Removal of Iron(II) as Magnetite from Acid Mine Water

Abstract

The High-Density Sludge (HDS) process is widely used for the treatment of acid mine water as it produces a sludge of high density. The aim of this study was the development of a process where iron in mine water can be removed as magnetite, to assist with rapid settling of sludge. It was concluded that Fe²⁺ can be removed as Fe₃O₄ (magnetite) by forming Fe(OH)₂ and Fe(OH)₃ in the mole ratio of 1:2. Magnetite can form in the absence or presence of gypsum. The settling rate of magnetite-rich sludge is substantially faster than that of ferric hydroxide-rich sludge. It is recommended that further studies be carried out on the separation of magnetite gypsum through magnetic separation.

1. Background

Acid mine drainage (AMD) is one of the major environmental problems in the mining industry. AMD is an acidic, metal-rich effluent that originates from the oxidative dissolution of metal sulphides in mines. When these metal sulphides are exposed to air and water, they can react to form sulphuric acid, which further accelerates the dissolution of metal sulphides and may release toxic metals. AMD is typically characterised by low pH values and high concentrations of heavy metals, which can cause serious harm to aquatic ecosystems. Iron (II) is one of the most common metals found in AMD. Iron (II) can easily be oxidised to form a precipitate of iron (III) hydroxide, which has a characteristic reddish orange colour and is known as “yellow boy” [1].

The detrimental impact of AMD on the ecosystem and infrastructure has been widely studied because of critical challenges that lie in the durability of materials exposed to acid environments. Sarmiento et al. [2] showed that AISI 1020 carbon steel and AW6060 aluminium are significantly compromised when exposed to acidic environments, and the article highlighted that both materials experience marked reductions in mechanical properties, including tensile strength and fatigue resistance. For instance, in a dynamic scenario with erosion–corrosion, the fatigue strength of carbon steel can diminish by nearly 50% after just a few days of exposure. This degradation not only affects the structural integrity of components used in mining and construction but also poses safety risks in applications like aircraft manufacturing, where material reliability is crucial [2]. Therefore, it is essential to address and remove these heavy metals from acidic environments to prevent corrosion-related failures, enhance material longevity, and minimise environmental impacts associated with metal leaching and pollution.

In the High-Density Sludge (HDS) process [3], Fe²⁺ is removed by raising the pH with Ca(OH)₂ to 7.2, where the rate of Fe²⁺-oxidation is fast, and Fe³⁺ precipitates as Fe(OH)₃. Magnetite can be recovered from iron-rich acid mine water by using Na₂CO₃ [4]. Fe(OH)₃ can be recovered separately from gypsum by dosing scale inhibitors to keep gypsum in a solution for the period needed for sludge separation [5].

The overall chemical reaction for the formation of Fe₃O₄ in an aqueous medium involves the following reactants:

Dissolution of iron salts

FeSO₄ → Fe²⁺ + SO₄²⁻

(1)

Fe₂(SO₄)₃ → 2Fe³⁺ + 3SO₄²⁻

(2)

Dissolution of CaCO₃ under acidic conditions

CaCO₃ + H₂SO₄ → CaSO₄ + CO₂ + H₂O

(3)

This reaction provides a source of CO₂, which escapes as a gas under acidic conditions, and OH⁻, which increases the pH of the solution.

Formation of Fe₃O₄

Under the right pH conditions, Fe²⁺ and Fe³⁺ ions react with hydroxide ions (OH) in water to form Fe₃O_4.

The overall reaction is as follows:

Fe²⁺ + 2Fe³⁺ + 8OH⁻ → Fe₃O₄ + 4H₂O

(4)

Fe₃O₄ is a mixed-valence iron oxide with a chemical formula that can be represented as Fe[Fe₂]O₄, indicating the presence of both Fe(ll) and Fe(lll) ions in its structure. The formation of Fe₃O₄ through co-precipitation involves the simultaneous presence and precipitation of Fe(ll) and Fe(lll) ions, which ideally requires a 1:2 molar ratio of Fe²⁺ to Fe³⁺ [6].

The objective of this study was to remove Fe²⁺ as Fe₃O₄ (magnetite) through oxidation with Fe³⁺. The following steps were followed: (i) Determine the optimum Fe³⁺/Fe²⁺ ratio. (ii) Determine the required alkali dosage. (iii) Select the most cost-effective alkali combination. (iv) Possible separation of magnetite from gypsum. The alkalis that can be used are CaCO₃, Ca(OH)₂, Na₂CO₃ and NaOH. The calcium alkalis offer the benefit of partial desalination, which is achieved through gypsum crystallisation, while sodium alkalis result in water with high Total Dissolved Solids (TDS). The benefit of sodium alkalis is that only magnetite will form, not gypsum.

2. Literature Review

2.1. Introduction

Acid mine drainage (AMD) is rich in metals such as Fe²⁺/Fe³⁺, Al³⁺, Mn²⁺ and SO₄²⁻ among other trace metals, and as such, mine water is rendered as a promising source of Fe-based minerals and sulphate. These elements can be recovered as valuable products for different applications. Magnetite (Fe₃O₄) is one of the commercially valuable minerals that can be synthesised by using Fe-rich salts [7].

2.2. Iron in Mine Water

2.2.1. Formation of Pyrite

Fool’s gold, or pyrite (FeS₂), is a common sulphide mineral that is found in large quantities and is essential for the occurrence of acid mine drainage (AMD). A basic geochemical process known as pyrite oxidation takes place when oxygen and water are introduced to rocks that contain pyrite; this usually happens because of mining operations. Sulphuric acid (H₂SO₄) and dissolved iron species are produced because of this process, which has a negative effect on the environment and water quality [8].

Pyrite is oxidised by a sequence of complex chemical reactions that result in the formation of soluble and insoluble compounds from solid pyrite. The following equation can be used to summarise the entire reaction: FeS₂ + 7O₂ + 2H₂O → 2FeSO₄ + 2H₂SO₄.

2.2.2. Treatment Methods

Acid mine water can be effectively treated using a variety of alkalis to recover magnetite. Alkalis like lime (Ca(OH)₂), limestone (CaCO₃), calcium oxide (CaO), hydrated lime (Ca(OH)₂), soda ash (Na₂CO₃) and caustic soda (NaOH) have been used to treat acid mine drainage (AMD) with promising results. These alkalis help neutralise acidic water and promote the precipitation of metals, including magnetite. By increasing the pH of the acid mine water, these alkalis help remove metals by adsorption, co-precipitation and precipitation processes [9,10].

In 1970, the Bethlehem Steel Corporation (Bethlehem, PA, USA) pioneered a way to handle effluent, known as the High-Density Sludge (HDS) process. This practice entailed combining a portion of solidified sludge materials with the discharged AMD that had been treated and adding limescale to the blended mixture. As a result of this technique, a denser sludge took form [3]. Operations under the HDS process included three stages, namely (i) pH correction/sludge monitoring, (ii) neutralisation/aeration and (iii) solid/liquid separation [11] (Figure 1).

2.3. Magnetite

The most prevalent magnetic mineral on Earth is magnetite (Fe₃O₄, also known as iron (II, III) oxide), which is present in sedimentary, metamorphic and igneous rocks. It is an inversed, black spinel with a well-formed octahedral crystal structure and crystalline structure. It is the most magnetic of all the natural iron oxides. An early type of magnetic compass was made of magnetite, which is found in some locations as naturally magnetised stone known as lodestone. Numerous sedimentary rocks contain magnetite, and massive amounts have been discovered in banded iron formations. Additionally, practically all igneous and metamorphic rocks include this mineral, particularly when it is present in the form of tiny grains. Magnetite- and ilmenite-rich grains that precipitate together from magma are found in many igneous rocks [13].

The ability of magnetite to be magnetic makes it the most magnetic mineral on Earth. The type and concentration of trace elements in magnetite depend on factors such as temperature, physicochemical properties, silica activity, and melt composition. Therefore, magnetite is an iron oxide composed of partially oxidised iron, with its metal component primarily made up of iron [13].

2.3.1. Chemical Reaction and Mechanisms

Iron (Fe²⁺) removal is a critical step in water treatment processes, especially in the context of AMD and industrial wastewater. The removal of Fe²⁺ can be effectively achieved through the formation of magnetite (Fe₃O₄), a process that involves oxidation, precipitation and crystallisation reactions. Understanding the underlying chemical mechanisms is essential for optimising this removal process, ensuring its efficiency and minimising the environmental impact of iron-laden effluent [14,15].

