Catalytic Applications of Natural Iron Oxides and Hydroxides: A Review

Abstract

Iron oxides and hydroxides (Fe-OH) extracted from natural sources have garnered significant attention for their diverse catalytic applications. This article provides a comprehensive overview of the catalytic potential of naturally occurring Fe-OH, focusing on the influence of natural sources and preparation methods on their morphological characteristics and application in heterogeneous catalysis. The unique physicochemical properties of these catalysts, including their high surface area, redox activity, and tunable surface chemistry, make them promising candidates for various catalytic processes. The review discusses key catalytic reactions facilitated by natural Fe-OH, such as advanced oxidation processes (AOPs), electrochemical applications, catalytic cracking, and biodiesel production. Furthermore, it highlights recent advancements and challenges in utilizing these materials as heterogeneous catalysts. By presenting an analysis of the catalytic potential of natural iron oxides, this review aims to stimulate further research about the use of these materials, which are widely distributed in the Earth’s crust.

1. Introduction

Iron oxides and hydroxides, both referred to here as Fe-OH, are ubiquitous earth materials with high adsorption capacities for toxic elements and degradation ability towards organic contaminants. By mass, iron is the fourth most abundant element in the Earth’s crust, being found in 16 Fe-OH [1]. Iron oxides are widely distributed in nature and can be found across various components of the global system, including the lithosphere, hydrosphere, biosphere, pedosphere, and atmosphere. They participate in the complex interactions between these environmental compartments [2]. Iron’s different oxidation states (ferric (Fe³⁺) and ferrous (Fe²⁺)) further influence its interactions within these systems. For instance, iron availability in the oceans affects the growth of phytoplankton, impacting the entire marine food chain.

Geological Abundance of Iron Oxide and Hydroxide Minerals

Iron constitutes a significant portion of minerals on our planet, ranging from 28 to 35% of Earth’s total mass. The Earth’s inner core primarily comprises solid iron (Fe) and nickel (Ni), while the outer core consists of iron alloyed with 10–20% of other elements, such as oxygen (O), sulfur (S), and potassium (K). In the Earth’s structure, the primitive mantle contains roughly 12% iron, whereas the crust holds 4.32% iron, ranking just below oxygen (47.2%), silicon (Si, 28.8%), and aluminum (Al, 7.96%) in abundance. Consequently, iron silicates and iron oxides emerge as the most prevalent minerals in the Earth’s crust [3].

Table 1. Naturally occurring iron oxide minerals. Reproduced with permission from Reference [3]. Copyright© 2016, John Wiley and Sons.

Name	Formula	Crystal Symmetry
Magnetite (Mt)	Fe3O4	Cubic spinel
Hematite (Ht)	α-Fe2O3	Rhombohedral
Goethite (Gt)	α-FeOOH	Orthorhombic
Lepidocrocite	γ-FeOOH	Orthorhombic
Ferrihydrite (Fh)	5Fe2O3⋅9H2O	Hexagonal
Maghemite (Mh)	γ-Fe2O3	Cubic spinel
Wüstite	FeO	Isometric-hexoctahedral
Akaganéite	β-FeOOH	Monoclinic
Feroxyhyte	δ′-FeOOH	Hexagonal
Bernalite	Fe(OH)3	Orthorhombic

enumerates the naturally occurring iron oxide and hydroxide minerals. Goethite, hematite, and magnetite are the most prevalent, often occurring as standalone rock formations in nature. Lepidocrocite, ferrihydrite, and maghemite are found in significantly smaller quantities but are still distributed across many locations. On the other hand, wüstite, akaganéite, feroxyhyte, and bernalite are comparatively rare, both in terms of their distribution and abundance [4].

The primary source of iron is the rocks known as banded iron formations (BIFs). These formations are named for their characteristic composition, featuring alternating layers of iron (hematite or magnetite) and silica. Iron can also be found in geological formations like greenschist, gabbro, basalt, amphibolite, and others [5].

Schwertmann and Fitzpatrick (1993) [6] describe various notable features of iron oxides found on Earth’s surface. These minerals are physicochemical environmental indicators with unique mineral characteristics and strong pigmenting abilities, even at low concentrations. They tend to have fine grain sizes (5–100 nm) and large surface areas, allowing functionalized surfaces to absorb a range of molecules. Iron oxides are also highly insoluble and may include other cations like aluminum in their crystal structure. Due to their limited crystallinity, advanced techniques, such as Mössbauer spectroscopy, X-ray absorption near-edge structure (XANES), thermogravimetric analysis (TGA), X-ray diffraction (XRD), infrared spectroscopy (IR), differential thermal analysis (DTA), transmission electron microscopy (TEM), and selective chemical dissolution, are essential for identification [6].

Over the past decade, numerous researchers have investigated the utilization of naturally derived iron oxides as heterogeneous catalysts, driven by their widespread viability, readily accessible sources, low production costs, appealing physicochemical properties, and inherent biocompatibility [7,8,9]. The success achieved in numerous applications has led to reviews of articles on the subject. However, to the best of our knowledge, these works have focused on specific processes or the use of specific minerals. Therefore, there is a need for a comprehensive and integrative view of the topic. Fenton’s reaction, a well-established catalytic process, utilizes ferrous and/or ferric ions to decompose hydrogen peroxide, generating potent oxidizing agents capable of degrading a wide range of organic and inorganic contaminants. This reaction exemplifies a catalytic application of iron oxides. Thus, it is only tangentially addressed, referring readers to the reviews already published. A search on Scopus with the search terms TITLE (* “iron” AND “Catalyst” AND “review”) yields 35 results, of which only 14 are about natural minerals of iron oxyhydroxides. Three of these address their use in Fenton, photo-Fenton, and electro-Fenton processes [10,11,12], with four on water–gas shift, syngas and Fischer–Tropsch synthesis [13,14,15,16], three on persulfate activation [17,18,19], one on peroxides [20], one on nanoparticles applied in different catalysts [21], and two on advanced oxidation [22,23]. Additionally, a review of the performance of magnetite [24] in Fenton-like processes is identified. In just the last two years, by using the search terms ’iron’, ’oxide’, ’catalysts’, and ’review’ in article titles, the SciFinder^® search platform has yielded additional examples focused on Fischer–Tropsch synthesis [25,26], persulfate activation [27], NOx reduction [28,29], oxygen reduction reaction [30], and selective oxidation of hydrogen sulfide [31]. However, none of these address the use of natural materials. However, other catalytic processes involving iron oxides have also been tested. Among these applications are hydrocracking of coal tar, biodiesel production, hydrogenation, ammonia synthesis, reduction of nitrates, heterogeneous catalytic ozonation, oxidation and reduction of inorganic ions, hydrogen production, and water dissociation, which have not been sufficiently disseminated and analyzed in the literature. The main catalytic application of Fe-OH and the iron species used in them are depicted in Figure 1.

To the authors’ knowledge, there has been no previous attempt to present, in a single document, most of the applications proposed in the literature for naturally derived Fe-OH. In this review, chemical industry research professionals will find a comprehensive compilation of the identified applications of Fe-OH, especially in the last two decades. The aim is to provide an integrative approach, emphasizing the abundance and versatility of natural iron catalysts, which makes them economically attractive and environmentally benign. This document presents basic information on the relative abundance of Fe-OH in the Earth’s crust, their occurrence in nature, pretreatments of the catalysts, and both common and uncommon applications. It is divided into three sections: the first one for the most abundant, the second one for those less abundant but distributed globally, and finally, the third one for those considered difficult to obtain.

2. Most Prevalent Iron Oxides and Hydroxides

Magnetite (Fe₃O₄) and hematite (α-Fe₂O₃) are among the most prevalent iron oxides, while goethite (α-FeOOH) is the most abundant iron hydroxide. In this section, we will explore principal characteristics and both the primary and less common catalytic applications of these iron oxides and hydroxides.

2.1. Magnetite (Fe₃O₄)

Magnetite is a widely distributed accessory mineral in numerous ore deposits and ranks among the most abundant oxide minerals in the continental crust [32]. It occurs globally and can form under diverse temperature and pressure conditions in igneous, metamorphic, and sedimentary settings [33].

At ambient conditions, they crystallize in the inverse spinel cubic structure, which belongs to the space group $Fd\overline{3}m$. Its distinctive structure is characterized by a combination of diverse anions and cations. In a conventional spinel arrangement, divalent cations exclusively occupy tetrahedral positions, while trivalent cations are confined to octahedral coordination. However, within an inverse spinel configuration, the octahedral sites accommodate both divalent and trivalent cations simultaneously [34]. Fe³⁺ ions are positioned in tetrahedral (A) sites coordinated by oxygen, while octahedral (B) sites contain an equal mixture of Fe³⁺ and Fe²⁺ ions. Below the magnetic transition temperature of 850 K, the spins of iron atoms in the A and B sublattices align in a ferrimagnetic configuration [35].

Fe₃O₄ is the most utilized among various iron oxides, including FeO, γ-Fe₂O₃, and FeTiO₃, because it exhibits a notable degree of spin polarization at the Fermi level, a high Curie temperature, and electrocatalytic effect, stability, and ease of synthesis. Furthermore, nano-sized Fe₃O₄ shows additional advantageous features beyond those inherited from its bulk form (see

Table 2. Characteristics of bulk Fe₃O₄ and Fe₃O₄ NPs. Reproduced with permission from Reference [36]. Copyright© 2023 MDPI/CC BY 4.0.

Characteristics	Bulk Fe3O4	Fe3O4 NPs
Magnetism	Ferromagnetism	Superparamagnetism or ferrimagnetism
Saturation magnetization (Ms, 300 K, emu/g)	~(84–100)	Depend on size, shape, and coating: ~(0.5–92)
Size controllability	Unachievable	Precisely controllable: ~(2–100) nm
Shape controllability (Spherical, cubic, rod, hollow, and 2D nanoplate)	Unachievable	Precisely controllable
Specific surface area (m2/g)	0.34 *	Depend on size, shape, and coating: ~(20–300)
Magnetothermal conversion (W/g)	Absent	~(100–2500)
Electrocatalytic activity	Achievable	Achievable

* For an average crystallite size between 38 and 62 µm.

), such as remarkable size and shape controllability, a large specific surface area, and excellent magnetothermal conversion capability. Magnetite nanoparticles have been utilized in a wide range of catalytic applications, such as drug delivery, cancer treatment, organic synthesis, environmental remediation, synthesis of biodiesel, and electrocatalysis [36]. Magnetite is widely used in advanced oxidation processes (AOPs) [24]; however, its catalytic applications include organic synthesis, electrocatalysis, and environmental remediation, among others [36].

The catalytic performance of magnetite nanoparticles is strongly influenced by their size. In general, smaller nanoparticles exhibit enhanced catalytic efficiency due to their higher surface area and greater number of active sites compared to larger particles. Magnetite-based nanocatalysts are gaining prominence as adaptable tools in heterogeneous catalysis, contributing to the advancement of sustainable reaction methodologies [37].

Numerous methods exist for synthesizing bulk and magnetite nanoparticles [38]. However, the extraction of magnetite nanoparticles from natural sources has received limited attention. While some studies have explored this area [39,40,41,42,43], the primary objective of these syntheses typically differs from catalytic applications. Furthermore, it has been reported that different natural sources of magnetite and preparation methods significantly influence the morphology and size of the particles.

Magnetite prepared from iron sand via coprecipitation shows a more uniform morphology and particle size distribution. In contrast, the ball milling method allows control of particle size based on milling time, with longer times resulting in smaller particles but more irregular shapes (Figure 2A). On the other hand, when the natural source is mineral magnetite, the particles have more regular shapes. Nevertheless, the particle size distribution is broader (Figure 2B,C).

Main catalytic applications. Magnetite’s capacity to generate hydroxyl radicals enables its utilization in various advanced oxidation processes (AOPs), including Fenton, electro-Fenton, photo-Fenton, and Fenton-like reactions. These processes share the common principle of involving the interaction between iron-based catalysts and H₂O₂ to generate highly reactive hydroxyl radicals for degrading organic pollutants in water. Variations arise in activation methods: in electro-Fenton, H₂O₂ is generated by O₂ reduction at the cathode, photo-Fenton utilizes light to generate ^•OH and Fe²⁺ from water and Fe³⁺, and Fenton-like reactions use alternative iron species or transition metals [47]. According to catalytic evaluations and findings performed on various iron oxides in Fenton reactions, magnetite is highly active for both H₂O₂ decomposition and O₂ production at neutral and basic pH [48]. However, due to the magnetic nature of magnetite, the reactor design must consider conditions to prevent particle agglomeration, which could inhibit its catalytic activity by limiting proper dispersion in the solution. To address this issue, magnetic stirring is omitted, using centrifugal [49] or peristaltic pumps [50] for the recirculation of contaminants.

Mineral magnetite has been used to obtain catalysts for the degradation of organic pollutants such as p-nitrophenol [1], textile dye [44,49,51], methylparaben [52], and cefotaxime [46]. The natural sources used for the preparation of these catalysts were mineral magnetite and ferruginous sands. Additionally, the preparation of the catalysts commonly involves magnetic and gravitational separation as well as sieving and ball milling. In some cases, chemical treatment with acid, mainly HCl, is preferred to remove impurities present in the natural magnetite source.

Typical ranges of dosage for the Fenton and Fenton-like reactions are between 0.1 and 0.5 g·L⁻¹, and optimal degradation occurs in acidic or neutral conditions, with pH values of 3 to 7. Although H₂O₂ is the primary oxidizing agent used, other species, such as S₂O₈²⁻ [49,52], can also be employed. Magnetite obtained from natural sources commonly contains other minerals, including Cr, Mn, Ni, Co, and Ti, among others. However, these impurities can enhance catalytic activity. For example, the comparison between synthetic magnetite and mineral magnetite [1] for p-nitrophenol (p-NP) degradation indicates that the combination of nano-sized magnetite and H₂O₂ does not exhibit significant degradation throughout the entire process, underscoring the superior catalytic efficacy of natural magnetite compared to its synthetic counterpart. This fact was attributed to the presence of exsolved spinel due to the incorporation of transition metals with variable valence.

The preparation method of catalysts from natural sources significantly influences control over particle size, shape, and purity. Mufti et al. [51] and Amiruddin et al. [44] prepared magnetite nanoparticles using different approaches (coprecipitation and ball milling) and tested them in photo-Fenton degradation and Fenton degradation of methylene blue (MB), respectively. The coprecipitation method resulted in smaller particle sizes compared to the ball milling method (11 nm vs. 24–49 nm), with more uniform particle size distribution and a high-purity crystal phase; also, S_BET was higher compared to the usual values reported for Fe₃O₄. However, it involved the use of high molar concentrations of HCl and NH₄OH, as well as thermal treatment, which are avoided in the milling method. Although the catalytic performance was evaluated under different conditions (photo-Fenton and Fenton), the results for MB degradation were similar (80% vs. 86.89%). However, the reaction times required to achieve these results differed significantly (30 min vs. 120 min). Similarly, in the Fenton-like degradation of acid orange 7 (AO7) using natural magnetite and persulfate (S₂O₈²⁻), high degradation efficiency (90%) was achieved within 120 min under optimized conditions, including 0.5 gL⁻¹ of catalyst and 0.2 mM of S₂O₈²⁻ under acidic pH [49]. The catalyst exhibited low iron leaching and the potential to be scalable, while the reaction time was longer compared to systems leveraging photo-Fenton activation.

