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Room Temperature Turkevich Method: Formation of Gold Nanoparticles at the Speed of Mixing Using Cyclic Oxocarbon Reducing Agents Nathaniel E. Larm, Jeremy B. Essner, Keagan Pokpas, James A. Canon, Nazeem Jahed, Emmanuel Iheanyichukwu Iwuoha, and Gary A. Baker J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b10536 • Publication Date (Web): 05 Feb 2018 Downloaded from http://pubs.acs.org on February 14, 2018

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The Journal of Physical Chemistry C is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

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The Journal of Physical Chemistry

Room Temperature Turkevich Method: Formation of Gold Nanoparticles at the Speed of Mixing Using Cyclic Oxocarbon Reducing Agents

Nathaniel E. Larm,a Jeremy B. Essner,a Keagan Pokpas,b James A. Canon,a,c Nazeem Jahed,b Emmanuel I. Iwuoha,b and Gary A. Bakera*

a

Department of Chemistry, University of Missouri, Columbia, MO, 65211, USA.

b

Department of Chemistry, University of the Western Cape, Bellville, South Africa

c

Southwest Baptist University, Bolivar, MO, 65613, USA.

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ABSTRACT We demonstrate a facile and reproducible means of producing quasi-spherical, colloidally stable gold nanoparticles (AuNPs) based on rapid room temperature mixing of aqueous solutions of HAuCl4 and a cyclic oxocarbon diacid (squaric acid, SA; croconic acid, CA; or rhodizonic acid, SR) or ascorbic acid (AA) as dual reducing and capping agent. Although these reducing agents generally produced larger particles than those derived from the classical Turkevich method (using citrate in boiling water) and achieved a lower nanoparticle size uniformity in our hands (i.e., 30.4 ± 8.6 nm, 33.1 ± 9.3 nm, 29.9 ± 6.3 nm, and 29.7 ± 7.6 nm for SA, AA, CA, and SR, respectively, compared with 15.8 ± 3.7 nm for citrate), the method is versatile and exceptionally convenient as fairly monodisperse AuNPs can be made “on-demand” within seconds by simple mixing in the absence of heating. A preliminary investigation into the effects of reaction parameters such as synthesis temperature and the molar ratio of reducing agent to HAuCl4 was carried out. The reagent molar ratio was found to play a pivotal role in the mean AuNP size and size distribution whereas reaction temperature (e.g., 5, 20, or 100 °C) only played a very minor role. Interestingly, CA- and SR-mediated reduction generated AuNPs displaying bimodal size distributions, with a large total fraction number of the total nanoparticle count being represented by small AuNPs in the 3.5 ± 1.9 nm (CA) and 5.1 ± 1.0 nm (SR) size regime. Cyclic and differential pulse voltammetry were conducted to gain insight into the redox chemistry of the cyclic oxocarbons as prospective reducing agents for general metal nanoparticle synthesis as well as to furnish additional evidence in support of a proposed mechanism for the overall oxidation process using squaric acid as a representative cyclic oxocarbon acid. Finally, the catalytic activities of the prepared AuNPs were evaluated using the borohydride-assisted reduction of 4-nitrophenol as a model reaction, exhibiting apparent rates of 2.0 × 10–3 s–1, 3.6 × 10–3 s–1, 1.9 × 10–3 s–1, and 13.8 × 10–3 s–1 for SA-, AA-, CA-, and SR-derived AuNPs, respectively (5 mol% catalyst). Notably, AuNPs generated using SR boasted a catalytic rate twice as high as that of Turkevich (citrate)-derived AuNPs at the same Au catalyst loading, an outcome we attribute to the prevalence of ultra-small (~5 nm) AuNPs produced in that sample. Overall, these findings open the possibility for “on-the-fly” nanomanufacturing methods (e.g., “glow stick”-inspired preparation) that allow the expedient, reproducible, and lowcost synthesis of metal nanoparticles with minimal environmental impact.

CORRESPONDING AUTHOR *Email: [email protected]

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INTRODUCTION Bulk gold, due to its alluring lustre, has been used for economical and decorative purposes since the beginning of civilization. In addition, although the origin of the dazzling optical properties was a mystery at the time, early civilizations also exploited the unique properties of nanoscale gold structures.1 Indeed, it wasn’t until the last two centuries that the nanoscale origin of these unique properties was uncovered, and only recently have the physicochemical properties of gold nanostructures been fundamentally investigated.2-3 It is now well appreciated that gold nanoparticles (AuNPs), due to their unique size- and morphology-dependent properties, hold promise in a wide range of applications such as catalysis,4-6 drug and gene delivery,7-8 cancer therapy,9-11 and electronics,12-13 to name a few. Historically, the reduction of Au+ or Au3+ salts for the formation of AuNPs has been accomplished by various techniques including vapor deposition,14 thermal decomposition,15 and chemical reduction in aqueous16-18 or nonaqueous19 media. However, these approaches often require tedious, complex, or manifold steps and also frequently call for the use of hazardous or toxic reagents. Therefore, there is considerable impetus to uncover greener and operationally-simpler approaches to AuNPs comparable in performance to those obtained from more conventional routes.20-26 First reported by Turkevich et al. in 1951,16 and further modified by Frens in 1973,27 the most widely employed and studied approach for making water-soluble AuNPs involves the reduction and subsequent stabilization (capping) of AuCl4– by citrate in boiling water. While this method has been widely known and followed for decades, the experimental parameters have been vigorously scrutinized in recent years28 in an effort to optimize colloid monodispersity and arrive at a deeper fundamental understanding of the

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reduction processes occurring within this synthetic system.29-33 Given the facile nature of this approach and the quite monodisperse AuNPs that result, the Turkevich method is often used as a benchmark when considering new synthetic approaches. There are limitations, however. While this approach is fairly green, it was determined by OjeaJiménez et al.31 that the thermal degradation of citrate is the key step in the formation and stabilization of AuNPs in the Turkevich method, making it a necessity to bring the AuCl4– solution (or the citrate solution in the so-called inverse Turkevich method)31 to boiling (or near boiling, >80 °C)16 during AuNP synthesis. This inefficient, energyintensive, and time-consuming necessary step may also introduce some error when assuming the final gold concentration (due to evaporative losses), leading to inaccuracy in concentration-dependent catalytic assessments. Therefore, there is considerable impetus to develop expedient and operationally-simple methods to rapidly synthesize stable and well-defined AuNPs from AuCl4–(aq) at room temperature in a single vessel. We considered two benign reducing agents in the current work: ascorbic acid and cyclic oxocarbon diacids,34-35 exemplified by squaric acid. Interestingly, the carbonyls in squarate (Scheme 1A) and the other cyclic oxocarbons are not ketonic but rather are similar to acid carbonyls (i.e., all four oxygen atoms should become equivalent through resonance). The ionized forms of these acids constitute symmetrical, electron-delocalized anions of composition CnOn2– which have considerable aromatic character. We note that typical methods utilizing ascorbic acid as a reducing agent for AuNP formation require heating,36 entail bi-phasic media,37 employ seed-mediated growth,38 or require a separate capping agent.39 Meanwhile, room temperature studies using ascorbic acid as a dual reducing/capping agent are typically lacking in scope.40 Likewise, although the synthesis of AuNPs using squaric acid was recently reported,17 many important open questions

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remain regarding the use of oxocarbons as reducing agents for metal nanoparticle synthesis. Herein, ascorbic acid (AA) and the cyclic oxocarbon acids squaric acid (SA), croconic acid (CA), and sodium rhodizonate (SR) were studied as dual reducing and capping agents for the facile and rapid aqueous synthesis of colloidally-stable AuNPs upon rapid mixing at room temperature. The localized surface plasmon resonance (LSPR) of the synthesized AuNPs was used as a tool to determine the impact of experimental parameters (temperature and the R value, defined as the molar ratio of reducing agent to gold) on colloid size and monodispersity. Analysis of transmission electron microscopy (TEM) micrographs allowed us to quantitatively assess how AuNP size and shape dispersity are controlled by choice of reducing/capping agent and experimental parameters. Nuclear magnetic resonance (NMR) spectroscopy provided insight into the oxidation pathways of SA during the reduction of AuCl4–, allowing us to propose a mechanism similar to that suggested for AA previously.41 Cyclic voltammetry and differential pulse voltammetry experiments were performed to elucidate the redox behavior of each reducing agent and provide further evidence in support of our proposed reduction mechanism. Additionally, the reaction pH was analyzed to determine the pHdependent speciation of Au3+ in solution and to follow reaction kinetics. Finally, the catalytic activities of as-synthesized and aged (at least 6 months) AuNPs prepared using these reductants were assessed using the borohydride-assisted reduction of 4-nitrophenol to 4-aminophenol as a model reaction for benchmarking catalytic activity.42-43 EXPERIMENTAL SECTION Materials and reagents. All experiments were carried out using Ultrapure Millipore water (H2O) purified to a resistivity of 18.2 MΩ·cm. Squaric acid (123447, 99%),

