REDUCING AGENT

PRIORITY CLAIM

This application claims priority under 35 U.S.C. § 119 to U.S. Provisional Patent Application Serial Number 61/757,247 filed January 28, 2013. The contents of which is incorporated by reference herein in its entirety.

STATEMENT OF GOVERNMENT INTEREST

The present invention was developed using funds from the Office of Naval Research under Grant No. N00014-07- 1-0758. The United States Government has rights to this invention. The present invention was also developed using funding from the Welch Foundation, Grant No. F-1254.

TECHNICAL FIELD

The present disclosure relates to methods of wet chemical synthesis using metallic aluminum (Al) as a reducing agent to reduce a chemical. The metallic aluminum is subjected to corrosion, which may be enhanced by the presence of anions, such as F ^~ , Cl ^~ , and Br-, in protonic solvents, such as water and ethylene glycol (EG). The present disclosure also relates to the use of such methods for synthesizing metallic nanomaterials and metal oxides. The disclosure further relates to the use of such methods to form metal-oxide, metal-carbon, and metal-carbon-oxide composites.

BACKGROUND

Metallic aluminum (Al) is a strong reducing agent with a standard reduction potential E°AI _/ AI = ^~ 1.662 V. The use of Al as a reducing agent is also advantageous because it may be transported and stored easily due to its solid form, participates mostly in environmentally benign reactions, and is very stable because in the presence of air or water, it is very rapidly covered with an aluminum oxide film that is typically a few nanometers thick. This protective film largely prevents reaction of the aluminum with other components of its environment.

The protective aluminum oxide film, however, also serves as a barrier to accessing Al as a reducing agent. As a result, Al is commonly used as a reducing agent only in a few methods that are able to overcome the aluminum oxide film. For example, Al may be used as a reducing agent at high temperatures, such as at 1450 °C in melted glass. Al is also used as a reducing agent in high-energy mechanical milling process. However, the aluminum oxide layer has thus far been a barrier for employing Al as a reducing agent in wet-chemical synthesis (i.e. synthesis done in a liquid phase) because water or other liquids tend to form an aluminum oxide layer at such rapid rates.

SUMMARY

The disclosure provides a method of reducing a chemical by reacting metallic aluminum with an aluminum oxide film, a pitting corrosion enhancer, and a chemical to be reduced in a protonic solvent for a length of time sufficient to produce a reduced chemical.

BRIEF DESCRIPTION OF THE DRAWINGS

The current specification contains color drawings. Copies of these drawings may be obtained from the USPTO. Embodiments of the present invention may be better understood through reference to the following figures in which:

FIGURE 1 provides UV-Vis spectra, photos, and an XRD pattern. The UV-

Vis spectra in FIGURE 1 A at 0 min were obtained for the aqueous Na ₂ PdCl ₄ solutions (10 mM, 0.1 mM) without the addition of Al foil. The UV-Vis spectra at 2 and 5 min were acquired for the supernatants by centrifuging the 10 mM Na ₂ PdCl ₄ aqueous solutions after the addition of Al foil for 2 and 5 min. FIGURE IB contains the images of the supernatants after reacting for 2 and 5 min, and FIGURE 1C is the XRD pattern of the precipitated particles after reacting for 5 min.

FIGURE 2A provides UV-Vis spectra and FIGURE 2B provides pH measurements of Na ₂ PdCl ₄ aqueous solutions with different concentrations.

FIGURE 3 provides SEM images of Al foil surfaces. The Al foils were contacted with a drop of 10 mM Na ₂ PdCl ₄ solution for 1 minutes (FIGURE 3 A) and 3 minutes (FIGURE 3B). The dark spots without Pd NPs (marked with circles) and with Pd NPs (marked with squares) in FIGURE 3A are blisters where pitting corrosion has occurred. The insets of FIGURES 3A and 3B show the magnified images of the Pd NPs on Al foil. FIGURE 3C shows Al foil contacted with a drop of 10 mM CuF ₂ , CuCl ₂ , CuBr ₂ , Cu(N0 ₃ ) ₂ , and CuS0 ₄ aqueous solution for 10 min.

FIGURE 4 provides a schematic illustrating the reaction mechanism (FIGURE 4A and results of the reactions of stated metal salts with Al foil in H ₂ 0 or EG at room temperature (FIGURE 4B). Metal particles were produced with the compounds indicated in green color and no metal particles were produced with the compounds indicated in red color.

