Application of magnesium hydroxide and barium hydroxide for the removal of metals and sulphate from mine water

 

 

V Bologo^*; JP Maree; F Carlsson

Department Environmental, Water and Earth Sciences, Tshwane University of Technology, Private Bag X680, Pretoria 0001, South Africa

 

 

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ABSTRACT

The proposed magnesium-barium-oxide process consists of metal removal with Mg(OH)₂, magnesium and sulphate removal with Ba(OH)₂ and calcium removal with CO₂. The raw materials, Mg(OH)₂ and Ba(OH)₂ are recovered from the BaSO₄ and Mg(OH)₂ sludges that are produced. Laboratory studies showed that metals are removed to low levels. This includes iron(II), the dominant metal ion in mine water, which is first oxidised to iron(III), whereafter it precipitates as Fe(OH)₃ resulting in residual levels of Fe(II) in the mine water of less than 20 mg/ℓ. Sulphate is also removed to less than 25 mg/ℓ. The final sulphate concentration is a function of the amount of Ba(OH)₂ dosed, as the amount of sulphate removed is stoichiometrically equivalent to the Ba(OH)₂ dosage. During CO₂-dosing, CaCO₃ is precipitated to the saturation level of CaCO₃.

Keywords: Magnesium hydroxide; barium hydroxide; sulphate removal; water treatment

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Introduction

Mining is a significant contributor to water pollution owing primarily to pyrite oxidation that generates potentially high levels of acidity, metal ions (mainly Fe), and sulphate (Reaction (1)) (Barbes and Romberger, 1968):

South Africa';s Witwatersrand Basin produces 340 Mℓ/d of mine-water, of which 50 Mℓ/d is produced by the Western Basin. Most of the underground mining operations in the Western Basin were closed by 1990. By 2004 the underground voids were filled by rising water levels and acid water started to decant at the surface. Contaminated void spaces increase with time and result in: (a) less water suitable for irrigation of crops, (b) less potable water available for the growing local communities, (c) the natural drying up of rivers, and (d) stream bed losses. The water of mining origin in the Western Basin contains, on average, 4 800 mg/ℓ sulphate, 800 mg/ℓ Fe(II), 100 mg/ℓ Fe(III), 230 mg/ℓ Mn, 11 mg/ℓ Zn, 18 mg/ℓ Ni, 5 mg/ℓ Co, 6 mg/ℓ Al, 150 mg/ℓ Mg, some radioactivity and 700 mg/ℓ free acidity (as CaCO₃). This acid mine drainage (AMD) source creates a major concern in that the Cradle of Humankind, a World Heritage Site, is not far from the decant point. As of March 2010, AMD started decanting into the Krugersdorp Nature Reserve at a rate of between 10 and 60 Mℓ/day, with the possibility of reaching the Sterkfontein Cave System which includes the Cradle of Humankind, where the earliest known hominid fossil remains were discovered and where paleontological excavations continue (Zvinowanda et al., 2010).

It is essential that a technically sound and cost-effective solution be found for the acid mine-water problem. South Africa is water-constrained and also has a mature mining economy with numerous mines having closed or in the process of being decommissioned. Should a solution be found that generates income, the AMD problem could be solved via economic principles rather than through government intervention by means of policies and legislated control measures.

Several processes have been considered for sulphate removal, e.g. biological removal, SAVMIN (etringite) (Ramsay, 1998), ecoDose, reverse osmosis (Chamber of Mines Research Organisation, 1988), and electrodialysis (Pulles et al., 1996). Soluble barium salts can also be used for sulphate removal and have certain advantages in that sulphate can be removed to specific levels owing to the low solubility of barium sulphate (Volman, 1984; Maree et al., 1989; Adlem et al., 1991). Soluble barium salts, such as barium sulphide, barium chloride and barium hydroxide, are potential candidates for the treatment regime and can be regenerated at the end of the process.

Kun (1972) studied the removal of sulphate with barium carbonate (BaCO₃). He identified 3 problems: a long retention time is required; high concentrations of soluble barium remain in the treated water when excess BaCO₃ is dosed over the required stoichiometric amount; and, the high cost of BaCO₃. Volman (1984) and Maree et al. (1989) overcame the cost problem by demonstrating that barium sulphate (BaSO₄) could be reduced efficiently and economically with coal under thermal conditions to produce barium sulphide (BaS). This compound can be used directly for the process or first converted to BaCO₃. Wilsenach (1986) demonstrated the economic viability of the process by calculating the cost of producing BaS from BaSO₄.

