Why do some metal ions spontaneously form nanoparticles in water microdroplets? Disentangling the contributions of the air–water interface and bulk redox chemistry

Abstract

Water microdroplets containing 100 μM HAuCl₄ have been shown to reduce gold ions into gold nanoparticles spontaneously. It has been suggested that this chemical transformation takes place exclusively at the air–water interface of microdroplets, albeit without mechanistic insights. We compared the fate of several metallic salts in water, methanol, ethanol, and acetonitrile in the bulk phase and microdroplet geometry (sprays). Experiments revealed that when HAuCl₄ (or PtCl₄) is added to bulk water (or methanol or ethanol), metal NPs appear spontaneously. Over time, the nanoparticles grow, evidenced by the bulk solutions' changing colors. If the bulk solution is sprayed pneumatically and microdroplets are collected, the NP size distribution is not significantly enhanced. We find that the reduction of metal ions is accompanied by the oxidation of water (or alcohols); however, these redox reactions are minimal in acetonitrile. This establishes that the spontaneous reduction of metal ions is (i) a bulk phase phenomenon in water and several non-aqueous solutions, (ii) minimally affected by the air–water interface or the microdroplet geometry, and (iii) is not limited to Au³⁺ ions and can be explained via the electrochemical series. These results advance our understanding of aquatic chemistry and liquids in general and should be relevant in soil chemistry, biogeochemistry, electrochemistry, and green chemistry.

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Introduction

Biochemistry, chemistry, and chemical engineering textbooks often limit the role of water in chemical reactions to dissolving/hydrating molecules and ions;^1–3 however, this view is rapidly evolving.^4–9 In the last two decades, experimental and computational studies have revealed that the skin of water – the air–water interface – and water–hydrophobe interfaces, in general, have anomalous features such as enhanced speciation of ions^10,11 and electrified interfaces.^12–15 Several studies have reported on orders of magnitude faster (e.g., 10²–10⁶×) reaction rates in water microdroplets compared with the bulk phase, thereby expanding the function of aqueous interfaces in chemical transformations beyond heat and mass transfer.^5,16–24 This exciting field – microdroplet chemistry – has emerged as a frontier in chemical science with potential implications for atmospheric aerosols,²⁵ sea sprays,²⁶ disease transmission,^27,28 life's origins,^16,19,29 the chemistry of clouds, fog, smog, and dew,^30–32 soil processes,³³ green chemistry,¹⁸ hospital disinfection,³⁴ the food industry,³⁵ and fizzy beverages;³⁶ also, one may expect similar trends in microbubbles/foam in fermenter broths,³⁷ wastewater treatment,^38,39 and electrochemistry.^40,41

While much excitement exists, several claims for purely interfacial effects being responsible for the rate enhancement have been challenged. This is because the vast majority of reports exploit electrospray ionization mass spectrometry (ESI-MS), which entails rapid solvent evaporation, solute concentration, temperature, electric field gradients, etc., conditions that are far from thermodynamic equilibrium and, therefore, not representative of common systems.^20,23,42,43 For instance, the emergence of superacid chemistry in the microdroplets of mildly acidic water (pH < 4)^44,45 and the formation of abiotic sugar phosphates in water microdroplets¹⁶ have been challenged recently^46–49 (for recent reviews on this subject, see ref. 15, 24, 50 and 51). Controversial reports on microdroplet chemistry are not limited to ESIMS alone. For instance, recently, Zare & co-workers claimed that the air–water interface of microdroplets spontaneously generates H₂O₂,^52–55 and the reported amounts varied from ∼1 ppm or 30 μM (in 2019 for pneumatically sprayed droplets)⁵² to 114 μM (in 2020 for condensed droplets)⁵⁴ to 180 μM (in 2021 for microdroplets condensed at 55% or 70% relative humidity).⁵³ As for the mechanistic insights, computer simulations of Head-Gordon & co-workers^56–58 have suggested the emergence of instantaneous ultrahigh electric fields at the air–water interface that may drive H₂O₂ formation, while the simulations of Ruiz-López and co-workers⁵⁰ and Gong and co-workers⁵⁹ suggested that the local electric field was lower at the air–water interface compared to the bulk. Cooks and co-workers have speculated that the air–water interface of microdroplets promotes the formation of water radical cations H₂O⁺* and anions H₂O⁻*, which constitute the primary redox species.²⁴ Colussi⁶⁰ also proposed a similar mechanism whereby spraying yielded a small fraction of oppositely charged droplets, i.e., comprising excess H₃O⁺ and OH⁻, and these droplets collided to form OH* and H* radicals.⁶⁰ Our team has challenged these reports^{24,27,52–57,60} by pinpointing the artifacts arising due to the ambient ozone contamination,⁶¹ and OH* radical formation due to ultrasonication;⁶² also, we have explained the appearance of trace (∼1 μM) H₂O₂ concentrations in pneumatic sprays based on the reduction of dissolved oxygen gas at water–solid interfaces.⁶³

