions of the

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vanadium samples before and after gas exposure were determined by XPS survey scans. It was

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found that all of the as-received vanadium foils analyzed (see SI Table S2) had consistent surface

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compositions with the following average (± standard deviation) atomic percentages: 10.0 ± 1.7%

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vanadium, 35.5 ± 3.4% oxygen, 49.6 ± 5.2% carbon, 2.7 ± 0.9% nitrogen, and 2.1 ± 0.7%

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calcium. In addition to the expected V and O, the as-received foil surfaces all contained about

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50% carbon and small amounts of nitrogen and calcium. As will be shown by XPS region scans,

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the surface vanadium on the as-received foils mostly existed as V2O5. Since direct studies on the

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interactions of vanadium foils with similar gases are rare in the literature, studies of V2O5

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catalysts will be used to compare with the results obtained from our experiments.

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Surface elemental compositions of the samples after 10-hour gas exposure are shown in

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Table 1. The surface composition of the as-received foil used for this set of experiments is also

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shown. Atomic concentrations below 0.5% are shown in brackets and are considered negligible.

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Table 1. Surface elemental compositions of the vanadium samples from the 10-hour exposure

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experiments.


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O 1s

V 2p3/2

C 1s

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N 1s

Ca 2p

Si 2p

9.0%

34.5%

52.0%

1.5%

3.2%

-

N2

23.5%

58.9%

17.6%

-

[0.4%]

-

2500 ppmv SO2

22.0%

59.3%

16.9%

-

-

1.8%

250 ppmv NO

25.2%

59.9%

14.9%

[0.4%]

-

-

50 ppmv NO2

22.1%

59.0%

18.9%

[0.2%]

-

-

as-received (Foil #4)

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10-hour gas exposure removed about 60–70% carbon and almost all nitrogen originally

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present on the vanadium foil. C 1s region scans (see SI Figure S6) show that the carbon

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contamination is from both inorganic (e.g., deposition from air) and organic (e.g., solvent used in

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foil manufacturing) sources, and that most of the organic carbon species has been removed by

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gas exposure. The small amount of silicon content found on the SO2-10h sample is most likely

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from the quartz wool support used in the exposure test.

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The nitrogen content on the NO-10h and NO2-10h samples is negligible, while small

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amounts of nitrogen were found on the NO-5h and NO2-5h samples (see SI Table S3). Based on

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literature evidence, the most probable explanation is that native nitrogen species on the as-

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received foils were not fully desorbed or decomposed during the 5-hour tests, and NO or NO2

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did not adsorb on the vanadium surface. Barbaray et al. studied NO2 adsorption on V2O5 at 10−2


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Torr (1.3 Pa) and 25–450 °C and found that the adsorbed NO2 had an XPS N 1s BE close to

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406.5 eV (reference: C 1s = 285 eV), and that the N 1s signal decreased with temperature and

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disappeared above 250 °C.23 In contrast, the N 1s peaks of the NO-5h and NO2-5h samples from

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our experiments are in the range of 400–402 eV (shown in SI Figure S4), and the exposure

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temperature was much higher than 250 °C, so the detected nitrogen is unlikely to be associated

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with NO or NO2. Several groups studied V2O5 or V2O5/TiO2 for SCR reactions and concluded

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that NO did not strongly adsorb on the catalyst surface at the temperature investigated, i.e., 150–

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400 °C.24-27 Yu et al. exposed V2O5-WO3/TiO2 to 2% NO/N2 at 350 °C and concluded that the

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adsorption of NO on the catalyst was unstable at high temperatures.28 Inomata et al. found that

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instead of adsorbing onto the catalyst, gaseous NO reacted with strongly adsorbed NH3 during

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NO reduction, and this has been confirmed by Gruber and Hermann using cluster models and

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DFT calculations.25, 29 Guo et al. performed NO adsorption on V2O5/TiO2 at room temperature

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and concluded that nitrates only formed on titania.30