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croconic acid disodium salt (391719, 97%), sodium rhodizonate dibasic (R1609, 97%), Lascorbic acid (05878, 99.9%), sodium citrate (C-7254, ≥98%), gold(I) chloride (481130, 99.9%), gold(III) chloride (520918, ≥99.9%), silver nitrate (209139, 99.0%), chloroplatinic acid hydrate (520896, ≥99.9%), potassium tetrachloropalladate(II) (205796, 98%), 4-nitrophenol (241326, ≥99%), sodium borohydride (480886, 99.99%), and sulfuric acid (339741, 99.999 %) were all purchased from Sigma-Aldrich (St. Louis, MO). Deuterium oxide (DLM-6-10X1; D, 99.96%) was acquired from Cambridge Isotope Laboratories, Inc (Tewksbury, MA). Characterization techniques. UV-Vis spectra were measured in 1-cm path length disposable PMMA cuvettes using a Cary Bio 50 UV-Vis spectrophotometer. Transmission electron microscopy (TEM) was conducted on carbon-coated copper grids (Ted Pella, Inc. 01814-F, support films, carbon type-B, 400 mesh copper grid) using a FEI Tecnai (F20) microscope operating at a 200 keV accelerating electron voltage. For the generation of the particle size histograms, 300–1000 individual AuNPs were analyzed.

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C nuclear magnetic resonance (NMR) spectroscopy was performed in D2O

using a Bruker Avance III 500 MHz spectrometer operating at 125 MHz. A Gamry Potentiostat/Galvanostat/ZRA,

Reference

600

workstation

was

used

for

all

electrochemical experiments. The instrument was operated by the Gamry Framework software and data were analyzed by the Gamry EChem Analyst analysis software purchased from Gamry Instruments (Warminster, PA, USA). Temperature studies. In a typical experiment, 1.0 mL of 5.0 mM HAuCl4 was added to 18.0 mL of H2O in a 100-mL glass round-bottom flask. This solution was brought to the desired temperature (placed in a cold room at ~5 °C, left on the bench top at ~20 °C, or heated to a boil in an oil bath) followed by addition of 1.0 mL of a 17.0 mM solution of

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the desired reducing agent under rapid magnetic stirring (750 rpm). The resulting solution was stirred for ~1 min, after which it was allowed to stir for an additional minute at room temperature. Samples were stored in 50-mL centrifuge tubes (Fisher, PPE, catalog number 06-443-20) in a lab drawer to protect them from light. Importantly, reducing agent stock solutions were freshly prepared prior to conducting experiments. Reductant:HAuCl4 (R value). AuNPs were synthesized at room temperature in 50-mL centrifuge tubes according to Table S1. All stability and most R value experiments used 10.0 mM reducing agent stocks and a 5.0 mM HAuCl4 stock. For experiments using R values of 25 and 50, a 50 mM reducing agent stock was used. Since these reductions are extremely rapid, equal volumes of reducing agent and HAuCl4 solution were quickly homogenized by vortex mixing (Thermo Scientific Vortex Maxi Mix II) to lessen the impact of order of addition. All samples were stored at room temperature in 50-mL centrifuge tubes, with protection from ambient light. pH studies. The reaction pH was measured using a glass pH electrode (Vernier, pHBTA) and recorded with LoggerPro computer software to determine the prevalent species of gold chloride present during reduction. The system was programmed to record the pH every 0.50 s. For each reduction, the electrode was allowed to equilibrate for approximately 20 s under magnetic stirring in 10 mL of 0.50 mM HAuCl4 solution at room temperature to ensure a stable baseline. Once equilibrated, 10 mL of reducing agent solution (R values varied) was added and the pH monitored for an additional 30 s (50 s total). The solutions were stirred via magnetic stirring throughout the entire analyses. Electrochemical studies. Solutions of the cyclic oxocarbons (SA, CA, SR) as well as commonly employed reducing agents (AA, SC, NaBH4) were prepared at concentrations of 1.7 mM in 0.1 M aqueous sulfuric acid (H2SO4) as the supporting electrolyte.

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Additionally, solutions of select metal salts (HAuCl4, AgNO3, AuCl, H2PtCl6 and K2PdCl4), common to metal NP synthesis, were prepared by dissolving the salts in 0.1 M H2SO4 for electrochemical evaluation. A conventional 50-mL electrochemical cell was utilized in all experiments. A standard three-electrode configuration was used with glassy carbon or platinum electrodes (GCE and PtE, respectively) as the working electrode (WE) and platinum wire as the auxiliary/counter electrode (CE). Electrode potentials were measured vs. a Ag/AgCl (saturated KCl) reference electrode (RE). Electrodes were purchased from CH Instruments, Inc (Austin, TX). CV curves were recorded between – 200 and +1500 mV at a sweep rate of 100 mV s–1. Prior to use, the REs and CEs were rinsed with water and WEs were polished successively with 1, 0.3, and 0.05 µm alumina powder (CH Instruments) on a wet polishing pad. The WEs were rinsed with water and then ultrasonicated for 1 min each in 5 mL of ethanol then in 5 mL of water. The cleaned electrodes were dried under N2 gas without any further electrode conditioning. All experiments were carried out at room temperature. 4-NP reduction catalysis. The sodium borohydride-assisted reduction of 4-nitrophenol (4-NP) was used as a model reaction to assess the catalytic activities of the as-synthesized AuNPs (within 4 h of synthesis) following a previously reported protocol.43 Specifically, 2.10 mL of 0.20 mM aqueous 4-NP and 0.90 mL of freshly prepared 100 mM aqueous NaBH4 were mixed in a 4-mL PMMA cuvette (1-cm path length), yielding a yellow solution (4-nitrophenolate ion; λmax = 400 nm). To trigger catalysis, 84 µL of 0.25 mM AuNP solution was added to yield a system containing 5 mol% Au relative to 4-NP, followed by gentle (so as to minimize bubble formation) manual inversion to initiate reaction. Final concentrations were as follows: [4-NP] = 0.136 mM; [NaBH4] = 29.1 mM; [Au] = 0.00728 mM. Reaction progress was monitored by periodically measuring

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full UV-Vis spectra from 200–500 nm every 15 s at a scanning speed of 24,000 nm min– 1

. To assess apparent reduction kinetics, the absorbance was monitored at 400 nm using

the Cary Kinetics application, collecting 10 data points per second. Control experiments were performed using classical Turkevich (citrate-capped) gold nanoparticles. The apparent rate constants were calculated from the linear correlation of ln(A0/At) vs. time using the absorbance values at 400 nm. RESULTS AND DISCUSSION UV-Vis spectroscopy and TEM analysis of synthesized AuNPs. Preliminary AuNP synthesis was performed following the Turkevich protocol of rapid injection of a reducing agent into a boiling solution of HAuCl4, substituting a cyclic oxocarbon diacid (i.e., squaric acid, SA; croconic acid, CA; or rhodizonic acid, SR) or ascorbic acid (AA) in place of the usual citrate. These initial studies use a final gold concentration of 0.25 mM with a reducing agent concentration of 0.85 mM to yield a molar ratio (R) of reducing agent to HAuCl4 of 3.4. The detailed experimental procedures followed are provided in the Experimental Section. Upon injection of the reducing agent into a stirring HAuCl4 solution, the mixture transformed from a pale yellow to a deep ruby-red color essentially instantaneously, indicating the rapid formation of AuNPs. It should be noted that this reaction requires several minutes in the conventional Turkevich reaction employing citrate as the reductant. Following injection of the reducing agent, the colloidal AuNP solution was removed from the hotplate and allowed to cool to room temperature while still stirring. Given our interest in developing convenient methods for preparing AuNPs “on-demand” (via simple mixing), parallel experiments were also performed on the benchtop under ambient conditions (~20 °C), as well as in a cold room maintained at ~5 °C, to elucidate the effects of temperature on AuNP synthesis. Although