In a recent feature by the Royal Society of Chemistry's magazine, Chemistry World, entitled, “Why is Microdroplet Chemistry Contentious?“,⁶⁴ the following excerpt appeared, “Zare's team made its hydrogen peroxide discovery while trying to synthesize gold nanostructures in microdroplet, armed with the knowledge that Michael Faraday had made gold films with lemon juice. They added sodium borohydride – which Zare calls modern lemon juice – to water microdroplets and found that the reaction occurred 100 000 times faster than in bulk liquid. As a control experiment, the team removed the reducing agent – the sodium borohydride – yet still produced gold nanoparticles. That was a shock to us, says Zare.” In response, here, we extend our investigation of H₂O₂ production in water microdroplets to the spontaneous reduction of Au³⁺(aq) to Au nanoparticles (Au NPs) containing chloroauric acid (HAuCl₄, 100 μM concentration).⁶⁵ Zare and colleagues have speculated that this chemical transformation occurs exclusively at the air–water interface. While it has been reported that gold nanoparticles can be formed in bulk water by supplying energy via microwaves,⁶⁶ how water microdroplets (i.e., the air–water interface) drive this transformation without an external energy source, e.g., electrical voltage or reducing agent, has baffled chemists. Here, we resolve this mystery by addressing the following questions:

(1) Is the spontaneous reduction of gold in water limited only to water microdroplets, or does it also occur in bulk water?

(2) If gold ions get spontaneously reduced inside water, does water get oxidized? Is there a peroxide connection?

(3) What are the effects of the concentration of Au³⁺ ions, the water pH, and dissolved oxygen gas on the NP formation?

(4) Do other metallic ions, e.g., Pt⁴⁺, Al³⁺, Fe³⁺, Gd³⁺, Fe²⁺, Cu²⁺, Mg²⁺, Ni²⁺, Zn²⁺, and Na⁺, also exhibit this behavior?

(5) Does the air–water interface have unique features to form Au NPs, or can other common liquids such as ethanol, methanol, and acetonitrile also drive the spontaneous reduction of metallic ions?

Our experiments reveal that when salts containing Au³⁺ (or Pt⁴⁺ or Fe³⁺) ions are added to water they undergo spontaneous reduction in the bulk phase as well as in microdroplets. In fact, in this study, whenever NPs formed for a specific salt–solvent system, they formed both in the bulk phase and in the sprayed microdroplets. There was no scenario wherein the spontaneous reduction of Au³⁺ ions occurred exclusively inside microdroplets but not in bulk.

Results

In this work, we used MilliQ Advantage 10 (18 MΩ cm) deionized water, albeit the results were identical with HPLC grade water. Microdroplets were generated via pneumatic nebulization following the protocols of Zare et al.⁶⁵ Briefly, the spray device comprised two concentric capillaries, and the solution was injected through a 0.1 mm-wide inner stainless-steel capillary at a rate of 25 μL min⁻¹. The solution was sheared via high-pressure dry N₂ gas flowing through the outer concentric capillary at 100 psi, forming a microdroplet stream (ESI Fig. S1†).

