An aqua regia-free chemical recovery and reprecipitation of Ir from IrO_x catalysts: optimisation of the extraction efficiency using surface response methodology

Abstract

The widespread deployment of hydrogen is essential in the energy transition away from fossil fuels. There is an increasing demand for iridium (Ir) in the proton-exchange membrane water electrolysis (PEMWE) industry, acting as the catalyst for the oxygen evolution reaction (OER). However, Ir is one of the rarest platinum group metals (PGMs) and its recycling requires significant energy and has considerable environmental impacts. Therefore, developing an innovative energy-efficient recycling method is essential to enhance the use of Ir. This study focuses on the development of a microwave-assisted leaching protocol of Ir from a commercial iridium oxide (IrO_x) catalyst and its reprecipitation in the form of a salt. A Box–Behnken optimisation of the Ir leaching efficiency was performed by varying the temperature (70–200 °C), chloride concentration ([Cl⁻]) (0–4 M) and hydrogen peroxide concentration ([H₂O₂]) (0–3% v/v). An optimum of 91% extraction efficiency was predicted and then confirmed experimentally, achieving an Ir leaching efficiency of 83 ± 10%, which corresponds to an Ir mass extraction of 769 mg_Ir g_IrOx⁻¹. The precipitation of the Ir leachate solution was then performed to obtain Ir in the form of ammonium hexachloroiridate ((NH₄)₂IrCl₆) with 94.6% purity.

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Sustainability spotlight

The transition to hydrogen energy demands increased use of iridium (Ir) for proton-exchange membrane water electrolysis (PEMWE) catalysts. However, Ir is one of the rarest platinum-group metals, with its extraction and recycling posing significant environmental and energy challenges. This study introduces an energy-efficient leaching method to enhance Ir recycling efficiency, achieving an 83 ± 10% Ir recovery with minimal energy consumption. Such innovation aligns with Sustainable Development Goal #12 of the United Nations and is crucial for sustainable Ir use in PEMWE devices.

Introduction

Iridium (Ir) is one of the rarest earth elements, with a natural abundance of 0.02 ng g⁻¹ and a global production of less than ten tons a year.¹ This precious metal is valued due to its high corrosion resistance, high melting point and high density, which makes it a valuable component for high temperature applications such as spark plugs, hydrogenation catalysts, steam reforming catalysts, tail gas purification, and electroluminescence.^2–4 Various studies have also demonstrated the biocompatibility of IrO_x nanoparticles (NPs), and the use of IrO₂ NPs for medical applications is under study.⁴ Finally, iridium oxides IrO₂/IrO_x are used as state-of-the-art electrocatalysts for the oxygen evolution reaction (OER) at the anode in proton-exchange membrane water electrolysers (PEMWEs), facilitating the production of green hydrogen (H₂).^5–7 According to the Johnson Matthey Pt-group metal (PGM) Market Report of May 2023, the global annual demand in 2023 for Ir was estimated to be 7.4 tons, with 3.3 tons coming only from the electrochemical industry.⁸ The limited supply of Ir and its growing demand highlight the necessity to develop efficient recycling processes from end-of-life (EoL) materials. According to Minke et al., the technical EoL recycling rates for Ir catalysts must reach a minimum of 90% to cover future Ir demand for the large-scale industrialisation of PEMWE devices.⁹

Hydrometallurgy is a well-researched method for the recovery of PGMs. It has an environmental advantage over pyrometallurgy due to its low process temperatures whilst maintaining high metal recovery values. A hydrometallurgical method is based on three key steps: (1) leaching the metal in appropriate aqueous media (alkaline or acidic) with the addition of different reagents (e.g. hydrogen peroxide (H₂O₂) to oxidize the metal atoms to an ionic form, hydrofluoric acid (HF) or sodium chloride (NaCl) to increase the complexing action);^10–12 (2) concentrating and purifying the metal in solution; (3) recovering the metal in pure form or as a salt.¹³ The leaching step is often the limiting process, especially for metals such as Ir, whose stability range encompasses most of the water domain at 25 °C, in the absence of a complexing agent.¹⁴ Aqua regia (HCl + HNO₃) remains the preferred solution to leach PGMs, despite high corrosivity.¹⁵ Other used leaching methods combine a high concentration of HCl and an oxidising agent such as hydrogen peroxide (H₂O₂)²² or sodium chlorate (NaClO₃).²³