The initial step in the formation of magnetite involves the oxidation of Fe²⁺ to Fe³⁺. This reaction typically occurs in the presence of dissolved oxygen or other oxidising agents, such as hydrogen peroxide. The reaction is pH-dependent, with higher pH values favouring the oxidation process [8]. In aqueous solutions, Fe²⁺ oxidises to Fe³⁺ via the following reaction:

4Fe²⁺ + O₂ + 6H₂O → 4Fe(OH)₃ + 8H⁺

(5)

The formation of Fe(OH)₃ as a precipitate is crucial as it eventually leads to the formation of magnetite through further reactions. The rate of oxidation is influenced by factors such as temperature, pH and the presence of catalytic species. Once Fe³⁺ is formed, the next step involves the nucleation and growth of magnetite. This occurs via the interaction of Fe³⁺ and Fe²⁺ ions in a solution, typically in an alkaline environment. The reaction can be represented as follows [16]:

Fe²⁺ + 2Fe³⁺ + 8OH⁻ → Fe₃O₄ + 4H₂O

(6)

This process is highly sensitive to pH, with the optimal formation of magnetite occurring around pH 9–11. The co-precipitation of Fe²⁺ and Fe³⁺ in this alkaline medium leads to the formation of Fe₃O₄ nanoparticles, which aggregate into larger magnetite particles. The kinetics of nucleation and growth depend on supersaturation levels and the presence of impurities that can either promote or inhibit the process [17].

2.3.2. Parameters Affecting Magnetite Formation

The formation of magnetite from Fe²⁺ and Fe³⁺ is highly dependent on various reaction conditions, including temperature, Fe²⁺ concentration, pH level and the oxidising agent.

pH Level

The pH of the solution plays a crucial role in the formation of magnetite (Fe₃O₄). Optimal magnetite formation generally occurs in an alkaline environment, with pH levels ranging from 9 to 11. At this pH, the co-precipitation of Fe²⁺ and Fe³⁺ leads to the formation of magnetite nanoparticles. The high pH facilitates the conversion of Fe²⁺ to Fe³⁺ and promotes the precipitation of Fe₃O₄. Deviations from this pH range can result in the formation of other iron oxides, such as hematite or goethite, instead of magnetite. This sensitivity to pH is attributed to the solubility of iron species and the formation kinetics in the magnetite phase [18,19].

Fe²⁺ Concentration

In magnetite (Fe₃O₄) formation, the concentration of Fe²⁺ plays a crucial role. Magnetite consists of both Fe²⁺ and Fe³⁺ ions, and its formation typically occurs through the oxidation of Fe²⁺ in aqueous solutions. A high concentration of Fe²⁺ can promote the formation of magnetite by providing a sufficient source of iron that can undergo oxidation and ultimately form the mixed-valence structure characteristic of magnetite. However, if the Fe²⁺ concentration is too high relative to Fe³⁺, it can lead to the formation of other iron oxide phases like hematite (Fe₂O₃) and goethite (HFeO₂) instead of magnetite. Therefore, the Fe²⁺/Fe³⁺ ratio must be carefully controlled to favour the formation of magnetite over other iron oxide minerals [15,20].

Temperature

Temperature affects both the kinetics of magnetite formation and the characteristics of the nanoparticles. Generally, higher temperatures accelerate the nucleation and growth of magnetite, leading to faster reactions and potentially larger particles. However, excessively high temperatures can also cause the formation of unwanted by-products or lead to the aggregation of nanoparticles [21]. Temperature control is essential to optimise the size and distribution of magnetite particles, as well as to maintain consistent reaction conditions. Studies have shown that moderate temperatures typically result in better control over the size and morphology of magnetite nanoparticles [20].

Uses

Other research developments looked at the application of magnetite in other important areas, such as magnetic resonance imaging, ferro-fluids for audio speakers, magnetic targeted drug delivery, magnetic recording media and as an adsorbent in the water treatment industry.

Magnetite is used in magnetic resonance imaging (MRI) applications due to its high magnetic moment and superparamagnetic properties. Magnetite-based nano-systems can be designed to enhance the contrast in MRI images, allowing for better visualisation of tumors and other tissues. For example, magnetite nanoparticles can be functionalised with chitosan and graphene oxide to enhance their magnetic properties and improve their ability to accumulate in tumors. This allows for more accurate imaging and targeted treatment of cancer [22,23].

Nanoparticles are widely used in water treatment due to their ease of separation and reusability. They can be used as adsorbents to remove heavy metals, dyes and other contaminants from wastewater. Magnetite-based composites, such as magnetite silica nanocomposites and magnetite–graphene nanocomposites, are particularly effective in water purification. These composites can be designed to enhance the adsorption capacity and stability of magnetite, making them more effective in removing pollutants from water [24].

They are also applied in magnetic targeted drug delivery systems to enhance the delivery of anticancer drugs to specific sites in the body. Magnetite nanoparticles can be functionalised with various molecules to enhance their targeting capabilities and improve their ability to accumulate in tumors. For example, magnetite nanoparticles can be coated with polyethylene glycol and dipalmitoyl phosphatidylcholine to enhance their stability and targeting efficiency. This allows for more effective delivery of drugs to cancer cells and reduced systemic toxicity [25].

Magnetite is used in ferrofluids, which are colloidal suspensions of magnetite nanoparticles in a carrier fluid. Ferrofluids exhibit unique properties such as magnetisation and self-healing, making them useful in various applications such as magnetic storage, magnetic separation and magnetic propulsion. Magnetite ferrofluids can be designed to have specific properties such as high magnetisation, low viscosity and high stability, making them useful in a wide range of applications [24].

2.3.3. Formation

The formation of magnetite (Fe₃O₄) nanoparticles has drawn a lot of attention in the last few decades because of its potent magnetic characteristics and compatibility with biological systems. They are perfect for a variety of biomedical applications because of their qualities. There are up to sixteen distinct types of iron oxide, hydroxide and oxyhydroxide, which makes it difficult to synthesise iron oxide nanoparticles. However, certain methods like hydrothermal synthesis, microemulsion, thermal breakdown and coprecipitation can be used to effectively synthesise magnetite nanoparticles, which have an inverted spinel structure [26].

Magnetite nanoparticles are produced using a variety of techniques, the most common of which is chemical precipitation, which provides reasonably accurate control over the size and structure of the particles. The synthesis of magnetite by chemical precipitation involves many techniques: (1) the precipitation of Fe³⁺ and Fe²⁺ ions simultaneously, (2) partial reduction of Fe³⁺ ions, and (3) partial oxidation of Fe²⁺ ions, succeeded by simultaneous precipitation [27].

When a base is present, the aqueous coprecipitation of Fe³⁺ and Fe²⁺ ions at a 2:1 ratio to produce magnetite nanoparticles usually takes place in an anaerobic environment with a pH range of 9–12. Interestingly, the coprecipitation approach does not require precursor complexes and does not produce any hazardous intermediates or solvents. Moreover, it functions at temperatures lower than 100 °C. Its reproducibility, scalability and environmentally friendly reaction conditions are what make it significant for industry. Nevertheless, this method results in particles that vary widely in size, likely due to the complex series of reactions involved in magnetite formation.

2.3.4. Market

Due to magnetite nanoparticles’ numerous industrial applications, the market for these particles has grown significantly in recent years. These nanoparticles are widely used in biomedical areas for drug delivery, enhancing contrast in magnetic resonance imaging (MRI) and treating hyperthermia in cancer patients due to their distinct magnetic characteristics and biocompatibility. Because of their enormous surface area and reactivity, they are also essential in environmental applications, including pollutant removal and water purification [28]. Ongoing research and development, which consistently find new applications and enhance synthesis techniques, ensure the production of quality nanoparticles with regulated size and functionality and accelerate market expansion even more. As a result, it is expected that the demand for magnetite nanoparticles will increase, indicating their increasing significance for environmental sustainability and technological breakthroughs. By 2029, the market is expected to have grown to a value of 2.807 billion Rands (ZAR), with a compound yearly growth rate (CAGR) of 11.8% [29].

3. Material and Methods

3.1. Feedstock

Artificial mine water solutions containing FeSO₄ (Promark chemicals, South Africa) and Fe₂(SO₄)₃ (Promark chemicals, South Africa) were used for the formation of magnetite. NaOH, CaCO₃ (Promark chemicals, South Africa) and CaO (Sigma-Aldrich, South Africa) were used as alkalis. The Fe₂(SO₄)₃ was associated with 5.4 H₂O. There is significance in computing x in Fe₂(SO₄)₃ xH₂O; hydrates are substances whose crystal structures contain molecules of water. There are several hydrates with unique properties depending on the quantity of water molecules (x) in the mixture. It is important to know the precise hydration state (value of x). Hydrates have an impact on the stoichiometry and results of chemical reactions. The specific composition of the reactants must be known for accurate calculations.