The effects of the air–water interface on NP formation were analyzed by spraying 100 μM aqueous HAuCl₄ solution via pneumatic nebulization, which yielded microdroplets of an average size of ∼20 μm (ESI Fig. S1†). In a typical experiment, microdroplets were formed and collected in clean glass vials; also, the effects of solvent evaporation during spraying were compensated by adding bulk water to the glass vials to restore the original (bulk) salt concentration. Then microdroplets of the solution were transferred onto transmission electron microscopy (TEM) grids and SiO₂/Si wafers for TEM and scanning electron microscopy (SEM), respectively. Both techniques revealed NP formation in sprays. Curiously, the transparent mother solution (100 μM aqueous HAuCl₄ stored at 22 °C inside a laboratory) changed its color to red in 2–3 days and blue in 2–3 weeks (Fig. 2a). The TEM analysis of samples drawn from the mother solutions revealed the presence of gold NPs (Fig. 2c and d, and S3†). Dynamic light scattering (DLS) showed the appearance of a bimodal nanoparticle size distribution in the freshly prepared solutions composed of NPs and aggregates (Fig. 2b and S11b†). X-ray photoelectron spectroscopy (XPS) compared the oxidation states of gold in the newly prepared and one-week-old solutions (Fig. 2e and S11(e),† Methods). The oxidation state of gold in the freshly prepared solutions was +3 and +1, which indicated that the reduction was spontaneous, whereas the dominant oxidation state of gold in the 1 week-old samples was 0. These results demonstrate that the reduction of gold ions occurs spontaneously in bulk water and is not exclusive to the microdroplet environment/geometry.

Next, we investigated whether the microdroplet environment, i.e., the air–water interface, considerably accelerated the rate of reduction and nanoparticle formation compared with the bulk phase. To this end, bulk solutions of the same concentration (100 μM HAuCl₄) but different ages, i.e., varying from seconds to minutes to hours, were pneumatically sheared into water microdroplets (Fig. 1). The flying microdroplets were then intercepted and collected (in the liquid state, i.e., before complete evaporation) and analyzed via DLS. The observed particle size distributions were compared against those of the mother solution (Fig. S4†). Results revealed that the microdroplets afforded no significant enhancement in the particle sizes over the bulk phase, demonstrating that microdroplets do not have any significant effect on the formation or growth of nanoparticles. In summary, the reduction of gold ions in bulk water is spontaneous, and microdroplets are not essential, or the size of microdroplets is not the driving force for the spontaneity of this reduction reaction.

Fig. 1 Experimental design composed of solutions made from a variety of metallic salts and solvents. Nanoparticle formation was probed in bulk solutions and sprayed microdroplets to disentangle the contributions of bulk redox chemistry, microdroplet geometry, and the air–water interface.

Since the formation of AuNPs from Au³⁺ is a reduction reaction, a parallel oxidation reaction must occur simultaneously. Therefore, to reveal the redox reaction's oxidation half-reaction, we performed solution-state NMR spectroscopy to investigate the possible oxidation products in a liquid state. Gas chromatography (GC) was performed to determine the evolution of the formed gases. NMR spectroscopy revealed the spontaneous formation of H₂O₂(aq) in the bulk HAuCl₄ solution, and GC revealed the gradual evolution of oxygen (O₂) gas in the headspace (Fig. 3a and S5†). The formation of H₂O₂ was also quantified using a Hydrogen Peroxide Assay Kit (HPAK) (Fig. 3b). The mechanistic insights into this chemistry are presented in the Discussion section.