However, metallic Ir is highly resistant to most acids and oxidizing agents, making hydrometallurgical methods largely ineffective for its dissolution. Instead, techniques such as alkaline fusion, molten salt chlorination, or alloy fragmentation have demonstrated high dissolution rates. The core principle behind these methods is to convert metallic Ir into a compound that can easily be dissolved. A summary of the methods existing to leach Ir is presented in Table 1 and the interested reader is referred to two recent review articles discussing the topic.^20,24 Yoshimura et al.²¹ proposed an innovative approach using “dry aqua regia” (a mixture of FeCl₃ and KCl), where Ir dissolves as Ir³⁺ ions. They achieved Ir recovery ratios of 28.3% and 45.7% through treatment with HNO₃ and H₂O₂, respectively. Similar methods employing HCl–NaCl–H₂O₂ systems have also been utilized to extract Pt and other PGMs.^25–27 Melke et al. introduced a method to recover PGMs from Membrane Electrode Assemblies (MEAs) used in PEMWEs.¹⁷ Their process involves separating the membrane from the catalyst layers, melting the resulting solid powder with sodium peroxide (Na₂O₂) at 550 °C to convert the metals into water-soluble complexes, and subsequently dissolving them in diluted hydrochloric acid (HCl). Each metal is then separated individually.

Table 1 Table summarizing methods reported in the literature for the leaching of metallic Ir

Techniques	Starting material	T [°C]	Time [h]	Chemicals used	Possible hazards	Yield	Ref.
Alkaline fusion/melt acid leaching	Metallic Ir	600	4	Na2O2 and HCl 3 M	Production of Cl2	100%	16 and 17
					Na2O2 is corrosive and toxic by inhalation
Molten salt chlorination	Metallic Ir	625	2	NaCl, Cl2, and aqua regia (HCl + HNO3)	Production of Cl2 and N2O	80%	18
					Use of concentrated acids
Alloy fragmentation	Metallic Ir	1300	1	CaCO3, HCl	Production of Cl2	100%	19
Alloy fragmentation	Metallic Ir	800–1500	0.2–11	Zn, Sn, Al, Ni, Cu, Mn, aqua regia or HCl + Cl2	Use/Production of Cl2	70–100%	20
					Alloys may be thermally unstable and ignite or explode
Dry aqua regia	Metallic Ir	350–400	2–8	FeCl3, HCl, HNO3, and H2O2	Production of Cl2 and N2O	45.7%	21

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Notably, all these methods require the use of concentrated acids and significant quantities of chemicals. These processes have substantial environmental impacts due to high energy consumption, extensive chemical use, and the emission of harmful gases (e.g., Cl₂ and N₂O). Moreover, it is important to keep in mind that all the studies presented here are applied to the leaching of metallic Ir. To the best of our knowledge, there is no specific method developed to recover Ir from its oxides, especially as NPs.⁴ Therefore, there is a pressing need to develop a more environmentally friendly IrO_x leaching process.

In this study, we describe the development of an environmentally friendly Ir leaching process from iridium oxides (IrO_x), designed to minimize environmental impact. Our approach avoids the use of aggressive reagents and operates under milder conditions compared to conventional hydrometallurgical methods. This was achieved by developing a microwave-assisted protocol for leaching commercial iridium oxide dihydrate (IrO₂·2H₂O) without the use of aqua regia, employing low acid concentrations and moderate temperatures. The leaching parameters were optimized through a systematic approach, beginning with a screening design followed by a Box–Behnken optimization design to identify the most influential factors and maximize Ir extraction efficiency. Following leaching, Ir ions were reprecipitated as a salt, concentrating them in solid form to facilitate recovery. This study represents a significant advancement in reducing energy consumption and reagent use while maintaining high efficiency in the Ir recovery processes.

Results and discussion

A commercial iridium oxide dihydrate catalyst (99.99%, Ir 73% min., Thermo Scientific Chemicals), referred to as IrO_x hereafter, was the starting material for all experiments performed in this work. Transmission electron microscopy (TEM) images, X-ray diffractograms (XRD) and X-ray photoelectron spectra (XPS) of this material are presented in Fig. S1.† X-ray diffraction reveals the presence of a metallic phase, but no Ir(0) is detected from the XPS fit, suggesting an Ir-core@IrO_x-shell structure in which metallic Ir is located in the core of the particles (XPS being a surface sensitive technique) and is coated with a poorly crystalline IrO_x shell. From the Scherrer equation, the crystallite size of the metallic fraction can be estimated to be 44 nm. Morphological heterogeneity poses a challenge for the Ir leaching process as the lixiviation proceeds on the surface of the IrO_x particles: larger particles are likely to dissolve more slowly than smaller ones. Chemical heterogeneity is a challenge, as the method developed will need to be able to dissolve Ir in various degrees of oxidation.