The catalytic degradation of methylparaben (MeP) using a Fenton-like process demonstrated the effectiveness of persulfate (SPS) activation by natural magnetite as a heterogeneous catalyst. Peroxymonosulfate and peroxydisulfate are common ions in aqueous effluents, whose presence can be leveraged to generate strong oxidants (SO₄.⁻, ^•OH, and ¹O₂) through appropriate catalysts [53]. The study achieved a high degradation efficiency of 99.5% and mineralization rates of up to 37% within 60 min under optimized conditions, including a catalyst loading of 0.1–0.3 g·L⁻¹ and SPS concentrations of 1–5 mmol·L⁻¹ under acidic pH. The catalyst exhibited excellent reusability, maintaining over 90% degradation efficiency after five cycles, with minimal decline in performance. This highlights its potential as a cost-effective and robust option for pollutant removal [52].

The preparation of composites of TiO₂/magnetite nanoparticles using natural magnetite and P25 as raw materials resulted in particle sizes of about 20–80 nm [46]. No significant increase in S_BET was observed. These particles were tested on Fenton, photo-Fenton, and photocatalysis processes. In the photo-Fenton under UV light and photocatalysis processes, the composites achieved complete degradation of cefotaxime, whereas in the Fenton reaction, the degradation was negligible. The higher degradation rate reached by mediated light degradation was attributed to the fact that Fe³⁺ in the magnetic nanoparticle could capture photo-generated electrons from TiO₂, which accelerated Fe^2+/Fe³⁺ cycles in the Fenton reaction, the separation of electron–hole pairs in photocatalysis, the decomposition of H₂O₂, and the production of OH radical, resulting in the effective removal of cefotaxime. Regarding its application in an electro-Fenton process, it was able to reach a degradation efficiency of about 40% on the degradation of gemcitabine at a pH of 3, which is related to ferrous ions in its structure [54].

In an initial approach to synthesizing nanoparticles from mineral magnetite, Munoz et al. [55] evaluated their catalytic activity in the catalytic wet peroxide oxidation (CWPO) of phenol. In this mode of the Fenton process (heterogeneous), the catalyst used to decompose H₂O₂ into ^•OH is an insoluble solid [56]. The preparation involved only sieving the raw material, which was obtained from Marphil, Madrid, Spain. This resulted in a Fe₃O₄ catalyst with an S_BET of 8 m²g⁻¹ and a saturation magnetization (Ms) of 77.7 emu g⁻¹. Under optimal reaction conditions (H₂O₂ = 500 mg·L⁻¹, Fe₃O₄ = 2 g·L⁻¹, T = 75 °C, pH = 3, and t = 4 h), the enhanced performance of the minerals was confirmed in the CWPO of phenol (100 mg·L⁻¹), achieving complete conversion of the target pollutant with a high degree of mineralization (70–80%). However, even though the process was effective, relatively low oxidation rates were observed, requiring up to 3.5 h of reaction time to reach the complete removal of the pollutant under ambient temperature and circumneutral pH. Therefore, mineral magnetite was modified through thermal treatments of controlled oxidation and reduction to obtain core–shell materials [57]. A slight growth in the magnetization of the material was observed with an increase in the reduction temperature. At a pH of 5, ambient temperature (25 °C), the stoichiometric amount of H₂O₂ for the complete mineralization of sulfamethoxazole (SMX), and a catalyst load of 1 g·L⁻¹ were selected. The initial concentration of SMX was fixed at 5 mg·L⁻¹. Reduced iron species exhibited greater activity in the CWPO of SMX, leading to complete removal within 1.5 h. The optimal catalyst was reduced at 400 °C. This material remained active even after five reaction cycles.

Recently, the CWPO of azole pesticides, using natural magnetite as a catalyst, has been investigated [50]. The study utilized a fixed-bed reactor packed with natural magnetite powder to eliminate a representative blend of azole pesticides. System performance was assessed by examining the effects of inlet flow rate (0.25–1 mL·min⁻¹), magnetite loading (2–8 g), H₂O₂ dosage (3.6–13.4 mg·L⁻¹), and initial pollutant concentration (100–1000 µg·L⁻¹) over 300 h of continuous operation. Under optimized conditions (T = 25 °C, pH = 5, flow rate = 0.5 mL·min⁻¹, [H₂O₂] = 6.7 mg·L⁻¹, and [Fe₃O₄] = 8 g), azole pesticide removal exceeded 80%. The catalytic system exhibited high stability over 500 h of operation, with limited iron leaching. Also, the efficiency of the catalyst was proven in the treatment of a real wastewater treatment plant (WWTP) effluent sample, spiked with a mixture of azole pesticides at 500 µg·L⁻¹. The process was less efficient in the WWTP effluent; this was attributed to the presence of organic and inorganic compounds in the treated water since they exhibit OH scavenging properties. On CWPO, H₂O₂ dosage and catalyst loading are the most significant variables that influence catalytic activity; even with low S_BET and high particle size, effective degradation can be reached. The reactions are described well by pseudo-first-order rate constants.

Non-common catalytic applications. The use of catalysts based on magnetite obtained from natural sources is not limited to environmental remediation applications. It has also been tested in the conversion of spent engine oil into reusable diesel-like fuel [45]. This process, known as catalytic pyrolysis, involves the use of solid catalysts to break down long organic matter chains into lighter hydrocarbons, promoting specific transformations such as dehydration, dehydrogenation, deoxygenation, and decarboxylation [58]. Natural magnetite was extracted from freshly mined raw ore through magnetic separation. Pyrolysis reactions were conducted at 500 °C with a residence time of 90 min. The overall yield was close to 100% in a thermal non-catalyzed run and decreased linearly with magnetite dosage. However, the results indicated that magnetite was selective to liquid fraction and could promote different product compositions due to cracking and oxidative dehydrogenation of the paraffinic and alkyl aromatics. The resultant catalytically derived liquid product has fuel properties comparable to commercial diesel. Recently, CaO/MgO/Fe₃O₄ prepared from iron sand and dolomite was tested as a catalyst for biodiesel production from waste cooking oil [57]. The catalyst calcinated at 1000 °C reached 95.64% yield. The produced biodiesel fulfilled the quality standards for density and viscosity.

Pyrolysis is a promising option in regions with limited petroleum and gas resources because it offers a solution by producing char, tar, and gases simultaneously. Char can be used for combustion or gasification, while tar can be processed into liquid fuels and chemicals, improving economic feasibility. However, coal-derived tar often contains heavy compounds that complicate processing and cause operational challenges. To address this, He et al. [59] proposed a coal pyrolysis method combined with iron ore reduction. Iron facilitates the breaking of C-C and C-O bonds. The study used magnetite with a chemical composition primarily consisting of Fe (62.65%) and minor components like SiO₂ and MgO. Pyrolysis experiments were conducted using a Pt-filament pyrolyzer, with volatiles adsorbed and later desorbed at 300 °C for GC/MS analysis. Coal and catalyst samples were heated to 700 °C under argon. Magnetite showed a low surface area (3.32 m²·g⁻¹) and primarily produced phenols and CO₂ but performed poorly compared to other iron ores like ilmenite and siderite. Also, the nanohybrid catalyst Fe@C, using natural magnetite as an iron source, was evaluated on Fischer–Tropsch synthesis, a collection of reactions for converting syngas (hydrogen and carbon monoxide) into liquid hydrocarbons [60]. The core–shell structure was confirmed by TEM analysis. The results indicated that CO conversion exceeded 98%, with C₅ hydrocarbon selectivity of 51.7% and stability over 180 h of steam [61].

Table 3. Summary of catalytic applications of natural magnetite (Fe₃O₄).

Catalytic Reaction	Natural Source	Pretreatment Method	SBET (m2·g−1)	Reaction Conditions	Results of Catalytic Evaluation	Rate Constant	Ref.
Fenton p-NP degradation	Ores	Crushing, grinding, and magnetic and gravity separation	0.54–2.46	Cat = 1.0 g·L−1; p-nitrophenol = 10 mg·L−1; pH = 7; H2O2 = 0.05 mol·L−1; NaOH = 0.1 mol·L−1	17.8–95.0% degradation	0.24, 0.19, and 0.12 μg·L−1·min−1	[1]
Photo-Fenton degradation of MB	Natural iron sands	Magnetic separation and coprecipitation	147.12	Cat = 2 g·L−1; MB = 10 mL of 20 ppm; UV irradiation	68.52–76.32% degradation	0.000194, 0.000262, and 0.000336 g·mg−1·min−1	[51]
Fenton degradation of MB	Beach sand	Drying, magnetic separation, and ball milling	ND	Cat = 1 g·L−1; MB = 15 mg·L−1; pH = 4.8; H2O2 = 150 mL L−1	86.79% degradation	ND	[44]
Photo-Fenton degradation of cefotaxime *	Natural magnetite	Ball milling method	2.92–5.64	Cat = 0.2 g·L−1; cefotaxime = 0.4 mmol·L−1; pH = 5.6; H2O2 = 10 mmol·L−1	100% degradation after 60 min	0.09 min−1	[46]
Fenton-like degradation of AO7	Natural magnetite	ND	ND	Cat = 0.5 mg·L−1; AO7 = 15 mg·L−1; pH = 5; S2${\mathrm{O}}_8^{-2}$ = 0.2 mM	75% degradation after 120 min	0.0019 min−1	[49]
Fenton-like degradation of MeP	Natural magnetite	As received	62.5	Cat = 0.3 g·L−1; MeP = 10 µmol·L−1; pH = 6.5; SPS = 5 mmol·L−1	90.2–99.5% degradation	0.0066 min−1	[52]
Electro-Fenton removal of Gemcitabine	Natural mineral	Crushing, sieving, and milling	ND	Pt anode; 140 mL·min−1 of air; 1 cm separation between electrodes; 0.05 mol·L−1 of Na2SO4	25–35% degradation	63 M−1s−1 min−1	[54]
CWPO of phenol	Natural iron minerals	Sieving	8.0	Cat = 1 g·L−1; phenol = 100 mg·L−1; pH = 3; H2O2 = 500 mg·L−1; T = 75 °C; t = 4 h	70–80% mineralization and 100% of conversion	11 L2 mg−1 gcat−1 min−1	[62]
CWPO of azole pesticide	Pristine magnetite	As received	7.5	Cat = 8 g; H2O2 = 6.7 mg·L−1; T = 25 °C; pH = 5; flow rate = 0.5 mL·min−1	Conversion of the pollutants > 95%	0.22 mL·gcat−1·min−1	[50]
CWPO of SMX	Pristine magnetite	As received	7.5	Cat = 1 g·L−1; H2O2 = 25 mg·L−1; T = 25 °C; pH = 5	50% mineralization and 100% removal	0.0263 mg·L−1·min−1	[55]
Fischer–Tropsch synthesis **	Natural magnetite	Sol–gel and high-temperature pyrolysis	73.4–157.7	Cat = 1 g; 2 g of quartz sand; T = 300 °C; P = 2 MPa; GHSV of 3000 mL·g−1·h−1; t = 180 h	93–98.5% conversion of CO	ND	[61]
Production of diesel by catalytic pyrolysis	Magnetite ore	Magnetic separation, crushing, grinding, and ball milling	ND	Cat = 3 wt%; T = 500 °C; t = 90 min	Selectivity towards the formation of hydrocarbons having fuel value in conformity with diesel fuel	ND	[45]
Cracking of coal	Mineral magnetite	As received	3.32	Cat = 1 mg; coal; 1 mg; T = 700 °C; Ar atmosphere	Selectivity to aliphatics	ND	[59]
Biodiesel production from waste cooking oil ***	Iron sand	Coprecipitation and wet impregnation	ND	Ratio of methanol to oil = 9:1; Cat = 1 wt%; T = 60 °C; t = 6 h; 1100 rpm	95.64% yield	ND	[63]

* TiO₂/Fe₃O₄, ** Fe@C, and *** CaO/MgO/Fe₃O₄; ND = No data; Cat = Catalyst.

summarizes the catalytic applications of hematite prepared from natural sources. Although its main applications are Fenton-type processes, its use for fuel production is also notable. Furthermore, the analysis of the elimination based on the main parameters in AOPs (Figure 3) shows that the specific surface is not a determining factor in the catalytic activity of magnetite; likewise, acidic or neutral conditions allow better performance. In addition, specific surface area is not a determining factor in the catalytic activity of magnetite, while acidic or neutral pH conditions enable better performance. Even with small specific surface areas and low catalyst concentrations, magnetite enables removal rates close to 100% in advanced oxidation processes.

2.2. Hematite (α-Fe₂O₃)

Hematite (α-Fe₂O₃) is the oldest recognized form of iron oxide, widely found in soils and rocks. It is more commonly found in sedimentary rocks, either as detrital particles or as precipitates. Hematite is abundant in environments with ample oxygen as the ferric oxyhydroxides initially formed age and transform into hematite. It typically coexists with other iron oxides like goethite. Hematite formation is favored by low water content and elevated temperatures [64]. When finely powdered, hematite exhibits a reddish hue, while in bulk form, it appears black or gray. Known for its exceptional stability, hematite often represents the final stage in the transformation of other iron oxides [2]. Its crystalline structure resembles corundum, with lattice parameters a = 5.0340 Å and c = 13.750 Å.

This mineral displays two magnetic transitions, specifically the Morin transition (T_M) around T = 260 K and the Curie transition at approximately 950 K. The manifestation of its weak ferromagnetic conduct stems from a slight misalignment in spin antiparallelism. Notably, a correlation is established between the T_M and the hematite particle size, with smaller particle sizes correlating to reduced T_M values. However, this relationship is heavily contingent upon preparation methodologies, crystalline lattice defects, and the incorporation of OH groups into the hematite structure [65].

Magnetic properties are intimately tied to nanoparticle morphology and size. For example, the magnetization of saturation decreases and coercivity increases with Mn doping, which are attributed to the presence of more manganese atoms at the grain boundaries [66]. The coercivity interval (H_c) spans from 31 to 530 Oe, while the remanent magnetization (M_r) ranges between 0.6 and 16 Am² kg⁻¹ [67,68,69]. Higher H_c and Mr values are closely linked to larger particle sizes. Conversely, smaller hematite particles of varied morphologies exhibit superparamagnetic behavior beyond the blocking temperature [70]. It is deemed a n-type semiconductor, characterized by a bandgap of 2.2 eV. The valence band comprises occupied 3d orbitals of Fe³⁺ alongside some non-bonding 2p orbitals of oxygen. Meanwhile, the conduction band is composed of vacant 3d orbitals of Fe³⁺ [71]. Regarding its specific surface area, it is recognized that this parameter is influenced by the synthesis method and calcination temperature. Hematite subjected to thermal treatment between 1173 and 1273 K yields a specific surface area below 5 m²·g⁻¹ due to particle sintering. Broadly, hematite produced at low calcination temperatures (<373 K) possesses specific surface areas ranging from 10 to 90 m²·g⁻¹, while those generated through Fe³⁺ hydrolysis are dependent on particle size and shape, encompassing values between 2 and 30 m²·g⁻¹ [2].