FIGURE 5 provides TEM images of Pd/C catalysts, particle size analysis, and SEM images of Al foil surfaces. Pd/C prepared in EG (FIGURE 5A), with PVP in water (FIGURE 5B), and with FeCl ₂ in EG (FIGURE 5C). FIGURE 5D shows particle size distribution histograms of Pd NPs on Pd/C-PVP and Pd/C-FeCl ₂ . FIGURE 5E presents SEM images of Al foil surfaces (left) before and (right) after the preparation of Pd/C-FeCl ₂ .

FIGURE 6 provides an XRD pattern (FIGURE 6A) and SEM and EDX mapping images (FIGURE 6B) of Cu ₂ Sb/C.

FIGURE 7 provides SEM image of the surface of pristine Al foil.

FIGURE 8 provides XRD patterns of synthesized Cu particles. The Cu particles were prepared from the aqueous solutions (20 mM) of CuF ₂ , CuCl ₂ , and CuBr ₂ . The * indicates the oxide phases, which are produced by the oxidation of Cu with air.

FIGURES 9A-9K provide XRD patterns of the indicated metal particles with Al as a reducing agent. The * indicates oxide phases, which are produced by the oxidation of the metal with air.

FIGURE 10 provides photographs of 10 mM Na ₂ PdCl ₄ EG solutions after the addition of Al foil for different times.

FIGURE 11 provides variation of reaction solution pH with the concentration of the Na ₂ PdCl ₄ EG solutions (FIGURE 11 A) and UV-Vis spectra of the Na ₂ PdCl ₄ EG solutions (FIGURE 1 IB) with different concentrations.

FIGURE 12 provides an XRD pattern (FIGURE 12 A) and an EDX spectrum

(FIGURE 12B) of Pd/C prepared in EG without a stabilizer (denoted as Pd/C-EG).

FIGURE 13 provides an XRD pattern (FIGURE 13 A) and an EDX spectrum (FIGURE 13B) of Pd/C prepared with PVP as a stabilizer (denoted as Pd/C-PVP).

FIGURE 14 provides an XRD pattern (FIGURE 14 A) and an EDX spectrum (FIGURE 14B) of Pd/C prepared with FeCl ₂ as a stabilizer (denoted as Pd/C-FeCl ₂ ).

FIGURE 15 provides electrochemical measurements. CVs (FIGURE 15 A) and FOR curves (FIGURE 15B) of the Pd/C catalysts and a commercial Pt/C catalyst, respectively, in N2-saturated 0.5 M H ₂ SO ₄ with the potential cycled in the range of 50 - 1200 mV and in N ₂ -saturated 0.5 M H ₂ S0 ₄ + 0.5 M HCOOH with the potential scanned from 100 to 1100 mV. Both were conducted at a scan rate of 50 mV S ¹ at room temperature.

DETAILED DESCRIPTION

The present disclosure provides methods of wet chemical synthesis using metallic aluminum (Al) as a reducing agent to reduce a chemical.

In methods of the present disclosure, in a liquid protonic solvent such as water or ethylene glycol, Al is exposed to an agent able to enhance pitting corrosion at the same time it is exposed to the chemical to be reduced. The chemical to be reduced may be a compound or mixture of compounds containing metal ions. The ability of Al to reduce any given chemical may be readily determined by reference to the standard reduction potential of Al and the standard reduction potential or similar property of the chemical to be reduced.

Agents able to enhance pitting corrosion in Al include anions, such as F ^~ , Cl ^~ ,

Br-, SCN-, C10 ₃ ^~ , C10 ₄ ^~ , and gluconate anions The lower limit of the standard reduction potential for metal compounds to be reduced with Al may be in the range of -0.138 to -0.257 V. The Al and anions may be present in aqueous solutions or on other solutions that allow pitting corrosion to occur. The rate at which pitting corrosion occurs may vary depending upon the anion, the solution, accessibility of Al, which may be influenced by anion size, and other factors able to affect the pitting corrosion process. The Al reduction reaction is dependent upon the availability of Al via the pitting corrosion. Conditions may also be chosen to facilitate reduction of the reduced chemical or to avoid damage to the chemical to be reduced or its reduction product. However, the methods of the present disclosure may be used with a wide variety of temperatures, pH, and atmospheres. For some reaction products or some variations of the method, however, temperature, pH, and atmosphere may be controlled.