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the lower reaction temperatures resulted in noticeably slower reaction rates, AuNP formation was still very rapid and was complete in a matter of seconds (Figure S1). Motivated by the fact that the CA and SR stock solutions were strikingly yellow and orange-red, respectively (indeed, Gmelin named croconic acid from the Greek κρόκος meaning “saffron” and the name rhodizonic derives from the Greek ῥοδίζω meaning “to tinge red”), UV-Vis absorption spectra of the AuNP solutions produced and those of the corresponding reducing agent solutions were measured to account for the presence of reducing agent absorption bands which may overlap with the LSPR band of AuNPs. Although CA displays minor absorbance at the blue edge of the AuNP LSPR band, SR shows very high natural absorbance features in the vicinity of the LSPR peak, with a peak absorbance at 490 nm (Figure S2). We selected squaric acid as a representative cyclic oxocarbon dibasic acid for initial investigation. Samples produced using SA as the reducing agent (denoted SAAuNPs) showed strong extinction profiles peaked near 530 nm, attributable to the LSPR of AuNPs (Figure 1A). The as-synthesized AuNPs prepared at room temperature and ~5 °C using SA showed very similar spectral features, while the 100 °C reaction resulted in a broadened and red-shifted LSPR band, indicating that higher temperatures may result in larger or more polydisperse AuNPs. In the 5 and 20 °C samples, a minor shoulder was evident near 630 nm (Figure 1A), a feature we attribute to the presence of a small population of anisotropic AuNPs (e.g., nanorods, plates). To assess the short-term temporal stability of the SA-AuNPs, samples were stored at room temperature in a laboratory drawer (i.e., protected from light) for 5 days, after which UV-Vis spectra were recollected. No obvious spectroscopic changes were noted after a 5-day storage period, pointing to excellent colloidal stability during ambient, short-term storage. The size,

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morphology, and size distribution of the prepared SA-AuNPs were characterized by transmission electron microscopy (TEM) imaging analysis. As shown in Figures 1B and S3, TEM analysis reveals that the samples primarily contain quasi-spherical AuNPs, with a minor contribution from anisotropic or polygonal structures (e.g., hexagonal or triangular plates, nanorods). The 5 and 20 °C samples displayed comparable size distributions (32.1 ± 9.6 and 30.4 ± 8.6 nm, respectively), which is in line with their comparable LSPR band shapes and maxima (526 and 527 nm, respectively). On the other hand, AuNPs made at 100 °C (29.2 ± 9.7 nm; Figure S3I) were slightly smaller than those prepared at lower temperatures, although the broadened LSPR band is consistent with a slightly more polydisperse sample. Given that the reagent molar ratio (R) of reducing agent to gold is well known to affect the size and dispersity of synthesized AuNPs,29 SA reduction to generate AuNPs from HAuCl4 was investigated for a wide range of R values (i.e., 0.10, 0.25, 0.50, 0.75, 1.0, 2.0, 3.0, 5.0, 10, 25, and 50). We note that, unless specifically stated otherwise, these and all subsequent experiments were carried out in aqueous solutions at room temperature (~20 °C). The resulting SA-AuNP solutions were analyzed using UV-Vis spectroscopy and TEM to determine the overall effect of the R value on the LSPR, as well as AuNP size, shape, and size distribution (see Figures 2, S4, and S5). Additionally, the SA-AuNP samples were spectroscopically monitored at various intervals over the course of the first week (Figure 2B,C) and after seven months of storage (Figure S4) to assess their colloidal stability. The AuNP extinction profiles in Figure 2A and the TEM micrographs in Figure S5 were acquired using one-week-old samples since, by this time, SA-AuNP samples did not vary appreciably with further aging, even for low R values. As shown in Figure S4 (see panels A and B), the as-prepared SA-AuNP samples show a

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striking dependence on R. As R increases from 0.10 to 2.0, the LSPR profile becomes more intense, shifts to the blue, and ultimately narrows, indicating the formation of welldefined pseudo-spherical AuNPs for R ≥ 3.0. The LSPR spectral width was analyzed by calculating the full-width at three-quarters of the maximum peak height ([email protected] max). Whereas SA-AuNPs prepared at R values below 1.0 present significant LSPR spectral evolution within the first 24 h, the LSPR peaks (Figure 2B) and [email protected] max values (Figure 2C) for SA-AuNPs prepared for R = 2–10 show very little change over the course of seven months of storage (Figure S4), indicating remarkable colloidal stability. Interestingly, very high reagent ratios (R = 25, 50) are also associated with decreased extinction intensities over time, suggesting a compromised colloidal stability. This behavior likely stems from decreased electrostatic stabilization of SA-AuNPs due to the large excess of SA in solution (i.e., higher ionic strength), resulting in some sedimentation. To further examine the effect of the molar ratio of SA to HAuCl4 on AuNP size and morphology, TEM analysis was conducted on one-week-old samples (Figure S5). For all R values investigated, the samples predominately comprise quasi-spherical AuNPs, with a small content of trigonal plates. In general, as R increases, the average particle size decreases and the size distribution narrows. We note that AuNPs as large as 100–200 nm can be observed for one-week-old SA-AuNP samples for both low (R = 0.25) and high (R = 25.0) molar ratios of SA to HAuCl4, accounting for the increased extinction values from 600–800 nm. Although a detailed discussion of the pathways to obtain AuNPs by oxocarbon35 reduction is beyond the scope of this paper, we propose a mechanism for the overall oxidation processes using SA as a representative cyclic oxocarbon dibasic acid. Initially,

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Au(III) is reduced to Au(I), and this process involves two steps (Scheme 1B): (i) a ligand exchange with squarate anion to form an intermediate complex, and (ii) a dechlorination step involving the two-electron reduction of Au(III). In the first step (i), the enolate oxygen of squarate anion attacks the Au(III) center via an SN2 type nucleophilic substitution reaction. Following electron rearrangement, reduction of Au(III) to Au(I) occurs, concomitant with concerted loss of a second chloride ion to yield cyclobutene1,2,3,4-tetraone (species I) as the oxidation product of squarate. At this point, the resulting Au(I) species can generate Au(0) by disproportionation reaction (shown in pathway iii) or Au(I) species can be potentially reduced by another available squarate molecule. The isolated oxidation products of squarate are likely to be the hydrolytic byproducts of the tetraketone form (I in Scheme 1) which itself is not isolable. For example, the octahydroxycyclobutane (II in Scheme 1) may be viewed as the tetrahydrate form of I. Indirect evidence of the presence of various hydroxylated derivatives of squarate was obtained by

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C nuclear magnetic resonance (NMR) spectroscopy,

conducted on samples with R values of 0.75 and 1.5 (Figure S6), reagent ratios specifically chosen to ensure essentially complete consumption of SA with high abundance of oxidation products. The

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C NMR spectra indicate that the oxidized

products of SA consist of a mixture of carbon species bearing C=O and C-OH moieties (125 MHz, D2O: δ 174.4 (C=O), 173.0 (C=O), 172.5 (C=O), 172.0 (C=O), 124.6 (C– OH), 96.3 (C–OH), and 94.2 (C–OH) ppm), confirming varying degrees of hydrolysis of the oxidized product cyclobutane-1,2,3,4-tetraone (compound I in Scheme 1). We must point out that this overall mechanism, while not definitive, is based on the experimental evidence available and the existing literature.44

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Similar to the previous discussion of SA-AuNPs, temperature- and R valuedependent studies were also conducted for AuNPs made using AA, CA, and SR. Contrary to the results of temperature-dependent SA studies, increasing the reaction temperature when employing AA (R = 3.4) as a reducing/capping agent in AuNP synthesis had no obvious effect on the LSPR profile for AA-AuNPs (Figure 3A). Consistent with this observation, TEM analysis (Figures 3B and S7) reveals statistically-equivalent mean nanoparticle sizes for AA-AuNP synthesis at 5 and 20 °C (33.1 ± 10.6 and 33.1 ± 9.3, respectively). Meanwhile, TEM analysis reveals that AA-AuNP synthesis performed at 100 °C yields a statistically-significant decrease in the average particle size to 24.8 ± 9.1 nm. Measured UV-Vis extinction profiles and the corresponding normalized spectra for as-prepared R-dependent AA-AuNPs and samples aged one day, one week, and seven months are provided in Figure S8. The corresponding LSPR shifts and peak width analysis conducted over the first week are summarized in panels B and C of Figure 4. For low molar ratios of AA to HAuCl4 (R = 0.10–0.50) the LSPR peak and the [email protected] max values for AA-AuNPs shift considerably over a one-week period following their preparation. We note that no discernible plasmon band is observed for R = 0.10. For R of 0.75 and higher, the LSPR peak position and spectral width remain nominally the same over the course of aging for one week. Results of TEM particle analysis for one-week-old AA-AuNP samples made at R values of 0.5, 2.0, and 10.0 are in accord with our UV-Vis experiments (Figure S9). That is, smaller and more monodisperse AA-AuNPs are generated when employing larger R values. Specifically, when proceeding from an AA to HAuCl4 molar ratio of 0.5 to 2.0 to 10.0, the mean AA-AuNP size (and the associated