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No sulfur species was detected by XPS on the samples exposed to SO2/N2, suggesting that

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SO2 did not physically or chemically adsorb onto the surface vanadium oxides. This observation

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can be corroborated by previous studies. Barbaray et al. studied SO2 adsorption on V2O5 powder

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at 25–400 °C and 0.01 Torr (1.3 Pa) using XPS and found that chemisorption of SO2 onto V2O5

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did not occur below 150 °C, and that the intensity of the S 2p line for the adsorbed SO2 peaked at

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300 °C and decreased afterwards.31 Le Bars and Auroux studied SO2 adsorption on V2O5-based

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catalysts at 353 K (80 °C) and found that SO2 did not adsorb on bulk V2O5 or V2O5/SiO2 and was

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instead selective for uncovered alumina on V2O5/γ-Al2O3.32 They also concluded that the surface

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of V2O5 shows a weakly acidic character.32 Guo et al. exposed V2O5/TiO2 catalysts to a mixture

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of 2700 ppm SO2, 5% O2, 0 or 4% water vapor in He at 380 °C for 24 hours and concluded that


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stable sulfate species formed on titania instead of vanadia.30 Yu et al. exposed V2O5-WO3/TiO2

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to 2% SO2 and 5% O2 in N2 at 350 °C and concluded that sulfate was formed with titanium in the

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catalyst.28 Similar observations can also be found in sulfate impregnation studies.24, 33 In addition

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to SCR reactions, V2O5 is also used as a catalyst for SO2 oxidation.3 Several studies on

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V2O5/activated carbon found that in SO2 oxidation reactions, SO2 reacts with surface V2O5 to

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form a VOSO4-like structure, and O2 is necessary for SO3 formation and catalyst regeneration

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via reactions with the VOSO4 intermediate.34-38 Since sulfur species were not detected by XPS

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survey scans, and no O2 was used with SO2 in the exposure tests, it is not very likely that SO2

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reacted with the vanadium foil surface to form SO3.

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Overall, the survey scans showed that adsorption of SO2, NO, and NO2 as well as the

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presence of thermally unstable species are not likely to be a concern for the purpose of carbon

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capture in high-temperature flue gas for extended time periods. Nevertheless, coal-derived flue

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gas is a complex mixture containing fly ash, O2, moisture, carbon monoxide (CO), hydrocarbons,

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mercury (Hg), and many acid gases, such as sulfur trioxide (SO3) and hydrogen chloride

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(HCl).39-40 Deposition of solid carbon species such as fly ash particles would be a problem for

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membranes in general, as observed by Eiberger et al. in their study on ceramic membranes.41 The

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testing described in this work also does not include species such as CO, SO3, or HCl. There can

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be potential impacts of both CO and acid gases on the membrane, and this could be a fruitful

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area for future research on vanadium and other metallic membranes for carbon capture.

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Ex situ XPS characterization: Region scans. This section focuses on the changes in

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vanadium oxidation state, which can be indicative of possible reactions on the surface as well as

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potential permeation of these gases into the bulk of the vanadium foil. The V 2p3/2 photoemission

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spectra of the samples from the 10-hour exposure tests are shown in Figure 1. The region scan


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for the corresponding as-received foil and the reference BEs are also shown. The V 2p3/2

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reference BEs are taken from the Handbook of X-ray Photoelectron Spectroscopy:42 metallic V

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(“V met.”) at 512.4 eV, V2O3 at 515.7 eV, VO2 at 516.3 eV, and V2O5 at 517.6 eV.

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The spectrum for the as-received foil in Figure 1 is representative of the spectra for all the as-

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received foils analyzed (see SI Figure S3). These spectra are asymmetric towards the high BE

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side, indicating that V2O5 is likely to be the dominant surface oxide on these foils. It has been

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shown that V2O5 is the most stable phase in the V–O system,43-46 especially at room temperature

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in dry air.47-48 This explains the consistency of surface oxide compositions across the as-received

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vanadium foils, since the foils had been exposed to dry air at room temperature, and the surface

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compositions would have reached equilibrium with air prior to XPS characterization.