Methods of the present disclosure may be used to produce reduced single metal products, such as Pt, Au, Pd, Ag, Cu, Sb, Bi, Sn or Pb. Alloys and intermetallics show many unique properties that pure metals do not, allowing these materials a wider range of applications and the ability to customize properties by tailoring the composition. Thus, the methods of the present disclosure can also be extended to prepare bimetallic and multi-component metallic materials within the range of the reducing capability. For example, the intermetallic Cu ₂ Sb has been synthesized previously due to its potential application as a high capacity lithium-ion battery anode material. These prior synthesis methods require long reaction times, high temperatures, specialized equipment, or hazardous and toxic chemicals. Using methods described herein, and in particular those set forth in Example 3, the intermetallic compound Cu ₂ Sb/C may be synthesized by the simple, environmentally friendly method of adding Al foil as the reducing agent to an EG solution of the metal chlorides at room temperature without the use of specialized equipment.

The anion in method of the present disclosure may be provided to the reaction solution as part of a compound with the chemical to be reduced or other compositions added to the reaction solution, such as stabilizers. For example, reduced Pd may be obtained by exposing Al with an aluminum oxide coating to an aqueous solution of Na ₂ PdCl ₄ . In this method, the Na ₂ PdCl ₄ compound dissociates in water and the CI ^"

2_ _| _

ions cause pitting corrosion of the Al, which in turn reduces the Pd to form metallic Pd.

Al used in the methods may be in any form, such as foil or particles. The Al may already exhibit some pitting corrosion prior to exposure to a solution containing an anion.

Al (and its residual aluminum oxide film) may be separated from the reduced chemical product by mechanical removal.

Resulting reduced chemical products may be present in small particles, such as nanoparticles (particles 1000 nm or less in average length or diameter). One or more stabilizers may be added to the reaction solution in order to produce smaller particles.

Methods of the present disclosure may be used to form metal-carbon products. In such methods, a carbon source, such as carbon black, may be added to the reaction solution. Reaction conditions may be such that the reduced metal is transferred to the carbon. In certain methods, ultrasonication and scrubbing or similar interventions to remove the reduced metal from the Al are not required. Metal-oxide-carbon products and metal-oxide products, such as MnO _x may be produced in a similar fashion. Additional oxides, such as A1 ₂ 0 ₃ may be added for metal-oxide or metal-oxide-carbon products. Agents may be added to the reaction solutions to obtain uniform dispersal of metal on its substrate in such products.

Products of the present reaction methods may be useful in electrochemical systems, such as batteries, fuel cells, and other energy storage and conversion devices. For example, they may be used as anodes, cathodes, or electrocatalysts.

Without limiting the invention to a particular method of action, in some embodiments, the pitting corrosion of Al with an aluminum oxide surface may occur by adsorption of the anion onto the aluminum oxide film, penetration of the anion through the aluminum oxide film via oxygen vacancies, and localized dissolution of Al below the aluminum oxide film to form H, which may travel through the

2_ _| _

aluminum oxide layer to react with Pd , resulting in H ₂ gas and reduced metallic Pd. The instant formation of oxide film in the protonic solvent ensures that Al does not have a chance to directly react with the metal ions for reduction. The non-conductive nature of the oxide film means that the electrons released from the electro-oxidation of Al cannot travel through the oxide film to participate in the electro-reduction of the metal ions on the top surface. For example, some reactive metals, such as Zn and Fe, do not follow this mechanism because their oxide films are semiconducting. Instead, H performs this function.

The details of this mechanism are provided in FIGURE 4A and are discussed in Examples 1 and 2. This mechanism is different than the redox reaction between Al and cations conventionally assumed to occur. However, the methods of the present disclosure are intended to include such a conventional redox reaction or any other reaction that does, in fact, occur when Al exposed to a pitting corrosion agent is used to reduce a chemical.

EXAMPLES

The present invention may be better understood through reference to the following examples. These examples are included to describe exemplary embodiments only and should not be interpreted to encompass the entire breadth of the invention.

Example 1 - Mechanism of Pitting Corrosion

The UV-Vis spectra in FIGURE 1 show that a 10 mM Na ₂ PdCl ₄ aqueous solution was reduced by the addition of Al foil with > 95 % conversion at room temperature within 5 min. The precipitated particles were confirmed as metallic Pd by the XPvD pattern (right inset of FIGURE 1). These results demonstrate that the aluminum oxide film on Al foil is overcome and Pd ions are reduced without requiring any assistance such as ultrasonication or mechanical force, or any particular conditions, such as pH, temperature or atmosphere.