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distribution width) decreases from 43.8 ± 27.5 nm to 32.2 ± 9.1 nm to 30.3 ± 6.0 nm, respectively. We next studied the synthesis of AuNPs using CA and SR as reducing agents at various temperatures (5, 20, and 100 °C), the results of which are summarized in Figure 5 and Figures S10–S13. For CA-AuNPs prepared at a reaction temperature of 20 °C, the LSPR profile was initially ill-defined and broad, extending from 500 to 800 nm (Figure 5A). Instead, CA-AuNPs made at 100 °C present a narrow profile peaked near 520 nm immediately upon synthesis, typical of well-defined spherical AuNPs. Meanwhile, for SR-AuNPs, UV-Vis spectra gave no discernible trend except an apparent shoulder near 540 nm when using SR as a reductant at 100 °C (Figure 5B). Unlike the other reducing agents, we note that SR presents inherent absorbance which completely overlaps with typical AuNP LSPR features. This excessive background renders UV-Vis analysis futile, making TEM analysis vital for corroborating successful SR-AuNP synthesis. Representative TEM images of CA- and SR-AuNPs generated at room temperature verify that both reagents function as dual reducing and capping agents in AuNP formation (Figure 5C and D). Additional images and associated particle size histograms are provided in Figures S12 and S13. Interestingly, in contrast to results with SA or AA, bimodal AuNP size distributions were obtained when employing either CA or SR as the reducing agent. Expressly, CA-AuNPs contain discrete 3.5 ± 1.9 nm and 29.9 ± 6.3 nm particles (Figure S12) and SR-AuNPs comprise 5.1 ± 1.0 nm and 29.7 ± 7.6 nm populations (Figure S13). The formation of these bimodally-distributed AuNPs is tentatively

attributed

to

a

secondary

nucleation

brought

about

by

oxidation/degradation/hydrolysis products present in solution during AuNP synthesis. A similar conclusion was drawn by Xia et al.18 during exploration of the Turkevich method,

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wherein it was determined that a thermal degradation by-product of SC (i.e., acetone) exerted a profound effect on AuNP synthesis in the presence of Ag+. This hypothesis is especially pertinent for SR, as rhodizonate is known to readily hydrolyze in water in the presence of oxygen.45 Figures 6 and 7 summarize UV-Vis analysis results for AuNPs made using CA and SR at varying molar ratios of reducing agent to HAuCl4. In both cases, as the R value increases, the LSPR band is seen to blue shift and narrow, similar to observations for SA and AA. Stability studies of CA-AuNPs and SR-AuNPs indicate that for R values of 2.0 or higher, the monitored LSPR peak positions and [email protected] max values remain steady during aging for one week. In contrast, use of lower R values is linked with sweeping temporal changes in the optical features of the AuNP samples. Monitoring of pH during reaction. To better understand the reaction conditions within these systems, pH was measured during reaction (see Experimental Section for details). Briefly, a pH electrode was allowed to equilibrate in 10 mL of a 0.50 mM HAuCl4 solution under stirring. After equilibration and establishing a stable 20 s baseline, 10 mL of a stock solution of SA or AA corresponding to a desired, final (diluted) R value (e.g., 0.50, 2.0, 10) was injected, with stirring maintained throughout the analysis (Figure S16A). In the cases of CA and SR, R was set at 3.4 (Figure S16B). The initial pH of each run (~3.0) corresponds to the pH of 0.50 mM HAuCl4. Thus, as reported by Goia et al., the primary gold species initially present should be AuCl4–,46 which is the most easily reduced species of gold chloride according to Fry et al.47 Addition of a reducing agent to the stirring HAuCl4 creates pH fluctuations which rapidly stabilize in a matter of seconds, indicating a rapid equilibrium. In general, higher R values result in a lower final pH for the final AuNP solutions, an intuitive result arising from an increased amount of the

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diacid species. For both SA and AA, an R value of 0.50 is associated with a small jump in pH from 3.0 to 3.1. Assuming near-quantitative consumption of the reducing agent for R = 0.5, it is apparent that the oxidation byproduct of SA has lower acidity than SA, consistent with our earlier-proposed mechanism in that neither the oxidized form of SA (cyclobutane-1,2,3,4-tetraone; I in Scheme 1) nor its hydrolysis products should present obvious acidity. For R values of 2.0 and 10 using AA (and for R = 2.0 with SA), the final pH stabilized at ~2.9. Using SA with R = 10, however, the pH precipitously drops to ~2.6 within a few seconds and continues to decrease slightly over the remainder of the monitoring window. This large decrement is attributed to the strong acidity of the excess SA in solution (pK1 ≈ 1.5);48 given that pK2 is approximately 3.4,48 at a pH of 2.6, the majority of SA molecules are expected to be mono-protonated, contributing a large inventory of H+ to the reaction medium. Indeed, the pH of an SA solution at a concentration equivalent to R = 10 is 2.68. On the other hand, AA is comparably much less acidic (pK1 = 4.17 and pK2 = 11.57),49 suggesting that any excess AA will remain fully protonated, having only a marginal influence on the final pH. Contrary to the SA and AA cases, reaction media involving both CA and SR (for R = 3.4) display a substantial rise in pH from ~3.0 to ca. 3.4 and 4.05, respectively. Interestingly, whereas the pH in the CA-AuNP system stabilizes within seconds, the pH of the SR-AuNP system steadily increases over the entire recording period, approaching pH ~4.1 after a 30 s period. The increase in pH in these two cases is not surprising, given that the dibasic (disodium) forms of the oxocarbons were used. In the case of SR (pK1 = 4.25 and pK2 = 4.72),45 excess SR free base will abstract protons from HAuCl4 and cause the pH to rise. At this point, it is not entirely clear why the kinetics were overtly slower for SR

17



compared to CA, although it may be linked to the known hydrolytic instability of the rhodizonate anion. Electrochemical analysis of the reducing agents. To better understand the behavior of SA, CA, and SR as chemical reducing agents and the reaction mechanisms at work within these systems, their electrochemical behavior was studied using cyclic voltammetry (CV) and differential pulse voltammetry (DPV). To date, few studies have been conducted on the electrochemical activity of SA,50-53 CA,51, 54 and SR.55 The majority of these works were centered on their oxidation at polycrystalline and single-crystal PtEs in both acidic aqueous solutions and aprotic media where two-electron electrochemical oxidations of SA and CA were proposed.50-51 Two oxidation peaks for SA and CA were observed on PtEs: (i) a pre-wave as a consequence of adsorption of oxidative products whose nature depends on the material being oxidized, and (ii) a main oxidation peak. Rodes et al. confirmed, by FTIR analysis, that the adsorption of SA leads to the formation of adsorbed CO on platinum surfaces and is responsible for the pre-wave formation.53 Further electrochemical characterization of SA, performed by Sant’Ana et al.50 on gold electrodes, reported a single anodic peak due to the irreversible oxidation of SA and a notable absence of by-product sorption on the electrode surface as observed at a PtE. Fabre et al. demonstrated similar results for the electrochemical and photooxidation of CA to oxalic acid on a glassy carbon electrode (GCE).54 These works further established a pH dependence of the oxidation of SA and CA in aqueous solutions. The electrochemical oxidation of 0.85 and 1.7 mM SA and CA was investigated at a bare GCE in acidic aqueous media (pH 2.0) between –200 and +1100 mV at a sweep rate of 100 mV s–1, and the resulting cyclic voltammograms are presented in Figures S17 and S18, respectively. It is worth noting that since SR hydrolyzes readily in aqueous

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solution, it produces inconsistent results; thus, any SR data provided herein are used solely for qualitative comparisons and should be taken as such (for more in-depth electrochemical analysis of SR, the reader is referred to work by Carbó et al.55). The oxidation process of SA at GCE (Figure S17) appears irreversible with one prominent, well-resolved anodic peak at 795.2 mV, due to the oxidation of SA, and the absence of a discernible cathodic peak in the reverse scan. A similar result was reported in the work by Sant’Ana et al.50 The absence of a cathodic peak may be attributed to the oxidation product of SA being used in a chemical reaction (e.g., hydrolysis as per the proposed mechanism of Scheme 1) and swept away from the electrode surface, making it unavailable for reduction at the relatively low scan rates used.56 The observed oxidation peak is directly related to the concentration of SA in solution, as demonstrated by an increase in peak current with increasing SA concentration without any significant peakshift. The electrochemical oxidation of SA occurs via a two-electron process: SA → SA•– + e– SA•– → SA + e– which is in good agreement with experimental results presented within the paper and the NMR data presented earlier (Figure S6). Oxidation peaks attributed to the unstable SA•– radical anion are not observed, as reported in aprotic media,52 due to fast oxidation of the unstable species in aqueous media. The reduction of the oxidized SA was further investigated by means of DPV (Figure S17 , inset). Successive oxidation and reduction scans were conducted, and the resulting plots were overlaid, which showed one large anodic peak (764.3 mV vs. Ag/AgCl) attributed to the oxidation of SA and a considerably smaller cathodic peak (749.1 mV vs. Ag/AgCl) due to the reduction of the oxidized SA previously adsorbed onto the electrode surface.