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The V 2p3/2 spectra of the N2-10h and SO2-10h samples are broadened and shifted to lower

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BEs, suggesting that VO2 may be the dominant surface oxide on these samples. In contrast, the V

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2p3/2 spectra of the NO-10h and NO2-10h samples showed a more prominent V2O5 component

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with BEs at around 517.6 eV, especially the NO-10h sample. In theory, the V 2p3/2 peak at

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around 517.6 eV could also be associated with NH4VO3, which has a V 2p3/2 BE of 517.8 eV

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(using adventitious carbon as BE reference).49 However, the presence of NH4VO3 is unlikely, as

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N 1s peaks would otherwise be detected on the NO-10h and NO2-10h samples by XPS region

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scans.

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A general observation in this study was that XPS survey and region scans of different

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locations on the same sample often yielded very similar spectra. This suggests highly

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homogeneous surface features. One exception, however, was the NO2-10h sample (see SI Figure

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S7). The spectrum of the NO2-10h sample shown in Figure 1 was chosen to represent the mixed


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nature (i.e., both VO2 and V2O5 components are clearly visible) of the surface oxides on this

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sample. The mixed surface features can be confirmed by SEM, as will be shown in the following

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section.

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Peak fitting was performed for the region scan spectra to obtain a more quantitative

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comparison of the changes in chemical state before and after gas exposure. Like many other

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transition metals, vanadium has a range of chemical states with overlapping BEs. Since Figure 1

272

showed that multiple vanadium oxidation states are present on the sample surfaces, much work

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in this paper has been dedicated to establishing a consistent peak fitting procedure suitable for

274

these samples.

275

While general peak fitting procedures and reference fitting parameters are available in the

276

literature, there exist uncertainties in XPS peak fitting for transition metals. Existing studies and

277

standard databases report peak positions (i.e., BEs) but often do not adequately address issues

278

such as peak widths and other complexities found in transition metal 2p spectra.50 In addition,

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peak fitting is more than a mathematical problem. A common pitfall in XPS peak fitting is to let

280

goodness-of-fit dictate the interpretation of the physics and chemistry of the systems studied. In

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reality, however, peak fitting should be an iterative process where mathematical parameters and

282

physical interpretations inform each other, and the latter should play a more decisive role than

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the former in contributing to the understanding of the system.

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This study does not venture to address the various complexities involved in the XPS analysis

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of transition metals. Rather, we hope to better understand the samples in our work by drawing

286

insights from existing literature on similar systems. For this purpose, a careful literature review

287

was conducted on XPS studies of vanadium-based catalysts and vanadium oxide films. A


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288

summary of the samples, data processing techniques, BE references, and V2p3/2 BEs and full

289

widths at half maximum (FWHMs) are summarized in SI Table S5. Some studies have

290

concluded that for vanadium oxides, the O 1s line associated with vanadium oxides (VOx) serves

291

as a better BE reference than the C 1s line associated with adventitious carbon, and that the BE

292

difference between the O 1s and V 2p3/2 lines is a more reliable parameter for determining the

293

oxidation state of vanadium.51-54 For this reason, the BE differences between the O 1s (VOx) and

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V 2p3/2 lines (denoted as “∆”), are also included in the summary.

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The averages (avg.), standard deviations (stdev.), and ranges of the V 2p3/2 BEs, ∆, and V

296

2p3/2 peak FWHMs for selected oxides from Table S5 are summarized in Table 2. It can be seen

297

from Table 2 that there exists a fairly wide spread in the reported BEs, ∆, and FWHMs, which

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highlights the uncertainties associated with XPS peak fitting for vanadium. Nevertheless, these

299

literature data serve as a useful guide for the peak fitting in this work, with some observations

300

and caveats discussed in the Supporting Information (SI Section S6).