The Ir dissolution reaction can be summarised in eqn (1), while the standard potential vs. the standard hydrogen electrode (SHE) of the redox couples are given in eqn (2) and (3).²⁸ These equations show that the presence of chloride (Cl⁻) ions weakens the metallic bonds by enabling the formation of a stable complex (IrCl₆²⁻).²⁹

IrO₂ + 4H⁺ + 6Cl⁻ → IrCl₆²⁻ + 2H₂O (1)

IrO₂ + 4H⁺ + 4e⁻ = Ir + 2H₂O E° = 0.926 V_SHE (2)

IrCl₆²⁻ + 4e⁻ = Ir + 6Cl⁻E° = 0.860 V_SHE (3)

The Gibbs free energy of eqn (1) calculated using the standard potentials of reactions eqn (2) and (3) is negative (ΔG° = −0.025 kJ mol⁻¹), indicating that it is spontaneous. However, thermodynamic feasibility does not guarantee fast Ir dissolution kinetics and preliminary leaching experiments at T = 40 °C for 24 hours in both acidic and alkaline media demonstrated limited Ir extraction efficiency. To accelerate the Ir dissolution kinetics, we used microwave heating. Microwave-assisted leaching has been used in previous studies for the recovery of PGMs from catalyst materials.^30,31 It is a highly efficient route as shown by Suoranta et al., who achieved a rhodium leaching efficiency of 79 ± 3% with microwave assisted leaching, compared to just 20 ± 1% with conventional heating.³¹

In this study, a microwave oven and polytetrafluoroethylene (PTFE) reactors were used to elevate the temperature and pressure of the solution, thereby kinetically favouring the production of IrCl₆²⁻. Details of the reactor vessels and conditions are provided in the ESI.† Due to experimental limitations of the Mars 6™ CEM microwave oven used, namely the propensity of the sapphire thermowell to leach at alkaline pH, acidic conditions were chosen for the lixiviation. Moreover, acidic conditions allow for ultraviolet-visible (UV-vis) titration of Ir in solution as IrCl₆²⁻. Indeed, Zhao et al.^32,33 demonstrated that in a high pH solution, IrCl₆²⁻ undergoes hydrolysis, where the chloride ligands are exchanged for hydroxide ligands, forming Ir(OH)₆²⁻, which is colourless (eqn (4)). Massué et al. suggested that the discolouration is attributed to the disappearance of strong ligands to metal charge transfer (LMCT) bands.³⁴ Therefore, it is important that the solutions are maintained at a low pH to accurately quantify Ir using UV-vis. The UV-vis spectrum displayed in Fig. S2a† is comparable to the literature.³⁵

IrCl₆²⁻ + 6OH⁻ → Ir(OH)₆²⁻ + 6Cl⁻ (4)

The IrCl₆²⁻ species exhibits most intensity in the region 600–400 nm with four distinctive peaks at λ = 488 nm (x), 434 nm (y), 418 nm (z), and 305 nm (w). H₂IrCl₆·xH₂O (99.9% trace metals basis, Merck) was used to make up standard solutions of Ir in a concentration range of 50–250 ppm. Linear plots of the absorbance at wavelengths x, y, z, and w against the concentration of Ir are displayed in the ESI (Fig. S2b).† Each line passes through the origin with an adjusted R-squared value >0.99 (Table S2†). The Beer–Lambert law states that the absorbance of light by a sample (A) is directly proportional to the concentration of the species in the sample (c) and the path length (l) through which the light passes. The molar extinction coefficient (ε/dm³ mol⁻¹ cm⁻¹) is determined from the slope of the line in the absorbance vs. concentration graph. The molar extinction coefficient values of the complex displayed in Table S2† are in the same order of magnitude as those in the literature.³⁵ The bands represent LMCT whereby the partially filled low-lying t_2g and e_g* molecular orbitals (MOs) of the Ir⁴⁺ cations accept electrons from the σ- and π-bonding MOs of the chloride ligands. The equation of the line at wavelength 488 nm was used to calculate the Ir concentration (C_Ir/mol L⁻¹) in the leachate as detailed in eqn (5), where A is the absorbance of the solution at 488 nm and 3897.7 (L mol⁻¹) corresponds to the slope (which corresponds to εl in the Beer–Lambert law) of the calibration curves of the absorbance against the concentration of Ir in H₂IrCl₆·xH₂O in standard solutions for λ = 488 nm (see Fig. S2b and Table S2†). The peak at this wavelength displays the highest absorbance intensity and is well-separated from other peaks. The maximum possible concentration of Ir in solution (C_Ir(max)/mol L⁻¹) was calculated from eqn (6) where m_IrOx is the initial mass of commercial IrO_x, 0.7361 is the fractional weight of Ir in IrO_x provided by the manufacturer, M_Ir/g mol⁻¹ is the molar mass of Ir, and V_L/L is the volume of the leachate. Finally, the percent extraction was calculated using eqn (7).