Trusler et al. (1988) developed an integrated process which includes the recovery of BaCO₃ and lime. They noted that incomplete sulphate removal with BaCO₃ resulted when sulphate is not completely associated with calcium ions. They overcame this problem by dosing lime for pre-removal of magnesium as Mg(OH)₂. Maree et al. (1989) have developed a barium carbon ate process (a 2-stage, fluidised-bed reactor system) to overcome the other problems identified by Kun (1972), i.e., the long retention time and the high barium concentration in the treated water. A disadvantage of the barium carbonate process, proposed by Trusler (1988) and improved by Maree et al. (1989), is the formation of a mixed precipitate of BaSO₄ and CaCO₃. Bosman et al. (1990) and Adlem et al. (1991) overcame this problem by pre-treating the AMD with lime to remove metals before the addition of barium sulphide. BaS is highly soluble in water and can therefore readily provide Ba²⁺ ions for fast interaction with SO₄^2- in sulphate-rich water. Maree et al. (1990) elucidated the ability of BaS to directly remove sulphate from acid waters. Other studies have demonstrated that BaS has an advantage of producing by-products like sulphur and NaHS, which can potentially be derived from the H₂S produced during the H₂S stripping stage, and CaCO₃ from the softening stage (Maree et al., 1989; Bosman et al., 1990; Adlem et al., 1991; Du Preez and Maree, 1994). Laboratory studies were carried out by Maree et al. (2004) to demonstrate that the integrated BaS process is technically and economically viable for sulphate removal. This process consists of 4 stages. In the thermal stage, BaSO₄ is reduced to BaS at 1 050ºC, in a kiln using coal as the reductant. In the sulphate removal stage, sulphate is precipitated as BaSO₄. In the stripping stage, the barium sulphide is dissolved in water and reacted with CO₂ from the kiln to remove or strip off the H₂S gas. In the softening stage, CaCO₃ is precipitated as a result of CO₂-stripping with air.

Hlabela et al. (2007) demonstrated that BaCO₃ can be used effectively for sulphate removal after pre-treatment of the AMD with lime for removal of metals, including magnesium, as metal hydroxides.

The purpose of the present study was to demonstrate, using a laboratory study, the performance of an alternative, integrated magnesium hydroxide and barium hydroxide process (MBO Process) for treatment of acid mine-water. The provisionally patented process (Maree and Louw, 2010) consists of the following stages (Fig. 1):

• Magnesium hydroxide, for neutralisation and metal removal

• Barium hydroxide, for sulphate and magnesium removal

• CO₂, for CaCO₃ precipitation

• Recovery of Ba(OH)₂ via BaS from the BaSO₄ produced

The specific aims were to demonstrate that:

• Metals can be completely removed with Mg(OH)₂

• Sulphate concentration can be lowered to less than 200 mg/ℓ

• Magnesium concentration can be lowered to less than 10 mg/ℓ with Ba(OH)₂

• The high pH water can be stabilised with CO₂

The processing of BaSO₄ and Mg(OH)₂ in the sludge to recover the process raw materials, Ba(OH)₂ and Mg(OH)₂, were investigated separately.

 

Materials and methods

Feedstock

Waters from a coal mine and a gold mine were used as feed waters containing 2 400 - 4 800 mg/ℓ sulphate (for chemical analyses, see Tables 1 and 2). Magnesium hydroxide and barium hydroxide were used for pH adjustment and sulphate removal, respectively. CO₂ gas (Afrox) was used for pH adjustment after sulphate removal.

 

 

The following chemicals were used in this study: Barium hydroxide (Ba(OH)₂.8H₂O, analytical reagent, Merck) and calcium hydroxide (97% Ca(OH)₂, Rochelle Chemicals, Johannesburg); Mg(OH)₂ (CP grade, May and Baker Ltd, Dagenham, England), air (HP compressed; Afrox, South Africa), and CO₂ (Afrox, South Africa).

Equipment

Neutralisation with Mg(OH)₂ and sulphate removal with Ba(OH)₂ were studied using stirred beakers.