The mechanism for [PdCl ₄ ] to overcome the dense oxide film on Al foil appears to be pitting corrosion via Cl ^~ . The inset in FIGURE 2 shows that the pH of deionized water sharply decreases from about 6 to < 4 on adding a small amount of Na ₂ PdCl ₄ . This suggests that the [PdCl ₄ ] complex releases free CI ^" by the replacement of the Cl ^~ ligand with OH- from water to form [PdCl _x (OH) _y ] , leading to an increase in H ⁺ concentration in the solution. It is also supported by a shifting of the peak position of the Pd d-d spin-forbidden transition to a lower value in the UV-Vis spectra in FIGURE 2 as the concentration of Na ₂ PdCl ₄ decreases.

To confirm the formation of pitting corrosion on the Al foil with the exclusion of stirring effects, two drops (20 μί) of a 10 mM Na ₂ PdCl ₄ aqueous solution were, respectively, placed on Al foil for 1 min and 3 min and then the Al foil was rinsed with water. The dark spots without Pd nanoparticles (NPs) and with Pd NPs on the Al foil shown in FIGURE 3A are blisters where pitting corrosion has occurred, causing the oxide film to bubble. The reactions may occur first at the defect sites on the pristine surface of the Al foil (FIGURE 7), where the positively charged hydrated aluminum oxide surface is easily attacked by Cl ^~ ion. As the contacting time increases, the corrosion below the aluminum oxide film at those sites propagates and more new sites start pitting. Accordingly, the already formed Pd particles grow and more Pd particles nucleate as shown in FIGURE 3B. When the pitting corrosion evolves to the stage of blistering, the high pressure hydrogen gas built up in the blisters ruptures the oxide film and causes the Pd particles to detach from the Al foil. New oxide film forms atop the Al instantly because of the fast reaction between Al and water (3 x 10 ^~9 s at 25 °C). New pitting corrosion thus occurs below the new oxide films and Pd particles are produced nearby as shown in the right inset of FIGURE 3B.

To confirm that pitting corrosion assists the use of Al as a reducing agent, Al foil was added into 20 mM aqueous solutions of CuF ₂ , CuCl ₂ , CuBr ₂ , CuS0 ₄ and Cu(N0 ₃ ) ₂ . F ^~ , CI , and Br are known to enhance the pitting corrosion of Al foil with different rates. S0 ₄ and N0 ₃ ^~ are known to not cause pitting corrosion of Al. Cu particles, confirmed by XRD (FIGURE 8), were produced within minutes in the solutions of CuF ₂ , CuBr ₂ , and CuCl ₂ . The SEM images in FIGURE 3C also show that the Al foil after being in contact with drops of CuF ₂ , CuBr ₂ , and CuCl ₂ solutions (10 mM) for 10 min had Cu particles. In contrast, there were no notable changes in the solutions of CuS0 ₄ and Cu(N0 ₃ ) ₂ even after a much longer time and no evidence of corrosion pits or Cu particles, indicating pitting corrosion, which is not caused by

2— 2+

S0 ₄ and N0 ₃ - , is responsible for the reduction to Cu to form metallic Cu. Another observation is that the number of Cu particles increases as the size of the anions decreases, i.e., the Al foil with CuF ₂ has the most Cu particles and that with CuBr ₂ has the least. This occurs because a smaller anion can more easily transport through the oxide film during pitting corrosion.