Fig. 3a presents H-NMR data comparing H₂O₂ formation in freshly prepared 0.3 M HAuCl₄ solution in MilliQ water against a freshly prepared 0.3 M HAuCl₄ solution in 200 μM standard H₂O₂ solution. The presence of only one peak in both the solutions, at 9.4 ppm, is observed by following the protocol of Bax and co-workers (detection limit ≥50 nM).⁶⁷ Using the same technique, we found that the concentration of HAuCl₄ increases in the aqueous solution, so does the spontaneous H₂O₂(aq) formation (Fig. 3b), i.e., they appear to be stoichiometric (see Discussion). We also examined H₂O₂(aq) formation in HAuCl₄ solutions prepared using water with dissolved oxygen and deoxygenated water. The dissolved oxygen had no measurable effect on H₂O₂ formation in these experiments (Fig. S6†). Thus, we confirmed that the formation of H₂O₂(aq) in bulk water is the oxidation half-reaction during the spontaneous formation of Au NPs.

Next, we investigated the effects of the solution pH and concentration of Au³⁺ ions on the formation of AuNPs and corresponding H₂O₂ produced. As the solution turned acidic, the formation of the AuNPs and H₂O₂ decreased (Fig. 4a–d). SEM imaging revealed that when a 200 μM HAuCl₄ solution was prepared using deionized water, or 0.1 M HCl, or 1 M HCl, the nanoparticles were well-formed (Fig. 4b), or smaller and fewer (Fig. 4c), or nearly absent (Fig. 4d), respectively. This indicates that gold ions are more stable in lower pH or highly acidic aqueous solutions. Following this observation, we investigated the effects of the gold ion concentration on the formation of Au NPs and H₂O₂. Experimental results revealed that gold ions could undergo complete reduction to form Au NPs in relatively diluted solutions (<0.05 M), whereas in concentrated solutions (>0.05 M), Au³⁺ primarily reduced to lower oxidation states (mostly Au⁺¹), while H₂O₂ was produced stoichiometrically. XPS pinpointed that as the salt concentration increased, the ratio of hydrated ions Au¹⁺/Au³⁺ decreased (Fig. 4e).

To further elucidate whether the mechanism underlying this chemical transformation, i.e., the spontaneous reduction of Au³⁺ in water, is unique or in common with other metal ions, aqueous solutions of a few common metal ions such as Pt⁴⁺, Al³⁺, Fe³⁺, Gd³⁺, Fe²⁺, Cu²⁺, Mg²⁺, Ni²⁺, Zn²⁺, and Na⁺ were studied. We found that two metal cations, Pt⁴⁺ and Fe³⁺, also produced H₂O₂ in water (Fig. 5c and d), while the remaining metallic ions did not produce H₂O₂ in bulk solutions (or microdroplets). Notably, considerably higher amounts of H₂O₂ were produced by Au³⁺ reduction than those produced via Pt⁴⁺ and Fe³⁺ reduction.

XPS analysis confirmed that Pt⁴⁺ and Fe³⁺ were spontaneously reduced in water similar to Au³⁺, yielding Pt²⁺ and Fe²⁺ (Fig. 5a and b), and the formation of Pt NPs was observed by TEM imaging (Fig. S7†). Notably, the reduction of Fe³⁺ to its zero-valent form was not observed. In other words, of all the metal ions studied herein, the pure metallic form of NPs was only observed in HAuCl₄ and PtCl₄ solutions.

Discussion

This section consolidates all the findings and presents mechanistic insights. Our experiments have revealed that Au³⁺ ions from HAuCl₄ get spontaneously reduced in bulk water, forming Au NPs and generating H₂O₂(aq). Spraying bulk solutions (freshly prepared or aged) and collecting the microdroplets reveal no significant enhancement in the size of AuNPs, suggesting that the microdroplet environment's effects, if any, are minor, such as facilitating solvent evaporation. Based on our experimental results, wherein Au³⁺ ions in bulk HAuCl₄ solutions spontaneously reduce to Au⁺ ions and AuNPs and drive the formation of H₂O₂(aq) (Fig. 2, 3, and S5†) and decrease in the solution pH, we suggest the following half-reactions:^40,68

Fig. 2 Spontaneous formation and growth of AuNPs in bulk 100 μM HAuCl₄ in water. (a) Bulk solution changing color due to the nucleation and growth of AuNPs. (b) Formation and growth of particles with time, as measured by DLS. (c) TEM micrograph showing the presence of AuNPs of different sizes. (d) TEM imaging showing an interlayer spacing of ∼0.235 nm corresponding to the Au(111) atomic plane of metallic gold Au(0). (e) XPS analysis showing the presence of Au³⁺ and Au⁺ in the freshly prepared solution and Au(0) in the 10-days-old solution, revealing the complete reduction of gold ions to AuNPs after some time.