Stock solutions of 200 mmol/L of Fe²⁺, Fe³⁺ and H₂SO₄ (Promark chemicals, South Africa) were prepared, and 200 mL of each solution was added to conical flask and then followed by 400 mL of H₂O, which was used as feed. The concentration of Fe²⁺ in the feed was 2289.9 mg/L and 4579.7 mg/L for Fe³⁺. Two alkalis were used for neutralisation, CaCO₃ and CaO.

3.2. Equipment

Studies were performed in 1000 mL flat bottom conical flask at room temperature. A stirrer plate with stirrer bar was used for stirring. A hotplate was used for heat treatment of the slurry to produce magnetite at different temperatures (25, 50, 100, 200 and 300 °C). A portable pH meter was used to measure the pH of the samples during experiments.

3.3. Procedure

OLI software (Studio 11.5) simulations were used to predict the water quality when Fe²⁺-rich water reacted with Fe³⁺ by controlling the pH when dosing (i) NaOH, (ii) Ca(OH)₂ and (iii) CaCO₃.

3.3.1. Neutralisation

• Preparation of the Synthetic Mine Water:
  • Amounts of 200 mL of Fe³⁺ solution and 200 mL of H₂SO₄ were mixed to achieve complete dissolution of Fe³⁺.
  • Amounts of 400 mL of distilled water and 200 mL of Fe²⁺ solutions were added to the mixture.
  • The combined solution was stirred at 250 rpm for 10 min to ensure complete mixing.
  • An amount of 50 mL of the feed solution was collected to measure the initial pH, Fe²⁺, Fe³⁺ and acidity.
• Neutralisation with CaCO₃:
  • A beaker study was conducted to measure the neutralisation rate with CaCO₃.
  • An amount of 160 mmol (16 g) of CaCO₃ was added to the feed solution to adjust the pH to 5.
  • The solution was allowed to settle, and the clear water was decanted to measure the pH, Fe²⁺, Fe³⁺ and acidity.
• Neutralisation with CaO:
  • The pH was raised to 7 by adding 60 mmol (3.36 g) of CaO.
  • The slurry was stirred, and samples were taken every 30 min for 3 h to measure the pH, Fe²⁺, Fe³⁺ and acidity, as the pH continued to rise with the dissolution of CaO.
• Magnetite Formation:
  • The de-watered slurry from the previous steps was heated at different temperatures to promote magnetite formation.

3.3.2. Gypsum Crystallisation

Preparation of Gypsum-Saturated Solution:

An amount of 4 g of CaSO₄ was dissolved in 1 L of distilled water.

The solution was stirred at 250 rpm for 24 h to achieve saturation.

Gypsum/Magnetite Separation:

Single and multiple gypsum/magnetite dosages (0.5 g) were tested in 100 mL of the saturated gypsum solution.

The mixture was stirred on a magnetic stirrer for 24 h to allow for the formation of gypsum particles that can settle for better separation of magnetite and gypsum.

The Fe₃O₄ that was attracted to the magnet was separated and dried at 100 °C.

The mass of the dry Fe₃O₄ product was determined.

3.3.3. Particle Size Distribution

• An amount of 0.5 mg of each product was added in a sample vial.
• An amount of 1.5 mL of distilled H₂O was added in a sample.
• The solution was sonicated for 10 min at 25 °C.
• After sonication, the upper liquid was used to measure particle size use Malvern Zetasizer Nano series.

3.3.4. Sludge Settling

• Preparation of Solutions:
  • Solutions for settling rates of Fe(OH)₃ and Fe₃O₄ sludge were prepared using the following concentrations: Fe²⁺: 4 mmol/L, Fe³⁺: 4 mmol/L and H₂SO₄: 4 mmol/L.
• Production of Fe(OH)₃ Sludge:
  • An amount of 1.60g of CaCO₃ was added to 1 L of synthetic mine water, which comprised 200 mL of Fe²⁺ solution, 200 mL of Fe³⁺ solution, 200 mL of H₂SO₄ and 400 mL of distilled water.
  • A volume of 500 mL of this slurry was taken for further processing.
• Flocculant Addition:
  • An amount of 10 mg/L of SEP-G-54 flocculant was added to the 500 mL slurry.
  • The mixture was rapidly stirred at 200 rpm for 30 s.
  • It was then mixed slowly at 30 rpm for 3 min.
• Settling of Fe(OH)₃ Sludge:
  • After mixing, the sludge was allowed to settle.
  • The time taken and the height of the settling process were recorded.
• Production of Fe₃O₄ Sludge:
  • For the Fe₃O₄ sludge, 1.60 g of CaCO₃ and 0.28 g of CaO were added to the same synthetic mine water solution.
  • The same procedure was followed as described above to measure the settling rate of the Fe₃O₄ sludge.

3.4. Experimental

The effect of Fe²⁺ concentration (40 mmol/L), Fe³⁺ concentration (40 mmol/L) and alkalis (304 mmol/L NaOH, 160 mmol/L CaCO₃, 60 mmol/L CaO) were investigated using OLI simulations. Kinetic laboratory studies were also conducted that focused on sludge production at temperatures of 25, 50, 100, 200 and 300 °C; gypsum crystallisation through dosing of alkalis (instant and multiple) over 24 h; particle size of the recovered pigments; and settling studies focusing on sludge composition (Fe(OH)₃/gypsum, Fe₃O₄/gypsum slurry) and flocculant concentration (SEP-G54).

3.5. Analytical

Samples were collected at various stages, filtered (Whatman no.1 paper-Cytiva, Marlborough, MA, USA) and assayed for Fe²⁺, Fe³⁺, pH and acidity concentrations using standard procedures [30]. Fe²⁺ concentrations were determined by adding filtered sample (5 mL), 0.1N H₂SO₄ (5 mL) and Zimmerman-Reinhardt reagent (5 mL) to an Erlenmeyer flask and titrating the solution with 0.05 N KMnO₄ until pale pink [30]. Fe³⁺ concentrations were determined by first determining the total iron by bringing the sample (10 mL) to a pH of 0.5, and then drops of SnCl₂ solution for reduction of Fe³⁺ to Fe²⁺ were added to agitate iron solution until the yellow colour disappeared, followed by one drop extra. The solution was allowed to cool for a few minutes, and then 10 mL of mercury(ll) chloride solution was quickly added, and when the precipitate was grey, a mixture of 100 mL water and 13 mL of Zimmermann-Reinhardt solution was added. The reduced sample was titrated with 0.05 N of KMnO₄ until a faint pink colour was persistent for 30 s. The concentration of Fe³⁺ = total iron—Fe²⁺.

X-ray diffraction (XRD) was applied in this study where powdered samples were analysed using a PANalytical X’ Pert Pro diffractometer (Netherlands), which was run using Cu Kα radiation as the X-ray source at wavelength of 1.5406 Å. Mineral phase recognition by XRD depends on the diffraction peaks at different 2θ values relating to the mineral’s d-spacing value. According to Bragg’s Law, as each mineral has a highly ordered crystalline structure, its d-spacing values are characterised by a set of diffraction peaks [31].

Morphological and elemental properties of the synthesised pigments were elucidated using Scanning Electron Microscope (SEM) (Model: Crossbeam 340 FIB-SEM with Secondary Electron and Backscattered Electron Detectors, Carl Zeiss, Oberkochen, Germany). equipped with the means to perform Energy-dispersive X-ray spectroscopy (EDS).

Particle sizes of the nanoparticles were analysed using light scattering instrument (Model: Zetasizer Nano ZS, Malvern Panalytical, Worcestershire, UK).

3.6. OLI Software Simulations

The OLI ESP software (Studio 11.5) program was used to predict the behaviour of metals dissolved in water during dosing of alkalis, like CaCO₃ and Ca(OH)₂. The solubility of CaSO₄ and metal hydroxides was identified as a function of temperature and concentration.

One of the most complex issues for researchers and students in the field of chemistry or advanced chemical engineering is the problems with pH calculations, solubility calculations and redox calculations. The easiest way to solve these problems is to write out the equilibrium expressions or the K equations and their corresponding equations, such as mole balance and charge balance while making simple assumptions along the way wherever useful and possible, and finally to combine the remaining equations to solve the resulting polynomial. The problems can be time-consuming even for a simple setup such as one acid in water or one dissolving salt. Environmental engineers and water engineers, among many, must deal with such problems daily in the field. For this purpose, they use various software packages ranging from simple to very detailed.

Most equilibrium modelling software packages are thought of as having two critical parts: the algorithm to solve any distinct industrial problem and the complete datasets that contain the thermodynamic constants or K values for all these problems. Additional constituents such as ionic strength corrections, adsorption equations and titration simulations, however, vary from program to program. During this study equilibrium modelling will also be used for predicting the chemical speciation during the formation of magnetite reactions.