Experimental procedure

Feed water was mixed in a beaker (5 ℓ), aerated for iron(II) oxidation and stirred (30 min) for completion of the oxidation/precipitation reactions. The sludge was allowed to settle and the clear water decanted. Ba(OH)₂ was mixed with the decant water (1000 mℓ glass beakers). Stabilisation of the Ba(OH)₂ treated water was achieved by passing CO₂ through the water.

Experimental programme

The following parameters were investigated: metal removal with Mg(OH)₂, magnesium and sulphate removal with Ba(OH)₂, and calcium removal with CO₂.

Analytical

Samples were collected regularly and filtered (Whatman No. 1). Sulphate, alkalinity, Fe(II), mixed liquor suspended solids (MSS), volatile suspended solids (VSS), acidity, and pH were determined according to standard procedures (APHA, 1989). Calcium and magnesium were assayed using atomic absorption spectrophotometry.

 

Results and discussion

Water quality and chemical reactions

The MBO study can be applied to the treatment of acid mine-water. Tables 1 and 2 show the chemical composition of the feed waters before and after treatment with Mg(OH)₂, Ba(OH)₂ and CO₂, for effluents from coal and gold mines, respectively. The total dissolved solids (TDS) content in the case of coal mine-water was 3 406 mg/ℓ in the feed water, 3 434 mg/ℓ after Mg(OH)₂ treatment, 1 286 mg/ℓ after Ba(OH)₂ treatment and 454 mg/ℓ after CO₂ addition (Table 1). Similarly, in the case of gold mine effluent, the overall TDS content was lowered from 6 954 mg/ℓ in the feed water to 416 mg/ℓ, following treatment with the MBO process (Table 2). During Mg(OH)₂ treatment, free acid and all of the metal concentrations, excluding that of calcium and magnesium, were lowered to below allowable limits for drinking water in South Africa (SANS 241:2006). The lowering of the TDS in the case of Mg(OH)₂ treatment was mainly due to the oxidation of Fe²⁺ to Fe³⁺ (Reaction (2)) and precipitation as Fe(OH)₃ (Reaction (3)). Similarly, Mn²⁺ was oxidised to Mn⁴⁺ and precipitated as MnO₂. The other metals, Al³⁺, Co²⁺, Zn²⁺, Pb²⁺, Co²⁺ and Ni²⁺, precipitated as metal hydroxides. This was owing to the low solubility-products for Fe(OH)(2.64 x 10^-39); Al(OH)(8.5 x 10^-23); Mn(OH)(2.06 x 10^-13); Cu(OH)₂ (2.20 x 10^-20); Zn(OH)₂ (7.71 x 10^-17); Pb(OH)₂ (1.42 x 10^-20); Co(OH)₂ (1.09 x 10^-15) and Ni(OH)₂ (5.47 x 10^-16) (Sillen and Martell, 1964; Lide, 1992).

The precipitated BaSO₄ and Mg(OH)₂ can be made to settle rapidly by dosing coagulants. Coagulant dosing can be avoided by controlling the precipitation process to yield particles with improved settling rates. Hlabela (2009) showed that a multiple point, in-line dosing system can produce flocs within the size range 10 - 100 µm. In the final step, the calcium concentration in the feed water was lowered from 454 mg/ℓ down to 40 mg/ℓ (as Ca⁺) by stabilising the water with CO₂ (Reaction (8)). During CO₂ addition for pH adjustment, CaCO₃ crystallisation occurred, as indicated by the decrease in the calcium concentration. According to Loewenthal et al. (1986), the saturation, under-saturation and super-saturation states of CaCO₃ are theoretically identified by the activity product of Ca²⁺ and CO₃^2- that is comparable to the solubility product constant, k_sp, for CaCO₃. Above the solubility product, CaCO₃ will precipitate from solution, and below, CaCO₃ will dissolve. The rate of precipitation or dissolution is described by Eq. (9), where k is the precipitation rate-constant, S, the surface area of CaCO₃ growth/dissolution sites, and the parameters in square brackets, activities in mol/ℓ. The rate of precipitation depends on k, S and the term within curly brackets. Although k is affected to an unknown extent by crystal structure, S is dependent on the size, mass and structure of the crystals. By controlling S and the degree of super-saturation (term in curly brackets), the rate of precipitation or dissolution of CaCO₃ can be determined (Loewenthal et al., 1986).

where:

k = precipitation rate-constant,

S = surface area of CaCO₃ growth/dissolution sites and the terms in square brackets, activities in mol/ℓ.