Fig. 3 Production of H₂O₂ during the reduction of Au³⁺ ions in bulk water. (a) NMR spectroscopy pinpoints the formation of H₂O₂. (b) Quantification of H₂O₂ in freshly prepared aqueous solutions as a function of HAuCl₄ concentration.

Fig. 4 Effects of the solution pH and concentration on the reduction of gold ions: (a) decrease in the rate of H₂O₂ production in bulk solutions as a function of HCl concentration (or pH) for a fixed HAuCl₄ concentration. Representative scanning electron micrographs (SEM) of gold NPs formed in a 200 μM HAuCl₄ solution formed with: (b) deionized water, (c) 0.1 M HCl, and (d) 1 M HCl solution. (e) XPS reveals the ratio of Au¹⁺/Au³⁺ in freshly prepared HAuCl₄ solutions of varying concentrations (100 μM, 1 mM, and 5 mM).

Reduction half-reactions:

AuCl₄⁻(aq) + 3e⁻ ⇌ Au(s) + 4Cl⁻, E° (volts): 1.002 (1)

AuCl₄⁻(aq) + 2e⁻ ⇌ AuCl₂⁻(aq) + 2Cl⁻, E° (volts): 0.92 (2)

Possible feasible oxidation half-reactions that may drive the above reduction reactions:

O₂(aq) + 2H₂O(aq) + 4e⁻ ⇌ 4OH⁻(aq), E° (volts): 0.40 (3)

H₂O₂ + H⁺ + e⁻ ⇌ HO* + H₂O, E° (volts): 0.39 (4)

O₂(aq) + 2H⁺ + 2e⁻ ⇌ H₂O₂, E° (volts): 0.69 (5)

The standard reduction potentials of eqn (1) and (2) are higher than those of eqn (3) and (4); therefore, it is possible that OH⁻(aq) ions (eqn (3)) and H₂O may be oxidized to form H₂O₂(aq) (eqn (4)), which later decomposes/oxidizes gradually with time and forms O₂(g) (eqn (5)) further reducing gold ions.

Out of curiosity, we also probed the behavior of HAuCl₄ in methanol, ethanol, and acetonitrile, to see if there was something special about water in the nanoparticle formation. We observed that Au³⁺ reduced faster in methanol (and ethanol) than in bulk water, and it was negligible/sluggish in acetonitrile (Fig. 6a and b). The formation of Au NPs in methanol was accompanied by the formation of H₂O₂ (Fig. S12†), methylal (CH₃OCH₂OCH₃), and dimethyl ether (CH₃OCH₃) (Fig. S8†). Note: our investigation of the spontaneous formation of gold nanoparticles in the organic solvents was merely a curiosity, and an in-depth investigation of reaction byproducts and underlying mechanisms falls out of the scope of this manuscript. Lastly, in a complementary experiment, we quantified the amount of the reducing agent (standard hydrogen peroxide, 30%) required to completely reduce a fixed concentration of gold solution (0.1 M HAuCl₄) in methanol, water, and acetonitrile. The volume percent of H₂O₂ required was 0.4 ± 0.1, 1.0 ± 0.2, and 35 ± 1, respectively.