A dependable aqueous equilibrium chemistry calculator must have an interactive and self-instructive interface for clarifying reactions, the ability to work with all kinds of common equilibrium reactions, a strong solution algorithm, expressive and easily understandable displays for results and the ability to produce results in multiple formats per different use requirements. Depending on the application it is being used for, it may be more sensible to include a database of reactions relevant to this specific purpose, specific models of complex systems or the ability to vary temperature, pressure and ionic strength. For example, the OLI Systems Analyzer that will be used in this study can perform all this with the additional option of customising every feature according to requirements and simulating them using the OLI ESP Program (11.5) [32].

The Stream Analyzer is OLI’s simplest and clearest answer to the electrolyte thermodynamic water problems of the oil and gas industry. This software features single-point equilibrium calculations, multiple-point survey calculations for calculating a complete trend analysis for characteristics like temperature, pressure, pH and composition effects, and simple mix and separate capability. The calculations provide vapour, liquid, solid and second liquid phase separations for a fully specialised model.

4. Results and Discussion

4.1. Magnetite Formation

The aim of this section was to determine the effect of the following conditions needed for magnetite (Fe₃O₄) formation: (i) ratio of Fe³⁺/Fe²⁺ (ii) alkali dosage (iii) O₂ concentration. In addition to the mentioned aims, the following aspects will also be investigated: (i) the effect of gypsum that forms in the case of calcium alkalis (CaCO₃ and CaO) and (ii) the separate recovery of Fe₃O₄ and CaSO₄·2H₂O.

4.1.1. Solubility of Fe²⁺ and Fe³⁺

This study was carried out by using FeSO₄·7H₂O as Fe²⁺ and Fe₂(SO₄)₃·4H₂O as Fe³⁺.

Table 1. Acid required to keep Fe₂(SO₄)₃ in solution (OLI simulation).

Fe2(SO4)3 [mmol]	H2SO4 [mmol]	pH	Fe(OH)3 (Bernalite)—Sol [mmol]	Fe3+ Liq1 [mmol]	SO42− Liq1 [mmol]
40.00	0.00	1.79	22.72	57.28	120.00
40.00	4.00	1.79	20.47	59.53	124.00
40.00	8.00	1.79	18.20	61.80	128.00
40.00	12.00	1.79	15.93	64.07	132.00
40.00	16.00	1.79	13.65	66.35	136.00
40.00	20.00	1.79	11.36	68.64	140.00
40.00	24.00	1.79	9.07	70.93	144.00
40.00	28.00	1.79	6.76	73.24	148.00
40.00	32.00	1.79	4.45	75.55	152.00
40.00	36.00	1.79	2.13	77.87	156.00
40.00	40.00	1.79	0.00	80.00	160.00

shows that 40 mmol Fe₂(SO₄)₃·4H₂O needs to be acidified with 40 mmol H₂SO₄ to be soluble Equation (7):

Fe₂(SO₄)₃ + H₂SO₄ → 2Fe³⁺ + 4SO₄²⁻ + 2H⁺

(7)

4.1.2. Removal of Fe²⁺ and Fe³⁺

Fe²⁺ and Fe³⁺ ions in mine water need to be removed as Fe(OH)₃ or as Fe₃O₄ magnetite. The removal of Fe²⁺ requires oxidation to Fe³⁺ Equation (8) to be precipitated as Fe(OH)₃ through neutralisation (Equations (9) and (10)). Alternatively, Fe²⁺ can be removed as Fe₃O₄ (magnetite) by reacting Fe(OH)₂ (Equations (10) and (11)) with Fe(OH)₃.

2Fe²⁺ + ½O₂ + 2H⁺ → 2Fe³⁺ + H₂O

(8)

Fe³⁺ + 3H₂O → Fe(OH)₃ + 3H⁺

(9)

Fe²⁺ + 2OH⁻ → Fe(OH)₂

(10)

Fe(OH)₂ + 2Fe(OH)₃ → Fe₃O₄ + 4H₂O

(11)

Iron removal needs alkali addition, such as CaCO₃ (R1050/t), Ca(OH)₂ (R3950/t), Na₂CO₃ (R10,000/t) and NaOH (R12,000/t). The use of calcium alkalis has the following benefits: (i) partial desalination is achieved through gypsum crystallisation (Equation (6)) and (ii) lower cost, as shown in brackets. Sodium alkalis have the benefit that only Fe(OH)₃ is produced. A disadvantage is that the presence of sodium adds salinity to the treated water, which makes it less suitable for crop production.

Ca²⁺ + SO₄²⁻ + 2H₂O → CaSO₄·2H₂O

(12)

Equations (13) and (14) show the reactions for Fe₃O₄ and Fe(OH)₃, respectively, from water containing 40 mmol/L FeSO₄ and 40 mmol/L Fe₂(SO₄)₃.

2FeSO₄ + 2Fe₂(SO₄)₃ + 16OH⁻ → 2Fe₃O₄ + 8SO₄²⁻ + 8H₂O

(13)

2FeSO₄ + 2Fe₂(SO₄)₃ + ½O₂ + 16OH⁻ + H₂O → 6Fe(OH)₃ + 8SO₄²⁻

(14)

4.1.3. Conditions Needed for Magnetite Formation

OLI simulations were carried out to predict the effect of various parameters on the formation of Fe₃O₄ in a solution containing 40 mmol/L Fe₂SO₄, 40 mmol/L Fe(SO₄)₃·4H₂O and 40 mmol/L H₂SO₄. The H₂SO₄ was added to keep Fe³⁺ in solution in the feed. The alkalis, CaCO₃ and Na₂CO₃ were used to investigate the effect of gypsum on the formation of Fe₃O₄. O₂ was varied to confirm that its presence will prevent the formation of Fe₃O₄, as Fe(OH)₃ will be formed. The Fe²⁺/Fe³⁺ mole ratio was varied to identify the ratio needed for maximum Fe₃O₄ formation.

Table 2. Fe₃O₄ with Na₂CO₃.

Reactants						Solid Products					Solution
Na2CO3 [mmol]	FeSO4 [mmol]	Fe2(SO4)3 [mmol]	H2SO4 [mmol]	O2 [mmol]	pH	Fe3O4 (Magnetite)—Sol [mmol]	Fe(OH)3 (Bernalite)—Sol [mmol]	CaSO4·2H2O (Gypsum) [mmol]	CaCO3 (Calcite)—Sol [mmol]	CO2—Vap [mmol]	Fe3+ Liq1 [mmol]	Fe2+ Liq1 [mmol]	Ca2+ Liq1 [mmol]	Na+ Liq1 [mmol]	S6+ Liq1 [mmol]	C4+ Liq1 [mmol]	H2O [mmol]
0	40	40	40	0	1.83	0.00	0.00	0.0	0.0	0.00	80.00	40.00	0.00	0	200	0.00	55,483
20	40	40	40	0	1.85	0.00	11.99	0.0	0.0	0.00	68.01	40.00	0.00	40	200	20.00	55,468
40	40	40	40	0	1.87	0.00	23.79	0.0	0.0	0.00	56.21	40.00	0.00	80	200	40.00	55,454
60	40	40	40	0	1.90	0.00	35.36	0.0	0.0	13.63	44.64	40.00	0.00	120	200	46.37	55,441
80	40	40	40	0	1.94	0.00	46.66	0.0	0.0	34.14	33.34	40.00	0.00	160	200	45.86	55,428
100	40	40	40	0	1.99	0.00	57.56	0.0	0.0	54.66	22.44	40.00	0.00	200	200	45.34	55,417
120	40	40	40	0	2.08	0.00	67.78	0.0	0.0	75.18	12.22	40.00	0.00	240	200	44.82	55,409
140	40	40	40	0	2.25	0.00	76.42	0.0	0.0	95.72	3.58	40.00	0.00	280	200	44.28	55,407
160	40	40	40	0	4.21	0.00	80.00	0.0	0.0	115.98	0.00	40.00	0.00	320	200	44.02	55,424
180	40	40	40	0	4.62	19.60	40.80	0.0	0.0	142.54	0.00	20.40	0.00	360	200	37.46	55,482
200	40	40	40	0	5.16	38.34	3.33	0.0	0.0	166.79	0.00	1.66	0.00	400	200	33.21	55,536

shows that 200 mmol Na₂CO₃ was needed to produce 38.34 mmol Fe₃O₄ at a pH of 5.16.

Table 3. Fe₃O₄ with CaCO₃.