Main catalytic applications. While numerous publications document the use of hematite for dye removal, the majority of these studies focus on laboratory-synthesized hematite. However, dye removal using hematite-based catalysts from natural sources has been tested by some research groups. Photo-Fenton [72], Fenton [73,74], Fenton-like [75], and photocatalysis reactions [66,76,77,78] are the main applications where hematite has been tested. In photocatalysis, photon absorption by a material excites electrons from the valence band to the conduction band. The resulting electrons and holes react with O₂ and H₂O, generating O₂^•− and ^•OH radicals [76]. Although the preparation of hematite catalysts can result in various morphologies, temperature is the key parameter influencing this characteristic. For instance, catalysts prepared from natural hematite (Figure 4A) and from iron ore (Figure 4B) underwent a calcination process at 700 °C, resulting in similar spherical morphologies.

Phenol is naturally found in oilfield-produced waters, and it is commonly used in sectors such as pesticides and pharmaceuticals. Due to its persistence in the environment, much research focuses on its degradation by different catalytic processes. In an interesting approach, a catalyst using natural clay with cassava starch as a porogenic agent and raw material was tested on phenol photodegradation [78]. The catalyst presented around 40% of α-Fe₂O₃. The maximum phenol photodegradation achieved was 59.15% at a cassava starch concentration of 7%. Even after eight recycling cycles, the degradation efficiency remained above 90%. However, the synthesis method is more complex compared to other preparation methods [74] used to obtain catalysts for similar applications. Iron-bearing mining rejected from the amethyst geode deposit was processed by grinding, ball milling, and sieving to produce a catalyst for the photo-Fenton degradation of phenol. After this processing, the final percentage of Fe₂O₃ was close to 25%. Although the material had low porosity, it demonstrated significant catalytic efficiency in phenol degradation, achieving mineralization and removal efficiencies of 78% and 96.5%, respectively. This performance is largely attributed to its iron content, which effectively facilitates OH radical generation. Both studies demonstrated innovative approaches for preparing Fe₂O₃-based catalysts using low-cost raw materials. Despite differences in synthesis methods and reaction conditions, both highlight sustainable strategies for wastewater treatment.

A similar organic compound, namely, the low-molecular-mass 4-nitrophenol (4-NP), which is involved in many chemical processes and commonly present in soils and in surface and ground waters, causes severe environmental impacts and health risks. Fe₂O₃ [79] and Cu/Fe₂O₃ [77] catalysts were evaluated on the photocatalytic reduction of 4-NP to 4-aminophenol (4-AP), a less harmful compound. The raw material (iron ore) was crushed, sieved, and thermally treated to prepare the catalysts. Even if goethite was the main phase of iron oxide in the raw material, after heat treatment at 450 °C, only peaks related to hematite were found in an XRD powder pattern. The morphology of the catalyst prepared from natural sources and from iron nitrate is spherical in both cases. TEM analysis revealed a significant difference in particle size, with values of 50 nm (from natural sources) and 20 nm (synthesized). Impregnation with Cu (5 and 20 wt%) and calcination at higher temperatures resulted in a smaller particle size (50 nm) and reduction in S_BET (38 m²·g⁻¹), compared to hematite calcinated at 700 °C (34 nm and 10 m²·g⁻¹, respectively). The same reaction conditions were applied to 4-NP reduction, and conversion reached >99% in both cases. The rate constant was higher using Cu/Fe₂O₃. This was attributed to the presence of the copper ferrite CuFe₂O₄.

Hematite catalysts from natural sources have also been successfully tested in dye degradation. For the oxidation and mineralization of methyl orange (MO), catalysts were prepared from Fe-sand [72]. The sand underwent a series of purification steps to enhance its catalytic properties. Degradation experiments were conducted in a cylindrical batch photoreactor under UV irradiation. A higher rate constant of 0.048 min⁻¹ was achieved for a dye concentration of 100 mg·L⁻¹. One of the major current challenges is the scaling up of catalysts to enable their use in treatment plants [80]. A first step towards this goal is to test them in the laboratory using wastewater samples from real effluents. Under optimized conditions, including 200 mL of textile effluent, 0.3 g of catalyst, pH of 2.5, and an initial H₂O₂ concentration of 200 mg·L⁻¹, the system achieved an 88% reduction in optical density and a 79% decrease in chemical oxygen demand (COD) after 60 min. One interesting property of using iron oxide catalysts is their magnetic response, which allows for easy recovery and subsequent reuse. In this case, the catalyst was able to reach a degradation of 89% after four cycles, with no obvious leaching. Also, hematite can be used as a catalyst in combination with clay for the photocatalytic degradation of dye using raw iron ore [76]. α-Fe₂O₃/bentonite was further evaluated through photocatalytic degradation on indigo carmine (IC) of a composite prepared by mechanical milling, followed by magnetic separation and purification with HCl. Photocatalytic evaluation was performed under UV and solar light. The α-Fe₂O₃/bentonite composite has high photocatalytic activity on the removal of IC dye compared to extracted α-Fe₂O₃. The maximum removal efficiency (almost 100%) of IC was achieved with the following optimal conditions: pH = 1, 250 mg of catalyst, and 5 mg·L⁻¹ of IC. The photocatalytic activity under solar light was higher than that of UV irradiation. The removal of 88.88 % of MB was achieved by Mg/Fe₂O₃ prepared from iron sand in a photo-Fenton reaction [66].

The growing interest in the degradation of emergent pollutants results in the evaluation of novel or well-known catalysts in processes related to them. Among these emerging contaminants are herbicides. The simultaneous catalytic degradation of 2,4-dichlorophenoxyacetic acid (2,4-D) and 2-methyl-4-chlorophenoxyacetic acid (MCPA) herbicides over persulfate (PS) activated by mineral hematite via Fenton-like processes resulted in 36% of mineralization under optimum conditions (0.5 g·L⁻¹ of catalyst; 200 mg·L⁻¹ of 2,4-D and MCPA; 0.025 M of PS; pH = 3; T = 50 °C; and t = 120 min) [73]. The FESEM analysis indicated that the catalyst was almost porous, providing active sites for the adsorption of both pollutants and oxidant molecules. The synergistic effect between PS and catalysts showed better performance in acidic environments and higher temperatures in the degradation of herbicides.

Due to its abundance on Earth’s surface, different sources can be used to obtain hematite. In this line of knowledge, the same research group evaluated the use of Hormuz Red Soil (HRS) as a catalyst for the degradation of diclofenac (DCF) via PS activation [75]. The catalyst beads were prepared by ionic gelation encapsulation using alginate beads. For comparison, HRS in powder form was evaluated too. The catalytic experiments were conducted in two modes: batch experiments for the activation of peroxymonosulfate (PMS) by HRS and supported HRS by alginate beads in flow-through reactors. X-ray fluorescence (XRF) indicated that hematite was the main oxide on HRS (59.30%), with SiO₂ as the second most abundant compound (16.72%). Fresh catalyst particles were spherical, while HRS beads displayed a rougher irregular surface due to HRS crosslinking with alginate, enhancing the surface area and active sites for improved catalytic performance. Additionally, 98.2% of removal (HRS = 5 mg·L⁻¹; 75 mg·L⁻¹ of PMS; 50 mg·L⁻¹ of DCF; and t = 10 min) was achieved with HRS in suspension with PMS. Complete degradation was achieved in HRS supported on alginate beads in batch-mode processes (five granules of HRS beads; 75 mg·L⁻¹ of PMS; 50 mg·L⁻¹ of DCF; t = 2 min; and neutral pH). The alginate-supported HRS demonstrated superior performance, achieving complete degradation in significantly shorter reaction times under neutral pH conditions. However, the batch experiment was also efficient, with nearly 100% removal. Also, the catalytic ozonation and peroxone-mediated removal of acetaminophen (ACT) by HRS resulted in the complete degradation of 50 mg·L⁻¹ of ACT at 10 min. The calcinated catalyst exhibited high stability and reusability in consecutive catalytic cycles [81].

At the degradation of gemcitabine via the electro-Fenton process, hematite had poor degradation efficiency (20%) [54]. The analysis of parameters in Fenton and Fenton-like reactions catalyzed by hematite indicates that a narrow interval of pH values is available (Figure 5). This trend indicates that higher values of all these parameters cause a decrease in the removal efficiency of different contaminants in reactions catalyzed by hematite.

Non-common catalytic applications. While Fenton and Fenton-like reactions constitute the most prevalent applications of hematite, its utility has also been explored in hydrogen reactions. However, the literature reveals inconsistent results across different research groups in this area. The preparation of catalysts Mo/Al₂O₃/Fe₂O₃ using bauxite was evaluated through the hydrocracking of high-temperature coal tar, which was used as the raw material [82]. Hydrocracking involves breaking long C-C chains using a catalyst, hydrogen, and high pressures or temperatures [83]. The higher percentage of hematite was 16.85 wt%, and its catalytic evaluation results demonstrated varying conversion ratios and product yields among different bauxite-derived catalysts. The superior catalytic activity of the best catalyst was attributed to the leaching of Fe₂O₃ and the enrichment of Al₂O₃ in the bauxite modified with HCl, leading to increased acidity and specific surface area. This optimal combination enhanced mass transfer and acidity, resulting in compromised performance in terms of high-temperature catalytic treatment (HTCT) conversion and liquid yield. This result indicates that the presence of Fe₂O₃ is not favorable to the hydrocracking of high-temperature coal tar. In contrast, Kairbekov and coworkers [84] investigated the use of modified iron-containing catalysts and the preliminary ozonization of coal for its hydrogenation. The study focused on catalysts derived from natural bauxite samples with varying Fe₂O₃ contents. Among the tested materials, bauxite samples exhibited superior performance with 23.7% Fe₂O₃, achieving a higher yield of liquid products compared to other bauxite samples and catalysts such as red mud and pyrite concentrate. This result highlights the importance of Fe₂O₃ content in the catalytic hydrogenation process. In catalytic hydrogenation, hydrogen gas is used to reduce unsaturated organic compounds, such as aromatic rings, alkenes, or carbonyl groups, in the presence of a catalyst [85]. Although the specific iron oxide phase was not identified, it is likely that the Fe₂O₃ was in the hematite phase as bauxite typically contains hematite, goethite, and ilmenite. Furthermore, the study noted that the conversion of iron oxides into finely dispersed pyrrhotine during the hydrogenation process, particularly after modification with elemental sulfur, enhanced the catalytic activity of the iron-containing materials.

According to Widayat et al. [86], Indonesia’s substantial demand for catalysts is primarily fulfilled through imports, underscoring the critical need to identify and utilize domestic natural resources for catalyst production. Additionally, biofuel, particularly biodiesel, is one of the most widely used energy sources in the country, offering advantages such as being renewable and biodegradable, having lower emission levels, and promoting complete combustion in engines. The preparation of hematite catalysts from sand is of interest to address these issues.

In 2019, Widayat and coworkers prepared magnetic hematite nanoparticles from iron sand and tested them as catalysts for biodiesel production from waste cooking oil. The preparation of catalysts included acid leaching with HCl, chemical coprecipitation, and calcination at varied temperatures (650–800 °C). The optimal catalyst was calcinated at 700 °C. The catalysts were able to support both esterification and transesterification; biodiesel yield reached 86.78%, with 87.88% of fatty acid methyl ester; yield optimization was found to depend on the calcination temperature. Furthermore, Widayat’s group explored the potential of bifunctional α-Fe₂O₃/CaO₂ catalysts derived from iron sand for biodiesel production using waste cooking oil [87]. In the study, the catalysts were synthesized by combining hematite obtained through chemical coprecipitation with CaO₂ prepared from sources such as Ca(OH)₂, CaCO₃, and limestone. The catalytic performance varied depending on the CaO₂ source, with the limestone-derived catalyst yielding the highest biodiesel output at 77.17%. However, the CaCO₃-derived catalyst met both Indonesian and European biodiesel standards. It was also discovered that the α-Fe₂O₃ component favored the esterification reaction and formation of diglycerides, while CaO₂ favored the transesterification reaction and formation of methyl esters.

Catalytic cracking is a widely employed chemical process that breaks down complex hydrocarbons into simpler molecules using a catalyst. This process facilities the cleavage of carbon–carbon bonds in large hydrocarbons at lower temperatures and pressures compared to non-catalytic methods, resulting in enhanced energy efficiency and cost-effectiveness. In heavy oil, the increase in lighter hydrocarbons and reduced resin–asphaltene components was achieved by natural hematite containing 41% Fe, 58% Si, and impurities of Al, K, S, Ca, and Cu using a hydrothermal catalytic system. The density of the oil decreased as resin–asphaltene content was reduced [88], and the results demonstrated the possibility of increasing the number of lighter hydrocarbons and reducing their density by a regular decrease in the amount of resin–asphaltene components.

Hematite can be obtained via the heat treatment of other iron oxides. The catalytic cracking performance of hematite derived from thermally treated natural limonite was evaluated using toluene as a model compound for biomass tar [89]. The limonite was crushed, sieved, and calcinated at 700 °C for 2 h under an air atmosphere. The as-prepared hematite achieved 95% toluene conversion at 700 °C. However, at temperatures exceeding 700 °C, the excessive consumption of lattice oxygen led to carbon deposit accumulation, negatively impacting the conversion rate. The evaluation of the stability of the catalysts indicated that toluene conversion was maintained above 80% at 600 min. CO₂ and CO production confirmed the role of hematite as an oxygen carrier. XRD analysis showed the transformation of hematite into magnetite, which occurred due to the consumption of active lattice oxygen and the deposition of carbon; because of this, the spent catalyst was regenerated by heat treatment in an air atmosphere, with the regenerated catalyst maintaining a toluene conversion rate above 85% for the first three cycles.

Natural hematite has also proven to be effective in the catalytic cracking of coal tar [90]. Fe/Al and Fe/Ni composite oxygen carriers (OCs) were prepared through mechanical mixing. The hematite block was first crushed into a fine powder using a jaw crusher and then mixed with γ-Al₂O₃ or NiO powder. After extrusion molding and granulation, the mixture was dried in an oven at 100 °C for 12 h. Compared to natural hematite, the addition of γ-Al₂O₃ helps reduce the carbon deposition, while the inclusion of NiO promotes it. The carbon deposits on the reduced OCs mainly consist of hard coke formed through the graphitization of carbon black at high temperatures. In addition, γ-Al₂O₃ helps improve and preserve the porosity of the OC particles, significantly reducing sintering.