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No significant peak shifts can be seen between the cyclic and differential pulse voltammograms, while the reduction peak is clearly observed in DPV. The scan rate dependence of the electrochemical oxidation of SA at the GCE surface was further investigated between 10–100 mV s–1, as shown in Figure 8A. An increase in peak current with increasing scan rate is exhibited along with a small shift to more positive oxidation potentials at higher scan speeds. Here, electron transfer rates are insufficient to maintain Nernstian equilibrium and a distinct deviation from reversibility is observed, where mass transport (diffusional processes) dominates over electron transfer, and oxidation of SA becomes more difficult. The direct proportionality, illustrated by the linear increase in peak current with the square root of the sweep rate (Figure 8A inset) verifies the diffusion control of the main SA oxidation wave (i.e., the process occurs via spontaneous transfer of electroactive species from regions of high concentration to low concentration). The diffusion coefficients of irreversible and quasi-reversible systems can be estimated from CV by using the Delahay equation (eq 1):57-59

= 0.4961 FAC∗

!"

#

$⁄

(1)

where ipa is the anodic peak current, v is the scan rate, n is the number of electrons participating in the reaction, F is Faraday’s constant, A is the surface area of the electrode, C∗ is the ion concentration, α is the transfer coefficient, DE is the diffusion coefficient, R is the universal gas constant, and T is the absolute temperature. The diffusion coefficient of SA is estimated to be 1.11 × 10–5 cm2 s–1 when the transfer coefficient is assumed to be 0.5. The calculated diffusion coefficient is consistent with reported literature values51, 58 reinforcing the results obtained here. For example, Sazou et

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al.51 reported a value of 1.30 × 10–5 cm2 s–1 at a PtE which matches well with the value determined in this work. Figure S18 shows the voltammetric profile corresponding to the electrochemical oxidation of CA at a bare GCE, between –200 and +1100 mV vs. a Ag/AgCl reference electrode. The recorded cyclic voltammograms comprise two distinct redox couples attributed to the oxidation and reduction processes of CA at a GCE surface. Two discernible peaks are observed in the forward anodic scan at 548.3 and 741.2 mV, respectively, illustrating the oxidation of CA in two discrete steps. The oxidation commences as a result of the deposition of oxide products at the electrode surface, a behavior not previously observed at a GCE surface, followed by the direct oxidation of CA at more positive potentials. Two separate reduction peaks are also displayed in the cathodic scan: one at 329.3 mV as a result of the reduction of the deposited oxide reaction products on the electrode surface as well as a weaker peak at 651.2 mV due to reduction of the oxidized CA. The appearance of the adsorption pre-wave at potentials more negative than the main oxidation wave demonstrates a prevalence for the strong adsorption of the reaction products. The nature of the deposited oxide product has not yet been studied, but it appears in the same region as CO oxidation, as previously reported.53 The appearance of an adsorption peak at potentials more negative than the main diffusion peak indicates strong adsorption of the reaction product. Sazou et al.51 and Fabre et al.54 reported possible mechanisms for the oxidation of CA which support our experimental findings: CA → CA + 2e A comparison of the oxidation profile of CA with SA yielded similar primary oxidation features in the 700–800 mV range. The reduction peak appears to be more

21



pronounced for CA, suggesting that the reduction of CA occurs more readily than its four-membered ring counterpart SA, while the appearance of a second redox couple attributed to adsorbed CA oxide products suggests that the nature of the two oxide products differ in SA and CA. An overall lowering in current densities is also observed for CA relative to SA. Investigation of the redox processes for CA by DPV produced comparable results to those found in CV studies: a distinct oxidation peak and two wellresolved reduction peaks. Figure 8B investigates the scan rate dependence of CA oxidation between 10–100 mV s–1 at a GCE. The peak current for the main oxidation wave shows a linear dependence on the square root of the scan rate, confirming diffusional control. The shape and size of the pre-peak at 548.3 mV improves with scan rate being measurable only at higher scan rates. The adsorption control of the pre-peak is noted by the linear relationship between peak current and scan rate. Similar findings have been reported in the literature for CA oxidation.51 The diffusion coefficient estimated from the CV experiments using eq 1 was 0.42 x 10–5 cm2 s–1 (assuming a transfer coefficient of 0.5), which is comparable to the value reported by Sazou et al. (0.55 x 10–5 cm2 s–1).51 A direct comparison of the DPV of SA and CA within the potential window of 0.0 to +1200 mV is provided in Figure S19. The redox profiles of the main oxidation waves of both SA and CA show a discernible, large oxidation peak between 600–800 mV along with a significantly smaller corresponding reduction peak, in the anodic and cathodic scans, respectively. A 62-mV shift to more negative potentials for CA compared to SA, illustrates that the five-membered CA ring is more easily oxidized than the fourmembered SA. The voltammetric profiles differ in the region between 200–400 mV, where CA shows an additional redox peak over SA. The adsorption of oxidized products

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on the glassy carbon surface is evident in CA oxidation, while noticeably absent for SA. This further confirms the strong adsorption of the reaction product produced in the CA oxidation as previously discussed. The oxidation of the adsorbed oxide is distinctly smaller and less well-resolved than those found in CV due to the slow scan speeds utilized. The peak-to-peak separation of all redox couples appear below 50 mV. The bulk of the reported oxocarbon electrochemistry performed to date has been conducted on PtE surfaces. To accurately confirm the validity of the results produced, the oxidation profiles of SA and CA at GCE surfaces were compared to PtE surfaces. The voltammograms recorded vs. a Ag/AgCl reference electrode for GCE and PtE surfaces are compared in Figure S20. The CV of SA at a platinum surface shows two well resolved anodic and two cathodic peaks within the potential window investigated with no change in oxidation potential of the diffusion process compared with glassy carbon. Oxidized products formed due to the oxidation of SA show strong adsorption on the platinum surfaces over GCEs. Similar features were reported in the work by Sazou et al.51 The oxidation profiles of CA on GCE and PtE surfaces are significantly more similar to one another; two redox couples appear in both cases without significant change in redox potentials. For both SA and CA, the platinum surfaces show lower peak current densities relative to glass carbon. The standard reduction potential (Eo) of a chemical species is a measure of its tendency to acquire electrons and be reduced under specific, standard conditions and is measured relative to the standard hydrogen electrode (SHE) reference electrode. It is intrinsic to a particular chemical species, predicting the directionality of a redox reaction when comparing two species of interest.59 Cyclic voltammetry was employed to study the competancy of SA and CA as reducing agents for the reduction of various metal salts in

23



acidic aqueous media at 20 °C. Toward this, oxidation potentials of 1.7 mM SA and CA were investigated by CV at bare GCEs in aqueous 0.1 M H2SO4 (pH 2.0) and compared with those of 0.5 mM HAuCl4 and AgNO3 solutions between 0.0 and +1400 mV (Figure S21). For consistency, solution concentrations were selected to be similar to those employed for AuNP synthesis. SA and CA show formal potentials of the main diffusion processes at 795 and 741 mV, respectively, under ambient conditions, as previously discussed. It may be expected that metal salts with reduction potentials more positive than SA or CA will undergo reduction in their presence, while those with reduction potentials more negative will not. A similar study was reported by Newman et al. for predicting the ability of organic amines to function as reducing agents in the synthesis of AuNPs.60 The redox profile of HAuCl4 shows a single oxidation peak, along with two separate reduction peaks in the reverse scan. The oxidation wave shows a fast, single-step oxidation from Au0 to Au3+, while reduction processes occur at a slower rate with two separate reduction peaks for Au3+ to Au+ and Au+ to Au0 conversions. The formal potential of Au3+ appears at 1001 mV. Silver nitrate (AgNO3) was also studied in the same potential window. A single redox couple is observed; a sharp anodic peak for Ag0 to Ag+ oxidation and a significantly smaller cathodic peak attributed to the reduction of Ag+ to Ag0. The formal potential of Ag+ was found to be 342 mV. A direct comparison of the formal potentials shows that the SA oxidation potential lies intermediate between those of Ag+ and Au3+, respectively, accounting for its ability to spontaneously reduce Au3+ but not Ag+ under ambient conditions, as observed in the course of these studies. Further investigation was performed by studying the redox profiles of reducing agents and metal salts commonly used in nanoparticle syntheses at GCEs. As shown in Figure S22, SC, AA, and sodium borohydride (NaBH4) all show single oxidation peaks in