Example 2 - Mechanism of Reduction Reaction

Water molecules transport through the aluminum oxide film and react with metallic Al beneath to produce H ₂ gas, in which H is generated as an intermediate. Accordingly, H ₂ gas, H, and metallic Al are all possible reducing agents in the reactions of the present methods. A controlled experiment with the passage of H ₂ gas overnight revealed that H ₂ cannot reduce an aqueous solution of Na ₂ PdCl ₄ . In contrast, H is more active than H ₂ and it can reduce metal oxides such as CuO at room temperature. As described in Example 1, Pd particles form on Al foil in the early stage of pitting corrosion, when the oxide film is not broken. Therefore, the electro- reduction of [PdCl _x (OH) _y ] in the solution outside of the aluminum oxide does not obtain electrons released from the Al electro-oxidation reaction beneath the non- conductive aluminum oxide film, but instead obtains electrons from the electro- oxidation reaction of H on the top of the aluminum oxide film. The H is produced as an intermediate of the pitting corrosion of Al below the oxide film. It cannot reduce aluminum oxide because of the very low standard reduction potential of E° _A I ₍ O _H)3/ AI = -2.31 V, but it can diffuse through to the top of the aluminum oxide film. The hydrogen electro-oxidation reaction continues on the top of the aluminum oxide film and the released electrons flow onto the surface of the formed Pd particles for further reduction of [PdCl _x (OH) _y ] , leading to the growth of the Pd particles. As mentioned above, the Pd particles will become detached when the aluminum oxide film beneath is broken and a new oxide film forms atop the metallic Al instantly (3 x 10 ⁹ s at 25°C). The whole process is shown by the schematic drawing in FIGURE 4A. It was observed that the Pd particles develop into dendrites (left insert in FIGURE 3B), suggesting that the growth rate is limited by the diffusion of the solute species from the bulk solution to the surface of the particles. It also suggests an abundant supply of the reductant, which is supported with the fast hydrogen electrooxidation reaction and the high mobility of electrons on and within the Pd particles.

Al reduces metal compounds with positive standard reduction potentials and anions that cause significant pitting corrosion of Al. Various reactions were performed to explore reduction using the present methods. Results are shown in FIGURE 4B. Some reactions were conducted in ethylene glycol (EG) to avoid a hydrolysis reaction in water. Most reactions were preformed with compounds with a positive reduction potential, but some metal compounds with negative reduction potential, such as FeCl ₃ (E° _Fe ³⁺ _/Fe = -0.037 V), PbCl ₂ (E ⁰ _Pb ²⁺ _/P b = "0.126 V), SnCl ₂ (E°sn ²⁺ /sn = -0.138 V), CoCl ₂ (E°co ²⁺ /Co = -0.28 V), NiCl ₂ (E° _NI ²⁺ / = ^" 0.257 V), and FeCl ₂ (E° _Fe ²⁺ _/Fe = -0.447 V), were also tested.

Metallic particles were produced in the solutions of PbCl ₂ and SnCl ₂ but none were found in the solutions of FeCl ₃ , CoCl ₂ , NiCl ₂ , and FeCl ₂ . The FeCl ₃ solution is a unique case in that it is reduced to FeCl ₂ owing to E° _Fe ²⁺ / _Fe ³⁺ = 0.77 V, which is more positive than E° _Fe ²⁺ / _Fe and therefore the favored reaction. The lower limit of the standard reduction potential for metal compounds to be reduced with Al is located in the range of -0.138 to -0.257 V. This indicates that H is the actual reducing agent because the reduction potential of H/H ⁺ is about -0.17 and -0.24 V, respectively, at pH = 3 and 4 as calculated by the Nernst equation. All produced metallic particles were confirmed by XRD (FIGURE 9).

Example 3 - Synthesis of Metallic Nanoparticles

Two carbon supported metallic nanomaterials were prepared: 40 wt.% Pd on carbon support (Pd/C), which may be used as an electrocatalyst for formic acid electro-oxidation in fuel cells, and intermetallic Cu ₂ Sb/C, which may be used as an anode material for lithium-ion batteries. Both Al foil and powder were used to prepare Pd/C. Al foil is discussed in the present example for the following reasons: the reduction rate with Al foil is faster than with Al powder because the latter has a larger ratio of aluminum oxide to aluminum; it is easier to separate Al foil from the reaction products; and in the stirring reaction solution, Al foil has a lower moving speed than carbon black while Al powder has a speed similar to carbon black, implying a larger relative moving speed between Al foil and carbon black and thus a larger scrubbing force to detach the Pd NPs from the Al surface.

Carbon black cannot be well dispersed in water, so the preparation of Pd/C was carried out in EG. The reduction of 10 mM Na ₂ PdCl ₄ by Al foil in EG was complete in about 1 h (FIGURE 10). The same mechanism occurs in EG solution as described above for water, as is shown by the change in pH and UV-Vis peaks with concentration (FIGURE 11). Accordingly, Pd/C was synthesized in EG with Al foil at room temperature for 3 h. The obtained Pd/C (denoted as Pd/C-EG), confirmed by XRD and EDX (FIGURE 12), has very large Pd particles as shown in FIGURE 5 A. Such large particles may be useful in some applications, but may not be preferred for use in electrochemical systems.