To first screen through multiple experimental factors (temperature, solid-to-liquid ratio (S/L), [H₂O₂], [H⁺] and [Cl⁻]) and determine those that are critical and those that are less significant towards the extraction efficiency of Ir in the leaching process, a two-level factorial design using Stat-Ease 360® was conducted. Factorial designs are statistical experimental designs that offer the ability to investigate the effects and the interactions of two or more independent factors. By manipulating the levels of the factors and measuring the resulting impact on the dependent variables, the main effects can be identified. Moreover, factorial designs allow a study of multiple factors simultaneously, reducing the number of experiments required as opposed to testing each factor separately. The ranges of each factor were chosen with careful consideration based on the results of preliminary experiments and what was experimentally achievable. Note that the factor range may affect the validity and interpretability of the results. For example, if the range is too narrow for a given factor, the study may fail to capture the full range in which the response occurs. On the other hand, a range too broad may infer practical restraints and a loss of precision where the response occurs.

Sulfuric acid (H₂SO₄) was selected to provide protons, and a range of 0.5 to 4.0 M for [H₂SO₄] was chosen. NaCl was used as a chloride source without further increasing the acidity of the leachate solution. [H⁺] and [Cl⁻] were chosen as separate factors to distinguish if one was more significant than the other and to investigate if it was possible to optimize the leaching process while retaining mild acidity. The upper limit of the S/L ratio was set at 0.25 g L⁻¹, while the lower limit for the S/L factor was set at 0.1 g L⁻¹. Due to experimental limitations (maximum mass of 2.5 mg in 10 mL in the reaction vessel) and practical restraints (increased error when weighing <2.5 mg), a broader range of S/L values and their effect on the Ir extraction efficiency could not be fully investigated. The T values were chosen within a range of 70 to 200 °C. The Mars 6™ CEM microwave handbook recommends the introduction of a second ramp at temperatures above 200 °C. To be consistent with the hold time, t of the reaction and avoid extreme temperatures and possible distortion of the vessels, 200 °C was chosen as the upper limit. Finally, H₂O₂ is an oxidising agent often used to aid the extraction of PGMs into solution.^15,27,36 It undergoes reduction into various species as shown in eqn (8)–(10),³⁷ aiding in the oxidation of metal species by accepting their electrons. For example, Kazemi et al.³⁶ reported higher platinum extraction with the addition of 3% v/v H₂O₂ in the leach solution according to eqn (11).

H₂O₂ + 2H⁺ + 2e⁻ = 2H₂O E° = 1.35 V_SHE (8)

H₂O₂ + H⁺+ e⁻ = HO˙ + H₂O E° = 0.39 V_SHE (9)

HO˙ + H⁺+ e⁻ = H₂O E° = 2.31 V_SHE (10)

Pt + 2H⁺ + 4Cl⁻ + H₂O₂ → PtCl₄²⁻ + 2H₂O (11)

Therefore, the concentration of H₂O₂ was considered an important factor to be tested for the Ir extraction efficiency. For safety reasons, a maximum concentration of 3% v/v H₂O₂ was selected. Higher concentrations would raise the risk of explosive H₂O₂ decomposition, facilitating degassing of the PTFE vessels. The absence of H₂O₂ (0% v/v) was chosen as the lower limit to test the full effect of this factor. Finally, before applying the surface design plan, the reaction time was set to 30 min. The kinetics of the reaction was then studied using the optimized parameters, as presented in the following.