15


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301 302

Figure 1. V 2p3/2 XPS spectra of the as-received vanadium foil and samples from the 10-hour

303

exposure tests.

304 305

Table 2. Summary of selected XPS fitting parameters of V 2p3/2 and O 1s spectra for selected

306

oxides form the literature (full summary in SI Table S5).

V 2p3/2 BE [eV]

∆ [eV]

min.–max. avg. ± stdev.


(range)

V

512.3 ± 0.3

V 2p3/2 FWHM [eV]

511.5–512.9


(range)

-

-

(range)

1.2 ± 0.5

0.7–2.0


16

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(1.4)

(1.3)

514.4–516.8 V 2O 3

VO2

V 2O 5

515.5 ± 0.6

14.2–15.2 14.6 ± 0.4

1.1–4.8 2.7 ± 1.5

(2.4)

(1.0)

(3.7)

514.8–516.8

12.7–15.0

1.0–4.5

515.9 ± 0.5

13.9 ± 0.5

2.5 ± 1.2

(2.0)

(2.3)

(3.5)

516.3–518.3

12.6–13.2

0.9–3.0

517.2 ± 0.4

12.9 ± 0.1 (2.0)

1.5 ± 0.4 (0.6)

(2.1)

307 308

The peak fitting results for the V 2p3/2 and O 1s spectra from the 10-hour exposure tests are

309

shown in Table 3. Three main vanadium oxides—V2O3, VO2, and V2O5—were considered for

310

fitting of the V 2p3/2 lines, but it was found that using just VO2 and V2O5 was adequate to

311

provide reasonably good fits. Previous discussions have ruled out the possibility that vanadium

312

was associated with sulfur- or nitrogen-containing species, e.g., VOSO4 or NH4VO3. These

313

species were therefore not considered in peak fitting. The O 1s spectra were deconvoluted into

314

two components: one at around 530 eV corresponding to VOx, and one at around 531–532 eV

315

corresponding to C=O bonds.42, 54 The ∆ values were calculated from the BEs of the O 1s (VOx)

316

and V 2p3/2 components. The oxide composition is represented by the peak area percentages (%

317

area) of the V 2p3/2 components relative to the entire V 2p3/2 spectra. The average surface

318

oxidation state was calculated from the assigned oxidation states and oxide composition.


17


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319 320

Table 3. XPS fitting parameters for the V 2p3/2 spectra of the vanadium samples after the 10-

321

hour exposure experiments.

oxide

FWHM BE [eV]

assigned

VO2

average % area

∆ [eV]

[eV]

516.1

1.1

oxidation state

16.8%

14.2 4.8

as-received (Foil #4) V 2O 5

517.3

1.4

83.2%

13.0

VO2

516.3

1.6

75.0%

13.8 4.3

N2 V 2O 5

517.8

1.6

25.0%

12.4

VO2

516.3

1.6

69.9%

13.9 4.3

2500 ppmv SO2 V 2O 5

517.7

1.5

30.2%

12.5

VO2

516.2

1.3

14.4%

14.3 4.9

250 ppmv NO

50 ppmv NO2

V 2O 5

517.6

1.2

85.6%

12.9

VO2

516.0

1.5

60.0%

14.0

4.4


18

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V 2O 5

517.4

1.5

40.0%

12.6

322 323

It should be noted that although the results in Table 3 were generated from consistent peak

324

fitting procedures based on literature evidence, uncertainties such as satellite peaks and

325

asymmetry have not been fully accounted for in these fits. Nevertheless, these results provide a

326

quantitative comparison of the changes in surface oxidation state as well as some insights into

327

the physical and chemical processes that induced these changes during the thermochemical

328

exposure.

