Acidic or alkaline groundwater or porewater, in natural settings or man-made wastes or operations, may be derived from a multitude of sources. Examples of these sources may include:                Oxidation of sulphide-containing soils to form acid sulphate soils (ASS) by natural processes (e.g. seasonal changes in groundwater level and/or oxygen status) or disturbance (e.g. during construction or excavation)        Industrial processes (e.g. pyrite oxidation, sulphuric acid production) with offsite loss via soil/groundwater infiltration        Discharge, escape and infiltration of acidic or alkaline surface waters from mining or extractive metallurgical operations        In-situ leaching of orebodies (e.g. uranium or copper ores)        Fluids derived from mineral processing (e.g. alkaline red mud via the Bayer process).        
As a consequence of the processes that lead to the formation of these waters, such waters may often be enriched in a variety of metals, metalloids and anions, the concentrations of which may exceed both ANZECC Soil and Water Quality guidelines (ANZECC/NHMRC, 1992).
A challenge exists to identify methods for remediation of the acidic and alkaline groundwater that are both cost-effective and environmentally robust with safe and efficient immobilization (and if appropriate, off-site disposal) of the contaminants after neutralization. Effective long-term management of acidic and alkaline groundwaters is also required to meet regulatory requirements.
Layered double hydroxides (LDH) are a class of both naturally occurring and synthetically produced materials characterised by a positively-charged mixed metal hydroxide layers separated by interlayers that contain water molecules and a variety of exchangeable anions. Layered double hydroxides most commonly formed by the coprecipitation of divalent (e.g. Mg2+, Fe2+) and trivalent (e.g. Al3+, Fe3+) metal cation solutions at moderate to high pH (Taylor, 1984, Vucelic et al, 1997, Shin et al, 1996).
Layered double hydroxide compounds may be represented by the general formula (1):M(1−X)2+Mx3+(OH)2An−yH2O  (1)
where M2+ and M3+ are divalent and trivalent metal ions, respectively and An− is the interlayer ion of valence n. The x value represents the proportion of trivalent metal ion to the proportion of total amount metal ion and y denotes variable amounts of interlayer water.
Common forms of layered double hydroxides comprise Mg2+ and Al3+ (commonly known as hydrotalcites) and Mg2+ and Fe3+ (known as pyroaurites), but other cations, including Ni, Zn, Mn, Ca, Cr and La, are known. The amount of surface positive charge generated is dependant upon the mole ratio of the metal ions in the lattice structure and the conditions of preparation as they affect crystal formation.
The formation of hydrotalcite (the most commonly synthesised LDH with carbonate as the principal “exchangeable” anion) may be most simply described by the following reaction:6MgCl2+2AlCl3+16NaOH+H2CO3→Mg6Al2(OH)16CO3.nH2O+2HCl
Typically, ratios of divalent to trivalent cations in hydrotalcites vary from 2:1 to 3:1. Other synthetic pathways to form hydrotalcite (and other LDH) include synthesis from Mg(OH)2 (brucite) and MgO (calcined magnesia) via neutralisation of acidic solutions (eg. Albiston et al, 1996). This can be described by the following reaction:6Mg(OH)2+2Al(OH)3+2H2SO4→Mg6Al2(OH)16SO4.nH2O+2H2O
A range of metals of widely varying concentrations may also be simultaneously coprecipitated, hence forming a polymetallic LDH. Hydrotalcites or LDH were first described over 60 years ago (Frondel, 1941, Feitknecht, 1942). Sometimes, they can also occur in nature as accessory minerals in soils and sediments (eg. Taylor and McKenzie, 1980). Layered double hydroxides may also be synthesised from industrial waste materials by the reaction of bauxite residue derived from alumina extraction (red mud) with seawater (eg. Thornber and Hughes, 1987), as described by the following reaction:6Mg(OH)2+2Al(OH)3+2Na2CO3→Mg6Al2(OH)16CO3.nH2O+2NaOHor by the reaction of lime with fly ash derived from fossil fuel (eg. coal fired power stations, Reardon and Della Valle, 1997).
Within the LDH a structure there are octahedral metal hydroxide sheets that carry a net positive charge due to limited substitution of trivalent for divalent cations as described above. As a consequence, it is possible to substitute a wide range of inorganic or organic anions into the LDH structure. These anions are often referred to as “interlayer anions” as they fit between the layers of hydroxide material. Layered double hydroxides are generally unstable below a pH of approximately 5 (Ookubu et al, 1993) but may act as buffers over a wide range of solution pH (Seida and Nakano, 2002). Layered double hydroxides, and in particular those that contain carbonate as the predominant anion have also been demonstrated to have a considerable capacity to neutralise a range of mineral acids via consumption of both the hydroxyl and carbonate anions contained within the LDH structure (eg. Kameda et al, 2003).