In recent years, catalysts from natural hematite have been tested on catalytic systems, which had not been evaluated before. Iron ore (72% hematite, 22% magnetite, and 6% goethite) was ball-milled as a precursor for the synthesis of a Ni-Fe catalyst for oxygen evolution reaction (OER) [91]. The OER involves the anodic electrooxidation of water in acidic media or OH⁻ in alkaline media to generate O₂. This reaction is ubiquitous in aqueous environments and typically requires significant overpotentials to proceed, which has driven extensive research into electrocatalysts [92]. Even at a low overpotential (280 mV), this electrocatalyst reaches a current density of 10 mA cm⁻² and retains its catalytic performance over prolonged electrolysis under alkaline pH. Synthetic hematite has been applied in various structures for the oxygen evolution reaction (OER), including metal–organic frameworks, carbon-supported materials, heterojunctions, and metal–nonmetal/anion-doped Fe₂O₃, primarily in alkaline media (KOH). Catalyst stability varies significantly, with no clear correlation to synthesis methods or composition. One of the cases reporting the lowest stability is Fe₂O₃/g-C₃N₄, which lasts only 10 min and is synthesized by a thermal method [93]. Stability durations of 5–12 h are observed for C-doped CoFe₂O₄/Fe₂O₃ [94], Co₃O₄/Fe₃O₄ [95], Fe/Fe₂O₃-Fe-N-doped C [96], FeS/Fe₂O₃ [97], Fe₂O₃/FeS [98], Fe₂O₃/CNT [99], and Fe₂O₃/FeP [100], synthesized by calcination, coprecipitation, pyrolysis, chemical etching/solvothermal, hydrothermal, coprecipitation, and hydrothermal methods, respectively. On the other hand, cases with acceptable operational durations (about 100 h) are reported for RuNi-Fe₂O₃/IF [101], WO₃/Fe₂O₃-NiO [102], and CoMo/Fe₂O₃ [103], synthesized via hydrothermal, chemical etching/decomposition, and hydrothermal/electrodeposition techniques, respectively. Finally, examples with extended durability include IrO₂–Fe₂O₃, which maintained its performance for 600 voltammetry cycles via thermal decomposition [104], and Fe₂O₃–MnO, synthesized via a sol–gel process with 1000 cycles [105].

Twelve minerals, which were washed, dried, milled, and sieved, were evaluated using the catalytic hydrolysis of microcystin-LR (MC-LR). Limonite displayed the best catalytic activity, while hematite resulted in poor yields [106]. Also, Fischer–Tropsch synthesis could be catalyzed by iron oxides. Hematite derived from iron ores allowed CO conversion of up to 80% when H₂ was the reducing gas. This conversion is higher than those achieved by commercial catalysts. However, it was achieved using relatively higher temperature and pressure (270 °C vs. 250 °C and 20.9 vs. 1.85 bar, respectively) [107].

Table 4. Summary of catalytic applications of natural hematite (α-Fe₂O₃).

Catalytic Reaction	Catalyst	Natural Source	Pretreatment Method	SBET (m2·g−1)	Reaction Conditions	Results of Catalytic Evaluation	Rate Constant	Ref.
Photocatalytic degradation of MB	Mg/α-Fe2O3	Natural sand	Magnetic separation, drying, and ball milling	ND	UV irradiation; t = 300 min	Removal of 88.8% of MB	ND	[66]
Photocatalytic degradation of phenol	α-Fe2O3/gelcasting porous ceramic	Natural clay	Grinding, sieving, and extraction by 3 M of H2SO4 at 80 °C	31.92	UV irradiation; T = 23 °C; pH = 8; t = 3 h; phenol = 10 mg·L−1	Removal of 57% of phenol; catalytic activity remains without changes after 8 cycles	ND	[78]
Photo-Fenton degradation of MO	α-Fe2O3/SiO2	Iron sand	Sieving, immersion in HCl solution, washing, and drying	11.16	Cat = 1.5 g·L−1; MO = 100 mg·L−1; H2O2 = 200 mg·L−1; pH = 3; UV irradiation	93% of degradation of MO. 89% after 4 cycles of reaction	0.048 min−1	[72]
Fenton degradation of (2,4-D) and MCPA	α-Fe2O3	Natural mineral	Crushing, sieving, washing, sonication, and calcination	60.4	Cat = 0.5 g·L−1; 2,4-D and MCPA = 200 mg·L−1; PS = 0.025 M; pH = 3; T = 50 °C; t = 120 min	36% of mineralization	MCPA = 0.0064; 2,4-D = 0.0059 min−1	[73]
Fenton and photo-Fenton oxidation of phenol in water	Mining reject	Mining reject	Ball milling and sieving	16.35	Cat = 0.75 g·L−1; phenol = 50 mg·L−1; H2O2 = 0.75 g·L−1; pH = 3	96.5% degradation at 180 min	0.0411 min−1	[74]
Degradation of DCF via PMS activation	α-Fe2O3	HRS	As received	5.17	Cat = 5 mg·L−1; DCF = 50 mg·L−1; PMS = 75 mg·L−1; t = 10 min; neutral pH	98.2% of degradation in 10 min	0.334 min−1	[75]
Photocatalytic degradation of IC	α-Fe2O3/bentonite	Iron ore	Magnetic separation, grinding, ball milling, and coprecipitation	ND	Cat = 250 mg; IC = 5 mg·L−1; pH = 1; solar light	100% of degradation after 2 h	ND	[76]
Photocatalytic reduction of 4 NP to 4 AP	Cu/Fe2O3	Natural iron ore rock	Crushing, sieving, calcination, and impregnation	10–42	Cat = 33.3 mg·L−1; 4 NP = 5 × 10−5 mol·L−1; NaBH4 = 12.5 mL (0.5 M); pH = 11.5; λ = 200–500 nm	>99% conversion in less than 1 min	2.34, 3.36, and 5.4 min−1	[77]
Photocatalytic reduction of 4 NP to 4 AP	α-Fe2O3	Natural iron ore rock	Crushing, sieving, heat treatment, and calcination	18–85	Cat = 33.3 mg·L−1; 4 NP = 100 mL (0.1 mM); λ = 220–550 nm; NaBH4 = 12.5 mL (0.5 M)	>99% conversion in less than 3 min	1.38 min−1	[79]
Catalytic ozonation and peroxone-mediated removal of ACT	α-Fe2O3	HRS	Calcination in air atmosphere	3.63	Cat = 1 g·L−1; ACT = 50 mg·L−1; O3 = 1.2 mg/min; t= 10 min; pH = 7	100% degradation	0.40 min−1	[81]
Hydrocracking of high-temperature coal tar	Mo/Al2O3-Fe2O3	Natural bauxite	Crushing, calcination, purification, washing, and wet impregnation	126.9–237.9	Cat = 1.3 g; sulfur powder = 0.4 g; HTCT = 42 g; T = 430 °C; P = 12.5 MPa; t = 90 min	The presence of Fe2O3 is not favorable to the hydrocracking of high-temperature coal tar	ND	[82]
Hydrogenation of coal	Fe2O3–SiO2–Al2O3–TiO2–MnO2	Natural bauxite	ND	ND	T = 380–440 °C; P = 3–5 MPa; sulfur additive = 0–2%; t = 90 min	Selectivity to liquid products	ND	[84]
Biodiesel production from waste cooking oil	α-Fe2O3	Iron sand	Magnetic separation, coprecipitation, and calcination	10.5–22.9	Esterification: methanol: waste cooking oil molar ratio of 5:1; H2SO4 = 1 wt%; T = 70 °C; t = 300 min. Transesterification: waste cooking oil: methanol molar ratio of 15:1; Cat = 1 wt%; T = 65 °C; t = 3 h	87.88% biodiesel yield	ND	[86]
Biodiesel production from waste cooking oil	Fe2O3/CaO2	Iron sand	Magnetic separation, coprecipitation, and calcination	ND	Waste cooking oil: methanol molar ratio of 1:15; Cat = 1 wt%; T = 65 °C; t = 3 h	97.04% biodiesel yield with a Fe/Ca ratio of 1:4	ND	[87]
Hydrothermal catalytic conversion of extra-heavy Ashal’chinskoe oil	α-Fe2O3	Natural hematite	As received	ND	T = 210, 230, and 300 °C; P = 2–18 MPa; t = 2 h	The possibility of increasing the number of lighter hydrocarbons in heavy oil and reducing its density by a regular decrease in the amount of resin–asphaltene components	ND	[88]
Catalytic cracking of toluene	α-Fe2O3	Natural limonite	Crushing, sieving, and calcination	12.4	Cat = 0.5–1.5 g; toluene = 1000 ppm; T = 500–800 °C; t = 60 min	95% of toluene conversion	ND	[89]
Catalytic cracking of coal tar	α-Fe2O3/γ-Al2O3 α-Fe2O3/NiO	Mineral hematite	Crushing, drying, and mechanical mixing	ND	T = 700–900 °C; consumption of coal tar = 0.3 kg·h−1; mass flow rate of coal tar = 5 g·min−1	Poor performance of hematite/NiO; the addition of γ-Al2O3 could effectively inhibit carbon deposition	ND	[90]
OER	Ni/iron ore	Iron ore	Ball milling and solution-assisted electrode preparation	ND	Cat = 150 mg; NaOH = 1 M	Achieves a current density of 10 mA cm−2 at a low overpotential of 280 mV; potentially scalable to industrial applications	ND	[91]
Catalytic hydrolysis of microcystin-LR peptides	α-Fe2O3	Mineral hematite	Washing, drying, milling, and sieving	ND	Cat = 20 mg; MC-LR = 1 mL (10 ppm); T = 60 °C	20.7% hydrolysis yield	ND	[106]
Fischer–Tropsch synthesis	α-Fe2O3	Raw iron ore	Grinding and sieving	59.0	Syngas at space velocity of 60 Nml·gcat−1; T = 270 °C; P = 20 bar	The CO-reduced catalyst exhibited the highest CO conversion of 94.1%, followed by the H2-reduced catalyst with a CO conversion of 80.1%, while the syngas-reduced catalyst showed the least CO conversion of 54.1%	ND	[107]

provides a summary of the catalytic applications of hematite, which can be derived from various sources. It is the most extensively utilized natural iron oxide in heterogeneous catalysis.

2.3. Goethite (α-FeOOH)

Goethite (α-FeOOH) is one of the most thermodynamically stable and abundant crystal structures of iron in minerals, sediments, and soils. It results from various processes (hydroxylation, oxidation, hydration, or decomposition) of iron silicates, magnetite, siderite, and pyrite [108]. Goethite adopts a compact hexagonal arrangement of OH⁻ and O₂⁻ anions, with Fe³⁺ at the center of the octahedra aligned parallel to the [001] direction [109]. Goethite is antiferromagnetic, with two Neel temperatures of 63 and 97 °C. Above this, it becomes paramagnetic, transitioning to hematite at temperatures above 400 °C [110]. It often exhibits a weak ferromagnetic component caused by a mismatch in the antiferromagnetic grid, associated with a higher number of defects such as hollows, boundaries, and grain rotations [111].

Main catalytic applications. Goethite has proven to be effective as a catalyst for removing organic contaminants from aqueous effluents, making it a promising natural candidate for water treatment applications due to its widespread abundance in the environment. Liu et al. [109] and Ruan et al. [22] conducted reviews of Fenton-like processes based on the literature on this topic. The direct formation of ¹O₂ and ^•OH has been identified, either independently or in the presence of H₂O₂ or UV [112], under slightly acidic pH conditions. Examples such as hydroquinone, 2-chlorophenol, oxalate, succinate, bisphenol A, and benzoate have been treated in this manner [113,114,115,116].

As a Fenton reagent, goethite demonstrates remarkable performance. De la Plata and colleagues [117] conducted a comparison between goethite and TiO₂ in heterogeneous photo-Fenton catalysis. The authors employed a spectrophotometric system specifically designed for highly precise measurement of the optical properties of catalysts, demonstrating that goethite has a significantly larger optical path length and the ability to effectively degrade 2-chlorophenol. In a more comprehensive study [118], it was demonstrated that the decomposition of various organic contaminants follows a trend: salicylic acid ≈ m-hydroxybenzoic acid > p-hydroxybenzoic acid ≈ benzoic acid > p-bipthalic acid > phenol > benzenesulfonic acid. This trend is analogous to the sorption pattern of these species on goethite, which, in turn, depends on the structure (substituent groups and number of rings).

Yeh et al. [119] conducted a comparison of the degradation of five chlorinated ethylenes and three aromatic hydrocarbons through Fenton-like processes. They observed that aromatics were more readily degradable than chlorinated compounds. However, removal was achieved within a few minutes in all cases. The type of electrolyte significantly influences ^•OH generation by goethite–H₂O₂. Lin et al. [120], demonstrated that the maximum generation occurs at pH 3, remaining substantial at pH 6–7. The most suitable electrolyte would be ClO₄⁻ as it does not impact radical formation, while other tested inorganic anions showed interference tendencies: H₂PO₄ > SO₄²⁻ > Cl⁻ > NO₃⁻. Cl⁻ and SO₄²⁻ more severely suppress ^•OH generation with an increasing concentration. The reason for this appears to be related to the adsorption and blocking of active sites by these anions. In general, based on the results reported in the literature [22], it can be concluded that for goethite, the photo-Fenton process yields better results than simple Fenton processes, and electro-Fenton is more effective than photo-Fenton. During toluene oxidation, Mn/goethite nanoparticles (Figure 6A) selectively produced CO₂ and water because no significant formation of an incomplete combustion product (CO) was detected [121]. In order to avoid changes in the catalyst structure, the catalytic performance was evaluated at a maximum temperature of 450 °C.

Natural goethite (Figure 6B) has been successfully applied to produce nanofibers for water disinfection under visible light, demonstrating remarkable efficiency in both dye degradation and bacterial inactivation. Specifically, it achieved 90% degradation of MB within 5 h. Its recyclability was confirmed by retaining a 65% degradation rate after five reaction cycles. The goethite used was sourced from an abandoned mine in the Taouz region of Errachidia Province, Morocco. Nanoparticles were synthesized by high-energy ball milling of goethite rock and wet milling of metallic iron rocks in a planetary ball mill. The process utilized stainless steel vials and iron balls, with a ball-to-powder mass ratio of 40:1, operating at a rotational speed of 450 rpm for 5 h [122].

More recently, highly dispersed natural goethite/Fe₃O₄ composites (Figure 6C) were synthesized via a one-step hydrothermal method, demonstrating a potential as catalysts for tetracycline removal. Under optimal conditions (pH of 3.0, tetracycline concentration of 100 mg·L⁻¹, catalyst dosage of 0.3 g·L⁻¹, and H₂O₂ concentration of 6.0 mM), a removal efficiency of 90.1% was achieved within 30 min. The system also exhibited strong durability, retaining a tetracycline elimination rate of 73.3% after five treatment cycles [123].