The rate of sulphate removal through BaSO₄ crystallisation is influenced by the sulphate concentration in solution and the BaSO₄ seed crystal concentration. It was observed that sulphate removal was stoichiometrically equivalent to the Ba(OH)₂ dosage, as indicated by [Ba²⁺ dosage]/[SO₄ removed] (mol/mol) ratios of 1.05, 0.97 and 1.02, respectively (Table 3). As expected, higher residual Ba²⁺ concentrations were in solution with lower residual SO₄² concentrations. This is due to the low solubility product of BaSO₄ and the rate of BaSO₄ crystallisation at low seed-crystal concentrations. The practical implication is that about 200 mg/ℓ free sulphate should remain in solution to limit Ba²⁺ concentration to less than 2 mg/ℓ in solution, the guideline value for drinking water as laid down by the USEPA (2011). Optimum conditions also need to prevail to allow rapid BaSO₄-crystallisation, such as providing a solid BaSO₄ presence in the reactor by means of sludge recirculation.

 

 

Barium ions react with sulphate ions forming a white precipitate of BaSO. Figure 2 shows an increase in the sulphate removal rate with increased barium hydroxide concentration (molar ratios of 0.8, 0.85, 0.9 and 1) in the solution. From chemical Reaction (7), 1 mole of barium reacts with 1 mole of sulphate. Therefore, it is clear that the more barium ions available in solution the more sulphate ions are removed from the solution. Sulphate removal followed normal stoichiometry. Therefore, it is sufficient to use a 1:1 mole ratio of dissolved barium to sulphate for almost complete removal of sulphate from wastewater.

 

 

The effect of BaSO₄ seeds added to AMD before the addition of Ba(OH)₂ was investigated. The results are shown in Fig. 3. From these results it was observed that the presence of a small amount of barium sulphate seed did not have any significant effects on the removal of sulphate.

 

 

Four different experiments were carried out using increasingly rapid stirring rates. The results are shown in Fig. 4. It was also observed (Fig. 4) that the removal-rate initially increased linearly and subsequently flattened to become abruptly constant, showing a complete removal of sulphate. The stirring rates in the range 250-800 r/min did not show any significant differences in the removal rates of sulphates.

 

 

An advantage of using Mg(OH)₂ for neutralisation and metal removal is that the precipitated metal hydroxides are not mixed with gypsum. A sludge of higher density was produced which settles to the bottom faster. The sludge from the Mg(OH)₂ stage consisted of Fe(OH)₃, Al(OH)₃, MnO₂ and Zn(OH)₂. The 'Ba(OH)₂ sludge'; consisted of 70% BaSO₄ and 30% Mg(OH)₂.

 

Conclusions

With the MBO study at a laboratory scale, it was demonstrated, that:

• Mg(OH)₂ precipitated metals as metal hydroxides. Iron(II) was first oxidised to iron(III) before being precipitated.

• Ba(OH)₂ precipitated sulphate as BaSO₄ and magnesium as Mg(OH)₂. Sulphate removal was stoichiometrically equivalent to the Ba(OH)₂ dosage.

• ₂ precipitated calcium as CaCO₃.

 

Acknowledgements

This work was made possible with the support of the following institutions: The National Research Foundation, which provided funding for Tshwane University of Technology projects on neutralisation and sulphate removal through their Technology and Human Resources for Industry Programme (THRIP); Rand Water, which funds activities falling under the Rand Water Chair in Water Utilisation at the Tshwane University of Technology, and Tshwane University of Technology, which provided substantial financial support for the research programme. The contributions of the following people are gratefully acknowledged: Dr Fritz Carlsson for editing the manuscript and Dr Caliphs Zvinowanda, Mr Kebede Kefeni and Mr Dave Katlekgo Mpholoane for assistance with the laboratory work.

 

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Received 16 February 2011; accepted in revised form 14 December 2011.

 

 

* To whom all correspondence should be addressed. +27 12 382 6320; fax: +27 12 382-6354; e-mail: bologov@tut.ac.za