Now we discuss why some of the metal ions formed NPs in water, while others did not. For instance, we found that Pt⁴⁺(aq) was reduced to Pt(s) NPs and Pt²⁺(aq), whereas Fe³⁺(aq) was reduced only to Fe²⁺(aq) ions and Fe⁰(s) was not observed (Fig. 5 and S7†). Our observation about Fe is substantiated by the NIST XPS database⁶⁹ and high-resolution XPS spectra (Fig. 5b) in which the Fe 2p peak is broad and asymmetric, suggesting at least two oxidation states of Fe in the solution (Fig. 5b). The bottom line is that the reduction potentials of Pt⁴⁺(aq) to Pt(s) and Pt²⁺(aq) and that of Fe³⁺(aq) to Fe²⁺(aq) were more thermodynamically favored than the water oxidation reactions ((eqn (3) & (4) and ESI Table S2†). As the reduction potentials of Fe³⁺(aq) and Fe²⁺(aq) to Fe(s), i.e., −0.037 and −0.44 V, are lower than those listed in eqn (3) and (4), it is not feasible for Fe³⁺(aq) or Fe²⁺(aq) to reduce to Fe(s) by water. Similarly, the spontaneous reduction of Al³⁺, Cu²⁺, Ni²⁺, Zn²⁺, and Mg²⁺ ions in water was also blocked (ESI Table S2†).

Fig. 5 Spontaneous reduction of metal ions, besides Au³⁺, in water, which yields H₂O₂(aq). (a) XPS results of 100 μM PtCl₄ in water showing the reduction of Pt⁴⁺ to Pt²⁺; Pt NPs observed via TEM images (Fig. S7†). (b) XPS results of FeCl₃ in water showing the spontaneous reduction of Fe³⁺ ions to Fe²⁺. (c) NMR spectra confirming the formation of H₂O₂(aq) in aqueous solutions of FeCl₃ and PtCl₄ in water (each at a 100 μM initial concentration). (d) Comparison of H₂O₂(aq) formed by different concentrations of Au³⁺, Pt⁴⁺, and Fe³⁺ in water.

Fig. 6 Reduction of gold ions in different solvents. (a) Reduction of Au³⁺ in water, methanol, and acetonitrile. (b) Time-dependence of Au³⁺ ion concentration in water, methanol, and acetonitrile (all starting with 100 μM HAuCl₄) follows the trend: MeOH > H₂O ≫ CH₃-CN. (c) As a reducing agent, the amount of concentrated (30% v/v) H₂O₂ required for the complete reduction of 0.1 M HAuCl₄ in water, methanol, and acetonitrile. Since the amount of H₂O₂ required follows the trend: MeOH < H₂O ≪ CH₃-CN, it demonstrates that gold ions can undergo spontaneous (complete) reduction in water and methanol but not in acetonitrile. Note: the behaviors of methanol and ethanol were similar (Fig. S12†).

Next, we discuss the connection of the present findings with our recent reports unraveling the factors and mechanisms underlying the spontaneous formation of ∼1 μM H₂O₂(aq) in water microdroplets.^61–63 We have established that H₂O₂(aq) formation is independent of the air–water interface or the microdroplet geometry. Instead, it entails the reduction of dissolved O₂(aq) at the numerous water–solid interfaces involved in such experiments, such as pipes, tubing, containers, vials, plates, etc. Conversely, if dissolved O₂(aq) is removed from water, then H₂O₂(aq) formation is not observed in bulk, films, or microdroplets within a limit of detection of 50 nM. While some reports have corroborated our findings,^70–74 others have disagreed.^75–78 In contrast, the spontaneous NP formation and H₂O₂(aq) generation in water (bulk, films, or microdroplets) are driven by the reduction potential of metal ions and their concentration, and can be explained simply via the electrochemical series.⁴⁰

Here we list some limitations of this study. First, we did not focus on the shape specificity of the nanoparticles/nanowires as a function of the bulk HAuCl₄ concentration or as a function of solvent evaporation during spraying. Experiments with bulk HAuCl₄ solutions revealed that as the salt concentration exceeded 200 μM, nanowires began to appear – in bulk and in sprayed microdroplets without detectable differences (Fig. S13†). We note that if the solvent evaporation during pneumatic spraying becomes excessive, a regime that we did not study, we anticipate the precipitation of HAuCl₄ salt out of water, i.e., we do not expect a new interfacial chemistry to arise such that Au³⁺ ions specifically start reducing at the air–water interface to form Au NPs. Lastly, we did not take any special precautions to eliminate the existing nanoparticle impurities in the stock salts that may have accelerated the nucleation of nanoparticles in bulk solutions.⁷⁹