Trisodium citrate and polyvinylpyrrolidone (PVP) were tested as stabilizers to reduce the particle size. The reactions were carried out in water because carbon black disperses well in water with these stabilizers. The aqueous Na ₂ PdCl ₄ solution with PVP became dark within a few minutes after the addition of Al foil and the reaction was complete in about 15 min. In contrast, the solution with trisodium citrate did not show a noticeable change in 2 h. The difference may reflect that PVP can interact with the metallic Pd atoms through its carbonyl group, while trisodium citrate

2_ _| _

complexes with the Pd ions. In addition, citrate can inhibit the pitting corrosion of Al by a competitive adsorption with Cl ^~ on the oxide film. FIGURES 5B and 5D show that the Pd/C (denoted as Pd/C-PVP) prepared with a 10: 1 ratio of PVP repeat unit to Na ₂ PdCl ₄ in water at room temperature for 1 h results in Pd NPs with a mean particle diameter of 4.9 nm (additional XRD and EDX data are provided in FIGURE 13), indicating a significant size reduction compared to that of Pd/C-EG shown in FIGURE 5A. However, the dispersion of Pd NPs on the carbon is irregular; some carbon black is not loaded with NPs. FeCl ₂ was also used as a stabilizer to synthesize Pd/C (denoted as Pd/C-FeCl ₂ ) in EG at room temperature for 3 h, which was confirmed by XRD and EDX results (FIGURE 14). The Pd NPs are stabilized by the adsorption of Fe ²⁺ . FIGURES 5C and 5D show that the Pd/C-FeCl ₂ prepared with a 10:1 ratio of Na ₂ PdCl ₄ to FeCl ₂ had Pd NPs uniformly distributed on the carbon black with a mean diameter of 6.5 nm. Although Pd/C-FeCl ₂ has larger Pd NPs than Pd/C-PVP, the carbon black is more uniformly loaded with Pd NPs. Electrochemical evaluation (FIGURE 15) shows that Pd/C-FeCl ₂ and Pd/C-PVP have high activity toward formic acid electro-oxidation. Pitting corrosion occurrence during the preparation of Pd/C is shown by the SEM images of Al foil surfaces in FIGURE 5E.

Other stabilizers able to adsorb on certain metal crystalline planes can be adopted to tune the particle shape only if it does not inhibit the pitting corrosion of Al. For example, cetyltrimethylammonium bromide is a widely used surfactant to synthesize cubic-like metal nanoparticles. It can provide bromide anions in an aqueous solution, which can activate the Al as a reducing agent.

Carbon-supported Cu ₂ Sb (denoted Cu ₂ Sb/C) was prepared using Al foil as a reducing agent in EG under an N ₂ atmosphere and confirmed by XRD and SEM-EDX mapping in FIGURE 6. The molar ratio of Cu to Sb in the product is 2.05 as determined by EDX, indicating complete reduction of both CuCl ₂ and SbCl ₃ . Similar to Pd/C, the carbon black works as a scrubber to remove the formed Cu ₂ Sb from Al foil. The quantitative reduction of Cu and Sb could be attributed to their similar reduction reaction rates. The crystallite size of Cu ₂ Sb is about 20 nm, as determined by the Scherrer equation. The small crystalline size, in the presence of a carbon matrix, has the potential to buffer the large volume expansion of the anode during charge/discharge of lithium-ion batteries.

Example 4: Material Preparation and Characterization Methods

The following preparation and characterization methods were used in connection with Examples 1-3.

Synthesis

Aluminum (Al) foil (24 μιη thick, Fisher Sci.) was cut into pieces of approximately 1 cm x 3 cm prior to the use in the reactions. Al powder (325 mesh, Alfa Aesar), metal compounds, polyvinylpyrrolidone (PVP, MW = 40000, MP Biomedicals), and FeCl ₂ -4H ₂ 0 were used as received. The solvent was deionized water or ethylene glycol (EG) for different metal compounds.

