Article pubs.acs.org/JPCC
Nucleation and Growth Mechanisms of an Electrodeposited Manganese Oxide Oxygen Evolution Catalyst Michael Huynh,†,‡ D. Kwabena Bediako,† Yi Liu,† and Daniel G. Nocera*,† †
Department of Chemistry and Chemical Biology, Harvard University, 12 Oxford Street, Cambridge, Massachusetts 02138, United States ‡ Department of Chemistry, Massachusetts Institute of Technology, 77 Massachusetts Avenue, Cambridge, Massachusetts 02139, United States S Supporting Information *
ABSTRACT: We investigate the mechanisms of nucleation and steady-state growth of a manganese oxide catalyst (MnOx) electrodeposited from Mn2+ solutions in a weakly basic electrolyte. Early catalyst growth was probed through chronoamperometry transients, which were fit to reveal a progressive nucleation mechanism for initial catalyst formation. Time-dependent atomic force microscopy snapshots of the electrode surface reveal a rapid increase in nucleus size together with a sluggish rise in coverage, which is also characteristic of progressive nucleation. Electrochemical kinetic studies of the catalyst growth yield a Tafel slope of approximately 2.3 × RT/2F and a rate law consisting of a second-order and inverse fourth-order dependence on [Mn2+] and proton activity, respectively. These results are consistent with a deposition mechanism involving rate-limiting disproportionation of aqueous Mn3+, resolving a longstanding ambiguity surrounding the deposition of manganese oxides under nonacidic conditions.
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INTRODUCTION First-row transition-metal oxides are materials of choice for energy storage and conversion applications.1−6 Manganese oxides have found particular use as the functional materials of capacitors7,8 and batteries9−11 owing to the redox versatility of manganese ion (from Mn2+ to Mn7+).12,13 Inspired by nature’s use of a Mn4Ca-oxo cubane in Photosystem II, there has been a resurgence of interest in using manganese oxides as catalysts to drive solar-to-fuels conversions derived from the oxygen evolving reaction (OER). The activity of manganese oxides toward OER has been observed to depend on the method of preparation of the oxide, whether that be by electrodeposition,14−18 thermal decomposition,19−21 or chemical precipitation.22,23 An emerging trend from these studies is that the presence of Mn3+ in the oxides confers higher catalytic activity.14−23 However, a description of the manganese oxidation state is incomplete without consideration of the dynamical processes attendant to manganese oxides. These dynamical processes find their origins in the redox versatility of manganese ion. For instance, synthetic tetranuclear-manganese clusters liberate Mn2+ ions during operation, which are then reoxidized to furnish a Mn3+/4+ oxide.16 Similarly, manganeseoxo complexes embedded in a Nafion membrane transform into a birnessite-like manganese oxide upon application of anodic potentials.23 Results such as these provide an imperative to understand the role of the manganese ion oxidation state during the formation of manganese oxide OER catalysts. However, little attention has been devoted to understanding © XXXX American Chemical Society
the deposition of manganese oxides in the context of OER catalysis. Previous deposition studies have largely been devoted to the production of electrolytic manganese dioxide (EMD) from highly acidic electrolytes (0.1 to 5.0 M H2SO4) owing to interest in this material as a battery electrode.24−30 These conditions do not reflect those for OER, where catalyst films are deposited or evaluated at neutral to alkaline pHs.14−23 Moreover, even for EMD formation, the deposition mechanism is unresolved. Cyclic voltammetry (CV) was unable to distinguish between deposition by an ECE (i.e., two oneelectron transfer steps separated by a chemical step) and a disproportionation (DISP) process because ECE and DISP CV traces are almost identical and can only be distinguished by a subtle trace-crossing effect upon variation of scan rate.31 To discern the mechanism of manganese oxide deposition in neutral to alkaline electrolytes of relevance to OER, we now provide a detailed study of both nucleation and growth pathways of an electrodeposited manganese oxide (denoted MnOx) by using a combination of electrochemical methods. CV and coulometry, in conjunction with X-ray diffraction data, suggest the overall formation of an amorphous MnOx with an Special Issue: Michael Grätzel Festschrift Received: February 19, 2014 Revised: March 28, 2014
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dx.doi.org/10.1021/jp501768n | J. Phys. Chem. C XXXX, XXX, XXX−XXX
The Journal of Physical Chemistry C
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average Mn oxidation state of +3.68. Application of chronoamperometry and atomic force microscopy (AFM) reveal the nature of catalyst nucleation in the stage of early film growth. Electrochemical kinetic experiments using Tafel and reaction-order plots establish the rate law for MnOx nucleation and growth and guided mechanistic proposals for steady-state catalyst deposition. Together, these studies are able to differentiate between the longstanding issue of ECE versus DISP as the mechanism for manganese oxide deposition by demonstrating the latter to be operative and initiated by a progressive nucleation process.