A number of studies have been conducted to investigate ways to exploit the anion exchange properties of LDH. These studies have focussed on the removal of phosphate and other oxyanions and humic substances from natural and wastewaters (Miyata, 1980, Misra and Perrotta, 1992, Amin and Jayson, 1996, Shin et al, 1996, Seida and Nakano, 2000). Phosphate is one of the many anions that may be exchanged into the interlayer space in LDH. Laboratory studies of phosphate uptake using synthetically prepared Mg—Al hydrotalcites and a range of initial dissolved phosphate concentrations indicate an uptake capacity of from ca. 25-30 mg P/g (Miyata, 1983, Shin et al, 1996) to ca. 60 mg P/g with uptake also influenced by initial phosphate concentration, pH (with maximum phosphate absorption near pH 7), degree of crystallinity and the hydrotalcite chemistry (Ookubo et al, 1993). A major obstacle to the use of hydrotalcites for phosphate removal in natural and/or wastewaters is the selectivity for carbonate over phosphate, with a selectivity series in the approximate order CO32−>HPO42−>>SO42−, OH−>F−>Cl−>NO3− (Miyata, 1980, 1983, Sato et al, 1986, Shin et al, 1986, Cavani et al, 1991). Many hydrotalcites are also synthesised with carbonate as the predominant anion and thus require anion exchange before they are exposed to phosphate. When carbonate is also combined with sulphate, nitrate and chloride (as might commonly occur in natural or wastewaters) the reduction of phosphate absorption to the hydrotalcite is further decreased (Shin et al, 1996).
A number of recent studies have focussed on the formation and study of synthetic LDH and their subsequent reactivity to a range of anions, particularly silicate (e.g. Depege et al, 1996) with a view to forming polymetallic aluminosilicates, which as potential precursors to clay materials, are thought to limit metal mobility and bioavailability (eg. Ford et al, 1999). A potential also exists for the co-precipitation of silicate and aluminate anions and another precursor of analogue of clay minerals.
Calcined magnesia (MgO) or its derivative, magnesium hydroxide (Mg(OH)2), possess considerable advantages over other alkalis such as slaked lime (Ca(OH)2) in the neutralisation of acids or acidic wastes1.
One of the most important advantages is the relatively small amount of calcined magnesia (MgO) magnesium hydroxide (Mg(OH)2) that is required. For the neutralisation of 1 tonne of 98% sulphuric acid, only 424 kg of 96% solid MgO, 613 kg of 96% solid Mg(OH)2 or 1005 kg of a 58% slurry of Mg(OH)2 are required. In comparison, almost 1600 kg of a 50% NaOH solution, 1645 kg of a 45% solution of Ca(OH)2, 3210 kg of a 33% slurry of Na2CO3 or 975 kg of CaCO3 are required to achieve neutralization of 1 tonne of 98% sulphuric acid.
The chemistry of calcined magnesia also confers a number of distinct advantages. Alkalis such as caustic soda or lime can be considered to neutralise by a one-step dissociation reaction that results in the formation of hydroxyl ions and an increase in the solution pH. In contrast, the neutralisation of acidic solutions by calcined magnesia can be considered to be a two-step reaction as magnesium hydroxide, the intermediate product in the neutralisation process is only slightly soluble in water. As a consequence neutralisation occurs as soluble hydroxide ions derived from magnesium hydroxide are consumed by the acid. Using sulphuric acid as the acid source, the neutralisation reactions can be summarised as follows:MgO+H2O→Mg(OH)2 Mg(OH)2→Mg2++2OH−H2SO4+Mg2++2OH−→MgSO4+2H2O
As a consequence of the production of hydroxide ions from the slightly soluble magnesium hydroxide, the neutralisation reaction occurs rapidly at low pH and slows appreciably as the pH increases. In addition, varying mineral grainsize can change the reactivity of MgO. In contrast, the neutralisation rate of lime and similar products do not vary appreciably as a function of pH. It is the slower reaction rate of calcined magnesia that results in the formation of denser slurries (eg. of mineral precipitates) relative to lime, thus reducing handling and disposal costs. In addition, the positive charge on magnesium-based alkalis at neutral to marginally alkaline pH attracts negatively charged particles (eg. humic substances, some colloids) often facilitating superior filtration of high-solids sludges.
Calcined magnesia is also appreciably safer to handle than a range of other alkalis such as caustic soda. Magnesia-based alkalis are virtually non-corrosive, only weakly exothermic and reactive and hence easy to handle, thus reducing safety concerns. These features contrast strongly with lime and other alkalis. An additional feature of calcined magnesia is the potential for the efficient (and often simultaneous) removal of a range of metals. The removal efficiency is related to the presence of a high pH immediately adjacent to the particle surface. This high pH zone can provide an ideal zone for the precipitation of metal hydroxides which may cement onto the surface of calcined magnesia substrate. Calcined magnesia has also been used in the simultaneous removal of ammonia and phosphate (principally in sewage) via the precipitation of struvite (MgNH4PO4).
Calcined magnesia also has a number of potential advantages relative to other remedial strategies such as sulphate reduction, the application of calcite and/or lime in removing metals and other ions from groundwater. Comparatively, sulphate reduction is a slow process with a long residence time often required (ca. weeks) for effective reduction of metal concentrations to take place. Hence, an unrealistically thick barrier may be required, particularly in areas where there is a high groundwater gradient or flow velocity, the latter perhaps due to groundwater extraction. Where organic matter and/or calcite are used for treatment applications the residual pH of ca. 6 after treatment while allowing the removal of trivalent metals as hydroxides is insufficient to allow the precipitation of many divalent metals ions which require a higher pH. In the case of lime, even moderate over application may result in a residual pH of ca. 12 resulting in a range of deleterious effects including to endemic micro- and macro-biota, with possible re-neutralisation of any discharged water required, particularly in sensitive environmental areas (see, for example, Cortina et al, 2003).
The applicant does not concede that the prior art discussed herein forms part of the common general knowledge.