Another application of goethite is its role as an efficient dissociator of ozone molecules, making it valuable in heterogeneous catalytic ozonation [124], as demonstrated in the case of nitrobenzene by Zhang, Ma and coworkers [125]. In heterogeneous catalytic ozonation, O₃ is added to a contaminated effluent, along with a solid catalyst that decomposes the ozone molecule into reactive oxygen species, which, in turn, degrade the pollutant [126]. Toxic inorganic agents such as AsO₃³⁻ can also be oxidized (to AsO₄³⁻) this way, being more effective in the presence of H₂O₂ and UV. Zhang et al. [127] utilized goethite to adsorb tartaric acid and reduce dissolved Cr⁶⁺ in water through this process. They demonstrated that the presence of goethite is crucial; in its absence, tartaric acid alone promotes only a 12% reduction in Cr⁶⁺ in 72 h. However, the reduction is 100% within 24 h and can occur in just 4 h when reducing the acid concentration, allowing for the co-adsorption of Cr⁶⁺. The authors suggest that goethite’s role is to provide efficient channels for electron transfer between the reactants. Similar results were obtained for the reduction and stabilization of Cr⁶⁺ in soil by using calcium polysulfide catalyzed by goethite and hematite, where the presence of these oxides enhanced the performance. The removal of Cr⁶⁺ improved with increasing iron oxide content, from 0 to 9 g kg⁻¹, but then declined as the concentration rose from 9 to 12 g kg⁻¹ [128]. Pelalak and colleagues synthesized plasma-treated goethite nanoparticles, which exhibited enhanced surface area and a higher density of surface hydroxyl groups. These nanoparticles were derived from natural goethite [129]. They enhanced the heterogeneous catalytic ozonation of the sulfasalazine (SSZ) antibiotic.

Non-common catalytic applications. In a study on chemical synthesis, Fernández-Marchante [130] investigated the catalytic activity of goethite in the Westinghouse process for hydrogen production (HP), specifically focusing on the transformation of H₂SO₄ into SO₂. The study compared the performance of goethite to that of other established catalysts. According to their results, goethite demonstrated superior behavior compared to SiC but was lower than that of Fe₂O₃ and similar to that of CuO. The advantage lies in its resistance to sintering, making catalyst deactivation less likely over the long term. Shehzad et al. [131] used goethite as a catalyst for water dissociation in catalytic bipolar membranes (CBMs), reducing the activation energy (calculated by DFT) of the process by 80% compared to state-of-the-art membranes. This, combined with the natural availability of goethite, significantly enhances the process both economically and energetically. They achieved excellent catalyst stability through a protection strategy involving the in situ growth of conductive polyaniline (PANI). Hu et al. [132] employed goethite nanoflowers as support for Rh nanoparticles in the reduction and hydrogenation of nitrostyrene to vinylaniline. A key advantage over alternative methodologies is the ability to achieve 100% conversion without compromising the integrity of the vinyl group. This is a significant improvement as the co-reduction of other unsaturated groups is a common challenge in heterogeneous catalysts, often necessitating lower conversions to prevent unwanted side reactions. As a hydrogen donor, the use of hydrazine hydrate avoids the challenges of handling H₂ directly, generating only water and N₂ as byproducts in this process. Regarding stability, the authors conducted six cycles of catalyst recovery and reuse, detecting no loss of activity. Composites of Mn, Co, Ni, Cu, and Zn oxides were employed by Kuncser et al. [133] in the oxidation of cyclooctene. Catalytic oxidation occurs in the presence of O₂, using isobutyraldehyde as a reducing agent and CH₃CN as a solvent. In all instances, selectivity towards the formation of the corresponding epoxide exceeded 99% in the composites. However, the conversion varied in the order of Mn > Fe > Co > Ni > Zn > Cu, ranging from 55% to 4%. The concentration of basic sites in the composites was determined by measuring the UV–Vis signal of acrylic acid, irreversibly adsorbed by the materials. The results show a complete correlation between conversion and the ratio of the basic sites to the acidic sites of the composites. In microbial fuel cells (MFCs), goethite has demonstrated its ability to enhance catalysis [134]; the mineral was extracted from mining mud, and its performance was compared to stainless steel alone. Raw goethite tripled the electrical power obtained with stainless steel (11 W m⁻³ compared to 3.5 W m⁻³). Furthermore, when goethite was treated at 550 °C, the power increased to 17.1 W m⁻³. Additionally, it substantially improved the current and current efficiency obtained while reducing COD. Goethite exhibited superior performance compared to specularite and hematite as a catalytic bed for the catalytic reforming of lignite-blended co-pyrolysis with corn straw [135]. The process significantly enhanced the formation of light aromatic hydrocarbons. Verma et al. [136] employed goethite for the synthesis of α-aminonitriles from N,N-dimethylaniline using atmospheric oxygen. The authors experimented with various iron salts and minerals, including Fe(OAc)₂, FeCl₂, FeSO₄, FeCl₃, Fe(NO₃)₃, Fe₂(SO₄)₃, Fe(acac)₂, Fe(acac)₃, FeO@GO, Fe₃O₄, FeO, Fe@g-C₃N₄, and goethite. However, none yielded a performance greater than 17%. Remarkably, when goethite assumed a lychee-like morphology, the yield increased to 96%, indicating a high dependence on maintaining this specific morphology.

Figure 6. SEM micrographs of α-FeOOH particles prepared from natural sources: (A) acid mine drainage; (B) goethite rocks; (C) mineral goethite/Fe₃O₄; and (D) mineral goethite. Reproduced with permission from References (A) [121], (B) [122], (C) [123], and (D) [137]. Copyright © 2012, 2023 American Chemical Society; 2020 MDPI/CC BY 4.0; and 2023 Elsevier, respectively.

Figure 6. SEM micrographs of α-FeOOH particles prepared from natural sources: (A) acid mine drainage; (B) goethite rocks; (C) mineral goethite/Fe₃O₄; and (D) mineral goethite. Reproduced with permission from References (A) [121], (B) [122], (C) [123], and (D) [137]. Copyright © 2012, 2023 American Chemical Society; 2020 MDPI/CC BY 4.0; and 2023 Elsevier, respectively.

A recent study [137] examined the sunlight-driven activity and wavelength dependency of goethite obtained from Bingzhou, Hunan Province (Figure 6D). Before conducting photochemical experiments, the mineral was crushed, milled, and sieved to attain a particle size of less than 48 μm (<300 mesh). Goethite displayed strong light absorption, extending to approximately 570 nm in the visible spectrum, suggesting a narrow band gap favorable for photocarrier generation. Regarding hydroxyl radical (^•OH) production, goethite achieved a rate of 0.648 nM min⁻¹, surpassing both hematite and magnetite. Its performance in H₂O₂ production was still lower than that of hematite.

A comparative study on ^•OH photochemical production between natural and synthetic iron minerals revealed much higher rates in natural minerals. Specifically, natural goethite, hematite, and magnetite showed 1.7-, 4.1-, and 4.9-fold increases in ^•OH production, respectively, compared to their synthetic counterparts. Interestingly, the relatively smaller improvement in natural goethite’s performance corresponds to its lower impurity levels compared to the other minerals.

Goethite exhibits optimal catalytic performance under slightly acidic to neutral pH conditions, where the generation of hydroxyl radicals (^•OH) and reactive oxygen species is enhanced.

Table 5. Summary of catalytic applications of natural goethite (α-FeOOH). Unless otherwise stated, the catalyst is goethite.

Catalytic Reaction	Natural Source	Pretreatment Method	SBET (m2·g−1)	Reaction Conditions	Results of Catalytic Evaluation	Rate Constant	Ref.
Oxidation of toluene *	Acid mine drainage	Sequential precipitation method and wetness impregnation	115.70	Cat = 1 g; T = 250–450 °C; toluene = 0.9 vol%; flow rate = 500 L·min−1; GHSV = 18,000 h−1	30% of toluene conversion; selectivity to CO2	ND	[121]
Photocatalytic degradation of MB **	Goethite rocks	Ball milling, stirring, sonication, and electro-spinning process	ND	Cat = 0.5 g; MB = 100 mL (10 ppm); t = 5 h; visible light and UV	90% of bleaching after 5 h of illumination	0.0141 min−1; similar in UV and visible light	[122]
Fenton-like TC degradation ***	Mineral goethite	Hydrothermal method	20.4	Cat = 0.3 g·L−1; pH = 3; TC = 100 mg·L−1; H2O2 = 6.0 mM	90.1% removal after 30 min of treatment	ND	[123]
Reduction and stabilization of Cr(VI) in soil	Natural goethite	As received	ND	Cat = 3, 6, 9, and 12 g; pH = 8.6 to 9	Goethite increased the apparent rate constant	3.33 × 10−6 kg·mg−1·min−1 to 8.33 × 10−6 kg·mg−·min−1	[128]
Ozonation of SSZ ****	Natural goethite	Crushing, milling, washing, and plasma process	29.65, 73.24 and 77.31, respectively	Cat = 1–9 mg·L−1; T = 25 °C; P = 1 atm; inlet flow = 1 L·h−1; Na2SO3 = 1 mL (0.01 M)	96.5% degradation efficiency reached by N2-goethite after 40 min of reaction	0.076 min−1	[129]
Catalytic reforming of volatiles from co-pyrolysis	Mineral goethite	Dehydration	18.05–140.28	15–20 mg of a mixture of raw lignite and corn straw; T = room temperature to 800 °C; Ar flow = 100 mL·min−1	Goethite significantly promotes the production of light aromatic hydrocarbons	ND	[135]
Production of ROS	Natural goethite	Mechanical crushing, milling, and sieving	ND	Cat = 0.1 g·L−1; LI = 10 mW cm−2; pH = 7; T = 25 °C	Natural goethite ROS production was 4.9-fold higher than the standard	0.648 nM min−1	[137]

* Mn/goethite; ** PAN/goethite nanofibers; *** Goethite/Fe₃O₄; **** Natural N₂-treated goethite; ND = No data; LI = Light intensity.

presents the catalytic applications of goethite prepared from natural sources, with advanced oxidation processes being the most prominent. However, as previously discussed, synthetic goethite has been successfully used in other types of reactions, such as reduction and selective hydrogenation. Therefore, it would be worthwhile to test the efficiency of natural goethite in the same reactions. Nonetheless, its transformation into hematite above 400 °C should be considered.

Although the most abundant oxides are widely used in catalytic reactions, their preparation from natural sources is less common. In applications where high phase purity is required, their use can be detrimental to the reaction due to impurities commonly found in the sources from which they are obtained. However, these impurities can enhance catalytic activity in other processes. Additionally, in highly reducing or oxidizing environments, a phase change may occur during the process, potentially leading to its deactivation. While its magnetic properties are widely known and catalysts are magnetically characterized in some research reports, such characterization is carried out with a focus on the recovery and reuse of the catalysts. In spite of this, the effect of magnetic properties on catalytic activity has rarely been studied.

3. Less Abundant but Widely Distributed Fe-OH

While less abundant than magnetite and hematite, iron oxide such as maghemite (γ-Fe₂O₃), lepidocrocite (γ-FeOOH), and ferrihydrite (Fe_8.2O_8.5(OH)_7.4 + 3H₂O) also exhibit notable catalytic properties. This section will explore some of their applications in catalysis.

3.1. Maghemite (γ-Fe₂O₃)

Maghemite is widely recognized as a common component of various soils, particularly in highly weathered toxic soils found in tropical and subtropical regions. Recent environmental magnetic studies have shown that maghemite is distributed globally, having been detected in tropical, subtropical, arid, and tundra regions. The formation of ultra-fine maghemite during pedogenesis is considered the primary factor behind magnetic enhancement in paleosols. Studies have detailed the pedogenic processes responsible for maghemite formation, revealing that pedogenic maghemite, often associated with hematite, acts as an intermediate phase in the transformation of ferrihydrite into hematite. Experimental findings demonstrate that when phosphated ferrihydrite is subjected to a temperature of 150 °C for 120 days, it forms hydromaghemite (with grain sizes between 10 and 30 nm). Over time, the grain size of hydromaghemite grows, transitioning from the superparamagnetic to the single-domain range. After four months, hydromaghemite further transforms into the more thermodynamically stable hematite [138]; at nanosizes (36–40 nm), it transforms into hematite at 576 °C [139].

Maghemite (γ-Fe₂O₃) possesses an inverse spinel cubic structure with a lattice parameter a = 8.3500 Å. Its nomenclature derives from the amalgamation of magnetite (Fe₃O₄) and hematite (α-Fe₂O₃), owing to substantial congruence in both structural attributes and chemical composition. Operating as a ferrimagnetic substance at ambient temperature, it exhibits saturation magnetization values as impressive as 90 emu g⁻¹. However, its stability is inferior to that of hematite. Characterized as an n-type semiconductor, it displays a band gap of 2.0 eV. Similarities with magnetite (Fe₃O₄) are clearly observable. The distinctive demarcation lies in the exclusive presence of the Fe³⁺ oxidation state within maghemite, while magnetite encompasses both Fe²⁺ and Fe³⁺. The distinction between maghemite and magnetite is often complex when using XRD and X-ray photoelectron spectroscopy (XPS). However, some methods have been proposed to differentiate them from one another [140,141]. Determining the Curie temperature of maghemite has been proven to be challenging due to its tendency to transform into hematite, even at temperatures such as 413 K. This transformational behavior is contingent upon precursor selection, synthesis methodology, and external pressure exerted [142]. In addition, its Curie temperature is estimated to be within the range of 820 to 986 K [143,144]; nanoparticles present a considerably lower value of about 545 K [145]. Notwithstanding, since this analysis was conducted via the thermogravimetric analysis (TGA) technique and the observed transition temperature closely corresponds to the maghemite-to-hematite phase transition, assigning it as the Curie temperature remains a significant challenge.

Concerning its band structure, ab initio calculations elucidate that the upper segment of the valence band is dominated by oxygen 2p orbitals, with occupied 3d levels of Fe situated 6 to 7 eV below the Fermi level. The lower tier of the conduction band is populated by unoccupied 3d levels of Fe (octahedrally coordinated). This classification designates maghemite as a charge transfer semiconductor, with the foremost excitation potentially corresponding to electron transfer from O²⁻ anions to octahedral Fe³⁺ cations [146].

Synthetic maghemite has been successfully tested in hydrogenation reactions, with remarkable yields, especially when used in conjunction with other oxides [85,147,148]. However, to the best of our knowledge, maghemite prepared from natural sources has been poorly explored as a catalyst. Vasanthakumar and Karvembu [149] reported the catalytic evaluation of maghemite, obtained from river sand (Figure 7A), based on the base-free transfer hydrogenation of furfural, levulinic acid, and o-vanillin. River sand was collected from the Cauvery River in Kallanai, Tamil Nadu, India, and γ-Fe₂O₃ was separated using simple magnetic separation with a bar magnet. The γ-Fe₂O₃ powder was then washed with acetone, ethanol, and water to remove impurities and was dried at 80 °C for 3 h. To compare the catalytic activity of natural γ-Fe₂O₃, synthetic γ-Fe₂O₃ was also prepared following existing literature protocols. Catalytic tests were performed under a vacuum at 80 °C for 3 h. The activated catalyst was added to a sealed tube containing 1 mmol of substrate in 2-propanol and placed in a preheated oil bath at 150 °C for a specific duration. After 30–90 min, depending on the substrate, the catalyst was separated with an Nd magnet, and the solution was centrifuged. Magnetic characterization showed that the γ-Fe₂O₃ from river sand had higher saturation and remanent magnetization than synthetic γ-Fe₂O₃, exhibiting stronger magnetic behavior. Catalytic tests revealed that river sand γ-Fe₂O₃ displayed superior catalytic activity, which remained effective even after 10 cycles. The base-free catalytic transformation of levulinic acid to γ-valerolactone, via the 4-hydroxypentanoic acid intermediate, was attributed to the presence of silica in the γ-Fe₂O₃ from river sand.