Table S3† displays the chosen factors and their ranges for the factorial design in order to screen factors of least impact on the Ir extraction efficiency. Table S4† displays the design plan and the response values for the extraction efficiency of Ir. The experimental results were analysed with Design Expert software and models were refined by eliminating statistically insignificant model terms, except for those required to support hierarchy. The p-value probability was used to identify significance for each term. If the p-value is less than 0.05 then the terms are significant for the model, and those between 0.05 and 0.1 were manually evaluated. All effects for models, simple, interaction, and quadratic terms, showed a p-value of less than 0.05, indicating that models were significant and represented the area of study. The suitability and adequacy of both models were measured using ANOVA factors and are summarized in Table S5.† Analyses were carried out with a significance level of 0.05. This value represents a 5% risk of incorrectly concluding that a difference exists when, in reality, no such difference is present. No transformation for the Y response was required to ensure that the model meets the assumptions required for the analysis of variance (ANOVA). Statistical factors were acceptable, as shown in Table S6.† The lack of fit (LOF) (p-value > 0.1) was not significant which implied good quality of the model correlation between the process variables and the responses. The multiple correlation coefficient adjusted R-squared superior to 0.85 and its closeness to predicted R-squared confirmed the accuracy of the models. The coefficient of variation (CV), an indicator of the degree of precision, up to 10% means that the model has good accuracy. To complete this analysis, studying the residuals (= predicted response values − experimental values) is crucial to understanding their behaviour. It mainly concerns 3 points which must be verified:

(1) The residuals follow a normal distribution, because the experimental error must follow a normal distribution; (2) the residuals are independent of the response, because the variance of the experimental error must be constant and thus independent of the response; (3) the residuals are independent of the order of the experiments, because the variance of the experimental error must be constant and thus independent of the answer.

All diagnostic tests were satisfactory, implying that the models were acceptable for optimization and validation. The final model relationship between the response (Ir extraction efficiency) and the continuous factors, after the refining step, was described by a second order polynomial (eqn (12)).

From the screening phase, it can be safely concluded that the temperature and [Cl⁻] have a positive effect on the leaching process (Table S5†). Interestingly, while an increase in [H⁺] should theoretically shift the equilibrium of (eqn (1)) to the right, the ANOVA results from the factorial design indicate that [H⁺] is not as significant a factor compared to others in determining extraction dissolution efficiency. Therefore, [H⁺] was set to 0.5 M for the optimization process. Moreover, there seems to be a negative effect of increasing the S/L ratio on the Ir extraction efficiency, which can be attributed to the higher mass of the IrO_x starting material within the same liquid volume, leading to reduced dispersion of solid particles in the solution. Nevertheless, the interactions between S/L and H₂O₂, as well as temperature and H₂O₂, also demonstrate a positive effect. During the optimization step, an S/L ratio of 0.25 g L⁻¹ was selected, following the vessel recommendations.

Three factors were chosen, namely, T, [Cl⁻] and [H₂O₂], to optimize the Ir extraction efficiency. To optimise multiple factors for a desired response, surface design plans can be utilised. The most common methods include Central Composite designs (CCDs), Box–Behnken designs (BBD), and optimal designs.³⁸ BBD was chosen as it requires fewer experimental runs.³⁶ This method has already been utilized to optimize the recovery of Pt by Ding et al.²⁷ BBDs have three levels for each factor and are built to fit a quadratic model. The model works best using a small range of the chosen factors and gives high prediction precision in the centre of the factor space. The accuracy of the input is critical for the dependability of the model. The chosen ranges of factors (T, [Cl⁻] and [H₂O₂]) are displayed in Table 2 and were based on preliminary experiments and the screening phase in which an optimum of the Ir extraction efficiency was expected. Table 3 displays the experimental plan and the responses that were used for the Box–Behnken optimisation model. 3-Dimensional Box–Behnken optimisation models for the Ir extraction efficiency are displayed in Fig. 1. Fig. 1b clearly shows that an increase in [H₂O₂] enhances the dissolution in Ir. However, there is insufficient evidence to determine an optimum [H₂O₂] for the Ir extraction efficiency under these conditions. While it appears that further increasing the [H₂O₂] beyond 3% v/v could improve the Ir extraction efficiency, the concentration was only tested up to 3% v/v due to the explosive risks associated with the decomposition of H₂O₂ into molecular oxygen and water.

Table 2 Experimental factors and levels for the Box–Behnken design used for optimising the Ir extraction efficiency from IrO_x under microwave conditions

Factor	Unit	Level
		−1	0	1
[Cl^−]	M	0	1	2
Temperature	°C	70	135	200
[H2O2]	%v/v	0	1.5	3

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Table 3 Design plan and responses for the Box–Behnken optimisation of Ir extraction efficiency for factors at given levels under microwave conditions. Reaction hold t = 30 min, S/L = 0.25 g L⁻¹, [H⁺] = 0.5 M