Reactants						Solid Products					Solution
CaCO3 [mmol]	FeSO4 [mmol]	Fe2(SO4)3 [mmol]	H2SO4 [mmol]	O2 [mmol]	pH	Fe3O4 (Magnetite)—Sol [mmol]	Fe(OH)3 (Bernalite)—Sol [mmol]	CaSO4·2H2O (Gypsum) [mmol]	CaCO3 (Calcite)—Sol [mmol]	CO2—Vap [mmol]	Fe3+ Liq1 [mmol]	Fe2+ Liq1 [mmol]	Ca2+ Liq1 [mmol]	Na+ Liq1 [mmol]	S6+ Liq1 [mmol]	C4+ Liq1 [mmol]	H2O [mmol]
0	40	40	40	0	1.83	0.00	0.00	0.0	0.0	0.00	80.00	40.0	0.00	0.0	200.0	0.00	55,483
20	40	40	40	0	1.85	0.00	11.11	4.0	0.0	0.00	68.89	40.0	16.03	0.0	196.0	20.00	55,463
40	40	40	40	0	1.86	0.00	22.71	24.6	0.0	0.00	57.29	40.0	15.38	0.0	175.4	40.00	55,408
60	40	40	40	0	1.87	0.00	34.06	45.2	0.0	12.48	45.94	40.0	14.77	0.0	154.8	47.52	55,353
80	40	40	40	0	1.89	0.00	45.15	65.8	0.0	32.57	34.85	40.0	14.19	0.0	134.2	47.43	55,300
100	40	40	40	0	1.92	0.00	55.90	86.4	0.0	52.66	24.10	40.0	13.60	0.0	113.6	47.34	55,247
120	40	40	40	0	1.96	0.00	66.14	107.0	0.0	72.75	13.86	40.0	13.02	0.0	93.0	47.25	55,197
140	40	40	40	0	2.08	0.00	75.33	127.6	0.0	92.84	4.67	40.0	12.42	0.0	72.4	47.16	55,152
160	40	40	40	0	4.10	0.00	80.00	148.2	0.0	112.70	0.00	40.0	11.79	0.0	51.8	47.32	55,128
180	40	40	40	0	4.51	19.73	40.53	167.6	0.0	139.60	0.00	20.3	12.42	0.0	32.4	40.37	55,147
200	40	40	40	0	5.08	38.92	2.16	184.6	0.0	170.00	0.00	1.1	15.36	0.0	15.4	34.85	55,169

shows that CaCO₃ was needed for the formation of 39.92 mmol Fe₃O₄. In the case of CaCO₃, 184.6 mmol gypsum formed. OLI simulations showed that Fe₃O₄ forms when only carbonate is used as the alkali.

Table 4. Fe(OH)₃ with CaCO₃ and O₂.

Reactants						Solid Products					Solution
CaCO3 [mmol]	FeSO4 [mmol]	Fe2(SO4)3 [mmol]	H2SO4 [mmol]	O2 [mmol]	pH	Fe3O4 (Magnetite)—Sol [mmol]	Fe(OH)3 (Bernalite)—Sol [mmol]	CaSO4·2H2O (Gypsum) [mmol]	CaCO3 (Calcite)—Sol [mmol]	CO2—Vap [mmol]	Fe3+ Liq1 [mmol]	Fe2+ Liq1 [mmol]	Ca2+ Liq1 [mmol]	Na+ Liq1 [mmol]	S6+ Liq1 [mmol]	C4+ Liq1 [mmol]	H2O [mmol]
0	40	40	40	10	1.79	0.00	16.13	0.0	0.0	0.00	103.87	0.0	0.00	0.0	200.00	0.00	55,452
20	40	40	40	10	1.81	0.00	27.07	2.1	0.0	0.00	92.93	0.0	17.91	0.0	197.91	20.00	55,437
40	40	40	40	10	1.81	0.00	38.88	22.7	0.0	7.44	81.12	0.0	17.26	0.0	177.26	32.56	55,380
60	40	40	40	10	1.81	0.00	50.56	43.3	0.0	27.52	69.44	0.0	16.68	0.0	156.68	32.48	55,324
80	40	40	40	10	1.82	0.00	62.05	63.9	0.0	47.60	57.95	0.0	16.10	0.0	136.10	32.40	55,269
100	40	40	40	10	1.82	0.00	73.33	84.5	0.0	67.68	46.67	0.0	15.55	0.0	115.55	32.32	55,214
120	40	40	40	10	1.83	0.00	84.35	105.0	0.0	87.75	35.65	0.0	15.04	0.0	95.04	32.25	55,160
140	40	40	40	10	1.84	0.00	95.05	125.4	0.0	107.81	24.95	0.0	14.60	0.0	74.60	32.19	55,108
160	40	40	40	10	1.87	0.00	105.30	145.7	0.0	127.87	14.70	0.0	14.29	0.0	54.29	32.13	55,058
180	40	40	40	10	1.95	0.00	114.69	165.8	0.0	147.91	5.31	0.0	14.24	0.0	34.24	32.09	55,013
200	40	40	40	10	3.99	0.00	120.00	185.0	0.0	167.64	0.00	0.0	15.02	0.0	15.02	32.36	54,993

showed that the presence of O₂ resulted in the formation of Fe(OH)₃ and eliminated the formation of Fe₃O₄.

Table 5. Effect of Fe²⁺ to Fe³⁺ mole ratio on Fe₃O₄ formation.

NaOH [mmol]	FeSO4 [mmol]	Fe2(SO4)3 [mmol]	H2SO4 [mmol]	O2 [mmol]	pH	Fe3O4 (Magnetite)—Sol [mmol]	Fe(OH)3 (Bernalite)—Sol [mmol]	FeSO4 [mmol]	Fe3+ Liq1 [mmol]	Fe2+ Liq1 [mmol]	S6+ Liq1 [mmol]	Na+ Liq1 [mmol]	H2O [mmol]
400	0	40	0.00	40	0	12.67	0.00	80.00	0	0	160	400	55,588
400	4	40	0.05	40	0	12.62	4.00	72.00	0	0	164	400	55,604
400	8	40	0.10	40	0	12.57	8.00	64.00	0	0	168	400	55,620
400	12	40	0.15	40	0	12.51	12.00	56.00	0	0	172	400	55,636
400	16	40	0.20	40	0	12.44	16.00	48.00	0	0	176	400	55,652
400	20	40	0.25	40	0	12.36	20.00	40.00	0	0	180	400	55,668
400	24	40	0.30	40	0	12.26	24.00	32.00	0	0	184	400	55,684
400	28	40	0.35	40	0	12.13	28.00	24.00	0	0	188	400	55,700
400	32	40	0.40	40	0	11.96	32.00	16.00	0	0	192	400	55,716
400	36	40	0.45	40	0	11.65	36.00	8.00	0	0	196	400	55,732
400	40	40	0.50	40	0	7.32	40.00	0.00	0	0	200	400	55,748

shows that the ratio of Fe²⁺/Fe³⁺ needs to be 0.5 for maximum Fe₃O₄ formation (40 mmol).

The findings as predicted with OLI simulations were evaluated through laboratory studies.

Table 6. Composition of reactants to form magnetite.

Parameter		Mol Mass	Fe	Final Sol	Stock Sol				Unit
		g/mol	mg/L Fe	mmol/L	mmol/L	g/L	mL/L	Volume/Mass
Total Volume	mL							1000	mL
Fe2(SO4)3·5.4H2O	mmol	496.90	2234	40	200	99.38		200.0	mL
H2SO4	mmol	98.00		40	200	19.60	10.87	200.0	mL
H2O								400.0	mL
FeSO4·7H2O	mmol	277.86	4468	40	200	55.57		200.0	mL
CaO	mmol	56.00		200	500			11.20	g CaO
CaCO3	mmol	100.00		200	500			20.00	g CaCO3
Na2CO3	mmol	106.00		200				21.20	g Na2CO3

shows how the chemical solutions were prepared with the aim of determining the effect of alkali type and alkali dosage on the formation of Fe₃O₄:

• Separate preparation of Fe(OH)₃ from Fe³⁺ and Fe(OH)₂ from Fe²⁺ for both Na₂CO₃, Ca(OH)₂.
• Separate preparation of Fe(OH)₃ from Fe³⁺ with CaCO₃ and Fe(OH)₂ from Fe²⁺ with Ca(OH)₂.
• Sequential formation of Fe(OH)₃ and Fe(OH)₂ by dosing the alkali to the Fe³⁺/Fe²⁺/H⁺ solution.

4.1.4. Magnetite Formation in the Absence of Gypsum

Table 7. Magnetite formation in the absence of gypsum.