Iron ore raw materials obtained through the catalytic reduction of NO with NH₃ resulted in excellent selective catalytic reduction (SCR) activity (above 80% at 170–350 °C) and N₂ selectivity (above 90% up to 250 °C) at low temperatures. However, the addition of H₂O and SO₂ in the feed gas showed some adverse effects on SCR activity.

For the preparation of catalysts, the iron ore raw materials were dried at 105 °C for 4 h and ground. Then, these ground samples were dried again at 105 °C for 5 h. Finally, they were calcinated at 250, 350, and 450 °C, respectively, for 6 h [152]. The characterization results indicated that the best catalytic performance was achieved by a catalyst comprised mainly of α-Fe₂O₃ and γ-Fe₂O₃. The utilization of natural maghemite as a catalyst presents challenges due to its inherent impurities. Natural maghemite typically contains substitutional metal cations such as Ti⁴⁺, Mg²⁺, or Al⁺³, among others, which can significantly influence its catalytic properties. Since γ-Fe₂O₃ is a metastable phase, the methods used to prepare catalysts based on this compound must be adequate to avoid the formation of the most thermodynamically stable hematite (α-Fe₂O₃) [153]. The summary results are presented in

Table 6. Summary of catalytic applications of less abundant and rare Fe-OH prepared from natural sources.

Catalytic Reaction	Catalyst	Mineral Source	Pretreatment Method	SBET (m2·g−1)	Reaction Conditions	Results of Catalytic Evaluation	Ref.
Transfer hydrogenation of furfural, levulinic acid, and o-vanillin	γ-Fe2O3	River sand	Magnetic separation, grinding, washing, and drying at 80 °C	22.3	Cat = 20 mg (7 mol %); T = 150 °C; 1 mmol of substrate in 2-propanol (3 mL); t = 30–90 min	γ-Fe2O3 recovered from the natural source exhibited superior activity compared to the synthetic counterpart under base-free conditions	[149]
Catalytic reduction of NO with NH3	α-Fe2O3/γ-Fe2O3	Iron ores	Drying, grinding, and calcination	22.8–42.5	NO = 500 ppm; NH3 = 500 ppm; O2 = 3 vol %; SO2 (when used) = 150 ppm; H2O (when used) = 5 vol%; N2 = 145 mL·min−1; GHSV = 20,000 h−1	Selectivity catalytic reduction activity above 80% at 170–350 °C and N2 selectivity (above 90% up to 250 °C) at low temperatures	[152]
Decomposition of hydrogen peroxide	Ferrihydrite	Acid mine drainage	Vacuum drying	270	Cat = 1 g·L−1; H2O2 = 0.02 M	Waste ferrihydrite shows the same catalytic activity for H2O2 decomposition as the commercial goethite-based catalyst	[150]
Photo-Fenton degradation of ENR and LEV	Wüstite	ND	ND	108.7	Cat = 10 mg·L−1; ENR = LEV = 0.05 mmol·L−1; H2O2 = 1.0 mmol·L−1; pH = 6.5; T = 35 °C; t = 180 min	100% and 55 % of antibiotic activity elimination in 180 min for ENR and LEV, respectively; complete antibiotic activity elimination for ENR in the next four recycling cycles	[151]

ND = No data.

.

3.2. Lepidocrocite (γ-FeOOH)

Lepidocrocite is named after its platy crystal shape and orange coloration [2]. It is commonly found in hydromorphic soils with alternating reducing and oxidizing conditions and can biogenically be obtained [154]. It is metastable, serving as a precursor to more magnetic phases like maghemite and magnetite [155], and plays a key role in the geochemical behavior of nutrients, heavy metals, and organic pollutants. It presents an orthorhombic crystal structure with lattice parameters a = 3.072(2), b = 12.516(3), c = 3.873(2) Å, and space group Cmcm [156]. In the lepidocrocite structure, ferric iron is coordinated in edge-shared octahedra, which form layers within [001] linked by hydrogen bridges. In contrast to other oxyhydroxides, lepidocrocite exhibits a low Neel temperature, which ranges between 50 and 70 K [157].

Up to the time of writing this article and to the best of the authors’ knowledge, lepidocrocite has only been used in its synthetic form in Fenton [158] and Fenton-like processes [159,160,161], photodegradation [162,163], water splitting [164], CO oxidation [154], arsenic adsorption [165], and the combination of H₂O₂ and PMS to produce ROS and remove chloramphenicol (CAP) [166]. However, no reports were found in the literature regarding its use in mineral form, likely because it is rare to find it in ore deposits [167]. The unique way to form lepidocrocite is the oxidation of Fe(OH)₂ by air in suspension at pH ≈ 7 [2]. When it is prepared in the presence of citric acid, the crystallization of lepidocrocite can be controlled by its concentration, which decreases with its increasing concentration [162]. Related to photocatalytic activity, visible light and the presence of ethylenediaminetetraacetic acid (EDTA) could accelerate the reaction rate of lepidocrocite [163].

Experimental laboratory results indicate that the morphology of catalysts influences the exposed facets and interaction with OH groups of iron oxides, having a direct effect on their reactivity. For example, the adsorption and degradation of Orange G (OG) on lath-like particles were mediated by the interaction with the µ-OH groups of the (010) facet, while in rod morphology, it was adsorbed laterally by the interactions with the µ-OH and µ₃-OH groups of the (010) and (001) facets [161].

When lepidocrocite VOOH hollow nanospheres were prepared by the hydrothermal method, a variation in the reaction temperature allowed the tuning of the surface area, modifying the catalytic activity. Also, it exhibited stability in alkaline media and achieved good performance in water splitting at a potential of 270 mV for OER [164].

The findings from the evaluation of the adsorption of arsenic suggest that pH affects it remarkably and the co-occurrence of aqueous Fe²⁺ and Fe³⁺ solids is beneficial for arsenic removal because Fe²⁺ can induce As³⁺ oxidation to As⁵⁺. Above pH = 8, As⁵⁺ can be more mobile than As³⁺; at low pH, the radical scavengers can compete with As³⁺ for the oxidants [165].

Theoretical results comparing goethite and lepidocrocite in the elimination of 5-bromosalicylic acid (BSA) via the Fenton process show that electrons are more easily transferred from the H₂O₂ molecule to the Fe atoms of α-FeOOH (goethite), enabling it to perform better in catalyzing the decomposition of H₂O₂. However, free radicals are more likely to desorb from γ-FeOOH (lepidocrocite), making the γ-FeOOH/H₂O₂ system more efficient in degrading BSA [168]. Lepidocrocite’s catalytic applications are versatile, but its metastability, crystal morphology, and interaction with functional groups affect its reactivity and performance.

3.3. Ferrihydrite (Fe_8.2O_8.5(OH)_7.4 + 3H₂O)

Ferrihydrite is a hydrated Fe(III) nano-oxide commonly found in various surface environments on Earth. Its presence has a significant impact on soil properties. Ferrihydrite was officially recognized as a mineral by the International Mineralogical Association (IMA) in 1975. In nature, it is commonly associated with silicates (with an atomic ratio of Si/Fe ranging from 0.17 to 0.37), and these silicates seem to directly influence its low crystallinity. Due to its small particle size and poor crystallinity, ferrihydrite exhibits crystal and compositional properties that are challenging to characterize [169]. The complexity of ferrihydrite’s structural analysis is also influenced by the presence of various other species in its composition, including 0.5–31.5 wt% SiO₂, 0.5–9.6 wt% Al₂O₃, and up to 7.2 wt% PO₄, along with As(V), which ranges from 0.7 to 28.3 wt% [170]. Initially, its formula was identified as 5Fe₂O₃ + 9H₂O, featuring a hexagonal unit cell with parameters of a = 0.508 nm and c = 0.94 nm and containing FeO₆ octahedra [171]. Eggleton and Fitzpatrick [172] proposed a trigonal cell structure for ferrihydrite, with two-thirds of the iron in octahedral coordination and one-third in tetrahedral coordination. However, other studies have questioned the presence of tetrahedral Fe [173]. Additional formulas such as Fe₅HO₈ + 4H₂O [174] and Fe₂O₃ + 2FeOOH + 6H₂O [175] have also been suggested. The number of X-ray diffraction peaks varies from two to five in natural ferrihydrites, while synthetic forms can show up to seven well-defined peaks, reflecting the degree of Fe atom ordering in the lattice.

Currently, one of the most accepted crystalline structures for two-line ferrihydrite is Fe_8.2O_8.5(OH)_7.4 + 3H₂O [176], which consists of 20% tetrahedral iron and approximately 80% octahedral iron, featuring a local δ-Keggin structure. Additionally, the structural depletion (SD) model proposed by Hiemstra in 2013 suggests a distinction between the nanoparticle’s core and surface. While the core is largely free of defects, the surface is depleted of tetrahedral structures. The ideal structure of ferrihydrite consists of three types of Fe: two hexacoordinated and one tetracoordinated, arranged in a δ-Keggin framework with proportions of 60%, 20%, and 20%, respectively. The general structure includes three layers of polyhedra, with the first hexacoordinated Fe located in the outer layers, the second hexacoordinated Fe in the middle (surrounding the tetracoordinated Fe), and the third octahedral Fe, which is highly distorted, located in the inner layer [177].

Ferrihydrite, owing to its widespread presence, common occurrence as a coating on particles, and reactive surface with a small particle size (1–7 nm), is widely recognized for its role in influencing the mobility of both inorganic and organic pollutants via sorption processes. It typically forms under near-neutral pH conditions in a variety of redox-active environments as well as in large sedimentary deposits where there is a significant supply of ferrous iron from groundwater and high oxidation rates [170]. Its chemical composition and crystal structure have been debated for over five decades, largely due to its extremely small particle size and the high degree of structural defects it incorporates [178]. Ab initio calculation results suggest that its reaction sites are mainly disposed along rows at the edges of sheets of iron octahedra. Molecular dynamics studies on nanoparticles of up to 10 nm show that highly reactive hydroxo groups on the surface are typically unbound, but they participate as hydrogen bond acceptors in a network with less reactive groups [179]. Magnetic properties reveal that ferrihydrite exhibits antiferromagnetism with a ferromagnetic-like moment at lower temperatures (100 K and 10 K), but it is paramagnetic at room temperature. As the crystallite size of ferrihydrite increases, both magnetization and coercivity decrease, a trend attributed to the pronounced surface effects found in fine-grained materials [180].

The main catalytic applications of ferrihydrite are related to the Fenton reaction [181], primarily conducted with synthetic ferrihydrite. In their review, Zhu et al. [181] highlighted that ferrihydrite consistently exhibits higher reactivity with H₂O₂ compared to other iron minerals. Specifically, they noted performance hierarchies, such as ferrihydrite > hematite > goethite, ferrihydrite > feroxyhyte ≈ magnetite > maghemite > goethite > hematite, and ferrihydrite > hematite > goethite. These comparisons underscore ferrihydrite’s superior catalytic potential. The use of ferrihydrite in the Fischer–Tropsch (FT) reaction has been studied by Kyu Seomoon [182]. By incorporating Al, Cu, and K into the synthesis of the catalyst and analyzing the FT reaction using gas chromatography–mass spectrometry (GC-MS) over 100 h, Seomoon observed a 25% increase in paraffin selectivity, attributed to reduced catalyst basicity. Increasing the H₂/CO ratio from 1 to 3 reduced CO₂ emissions by 40% and enhanced hydrocarbon formation. Similarly, Sumit Bali et al. [183] doped ferrihydrite with Al and Cu, achieving slightly better FT performance than a commercial fixed-bed catalyst, securing a patent for their synthesis method. However, there are some reports on the use of natural minerals as catalysts. The decomposition of hydrogen peroxide was tested using ferrihydrite obtained directly from acid mine drainage in the Zlaté Hory underground mine district (Czech Republic). In this case, the only further treatment was vacuum drying of the collected solid. The dried sample was gently powdered to minimize mechanical disruption (Figure 7B) [150]. Physicochemical characterization revealed a composition of mainly iron oxyhydroxides (approximately 96 wt%), with ferrihydrite as the principal phase, along with traces of goethite and possibly schwertmannite. The constant rate (6.59 min⁻¹) was slightly higher than that of commercial goethite, with the advantage of using extremely low-cost precursors or even industrial waste.

Ferrihydrite has been extensively used to adsorb heavy metals and other ions. The adsorption mechanisms for metals such as lead [184], copper [185], molybdenum [186], chromium [187,188], and zinc [189] occur via inner-sphere complexation, typically forming bidentate or bidentate–monodentate complexes depending on the pH and concentration. For radium and barium, tetradentate bonds with a single surface functional group are feasible [190]. Actinides like thorium [191], plutonium [192], and uranium [193] also prefer bidentate adsorption, with the latter two being extensively studied for removal from industrial systems through ferrihydrite nanoparticle precipitation.

Sorption capacity is strongly influenced by temperature and pH. Anion adsorption can significantly influence cation sorption by competing for available binding sites on the surface, thereby reducing the capacity for cation uptake. Carbonates alter ferrihydrite solubility [194], while phosphate addition enhances Cd²⁺ sorption by shifting from bidentate to ternary complex formation [195]. Zhu et al. [196] observed that ligands’ ability to inhibit arsenic (As³⁺ and As⁵⁺) sorption follows the sequence SeO₄ ≈ SO₄ < oxalic acid ≈ malic acid < citric acid < SeO₃ << PO₄. Additionally, the content of other metals in ferrihydrite’s lattice, such as aluminum, affects the sorption of chromates, selenates, and sulfates, as demonstrated by Johnston and Chrysochoou [187].

Maghemite, lepidocrocite, and ferrihydrite are considered less abundant oxides. Maghemite can only be distinguished from magnetite through characterizations that identify the oxidation states of iron. However, in the absence of such results, its catalytic activity may be mistakenly attributed to magnetite. Additionally, it is common to find traces of maghemite during hematite synthesis, making the evaluation of its performance as a catalyst a challenge.

Lepidocrocite, on the other hand, is rarely found in ore deposits as it is a metastable phase that serves as a precursor to other magnetic phases. Nevertheless, its synthesis can be achieved under controlled laboratory conditions by adjusting factors such as pH and the presence of additives. The morphology of the obtained particles is highly related to their catalytic activity.

As for ferrihydrite, it is found in nanometric sizes and, compared to other oxides and hydroxides, its recognition as a mineral is relatively recent. Its nanometric size and lack of crystallinity make it difficult to even determine its chemical formula. However, its capacity to adsorb heavy metals has been well established.

4. Rare Iron Oxides and Hydroxides

This section explores the catalytic applications of several rare iron oxides, including wüstite (FeO), akaganéite (β-FeOOH, feroxyhyte (δ′-FeOOH), and bernalite (Fe(OH)₃). The unique properties of these iron oxides and their potential catalytic applications will be discussed.