329

The surface oxides on the as-received foil shown in Table 3 consist of 83% V2O5 and 17%

330

VO2, which confirms the observation from the corresponding spectrum in Figure 1 that V2O5

331

was the dominant surface oxide. The average surface oxidation states of the N2-10h and SO2-10h

332

samples are both 4.3, close to those of the N2-5h and SO2-5h samples (see SI Table S4). The NO-

333

10h sample has a high average surface oxidation state of 4.9, consistent with the peak position

334

and shape of the corresponding spectrum in Figure 1. This value is also similar to the average

335

surface oxidation state of the NO-5h sample, which is 4.8. It can be concluded that V2O5 is the

336

dominant oxide on both NO samples. The NO2-10h sample has an average surface oxidation

337

state of 4.4, which is lower than that of the NO2-5h sample, 4.6. This can also be inferred by

338

comparing the corresponding spectra (in Figure S5 and Figure 1). This difference highlights the

339

mixed surface features resulting from exposure to NO2. The ∆ values in Table 3 are 14.0 ± 0.2

340

eV for the fitted VO2 peaks and 12.6 ± 0.2 eV for the fitted V2O5 peaks. The V 2p3/2 BE of the

341

fitted VO2 and V2O5 peaks are 516.2 ± 0.1 eV and 517.6 ± 0.2 eV, respectively. These values

342

agree with the literature data in Table 2.


19


Page 20 of 30

343

The lowered surface oxidation states could have resulted from thermal decomposition of

344

V2O5 to lower oxides at high temperature and low O2 partial pressure conditions. This hypothesis

345

is supported by both experimental and theoretical evidence from the literature. Su and Schlögl

346

observed thermal reduction of V2O5 crystals to V2O3 via VO2 when the crystals were heated up

347

to 600 °C in a high vacuum of 10−7 Torr (1.3 × 10−5 Pa).55 Based on thermodynamic calculations,

348

Zhang et al. predicted that V2O5 decomposes at 808 K (535 °C) in an inert atmosphere, while

349

Gulbransen and Andrew predicted that V2O5 decomposes to VO2 and O2 below an oxygen partial

350

pressure of 10−6 mbar above 500 °C.48, 56-57 The phase diagram simulated from computational

351

thermodynamics by de Castro et al. indicates that VO2 is the most stable phase at high

352

temperatures and low pressures. In particular, at 873 K (600 °C), O2 partial pressure would need

353

to be on the order of 1 Pa (10−5 bar) for V2O5 to be stable.47 From in situ XRD measurements of

354

controlled oxidation of vanadium thin films, Rampelberg et al. observed that VO2 was present

355

throughout the temperature and O2 partial pressure ranges investigated (430–615 °C, 0.2–200

356

mbar), whereas V2O5 only appeared at higher O2 partial pressures (20–200 mbar).58 Since N2

357

fluxes through vanadium foils were only detected at above a pressure difference of 2 bar in

358

previous permeation experiments, and the detected fluxes were comparatively small,4, 8 N2 can be

359

considered chemically inert to the vanadium foils in the current experimental setup, where the

360

foil was not sealed or pressurized. The purge gas Ar is also considered chemically inert,

361

independent of the pressure difference across the foils. Therefore, thermal decomposition is

362

likely to be a mechanism that resulted in the lowering of oxidation states after exposure to N2.

363

Thermal decomposition is also thought to have contributed to the lowered oxidation states on the

364

SO2 samples for two reasons. First, the lineshapes and peak fitting results for the SO2 samples

365

are similar to those for the N2 samples. Second, since sulfur-containing species were not detected


20

Page 21 of 30


366

by XPS after exposure to SO2/N2, it is less likely that SO2 has adsorbed to the vanadium surface

367

and reduced the surface oxides. This conclusion may change depending on the SO2 concentration

368

used in gas exposure as well as the sample condition prior to exposure, as shown by a recent

369

study using deposited vanadium films and 50% SO2/N2.59

370

Exposure to NO/N2 and NO2/N2, especially the former, has resulted in higher surface

371

oxidation states compared to exposure to N2 and SO2/N2. Here, oxidation by NO or NO2 may