Run	Temperature (°C)	[Cl^−] (M)	[H2O2] (M)	Response in Ir extraction efficiency (%)	Response in Ir mass extraction (mgIr gIrOx^−1)
1	135	2	0	35	254
2	135	1	1.5	19	137
3	200	1	0	9	63
4	70	0	1.5	0	0
5	135	0	0	0	0
6	135	1	1.5	41	302
7	200	1	3	91	109
8	70	1	3	15	666
9	70	1	0	5	34
10	70	2	1.5	9	64
11	135	1	1.5	24	173
12	200	0	1.5	0	0
13	135	0	3	0	0
14	200	2	1.5	11	84
15	135	1	1.5	24	172
16	135	1	1.5	27	198

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Fig. 1 (a) Side view of the optimisation model for the Ir extraction efficiency as a function of [Cl⁻] and temperature. [H₂O₂] = 3% v/v, S/L ratio = 0.25 g L⁻¹. (b) Side view of the optimisation model for the Ir extraction efficiency as a function of [Cl⁻] and [H₂O₂]. T = 135 °C and S/L ratio = 0.25 g L⁻¹. Graduated coloured areas from cool blue for lower response levels to red for higher values are presented in 3D graphs. Blue and red areas represent the 25% and 75% ranges of the maximal value obtained in the representation, respectively.

A striking result in Table 3 is the fundamental role played by [Cl⁻], as we observe an Ir extraction efficiency of 0% in each run where [Cl⁻] = 0 M. According to Fig. 1a, there is an optimum in [Cl⁻] for the extraction efficiency of Ir at [Cl⁻] = 1.6 M at T = 139 °C and [H₂O₂] = 3% v/v. The increase in extraction efficiency up to [Cl⁻] of 1.6 M can be explained by the shifting of the equilibrium to the right in eqn (1) according to Le Chatelier's principle. Above 1.6 M, the decrease in extraction efficiency can be explained by the evolution of chlorine (Cl₂) gas from the oxidation of Cl⁻ ions (eqn (13)).³⁹ Although Cl₂ evolution was not confirmed (as it was not investigated in this study), the release of Cl₂ gas from the reaction vessel could reduce the Cl⁻ concentration in solution, leading to a decrease in Ir extraction efficiency due to less reagent being available for the formation of IrCl₆²⁻. Further investigations, such as implementing the measure of the Cl₂ concentration during the leaching, should be investigated to confirm this hypothesis.

2Cl⁻ → Cl_2(g) + 2e⁻ (13)

From the optimisation model of the Ir extraction efficiency as a function of temperature and [Cl⁻] (Fig. 1a), there is also an optimum in the temperature at 139 °C. The increase in Ir extraction efficiency up to 139 °C can be explained by the increase in the rate constant for the reaction according to Arrhenius' law. The decrease in extraction efficiency at temperatures above 139 °C can be explained by the competition of other reactions, namely, the evolution of Cl₂ and the decomposition of H₂O₂, thereby decreasing the concentration of complexing and oxidising agent in solution, respectively. Indeed, it is well-established that increasing the temperature accelerates the decomposition rate of H₂O₂ (see eqn (8) and (9)).^40,41 It has also been demonstrated that Ir, under oxidative conditions, can catalyse the evolution of Cl₂ ^42,43 and that increasing temperature enhances the reaction rate. However, this is a hypothesis and further research studies need to be conducted to better understand the temperature maximum point.

Confirmation of the optimisation model is usually done at or near the numerical optimisation. Therefore, microwave leaching conditions of T = 139 °C, [Cl⁻] = 1.6 M, and [H₂O₂] = 3% v/v were selected with constant values for [H⁺] = 0.5 M, S/L = 0.25 g L⁻¹, and hold time t = 30 min. These conditions resulted in an Ir extraction efficiency of 83 ± 10% which falls in the range for the system model statistics in the ESI (Table S7).† This corresponds to an Ir mass extraction of 769 mg_Ir g_IrOx⁻¹. To determine the kinetics of the Ir dissolution reaction under microwave conditions and confirm the chosen reaction time, reactions were performed altering the hold t of the reaction. The results are displayed in Table 4. The concentration of Ir in solution against the hold t of the reaction is shown in Fig. 2.