4.1. Wüstite (FeO)

Under ambient conditions, wüstite (FeO) crystallizes in a highly defective form of the ideal NaCl-type structure, with a nonstoichiometric formula of Fe_1−xO [197]. It is metastable and tends to decay into a two-phase mixture of α-Fe and magnetite Fe₃O₄. In addition, cubic Fe_1−xO can be synthesized at ambient pressure for iron deficiencies ranging from 0.05 < x < 0.15 at 560 °C [198,199]. The number of vacant octahedral sites is twice the amount of iron deficiency (x), according to the neutron diffraction data, implying a fraction approximately equal to x of iron located at interstitial tetrahedral sites. Further analysis of several samples reveals a direct linear relationship between the stoichiometry of Fe_1-xO and its lattice parameter, given by the equation a (Å) = 3.856 + 0.478(1 − x).

Wüstite exhibits antiferromagnetic ordering as temperatures descend below approximately 195 K. The nature of the magnetic order–disorder transition undergoes modulation in response to variations in the composition of Fe_1−xO, displaying heightened cooperativity in the presence of nearly stoichiometric FeO. The Fe²⁺ ion spins align antiferromagnetically along the (111) plane, with the predominant magnetic vector aligned parallel to the [111] direction. However, this axial component (3.8 µB for Fe_0.99O) is accompanied by a less perpendicular component (1.3 µB for Fe_0.99O). The magnitude of the ordered magnetic moment experiences a marked reduction in samples with decreased iron content. The paramagnetic phase assumes the structural attributes of a NaCl-type arrangement (space group $Fm\overline{3}m$), while the lower temperature phase exhibits a rhombohedral distortion (space group R3) [200].

Liu and colleagues [201] have conducted a review of the applications of wüstite in ammonia synthesis. The most successful ammonia synthesis method is the Haber–Bosch process, in which atmospheric nitrogen is converted into ammonia in the presence of a catalyst, high pressures, and high temperatures [202]. According to their findings, wüstite demonstrates ideal performance in terms of stability, activity, and cost. Its two industrial-level competitors are catalysts based on magnetite and ruthenium. The Ru/C catalyst can be up to 20 times more active than the magnetite-based catalyst (Haber catalyst), but in comparison to wüstite, it lacks competitiveness. Pernicone et al. [203] compared the performance of wüstite A301 to that of the best Ru/C catalyst [204], obtaining equivalent results. This positions wüstite as the most attractive option due to its abundance and cost-effectiveness. Its main challenges still lie in its production as wüstite does not exist naturally. In the laboratory, it is obtained by the decomposition of ferrous oxalate, resulting in a mixture with Fe₃O₄ and Fe that lacks the high-purity properties of wüstite. Industrially, it is obtained by reducing natural magnetite, using Fe as a reducing agent, directly in the smelting furnace at T > 834.5 K, incurring a significant energy cost.

Wüstite, with a bandgap of 2.1 eV (Figure 7C), has been utilized as a photocatalyst for removing fluoroquinolones from water [151]. When combined with H₂O₂, it functions as a photo-Fenton heterogeneous catalytic system, facilitating the generation of reactive oxygen species (ROS), including O₂^•−, ¹O₂, and ^•OH. The complete elimination of microbial activity is achieved within 3 h, accompanied by the mineralization of 21% of enrofloxacin. The catalyst’s performance proves reproducible for at least four treatment cycles following the recovery and drying of wüstite in an oven at 105 °C. The reduction of nitrates (NO₃⁻) is another application in which wüstite has been implicated [205]. Nearly stoichiometric conversion from NO₃⁻ to NH₄⁺ is attained, with NO₂⁻ identified as an intermediate. In this case, it is not a catalysis but rather a heterogeneous redox reaction as FeO serves as a reactant, forming Fe₃O₄ as a product. Advanced oxidation processes represent a field where iron oxides and hydroxides have been extensively applied [206]. In a study by Expósito and colleagues [207], the performance of wüstite was compared to that of other iron oxides in the electro-Fenton removal of aniline. Wüstite and magnetite, along with hematite, goethite, and FeSO₄, exhibited approximately 80% total organic carbon (TOC) removal at a rate similar to FeSO₄ (homogeneous phase). Furthermore, the decay pattern and final concentrations of H₂O₂ and Fe²⁺ also mirrored the homogeneous phase process. This suggests that the mechanism of the degradation of wüstite aligns with the homogeneous phase, initiated by dissolved iron ions. The advantage lies in the minerals’ ability to self-regulate the supply of iron ions to the solution based on the process’s requirements. Also, FeO catalysis based on the catalytic degradation of chloramphenicol (CAP) by water falling film dielectric barrier discharge (WFFDBD) enhanced the degradation from 69.5% to 97.7% in 8 min by just using WFFDBD. The reaction conditions are shown in Table 6. DFT studies of the catalytic mechanism showed that the key step in the catalytic reactions was the electron transfer from FeO to the antibonding orbitals of dissolved oxygen, H₂O₂, and O₃, which facilitated their transformation and decomposition. This process converted dissolved oxygen into O₂^•−, which was then adsorbed onto FeO through electron transfer, crucial to the catalytic reaction. Additionally, oxygen vacancies increased the charge density of Fe, enhancing electron transfer [208].

The metastable nature of wüstite, its tendency to decompose into magnetite and hematite, and the difficulty in obtaining high-purity forms are the main limitations to its widespread use. Here, wüstite is effective for ammonia synthesis. However, its effectiveness in AOPs is dependent on acidic conditions (pH~3).

4.2. Akaganéite (β-FeOOH)

Akaganeite (β-FeOOH) derives its name from the Akagané mine in Japan, where its initial discovery occurred in 1962 [209]. Widely acknowledged as a corrosion byproduct of steel [210], it stands as the predominant iron oxide in both soils and geothermal brines [211]. Showing a hollandite-like structure, its empirical formula is Fe_7.6Ni_0.4O_6.35(OH)_9.65Cl_1.25, discerned from the corrosion crust formed on the Camp del Cielo meteorite. Exhibiting monoclinic symmetry (space group I2/m), its lattice parameters are characterized by a = 10.600(2) Å, b = 3.0339(5) Å, c = 10.513(2) Å, and β = 90.24(2)°. Within akaganeite, Fe³⁺ ions occupy octahedral sites, forming a double-chain structure that shares corners, thereby allowing ample space for the inclusion of large tunnels featuring square cross-sections, suitable for hosting Cl and H atoms [212]. This compound is recognized as an antiferromagnetic material, manifesting a Neel temperature of 259 K, coupled with ferromagnetic-type hysteresis [213].

Akaganeite is particularly characterized by its ability to remove contaminants through adsorption. According to Zhao et al.’s review [214], the successful adsorption of As³⁺, As⁵⁺, Cr⁶⁺, Zn²⁺, U⁶⁺, Cd²⁺, Cs⁺, and PO₄³⁻ has been explored previously (for further details on the removal of these contaminants, please refer to that work and the references cited therein). Briefly, the morphology (size and shape) of akaganeite largely determines its properties. Nanocrystals with various shapes, such as rods, bundles, needles, fibrils, and cigars, have been obtained. These depend on the preparation method, usually involving hydrolysis and the use of different surfactant templates such as polymers, zeolites, activated carbon, and clays. Using ATR-FTIR and liquid chromatography, Marsac et al. [215] demonstrated that in the case of oxolinic acid removal, the removal occurs via the adsorption of the contaminant and not by oxidation on akaganeite. This adsorption occurred via metal complexes under slightly acidic pH and via hydrogen-bonded complexes under alkaline pH. They even determined that the most reactive crystal plane is (010).

Additionally, degradation applications have been found. In the laboratory, starting from FeCl₃, in the presence of glucose or the surfactant Brij30 at 60 °C for 3 days, Xiong et al. [216] determined that it can remove 100% of methyl orange in the presence of H₂O₂ and when pH = 4.5 using photo-Fenton reactions. This capacity is reduced to 86% after four working cycles. Polyaniline degrades acidic orange II even better than akaganeite alone [217] in Fenton processes. The β-FeOOH@GO catalyst, evaluated in a photo-Fenton-like system using UV irradiation, achieved 99.7% decolorization of MB after 60 min. The pseudo-first-order rate constant was 0.632 min⁻¹. The proposed degradation pathway of MB predominantly proceeded with the rupture of phenothiazine ring oxides with ^•OH, HO^2•−, and singlet oxygen (¹O₂) radicals [218].

4.3. Feroxyhyte (δ′-FeOOH)

It is a relatively uncommon iron oxide mineral and one of the few phases in the Fe₂O₃-H₂O system [219]. The feroxyhyte structure is composed of O²⁻ and OH⁻ ions arranged in the closest packing, with the Fe³⁺ ions statistically distributed in half of the octahedral positions. It has a hexagonal cell, with parameters a = 2.93 Å and c = 4.6 Å. It is completely converted to goethite in 6 h at a temperature of 60 °C and is partially transformed into hematite at 80 °C. Feroxyhyte is ferrimagnetic with a low Curie temperature (155 °C). The widely varying degree of the order leads to the formation of both magnetic and nonmagnetic varieties. The specific magnetic susceptibility reached 5000 × 10⁻⁶ cm³·g⁻¹ and only 400 × 10⁻⁶ cm³·g⁻¹ in the studied coarse crystalline and fine crystalline samples of synthetic feroxyhyte, respectively. Magnetic susceptibility was even lower in the finer crystals. In natural occurrences, it is always fine-grained, although samples with larger particle sizes and better crystallinity can be prepared in the laboratory [220].

Feroxyhite has demonstrated strong catalytic performance in various applications, although reported uses primarily involve laboratory-synthesized materials, often as nanoparticles. Among FeOOH polymorphs, feroxyhite stands out for its sufficient magnetic properties, enabling easy recovery in aqueous media. It can be synthesized through a single-step reaction—a critical factor for scalability in laboratory chemical processes [80]—and features a significantly higher proportion of surface OH groups compared to common magnetic iron oxides [221]. Recent studies highlight its use in photo-Fenton reactions, water splitting, and metal adsorption [222,223,224].

Weiping Du [225] studied the photodegradation of orange II using synthetic hematite, maghemite, magnetite, goethite, lepidocrocite, and feroxyhite. According to their data, hydroxides outperformed iron oxides, contrary to the authors’ expectations. Among hydroxides, feroxyhite showed the lowest performance. The mechanism involved dye adsorption, significantly enhanced by H₂O₂, AgNO₃, and NaF through different pathways. Pinto [226] utilized synthetic feroxyhite to decompose H₂O₂ into radicals via the Haber–Weiss mechanism, effectively degrading MB and IC. They demonstrated that the adsorption of anionic dyes is more effective than that of cationic dyes. Here, the degradation reaction occurred almost entirely on the surface, primarily through the generation of ^•OH radicals. Li et al. [227] developed δ-FeOOH/BiOBr_xI_1-x for the visible-light photodegradation of tetracycline, demonstrating improved charge transfer, synergy with H₂O₂, and enhanced recyclability. Radical species O₂^•− and ^•OH were identified as key players.

Chen Wang [228], in a review of FeOOH applications, highlighted that δ-FeOOH exhibits the smallest band gap (1.94 eV) and the largest specific surface area among FeOOH crystalline phases, making its band edge positions ideal for oxidizing H₂O to ^•OH radicals. Pereira [229] first applied synthetic feroxyhite for water splitting and H₂ production, citing its 2.2 eV band gap as being sufficient for water dissociation while absorbing visible light. The photocurrent generated was tenfold that of P25 (TiO₂). Da Silva Rocha et al. [230] demonstrated that Ni²⁺ doping in synthetic δ-FeOOH enhanced photocatalytic activity, conductivity, and charge transfer, while surface Ni(OH)₂ improved charge separation.

When modified with other materials, one of the most common applications of δ-FeOOH is the removal of heavy metals from water. Jing Hu [231] used δ-FeOOH-coated γ-Fe₂O₃ for Cr⁶⁺ adsorption via an outer-sphere complexation mechanism. In a Fe/Mn oxy-hydroxide (δ-Fe_0.76Mn_0.24OOH), Tresintsi et al. [232] found it to be highly efficient for As³⁺ and As⁵⁺ adsorption. Pinakidou [233] employed amorphous tetravalent manganese feroxyhyte nanoparticles (TMFx) to remove As⁵⁺. Later, Kokkinos et al. [234] reported TMFx as being effective in eliminating Cd, Hg, and Ni.

The versatility of these materials is also evident in their ability to immobilize phosphates with the aid of Fe²⁺, achieving up to 94% removal under acidic pH [221]. Additionally, they have biomedical applications, such as controlled heat release under an AC magnetic field [235]. Recently, Lacerda and coworkers proposed, using ab initio calculations, that feroxyhyte doped with niobium can function as a bifunctional catalyst. Their results suggest that doping δ-FeOOH layers leads to stronger van der Waals interactions, enhancing the catalyst’s thermal stability. Additionally, the incorporation of Nb enhances Brönsted acid characteristics and introduces Lewis acidic sites on the catalyst’s surface [236].

4.4. Bernalite (Fe(OH)₃)

Bernalite, Fe(OH)₃, is one of the more recently identified iron oxyhydroxide minerals, recognized in 1992. It is the scarcest iron mineral and has been documented in only a handful of locations [4], including New South Wales, Australia, and the central Black Forest, Germany [237]. McCammon et al. deduced [238] its empirical formula to be [(H₂O)_0.04(CO₂)_0.03Pb_0.01]Fe³⁺_0.93Si_0.06Zn_0.01)(OH)_2.95O_0.04. This mineral exhibits a protonated octahedral framework (POF), adopting a distorted perovskite-type structure (ReO₃) within the Pmmn space group [239]. Its tilt system of a + b + c⁻ arises from the topology of hydrogen bonds, resulting in alternating layers of Fe(1) and Fe(2) octahedra with an inverted tilt along the [001] direction, n-glides passing through the Fe atoms, and notable anisotropy in hydrogen bond topology [239]. With a slight deviation from the collinear antiferromagnet, bernalite manifests a weak ferromagnetic moment [240]. To date, there exists scant specific information regarding its application as a catalyst, neither within Scopus nor Google Scholar. Scopus records a mere seven references with “bernalite” in the title, with only two in the last two decades, the remainder hailing from the 1990s. Of these, four discuss structural or magnetic properties.

Recently, a potential use for this mineral was serendipitously uncovered. Under conditions of pH = 10, in the presence of arsenates, goethite undergoes partial transformation into bernalite. Remarkably, bernalite exhibits 2.18 times greater arsenate sorption than goethite, achieved through the formation of monodentate mononuclear (MM) complexes. The goethite used in these studies was sourced from rod-shaped nano-goethite (US3162) crystals obtained from US Research Nanomaterials (USA) [241] or Sigma-Aldrich (St. Louis, MO, USA) [242]. In another possible application, a composite comprising bernalite, Fe₂O₃, and Ru was synthesized through microwave-assisted heating in deionized water. This resultant material demonstrates the potential for oxygen reduction and hydrogen oxidation reactions. In spite of that, discerning the catalytic role of each component remains challenging [243].