Reactants				Feed	Treated
Conditions
Reagents	Conc.	mol Mass	Conc.
	mmol/L	g	g/L
Fe2(SO4)3·5.4H2O	34.78	496.90	17.28
H2SO4	34.78	98.00	3.41
FeSO4·7H2O	34.78	277.85	9.66
NaOH for Fe3+ and Fe2+	304.35	40.00	12.17
NaOH dosage (g/L)				0.00	12.17	12.18	12.21	12.24	12.28
Results
pH				1.80	7.00	8.20	9.08	10.12	11.08
Fe2+ (mg/L)				2290	502.7	167.6	111.7	111.7	55.9
Fe3+ (mg/L)				4580	0.00	0.00	0.00	0.00	0.00
Acidity (mg/L CaCO3)				19,800	900	200	0.00	0.00	0.00
Temperature					Magnetic strength after heating (0—none; 10—strong)
25 °C dry					0.00	10.00	10.00	10.00	10.00
50 °C dry					0.00	10.00	10.00	10.00	10.00
100 °C dry					10.00	10.00	10.00	10.00	10.00
200 °C dry					10.00	10.00	10.00	10.00	10.00
300 °C dry					10.00	10.00	10.00	10.00	10.00
Temperature					Colour
25 °C dry					Brown	Black	Black	Black	Black
50 °C dry					Brown	Black	Black	Black	Black
100 °C dry					Brown	Brown	Dark brown	Black	Black
200 °C dry					Brown	Brown	Brown	Brown	Dark brown
300 °C dry					Brown	Brown	Brown	Brown	Brown

shows the results when Fe₃O₄ was formed from Fe(OH)₃ and Fe(OH)₂. Fe(OH)₃ was produced from Fe₂(SO₄)₃ and NaOH, and Fe(OH)₂ was produced from FeSO₄ and NaOH. All the reagents were put in the same vessel. NaOH was dosed with the aim first to raise the pH to 5.3 and to precipitate Fe³⁺ as Fe(OH)₃, followed by dosing more NaOH to precipitate Fe²⁺ as Fe(OH)₂ in the pH range of 7.0 to 11.1.

The following observations were made:

• No gypsum was formed due to the absence of calcium alkalis.
• The Fe(OH)₃ and Fe(OH)₂ sludge reacted to form Fe₃O₄ after dewatering and heated to 50 °C, 100 °C, 200 °C and 300 °C at pH 8.2 and higher. At pH 7, magnetic sludge was only produced when heated to 100 °C and higher. The combined Fe(OH)₃/Fe(OH)₂ sludge was not magnetic immediately after mixing, only after the sludge was dewatered and heated. The reason for this needs a special investigation.

At pH 7, when the slurry was dried at 25 °C and 50 °C, the recovered magnetite was not magnetic at this lower pH value. This was due to the absence of Fe(OH)₂. It only forms after pH 7.6, as predicted by OLI simulations (Figure 2). The magnetite produced from pH 8–11 and its magnetic strength is shown in (Figure 3). It is thus concluded that the solid, Fe(OH)₂, is needed as a reactant for the formation of Fe₃O₄.

Table 8. Effect of pH on the formation of goethite and magnetite.

Compound	Formula	pH
		8	11
Gypsum	CaSO4·2H2O	1.9	1.5
Magnetite	Fe3O4	8.4	21.2
Hematite	Fe2O3	0.0	0.0
Thenardite	Na2SO4	53.0	57.3
Goethite	HFeO2	36.7	20.0
Total		100.0	100.0
Colour		Black	Black

shows the effect of pH (8 and 11) on the formation of magnetite and goethite when Fe²⁺ is reacted with Fe³⁺ in the absence of gypsum. It was noted that magnetite formation was favoured at pH 11, while goethite was the main product at pH 8. The effect of pH on the formation of magnetite will be further investigated. The high Na₂SO₄ (53.0 g/L and 57.3 g/L) can be ascribed to incomplete washing of the sludge. The low gypsum concentrations should be ignored as no Ca²⁺ was present in the reactants.

Removal of Na₂SO₄ from Fe₃O₄

The pH of 11 was selected as the optimal value because magnetite particles formed even when the sludge was air-dried at room temperature. The magnetite recovered at pH 11 (dried at 25 and 100 °C) was thoroughly washed multiple times with deionised water by centrifuging at 4500 rpm for 10 min to ensure the complete removal of Na₂SO₄. The results presented below show the SEM-EDS after washing.

The micrographs of Fe₃O₄ (magnetite) particles, taken at 25 °C and 100 °C drying conditions in Figure 4 demonstrate notable morphological differences. At room temperature (25 °C), the magnetite exhibits a porous, aggregated structure made up of loosely packed nanoparticles. Lower magnification reveals large, irregular clusters, while higher magnification highlights the rough texture and intact porous network. Conversely, the magnetite dried at 100 °C shows a more compact and less porous structure, characterised by denser aggregates and fewer gaps between particles. At higher magnifications, partially fused nanoparticles with smoother surfaces are observed, likely a result of sintering or interparticle bonding due to the increased drying temperature. Irregular morphology is common when using precipitation [33]. The SEM/EDS shows that Na₂SO₄ was completely washed off. The C present was from the carbon paper that was used. The recovered magnetite at pH 11 indicates a nanoscale material with a particle size of 143.6 d.nm, as shown in Figure 5.

4.1.5. Fe₃O₄/Fe₂O₃ Formation in the Presence of Gypsum

CaCO₃ as the Alkali

OLI simulations indicated that only CaCO₃ is needed for the formation of magnetite (Fe₃O₄) from Fe²⁺ and Fe³⁺.

Table 9. Fe(OH)₃ and gypsum formation with CaCO₃.

Parameter	Feed	Treated
Water Quality
Time (h)	0.0	0.3	1	3	4	5
CaCO3 (g/L)		20.00
pH	1.3	6.1	6.2	5.4	5.9	6.1
Fe2+	2457	2457	2457	2569	2457	2457
Fe3+	4580	0	0	0	0	0
Acidity	19,800.0
Sludge colour		Brown	Brown	Brown	Brown	Brown

showed that when CaCO₃ is used as the alkali for the treatment of a solution that contained Fe₂(SO₄)₃, FeSO₄ and H₂SO₄, only Fe₂O₃ formed, and no Fe₃O₄. The CaCO₃ dosage was sufficient to neutralise the H₂SO₄ and precipitate Fe³⁺ as Fe(OH)₃ and Fe²⁺ as Fe(OH)₂. Strict conditions were created to prevent Fe²⁺ oxidation. The presence of O₂ was eliminated by closing the 1 L conical flask with parafilm and by putting the flask in a 5 L nitrogen filled beaker. Gypsum was formed due to the presence of calcium alkalis. The settled Fe(OH)₃ and CaSO₄·2H₂O sludge was stirred for 10 min at room temperature, dewatered and heated to 50 °C, 100 °C, 200 °C and 300 °C. Fe₃O₄ (magnetite) can only be formed when Fe(OH)₂ is formed and when Ca(OH)₂ is dosed. Only CaCO₃ could not result in Fe₃O₄ formation as it cannot produce the essential Fe(OH)₂ when the pH is raised to 7.6 and higher, as shown in Figure 2. This finding is in contradiction to that predicted by OLI simulations.

Figure 6 showed that upon heating, the colour changed to yellow at 100 °C, which is an indication of goethite formation, and to red at 300 °C, which is an indication of hematite formation.

Table 10. Effect of temperature on the formation of goethite and hematite (XRD Wt%).

Compound	Formula	Temperature (°C)
		100	300
Gypsum	CaSO4·2H2O	2.4	0.3
Magnetite	Fe3O4	0.0	0.5
Calcite	CaCO3		12.1
Bassanite	HCaSO4.5	86.3	43.4
Anhydrite	CaSO4		37.6
Hematite	Fe2O3		6.1
Thenardite	Na2SO4	0.0
Goethite	HFeO2	11.3
Total		100.0	100.0
Colour		Yellow	Red

shows the effect of temperature (100 °C and 300 °C) on the chemical composition of the sludge. Goethite was the main ion compound at 100 °C (11.3%) and hematite at 300 °C (6.1%). It was noted that at 100 °C, the CaSO₄ was only present as bassanite, while at 300 °C, it appeared as bassanite (HCaSO₄) and anhydrite (CaSO₄). The SEM images in Figure 7 of goethite and hematite containing gypsum reveal rough, textured surfaces typical of their crystalline structures, shaped by the presence of gypsum. In the case of goethite, the interaction with gypsum encourages agglomeration, creating irregular porous features, different degrees of particle clustering and possible gypsum intercalation within the goethite matrix, resulting in layered structures or voids. Likewise, for hematite, the interaction with gypsum promotes particle aggregation, resulting in clusters, layered or porous formations and a diverse range of particle sizes, including both larger aggregates and smaller particles. The XRD spectra in Figure 8 and Figure 9 confirmed the presence of goethite and hematite with good crystallinity.

CaCO₃ and CaO as the Alkali

Table 11. Magnetite and gypsum formation with CaCO₃ and CaO.