Table 4 Ir dissolution in microwave-assisted leaching experiments, where the holding time at the desired temperature was investigated. The conditions were T = 139 °C, [Cl⁻] = 1.6 M, [H₂O₂] = 3% v/v, S/L = 0.25 g L⁻¹ and ramp time = 20 min

Hold time (min)	Absorbance at 488 nm	Concentration of Ir ( ×10^−4 M)	Max. Ir concentration ( ×10^−4 M)	Ir extraction efficiency (%)
0	1.7 ± 0.4	4.4 ± 1.3	9.6 ± 0.2	46 ± 14
20	2.6 ± 0.4	6.7 ± 1.3	9.6 ± 0.2	70 ± 13
30	3.0 ± 0.2	7.7 ± 0.8	9.6 ± 0.2	80 ± 9
59	3.0 ± 0.2	7.8 ± 0.7	9.6 ± 0.2	82 ± 7

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Fig. 2 Kinetic model of the extraction of Ir into solution at various holding times. T = 139 °C, [Cl⁻] = 1.6 M, [H₂O₂] = 3% v/v, S/L = 0.25 g L⁻¹. The measurements were conducted three times, and the average results are presented, with error bars representing the standard deviation of the three replicates.

Each experiment involved a ramp t of 20 min to the desired temperature. Increasing the hold t of the reaction from 0 to 30 min results in greater Ir extraction efficiency. However, beyond 30 min, the increase in Ir extraction efficiency was not statistically significant. This indicates that performing these leaching experiments for longer than 30 min is unnecessary. By limiting the hold time t to a maximum of 30 min, it is possible to limit the energy consumption linked to Ir leaching. The observed plateau in the leaching efficiency after 30 min could be associated with the deactivation of the leachate during the experiment (through H₂O₂ decomposition and Cl₂ evolution). This could also be explained by the presence of metallic Ir in the sample.

Finally, after being extracted into the aqueous phase, Ir was precipitated in the form of a salt by the addition of NH₄Cl and filtered out of solution, following the protocol introduced by Yoshimura et al.²¹ The reaction scheme is summarised in eqn (14) where (NH₄)₂IrCl₆ is the formed precipitate. An IrCl₆²⁻ containing solution obtained from the leaching step under optimum conditions (T = 139 °C, [Cl⁻] = 1.6 M, [H₂O₂] = 3% v/v, [H⁺] = 0.5 M, S/L = 0.25 g L⁻¹, and t = 30 min) was used as the starting material. The leachate was reacted with NH₄Cl at the highest concentration possible without saturating the solution. This approach was used to shift the equilibrium in eqn (14) to the right, maximizing the formation of the desired precipitate while avoiding the introduction of undissolved NH₄Cl impurities.

2[NH₄]_(aq)⁺ + IrCl_6(aq)²⁻ → (NH₄)₂IrCl_6(s) (14)

Fig. S3† displays the process of the precipitation of (NH₄)₂IrCl₆ salt. The change in colour of the solution from deep orange to clear confirms the decrease in the IrCl₆²⁻ concentration in solution and the transfer of Ir into the solid phase. The obtained precipitate was analysed by XRD as shown in Fig. 3 and the pattern was compared to that of Reig-i-Plessis et al. and Yoshimura et al. who analysed the same crystal structure.^21,44 The Bragg peak positions of (NH₄)₂IrCl₆ and NH₄Cl were indicated with markers from the expected patterns of (NH₄)₂IrCl₆ and NH₄Cl taken from the International Crystal Structure Database (ICSD).⁴⁵ An NH₄Cl minor phase is present because the product was not washed after filtration in order to prevent re-dissolving (NH₄)₂IrCl₆; therefore NH₄Cl may have recrystallised out of solution and remained on the filter paper. The main phase is (NH₄)₂IrCl₆ with 94.6% purity. Finally, the product was analysed by scanning electron microscopy coupled with energy dispersive X-ray spectroscopy (SEM-EDS). Fig. 4 displays the octahedral crystal structure of (NH₄)₂IrCl₆. The energy-dispersive spectrum displayed in Fig. S4† confirms the absence of other impurities in the precipitate sample. The atomic ratios of Ir, Cl and nitrogen (N) were determined from the normal mass as determined using the area under the peaks (Ir, 12 at%; Cl, 69 at%; N, 19 at%). The approximate 1 : 6 ratio of atomic Ir to Cl and 1 : 2 atomic ratio of Ir to N also confirm the presence of (NH₄)₂IrCl₆.

Fig. 3 X-ray diffraction pattern of the recovered precipitate (NH₄)₂IrCl₆.

Fig. 4 Scanning electron microscopy images of the (NH₄)₂IrCl₆ sample precipitate, magnified (a) 500 and (b) 2k times.