A solution of the desired metal compounds (AgF, AgN0 ₃ , BiCl ₃ /EG, CoCl ₂ -6H ₂ 0, Cu(N0 ₃ ) ₂ -3H ₂ 0, CuBr ₂ , CuCl ₂ , CuF ₂ , CuS0 ₄ -5H ₂ 0, FeCl ₂ -4H ₂ 0, FeCl ₃ -6H ₂ 0, H ₂ PtCl ₆ -6H ₂ 0, HAuCl ₄ xH ₂ 0, InCl ₃ xH ₂ 0, K ₂ PtCl ₄ , Na ₂ PdCl ₄ , NiCl ₂ -6H ₂ 0, PbCl ₂ , PdCl ₂ , SbCl ₃ /EG, and SnCl ₂ -2H ₂ 0/EG) was prepared in either deionized water (CuCl ₂ and all others) or EG (BiCl ₃ , CuCl ₂ , InCl ₃ , PbCl ₂ , SbCl ₃ , and SnCl ₂ ). In most cases, a 20 mM solution was prepared. The required amount of Al foil was then added to the stirring solution to result in a molar ratio between Al and the metal compound of 10: 1. The produced metallic powders were separated by centrifugation followed by washing in either deionized water or isopropanol.

The typical synthesis for Pd/C (40 wt.%) in EG (denoted as Pd/C-EG) proceeds as follows. Carbon black (90 mg, Vulcan XC-72R, Cabot) was dispersed in 30 mL EG by vigorously stirring for at least 3 h. A EG solution of Na ₂ PdCl ₄ (5.638 mL, 0.1 M) was added drop-wise into the carbon dispersion and stirred for 1 h. Al foil, at a 10: 1 molar ratio, was then added and the suspension was kept stirring for 3 h at room temperature in air. To remove a majority of the unreacted Al, the suspension was kept sedentary for 15 min. The upper suspension was then centrifuged at 20,000 g for 15 min. The precipitated Pd/C-EG powders were then redispersed in a solution of 20 mL water and 5 mL ethanol, into which 1.0 M sulfuric acid was added to result in a final concentration of 0.02 M and stirred overnight to remove the residual Al and aluminum oxide. The Al-free Pd/C-EG powder was then centrifuged and washed in a water/ethanol solution by stirring for 1 h a total of three times. Lastly, the washed sample was dried under vacuum at room temperature.

To prepare Pd/C employing FeCl ₂ as a stabilizer (denoted as Pd/C-FeCl ₂ ),

FeCl ₂ -4H ₂ 0 was added to an EG solution of Na ₂ PdCl ₄ (0.1 M) at a 10: 1 molar ratio of Na ₂ PdCl ₄ to FeCl ₂ and the solution was stirred for 1 h before being added to the carbon black dispersion. As for the Pd/C prepared with PVP as a stabilizer (denoted as Pd/C-PVP), carbon black was dispersed in water with PVP at a 10:1 molar ratio of PVP repeat unit to Na ₂ PdCl ₄ by stirring for 3 h, and then aqueous Na ₂ PdCl ₄ solution (0.1 M) was added and stirred for 1 h.

The intermetallic Cu ₂ Sb supported on carbon black (acetylene black, 100 % compressed, Alfa Aaser) was prepared in a similar manner to Pd/C-EG. Briefly, 100 mg carbon was dispersed in EG to result in a 20 wt. % composite with Cu ₂ Sb. After dissolving 432.2 mg CuCl ₂ and 366.7 mg SbCl ₃ , Al foil was added at a 25: 1 molar ratio of Al to Cu + Sb and the reaction took place under an N ₂ -atmosphere for 4 h. After straining the unreacted Al, the mixture was filtered and washed with isopropanol before drying in a vacuum oven overnight.

Physical characterization

The pH of the solutions was measured by a Corning 313 pH/Temperature meter. The concentrations of Na ₂ PdCl ₄ solutions were characterized by the Pd d-d spin-forbidden transition peaks between 420 - 450 nm in UV-Vis spectra, recorded with a Cary 5000 UV-Vis-NIR spectrophotometer (Varian) using ultramicro UV- cuvettes (Fisher Sci.). To determine the reaction conversion with time, Al foil was added to a 10 mM Na ₂ PdCl ₄ solution with a 10:1 molar ratio of Al to Na ₂ PdCl ₄ for different times. The reaction solution was centrifuged at 10,000 g for 5 min, and the supernatant was characterized by UV-Vis.