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EXPERIMENTAL SECTION A. Materials. MnCl2·4H2O (99.995% trace metals basis, Strem), phosphoric acid (85 wt % in H2O, 99.99% trace metals basis, Aldrich), KOH (88%, Macron), and KNO3 (99.0 to 100.5%, Macron) were used as received. Methylphosphonic acid (98%, Aldrich) was twice recrystallized from boiling acetonitrile (HPLC grade, Aldrich). Electrodes and glassware were rinsed with Type-I water (EMD Millipore, 18.2 MΩ-cm resistivity), which was also used to prepare solutions. B. Solution Preparation. Stock solutions of 10 mM Mn2+ and 1.0 M methylphosphonate (MePi, pKa1 = 2.35, pKa2 = 7.1,32 preadjusted to target pH with KOH) were diluted with type-I reagent-grade water to prepare electrolytes for deposition studies. The concentrated stock MePi stock solution was first diluted before adding Mn2+, and the electrolyte was permitted to equilibrate for ca. 8 h prior to experiments. Additional KNO3 was added to maintain ∼2 M ionic strength (μ). C. General Electrochemical Methods. A CH Instruments 760D bipotentiostat, a BASi Ag/AgCl reference electrode (filled with saturated KCl), and a Pt-mesh counter electrode were employed for all electrochemical experiments with ohmic potential loss correction (iR compensation). The electrochemical cell was a standard three-electrode configuration with a porous glass frit separating the working and auxiliary compartments. Glassware and cells were cleaned by soaking in aqua regia and rinsing with reagent-grade water. All measurements were performed at ambient temperature (23 ± 1 °C), and electrode potentials were converted to the normal hydrogen electrode (NHE) scale using E(NHE) = E(Ag/AgCl) + 0.197 V. Unless otherwise specified, all reported potentials are referenced to NHE. D. Cyclic Voltammetry. Solutions of 0.5 mM Mn2+ with 50 mM of MePi at pH 3 and 8 were prepared. For each quiescent solution, cyclic voltammograms (CVs) were taken at 10 and 50 mV/s scan rates on an immersed 1 cm2 fluorine-doped tinoxide-coated glass (FTO, TEC-7 purchased from Hartford Glass). All FTO slides were sonicated in acetone and rinsed with reagent-grade water prior to use, and a 1 cm2 electrode area was produced by using a 0.5 cm wide strip of Scotch tape as a mask. The potential range of the scans depended on the pH of the electrolyte (Figure 1), and CVs were permitted to cycle 100 times to visualize the progression of the peaks over time. E. Coulometric Titration. Films of electrodeposited manganese oxide (MnOx) at loadings of 6, 32, 100, and 250 mC/cm2 were prepared on 1 cm2 FTO by controlled potential deposition at 0.577 V in quiescent solutions of 0.5 mM Mn2+ with 50 mM MePi buffer at pH 8.0. The films were dipped in reagent-grade water to remove any remaining solution, dried under ambient conditions, and the Scotch tape used to mask the FTO was carefully removed from the plates. Each plate was
Figure 1. Cyclic voltammograms of a 1 cm2 FTO electrode in 0.5 mM Mn2+ and 50 mM MePi solution at (a) pH 3.0 and (b) pH 8.0, 50 mV/ s scan rate showing the 1st (blue line) and 200th (red line) cycles. (c) Representative CV at pH 8.0 at 10 mV/s (blue line). Region I is attributed to water oxidation while region II denotes MnOx deposition. For comparison, background CV of same electrolyte but without Mn2+ (gray dashed line ).