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the anodic scan at 1670, 544, and 488 mV respectively, without any reduction peaks observed in the reverse cathodic scan. Similar findings for citric acid,61 AA,62-63 and sodium borohydride64 have previously been reported. The half-wave potential for AA, however, may fluctuate between 200 and 600 mV vs. Ag/AgCl reference electrode depending on the electrode cleaning procedure. These findings support those reported in the literature. Namely, AA and NaBH4, whose half-wave potentials lie at potentials less positive than silver and gold, have previously been reported to readily reduce Ag+ and Au3+.65-66 Sodium citrate reduction of silver and gold, however, does not occur spontaneously at ambient temperature since the key step in the reduction of Au3+ in the Turkevich method is the thermal degradation of citrate which requires temperatures ≥80 °C.67 Figure 9 summarizes these electrochemical findings, showing the reduction potentials of SA, AA, CA, and SR alongside metal salts relevant to nanoparticle synthesis. These results accord with our experimental findings and suggest that the current cyclic oxocarbon reducing agents are unlikely to spontaneously reduce Ag+, Pd2+, and Pt4+. Catalytic activity of synthesized AuNPs towards 4-nitrophenol reduction. In a final set of experiments, the catalytic potential of AA-, SA-, CA-, and SR-AuNPs (synthesized at an R value of 3.4) were evaluated using borohydride-assisted reduction of 4nitrophenol (4-NP) to 4-aminophenol (4-AP) as a model reaction.68-69 The resulting catalytic rates were compared to SC-AuNPs (i.e., Turkevich method, R = 3.4) as a benchmark. The reaction kinetics for the catalytic conversion were monitored by following the loss in absorbance at 400 nm (i.e., 4-nitrophenolate) using UV-Vis spectroscopy (Figure S23). The apparent rate constant (kapp) for each AuNP catalyst was

25



estimated from the slope of the linear correlation of ln(A0/At) vs. time, where A0 is the initial absorbance and At is the time-dependent absorbance. In decreasing order of activity, the kapp values for the as-synthesized SR-, SC-, AA-, SA-, and CA-AuNP catalysts were 1.38 (±0.03) × 10–2 s–1, 7.07 (±0.29) × 10–3 s–1, 3.55 (±0.07) × 10–3 s–1, 2.04 (±0.33) × 10–3 s–1, and 1.87 (±0.43) × 10–3 s–1, respectively (Figure 10). Catalytically, AA-AuNPs and SA-AuNPs show catalytic rates about 2- and 3.5-fold lower than the classical Turkevich-derived SC-AuNPs, respectively. We attribute this reduced activity primarily to the fact that the Turkevich reaction produces much smaller AuNPs. Consider, for example, the fact that when comparing 30.4 nm SA-AuNPs with 15.8 nm SC-AuNPs, the latter presents a 1.9-fold higher surface area (catalytic area) for an identical catalyst loading (mol% Au), based on simple geometric considerations (assuming spherical AuNPs), fully accounting for the halved catalytic rate in going from SC-AuNPs to SA-AuNPs. In comparison, the SR-AuNPs performed particularly well for 4-NP reduction, displaying an apparent catalytic rate double that of the SC-AuNPs. This drastic improvement in kapp can be explained on similar grounds. Indeed, we attribute the catalytic improvement to the presence of tiny (~5 nm) AuNPs within the bimodal population of SR-AuNPs. The CA-AuNPs also contain ultrasmall AuNPs (3.5 nm) and might be expected to perform as even better catalysts. However, the activity of CAAuNPs is probably compromised by nanoparticle agglomeration and shielding by an indefinite organic matrix (TEM images in Figure S12 provide evidence for this), limiting the AuNP surface availability for catalytic reaction. Control experiments were performed to rule out contribution to the catalytic activity from the reducing agents themselves. Aqueous solutions of SA, AA, CA, and SR were thus tested directly for BH4-assisted 4-NP reduction over a 30-min monitoring

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period (Figure S24). In all cases, the 400 nm peak for the 4-nitrophenolate remained unchanged, verifying that any contribution to the catalytic reduction of 4-NP from the agents themselves is negligible. We note that analysis of the CA sample does reveal timedependent spectral features in the 260–395 nm region, with decreased absorbance over time in the 325–395 nm window and the genesis of a new peak at ~285 nm. This result reveals the aqueous-solution decomposition of the croconate species. Aging studies reveal that six-month-old SA-, AA-, and CA-AuNPs samples maintain the catalytic activities of the original materials (Figure S25). In contrast, kapp for 4-NP reduction was reduced nearly 2.5-fold to 5.62 (±0.85) × 10–3 s–1 after six months of storage (in glass vials reserved in a laboratory drawer) for SR-AuNPs. Given the aforementioned difficulty in characterizing SR-AuNPs using UV-Vis due to the obscuring nature of the SR absorbance, TEM imaging was conducted on a six-month-old SR-AuNP sample. The results of TEM imaging reveal that after aging, what was once a bimodal distribution of nominally 5 and 30 nm AuNPs had now merged into a singular distribution having a mean particle size of 6.4 ± 1.7 nm (Figure S14), concomitant with a high degree of nanoparticle aggregation. The ripening in AuNP size and the colloid association account for the significant decrease in catalytic activity, although it is worth noting that even the aged SR-AuNP sample displays a catalytic activity that is 80% as high as that seen for fresh SC-AuNPs prepared using the classical Turkevich method. Lastly, we conducted a series of tests on one-year-old samples to assess the retention of catalytic activity after very long-term storage, as well as to evaluate the potential effects of surface ligand exchange on catalytic performance. We note that the AuNP dispersions used herein are colloidally stable after one year of ambient storage (laboratory drawer). In these experiments, 5 mL of SA-AuNP and SR-AuNP solutions (R

27



= 3.4, [Au = 0.25 mM) were separately dialyzed against 1.0 L of 0.85 mM aqueous sodium citrate for 2 days at 4 °C under slow stirring using 1-kDa molecular weight cutoff dialysis tubing. The LSPR bands of the dialyzed and parent solutions were measured and their catalytic activities determined for 4-NP reduction (Figure S26). The citrateexchanged SA-AuNPs and SR-AuNPs are denoted as SA-SC-AuNPs and SR-SC-AuNPs, respectively. SA-SC-AuNPs showed slight peak broadening, a 2 nm red shift, and an increase in intensity in their LSPR (~5%) when compared to their parent SA-AuNPs, indicating minor changes in colloid association as a result of capping ligand exchange.70 Catalytic profiles reveal kapp values of 2.25 (±0.21) × 10–3 and 2.38 (±0.21) × 10–3 s–1 for one-year-old SA-AuNPs and SA-SC-AuNPs, respectively, comparable to the apparent rate measured using freshly made SA-AuNPs (2.04 (±0.33) × 10–3 s–1). Conversely, oneyear-old SR-AuNPs show a greater change following ligand exchange. The SR-SCAuNPs experience a 5 nm red shift, a ~3% decrease in intensity, and slightly more peak broadening in their LSPR when compared to their parent particles. Rates of 4-NP reduction were mildly enhanced, with kapp values of 7.53 (±0.29) × 10–3 and 8.76 (±0.12) × 10–3 s–1 for SR-AuNPs and SR-SC-AuNPs, respectively. While these values are similar, both solutions present a catalytic rate less than the rate determined for fresh SR-AuNPs (1.38 (±0.03) × 10–2 s–1), likely caused by coarsening of the AuNPs upon aging. It can be concluded that ligand exchange exerts relatively little effect on the catalytic activities of the AuNPs, with the predominant factor being the size, morphology, and dispersity of the AuNPs initially formed. CONCLUSIONS In conclusion, the room temperature reduction of AuCl4– using various cyclic oxocarbon diacids (squaric acid, SA; croconic acid, CA; or rhodizonic acid, SR) as well

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as ascorbic acid (AA) was shown to result in colloidally-stable and catalytically-active AuNPs. In general, these reducing agents produced larger (~2-fold) particles than those obtained when employing the benchmark Turkevich method, with a slight increase in polydispersity as well (15.8 ± 3.7 nm for SC compared to 30.4 ± 8.6, 33.1 ± 9.3, 29.9 ± 6.3, and 29.7 ± 7.6 nm for SA, AA, CA, and SR, respectively). The reaction temperature and the ratio of reducing agent to gold were preliminarily studied and both were found to impact particle dispersity and stability. Interestingly, the croconate and rhodizonate reagents produced AuNPs having a bimodal size distribution, with a significant portion of TEM-counted nanoparticles having ultrasmall sizes of ca. 3.5 nm and 5 nm, respectively. Electrochemical cyclic voltammetry measurements were conducted to gain a fundamental understanding of gold reduction by these oxocarbon agents and were instructive for making predictions on which metal salts might be compatible with the use of these reducing agents in future nanosynthetic work. Finally, the catalytic activities of the prepared AuNPs were evaluated using borohydride-assisted reduction of 4-nitrophenol as a model catalytic reaction, showing rates of 2.0 × 10–3 s–1, 3.6 × 10–3 s–1, 1.9 × 10–3 s–1, and 13.8 × 10–3 s–1 for SA, AA, CA, and SR, respectively. Notably, SR-AuNPs presented a catalytic efficiency that was essentially twice that for citrate-capped AuNPs made from the classical Turkevich method at the same gold content, a result stemming from the presence of ultrasmall (~5 nm) AuNPs in the SR-AuNP sample. Remarkably, the synthesized particles are colloidally very stable, with SA-, AA-, and CA-AuNPs retaining their full catalytic proficiency even after storage in a laboratory drawer for one full year. Overall, these findings provide a path forward in nanoscale manufacturing, offering a means to generate fairly uniform AuNPs by a reliable, convenient, and scalable mixing protocol that, unlike the classical Turkevich method, requires no heating step. It is