X-ray diffraction (XRD) patterns of the samples were recorded with a Philips APD 3520 diffractometer with Cu Ka radiation (λ = 0.15418 nm) and analyzed with the JADE 9.0 software package (Rigaku). The scanning electron microscopy (SEM) images for Al foils and Cu ₂ Sb/C were taken, respectively, with JEOL-JSM5610 and Hitachi S5500. The compositions of the samples were determined using INC A software (Oxford Instruments) to quantify the energy-dispersive X-ray spectroscopy (EDS) spectra obtained by an EDS attachment (Oxford instruments) on the JEOL SEM. Transmission electron microscope (TEM) images were acquired with a JEOL 201 OF operated at 200 keV to characterize the morphology of the samples. The particle size distributions were obtained by analyzing the TEM images using ImageJ software (NIH).

Electrochemical measurements

Electrochemical experiments were carried out on a potentiostat (Autolab PGSTAT302N, Eco Chemie B.V.) in a three-electrode configuration: a Pt mesh as the counter electrode, a glassy carbon rotating disk electrode (RDE, 5 mm diameter, Pine Instrument) as the working electrode, and a saturated calomel electrode as the reference electrode. All potentials in this study are reported against the reversible hydrogen electrode (RHE).

The catalyst ink (5.0 mg _cat ml/ ¹ ) was prepared by dispersing the Pd/C catalyst in a 1 :5 (v/v) deionized water/isopropanol mixture by sonication for 15 min. A certain volume of the ink was cast onto the freshly polished RDE to give a loading of 51.0 μgmetal cm and dried at room temperature. The metal loadings of Pd/C-EG, Pd/C- PVP, and Pd/C-FeCl ₂ were, respectively, 38.8, 32.6, and 38.3 wt.%, as determined by burning off the carbon black and oxidizing Pd to PdO at 550 °C for 4 h in air.

Cyclic voltammetry experiments were conducted in an N ₂ -saturated 0.5 M H ₂ SO ₄ solution (ACS, Fisher Sci.). The 6th cyclic voltammograms in the range of 0.05 - 1.2 V at a scan rate of 50 mV s ^_1 are presented. The catalyst-coated RDE was mounted onto an interchangeable RDE holder (Pine Instruments) and immersed in an N ₂ -saturated solution of 0.5 M H ₂ SO ₄ and 0.5 M formic acid. Formic acid electrooxidation experiments were performed with a rotation rate of 1000 rpm and a potential sweep between 0.1 and 1.1 V at a scan rate of 50 mV s ^-1 . The second forward sweep voltammograms are presented.

Electrochemically active surface area

The electrochemically active surface areas (ECS As) of the commercial Pt/C catalyst and the Pd/C catalysts were characterized as follows.

The charge of monolayer hydrogen underpotential deposition (HU _PD ) occurring on Pt NPs in the range of 0.05 - 0.35 V was used to determine the ECSA of Pt/C with the equation

-4

ECSA _P [m ² g' ¹ ] = ^{HupD char} g ^e ^C]/210 ^C - cm _Pt ]) x lO

Pt mass [g]

As for the Pd/C catalysts, since the charge of the HU _PD cannot be accurately determined because of the interference of hydrogen absorption in Pd, the ECSA was estimated by measuring the reduction charge of surface PdO in the range of 0.6 - 0.9 V with a charge density of 424 μC cm with the equation,

_{R 2} _ _{l n} (PdO reduction charge [ uC] / 424 [uC · cm _p ² 1) x 10 ^-4

ECSA _pd [m · g ] = — —

Pd mass [g]

Formic acid oxidation activity

The ECS As of the Pd/C-EG, Pd/C-PVP, Pd/C-FeCl ₂ , and the commercial Pt/C (40 wt.%, Johnson Matthey) are, respectively, 14.3, 45.5, 33.4, and 53.3 m ² g ^_1 from the cyclic voltammograms (CVs) shown in FIGURE 14. The formic acid oxidation reaction (FOR) curves in FIGURE 14 show that the prepared Pd/Cs have lower onset potentials and much higher activities compared to the commercial Pt/C, consistent with the literature data. This validates the high performances of the Pd/C catalyst prepared by this method.

Although only exemplary embodiments of the invention are specifically described above, it will be appreciated that modifications and variations of these examples are possible without departing from the spirit and intended scope of the invention. For example, throughout the specification particular measurements are given. It would be understood by one of ordinary skill in the art that in many instances, particularly outside of the examples, other values similar to, but not exactly the same as the given measurements may be equivalent and may also be encompassed by the present invention.