372

have been a more competitive mechanism than thermal decomposition, resulting in overall

373

higher surface oxidation states. The oxidizing nature of NO can be indirectly verified by several

374

studies in the literature. Catalysis studies of V2O5/TiO2 by Topsøe et al. in the temperature range

375

of 375–625 K (202–352 °C) concluded that vanadium surface pre-adsorbed with NH3 was more

376

readily regenerated in NO than in O2, with the former providing a more oxidizing atmosphere.27,

377

60

378

(1.3–1333.2 Pa) and found that NO adsorbed on Fe3O4 caused oxidation of the metal oxide

379

surface via back donation of electrons.61 Dianis and Lester studied NO adsorption on NiO,

380

Co3O4, and graphite at −100 °C and 2 × 10−5 Torr (0.003 Pa) and concluded that back donation

381

occurred from the metal oxides to the adsorbed NO.62 An adsorption study of NO and N2O on

382

Al(100) single crystals at 80 K by Pashutski and Folman found that there was a negative charge

383

transfer from the metal to the adsorbed gases, and that NO was more reactive on Al(100) than O2

384

or N2O.63 A caveat is that the experimental conditions in these studies as well as the metals or

385

metal oxides used are quite different from the current study. Future work should include more

386

focused studies on NO and NO2 to further test and refine the hypothesis in this work.

Contour and Mouvier conducted adsorption experiments at −100–200 °C and 10−2–10 Torr

387

Ex situ SEM analyses. The surface SEM images of the samples from the 10-hour exposure

388

experiments and the corresponding as-received foil are shown in Figure 2 (see SI Section S7 for


21


Page 22 of 30

389

selected cross-sectional SEM images). Two surface features are common among the annealed

390

samples. The N2-10h and SO2-10h samples have dense, irregularly shaped grains with lateral

391

sizes of 1–2 µm. Similar features have also been observed for VO2 films and crystals in the

392

literature.58, 64-67 The NO-10h samples have plate-like crystals with widths below 1 µm. Similar

393

features can also be found for V2O5 films or features associated with V2O5 in the literature.44-45,

394

68-70

395

crystals, as shown in Figure 2(e). This confirms the observation from the XPS scans that the

396

surface oxide composition on this sample is rather inhomogeneous. In general, one can see that

397

the SEM images consistently confirm the XPS results discussed in the previous section.

The NO2-10h sample has mixed features of both the irregularly shaped grains and plate-like

398 399

Figure 2. SEM images of (a) an as-received vanadium foil (Foil #4) and the samples exposed to

400

(b) N2, (c) 2500 ppmv SO2/N2, (d) 250 ppmv NO/N2, and (e) 50 ppmv NO2/N2 in the 10-hour

401

exposure experiments.

402


22

Page 23 of 30


403

It can be inferred from the XPS region scans and SEM results that while vanadium and its

404

surface oxides showed certain chemical resistance to flue gas components such as SO2 and H2O,

405

an oxidizing environment could be a concern for the long-term structural integrity and

406

performance of the vanadium membrane as well as other metallic membranes. For metallic

407

membrane systems to be implemented for carbon capture in high-temperature flue gas, future

408

research should investigate potential process modifications to reduce the oxidant levels in flue

409

gas and material improvements such alloying to enhance the resistance of membranes to

410

oxidation.