Parameter	Unit	Instant Dosing						Multiple Dosing
		Feed	CaCO3 and CaO					Feed	CaCO3				CaO
Dosing		0	1	2	3	4	5	0	1	2	3	4	5	6	7	8
Time	min		0	30	60	120	150		0	30	60	90	120	150	180	181
CaCO3	g		16.0						4.0	4.0	4.0	4.0
CaO	g			3.36									1.12	1.12	1.12
pH		1.78	1.78	5.01	6.37	6.61	7.13	1.78	1.78	2.47	2.50	5.65	6.01	6.54	6.92	7.20
Fe2+	mg/L	2402	2402	2290	1173	614	279	2402	2402	2346	2346	2346	2066	1396	559	223
Fe3+	mg/L	4580	4580	0	0	0	0	4580	3407	614	167	0	0	0	0	0
Acidity	mg/L	19,800	4000	3600	3100	2700	800	19,800	19,400	12,000	9600	2900	2500	1300	800	600
Temperature	°C	Magnetic strength after heating (0—none; 10—strong) pH 7.1						Magnetic strength after heating (0—none; 10—strong) pH 7.2
50		6.00						6.00
100		8.00						8.00
200		8.00						8.00
300		8.00						8.00
Temperature	°C	Colour						Colour
50		Brown						Brown
100		Black-grey						Black-grey
200		Grey						Grey
300		Brown						Brown

shows that the CaCO₃ dosage was sufficient to neutralise the H₂SO₄, precipitate Fe³⁺ as Fe(OH)₃ and CaO to precipitate Fe²⁺ as Fe(OH)₂. Gypsum was formed due to the presence of Ca²⁺. The settled Fe(OH)₃, Fe(OH)₂ and CaSO₄·2H₂O sludge was stirred for 10 min at room temperature, dewatered and heated to 50 °C, 100 °C, 200 °C and 300 °C. The Fe(OH)₃/Fe(OH)₂/CaSO₄·2H₂O sludge was not magnetic immediately after mixing, only after the sludge was dewatered and heated.

Table 12. Magnetite combined with gypsum.

Compound	Formula	Composition (%)
		Instant	Multiple
Gypsum	CaSO4·2H2O	92.2	91.0
Magnetite	Fe3O4	6.7	6.5
Calcite	CaCO3	0.8	1.3
Bassanite	HCaSO4.5	0.0	1.0
Anhydrite	CaSO4	0.3	0.1
Hematite	Fe2O3	0.0	0.1
Thenardite	Na2SO4	0.0	0.0
Goethite	HFeO2	0.0	0.0

shows the results of when Fe₃O₄ was formed from a solution that contained Fe₂(SO₄)₃, FeSO₄ and H₂SO₄ when treated with CaO.

Table 13. Single and multiple dosing for sludge production.

Parameter	Unit	Instant Dosing						Multiple Dosing
		Feed	CaCO3 and CaO					Feed	CaCO3				CaO
Dosing		0	1	2	3	4	5	0	1	2	3	4	5	6	7	8
Time	min		0	30	60	120	150		0	30	60	90	120	150	180	181
CaCO3	g		16.0						4.0	4.0	4.0	4.0
CaO	g			3.36									1.12	1.12	1.12
pH		1.78	1.78	5.01	6.37	6.61	7.13	1.78	1.78	2.47	2.50	5.65	6.01	6.54	6.92	7.20
Fe2+	mg/L	2402	2402	2290	1173	614	279	2402	2402	2346	2346	2346	2066	1396	559	223
Fe3+	mg/L	4580	4580	0	0	0	0	4580	3407	614	167	0	0	0	0	0
Acidity	mg/L	19,800	4000	3600	3100	2700	800	19,800	19,400	12,000	9600	2900	2500	1300	800	600

show the sludge composition after heating to 100 °C. It was shown that the iron was magnetite (6.7% and 6.5%) with no other iron compounds present. The amount of magnetite formed was not affected by instant or multiple dosing. The question is whether it has an influence on magnetic separation.

The micrographs of Fe₃O₄ (magnetite) particles for both instant and multiple dosing, as shown in Figure 10, reveal clear morphological variations between instant and multiple alkali dosing methods. Instant dosing produces a compact structure with densely packed aggregates and larger clusters, whereas multiple dosing results in a more varied morphology, featuring finer gypsum crystals and a broader range of particle sizes. These observations emphasise the significant influence of dosing techniques on the properties of the magnetite–gypsum composite. Irregular morphology is common when using the precipitation method [33]. Gypsum showed a more elongated, fibrous and needle-like morphology, which is consistent with its natural tendency to form acicular (needle-shaped) or prismatic crystals [34,35].

The XRD results in Figure 11 and Figure 12 show the presence of characteristic magnetite peaks. The use of CaCO₃ and CaO provided a favourable environment for the precipitation of magnetite from the Fe²⁺/Fe³⁺ solution. The most prominent peaks for magnetite appear around 35° (2θ = 35.076°), and other peaks at 41.385°, 50.448°, 67.228° and 74.124°. These peaks are associated with the characteristic reflections of magnetite. The presence of gypsum and anhydrite compounds was the result of crystallisation as its solubility level was exceeded. The recovered magnetite was highly crystalline, as indicated by the sharp and intense peaks. The presence of other crystalline phases, such as gypsum, calcite and anhydrite, pointed to a well-ordered crystal structure for these byproducts. This confirms that these alkalis can be effective for the recovery of magnetite, with the byproducts being stable calcium sulphate minerals.

Importance of Hydroxide in Fe₃O₄ Formation

The hydroxide ion (OH⁻) is pivotal in the formation of magnetite (Fe₃O₄) from iron salts, influencing the entire synthesis process from nucleation to the final properties of the nanoparticles. One of the primary roles of OH⁻ is in controlling the pH of the reaction medium. The pH determines the speciation of iron ions and the subsequent precipitation of iron hydroxides. In a basic environment, OH⁻ ions react with Fe²⁺ and Fe³⁺ ions to form Fe(OH)₂ and Fe(OH)₃, which are precursors to magnetite. The optimal pH for magnetite formation is typically between 8 and 11. At this pH range, the solubility of iron hydroxides decreases, facilitating their precipitation and subsequent transformation into magnetite through a series of redox reactions [36,37].

Beyond pH control, OH⁻ ions directly influence the nucleation and growth of magnetite nanoparticles. A higher concentration of OH⁻ ions accelerates the nucleation process, leading to the formation of numerous small nuclei, which then grow into nanoparticles. The availability of OH⁻ ensures a consistent supply of reactants for these nucleation sites, promoting uniform particle size and morphology. Conversely, insufficient OH⁻ can result in incomplete precipitation and the formation of iron oxides other than magnetite, such as goethite or lepidocrocite, which have different magnetic properties [13,38].

OH⁻ ions played a crucial role in maintaining the stoichiometry and crystallinity of magnetite. The balance between Fe²⁺ and Fe³⁺ in the presence of OH⁻ ions is essential for forming the mixed-valence compound Fe₃O₄. An optimal OH⁻ concentration ensures that the redox reactions proceed correctly, leading to the formation of stoichiometric magnetite. Moreover, OH⁻ ions help in achieving high crystallinity by providing the necessary conditions for the ordered arrangement of iron and oxygen atoms in the crystal lattice. High crystallinity is directly related to better magnetic properties, as it reduces defects and enhances magnetic domain alignment [13].

Finally, OH⁻ ions aid in the removal of impurities and by-products during the synthesis process. During the formation of magnetite, various side reactions can occur, producing soluble by-products. OH⁻ ions can help in precipitating these impurities, which can then be removed through washing. This purification step is crucial for obtaining pure magnetite with desired magnetic properties. In summary, the role of OH⁻ ions in magnetite synthesis is multifaceted, encompassing pH control, nucleation and growth regulation; stoichiometry and crystallinity maintenance; and impurity removal, all of which are critical for producing high-quality magnetite nanoparticles [38].

4.2. Separation of Magnetite and Gypsum

When calcium alkalis are used for the neutralisation of mine water, rich in sulphate, gypsum co-precipitates with iron hydroxides. The purpose of this investigation was to recover the magnetite separately from gypsum through magnetic separation. It was argued that large gypsum crystals will favour the magnetic separation of the two compounds. Gypsum of small-sized particles was produced by dosing the calcium alkali in a single dosage. Gypsum with particles with large sizes was produced by dosing the calcium alkali in splitting the dosage in a number of small dosages over time. Iron-rich water (40 mmol/L FeSO₄, 40 mmol/L Fe₂(SO₄)₃, 40 mmol/L H₂SO₄) was neutralised with 160 mol/L CaCO₃ followed by 40 mmol/L Ca(OH)₂. In the case of a single dosage, the CaCO₃ was dosed at time 0 and the Ca(OH)₂ after 30 min. Samples were collected over a period of 180 min and analysed for pH, Fe²⁺ and acidity. In the case of multiple dosages, the CaCO₃ and CaO were split in four and dosed after 0, 30, 60 and 90 min. The CaCO₃ dosing was carried out first for the removal of Fe³⁺ as Fe(OH)₃, followed by CaO dosing for the removal of Fe²⁺ as Fe(OH)₂.