In summary, we presented a novel method for leaching IrO_x NPs using microwave-assisted heating in the presence of Cl⁻ under acidic and oxidative conditions. Our results demonstrate that this approach achieves satisfactory Ir leaching efficiency under milder conditions compared to those reported in the literature for metallic Ir leaching. Table 5 summarizes the key properties of this method using the same criteria as those in Table 1, allowing for a direct comparison with previously reported methods. To ensure a fair comparison between the IrO_x leaching methods, a life cycle assessment of the methods discussed in the literature and the one described here should be implemented, but this falls outside the scope of this work. Importantly, our method operates at significantly lower temperature (139 °C) than other techniques (>400 °C) and offers a faster protocol, which helps reduce energy consumption. Additionally, the acid concentration used is lower than that in existing processes for Ir (not IrO_x) dissolution. Nevertheless, our method still relies on Cl⁻ ions, which may lead to Cl₂ production—albeit in lower concentrations than in other methods—as well as H₂O₂, a corrosive chemical. Finally, the authors would like to emphasize that this method was specifically developed for leaching IrO_x particles and has not been optimized for metallic Ir. Further work is ongoing to assess the effects of IrO_x particle size, crystallinity, and speciation (mean Ir oxidation state).

Table 5 Main properties of the IrO_x leaching process described in this work. This allows for a comparison of the other Ir leaching processes described in the literature and summarized in Table 1

Technique	Starting material	T [°C]	Time [h]	Chemicals used	Possible hazards	Yield
Microwave assisted leaching	IrOx NPs	139 °C	0.5	H2SO4 0.5 M, NaCl 1.6 M, and H2O2 3% v/v	Potential production of Cl2	83 ± 10%
					H2O2 is corrosive and toxic by inhalation

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Conclusions

In this study, we introduced a microwave-assisted Ir leaching protocol from commercial IrO₂·2H₂O and its precipitation from the aqueous phase to recuperate Ir in the form of a salt (NH₄)₂IrCl₆. A Box–Behnken optimisation was conducted using an experimental plan with the following factors: temperature, [Cl⁻] and [H₂O₂]. An optimum Ir extraction efficiency from IrO_x was predicted from the model at 91% for conditions of T = 139 °C, [Cl⁻] = 1.6 M, [H₂O₂] = 3% v/v, where [H⁺] = 0.5 M, hold time t = 30 min, and S/L = 0.25 g L⁻¹. The experimental Ir extraction efficiency measured under these conditions was 83 ± 10%, highlighting a good correlation between theory and experiment and demonstrating the possibility to reach a high recovery efficiency of Ir from IrO_x. Furthermore, Ir in the form of IrCl₆²⁻ was precipitated from the solution as (NH₄)₂IrCl₆ with a purity of 94.6%, as confirmed by X-ray diffraction and energy-dispersive X-ray spectroscopy combined with scanning electron microscopy. This demonstrates the potential to recover leached Ir in a suitable form for revalorization.

Data availability

Data for this article, including the kinetic model of the extraction of Ir into solution (Fig. 2), X-ray diffraction pattern of the recovered precipitate (NH₄)₂IrCl₆ (Fig. 3), X-ray photoelectron spectra (Ir 4f band) and X-ray diffraction pattern of the commercial IrO₂·2H₂O catalyst (Fig. S1b and c†), UV-vis spectrum of IrCl₆²⁻ species in HCl after microwave-assisted leaching of IrO₂·2H₂O, calibration curves of the absorbance against the concentration of Ir in H₂IrCl₆·xH₂O in standard solutions (Fig. S2a and b†), and energy-dispersive X-ray spectrum of the (NH₄)₂IrCl₆ sample precipitate (Fig. S4†) are available at the Open Science Framework at https://osf.io/wkqzs/files/osfstorage/678a498b206c7757dd282b4d. Note: the data required to build Fig. 1 are available in Tables 2 and 3.

Author contributions

Sarah Turnbull: investigation; sample collection and analysis; formal analysis; drawing; writing – original draft. Delphine Clauss: conceptualization; sample collection and analysis; supervision; validation; writing – original draft, review & editing. Vincent Martin: sample collection and analysis; investigation; data curation. Jean-Pierre Magnin: conceptualization; methodology; validation. Laetitia Dubau: conceptualization; investigation; validation. Frédéric Maillard: conceptualization; data curation; funding acquisition; supervision; validation; writing – review & editing.

Conflicts of interest

There are no conflicts to declare.

Acknowledgements

This work was supported by the “France 2030” government investment plan managed by the French National Research Agency, under the reference “ANR-22-PEHY-0011”, in the frame of the MATHYLDE project.

Notes and references

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Footnotes

† Electronic supplementary information (ESI) available. See DOI: https://doi.org/10.1039/d5su00038f

‡ These two authors contributed equally.

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