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inferred through the present results that squaric acid, for example, provides an interesting alternative to the use of citrate for very rapid, ambient-temperature “on-the-fly” AuNP production, providing incentive to learn how to better control nucleation and growth to achieve more uniform and smaller colloids. Simple manual injection of HAuCl4 solution into an equal volume of reducing agent was followed in the current work. In this regard, we note that particle size dispersion in the current work is limited by the fact that macroscale batch reaction (in which mixing is turbulent and not well-defined) results in concentration gradients, leading to polydispersity in nanoparticle synthesis. Particularly promising in this regard is the use of microfluidic droplet flows to eliminate concentration dispersion for high-fidelity metal nanoparticle synthesis, as illustrated in work by Malmstadt, Brutchey and co-workers.71 Future efforts will seek to improve upon the quality of AuNPs produced using cyclic oxocarbon reducing agents by using microfluidic technology to achieve efficient mixing to promote a more homogeneous reaction environment.

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SUPPORTING INFORMATION The Supporting Information is available free of charge on the ACS Publications website at DOI: Temperature- and time-dependent extinction spectra and TEM images and histograms for SA-, CA-, SR-, and AA-AuNPs synthesized at different R values; additional electrochemical analyses and catalytic comparisons (PDF)

ACKNOWLEDGMENTS The authors would like to acknowledge the partnership between the University of Missouri System and the University of the Western Cape (UWC), created in 1986, for providing funds allowing K.P. to visit the University of Missouri-Columbia and lead the electrochemistry portion of this research.

■ AUTHOR INFORMATION Corresponding Author *E-mail: [email protected]. ORCID Gary Baker: 0000-0002-3052-7730 Emmanuel Iwuoha: 0000-0001-6102-0433 Jeremy Essner: 0000-0002-2500-7968 Nathaniel Larm: 0000-0002-3369-4980 Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes The authors declare no competing financial interest.

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26. Kharissova, O. V.; Dias, H. V. R.; Kharisov, B. I.; Pérez, B. O.; Pérez, V. M. J. The Greener Synthesis of Nanoparticles. Trends Biotechnol. 2013, 31, 240-248. 27. Frens, G. Controlled Nucleation for the Regulation of the Particle Size in Monodisperse Gold Suspensions. Nature Physical Science 1973, 241, 20. 28. Poonthiyil, V.; Golovko, V. B.; Fairbanks, A. J. Control of Gold Nanostructure Morphology by Variation of Temperature and Reagent Ratios in the Turkevich Reaction. Aust. J. Chem. 2015, 68, 858-862. 29. Kimling, J.; Maier, M.; Okenve, B.; Kotaidis, V.; Ballot, H.; Plech, A. Turkevich Method for Gold Nanoparticle Synthesis Revisited. J. Phys. Chem. B 2006, 110, 1570015707. 30. Kumar, S.; Gandhi, K. S.; Kumar, R. Modeling of Formation of Gold Nanoparticles by Citrate Method†. Ind. Eng. Chem. Res. 2007, 46, 3128-3136. 31. Ojea-Jiménez, I.; Bastús, N. G.; Puntes, V. Influence of the Sequence of the Reagents Addition in the Citrate-Mediated Synthesis of Gold Nanoparticles. J. Phys. Chem. C 2011, 115, 15752-15757. 32. Ojea-Jiménez, I.; Romero, F. M.; Bastús, N. G.; Puntes, V. Small Gold Nanoparticles Synthesized with Sodium Citrate and Heavy Water: Insights Into the Reaction Mechanism. J. Phys. Chem. C 2010, 114, 1800-1804. 33. Schulz, F.; Homolka, T.; Bastús, N. G.; Puntes, V.; Weller, H.; Vossmeyer, T. Little Adjustments Significantly Improve the Turkevich Synthesis of Gold Nanoparticles. Langmuir 2014, 30, 10779-10784. 34. O'Neil, M. J. The Merck Index - an Encyclopedia of Chemicals, Drugs, and Bilogicals; Merck and Co., Inc.: Whitehouse Station, NJ, 2006. 35.

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36. D'Souza, S. L.; Pati, R. K.; Kailasa, S. K. Ascorbic Acid Functionalized Gold Nanoparticles as a Probe for Colorimetric and Visual Read-out Determination of Dichlorvos in Environmental Samples. Anal. Methods 2014, 6, 9007-9014. 37. Sun, K.; Qiu, J.; Liu, J.; Miao, Y. Preparation and Characterization of Gold Nanoparticles Using Ascorbic Acid as Reducing Agent in Reverse Micelles. J. Mater. Sci. 2009, 44, 754-758. 38. Jana, N. R.; Gearheart, L.; Murphy, C. J. Seed-Mediated Growth Approach for Shape-Controlled Synthesis of Spheroidal and Rod-Like Gold Nanoparticles Using a Surfactant Template. Adv. Mater. 2001, 13, 1389-1393.

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39. Khan, Z.; Singh, T.; Hussain, J. I.; Hashmi, A. A. Au(Iii)–Ctab Reduction by Ascorbic Acid: Preparation and Characterization of Gold Nanoparticles. Colloids Surf., B 2013, 104, 11-17. 40. Tyagi, H.; Kushwaha, A.; Kumar, A.; Aslam, M. Ph-Dependent Synthesis of Stabilized Gold Nanoparticles Using Ascorbic Acid. Int. J. of Nanosci. 2011, 10, 857860. 41. Grinstead, R. R. The Oxidation of Ascorbic Acid by Hydrogen Peroxide. Catalysis by Ethylenediaminetetraacetato-Iron(Iii). J. Am. Chem. Soc. 1960, 82, 34643471. 42. Ciganda, R.; Li, N.; Deraedt, C.; Gatard, S.; Zhao, P.; Salmon, L.; Hernandez, R.; Ruiz, J.; Astruc, D. Gold Nanoparticles as Electron Reservoir Redox Catalysts for 4Nitrophenol Reduction: A Strong Stereoelectronic Ligand Influence. Chem. Commun. 2014, 50, 10126-10129. 43. Herves, P.; Perez-Lorenzo, M.; Liz-Marzan, L. M.; Dzubiella, J.; Lu, Y.; Ballauff, M. Catalysis by Metallic Nanoparticles in Aqueous Solution: Model Reactions. Chem. Soc. Rev. 2012, 41, 5577-5587. 44. Cava, M. P., Mitchell, M. J. Cyclobutadiene and Related Compounds; Academic Press, Inc.: London, UK, 1967. 45. Patton, E.; West, R. New Aromatic Anions. Viii. Acidity Constants of Rhodizonic Acid. J. Phys. Chem. 1970, 74, 2512-2518. 46. Goia, D.; Matijević, E. Tailoring the Particle Size of Monodispersed Colloidal Gold. Colloids Surf., A 1999, 146, 139-152. 47. Fry, F. H.; Hamilton, G. A.; Turkevich, J. The Kinetics and Mechanism of Hydrolysis of Tetrachloroaurate(Iii). Inorg. Chem. 1966, 5, 1943-1946. 48. Gelb, R. I. Conductometric Determination of Pka Values. Oxalic and Squaric Acids. Anal. Chem. 1971, 43, 1110-1113. 49. Williams, M. The Merck Index: An Encyclopedia of Chemicals, Drugs, and Biologicals, 15th Edition Edited by M.J. O'neil , Pb - Royal Society of Chemistry , Cambridge, Uk Isbn 9781849736701; 2708 Pages. April 2013. Drug Dev. Res. 2013, 74, 339-339. 50. Sant'Ana, A. C.; Santos, P. S.; Temperini, M. L. A. The Adsorption of Squaric Acid and Its Derived Species on Silver and Gold Surfaces Studied by Sers. J. Electroanal. Chem. 2004, 571, 247-254.