411

In situ transmission XRD. Throughout the exposure experiments at 500 °C, no new phases

412

in the vanadium membranes were detected by in situ transmission XRD measurements,

413

independent of the gas or gas mixture used. The complete results of the in situ XRD experiments

414

are shown in the Supporting Information (SI Section S7). For comparison, the samples from the

415

previously described 10-hour exposure experiments at 600 °C were analyzed using ex situ XRD,

416

and preliminary results (not included here) showed evidence that vanadium oxides were present

417

in the bulk the these samples. Since the main differences between the two sets of samples are

418

exposure temperature and duration, these observations confirmed the findings that vanadium

419

oxide film growth is promoted by increasing temperature and reaction time.1,

420

hypothesize that the bulk vanadium oxides formed due to the diffusion or dissolution of oxygen

421

from the surface oxides to the metal bulk, with the diffusion driven by the difference in the

422

oxygen content between the surface and the bulk. This hypothesis can particularly be supported

423

by the fact that bulk oxidization was observed even when vanadium was only exposed to non-

424

oxidizing gases, e.g., N2. In this case, the oxygen in the surface oxides was the only possible

425

source of oxygen that could have resulted in bulk oxides. The previously discussed reduction of

56, 71-73

We


23


Page 24 of 30

426

surface V2O5 can also be attributed to oxygen diffusion or dissolution into the bulk, which has

427

been observed in several previous studies on vanadium oxides.71, 74-75 It should be emphasized

428

that the in situ XRD experiments in this study only aimed at examining the vanadium bulk. For

429

investigations of thin films, readers are referred to a recent in situ XRD study by Rampelberg et

430

al. on controlled oxidation of vanadium thin films over a range of O2 partial pressures and

431

temperatures as well as controlled reduction of V2O5 thin films in H2.58

432

If oxygen diffusion or dissolution from the surface oxides has indeed contributed to

433

vanadium oxidation in the membrane bulk, future studies on vanadium membranes for carbon

434

capture should investigate ways to reduce the operating temperature of the membrane system to

435

slow down bulk oxidation. A lower operating temperature can also potentially reduce the carbon

436

capture energy use of the vanadium membrane7 and is therefore also preferred from the process

437

design perspective. In general, mechanical stability and integrity are always an issue with any

438

membrane, as is long-term exposure to high temperatures. These will be investigated in future

439

research. In addition, the actual N2 permeation performance of the vanadium membrane will be

440

examined in simulated gas mixtures and correlated with the analytical work reported in this

441

paper. These permeation experiments can especially be benefited by coupling with in situ or

442

operando XRD characterization, which can effectively track the influences of these gases on the

443

membrane with time and potentially reveal transient structural or phase changes in the

444

membrane.


24

Page 25 of 30


445

ASSOCIATED CONTENT

446

Supporting Information

447

The Supporting Information is available free of charge on the ACS Publications website at DOI:

448

.

449

Additional descriptions of experimental design, literature summary of XPS peak fitting

450

parameters, and additional results and discussions (PDF)

451 452

AUTHOR INFORMATION

453

Corresponding Author

454

*Address: Chemical and Biological Engineering Department, Colorado School of Mines, 1613

455

Illinois Street, Golden, CO 80401. Email: [email protected].

456

Notes

457

The authors declare no competing financial interest.

458 459

ACKNOWLEDGMENTS

460

This work is supported by the National Science Foundation under Grant No. 1263991. The XPS

461

and SEM analyses were performed at the Stanford Nano Shared Facilities (SNSF), supported by

462

the National Science Foundation under award ECCS-1542152. The XRD experiments were

463

carried out at the Stanford Synchrotron Radiation Lightsource, SLAC National Accelerator

464

Laboratory, which is supported by the US Department of Energy, Office of Science, Office of

465

Basic Energy Sciences under Contract No. DE-AC02-76SF00515. The authors would like to


25


Page 26 of 30

466

thank Dr. Laura Schelhas at SLAC, Chuck Hitzman, Dr. Juliet Jamtgaard, and Richard Chin at

467

the SNSF for their helpful discussions and insights on XPS and SEM analyses, and Ron Marks

468

and Tim J. Dunn at the SSRL for providing beamline support.

469 470

REFERENCES

471 472 473 474 475 476 477 478 479 480 481 482 483 484 485 486 487 488 489 490 491 492 493 494 495 496 497 498 499 500 501 502 503 504 505

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Vanadium As a Potential Membrane Material for Carbon Capture

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