Table 14. Separation of gypsum and Fe₃O₄.

Parameter	Instant	Multiple
Time (h)	24	24
Feed	Fe3O4 Gypsum	Fe3O4 Gypsum
Total mass in 100 sat gypsum solution (Mg)	500.00	500.00
CaSO4·2H2On (Mg)	370.00	370.00
Fe3O4 (mg)	130.00	130.00
Dry product mass (mg) (100 °C)	84.90	149.50
Magnetite (XRD results) (%)	6.50	11.70
Magnetic	Yes	yes

shows how Fe²⁺ and acidity were removed as CaCO₃ and CaO were dosed for the two dosing methods. It was noticed that the gypsum particle size in the case of multiple dosing was larger than in the case of a single dosage. That was due to the growth of larger crystals at lower levels of over-saturation in the case of multiple dosages.

The single and multiple gypsum/magnetite dosages (0.5 g) were tested for separation by placing them in a 100 mL saturated gypsum solution to avoid dissolution of the gypsum component (

Table 15. XRD results (Wt %).

Compound	Formula	Dosing
		Instant (%)	Multiple (%)
Gypsum	CaSO4·2H2O	91.0	82.1
Magnetite	Fe3O4	6.5	11.7
Calcite	CaCO3	1.3	5.8
Bassanite	HCaSO4.5	1.0	0.0
Anhydrite	CaSO4	0.1	0.1
Hematite	Fe2O3	0.1	0.3

). A magnet was used to attract the Fe₃O₄. The mass was determined and analysed with XRD for the solid species on the magnet.

Table 16. Production of Fe₃O₄ sludge.

Parameter	mmol/L	g	L	Mass (g)	Vol (mL)	Stock	g/L	Stock Volume
H2O								400
H2SO4	4.0	98.0	1.0	0.39	0.22	5.0	1.96	200
Fe2(SO4)3·5.4H2O	4.0	496.9	1.0	1.99		5.0	9.94	200
FeSO4·7H2O	4.0	277.9	1.0	1.11		5.0	5.56	200
CaO for Fe2+	5.0	56.0	1.0	0.28			0.28
Total Volume (mL)								1000
Products
Fe3O4	4.0	231.6	1.00	0.93		5.0		200
Gypsum	21.0	172.0	1.00	3.61		5.0
Total		403.6	1.00	4.54

shows that more magnetite was attracted by the magnet in the case of multiple dosages (11.7%) than in the case of single doses (6.5%). This can be ascribed to larger gypsum particles that formed in the case of multiple dosing and that settled. This aspect should be further investigated as it may have the potential for full-scale application if a good separation is achieved. The focus is now on the production of magnetite-rich sludge and its settling.

The particle size of the products was measured as shown in Figure 13, and the presence of gypsum in the magnetite sample had a significant impact on particle size measurements, yielding Z-average values of 1588 d.nm (PDI 0.377) for instant dosing and 2064 d.nm (PDI 0.422) for multiple dosing, which indicate a polydisperse system. Gypsum promotes particle aggregation by connecting magnetite particles or modifying surface stability. XRD analysis showed a decrease in gypsum intensity after multiple dosing treatment and an increase after instant dosing, aligning with reduced particle sizes of 1404 d.nm (multiple dosing) and increased 2823 d.nm (instant dosing). The gradual addition of alkalis during multiple dosing resulted in the formation of large gypsum crystals that remained after heat treatment. During saturated gypsum treatment, these crystals settled at the bottom, separating from the magnetite during the recovery process.

4.3. Settling Studies

Fe(OH)₃ sludge settled slowly. The High-Density Sludge (HDS) process was developed to produce a sludge that settles at a fast rate. Lime was used in the HDS process to condition the sludge at pH 12. The conditioned sludge was then contacted with the AMD to form High-Density Sludge. The focus of this investigation was to compare the settling rate of magnetite-rich sludge, as described in the previous sections, with the settling rate of Fe(OH)₃-rich sludge. Fe₃O₄ and Fe(OH)₃-rich sludges were prepared as calculated in

Table 17. Production of Fe(OH)₃-rich sludge.

Parameter	mmol/L	g	L	Mass	Vol	Stock	g/L	Stock Volume
				g	mL
H2O								400
H2SO4	4.0	98.0	1.0	0.39	0.22	5.0	1.96	200
Fe2(SO4)3·5.4H2O	4.0	496.9	1.0	1.99		5.0	9.94	200
FeSO4·7H2O	0.0	277.9	1.0	0.00		5.0	0.00	200
CaO for Fe2+	0.0	56.0	1.0	0.00			0.00
Total Volume (mL)								1000
Products
Fe3O4	4.0	231.6	1.00	0.93
Gypsum	16.0	172.0	1.00	2.75
Total		403.6	1.00	3.68

. Settling was supported by dosing 10 mg/L SEP-G-54 flocculant to the waters with the sludge. After dosing with the flocculant, rapid stirring (200 rpm) was applied for 30 s, followed by slow mixing (30 rpm) for 3 min and sludge settling.

Table 18. Settling rate of Fe₃O₄/gypsum-rich sludge.

Time	Height	Rate
Min	Mm	m/h
0.0	166
0.5	163	0.40
1.2	156	0.50
1.5	149	0.70
2.1	132	0.97
3.0	122	0.88
3.5	115	0.88
4.0	108	0.87
4.4	97	0.95
5.0	87	0.94
6.0	73	0.93
Average		0.80

and

Table 19. Settling rate of Fe(OH)₃/gypsum-rich sludge.

Time	Height	Rate
Min	mm	m/h
0	166
2	160	0.18
4	149	0.26
7	143	0.20
9	136	0.20
13	125	0.19
15	115	0.20
18	104	0.21
20	99	0.20
24	80	0.22
27	70	0.21
Average		0.21

show the settling rate of Fe₃O₄/gypsum and Fe(OH)₃/gypsum-rich sludges. It was noted from Figure 14, Table 18 and Table 19 that the Fe₃O₄/gypsum sludge settled faster than the Fe(OH)₃/gypsum, namely 0.80 m/h versus 0.21 m/h. This can be ascribed to the higher density of Fe₃O₄ particles compared to Fe(OH)₃ particles and the formation of larger aggregates. In terms of hydration, Fe(OH)₃ is more hydrated and has a higher water content, making the particles less dense and more difficult to settle. This hydration slows down the settling process compared to Fe₃O₄ particles. Figure 15 shows the settled sludge for Fe(OH)₃ and Fe₃O₄-rich sludge.

4.4. Novelty

The novelty of the removal of iron from iron-rich mine water lies in the removal of the Fe²⁺ portion without the need to oxidise it to Fe³⁺, as is the case with the HDS process. Another novel aspect is the higher settling rate of the magnetite-rich sludge.

5. Conclusions and Recommendations

The following conclusions were made from this study: (i) Fe²⁺ can be removed as Fe₃O₄ (magnetite) by forming Fe(OH)₂ and Fe(OH)₃ in the mole ratio of 1:2 by precipitating Fe³⁺ with a carbonate or a hydroxide and by precipitating Fe²⁺ as a hydroxide, (ii) magnetite can form in the absence or presence of gypsum, (iii) the settling rate of magnetite-rich sludge is substantially faster than that of ferric hydroxide-rich sludge and (iv) separate recovery of magnetite and gypsum may be possible after further investigations. It is recommended that the following aspects be further investigated due to their practical value: (i) the effect of pH on magnetite formation, (ii) sludge settling and (iii) the separation of magnetite from gypsum through magnetic separation.

The optimal parameters for the recovery of magnetite involve using nitrogen gas to prevent the oxidation of Fe²⁺ to Fe³⁺ in solution. A stirring rate of 250 rpm effectively balanced particle dispersion. When NaOH was used as an alkali, optimal conditions were achieved at a pH of 11 and a temperature of 25 °C and NaOH (304.4 mmol/L). For instant and multiple dosing, the ideal conditions are a pH of 7.2 and a temperature of 100 °C, along with concentrations of CaCO₃ (160 mmol/L) and CaO (60 mmol/L). The Fe²⁺ removal rate for instant and multiple dosing was found to be 88.4% and 90.7%, respectively, while for NaOH it was found to be 97.6%. All methods demonstrate high Fe²⁺ removal rates, showcasing their effectiveness in treatment processes.