placed in a polypropylene test tube (Falcon) and sonicated in a small aliquot of 69.7% nitric acid (Fluka Analytical, TraceSelect) and left to dissolve overnight. The blank FTO plate was removed from the tube, and reagent-grade water was added to dilute the existing solution to 5% v/v nitric acid concentration. Samples were analyzed for manganese content by inductively coupled plasma atomic emission spectrometery (ICP-AES; HORIBA Jobin ACTIVA instrument) and quadrupole inductively coupled plasma mass spectrometry (ICP-MS; Elemental Analysis, Inc., Lexington, KY); see Table S1 in the Supporting Information. Manganese standards were prepared from solutions purchased from Fluka, designated suitable for ICP analysis. F. Potential Step Chronoamperometry. Potential step chronoamperometry traces were recorded in 0.5 mM Mn2+, 50 mM MePi, and 1.90 M KNO3 solution with the following electrodes: 2 mm diameter platinum disk button electrode (CH Instruments), 3 mm diameter glassy carbon (GC) disk button electrode (CH Instruments), 1 cm2 TEC-7 FTO on glass slide (Hartford Glass), and 0.15 mm2 highly oriented pyrolytic graphite (HOPG, grade SPI-2, SPI Supplies). The Pt and GC button electrodes were polished to a mirror shine for 60 s with 0.05 μm alumina (CH Instruments) and sonicated for 60 s in reagent-grade water. Prior to each experiment, the Pt button electrodes were cycled ∼50 times from −0.2 to 1.5 V in 0.5 M sulfuric acid (99.999% trace metals basis, Aldrich) and then rinsed with reagent-grade water. GC button electrodes were repolished with alumina before each experiment. Electrodes were initially polarized at 0.397 V for 5 s to charge only the double-layer and then stepped to a more anodic potential sufficient for catalyst nucleation for 60 s (Figure 2 and Figure B
dx.doi.org/10.1021/jp501768n | J. Phys. Chem. C XXXX, XXX, XXX−XXX
The Journal of Physical Chemistry C
Article
between independent experiments. In all cases, after termination of electrolysis, the electrode was immediately removed from solution, rinsed in reagent-grade water, and dried in preparation for imaging by AFM. AFM images were collected using a Veeco Nanoscope Dimension 4100 (Veeco Instruments, Santa Barbara, CA) operating in tapping mode with a Veeco silicon nitride probe with a resonance frequency of ∼200 kHz and average tip radius of 3 nm. All measurements were performed in air and at room temperature. Nucleus size, height, number, and percent coverage values were determined using Nanoscope V5.31 software (Figure 4). Figure 2. Potential step chronoamperograms of a platinum disk electrode in a solution of 0.5 mM Mn2+ and 50 mM MePi with 1.90 M KNO3 electrolyte at pH 8. Step potentials of 0.517 (red dashed dotted line), 0.522 (yellow dashed line), 0.527 (green dashed line), 0.532 (blue line), and 0.537 V (purple dashed line) were applied. All experiments were preceded by an initial 5 s pulse at 0.397 V (not shown) to partially precharge the double layer.
S1 in the Supporting Information). In all cases, data points were sampled every 0.03 s, and iR compensation was applied. Multiple independent experiments on fresh electrodes using the same potential step parameters reproduced the values of peak current (jmax) and time of peak current (tmax) within 5%. G. Atomic Force Microscopy of Nucleation. AFM employed HOPG electrodes (Grade SPI-2, SPI Supplies) prepared by cleaving a thin layer of HOPG from a bulk slab using double-sided Scotch tape and mounting onto a glass slide. The HOPG layer was contacted using a toothless alligator clip, and a 0.1 to 0.15 cm2 area was immersed in solution. A series of partially nucleated catalyst islands on HOPG were prepared by potential step chronoamperometry with a step potential of 0.637 V and times of 0.2, 0.5, 1.0, 2.0, 4.0, and 8.0 × tmax (for tmax ≈ 13 s under these conditions) in 0.5 mM Mn2+, 50 mM MePi, and 1.90 M KNO3 solution (Figure 3 and Figure S2 in the Supporting Information). For longer electrolysis times, the amperometric trace was monitored, and electrolysis was manually terminated at the appropriate time relative to the observed tmax in each run. This procedure served to limit errors due to slight variations (