In summary, our report refutes the claims that water microdroplets or the air–water interface has unique properties that can drive the spontaneous reduction of gold ions. This chemical reduction takes place in the bulk phase, and it is not limited to (i) gold ions, or (ii) water as a solvent, or (iii) microdroplets or the air–water interface. The entire phenomenon can be explained by the standard electrochemical series, i.e., cations with very high reduction potential such as Au³⁺ and Pt⁴⁺ are so unstable that they can oxidize the solvent (water, methanol, and ethanol) to form NPs and H₂O₂(aq) and other byproducts in the case of alcohols. These findings advance our understanding of aquatic chemistry and warrant caution in investigating microdroplet chemistry and its environmental and practical relevance.

Methods

Chemicals

Gold(III) chloride trihydrate (HAuCl₄·3H₂O, 520918), platinum(IV) chloride (PtCl₄, 206113), iron(III) chloride hexahydrate (FeCl₃·6H₂O, 10025-77-1), iron(II) chloride (FeCl₂, 37287-0), aluminum(III) chloride (AlCl₃, 06220), zinc chloride (ZnCl₂, 211273), magnesium chloride hexahydrate (MgCl₂·6H₂O, 13152), sodium chloride (NaCl, 7647145), ammonium chloride (NH₄Cl, 213330), and copper chloride (CuCl₂, 7447394) were used. The following liquids were also used: deuterium oxide (D₂O, 3000007892), methanol-D4+0.03%TMS, acetonitrile-D3 (CAS 2206-26-0), acetonitrile (HPLC grade purchased from Fisher Scientific, batch 1072451), methanol (HPLC LC-MS grade), standard hydrogen peroxide (H₂O₂) 30% (Cat. 270733), HPLC grade water (Cat. 2594649), and deionized water (DI) obtained from a MilliQ Advantage 10 set-up (with resistivity 18.2 MΩ cm).

Characterization of NPs

SEM and TEM analyses were performed to characterize the NPs in the bulk solutions. For SEM, a small drop of liquid containing NPs was drop-coated on a silicon wafer substrate, and the solution was dried at room temperature. TEM images were acquired using a Titan ST Image Corrected (Thermo Fisher) instrument operated at a 300 kV acceleration voltage. EDS spectra and maps were acquired in the STEM mode using a four-quadrant SuperX EDS detector. For sample preparation, a drop of solution was placed on an ultra-thin carbon-coated holey copper grid, blotted with filter paper, and dried under ambient conditions. The presence and growth of NPs were examined via DLS. The growth of Au NPs was observed in 100 μM HAuCl₄ bulk solution at varying time scales.

XPS measurements

For the XPS studies of gold ion reduction in water, 100 μM HAuCl₄ solution was used. The drop coating method was employed to prepare the samples using silicon wafers, and Baer's method was closely followed.⁸⁰ A Kratos Axis Supra instrument equipped with a monochromatic Al Kα X-ray source (hν = 1486.6 eV) operating at a power of 75 W and under UHV conditions in the range of ∼10⁻⁹ mbar was used to obtain the data. All the spectra were recorded in hybrid mode, using magnetic and electrostatic lenses and an aperture slot of 300 μm × 700 μm. The high-resolution spectra were acquired at a fixed analyzer pass energy of 20 eV. The samples were mounted in a floating mode to avoid differential charging.

Quantification of H₂O₂

H₂O₂ concentration in all the diluted salt solutions was quantified using a hydrogen peroxide assay kit (HPAK). This is based on the principle that when hydrogen peroxide maintains contact with the AbIR peroxidase indicator, it causes fluorescence. Its maximum emission and excitation wavelengths are 674 and 647 nm, respectively. The samples were analyzed by mixing 50 μL of HPAK reaction mixture with 50 μL of samples in a 96-well black/transparent bottom microtiter plate using a SpectraMax M3 microplate reader. For the fluorescence reading, the SoftMax Pro 7 software was used. H₂O₂ in the samples was quantified using the calibration curve obtained from the standard samples on the same day.