Wüstite, akaganéite, feroxyhyte, and bernalite are among the least abundant iron oxyhydroxides, limiting their use in large-scale applications due to the rarity of natural sources. Their catalytic applications are less explored compared to more common iron oxides. Nevertheless, these materials exhibit properties suitable for advanced oxidation processes (AOPs), particularly in acidic conditions and systems like Fenton, photo-Fenton, and electro-Fenton reactions. Their optical band gaps, i.e., 1–2.1 eV for wüstite [151,198], 1.94 eV for feroxyhyte, and 2.1 eV for akaganéite [244], make them suitable for solar light-driven photocatalysis. Excluding wüstite, their crystal structure and chemical composition provide a high affinity for adsorbing heavy metals and organic pollutants. Notable differences include akaganéite’s superior chemical stability and wüstite’s potential in ammonia synthesis. Bernalite remains underexplored due to its rarity, although recent studies hint at promising environmental and electrochemical applications.

Production challenges include synthesis complexity and stability, especially for large-scale use. With magnetic properties, feroxyhyte is easier to handle in aqueous environments, while wüstite’s instability and bernalite’s scarcity hinder broader adoption. These contrasts highlight the need for innovative applications that optimize cost-effectiveness.

5. Future Directions and Perspectives on Catalytic Systems

The reaction mechanism of Fe-OH catalysts from natural sources depends on the evaluated reaction, conditions, and impurities. For example, in photocatalytic processes with Fe₂O₃ over gelcasting porous ceramic [78] (cf. Figure 8A), UV-induced electron excitation produces ROS that attack phenolic compounds. In contrast, persulfate (PS) activation (cf. Figure 8D), used for the catalytic degradation of 2,4-D and MCPA [73], occurs via Fe²⁺-generating SO₄^•− radicals both in homogeneous and heterogeneous phases. This dual-phase reaction provides enhanced degradation efficiency, with higher pH values compared to the more pH-dependent process photocatalytic, Fenton, and photo-Fenton processes.

The mechanism of the catalytic ozonation process mediated by HRS [81] (Figure 8B) involves three main pathways for the degradation of acetaminophen (ACT): (i) the direct oxidation of ACT by dissolved ozone (O₃) molecules in the bulk solution; (ii) the radical-based degradation of ACT by hydroxyl radicals (^•OH) generated in the solution, and (iii) the indirect oxidation through the interaction of O₃ with the catalyst surface, which produces ^•OH and superoxide radicals (O₂^•−). The presence of ions like phosphate and carbonate can hinder the process by occupying active sites on the catalyst or reacting with ^•OH. In turn, two key mechanisms drive SSZ ozonation using goethite as a catalyst [129] (cf. Figure 8C): (i) the reductive degradation on the catalyst surface and (ii) the direct oxidation by non-adsorbed O₂^•− radicals. Additionally, plasma treatment on α-FeOOH enhances the production of ^•OH and O₂^•−, improving the ozonation process.

Figure 8E shows the mechanism of the photocatalytic reduction of 4-NP over Cu/α-Fe₂O₃ [77]. Here, we can observe that this follows three main steps: (i) the diffusion and adsorption of 4-NP onto the metal catalyst surface, (ii) electron transfer from BH₄⁻ to 4-NP, mediated by Fe and Cu metal particles, and (iii) the reduction of the nitro group by hydride transfer from a metal–hydride complex. This leads to the formation of 4-hydroxylaminophenol as a stable intermediate, which undergoes three hydro-deoxygenation reactions to produce 4-AP. Finally, 4-AP desorbs from the catalyst surface.

In transfer hydrogenation (cf. Figure 8F), the γ-Fe₂O₃ catalyst facilitates the initial activation of levulinic acid (LA) by interacting with its carbonyl carbon [149]. This interaction enhances the reactivity of LA, allowing it to interact with 2-propanol, which donates hydrogen to form 4-hydroxypentanoic acid as an intermediate. The iron catalyst promotes hydrogen transfer. For example, 4-hydroxypentanoic acid undergoes dehydration, which leads to the formation of Gamma-valerolactone (GVL). The Fe-OH minerals from natural sources can act as catalysts for biodiesel production. The mechanism proposed in Figure 8G [87] initiates with the adsorption of the carbonyl group in free fatty acids (FFAs) onto the acidic site and the adsorption of methanol onto the basic site of the catalyst. These interactions generate a carbocation and an oxygen anion, respectively. A nucleophilic attack occurs; the carbocation reacts with the hydroxyl group in glycerol on the acidic site, while the oxygen anion in methanol reacts with the carbonyl group in triglycerides on the basic site. This leads to the transesterification reaction, breaking down the three fatty acid chains of the triglyceride, yielding biodiesel molecules and glycerol as a byproduct. Simultaneously, the esterification of glycerol with FFAs on the acidic site produces water and diglycerides. Finally, the reaction products desorb, and the catalytic cycle restarts.

The separation of Fe-OH from other minerals and their purification can be difficult to achieve due to their textural intergrowths. However, the impurities present in Fe-OH catalysts can either inhibit or enhance catalytic reactions. There are few studies comparing synthetic catalysts to those prepared from natural sources. Based on existing comparative results [1,149], Fe-OHs from natural sources are more effective than their synthetic counterparts, but their catalytic activity is much more complex. For example, catalysts from different sources of mineral Fe₃O₄ with similar impurity content showed marked differences in the degradation of p-NP, indicating that the purity of the mineral might not be the determining factor for its catalytic activity. Here, the key factor in the performance is related to the surface site density of the parameter.

Maghemite prepared from river sand exhibited higher biomass conversion by transfer hydrogenation [149]. This was attributed to its crystallinity and higher pore size, allowing more interaction with the reactants. The presence of impurities of silica and metal carbonates also promotes the Lewis acidic sites. Therefore, one of the challenges for the use of these types of catalysts is understanding the effect of impurities on catalytic activity. A greater number of comparative studies between synthetic catalysts and those prepared from natural sources are needed, focusing on both their physicochemical characteristics and catalytic performances.

On the other hand, catalytic evaluation requires an appropriate experimental design as magnetic stirring could cause poor dispersion of particles in the solution, while mechanical stirring could hinder the scaling up of reactors to pilot or large-scale levels. Figure 9 shows different catalytic systems where Fe-OH prepared from natural sources were used as catalysts. At the laboratory scale, magnetic stirring (Figure 9C,F) can be used when there is a decrease in the magnetic response of catalysts. However, the use of pumps (Figure 9A,B,F) or gas carriers (Figure 9D,E) allows the continuous flow of solutions passing through the catalyst bed. To facilitate real-world scaling, we propose adapting and optimizing proven experimental reactor designs from analogous catalytic processes [80,245,246].

The scalability of catalysts largely depends on their stability, lifespan, and resistance to poisoning. One of the most remarkable characteristics of Fe-OH is their ability to be recovered using magnetic fields, enabling their reuse. Figure 10 depicts conversion versus the number of cycles reached by natural Fe-OH. There is more information available on the reuse cycles of the most abundant iron oxide (magnetite (Figure 10A) and hematite (Figure 10B)) and hydroxide catalysts (goethite, Figure 10C). In contrast, information about less abundant and rare Fe-OH (maghemite and wüstite, Figure 10D) is rarely found.

According to reported results, magnetite prepared from natural sources is a stable catalyst in the photo-Fenton degradation of cefotaxime [46], CWPO degradation of phenol [62], and Fenton-like degradation of MeP reactions [52], showing less than 10% of deactivation until after five reused cycles. The primary cause of deactivation appears to be the partial oxidation of the catalyst surface, leading to a progressive decrease in the Fe²⁺/Fe³⁺ ratio [57]. Nevertheless, to understand the decrease in reaction performance, it is necessary to characterize the catalysts after being recovered from the reaction. Comparisons between the physicochemical properties before and after the catalytic performance would help elucidate the mechanism of catalyst deactivation. The reuse results of hematite indicate that its conversion capacity decreases only slightly in the photocatalytic reduction of 4-NP to 4-AP [77], the ozonation of ACT [81], and the photo-Fenton degradation of phenol [74]. It also demonstrates stability in the cracking of toluene [89], where the reaction conditions are more aggressive than in AOPs. In this case, the catalysts’ deactivation is attributed to carbon deposits, the aggregation of active metal particles, and poisonous impurities. This decline could be due to (i) the loss of Fe mass and/or iron leaching from the surface after each usage cycle, (ii) the reduction in the activation rate of PS molecules by the catalyst, and, as a result, (iii) the reduction in the efficiency of the heterogeneous catalytic degradation process. However, the iron loss was minimal, which confirmed the stability of the catalyst. Further characterization is needed to better understand the deactivation mechanism.

Figure 10. Conversion vs. recycled use of iron oxides and hydroxides: (A) magnetite, (B) hematite, (C) goethite, and (D) maghemite and wüstite. Elaborated with data from References [46,49,51,52,55,57,62], [72,73,74,77,81,89], [122,123,129], and [149,151], respectively. Dash lines are guides for the eye.

Figure 10. Conversion vs. recycled use of iron oxides and hydroxides: (A) magnetite, (B) hematite, (C) goethite, and (D) maghemite and wüstite. Elaborated with data from References [46,49,51,52,55,57,62], [72,73,74,77,81,89], [122,123,129], and [149,151], respectively. Dash lines are guides for the eye.

The decrease in catalytic conversion mediated by goethite is more evident (cf. Figure 10C) compared to magnetite and hematite, being more stable on the ozonation of SSZ using plasma-treated goethite (shown by pink triangles) [129]. Natural goethite (shown by blue triangles) interestingly exhibited a higher drop in conversion. Using plasma treatment, the surface enhancement of α-FeOOH can increase the density of hydroxyl groups at the surface.

Related to less abundant and rare iron oxides, the scant available results about reuse indicated that they are highly stable during transfer hydrogenation [149] and the photo-Fenton degradation of ENR [151]. The recovered catalyst (maghemite) from transfer hydrogenation remained stable and did not lose crystallinity, phase purity, and activity, even after extensive usage. Meanwhile, the surface area of the goethite decreased compared to that of the fresh catalyst [149]. This is consistent with the discussion throughout this review, which suggests that surface area is not the determining factor in the activity of these catalysts. In the photo-Fenton degradation of ENR, the reuse of wüstite resulted in the release of iron, but after the first cycle, iron leaching decreases, and the macroscopic solid undergoes disaggregation. Additionally, the first cycle shows a reduction in catalyst particle size. Furthermore, the observed disaggregation and particle size reduction suggest alterations to the solid surface throughout reuse cycles, potentially impacting the catalyst’s interaction with light and H₂O₂ [151].

Generally, the reuse of Fe-OH catalysts does not require thermal treatments to reactivate the recovered catalysts, which reduces energy consumption and preparation costs. Additionally, the fact that they can be pretreated (or used directly without the need for thermal or chemical treatments) from inexpensive, accessible, and widely available sources may contribute to their scalability to real-world applications. One of the challenges to overcome in AOPs is the need for acidic pH conditions for these catalysts to remain active. One potential approach is exploring the use of oxidizing agents other than H₂O₂, such as PS, which has been shown to require nearly neutral pH conditions for effective catalytic reactions. In catalytic cracking, high temperatures are required, which can lead to the excessive consumption of network oxygen, causing catalyst deactivation. A thermal treatment in an oxidizing atmosphere could help reactivate it. However, this would increase energy consumption, potentially hindering its large-scale application.

Natural Fe-OH catalysts present challenges that must be addressed to enable their practical application. Their composition and mineralogical structure can vary depending on the source, affecting catalytic activity and reproducibility. Also, impurities may either inhibit or enhance catalytic performance, making it crucial to understand their role. In addition, AOP processes usually require acidic pH conditions, which limit natural Fe-OH applications in environmental samples or industrial waters. Another key challenge is the stability and reusability of these catalysts as iron leaching, surface fouling, or structural changes can lead to deactivation over multiple cycles. Furthermore, natural catalysts are widely available and cost-effective, but their efficiency must be comparable to synthetic catalysts to justify large-scale implementation without excessive energy-consuming treatments. Proper experimental design is also essential since factors such as stirring methods can impact catalyst dispersion and overall performance. To overcome these challenges, more comparative studies between natural and synthetic catalysts are needed, along with strategies to enhance stability, efficiency, and operational conditions.

6. Conclusions

While iron oxides are abundantly available in the Earth’s crust, their utilization as catalysts has been largely overshadowed by the prevalence of laboratory-synthesized materials. A key advantage of iron oxide catalysts lies in their magnetic properties, facilitating easy recovery and potential for reuse. However, crucial knowledge gaps remain. First, the impact of repeated reuse cycles on the magnetic properties of these catalysts and their subsequent influence on catalytic performance requires further investigation. Second, a comprehensive comparative cost analysis between the synthesis of iron oxide catalysts and their extraction and preparation from natural sources is currently lacking. The main applications of iron oxides are AOPs, electrochemical applications (water splitting and OER), FTS, and ammonia synthesis. Hematite (α-Fe₂O₃) is the most widely used naturally sourced iron oxide, probably because it usually represents the final stage in the transformation of other iron oxides. However, AOPs are operative only in a narrow interval of pH compared to other iron oxides. The next most used iron oxides are magnetite (Fe₃O₄) and goethite (α-FeOOH). The catalytic activity of magnetite has been related to the presence of Fe³⁺ and Fe²⁺ in its structure, where the oxidizing conditions can transform it into hematite. On the other hand, goethite can participate in redox reactions through the interconversion of Fe³⁺; and Fe²⁺, facilitating electron transfer in catalytic processes. This results in reactions where reactive species such as hydroxyl radicals (^•OH) are necessary. Additionally, goethite exhibits the highest specific surface area among the three most used iron oxides.

Although conventional catalyst preparation often involves acid purification and thermal treatments, simpler approaches utilizing magnetic separation, grinding, sieving, and milling have demonstrated promising catalytic activity in some evaluations. Despite that, these methods may have limitations when precise control over particle size and morphology is critical as achieving consistent shape and size distribution through these techniques can be challenging.

The preparation and evaluation of catalytic applications of less abundant natural iron oxides, such as maghemite (γ-Fe₂O₃), lepidocrocite (γ-FeOOH), and ferrihydrite (5Fe₂O₃⋅9H₂O), face diverse challenges, as follows: natural maghemite is not readily available in a pure form, and lepidocrocite is rarely found in natural ore deposits, limiting its accessibility for catalytic applications. On the one hand, ferrihydrite exhibits poor crystallinity, a small particle size, and structural defects that make its physicochemical characterization and reproducibility in catalytic processes difficult. On the other hand, the so-called rare iron oxides, i.e., wüstite (FeO), akaganeite (β-FeOOH), feroxyhyte (δ′-FeOOH), and bernalite (Fe(OH)₃), have been employed in their synthetic forms in few catalytic reactions. This may be related to their metastability, synthesis complexity, and dependence on specific operational conditions.

Despite the potential complexities associated with the preparation and evaluation of catalysts derived from natural sources, their abundance on the Earth’s surface, coupled with their unique physicochemical properties and low-cost availability, presents a compelling opportunity to address environmental and energy-related challenges. Regardless, optimizing their catalytic performance necessitates a thorough understanding of the specific characteristics of each iron oxide to enable the development of effective strategies for scaling up these processes to pilot and industrial levels.