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63. Yang, L.; Liu, D.; Huang, J.; You, T. Simultaneous Determination of Dopamine, Ascorbic Acid and Uric Acid at Electrochemically Reduced Graphene Oxide Modified Electrode. Sensor Actuat. B-Chem. 2014, 193, 166-172. 64. Celikkan, H.; Aydin, H.; Aksu, M. L. The Electroanalytical Determination of Sodium Borohydride Using a Gold Electrode. Turk. J. Chem. 2005, 29, 519-524. 65. Qin, Y.; Ji, X.; Jing, J.; Liu, H.; Wu, H.; Yang, W. Size Control over Spherical Silver Nanoparticles by Ascorbic Acid Reduction. Colloids Surf., A 2010, 372, 172-176. 66. Pal, A.; Shah, S.; Devi, S. Preparation of Silver, Gold and Silver–Gold Bimetallic Nanoparticles in W/O Microemulsion Containing Tritonx-100. Colloids Surf., A 2007, 302, 483-487. 67. Grasseschi, D.; Ando, R. A.; Toma, H. E.; Zamarion, V. M. Unraveling the Nature of Turkevich Gold Nanoparticles: the Unexpected Role of the Dicarboxyketone Species. RSC Adv. 2015, 5, 5716-5724. 68. Bhawawet, N.; Essner, J. B.; Wagle, D. V.; Baker, G. A. Ionic Liquid Anion Controlled Nanoscale Gold Morphology Grown at a Liquid Interface. Langmuir 2017, 33, 6029-6037. 69. Aditya, T.; Pal, A.; Pal, T. Nitroarene Reduction: A Trusted Model Reaction to Test Nanoparticle Catalysts. Chem. Commun. 2015, 51, 9410-9431. 70. Ghosh, S. K.; Nath, S.; Kundu, S.; Esumi, K.; Pal, T. Solvent and Ligand Effects on the Localized Surface Plasmon Resonance (Lspr) of Gold Colloids. J. Phys. Chem. B 2004, 108, 13963-13971. 71. Lazarus, L. L.; Riche, C. T.; Marin, B. C.; Gupta, M.; Malmstadt, N.; Brutchey, R. L. Two-Phase Microfluidic Droplet Flows of Ionic Liquids for the Synthesis of Gold and Silver Nanoparticles. ACS Appl. Mater. Interfaces 2012, 4, 3077-3083.

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SCHEME 1: (A) Ascorbic Acid and Cyclic Oxocarbon Diacid Reducing Agents and (B) Proposed Mechanism for (i) Ligand Exchange, (ii) Au(III) Reduction, (iii) Au(I) Disproportionation, and (iv) Oxidation Product Hydrolysis for Squarate as Reducing Agent

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Figure 1. (A) Extinction profiles of AuNPs resulting from synthesis at 5, 20, and 100 °C using squaric acid as the reducing agent (R = 3.4). AuNPs were analyzed within 1 h of synthesis. Increasing temperature results in attenuation and broadening of the LSPR profile, with minor red shifting. Additionally, the 5 and 20 °C reductions displayed a small shoulder near 630 nm which derives from the presence of trigonal plates and nanorods observed in TEM studies. (B) Representative TEM image of SA-AuNPs synthesized at 20 °C.

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Figure 2. (A) Extinction spectra for SA-AuNPs made at various reductant-to-gold reagent ratios (R) measured 48 h after their preparation. After one week, all SA-AuNP solutions remain essentially unchanged, regardless of R value. To further assess AuNP evolution, the resulting (B) LSPR maxima (nm) and (C) full width at 0.75 max values were extracted and plotted versus time. The LSPR and [email protected] max values for R above 2.0 remain relatively constant over the monitored aging period. Low R values show marked changes over time, particularly within the first 48 h following synthesis, indicating a high degree of colloidal instability and alluding to likely morphological changes. The legend in (A) applies to panels (B) and (C) as well.

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Figure 3. (A) Extinction profiles of AuNPs synthesized at 5, 20, and 100 °C using ascorbic acid as reducing agent (R = 3.4) and analyzed within 1 h of synthesis. The LSPR bands are completely insensitive to reaction temperature. (B) Representative TEM image of AA-AuNPs synthesized at room temperature (20 °C).

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Figure 4. (A) Extinction spectra measured for AA-AuNPs made using various R values 48 h after their preparation. The LSPR peak and [email protected] max values for R > 0.75 remain relatively constant over the monitored aging period. Lower R values, however, are associated with intense spectral shifts over time (particularly, for R = 0.25 and 0.50), indicating ripening arising from colloidal instability. The legend in (A) applies to panels (B) and (C) as well.

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Figure 5. Extinction spectra of as-synthesized AuNPs generated at 5, 20, and 100 °C using (A) croconic acid (CA) and (B) sodium rhodizonate (SR) measured 1 h after synthesis (R = 3.4 in all cases). For the CAAuNPs, as the temperature increases, the LSPR band blue shifts and narrows significantly, indicating the formation of smaller and more monodisperse AuNPs. In the case of SR-AuNPs, the LSPR band is completely masked by the prominent intrinsic absorption properties of SR itself. However, at a reduction temperature of 100 °C, a clear shoulder near 540 nm can be observed which we ascribe to the formation of AuNPs. Representative TEM images of (C) CA-AuNPs and (D) SR-AuNPs, demonstrating that AuNPs do indeed form when employing both CA and SR as the reducing agent at room temperature. Interestingly, both systems resulted in bimodal AuNP size distributions with a significant population of ultrasmall AuNPs (CA-AuNPs: 3.5 ± 1.9 nm and 29.9 ± 6.3 nm; SR-AuNPs: 5.1 ± 1.0 nm and 29.7 ± 7.6 nm).

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Figure 6. (A) Extinction spectra of CA-AuNPs for various R values measured 48 h after synthesis. The corresponding time-dependent (B) LSPR peak (nm) and (C) full width at 0.75 max values are also provided. The LSPR spectral features remain relatively constant during aging for R > 1.0. Lower R values, however, are associated with significant shifts, especially within the first 48 h, suggesting colloid instability and attendant morphological alterations. The legend in (A) applies to panels (B) and (C).

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Figure 7. (A) Extinction spectra of SR-AuNPs for various R values measured 48 h after synthesis. The corresponding time-dependent (B) LSPR peak (nm) and (C) full width at 0.75 max values are also provided. The LSPR spectral features remain relatively constant during aging for R > 0.5. Lower R values, however, are associated with significant shifts, especially within the first 48 h, suggesting colloid instability and attendant morphological alterations. The legend in (A) applies to panels (B) and (C). The asterisk (*) denotes results which were indeterminable or highly uncertain due to interfering background absorbance from SR. The legend in (A) also applies to panels (B) and (C).

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Figure 8. Cyclic voltammograms of 1.7 mM (A) squarate and (B) croconate at a bare GCE in 0.1 M H2SO4 (pH 2.0) solution for scan rates between 10 and 100 mV s–1. Insets: plots of peak current vs. scan rate1/2 (●) and scan rate (▲) for the main oxidation waves. Main oxidation wave potentials of 0.795 V and 0.741 V (vs. Ag/AgCl RE) are demonstrated for the two-electron oxidation of squarate and croconate, respectively, along with a clear diffusion-controlled process.

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Figure 9. (A) Schematic representation of the redox potentials extracted from CV measurements for the studied reducing agents and HAuCl4. We include for reference the reduction potentials for commonlyemployed weak (citrate) and strong reducing agents (sodium borohydride, NaBH4) as well as other Au, Ag, Pd, and Pt metal salts popularly employed in metal nanoparticle synthesis. These results indicate that the reducing agents employed in this study are well suitable for the spontaneous reduction of Au salts at room temperature, but may not be pertinent for the low-temperature synthesis of other metal nanoparticles, in accord with our preliminary observations.

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Figure 10. Plots showing the catalytic performance of various AuNPs (R = 3.4) using NaBH4-assisted reduction of 4-nitrophenol (4-NP) as a model reaction. Each catalyst was tested at least five separate times within the first few days of AuNP synthesis using 5 mol% Au catalyst (relative to the 4-NP substrate). The apparent rate constants (kapp) were determined from the slopes of linear fits of ln(A0/At) versus time plots where A0 and At denote the initial and time-dependent absorbance, respectively, of the 4-nitrophenolate ion at 400 nm. The lower apparent rates for AuNPs made using croconic, squaric, and ascorbic acid compared with those made from the Turkevich (citrate) route may be attributed to the significantly larger nanoparticle size. Conversely, rhodizonate-based AuNPs displayed an apparent rate roughly twice that of the citrate AuNPs, presumably a result of the presence of ultrasmall AuNPs (2–6 nm) in SR-AuNPs.

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Room Temperature Turkevich Method: Formation ... - ACS Publications

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