Peroxide test strips for semi-quantitative analysis. For the qualitative estimation of H₂O₂ in the aqueous samples, peroxide test strips (Baker Test Strips purchased from VWR International) with a detection limit of 1 ppm were used. These strips use a colorimetric reagent that changes to blue when in contact with H₂O₂.

NMR spectroscopy for H₂O₂(aq) quantification

NMR measurements were performed on a 700 MHz Bruker Avance Neo NMR spectrometer equipped with a 5 mm Z-axis gradient TXO Cryoprobe at 275 K. For water samples, a 6 ms Gaussian 90-degree pulse was applied to excite the proton of hydrogen peroxide, followed by a 50 ms acquisition and a 1 ms recycle delay. For acetonitrile samples, a W5 binomial pulse sequence with gradients or a routine 1D proton pulse was applied for solvent suppression or non-suppression measurements, respectively. Peak quantification and peak position confirmation of H₂O₂ in salt solutions were performed by comparing the results obtained for standard H₂O₂ samples.

GC-MS experiment

The methanol sample was extracted by adding 2 mL each of water and hexane to initiate phase separation. Approximately 1 μL of the top hexane layer was analyzed using a single quadrupole GC-MS system (Agilent 7890 GC/5975C MSD) coupled to an EI source with an ionization energy of 70 eV. The ion source and mass analyzer temperatures were kept at 230 °C and 150 °C, respectively, with a solvent delay of 0 min. The mass analyzer was tuned according to the manufacturer's instructions, and the scan was fixed at 15–200 Da. A DB-5MS fused silica capillary column (30 m × 0.25 mm I.D., 0.25 μm film thickness; Agilent J&W Scientific, Folsom, CA) was used for chromatographic separation; the column contained a 5% phenyl and 95% methylpolysiloxane cross-linked stationary phase. Furthermore, helium was the carrier gas at a constant flow rate of 1.0 mL min⁻¹. The sample was injected into a 6890A gas chromatograph (Agilent, USA). The oven program was set to 30 °C and held for 10 min; the temperature was ramped at 20 °C min⁻¹ to 260 °C with a 0 min hold time. The GC inlet temperature was set at 250 °C, and the transfer line temperature to the MS EI source was kept at 320 °C. The sample was injected using an autosampler equipped with a 10 μL syringe into a split/splitless inlet with a split ratio 10 : 1.

Oxygen gas analysis

Headspace analysis of the aqueous HAuCl₄ solution was performed using a gas analyzer (Model 310C, SRI, USA) equipped with a thermal conductivity detector (TCD). About 1 mL of the headspace above the sample was withdrawn using a gas-tight syringe (Hamilton, USA) and injected into a 6-foot molecular sieve 13× packed column. The oven temperature of the column was set to 100 °C with argon as the carrier gas.

Data availability

All data needed to evaluate the conclusions in the paper are present in the paper or the ESI.†

Author contributions

ME and HM designed the experiments, which ME performed. NK and ME performed LC-MS, GC, and ICP-OES experiments. XG and ME performed NMR spectroscopy experiments. NW and ME performed XPS studies. HM and ME wrote the manuscript.

Conflicts of interest

The authors declare no competing interests.

Acknowledgements

HM acknowledges KAUST for funding (Grant No. BAS/1/1070-01-01). The co-authors thank Dr Adair Gallo, Dr Peng Zhang, and Ms Nayara Musskopf for building a robust experimental setup in HM's laboratory; Dr Ryo Mizuta, Scientific Illustrator at KAUST, for preparing illustrations in Fig. 1; Dr Valentina-Elena Musteata, scientist from IAC at KAUST for her assistance with TEM; Mohammed Abdulwahab ElShaer from WDRC at KAUST for his assistance with DLS and Usman Sharif from Laboratory Equipment Maintenance (LEM) team at KAUST for his contribution to maintaining the XPS instrument at the best operating conditions.

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Footnote

† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d4sc03217a

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